# Glossary


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Glossary Entries
Word(s) Definition Image Caption Link Source
base ionization constant An equilibrium constant for the reaction of a weak base (B) with water, in which the concentration of water is treated as a constant. Abbreviation: Kb
Chemistry The study of matter and the changes that material substances undergo.
biochemistry The application of chemistry to the study of biological processes.
scientific method The procedure that scientists use to search for answers to questions and solutions to problems.
hypothesis A tentative explanation for scientific observations that puts the system being studied into a form that can be tested.
law of definite proportions A chemical substance always contains the same proportions of elements by mass.
theory A statement that attempts to explain why nature behaves the way it does.
matter Anything that occupies space and has mass.
weight A force caused by the gravitational attraction that operates on an object. The weight of an object depends on its location (c.f. mass).
physical change A change of state that does not affect the chemical composition of a substance.
homogeneous A mixture in which all portions of a material are in the same state, have no visible boundaries, and are uniform throughout.
heterogeneous A mixture in which a material is not completely uniform throughout.
Distillation A physical process used to separate homogeneous mixtures (solutions) into their component substances. Distillation makes use of differences in the volatilities of the component substances.
Crystallization A physical process used to separate homogeneous mixtures (solutions) into their component substances. Crystallization separates mixtures based on differences in their solubilities.
chemical change A process in which the chemical composition of one or more substances is altered.
Chemical properties The characteristic ability of a substance to react to form new substances.
Intensive properties A physical property that does not depend on the amount of the substance and physical state at a given temperature and pressure.
density (d) An intensive property of matter, density is the mass per unit volume (usually expressed in g/cm3). At a given temperature, the density of a substance is a constant.
atoms The fundamental, individual particles of which matter is composed.
transmutation The process of converting one element to another.
combustion The burning of a material in an oxygen atmosphere.
law of conservation of mass In any chemical reaction, the mass of the substances that react equals the mass of the products that are formed.
law of multiple proportions When two elements form a series of compounds, the ratios of the masses of the second element that are present per gram of the first element can almost always be expressed as the ratios of integers. (The same law holds for the mass ratios of compounds forming a series that contains more than two elements.)
neutrons A subatomic particle with no charge that resides in the nucleus of almost all atoms.
radioactivity The spontaneous emission of energy rays (radiation) by matter.
nucleus The central core of an atom where protons and any neutrons reside.
periodic table A chart of the chemical elements arranged in rows of increasing atomic number so that the elements in each column (group) have similar chemical properties.
mass number (A) The number of protons and neutrons in the nucleus of an atom of an element.
ions A charged particle produced when one or more electrons is removed from or added to an atom or molecule.
atomic mass unit (amu) One-twelfth of the mass of one atom of $C12$; .
transition elements Any element in groups 3–12 in the periodic table. All of the transition elements are metals.
semimetals Any element that lies adjacent to the zigzag line in the periodic table that runs from boron to astatine. Semimetals (also called metalloids) exhibit properties intermediate between those of metals and nonmetals.
lustrous Having a shiny appearance. Metals are lustrous, whereas nonmetals are not.
noble gases Any element in group 18 of the periodic table. All are unreactive monatomic gases at room temperature and pressure.
monatomic A species containing a single atom.
essential elements Any of the 19 elements that are absolutely required in the human diet for survival. An additional seven elements are thought to be essential for humans.
volume The amount of space occupied by a sample of matter.
Système internationale d’unités (or SI) A system of units based on metric units that requires measurements to be expressed in decimal form. There are seven base units in the SI system.
scientific notation A system that expresses numbers in the form N × 10n, where N is greater than or equal to 1 and less than 10 (1 ≤ N < 10) and n is an integer that can be either positive or negative (100 = 1). The purpose of scientific notation is to simplify the manipulation of numbers with large or small magnitudes.
significant figures Numbers that describe the value without exaggerating the degree to which it is known to be accurate.
exact numbers An integer obtained either by counting objects or from definitions (e.g., 1 in. = 2.54 cm). Exact numbers have infinitely many significant figures.
precise Multiple measurements give nearly identical values.
Chemistry The study of matter and the changes that material substances undergo.
biochemistry The application of chemistry to the study of biological processes.
scientific method The procedure that scientists use to search for answers to questions and solutions to problems.
hypothesis A tentative explanation for scientific observations that puts the system being studied into a form that can be tested.
law of definite proportions A chemical substance always contains the same proportions of elements by mass.
theory A statement that attempts to explain why nature behaves the way it does.
matter Anything that occupies space and has mass.
weight A force caused by the gravitational attraction that operates on an object. The weight of an object depends on its location (c.f. mass).
physical change A change of state that does not affect the chemical composition of a substance.
homogeneous A mixture in which all portions of a material are in the same state, have no visible boundaries, and are uniform throughout.
heterogeneous A mixture in which a material is not completely uniform throughout.
Distillation A physical process used to separate homogeneous mixtures (solutions) into their component substances. Distillation makes use of differences in the volatilities of the component substances.
Crystallization A physical process used to separate homogeneous mixtures (solutions) into their component substances. Crystallization separates mixtures based on differences in their solubilities.
chemical change A process in which the chemical composition of one or more substances is altered.
Chemical properties The characteristic ability of a substance to react to form new substances.
Intensive properties A physical property that does not depend on the amount of the substance and physical state at a given temperature and pressure.
density (d) An intensive property of matter, density is the mass per unit volume (usually expressed in g/cm3). At a given temperature, the density of a substance is a constant.
atoms The fundamental, individual particles of which matter is composed.
transmutation The process of converting one element to another.
combustion The burning of a material in an oxygen atmosphere.
law of conservation of mass In any chemical reaction, the mass of the substances that react equals the mass of the products that are formed.
law of multiple proportions When two elements form a series of compounds, the ratios of the masses of the second element that are present per gram of the first element can almost always be expressed as the ratios of integers. (The same law holds for the mass ratios of compounds forming a series that contains more than two elements.)
neutrons A subatomic particle with no charge that resides in the nucleus of almost all atoms.
radioactivity The spontaneous emission of energy rays (radiation) by matter.
nucleus The central core of an atom where protons and any neutrons reside.
periodic table A chart of the chemical elements arranged in rows of increasing atomic number so that the elements in each column (group) have similar chemical properties.
mass number (A) The number of protons and neutrons in the nucleus of an atom of an element.
ions A charged particle produced when one or more electrons is removed from or added to an atom or molecule.
atomic mass unit (amu) One-twelfth of the mass of one atom of $C12$; .
transition elements Any element in groups 3–12 in the periodic table. All of the transition elements are metals.
semimetals Any element that lies adjacent to the zigzag line in the periodic table that runs from boron to astatine. Semimetals (also called metalloids) exhibit properties intermediate between those of metals and nonmetals.
lustrous Having a shiny appearance. Metals are lustrous, whereas nonmetals are not.
noble gases Any element in group 18 of the periodic table. All are unreactive monatomic gases at room temperature and pressure.
monatomic A species containing a single atom.
essential elements Any of the 19 elements that are absolutely required in the human diet for survival. An additional seven elements are thought to be essential for humans.
volume The amount of space occupied by a sample of matter.
Système internationale d’unités (or SI) A system of units based on metric units that requires measurements to be expressed in decimal form. There are seven base units in the SI system.
scientific notation A system that expresses numbers in the form N × 10n, where N is greater than or equal to 1 and less than 10 (1 ≤ N < 10) and n is an integer that can be either positive or negative (100 = 1). The purpose of scientific notation is to simplify the manipulation of numbers with large or small magnitudes.
significant figures Numbers that describe the value without exaggerating the degree to which it is known to be accurate.
exact numbers An integer obtained either by counting objects or from definitions (e.g., 1 in. = 2.54 cm). Exact numbers have infinitely many significant figures.
precise Multiple measurements give nearly identical values.
chemical bonds An attractive interaction between atoms that holds them together in compounds.
covalent bond The electrostatic attraction between the positively charged nuclei of the bonded atoms and the negatively charged electrons they share.
polyatomic Molecules that contain more than two atoms.
molecular formula A representation of a covalent compound that consists of the atomic symbol for each component element (in a prescribed order) accompanied by a subscript indicating the number of atoms of that element in the molecule. The subscript is written only if the number is greater than 1.
inorganic compounds An ionic or covalent compound that consists primarily of elements other than carbon and hydrogen.
structural formulas A representation of a molecule that shows which atoms are bonded to one another and, in some cases, the approximate arrangement of atoms in space.
triple bond A chemical bond formed when two atoms share three pairs of electrons.
anions An ion that has fewer protons than electrons, resulting in a net negative charge.
monatomic ions An ion with only a single atom.
formula unit The absolute grouping of atoms or ions represented by the empirical formula.
binary ionic compound An ionic compound that contains only two elements, one present as a cation and one as an anion.
Polyatomic ions A group of two or more atoms that has a net electrical charge.
waters of hydration The loosely bound water molecules in hydrate compounds. These waters of hydration can often be removed by simply heating the compound.
aromatic hydrocarbons An unsaturated hydrocarbon consisting of a ring of six carbon atoms with alternating single and double bonds.
cyclic hydrocarbon A hydrocarbon in which the ends of the carbon chain are connected to form a ring of covalently bonded carbon atoms.
aliphatic hydrocarbons Alkanes, alkenes, alkynes, and cyclic hydrocarbons (hydrocarbons that are not aromatic).
R The abbreviation used for alkyl groups and aryl groups in general formulas and structures.
alcohol A class of organic compounds obtained by replacing one or more of the hydrogen atoms of a hydrocarbon with an −OH group.
bases A substance that produces one or more hydroxide ions $(OH−)$ and a cation when dissolved in aqueous solution, thereby forming a basic solution.
oxoacids An acid in which the dissociable $H+$ ion is attached to an oxygen atom of a polyatomic anion.
carboxylic acids An organic compound that contains an −OH group covalently bonded to the carbon atom of a carbonyl group. The general formula of a carboxylic acid is $RCO2H$. In water a carboxylic acid dissociates to produce an acidic solution.
amine An organic compound that has the general formula $RNH2$, where R is an alkyl group. Amines, like ammonia, are bases.
reforming The second process used in petroleum refining, which is the chemical conversion of straight-chain alkanes to either branched-chain alkanes or mixtures of aromatic hydrocarbons.
octane rating A measure of a fuel’s ability to burn in a combustion engine without knocking or pinging (indications of premature combustion). The higher the octane rating, the higher quality the fuel.
chemical bonds An attractive interaction between atoms that holds them together in compounds.
covalent bond The electrostatic attraction between the positively charged nuclei of the bonded atoms and the negatively charged electrons they share.
polyatomic Molecules that contain more than two atoms.
molecular formula A representation of a covalent compound that consists of the atomic symbol for each component element (in a prescribed order) accompanied by a subscript indicating the number of atoms of that element in the molecule. The subscript is written only if the number is greater than 1.
inorganic compounds An ionic or covalent compound that consists primarily of elements other than carbon and hydrogen.
structural formulas A representation of a molecule that shows which atoms are bonded to one another and, in some cases, the approximate arrangement of atoms in space.
triple bond A chemical bond formed when two atoms share three pairs of electrons.
anions An ion that has fewer protons than electrons, resulting in a net negative charge.
monatomic ions An ion with only a single atom.
formula unit The absolute grouping of atoms or ions represented by the empirical formula.
binary ionic compound An ionic compound that contains only two elements, one present as a cation and one as an anion.
Polyatomic ions A group of two or more atoms that has a net electrical charge.
waters of hydration The loosely bound water molecules in hydrate compounds. These waters of hydration can often be removed by simply heating the compound.
aromatic hydrocarbons An unsaturated hydrocarbon consisting of a ring of six carbon atoms with alternating single and double bonds.
cyclic hydrocarbon A hydrocarbon in which the ends of the carbon chain are connected to form a ring of covalently bonded carbon atoms.
aliphatic hydrocarbons Alkanes, alkenes, alkynes, and cyclic hydrocarbons (hydrocarbons that are not aromatic).
R The abbreviation used for alkyl groups and aryl groups in general formulas and structures.
alcohol A class of organic compounds obtained by replacing one or more of the hydrogen atoms of a hydrocarbon with an −OH group.
bases A substance that produces one or more hydroxide ions $(OH−)$ and a cation when dissolved in aqueous solution, thereby forming a basic solution.
oxoacids An acid in which the dissociable $H+$ ion is attached to an oxygen atom of a polyatomic anion.
carboxylic acids An organic compound that contains an −OH group covalently bonded to the carbon atom of a carbonyl group. The general formula of a carboxylic acid is $RCO2H$. In water a carboxylic acid dissociates to produce an acidic solution.
amine An organic compound that has the general formula $RNH2$, where R is an alkyl group. Amines, like ammonia, are bases.
reforming The second process used in petroleum refining, which is the chemical conversion of straight-chain alkanes to either branched-chain alkanes or mixtures of aromatic hydrocarbons.
octane rating A measure of a fuel’s ability to burn in a combustion engine without knocking or pinging (indications of premature combustion). The higher the octane rating, the higher quality the fuel.
molecular mass The sum of the average masses of the atoms in one molecule of a substance, each multiplied by its subscript.
formula mass The sum of the atomic masses of all the elements in the empirical formula, each multiplied by its subscript.
mole (mol) The quantity of a substance that contains the same number of units (e.g., atoms or molecules) as the number of carbon atoms in exactly 12 g of isotopically pure carbon-12.
Avogadro’s number The number of units (e.g., atoms, molecules, or formula units) in 1 mol: $6.022×1023$.
molar mass The mass in grams of 1 mol of a substance.
percent composition The percentage of each element present in a pure substance. With few exceptions, the percent composition of a chemical compound is constant (see law of definite proportions).
product(s) The final compound(s) produced in a chemical reaction.
coefficient A number greater than 1 preceding a formula in a balanced chemical equation and indicating the number of atoms, molecules, or formula units of a reactant or a product.
mole ratio The ratio of the number of moles of one substance to the number of moles of another, as depicted by a balanced chemical equation.
stoichiometric quantity The amount of product or reactant specified by the coefficients in a balanced chemical equation.
limiting reactant The reactant that restricts the amount of product obtained in a chemical reaction.
theoretical yield The maximum amount of product that can be formed from the reactants in a chemical reaction, which theoretically is the amount of product that would be obtained if the reaction occurred perfectly and the method of purifying the product were 100% efficient.
percent yield The ratio of the actual yield of a reaction to the theoretical yield multiplied by 100 to give a percentage.
oxidation–reduction reactions A chemical reaction that exhibits a change in the oxidation states of one or more elements in the reactants that has the general form oxidant + reductant → reduced oxidant + oxidized reductant.
oxidation The loss of one or more electrons in a chemical reaction. The substance that loses electrons is said to be oxidized.
reduction The gain of one or more electrons in a chemical reaction. The substance that gains electrons is said to be reduced.
oxidation state The charge that each atom in a compound would have if all its bonding electrons were transferred to the atom with the greater attraction for electrons.
reductants (or reducing agents) A compound that is capable of donating electrons; thus it is oxidized.
combustion reaction An oxidation–reduction reaction in which the oxidant is $O2$.
catalysis The acceleration of a chemical reaction by a catalyst.
heterogeneous catalyst A catalyst that is in a different physical state than the reactants.
ozone layer A concentration of ozone in the stratosphere (about 1015 ozone molecules per liter) that acts as a protective screen, absorbing ultraviolet light that would otherwise reach the surface of the earth, where it would harm plants and animals.
Ozone An unstable form of oxygen that consists of three oxygen atoms bonded together (O3). A layer of ozone in the stratosphere helps protect the plants and animals on earth from harmful ultraviolet radiation. Ozone is responsible for the pungent smell we associate with lightning discharges and electric motors. It is also toxic.
troposphere The lowest layer of the atmosphere, the troposphere extends from earth’s surface to an altitude of about 11–13 km (7–8 miles). The temperature of the troposphere decreases steadily with increasing altitude.
ultraviolet light High-energy radiation that cannot be detected by the human eye but can cause a wide variety of chemical reactions that are harmful to organisms.
chemical reactions A process in which a substance is converted to one or more other substances with different compositions and properties.
molecular mass The sum of the average masses of the atoms in one molecule of a substance, each multiplied by its subscript.
formula mass The sum of the atomic masses of all the elements in the empirical formula, each multiplied by its subscript.
mole (mol) The quantity of a substance that contains the same number of units (e.g., atoms or molecules) as the number of carbon atoms in exactly 12 g of isotopically pure carbon-12.
Avogadro’s number The number of units (e.g., atoms, molecules, or formula units) in 1 mol: $6.022×1023$.
molar mass The mass in grams of 1 mol of a substance.
percent composition The percentage of each element present in a pure substance. With few exceptions, the percent composition of a chemical compound is constant (see law of definite proportions).
product(s) The final compound(s) produced in a chemical reaction.
coefficient A number greater than 1 preceding a formula in a balanced chemical equation and indicating the number of atoms, molecules, or formula units of a reactant or a product.
mole ratio The ratio of the number of moles of one substance to the number of moles of another, as depicted by a balanced chemical equation.
stoichiometric quantity The amount of product or reactant specified by the coefficients in a balanced chemical equation.
limiting reactant The reactant that restricts the amount of product obtained in a chemical reaction.
theoretical yield The maximum amount of product that can be formed from the reactants in a chemical reaction, which theoretically is the amount of product that would be obtained if the reaction occurred perfectly and the method of purifying the product were 100% efficient.
percent yield The ratio of the actual yield of a reaction to the theoretical yield multiplied by 100 to give a percentage.
oxidation–reduction reactions A chemical reaction that exhibits a change in the oxidation states of one or more elements in the reactants that has the general form oxidant + reductant → reduced oxidant + oxidized reductant.
oxidation The loss of one or more electrons in a chemical reaction. The substance that loses electrons is said to be oxidized.
reduction The gain of one or more electrons in a chemical reaction. The substance that gains electrons is said to be reduced.
oxidation state The charge that each atom in a compound would have if all its bonding electrons were transferred to the atom with the greater attraction for electrons.
reductants (or reducing agents) A compound that is capable of donating electrons; thus it is oxidized.
combustion reaction An oxidation–reduction reaction in which the oxidant is $O2$.
catalysis The acceleration of a chemical reaction by a catalyst.
heterogeneous catalyst A catalyst that is in a different physical state than the reactants.
ozone layer A concentration of ozone in the stratosphere (about 1015 ozone molecules per liter) that acts as a protective screen, absorbing ultraviolet light that would otherwise reach the surface of the earth, where it would harm plants and animals.
Ozone An unstable form of oxygen that consists of three oxygen atoms bonded together (O3). A layer of ozone in the stratosphere helps protect the plants and animals on earth from harmful ultraviolet radiation. Ozone is responsible for the pungent smell we associate with lightning discharges and electric motors. It is also toxic.
troposphere The lowest layer of the atmosphere, the troposphere extends from earth’s surface to an altitude of about 11–13 km (7–8 miles). The temperature of the troposphere decreases steadily with increasing altitude.
ultraviolet light High-energy radiation that cannot be detected by the human eye but can cause a wide variety of chemical reactions that are harmful to organisms.
polar bond A chemical bond in which there is an unequal distribution of charge between the bonding atoms.
hydrated ions Individual cations and anions that are each surrounded by their own shell of water molecules.
strong electrolytes An electrolyte that dissociates completely into ions when dissolved in water, thus producing an aqueous solution that conducts electricity very well.
weak electrolytes A compound that produces relatively few ions when dissolved in water, thus producing an aqueous solution that conducts electricity poorly.
ketones A class of organic compounds with the general form RC(O)R’, in which the carbon atom of the carbonyl group is bonded to two alkyl groups (c.f. aldehyde). The alkyl groups may be the same or different.
concentration The quantity of solute that is dissolved in a particular quantity of solvent or solution.
molarity (M) A common unit of concentration that is the number of moles of solute present in exactly 1 L of solution $(mol/L).$
stock solution A commercially prepared solution of known concentration.
overall chemical equation A chemical equation that shows all the reactants and products as undissociated, electrically neutral compounds.
complete ionic equation A chemical equation that shows which ions and molecules are hydrated and which are present in other forms and phases.
net ionic equation A chemical equation that shows only those species that participate in the chemical reaction.
precipitate The insoluble product that forms in a precipitation reaction.
base A substance that produces one or more hydroxide ions $(OH−)$ and a cation when dissolved in aqueous solution, thereby forming a basic solution.
triprotic acid A compound that can donate three protons per molecule in separate steps.
equilibrium The point at which the rates of the forward and reverse reactions become the same, so that the net composition of the system no longer changes with time.
weak bases A base in which only a fraction of the molecules react with water to produce $OH−$ and the corresponding cation.
hydronium ion The $H3O+$ ion, represented as $H+(aq).$
amphoteric When substances can behave as both an acid and a base.
conjugate acid–base pairs An acid and a base that differ by only one hydrogen ion. All acid–base reactions involve two conjugate acid–base pairs, the Brønsted–Lowry acid and the base it forms after donating its proton, and the Brønsted–Lowry base and the acid it forms after accepting a proton.
salt The general term for any ionic substance that does not have $OH−$ as the anion or $H+$ as the cation.
pH scale A logarithmic scale used to express the hydrogen ion $(H+)$ concentration of a solution, making it possible to describe acidity or basicity quantitatively.
neutral solution A solution in which the total positive charge from all the cations is matched by an identical total negative charge from all the anions.
pH The negative base-10 logarithm of the hydrogen ion concentration:
indicators An intensely colored organic molecule whose color changes dramatically depending on the pH of the solution.
acid rain Precipitation that is dramatically more acidic because of human activities.
oxidation state method A procedure for balancing oxidation–reduction (redox) reactions in which the overall reaction is conceptually separated into two parts: an oxidation and a reduction.
single-displacement reactions A chemical reaction in which an ion in solution is displaced through oxidation of a metal.
inert metals The metals at the bottom of the activity series, which have the least tendency to be oxidized.
equivalence point The point in a titration where a stoichiometric amount (i.e., the amount required to react completely with the unknown) of the titrant has been added.
standard solution A solution whose concentration is precisely known.
endpoint The point in a titration at which an indicator changes color.
aqueous solution A solution in which water is the solvent.
polar bond A chemical bond in which there is an unequal distribution of charge between the bonding atoms.
hydrated ions Individual cations and anions that are each surrounded by their own shell of water molecules.
strong electrolytes An electrolyte that dissociates completely into ions when dissolved in water, thus producing an aqueous solution that conducts electricity very well.
weak electrolytes A compound that produces relatively few ions when dissolved in water, thus producing an aqueous solution that conducts electricity poorly.
ketones A class of organic compounds with the general form RC(O)R’, in which the carbon atom of the carbonyl group is bonded to two alkyl groups (c.f. aldehyde). The alkyl groups may be the same or different.
concentration The quantity of solute that is dissolved in a particular quantity of solvent or solution.
molarity (M) A common unit of concentration that is the number of moles of solute present in exactly 1 L of solution $(mol/L).$
stock solution A commercially prepared solution of known concentration.
overall chemical equation A chemical equation that shows all the reactants and products as undissociated, electrically neutral compounds.
complete ionic equation A chemical equation that shows which ions and molecules are hydrated and which are present in other forms and phases.
net ionic equation A chemical equation that shows only those species that participate in the chemical reaction.
precipitate The insoluble product that forms in a precipitation reaction.
base A substance that produces one or more hydroxide ions $(OH−)$ and a cation when dissolved in aqueous solution, thereby forming a basic solution.
triprotic acid A compound that can donate three protons per molecule in separate steps.
equilibrium The point at which the rates of the forward and reverse reactions become the same, so that the net composition of the system no longer changes with time.
weak bases A base in which only a fraction of the molecules react with water to produce $OH−$ and the corresponding cation.
hydronium ion The $H3O+$ ion, represented as $H+(aq).$
amphoteric When substances can behave as both an acid and a base.
conjugate acid–base pairs An acid and a base that differ by only one hydrogen ion. All acid–base reactions involve two conjugate acid–base pairs, the Brønsted–Lowry acid and the base it forms after donating its proton, and the Brønsted–Lowry base and the acid it forms after accepting a proton.
salt The general term for any ionic substance that does not have $OH−$ as the anion or $H+$ as the cation.
pH scale A logarithmic scale used to express the hydrogen ion $(H+)$ concentration of a solution, making it possible to describe acidity or basicity quantitatively.
neutral solution A solution in which the total positive charge from all the cations is matched by an identical total negative charge from all the anions.
pH The negative base-10 logarithm of the hydrogen ion concentration:
indicators An intensely colored organic molecule whose color changes dramatically depending on the pH of the solution.
acid rain Precipitation that is dramatically more acidic because of human activities.
oxidation state method A procedure for balancing oxidation–reduction (redox) reactions in which the overall reaction is conceptually separated into two parts: an oxidation and a reduction.
single-displacement reactions A chemical reaction in which an ion in solution is displaced through oxidation of a metal.
inert metals The metals at the bottom of the activity series, which have the least tendency to be oxidized.
equivalence point The point in a titration where a stoichiometric amount (i.e., the amount required to react completely with the unknown) of the titrant has been added.
standard solution A solution whose concentration is precisely known.
endpoint The point in a titration at which an indicator changes color.
thermochemistry A branch of chemistry that describes the energy changes that occur during chemical reactions.
chemical energy One of the five forms of energy, chemical energy is stored within a chemical compound because of a particular arrangement of atoms. The other four forms of energy are radiant, thermal, nuclear, and electrical.
kinetic energy (KE) Energy due to the motion of an object: $KE=12mv2,$ where $m$ is the mass of the object and $v$ is its velocity.
law of conservation of energy The total amount of energy in the universe remains constant. Energy can be neither created nor destroyed, but it can be converted from one form to another.
mechanical work The energy required to move an object a distance $d$ when opposed by a force $F$: $w=F×d.$
heat (q) Thermal energy that can be transformed from an object at one temperature to an object at another temperature.
joule (J) The SI unit of energy:
calories (cal) A non-SI unit of energy: 1 cal = 4.184 J exactly.
surroundings All the universe that is not the system; that is, system + surroundings = universe.
isolated system A system that can exchange neither energy nor matter with its suroundings.
state function A property of a system whose magnitude depends on only the present state of the system, not its previous history.
exothermic A process in which heat $(q)$ is transferred from a system to its surroundings.
endothermic A process in which heat $(q)$ is transferred to a system from its surroundings.
enthalpy (H) The sum of a system’s internal energy $E$ and the product of its pressure $P$ and volume $V$: $H=E+PV.$
change in enthalpy (ΔH) At constant pressure, the amount of heat transferred from the surroundings to the system or vice versa: .
enthalpy of reaction (ΔHrxn) The change in enthalpy that occurs during a chemical reaction.
Hess’s law The enthalpy change $(ΔH)$ for an overall reaction is the sum of the $ΔH$ values for the individual reactions.
Enthalpy of formation (ΔHf) The enthalpy change for the formation of 1 mol of a compound from its component elements.
standard enthalpies of formation ($ΔHfο$) The enthalpy change for the formation of 1 mol of a compound from its component elements when the component elements are each in their standard states. The standard enthalpy of formation of any element in its most stable form is zero by definition.
standard enthalpy of reaction ($ΔHrxnο$) The enthalpy change that occurs when a reaction is carried out with all reactants and products in their standard state.
Calorimetry A set of techniques used to measure enthalpy changes in chemical processes.
heat capacity (C) The amount of energy needed to raise the temperature of an object 1°C. The units of heat capacity are joules per degree Celsius $(J/°C).$
specific heat (Cs) The amount of energy needed to increase the temperature of 1 g of a substance by 1°C. The units of $Cs$ are $J/(g•°C).$
constant-pressure calorimeter A device used to measure enthalpy changes in chemical processes at constant pressure.
bomb calorimeter A device used to measure energy changes in chemical processes.
Calorie A unit used to indicate the caloric content of food. It is equal to 1 kilocalorie (1 kcal).
Coal A complex solid material derived primarily from plants that died and were buried hundreds of millions of years ago and were subsequently subjected to high temperatures and pressures. It is used as a fuel.
greenhouse effect The phenomenon in which substances absorb thermal energy radiated by Earth, thus trapping thermal energy in the atmosphere.
carbon cycle The distribution and flow of carbon throughout the planet.
greenhouse gases A substance that absorbs thermal energy radiated by Earth, thus trapping thermal energy in the atmosphere.
thermochemistry A branch of chemistry that describes the energy changes that occur during chemical reactions.
chemical energy One of the five forms of energy, chemical energy is stored within a chemical compound because of a particular arrangement of atoms. The other four forms of energy are radiant, thermal, nuclear, and electrical.
kinetic energy (KE) Energy due to the motion of an object: $KE=12mv2,$ where $m$ is the mass of the object and $v$ is its velocity.
law of conservation of energy The total amount of energy in the universe remains constant. Energy can be neither created nor destroyed, but it can be converted from one form to another.
mechanical work The energy required to move an object a distance $d$ when opposed by a force $F$: $w=F×d.$
heat (q) Thermal energy that can be transformed from an object at one temperature to an object at another temperature.
joule (J) The SI unit of energy:
calories (cal) A non-SI unit of energy: 1 cal = 4.184 J exactly.
surroundings All the universe that is not the system; that is, system + surroundings = universe.
isolated system A system that can exchange neither energy nor matter with its suroundings.
state function A property of a system whose magnitude depends on only the present state of the system, not its previous history.
exothermic A process in which heat $(q)$ is transferred from a system to its surroundings.
endothermic A process in which heat $(q)$ is transferred to a system from its surroundings.
enthalpy (H) The sum of a system’s internal energy $E$ and the product of its pressure $P$ and volume $V$: $H=E+PV.$
change in enthalpy (ΔH) At constant pressure, the amount of heat transferred from the surroundings to the system or vice versa: .
enthalpy of reaction (ΔHrxn) The change in enthalpy that occurs during a chemical reaction.
Hess’s law The enthalpy change $(ΔH)$ for an overall reaction is the sum of the $ΔH$ values for the individual reactions.
Enthalpy of formation (ΔHf) The enthalpy change for the formation of 1 mol of a compound from its component elements.
standard enthalpies of formation ($ΔHfο$) The enthalpy change for the formation of 1 mol of a compound from its component elements when the component elements are each in their standard states. The standard enthalpy of formation of any element in its most stable form is zero by definition.
standard enthalpy of reaction ($ΔHrxnο$) The enthalpy change that occurs when a reaction is carried out with all reactants and products in their standard state.
Calorimetry A set of techniques used to measure enthalpy changes in chemical processes.
heat capacity (C) The amount of energy needed to raise the temperature of an object 1°C. The units of heat capacity are joules per degree Celsius $(J/°C).$
specific heat (Cs) The amount of energy needed to increase the temperature of 1 g of a substance by 1°C. The units of $Cs$ are $J/(g•°C).$
constant-pressure calorimeter A device used to measure enthalpy changes in chemical processes at constant pressure.
bomb calorimeter A device used to measure energy changes in chemical processes.
Calorie A unit used to indicate the caloric content of food. It is equal to 1 kilocalorie (1 kcal).
Coal A complex solid material derived primarily from plants that died and were buried hundreds of millions of years ago and were subsequently subjected to high temperatures and pressures. It is used as a fuel.
greenhouse effect The phenomenon in which substances absorb thermal energy radiated by Earth, thus trapping thermal energy in the atmosphere.
carbon cycle The distribution and flow of carbon throughout the planet.
greenhouse gases A substance that absorbs thermal energy radiated by Earth, thus trapping thermal energy in the atmosphere.
wave A periodic oscillation that transmits energy through space.
speed (v) The distance traveled by a wave per unit time.
speed of light (c) The speed with which all forms of electromagnetic radiation travel in a vacuum.
blackbody radiation Electromagnetic radiation whose wavelength and color depends on the temperature of the object.
quantum The smallest possible unit of energy. Energy can be gained or lost only in integral multiples of a quantum.
photoelectric effect A phenomenon in which electrons are ejected from the surface of a metal that has been exposed to light.
photons A quantum of radiant energy, each of which possesses a particular energy $E$ given by $E=hν.$
line spectrum A spectrum in which light of only a certain wavelength is emitted or absorbed, rather than a continuous range of wavelengths.
excited state Any arrangement of electrons that is higher in energy than the ground state.
absorption spectrum A spectrum produced by the absorption of light by ground-state atoms.
wave–particle duality A principle that matter and energy have properties typical of both waves and particles.
fundamental The lowest-energy standing wave.
overtones The vibration of a standing wave that is higher in energy than the fundamental vibration.
nodes The point where the amplitude of a wave is zero.
Heisenberg uncertainty principle A principle stating that the uncertainty in the position of a particle $(Δx)$ multiplied by the uncertainty in its momentum $[Δ(mv)]$ is greater than or equal to Planck’s constant $(h)$ divided by 4π: $Δx[Δ(mv)]≥h/4π.$
quantum mechanics A theory developed by Erwin Schrödinger that describes the energies and spatial distributions of electrons in atoms and molecules.
wave function (Ψ) A mathematical function that relates the location of an electron at a given point in space to the amplitude of its wave, which corresponds to its energy.
principal quantum number (n) One of three quantum numbers that tells the average relative distance of an electron from the nucleus.
principal shell All the wave functions that have the same value of $n$ because those electrons have similar average distances from the nucleus.
azimuthal quantum number (l) One of three quantum numbers that discribes the shape of the region of space occupied by an electron.
subshell A group of wave functions that have the same values of $n$ and $l.$
magnetic quantum number (ml) One of three quantum numbers that describes the orientation of the region of space occupied by an electron with respect to an applied magnetic field.
atomic orbital A wave function with an allowed combination of $n$, $l.$, and $ml$ quantum numbers.
electron density Electron distributions that are represented as standing waves.
orbital energies A particular energy associated with a given set of quantum numbers.
degenerate Having the same energy.
electron shielding The effect by which electrons closer to the nucleus neutralize a portion of the positive charge of the nucleus and thereby decrease the attractive interaction between the nucleus and an electron father away.
electron spin The magnetic moment that results when an electron spins. Electrons have two possible orientations (spin up and spin down), which are described by a fourth quantum number (ms).
Pauli exclusion principle A principle stating that no two electrons in an atom can have the same value of all four quantum numbers.
electron configuration The arrangement of an element’s electrons in its atomic orbitals.
aufbau principle The process used to build up the periodic table by adding protons one by one to the nucleus and adding the corresponding electrons to the lowest-energy orbital available without violating the Pauli exclusion principle.
Hund’s rule A rule stating that the lowest-energy electron configuration for an atom is the one that has the maximum number of electrons with parallel spins in degenerate orbitals.
valence electrons Electrons in the outermost shell of an atom.
f block The elements in the periodic table in which the (n − 2)f orbitals are being filled.
wave A periodic oscillation that transmits energy through space.
speed (v) The distance traveled by a wave per unit time.
speed of light (c) The speed with which all forms of electromagnetic radiation travel in a vacuum.
blackbody radiation Electromagnetic radiation whose wavelength and color depends on the temperature of the object.
quantum The smallest possible unit of energy. Energy can be gained or lost only in integral multiples of a quantum.
photoelectric effect A phenomenon in which electrons are ejected from the surface of a metal that has been exposed to light.
photons A quantum of radiant energy, each of which possesses a particular energy $E$ given by $E=hν.$
line spectrum A spectrum in which light of only a certain wavelength is emitted or absorbed, rather than a continuous range of wavelengths.
excited state Any arrangement of electrons that is higher in energy than the ground state.
absorption spectrum A spectrum produced by the absorption of light by ground-state atoms.
wave–particle duality A principle that matter and energy have properties typical of both waves and particles.
fundamental The lowest-energy standing wave.
overtones The vibration of a standing wave that is higher in energy than the fundamental vibration.
nodes The point where the amplitude of a wave is zero.
Heisenberg uncertainty principle A principle stating that the uncertainty in the position of a particle $(Δx)$ multiplied by the uncertainty in its momentum $[Δ(mv)]$ is greater than or equal to Planck’s constant $(h)$ divided by 4π: $Δx[Δ(mv)]≥h/4π.$
quantum mechanics A theory developed by Erwin Schrödinger that describes the energies and spatial distributions of electrons in atoms and molecules.
wave function (Ψ) A mathematical function that relates the location of an electron at a given point in space to the amplitude of its wave, which corresponds to its energy.
principal quantum number (n) One of three quantum numbers that tells the average relative distance of an electron from the nucleus.
principal shell All the wave functions that have the same value of $n$ because those electrons have similar average distances from the nucleus.
azimuthal quantum number (l) One of three quantum numbers that discribes the shape of the region of space occupied by an electron.
subshell A group of wave functions that have the same values of $n$ and $l.$
magnetic quantum number (ml) One of three quantum numbers that describes the orientation of the region of space occupied by an electron with respect to an applied magnetic field.
atomic orbital A wave function with an allowed combination of $n$, $l.$, and $ml$ quantum numbers.
electron density Electron distributions that are represented as standing waves.
orbital energies A particular energy associated with a given set of quantum numbers.
degenerate Having the same energy.
electron shielding The effect by which electrons closer to the nucleus neutralize a portion of the positive charge of the nucleus and thereby decrease the attractive interaction between the nucleus and an electron father away.
electron spin The magnetic moment that results when an electron spins. Electrons have two possible orientations (spin up and spin down), which are described by a fourth quantum number (ms).
Pauli exclusion principle A principle stating that no two electrons in an atom can have the same value of all four quantum numbers.
electron configuration The arrangement of an element’s electrons in its atomic orbitals.
aufbau principle The process used to build up the periodic table by adding protons one by one to the nucleus and adding the corresponding electrons to the lowest-energy orbital available without violating the Pauli exclusion principle.
Hund’s rule A rule stating that the lowest-energy electron configuration for an atom is the one that has the maximum number of electrons with parallel spins in degenerate orbitals.
valence electrons Electrons in the outermost shell of an atom.
f block The elements in the periodic table in which the (n − 2)f orbitals are being filled.
triads A set of three elements that have similar properties.
octaves A group of seven elements, corresponding to the horizontal rows in the main group elements (not counting the noble gases, which were unknown at the time).
molar volume The molar mass of an element divided by its density.
covalent atomic radius (rcov) Half the distance between the nuclei of two like atoms joined by a covalent bond in the same molecule.
van der Waals atomic radius(rvdW) Half the internuclear distance between two nonbonded atoms in the solid.
ionic radius The radius of a cation or anion.
isoelectronic series A group of ions or atoms and ions that have the same number of electrons and thus the same ground-state electron configuration.
ionization energy(I) The minimum amount of energy needed to remove an electron from the gaseous atom in its ground state:
pseudo noble gas configurations The $(n−1)d10$ and similar electron configurations that are particularly stable and are often encountered in the heavier $p$-block elements.
electron affinity(EA) The energy change that occurs when an electron is added to a gaseous atom:
electronegativity The relative ability of an atom to attract electrons to itself in a chemical compound.
nanotubes One of at least four allotropes of carbon that are cylinders of carbon atoms and are intermediate in structure between graphite and the fullerenes.
pnicogens The elements in group 15 of the periodic table.
chalcogens The elements in group 16 of the periodic table.
transition metals Any element in groups 3–12 in the periodic table. All of the transition elements are metals.
actinides Any of the 14 elements between $Z=90$ (thorium) and $Z=103$ (lawrencium).
amplification mechanism A process by which elements that are present in trace amouts can exert large effects on the health of an organism.
Essential trace elements Elements that are required for the growth of most organisms.
macrominerals Any of the six essential elements (Na, Mg, K, Ca, Cl, and P) that provide essential ions in body fluids and form the major structural components of the body.
ion pumps A complex assembly of proteins that selectively transports ions across cell membranes toward the side with the higher concentration.
group-transfer reactions A reaction in which a recognizable functional group is transferred from one molecule to another.
triads A set of three elements that have similar properties.
octaves A group of seven elements, corresponding to the horizontal rows in the main group elements (not counting the noble gases, which were unknown at the time).
molar volume The molar mass of an element divided by its density.
covalent atomic radius (rcov) Half the distance between the nuclei of two like atoms joined by a covalent bond in the same molecule.
van der Waals atomic radius(rvdW) Half the internuclear distance between two nonbonded atoms in the solid.
ionic radius The radius of a cation or anion.
isoelectronic series A group of ions or atoms and ions that have the same number of electrons and thus the same ground-state electron configuration.
ionization energy(I) The minimum amount of energy needed to remove an electron from the gaseous atom in its ground state:
pseudo noble gas configurations The $(n−1)d10$ and similar electron configurations that are particularly stable and are often encountered in the heavier $p$-block elements.
electron affinity(EA) The energy change that occurs when an electron is added to a gaseous atom:
electronegativity The relative ability of an atom to attract electrons to itself in a chemical compound.
nanotubes One of at least four allotropes of carbon that are cylinders of carbon atoms and are intermediate in structure between graphite and the fullerenes.
pnicogens The elements in group 15 of the periodic table.
chalcogens The elements in group 16 of the periodic table.
transition metals Any element in groups 3–12 in the periodic table. All of the transition elements are metals.
actinides Any of the 14 elements between $Z=90$ (thorium) and $Z=103$ (lawrencium).
amplification mechanism A process by which elements that are present in trace amouts can exert large effects on the health of an organism.
Essential trace elements Elements that are required for the growth of most organisms.
macrominerals Any of the six essential elements (Na, Mg, K, Ca, Cl, and P) that provide essential ions in body fluids and form the major structural components of the body.
ion pumps A complex assembly of proteins that selectively transports ions across cell membranes toward the side with the higher concentration.
group-transfer reactions A reaction in which a recognizable functional group is transferred from one molecule to another.
ionic bonding A type of chemical bonding in which positively and negatively charged ions are held together by electrostatic forces.
bond energy The enthalpy change that occurs when a given bond in a gaseous molecule is broken.
bond distance(r0) The optimal internuclear distance between two bonded atoms.
melting point The temperature at which the individual ions in a lattice or the individual molecules in a covalent compound have enough kinetic energy to overcome the attractive forces that hold them together in the solid.
hardness The resistance of ionic materials to scratching or abrasion.
Born–Haber cycle A thermochemical cycle that decribes the process in which an ionic solid is conceptually formed from its component elements in a stepwise manner.
enthalpy of sublimation(ΔHsub) The enthalpy change that accompanies the conversion of a solid directly to a gas.
Lewis electron dot symbols A system that can be used to predict the number of bonds formed by most elements in their compounds.
octet rule The tendency for atoms to lose, gain, or share electrons to reach a total of eight valence electrons.
coordinate covalent bond A covalent bond in which both electrons come from the same atom.
formal charge The difference between the number of valence electrons in a free atom and the number of electrons assigned to it in a particular Lewis electron structure.
resonance structures A Lewis electron structure that has different arrangements of electrons around atoms whose positions do not change.
expanded-valence molecules A compound with more than an octet of electrons around an atom.
Lewis acid Any species that can accept a pair of electrons.
Electron-deficient molecules A compound that has less than an octet of electrons around one atom.
adduct The product of a reaction between a Lewis acid and a Lewis base with a coordinate covalent bond.
bond order The number of electron pairs that hold two atoms together.
polar covalent bonds A covalent bond in which the electrons are shared unequally between the bonded atoms.
dipole moment The product of the partial charge $Q$ on the bonded atoms and the distance $r$ between the partial charges: $µ=Qr$, where $Q$ is measured in coulombs (C) and $r$ in meters (m).
ionic bonding A type of chemical bonding in which positively and negatively charged ions are held together by electrostatic forces.
bond energy The enthalpy change that occurs when a given bond in a gaseous molecule is broken.
bond distance(r0) The optimal internuclear distance between two bonded atoms.
melting point The temperature at which the individual ions in a lattice or the individual molecules in a covalent compound have enough kinetic energy to overcome the attractive forces that hold them together in the solid.
hardness The resistance of ionic materials to scratching or abrasion.
Born–Haber cycle A thermochemical cycle that decribes the process in which an ionic solid is conceptually formed from its component elements in a stepwise manner.
enthalpy of sublimation(ΔHsub) The enthalpy change that accompanies the conversion of a solid directly to a gas.
Lewis electron dot symbols A system that can be used to predict the number of bonds formed by most elements in their compounds.
octet rule The tendency for atoms to lose, gain, or share electrons to reach a total of eight valence electrons.
coordinate covalent bond A covalent bond in which both electrons come from the same atom.
formal charge The difference between the number of valence electrons in a free atom and the number of electrons assigned to it in a particular Lewis electron structure.
resonance structures A Lewis electron structure that has different arrangements of electrons around atoms whose positions do not change.
expanded-valence molecules A compound with more than an octet of electrons around an atom.
Lewis acid Any species that can accept a pair of electrons.
Electron-deficient molecules A compound that has less than an octet of electrons around one atom.
adduct The product of a reaction between a Lewis acid and a Lewis base with a coordinate covalent bond.
bond order The number of electron pairs that hold two atoms together.
polar covalent bonds A covalent bond in which the electrons are shared unequally between the bonded atoms.
dipole moment The product of the partial charge $Q$ on the bonded atoms and the distance $r$ between the partial charges: $µ=Qr$, where $Q$ is measured in coulombs (C) and $r$ in meters (m).
valence-shell electron-pair repulsion (VSEPR) model A model used to predict the shapes of many molecules and polyatomic ions, based on the idea that the lowest-energy arrangement for a compound is the one in which its electron pairs (bonding and nonbonding) are as far apart as possible.
molecular geometry The arrangement of the bonded atoms in a molecule or a polyatomic ion in space.
valence bond theory A localized bonding model that assumes that the strength of a covalent bond is proportional to the amount of overlap between atomic orbitals and that an atom can use different combinations of atomic orbitals (hybrids) to maximize the overlap between bonded atoms.
promotion The excitation of an electron from a filled $ns2$ atomic orbital to an empty $np$ or $(n−1)d$ valence orbital.
hybrid atomic orbitals New atomic orbitals formed from the process of hybridization.
sp hybrid orbital The two equivalent hybrid orbitals that result when one $ns$ orbital and one $np$ orbital are combined (hybridized). The two $sp$ hybrid orbitals are oriented at 180° from each other. They are equivalent in energy, and their energy is between the energy values associated with pure $s$ and pure $p$ orbitals.
sp2 hybrid atomic orbitals The three equivalent hybrid orbitals that result when one $ns$ orbital and two $np$ orbitals are combined (hybridized). The three $sp2$ hybrid orbitals are oriented in a plane at 120° from each other. They are equivalent in energy, and their energy is between the energy values associated with pure $s$ and pure $p$ orbitals.
sp3 hybrid atomic orbitals The four equivalent hybrid orbitals that result when one $ns$ orbital and three $np$ orbitals are combined (hybridized). The four $sp3$ hybrid orbitals point at the vertices of a tetrahedron, so they are oriented at 109.5° from each other. They are equivalent in energy, and their energy is between the energy values associated with pure $s$ and pure $p$ orbitals.
sp3d hybrid orbitals The five hybrid orbitals that result when one $ns,$ three $np,$ and one $(n−1)d$ orbitals are combined (hybridized).
sp3d2 hybrid orbitals The six equivalent hybrid orbitals that result when one $ns$, three $np$, and two $(n−1)d$ orbitals are combined (hybridized).
molecular orbital theory A delocalized bonding model in which molecular orbitals are created from the linear combination of atomic orbitals (LCAOs).
linear combinations of atomic orbitals (LCAOs) Molecular orbitals created from the sum and the difference of two wave functions (atomic orbitals).
sigma (σ) orbital A bonding molecular orbital in which the electron density along the internuclear axis and between the nuclei has cylindrical symmetry.
sigma star (σ*) orbital An antibonding molecular orbital in which there is a region of zero electron probability (a nodal plane) perpendicular to the internuclear axis.
antibonding molecular orbital A molecular orbital that forms when atomic orbitals or orbital lobes of opposite sign interact to give decreased electron probability between the nuclei due to destructuve reinforcement of the wave functions.
energy-level diagram A schematic drawing that compares the energies of the molecular orbitals (bonding, antibonding, and nonbonding) with the energies of the parent atomic orbitals.
homonuclear diatomic molecule A molecule that consists of two atoms of the same element.
bond order One-half the net number of bonding electrons in a molecule.
pi star (π*) orbital An antibonding molecular orbital formed from the difference of the side-to-side interactions of two or more parallel $np$ atomic orbitals, creating a nodal plane perpendicular to the internuclear axis.
law of conservation of orbitals A law that states that the number of molecular orbitals produced is the same as the number of atomic orbitals used to create them.
heteronuclear diatomic molecules A molecule that consists of two atoms of different elements.
nonbonding molecular orbitals A molecular orbital that forms when atomic orbitals or orbital lobes interact only very weakly, creating essentially no change in the electron probability density between the nuclei.
valence-shell electron-pair repulsion (VSEPR) model A model used to predict the shapes of many molecules and polyatomic ions, based on the idea that the lowest-energy arrangement for a compound is the one in which its electron pairs (bonding and nonbonding) are as far apart as possible.
molecular geometry The arrangement of the bonded atoms in a molecule or a polyatomic ion in space.
valence bond theory A localized bonding model that assumes that the strength of a covalent bond is proportional to the amount of overlap between atomic orbitals and that an atom can use different combinations of atomic orbitals (hybrids) to maximize the overlap between bonded atoms.
promotion The excitation of an electron from a filled $ns2$ atomic orbital to an empty $np$ or $(n−1)d$ valence orbital.
hybrid atomic orbitals New atomic orbitals formed from the process of hybridization.
sp hybrid orbital The two equivalent hybrid orbitals that result when one $ns$ orbital and one $np$ orbital are combined (hybridized). The two $sp$ hybrid orbitals are oriented at 180° from each other. They are equivalent in energy, and their energy is between the energy values associated with pure $s$ and pure $p$ orbitals.
sp2 hybrid atomic orbitals The three equivalent hybrid orbitals that result when one $ns$ orbital and two $np$ orbitals are combined (hybridized). The three $sp2$ hybrid orbitals are oriented in a plane at 120° from each other. They are equivalent in energy, and their energy is between the energy values associated with pure $s$ and pure $p$ orbitals.
sp3 hybrid atomic orbitals The four equivalent hybrid orbitals that result when one $ns$ orbital and three $np$ orbitals are combined (hybridized). The four $sp3$ hybrid orbitals point at the vertices of a tetrahedron, so they are oriented at 109.5° from each other. They are equivalent in energy, and their energy is between the energy values associated with pure $s$ and pure $p$ orbitals.
sp3d hybrid orbitals The five hybrid orbitals that result when one $ns,$ three $np,$ and one $(n−1)d$ orbitals are combined (hybridized).
sp3d2 hybrid orbitals The six equivalent hybrid orbitals that result when one $ns$, three $np$, and two $(n−1)d$ orbitals are combined (hybridized).
molecular orbital theory A delocalized bonding model in which molecular orbitals are created from the linear combination of atomic orbitals (LCAOs).
linear combinations of atomic orbitals (LCAOs) Molecular orbitals created from the sum and the difference of two wave functions (atomic orbitals).
sigma (σ) orbital A bonding molecular orbital in which the electron density along the internuclear axis and between the nuclei has cylindrical symmetry.
sigma star (σ*) orbital An antibonding molecular orbital in which there is a region of zero electron probability (a nodal plane) perpendicular to the internuclear axis.
antibonding molecular orbital A molecular orbital that forms when atomic orbitals or orbital lobes of opposite sign interact to give decreased electron probability between the nuclei due to destructuve reinforcement of the wave functions.
energy-level diagram A schematic drawing that compares the energies of the molecular orbitals (bonding, antibonding, and nonbonding) with the energies of the parent atomic orbitals.
homonuclear diatomic molecule A molecule that consists of two atoms of the same element.
bond order One-half the net number of bonding electrons in a molecule.
pi star (π*) orbital An antibonding molecular orbital formed from the difference of the side-to-side interactions of two or more parallel $np$ atomic orbitals, creating a nodal plane perpendicular to the internuclear axis.
law of conservation of orbitals A law that states that the number of molecular orbitals produced is the same as the number of atomic orbitals used to create them.
heteronuclear diatomic molecules A molecule that consists of two atoms of different elements.
nonbonding molecular orbitals A molecular orbital that forms when atomic orbitals or orbital lobes interact only very weakly, creating essentially no change in the electron probability density between the nuclei.
pressure(P) The amount of force $(F)$ exerted on a given area $(A)$ of surface: $P=F/A.$
pascal (Pa) The SI unit for pressure. The pascal is newtons per square meter: $N/m2.$
barometer A device used to measure atmospheric pressure.
atmosphere (atm) Also referred to as standard atmospheric pressure, it is the atmospheric pressure required to support a column of mercury exactly 760 mm tall.
manometers A device used to measure the pressures of samples of gases contained in an apparatus.
Boyle’s law A law that states that at constant temperature, the volume of a fixed amount of a gas is inversely proportional to its pressure.
absolute zero (0 K) The lowest possible temperature that can be theoretically achieved; it corresponds to −273.15°C.
Charles’s law A law that states that at constant pressure, the volume of a fixed amount of gas is directly proportional to its absolute temperature (in kelvins).
Avogadro’s law A law that states that at constant temperature and pressure, the volume of a sample of gas is directly proportional to the number of moles of gas in the sample.
gas constant A proportionality constant that is used in the ideal gas law.
ideal gas law A law relating pressure, temperature, volume, and the amount of an ideal gas.
ideal gas A hypothetical gaseous substance whose behavior is independent of attractive and repulsive forces.
standard temperature and pressure (STP) The conditions 0°C (273.15 K) and 1 atm pressure for a gas.
standard molar volume The volume of 1 mol of an ideal gas at STP (0°C and 1 atm pressure), which is 22.41 L.
partial pressure The pressure a gas in a mixture would exert if it were the only one present (at the same temperature and volume).
Dalton’s law of partial pressures A law that states that the total pressure exerted by a mixture of gases is the sum of the partial pressures of component gases.
mole fraction (X) The ratio of the number of moles of any component of a mixture to the total number of moles of all species present in the mixture.
kinetic molecular theory of gases A theory that describes, on the molecular level, why ideal gases behave the way they do.
root mean square (rms) speed(vrms) The speed of a gas particle that has average kinetic energy.
Boltzmann distributions A curve that shows the distribution of molecular speeds at a given temperature.
effusion The escape of a gas through a small (usually microscopic) opening into an evacuated space.
Graham’s law A law that states that the rate of effusion of a gaseous substance is inversely proportional to the square root of its molar mass.
mean free path The average distance traveled by a molecule between collisions.
van der Waals equation A modification of the ideal gas law designed to describe the behavior of real gases by explicitly including the effects of molecular volume and intermolecular forces.
Liquefaction The condensation of gases into a liquid form.
cryogenic liquids An ultracold liquid formed from the liquefaction of gases.
pressure(P) The amount of force $(F)$ exerted on a given area $(A)$ of surface: $P=F/A.$
pascal (Pa) The SI unit for pressure. The pascal is newtons per square meter: $N/m2.$
barometer A device used to measure atmospheric pressure.
atmosphere (atm) Also referred to as standard atmospheric pressure, it is the atmospheric pressure required to support a column of mercury exactly 760 mm tall.
manometers A device used to measure the pressures of samples of gases contained in an apparatus.
Boyle’s law A law that states that at constant temperature, the volume of a fixed amount of a gas is inversely proportional to its pressure.
absolute zero (0 K) The lowest possible temperature that can be theoretically achieved; it corresponds to −273.15°C.
Charles’s law A law that states that at constant pressure, the volume of a fixed amount of gas is directly proportional to its absolute temperature (in kelvins).
Avogadro’s law A law that states that at constant temperature and pressure, the volume of a sample of gas is directly proportional to the number of moles of gas in the sample.
gas constant A proportionality constant that is used in the ideal gas law.
ideal gas law A law relating pressure, temperature, volume, and the amount of an ideal gas.
ideal gas A hypothetical gaseous substance whose behavior is independent of attractive and repulsive forces.
standard temperature and pressure (STP) The conditions 0°C (273.15 K) and 1 atm pressure for a gas.
standard molar volume The volume of 1 mol of an ideal gas at STP (0°C and 1 atm pressure), which is 22.41 L.
partial pressure The pressure a gas in a mixture would exert if it were the only one present (at the same temperature and volume).
Dalton’s law of partial pressures A law that states that the total pressure exerted by a mixture of gases is the sum of the partial pressures of component gases.
mole fraction (X) The ratio of the number of moles of any component of a mixture to the total number of moles of all species present in the mixture.
kinetic molecular theory of gases A theory that describes, on the molecular level, why ideal gases behave the way they do.
root mean square (rms) speed(vrms) The speed of a gas particle that has average kinetic energy.
Boltzmann distributions A curve that shows the distribution of molecular speeds at a given temperature.
effusion The escape of a gas through a small (usually microscopic) opening into an evacuated space.
Graham’s law A law that states that the rate of effusion of a gaseous substance is inversely proportional to the square root of its molar mass.
mean free path The average distance traveled by a molecule between collisions.
van der Waals equation A modification of the ideal gas law designed to describe the behavior of real gases by explicitly including the effects of molecular volume and intermolecular forces.
Liquefaction The condensation of gases into a liquid form.
cryogenic liquids An ultracold liquid formed from the liquefaction of gases.
van der Waals forces The intermolecular forces known as dipole–dipole interactions and London dispersion forces.
dipole–dipole interactions A kind of intermolecular interaction (force) that results between molecules with net dipole moments.
London dispersion forces A kind of intermolecular interaction (force) that results from temporary fluctuations in the electron distribution within atoms and nonpolar molecules.
induced dipole A short-lived dipole moment that is created in atoms and nonpolar molecules adjacent to atoms or molecules with an instantaneous dipole moment.
polarizability The ease of deformation of the electron distribution in an atom or molecule.
hydrogen bonds An unusually strong dipole-dipole interaction (intermolecular force) that results when hydrogen is bonded to very electronegative elements, such as O, N, and F.
surface tension The energy required to increase the surface area of a liquid by a certain amount. Surface tension is measured in units of energy per area (e.g., $J/m2$).
surfactants Substances (surface-active agents), such as soaps and detergents, that disrupt the attractive intermolecular interactions between molecules of a polar liquid, thereby reducing the surface tension of the liquid.
capillary action The tendency of a polar liquid to rise against gravity into a small-diameter glass tube.
meniscus The upper surface of the liquid in a tube.
Viscosity (η) The resistance of a liquid to flow.
vapor pressure The pressure created over a liquid by the molecules of a liquid substance that have enough kinetic energy to escape to the vapor phase.
evaporation (or vaporization) The physical process by which atoms or molecules in the liquid phase enter the gas or vapor phase.
condensation The physical process by which atoms or molecules in the vapor phase enter the liquid phase.
equilibrium vapor pressure The pressure exerted by a vapor in dynamic equilibrium with its liquid.
nonvolatile liquids A liquid with a relatively low vapor pressure.
Clausius–Clapeyron equation A linear relationship that expresses the nonlinear relationship between the vapor pressure of a liquid and temperature: ln $P=−ΔHvap/RT+C,$ where $P$ is pressure, $ΔHvap$ is the heat of vaporization, $R$ is the universal gas constant, $T$ is the absolute temperature, and C is a constant. The Clausius–Clapeyron equation can be used to calculate the heat of vaporization of a liquid from its measured vapor pressure at two or more temperatures.
normal boiling point The temperature at which a substance boils at a pressure of 1 atm.
phase changes A change of state that occurs when any of the three forms of matter (solids, liquids, and gases) is converted to either of the other two.
fusion (or melting) The conversion of a solid to a liquid.
enthalpy of sublimation (ΔHsub) The enthalpy change that accompanies the conversion of a solid directly to a gas.
specific heat (Cs) The number of joules required to raise the temperature of 1 g of a substance by 1°C.
superheated liquid An unstable liquid at a temperature and pressure at which it should be a gas.
seed crystal A solid sample of a substance that can be added to a supercooled liquid or a supersaturated solution to help induce crystallization.
critical point The combination of the critical temperature and the critical pressure of a substance.
supercritical fluid The single, dense fluid phase that exists above the critical temperature of a substance.
molten salt A salt that has been heated to its melting point.
ionic liquids Ionic substances that are liquids at room temperature and pressure and that consist of small, symmetrical anions combined with larger, symmetrical organic cations that prevent the formation of a highly organized structure.
phase diagram A graphic summary of the physical state of a substance as a function of temperature and pressure in a closed system.
triple point The point in a phase diagram where the solid/liquid, liquid/gas, and solid/gas lines intersect; it represents the only combination of temperature and pressure at which all three phases are in equilibrium and can therefore exist simultaneously.
liquid crystals A substance that exhibits phases that have properties intermediate between those of a crystalline solid and a normal liquid and possess long-range molecular order but still flow.
anisotropic An arrangement of molecules in which their properties depend on the direction they are measured.
cholesteric phase One of three different ways that most liquid crystals can orient themselves. The molecules are arranged in planes (similar to the smectic phase), but each layer is rotated by a certain amount with respect to those above and below it, giving it a helical structure.
van der Waals forces The intermolecular forces known as dipole–dipole interactions and London dispersion forces.
dipole–dipole interactions A kind of intermolecular interaction (force) that results between molecules with net dipole moments.
London dispersion forces A kind of intermolecular interaction (force) that results from temporary fluctuations in the electron distribution within atoms and nonpolar molecules.
induced dipole A short-lived dipole moment that is created in atoms and nonpolar molecules adjacent to atoms or molecules with an instantaneous dipole moment.
polarizability The ease of deformation of the electron distribution in an atom or molecule.
hydrogen bonds An unusually strong dipole-dipole interaction (intermolecular force) that results when hydrogen is bonded to very electronegative elements, such as O, N, and F.
surface tension The energy required to increase the surface area of a liquid by a certain amount. Surface tension is measured in units of energy per area (e.g., $J/m2$).
surfactants Substances (surface-active agents), such as soaps and detergents, that disrupt the attractive intermolecular interactions between molecules of a polar liquid, thereby reducing the surface tension of the liquid.
capillary action The tendency of a polar liquid to rise against gravity into a small-diameter glass tube.
meniscus The upper surface of the liquid in a tube.
Viscosity (η) The resistance of a liquid to flow.
vapor pressure The pressure created over a liquid by the molecules of a liquid substance that have enough kinetic energy to escape to the vapor phase.
evaporation (or vaporization) The physical process by which atoms or molecules in the liquid phase enter the gas or vapor phase.
condensation The physical process by which atoms or molecules in the vapor phase enter the liquid phase.
equilibrium vapor pressure The pressure exerted by a vapor in dynamic equilibrium with its liquid.
nonvolatile liquids A liquid with a relatively low vapor pressure.
Clausius–Clapeyron equation A linear relationship that expresses the nonlinear relationship between the vapor pressure of a liquid and temperature: ln $P=−ΔHvap/RT+C,$ where $P$ is pressure, $ΔHvap$ is the heat of vaporization, $R$ is the universal gas constant, $T$ is the absolute temperature, and C is a constant. The Clausius–Clapeyron equation can be used to calculate the heat of vaporization of a liquid from its measured vapor pressure at two or more temperatures.
normal boiling point The temperature at which a substance boils at a pressure of 1 atm.
phase changes A change of state that occurs when any of the three forms of matter (solids, liquids, and gases) is converted to either of the other two.
fusion (or melting) The conversion of a solid to a liquid.
enthalpy of sublimation (ΔHsub) The enthalpy change that accompanies the conversion of a solid directly to a gas.
specific heat (Cs) The number of joules required to raise the temperature of 1 g of a substance by 1°C.
superheated liquid An unstable liquid at a temperature and pressure at which it should be a gas.
seed crystal A solid sample of a substance that can be added to a supercooled liquid or a supersaturated solution to help induce crystallization.
critical point The combination of the critical temperature and the critical pressure of a substance.
supercritical fluid The single, dense fluid phase that exists above the critical temperature of a substance.
molten salt A salt that has been heated to its melting point.
ionic liquids Ionic substances that are liquids at room temperature and pressure and that consist of small, symmetrical anions combined with larger, symmetrical organic cations that prevent the formation of a highly organized structure.
phase diagram A graphic summary of the physical state of a substance as a function of temperature and pressure in a closed system.
triple point The point in a phase diagram where the solid/liquid, liquid/gas, and solid/gas lines intersect; it represents the only combination of temperature and pressure at which all three phases are in equilibrium and can therefore exist simultaneously.
liquid crystals A substance that exhibits phases that have properties intermediate between those of a crystalline solid and a normal liquid and possess long-range molecular order but still flow.
anisotropic An arrangement of molecules in which their properties depend on the direction they are measured.
cholesteric phase One of three different ways that most liquid crystals can orient themselves. The molecules are arranged in planes (similar to the smectic phase), but each layer is rotated by a certain amount with respect to those above and below it, giving it a helical structure.
amorphous solid A solid with no particular structural order.
glass An amorphous, translucent solid. A glass is a solid that has been cooled too quickly to form ordered crystals.
unit cell The smallest repeating unit of a crystal lattice.
face-centered cubic (fcc) A cubic unit cell with eight component atoms, molecules, or ions located at the corners of a cube plus an identical component in the center of each face of the cube.
cubic close-packed (ccp) structure One of two variants of the close-packed arrangement—the most efficient way to pack spheres in a lattice—in which the atomic positions alter from layer to layer in an ABCABC… pattern.
coordination number The number of nearest neighbors in a solid structure.
cesium chloride structure The unit cell for many ionic compounds with relatively large cations and a 1:1 cation:anion ratio.
tetrahedral holes One of two kinds of holes in a face-centered cubic array of atoms or ions (the other is an octahedral hole). Tetrahedral holes are located between an atom at a corner and the three atoms at the centers of the adjacent faces of the face-centered cubic unit cell. An atom or ion in a tetrahedral hole has a coordination number of 4.
sodium chloride structure The solid structure that results when the octahedral holes of an fcc lattice of anions are filled with cations.
zinc blende structure The solid structure that results when half of the tetrahedral holes in an fcc lattice of anions are filled with cations with a 1:1 cation:anion ratio and a coordination number of 4.
perovskite structure A structure that consists of a bcc array of two metal ions, with one set (M) located at the corners of the cube, and the other set (M′) in the centers of the cube.
x-ray diffraction An technique used to obtain information about the structures of crystalline substances by using x-rays.
Bragg equation The equation that describes the relationship between two x-ray beams diffracted from different planes of atoms:
defects Errors in an idealized crystal lattice.
vacancy A point defect that consists of a single atom missing from a site in a crystal.
interstitial impurity A point defect that results when an impurity atom occupies an octahedral hole or a tetrahedral hole in the lattice between atoms.
substitutional impurity A point defect that results when an impurity atom occupies a normal lattice site.
edge dislocation A crystal defect that results from the insertion of an extra plane of atoms into part of the crystal lattice.
Deformation A distortion that occurs when a dislocation moves through a crystal.
pinning A process that increases the mechanical strength of a material by introducing multiple defects into a material so that the presence of one defect prevents the motion of another.
grain boundary The place where two grains in a solid intersect.
Work hardening The practice of introducing a dense network of dislocations throughout a solid, making it very tough and hard.
Schottky defects A coupled pair of vacancies—one cation and one anion—that maintains the electrical neutrality of an ionic solid.
Frenkel defect A defect in an ionic lattice that occurs when one of the ions is in the wrong position.
solid electrolytes A solid material with a very high electrical conductivity.
nonstoichiometric compounds A solid that has intrinsically variable stoichiometries without affecting the fundamental structure of the crystal.
ionic solid A solid that consists of positively and negatively charged ions held together by electrostatic forces.
Molecular solids A solid that consists of molecules held together by relatively weak forces, such as dipole-dipole interactions, hydrogen bonds, and London dispersion forces.
Covalent solids A solid that consists of two- or three-dimensional networks of atoms held together by covalent bonds.
metallic solids A solid that consists of metal atoms held together by metallic bonds.
electron sea Valence electrons that are delocalized throughout a metallic solid.
interstitial alloy An alloy formed by inserting smaller atoms into holes in the metal lattice.
intermetallic compounds An alloy that consists of certain metals that combine in only specific proportions and whose properties are frequently quite different from those of their constituent elements.
band theory A theory used to describe the bonding in metals and semiconductors.
bandwidth The difference in energy between the highest and lowest energy levels in an energy band.
band gap The difference in energy between the highest level of one energy band and the lowest level of the band above it, which represents a set of forbidden energies that do not correspond to any allowed combinations of atomic orbitals.
overlapping bands Molecular orbitals derived from two or more different kinds of valence electrons that have similar energies.
electrical insulators A material that conducts electricity poorly because its valence bands are full.
conduction band The band of empty molecular orbitals in a semiconductor.
semiconductors A substance such as Si and Ge that has a conductivity between that of metals and insulators.
n-type semiconductor A semiconductor that has been doped with an impurity that has more valence electrons than the atoms of the host lattice.
p-type semiconductor A semiconductor that has been doped with an impurity that has fewer valence electrons than the atoms of the host lattice.
Meissner effect The phenomenon in which a superconductor completely expels a magnetic field from its interior.
Cooper pairs Pairs of electrons that migrate through a superconducting material as a unit.
high-temperature superconductors A material that becomes a superconductor at temperatures greater than 30 K.
Plastic The property of a material that allows it to be molded into almost any shape.
enzymes Catalysts that occur naturally in living organisms and that catalyze biological reactions.
fibers A particle of a synthetic polymer that is more than 100 times longer than it is wide.
pyrolysis A high-temperature decomposition reaction that can be used to form fibers of synthetic polymers.
ceramic Any nonmetallic inorganic solid that is strong eneough to be used in structural applications.
sintering A process that fuses the grains of a ceramic into a dense, strong material. Sintering is used to produce high-strength ceramics.
sol-gel process A process used to manufacture ceramics by producing fine powders of ceramic oxides with uniformly sized particles.
Superalloys A high-strength alloy based on cobalt, nickel, and iron, often of complex composition, that is used in applications that require mechanical strength, high surface stability, and resistance to high temperatures.
Composite materials A material that consists of at least two distinct phases: the matrix (which constitutes the bulk of the material) and fibers or granules that are embedded within the matrix.
polymer-matrix composite A compositie that consists of reinforcing fibers embedded in a polymer matrix.
Metal-matrix composites A composite that consists of reinforcing fibers embedded in a metal or a metal alloy matrix.
Ceramic-matrix composites A composite consisting of reinforcing fibers embedded in a ceramic matrix.
amorphous solid A solid with no particular structural order.
glass An amorphous, translucent solid. A glass is a solid that has been cooled too quickly to form ordered crystals.
unit cell The smallest repeating unit of a crystal lattice.
face-centered cubic (fcc) A cubic unit cell with eight component atoms, molecules, or ions located at the corners of a cube plus an identical component in the center of each face of the cube.
cubic close-packed (ccp) structure One of two variants of the close-packed arrangement—the most efficient way to pack spheres in a lattice—in which the atomic positions alter from layer to layer in an ABCABC… pattern.
coordination number The number of nearest neighbors in a solid structure.
cesium chloride structure The unit cell for many ionic compounds with relatively large cations and a 1:1 cation:anion ratio.
tetrahedral holes One of two kinds of holes in a face-centered cubic array of atoms or ions (the other is an octahedral hole). Tetrahedral holes are located between an atom at a corner and the three atoms at the centers of the adjacent faces of the face-centered cubic unit cell. An atom or ion in a tetrahedral hole has a coordination number of 4.
sodium chloride structure The solid structure that results when the octahedral holes of an fcc lattice of anions are filled with cations.
zinc blende structure The solid structure that results when half of the tetrahedral holes in an fcc lattice of anions are filled with cations with a 1:1 cation:anion ratio and a coordination number of 4.
perovskite structure A structure that consists of a bcc array of two metal ions, with one set (M) located at the corners of the cube, and the other set (M′) in the centers of the cube.
x-ray diffraction An technique used to obtain information about the structures of crystalline substances by using x-rays.
Bragg equation The equation that describes the relationship between two x-ray beams diffracted from different planes of atoms:
defects Errors in an idealized crystal lattice.
vacancy A point defect that consists of a single atom missing from a site in a crystal.
interstitial impurity A point defect that results when an impurity atom occupies an octahedral hole or a tetrahedral hole in the lattice between atoms.
substitutional impurity A point defect that results when an impurity atom occupies a normal lattice site.
edge dislocation A crystal defect that results from the insertion of an extra plane of atoms into part of the crystal lattice.
Deformation A distortion that occurs when a dislocation moves through a crystal.
pinning A process that increases the mechanical strength of a material by introducing multiple defects into a material so that the presence of one defect prevents the motion of another.
grain boundary The place where two grains in a solid intersect.
Work hardening The practice of introducing a dense network of dislocations throughout a solid, making it very tough and hard.
Schottky defects A coupled pair of vacancies—one cation and one anion—that maintains the electrical neutrality of an ionic solid.
Frenkel defect A defect in an ionic lattice that occurs when one of the ions is in the wrong position.
solid electrolytes A solid material with a very high electrical conductivity.
nonstoichiometric compounds A solid that has intrinsically variable stoichiometries without affecting the fundamental structure of the crystal.
ionic solid A solid that consists of positively and negatively charged ions held together by electrostatic forces.
Molecular solids A solid that consists of molecules held together by relatively weak forces, such as dipole-dipole interactions, hydrogen bonds, and London dispersion forces.
Covalent solids A solid that consists of two- or three-dimensional networks of atoms held together by covalent bonds.
metallic solids A solid that consists of metal atoms held together by metallic bonds.
electron sea Valence electrons that are delocalized throughout a metallic solid.
interstitial alloy An alloy formed by inserting smaller atoms into holes in the metal lattice.
intermetallic compounds An alloy that consists of certain metals that combine in only specific proportions and whose properties are frequently quite different from those of their constituent elements.
band theory A theory used to describe the bonding in metals and semiconductors.
bandwidth The difference in energy between the highest and lowest energy levels in an energy band.
band gap The difference in energy between the highest level of one energy band and the lowest level of the band above it, which represents a set of forbidden energies that do not correspond to any allowed combinations of atomic orbitals.
overlapping bands Molecular orbitals derived from two or more different kinds of valence electrons that have similar energies.
electrical insulators A material that conducts electricity poorly because its valence bands are full.
conduction band The band of empty molecular orbitals in a semiconductor.
semiconductors A substance such as Si and Ge that has a conductivity between that of metals and insulators.
n-type semiconductor A semiconductor that has been doped with an impurity that has more valence electrons than the atoms of the host lattice.
p-type semiconductor A semiconductor that has been doped with an impurity that has fewer valence electrons than the atoms of the host lattice.
Meissner effect The phenomenon in which a superconductor completely expels a magnetic field from its interior.
Cooper pairs Pairs of electrons that migrate through a superconducting material as a unit.
high-temperature superconductors A material that becomes a superconductor at temperatures greater than 30 K.
Plastic The property of a material that allows it to be molded into almost any shape.
enzymes Catalysts that occur naturally in living organisms and that catalyze biological reactions.
fibers A particle of a synthetic polymer that is more than 100 times longer than it is wide.
pyrolysis A high-temperature decomposition reaction that can be used to form fibers of synthetic polymers.
ceramic Any nonmetallic inorganic solid that is strong eneough to be used in structural applications.
sintering A process that fuses the grains of a ceramic into a dense, strong material. Sintering is used to produce high-strength ceramics.
sol-gel process A process used to manufacture ceramics by producing fine powders of ceramic oxides with uniformly sized particles.
Superalloys A high-strength alloy based on cobalt, nickel, and iron, often of complex composition, that is used in applications that require mechanical strength, high surface stability, and resistance to high temperatures.
Composite materials A material that consists of at least two distinct phases: the matrix (which constitutes the bulk of the material) and fibers or granules that are embedded within the matrix.
polymer-matrix composite A compositie that consists of reinforcing fibers embedded in a polymer matrix.
Metal-matrix composites A composite that consists of reinforcing fibers embedded in a metal or a metal alloy matrix.
Ceramic-matrix composites A composite consisting of reinforcing fibers embedded in a ceramic matrix.
miscible Capable of forming a single homogeneous phase, regardless of the proportions with which the substances are mixed.
hydration The process of surrounding solute particles with water molecules.
entropy(S) The degree of disorder in a thermodynamic system. The greater the number of possible microstates for a system, the higher the entropy.
solubility A measure of the how much of a solid substance remains dissolved in a given amount of a specified liquid at a specified temperature and pressure.
saturated A solution with the maximum possible amount of a solute under a given set of conditions.
seed crystal A solid sample of a substance that can be added to a supercooled liquid or a supersaturated solution to help induce crystallization.
hydrophobic A substance that repels water. Hydrophobic substances do not interact favorably with water.
amalgams A solution (usually a solid solution) of a metal in liquid mercury.
dielectric constant (ε) A constant that expresses the ability of a bulk substance to decrease the electrostatic forces between two charged particles.
crown ethers Cyclic polyether with four or more oxygen atoms separated by two or three carbon atoms. All crown ethers have a central cavity that can accommodate a metal ion coordinated to the ring of oxygen atoms.
Cryptands Consisting of three $(–OCH2CH2O–)n$ chains connected by two nitrogen atoms, cryptands have a central cavity that can encapsulate a metal ion coordinated to the oxygen and nitrogen atoms.
concentration The quantity of solute that is dissolved in a particular quantity of solvent or solution.
molality (m) The number of moles of solute present in exactly 1 kg of solvent.
parts per billion (ppb) Micrograms of solute per kilogram of solvent.
parts per thousand (ppt) Grams of solute per kilogram of solvent, primarily used in the health sciences.
fractional crystallization The separation of compounds based on their relative solubilities in a given solvent.
Henry’s law An equation that quantifies the relationship between the pressure and the solubility of a gas: $C=kP.$
colligative properties A property of a solution that depends primarily on the number of solute particles rather than the kind of solute particles.
Raoult’s law An equation that quantifies the relationship between solution composition and vapor pressure: $PA=XAPA0.$
ideal solution A solution that obeys Raoult’s law.
boiling point elevation (ΔTb) The difference between the boiling point of a solution and the boiling point of the pure solvent.
freezing point depression (ΔTf) The difference between the freezing point of a pure solvent and the freezing point of the solution.
osmosis The net flow of solvent through a semipermeable membrane.
osmotic pressure (Π) The pressure difference between the two sides of a semipermeable membrane that separates a pure solvent from a solution prepared from the same solvent.
dialysis A process that uses a semipermeable membrane with pores large enough to allow small solute molecules and solvent molecules to pass through but not large solute molecules.
reverse osmosis A process that uses the application of an external pressure greater than the osmotic pressure of a solution to reverse the flow of solvent through the semipermeable membrane.
van’t Hoff factor(i) The ratio of the apparent number of particles in solution to the number predicted by the stoichiometry of the salt.
ion pairs A cation and anion that are in intimate contact in solution rather than separated by solvent and that migrates in solution as a single unit.
colloid A heterogeneous mixture of particles with diameters of about 2–500 nm that are distributed throughout a second phase and do not separate from the dispersing phase on standing.
aerosols A dispersion of solid or liquid particles in a gas.
Tyndall effect The phenomenon of scattering a beam of visible light.
Emulsions A dispersion of one liquid phase in another liquid with which it is immiscible.
micelles A spherical or cylindrical aggregate of detergents or soaps in water that minimizes contact between the hydrophobic tails of the detergents or soaps and water.
bilayers A two-dimensional sheet consisting of a double layer of phospholipid molecules arranged tail to tail.
cell A collection of molecules, capable of reproducing itself, that is surrounded by a phospholipid bilayer.
solutions A homogeneous mixture of two or more substances in which the substances present in lesser amounts (the solutes) are dispersed uniformly throughout the substance present in greater amount (the solvent).
miscible Capable of forming a single homogeneous phase, regardless of the proportions with which the substances are mixed.
hydration The process of surrounding solute particles with water molecules.
entropy(S) The degree of disorder in a thermodynamic system. The greater the number of possible microstates for a system, the higher the entropy.
solubility A measure of the how much of a solid substance remains dissolved in a given amount of a specified liquid at a specified temperature and pressure.
saturated A solution with the maximum possible amount of a solute under a given set of conditions.
seed crystal A solid sample of a substance that can be added to a supercooled liquid or a supersaturated solution to help induce crystallization.
hydrophobic A substance that repels water. Hydrophobic substances do not interact favorably with water.
amalgams A solution (usually a solid solution) of a metal in liquid mercury.
dielectric constant (ε) A constant that expresses the ability of a bulk substance to decrease the electrostatic forces between two charged particles.
crown ethers Cyclic polyether with four or more oxygen atoms separated by two or three carbon atoms. All crown ethers have a central cavity that can accommodate a metal ion coordinated to the ring of oxygen atoms.
Cryptands Consisting of three $(–OCH2CH2O–)n$ chains connected by two nitrogen atoms, cryptands have a central cavity that can encapsulate a metal ion coordinated to the oxygen and nitrogen atoms.
concentration The quantity of solute that is dissolved in a particular quantity of solvent or solution.
molality (m) The number of moles of solute present in exactly 1 kg of solvent.
parts per billion (ppb) Micrograms of solute per kilogram of solvent.
parts per thousand (ppt) Grams of solute per kilogram of solvent, primarily used in the health sciences.
fractional crystallization The separation of compounds based on their relative solubilities in a given solvent.
Henry’s law An equation that quantifies the relationship between the pressure and the solubility of a gas: $C=kP.$
colligative properties A property of a solution that depends primarily on the number of solute particles rather than the kind of solute particles.
Raoult’s law An equation that quantifies the relationship between solution composition and vapor pressure: $PA=XAPA0.$
ideal solution A solution that obeys Raoult’s law.
boiling point elevation (ΔTb) The difference between the boiling point of a solution and the boiling point of the pure solvent.
freezing point depression (ΔTf) The difference between the freezing point of a pure solvent and the freezing point of the solution.
osmosis The net flow of solvent through a semipermeable membrane.
osmotic pressure (Π) The pressure difference between the two sides of a semipermeable membrane that separates a pure solvent from a solution prepared from the same solvent.
dialysis A process that uses a semipermeable membrane with pores large enough to allow small solute molecules and solvent molecules to pass through but not large solute molecules.
reverse osmosis A process that uses the application of an external pressure greater than the osmotic pressure of a solution to reverse the flow of solvent through the semipermeable membrane.
van’t Hoff factor(i) The ratio of the apparent number of particles in solution to the number predicted by the stoichiometry of the salt.
ion pairs A cation and anion that are in intimate contact in solution rather than separated by solvent and that migrates in solution as a single unit.
colloid A heterogeneous mixture of particles with diameters of about 2–500 nm that are distributed throughout a second phase and do not separate from the dispersing phase on standing.
aerosols A dispersion of solid or liquid particles in a gas.
Tyndall effect The phenomenon of scattering a beam of visible light.
Emulsions A dispersion of one liquid phase in another liquid with which it is immiscible.
micelles A spherical or cylindrical aggregate of detergents or soaps in water that minimizes contact between the hydrophobic tails of the detergents or soaps and water.
bilayers A two-dimensional sheet consisting of a double layer of phospholipid molecules arranged tail to tail.
cell A collection of molecules, capable of reproducing itself, that is surrounded by a phospholipid bilayer.
reaction rates The changes in concentrations of reactants and products with time.
average reaction rate The reaction rate calculated for a given time interval from the concentrations of either the reactant or one of the products at the beginning of the interval time $(t0)$ and at the end of the interval $(t1).$
instantaneous rate The reaction rate of a chemical reaction at any given point in time.
rate laws Mathematical expressions that describe the relationships between reactant rates and reactant concentrations in a chemical reaction.
integrated rate law A rate law that expresses the reaction rate in terms of the initial concentration $([R0])$ and the measured concentration of one or more reactants ([R]) after a given amount of time $(t).$
rate constant A proportionality constant whose value is characteristic of the reaction and the reaction conditions and whose numerical value does not change as the reaction progresses under a given set of conditions.
reaction order Numbers that indicate the degree to which the reaction rate depends on the concentration of each reactant.
zeroth-order reaction A reaction whose rate is independent of concentration.
enzyme A catalyst that occurs naturally in living organisms and catalyzes biological reactions.
first-order reaction A reaction whose rate is directly proportional to the concentration of one reactant.
second-order reaction A reaction whose rate is proportional to the square of the concentration of the reactant (for a reaction with the general form 2A → products) or is proportional to the product of the concentrations of two reactants (for a reaction with the general form A + B → products).
reaction mechanisms The sequence of events that occur at the molecular level during a reaction.
half-life The period of time it takes for the concentration of a reactant to decrease to one-half its initial value.
activity (A) The decrease in the number of a radioisotope’s nuclei per unit time: $A=−ΔN/Δt.$
elementary reaction Each of the complex series of reactions that take place in a stepwise fashion to convert reactants to products.
intermediate A species in a reaction mechanism that does not appear in the balanced chemical equation for the overall reaction.
molecularity The number of molecules that collide during any step in a reaction mechanism.
rate-determining step The slowest step in a reaction mechanism.
chain reactions A reaction mechanism in which one or more elementary reactions that contain a highly reactive species repeat again and again during the reaction process.
radicals Species that have one or more unpaired valence electrons.
activation energy (Ea) The energy barrier or threshold that corresponds to the minimum amount of energy the particles in a reaction must have to react when they colllide.
transition state Also called the activated complex, the arrangement of atoms that first forms when molecules are able to overcome the activation energy and react.
steric factor (p) The fraction of orientations of particles that result in a chemical reaction.
frequency factor A constant in the Arrhenius equation, it converts concentrations to collisions per second.
Arrhenius equation An expression that summarizes the collision model of chemical kinetics: $k=Ae−Ea/RT.$
catalysts A substance that participates in a reaction and causes it to occur more rapidly but that can be recovered unchanged at the end of the reaction and reused. Catalysts may also control which products are formed in a reaction.
heterogeneous catalysis A catalytic reaction in which the catalyst is in a different phase from the reactants.
homogeneous catalysis A catalytic reaction in which the catalyst is uniformly dispersed throughout the reactant mixture to form a solution.
substrate The reactant in an enzyme-catalyzed reaction.
Enzyme inhibitors Substances that decrease the reaction rate of an enzyme-catalyzed reaction by binding to a specific portion of the enzyme, thus slowing or preventing a reaction from occurring.
reaction rates The changes in concentrations of reactants and products with time.
average reaction rate The reaction rate calculated for a given time interval from the concentrations of either the reactant or one of the products at the beginning of the interval time $(t0)$ and at the end of the interval $(t1).$
instantaneous rate The reaction rate of a chemical reaction at any given point in time.
rate laws Mathematical expressions that describe the relationships between reactant rates and reactant concentrations in a chemical reaction.
integrated rate law A rate law that expresses the reaction rate in terms of the initial concentration $([R0])$ and the measured concentration of one or more reactants ([R]) after a given amount of time $(t).$
rate constant A proportionality constant whose value is characteristic of the reaction and the reaction conditions and whose numerical value does not change as the reaction progresses under a given set of conditions.
reaction order Numbers that indicate the degree to which the reaction rate depends on the concentration of each reactant.
zeroth-order reaction A reaction whose rate is independent of concentration.
enzyme A catalyst that occurs naturally in living organisms and catalyzes biological reactions.
first-order reaction A reaction whose rate is directly proportional to the concentration of one reactant.
second-order reaction A reaction whose rate is proportional to the square of the concentration of the reactant (for a reaction with the general form 2A → products) or is proportional to the product of the concentrations of two reactants (for a reaction with the general form A + B → products).
reaction mechanisms The sequence of events that occur at the molecular level during a reaction.
half-life The period of time it takes for the concentration of a reactant to decrease to one-half its initial value.
activity (A) The decrease in the number of a radioisotope’s nuclei per unit time: $A=−ΔN/Δt.$
elementary reaction Each of the complex series of reactions that take place in a stepwise fashion to convert reactants to products.
intermediate A species in a reaction mechanism that does not appear in the balanced chemical equation for the overall reaction.
molecularity The number of molecules that collide during any step in a reaction mechanism.
rate-determining step The slowest step in a reaction mechanism.
chain reactions A reaction mechanism in which one or more elementary reactions that contain a highly reactive species repeat again and again during the reaction process.
radicals Species that have one or more unpaired valence electrons.
activation energy (Ea) The energy barrier or threshold that corresponds to the minimum amount of energy the particles in a reaction must have to react when they colllide.
transition state Also called the activated complex, the arrangement of atoms that first forms when molecules are able to overcome the activation energy and react.
steric factor (p) The fraction of orientations of particles that result in a chemical reaction.
frequency factor A constant in the Arrhenius equation, it converts concentrations to collisions per second.
Arrhenius equation An expression that summarizes the collision model of chemical kinetics: $k=Ae−Ea/RT.$
catalysts A substance that participates in a reaction and causes it to occur more rapidly but that can be recovered unchanged at the end of the reaction and reused. Catalysts may also control which products are formed in a reaction.
heterogeneous catalysis A catalytic reaction in which the catalyst is in a different phase from the reactants.
homogeneous catalysis A catalytic reaction in which the catalyst is uniformly dispersed throughout the reactant mixture to form a solution.
substrate The reactant in an enzyme-catalyzed reaction.
Enzyme inhibitors Substances that decrease the reaction rate of an enzyme-catalyzed reaction by binding to a specific portion of the enzyme, thus slowing or preventing a reaction from occurring.
equilibrium constant (K) The ratio of the rate constants for the forward reaction and the reverse reaction; that is, $K=kf/kr.$ It is also the equilibrium constant calculated from solution concentrations: $K=[C]c[D]d/[A]a[B]b$ for the general reaction $aA+bB⇌cC+dD,$ in which each component is in solution.
law of mass action For the general balanced chemical equation $aA+bB⇌cC+dD,$ the equilibrium constant expression is $K=[C]c[D]d/[A]a[B]b.$
equilibrium constant expression For a balanced chemical equation, the ratio is $[C]c[D]d/[A]a[B]b$ for the general reaction $aA+bB⇌cC+dD.$
Kp An equilibrium constant expressed as the ratio of the partial pressures of the products and reactants, each raised to its coefficient in the chemical equation.
heterogeneous equilibrium An equilibrium in which the reactants of an equilibrium reaction, the products, or both are in more than one phase.
reaction quotient (Q) A quantity derived from a set of values measured at any time during the reaction of any mixture of reactants and products, regardless of whether the system is at equilibrium: $Q=[C]c[D]d/[A]a[B]b$ for the general balanced chemical equation $aA+bB⇌cC+dD.$
reaction quotient (Qp) A quantity derived from a set of values measured at any time during the reaction of any mixture of reactants and products in the gas phase, regardless of whether the system is at equilibrium: $Qp=(PC)c(PD)d/(PA)a(PB)b$ for the general balanced chemical equation $aA+bB⇌cC+dD.$
Le Châtelier’s principle If a stress is applied to a system at equilibrium, the composition of the system will change to relieve the applied stress.
thermodynamic control The altering of reaction conditions so that a single desired product or set of products is present in significant quantities at equilibrium.
chemical equilibrium The point at which the forward and reverse reaction rates become the same so that the net composition of the system no longer changes with time.
equilibrium constant (K) The ratio of the rate constants for the forward reaction and the reverse reaction; that is, $K=kf/kr.$ It is also the equilibrium constant calculated from solution concentrations: $K=[C]c[D]d/[A]a[B]b$ for the general reaction $aA+bB⇌cC+dD,$ in which each component is in solution.
law of mass action For the general balanced chemical equation $aA+bB⇌cC+dD,$ the equilibrium constant expression is $K=[C]c[D]d/[A]a[B]b.$
equilibrium constant expression For a balanced chemical equation, the ratio is $[C]c[D]d/[A]a[B]b$ for the general reaction $aA+bB⇌cC+dD.$
Kp An equilibrium constant expressed as the ratio of the partial pressures of the products and reactants, each raised to its coefficient in the chemical equation.
heterogeneous equilibrium An equilibrium in which the reactants of an equilibrium reaction, the products, or both are in more than one phase.
reaction quotient (Q) A quantity derived from a set of values measured at any time during the reaction of any mixture of reactants and products, regardless of whether the system is at equilibrium: $Q=[C]c[D]d/[A]a[B]b$ for the general balanced chemical equation $aA+bB⇌cC+dD.$
reaction quotient (Qp) A quantity derived from a set of values measured at any time during the reaction of any mixture of reactants and products in the gas phase, regardless of whether the system is at equilibrium: $Qp=(PC)c(PD)d/(PA)a(PB)b$ for the general balanced chemical equation $aA+bB⇌cC+dD.$
Le Châtelier’s principle If a stress is applied to a system at equilibrium, the composition of the system will change to relieve the applied stress.
thermodynamic control The altering of reaction conditions so that a single desired product or set of products is present in significant quantities at equilibrium.
amphiprotic Substances that can behave as either an acid or a base in a chemical reaction, depending on the nature of the other reactant(s).
ion-product constant of liquid water (Kw) An equilibrium constant for the autoionization of water, $2H2O(l)$$H3O+(aq)$ + $OH−(aq),$ in which the concentration of water is treated as a constant: $Kw$ = $[H3O+][OH−]$ = $1.006×10−14.$
conjugate acid–base pair An acid and a base that differ by only one hydrogen ion.
acid ionization constant (Ka) An equilibrium constant for the ionization (dissociation) of a weak acid (HA) with water, $HA(aq)$ + $H2O(l)$$H3O(aq)$ + $A−(aq),$ in which the concentration of water is treated as a constant: $Ka$ = $[H3O+][A−]/[HA].$
base ionization constant (Kb) An equilibrium constant for the reaction of a weak base (B) with water, $B(aq)$ + $H2O(l)$$BH+(aq)$ + $OH−(aq),$ in which the concentration of water is treated as a constant: $Kb$ = $[BH+][OH−]/[B].$
leveling effect The phenomenon that makes $H3O+$ the strongest acid that can exist in water. Any species that is a stronger acid than $H3O+$ is leveled to the strength of $H3O+$ in aqueous solution.
hydrolysis reactions A chemical reaction in which a salt reacts with water to yield an acidic or a basic solution.
titration curve A plot of the pH of the solution being titrated versus the amount of acid or base (of known concentration) added.
equivalence point The point in a titration where a stoichiometric amount of the titrant has been added.
midpoint The point in an acid–base titration at which exactly enough acid (or base) has been added to neutralize one-half of the base (or the acid) originally present: $[HA]=[A−].$
acid–base indicator A compound added in small amounts to an acid–base titration to signal the equivalence point by changing color.
Buffers Solutions that maintain a relatively constant pH when an acid or a base is added.
common ion effect The shift in equilibrium that results when a strong electrolyte containing one ion in common with a reaction system that is at equilibrium is added to the system.
buffer capacity The amount of strong acid or strong base that a buffer solution can absorb before the pH changes dramatically.
Henderson-Hasselbalch equation A rearranged version of the equilibrium constant expression that provides a direct way to calculate the pH of a buffer solution: pH = $pKa$ + log([base]/[acid]).
amphiprotic Substances that can behave as either an acid or a base in a chemical reaction, depending on the nature of the other reactant(s).
ion-product constant of liquid water (Kw) An equilibrium constant for the autoionization of water, $2H2O(l)$$H3O+(aq)$ + $OH−(aq),$ in which the concentration of water is treated as a constant: $Kw$ = $[H3O+][OH−]$ = $1.006×10−14.$
conjugate acid–base pair An acid and a base that differ by only one hydrogen ion.
acid ionization constant (Ka) An equilibrium constant for the ionization (dissociation) of a weak acid (HA) with water, $HA(aq)$ + $H2O(l)$$H3O(aq)$ + $A−(aq),$ in which the concentration of water is treated as a constant: $Ka$ = $[H3O+][A−]/[HA].$
leveling effect The phenomenon that makes $H3O+$ the strongest acid that can exist in water. Any species that is a stronger acid than $H3O+$ is leveled to the strength of $H3O+$ in aqueous solution.
hydrolysis reactions A chemical reaction in which a salt reacts with water to yield an acidic or a basic solution.
titration curve A plot of the pH of the solution being titrated versus the amount of acid or base (of known concentration) added.
equivalence point The point in a titration where a stoichiometric amount of the titrant has been added.
midpoint The point in an acid–base titration at which exactly enough acid (or base) has been added to neutralize one-half of the base (or the acid) originally present: $[HA]=[A−].$
acid–base indicator A compound added in small amounts to an acid–base titration to signal the equivalence point by changing color.
Buffers Solutions that maintain a relatively constant pH when an acid or a base is added.
common ion effect The shift in equilibrium that results when a strong electrolyte containing one ion in common with a reaction system that is at equilibrium is added to the system.
buffer capacity The amount of strong acid or strong base that a buffer solution can absorb before the pH changes dramatically.
Henderson-Hasselbalch equation A rearranged version of the equilibrium constant expression that provides a direct way to calculate the pH of a buffer solution: pH = $pKa$ + log([base]/[acid]).
solubility product (Ksp) The equilibrium constant expression for the dissolution of a sparingly soluble salt that includes the concentration of a pure solid, which is a constant.
ion product (Q) A quantity that has precisely the same form as the solubility product for the dissolution of a sparingly soluble salt, except that the concentrations used are not necessarily equilibrium concentrations.
ion pair A cation and an anion that are in intimate contact in solution rather than separated by solvent. An ion pair can be viewed as a species that is intermediate between the ionic solid and the completely dissociated ions in solution.
ligands An ion or a molecule that contains one or more pairs of electrons that can be shared with the central metal in a metal complex.
formation constant (Kf) The equilibrium constant for the formation of a complex ion from a hydrated metal ion; that is, for the reaction $aA+bB⇌cC+dD,$ $Kf=[C]c[D]d/[A]a[B]b.$
Acidic oxides An oxide that reacts with water to produce an acidic solution or dissolves in aqueous base.
amphoteric oxides An oxide that can dissolve in acid to produce water and dissolve in base to produce a soluble complex.
qualitative analysis A procedure for determining the identity of metal ions present in a mixture that does not include information about their amounts.
solubility product (Ksp) The equilibrium constant expression for the dissolution of a sparingly soluble salt that includes the concentration of a pure solid, which is a constant.
ion product (Q) A quantity that has precisely the same form as the solubility product for the dissolution of a sparingly soluble salt, except that the concentrations used are not necessarily equilibrium concentrations.
ion pair A cation and an anion that are in intimate contact in solution rather than separated by solvent. An ion pair can be viewed as a species that is intermediate between the ionic solid and the completely dissociated ions in solution.
ligands An ion or a molecule that contains one or more pairs of electrons that can be shared with the central metal in a metal complex.
formation constant (Kf) The equilibrium constant for the formation of a complex ion from a hydrated metal ion; that is, for the reaction $aA+bB⇌cC+dD,$ $Kf=[C]c[D]d/[A]a[B]b.$
Acidic oxides An oxide that reacts with water to produce an acidic solution or dissolves in aqueous base.
amphoteric oxides An oxide that can dissolve in acid to produce water and dissolve in base to produce a soluble complex.
qualitative analysis A procedure for determining the identity of metal ions present in a mixture that does not include information about their amounts.
state function A property of a system whose magnitude depends on only the present state of the system, not its previous history.
internal energy (E) A state function that is the sum of the kinetic and potential energies of all a system’s components.
first law of thermodynamics The energy of the universe is constant: $ΔEuniverse$ = $ΔEsystem$ + $ΔEsurroundings$ = 0.
enthalpy (H) A state function that is the sum of the system’s internal energy $E$ and the product of its pressure $P$ and volume $V:$ $H=E+PV.$
entropy (S) The degree of disorder in a thermodynamic system, which is directly proportional to the possible number of microstates.
irreversible process A process in which the intermediate states between the extremes are not equilibrium states, so change occurs spontaneously in only one direction.
second law of thermodynamics The entropy of the universe remains constant in a reversible process, whereas the entropy of the universe increases in an irreversible (spontaneous) process.
third law of thermodynamics The entropy of any perfectly ordered, crystalline substance at absolute zero is zero.
standard molar entropy (S°) The entropy of 1 mol of a substance at a standard temperature of 298 K.
Gibbs free energy (G) A state function that is defined in terms of three other state functions—namely, enthalpy $(H),$ entropy $(S),$ and temperature $(T):$ $G=H−TS.$
standard free-energy change (ΔG°) The change in free energy when one substance or a set of substances in their standard states is converted to one or more other sustances, also in their standard states: $ΔG°=ΔH°−TΔS°.$
standard free energy of formation $(ΔGf°)$ The change in free energy that occurs when 1 mol of a substance in its standard state is formed from the component elements in their standard states: $ΔG°f=ΔH°f−TΔS°f.$
photosynthesis The fundamental reaction by which all green plants and algae obtain energy from sunlight in which $CO2$ is photochemically reduced to a carbon compound such as glucose. Oxygen in water is concurrently oxidized to $O2.$
respiration A process by which chemotrophs obtain energy from their environment; the overall reaction of respiration is the reverse of photosynthesis. Respiration is the combustion of a carbon compound such as glucose to $CO2$ and water.
fermentation A process used by some chemotrophs to obtain energy from their environment; a chemical reaction in which both the oxidant and the reductant are organic compounds.
thermodynamics The study of the interrelationships among heat, work, and the energy content of a system at equilibrium.
state function A property of a system whose magnitude depends on only the present state of the system, not its previous history.
internal energy (E) A state function that is the sum of the kinetic and potential energies of all a system’s components.
first law of thermodynamics The energy of the universe is constant: $ΔEuniverse$ = $ΔEsystem$ + $ΔEsurroundings$ = 0.
enthalpy (H) A state function that is the sum of the system’s internal energy $E$ and the product of its pressure $P$ and volume $V:$ $H=E+PV.$
entropy (S) The degree of disorder in a thermodynamic system, which is directly proportional to the possible number of microstates.
irreversible process A process in which the intermediate states between the extremes are not equilibrium states, so change occurs spontaneously in only one direction.
second law of thermodynamics The entropy of the universe remains constant in a reversible process, whereas the entropy of the universe increases in an irreversible (spontaneous) process.
third law of thermodynamics The entropy of any perfectly ordered, crystalline substance at absolute zero is zero.
standard molar entropy (S°) The entropy of 1 mol of a substance at a standard temperature of 298 K.
Gibbs free energy (G) A state function that is defined in terms of three other state functions—namely, enthalpy $(H),$ entropy $(S),$ and temperature $(T):$ $G=H−TS.$
standard free-energy change (ΔG°) The change in free energy when one substance or a set of substances in their standard states is converted to one or more other sustances, also in their standard states: $ΔG°=ΔH°−TΔS°.$
standard free energy of formation $(ΔGf°)$ The change in free energy that occurs when 1 mol of a substance in its standard state is formed from the component elements in their standard states: $ΔG°f=ΔH°f−TΔS°f.$
photosynthesis The fundamental reaction by which all green plants and algae obtain energy from sunlight in which $CO2$ is photochemically reduced to a carbon compound such as glucose. Oxygen in water is concurrently oxidized to $O2.$
respiration A process by which chemotrophs obtain energy from their environment; the overall reaction of respiration is the reverse of photosynthesis. Respiration is the combustion of a carbon compound such as glucose to $CO2$ and water.
fermentation A process used by some chemotrophs to obtain energy from their environment; a chemical reaction in which both the oxidant and the reductant are organic compounds.
electrochemistry The study of the relationship between electricity and chemical reactions.
oxidant A substance that is capable of accepting electrons and in the process is reduced.
half-reactions Reactions that represent either the oxidation half or the reduction half of an oxidation–reduction (redox) reaction.
electrochemical cell An apparatus that generates electricity from a spontaneous oxidation–reduction (redox) reaction or, conversely, uses electricity to drive a nonspontaneous redox reaction.
cathode One of two electrodes in an electrochemical cell, it is the site of the reduction half-reaction.
salt bridge A U-shaped tube inserted into both solutions of a galvanic cell that contains a concentrated liquid or gelled electrolyte and completes the circuit between the anode and the cathode.
potential (Ecell) Related to the energy needed to move a charged particle in an electric field, it is the difference in electrical potential beween two half-reactions.
standard cell potential The potential of an electrochemical cell measured under standard conditions (1 M for solutions, 1 atm for gases, and pure solids or pure liquids for other substances) and at a fixed temperature (usually 298 K).
standard hydrogen electrode (SHE) The electrode chosen as the reference for all other electrodes, which has been assigned a standard potential of 0 V and consists of a Pt wire in contact with an aqueous solution that contains 1 M in equilibrium with $H2$ gas at a pressure of 1 atm at the Pt-solution interface.
standard electrode potential The potential of a half-reaction measured against the SHE under standard conditions.
reference electrode An electrode in an galvanic cell whose potential is unaffected by the properties of the solution.
silver–silver chloride electrode A reference electrode that consists of a silver wire coated with a very thin layer of AgCl and dipped into a chloride ion solution with a fixed concentration.
saturated calomel electrode (SCE) A reference electrode that consists of a platinum wire inserted into a moist paste of liquid mercury (calomel; $Hg2Cl2$) and KCl in an interior cell, which is surrounded by an aqueous KCl solution.
glass electrode An electrode used to measure the $H+$ ion concentration of a solution and consisting of an internal Ag/AgCl electrode immersed in a 1 M HCl solution that is separated from the solution by a very thin glass membrane.
Ion-selective electrodes An electrode whose potential depends on only the concentration of a particular species in solution.
amperes (A) The fundamental SI unit of electric current; it is defined as the flow of 1 C/s past a given point: 1A = 1 C/s.
faraday (F) The charge on 1 mol of electrons; it is obtained by multiplying the charge on the electron by Avogadro’s number.
Nernst equation An equation for calculating cell potentials $(Ecell)$ under nonstandard conditions; it can be used to determine the direction of spontaneous reaction for any redox reaction under an conditions:
concentration cell An electrochemical cell in which the anode and the cathode compartments are identical except for the concentration of a reactant.
fuel cell A galvanic cell that requires a constant external supply of one or more reactants to generate electricity.
Leclanché dry cell A battery consisting of an electrolyte that is an acidic water-based paste containing $MnO2,$ $NH4Cl,$ $ZnCl2,$ graphite, and starch.
alkaline battery A battery that consists of a Leclanché cell adapted to operate under alkaline (basic) conditions.
lithium–iodine battery A battery that consists of an anode of lithium metal and a cathode containing a solid complex of $I2,$ with a layer of solid LiI in between that allows the diffusion of $Li+$ ions.
nickel–cadmium A type of battery that consists of a water-based cell with a cadmium anode and a highly oxidized nickel cathode.
lead–acid battery A battery consisting of a plate or grid of spongy lead metal, a cathode containing powdered $PbO2,$ and an electrolyte that is usually an aqueous solution of $H2SO4.$
Corrosion A galvanic process by which metals deteriorate through oxidation—usually but not always to their oxides.
sacrificial electrodes An electrode containing a more reactive metal that is attached to a metal object to inhibit that object’s corrosion.
electrolysis An electrochemical process in which an external voltage is applied to an electrolytic cell to drive a nonspontaneous reaction.
overvoltage The voltage that must be applied in electrolysis in addition to the calculated (theoretical) value to overcome factors such as a high activation energy and the formation of bubbles on a surface.
electroplating A process in which a layer of a second metal is deposited on the metal electrode that acts as the cathode during electrolysis.
electrochemistry The study of the relationship between electricity and chemical reactions.
oxidant A substance that is capable of accepting electrons and in the process is reduced.
half-reactions Reactions that represent either the oxidation half or the reduction half of an oxidation–reduction (redox) reaction.
electrochemical cell An apparatus that generates electricity from a spontaneous oxidation–reduction (redox) reaction or, conversely, uses electricity to drive a nonspontaneous redox reaction.
cathode One of two electrodes in an electrochemical cell, it is the site of the reduction half-reaction.
salt bridge A U-shaped tube inserted into both solutions of a galvanic cell that contains a concentrated liquid or gelled electrolyte and completes the circuit between the anode and the cathode.
potential (Ecell) Related to the energy needed to move a charged particle in an electric field, it is the difference in electrical potential beween two half-reactions.
standard cell potential The potential of an electrochemical cell measured under standard conditions (1 M for solutions, 1 atm for gases, and pure solids or pure liquids for other substances) and at a fixed temperature (usually 298 K).
standard hydrogen electrode (SHE) The electrode chosen as the reference for all other electrodes, which has been assigned a standard potential of 0 V and consists of a Pt wire in contact with an aqueous solution that contains 1 M in equilibrium with $H2$ gas at a pressure of 1 atm at the Pt-solution interface.
standard electrode potential The potential of a half-reaction measured against the SHE under standard conditions.
reference electrode An electrode in an galvanic cell whose potential is unaffected by the properties of the solution.
silver–silver chloride electrode A reference electrode that consists of a silver wire coated with a very thin layer of AgCl and dipped into a chloride ion solution with a fixed concentration.
saturated calomel electrode (SCE) A reference electrode that consists of a platinum wire inserted into a moist paste of liquid mercury (calomel; $Hg2Cl2$) and KCl in an interior cell, which is surrounded by an aqueous KCl solution.
glass electrode An electrode used to measure the $H+$ ion concentration of a solution and consisting of an internal Ag/AgCl electrode immersed in a 1 M HCl solution that is separated from the solution by a very thin glass membrane.
Ion-selective electrodes An electrode whose potential depends on only the concentration of a particular species in solution.
amperes (A) The fundamental SI unit of electric current; it is defined as the flow of 1 C/s past a given point: 1A = 1 C/s.
faraday (F) The charge on 1 mol of electrons; it is obtained by multiplying the charge on the electron by Avogadro’s number.
Nernst equation An equation for calculating cell potentials $(Ecell)$ under nonstandard conditions; it can be used to determine the direction of spontaneous reaction for any redox reaction under an conditions:
concentration cell An electrochemical cell in which the anode and the cathode compartments are identical except for the concentration of a reactant.
fuel cell A galvanic cell that requires a constant external supply of one or more reactants to generate electricity.
Leclanché dry cell A battery consisting of an electrolyte that is an acidic water-based paste containing $MnO2,$ $NH4Cl,$ $ZnCl2,$ graphite, and starch.
alkaline battery A battery that consists of a Leclanché cell adapted to operate under alkaline (basic) conditions.
lithium–iodine battery A battery that consists of an anode of lithium metal and a cathode containing a solid complex of $I2,$ with a layer of solid LiI in between that allows the diffusion of $Li+$ ions.
nickel–cadmium A type of battery that consists of a water-based cell with a cadmium anode and a highly oxidized nickel cathode.
lead–acid battery A battery consisting of a plate or grid of spongy lead metal, a cathode containing powdered $PbO2,$ and an electrolyte that is usually an aqueous solution of $H2SO4.$
Corrosion A galvanic process by which metals deteriorate through oxidation—usually but not always to their oxides.
sacrificial electrodes An electrode containing a more reactive metal that is attached to a metal object to inhibit that object’s corrosion.
electrolysis An electrochemical process in which an external voltage is applied to an electrolytic cell to drive a nonspontaneous reaction.
overvoltage The voltage that must be applied in electrolysis in addition to the calculated (theoretical) value to overcome factors such as a high activation energy and the formation of bubbles on a surface.
electroplating A process in which a layer of a second metal is deposited on the metal electrode that acts as the cathode during electrolysis.
nuclide An atom with a particular number of nucleons.
strong nuclear force An extremely powerful but very short-range attractive force between nucleons that keeps the nucleus of an atom from flying apart (due to electrostatic repulsions between protons).
superheavy elements An element with an atomic number near the magic number of 126.
nuclear transmutation reaction A nuclear reaction in which a nucleus reacts with a subatomic particle or another nuleus to give a product nucleus that is more massive than the starting material.
alpha (α) particle A helium nucleus: $4He.$
beta (β) decay A nuclear decay reaction in which a neutron is converted to a proton and a high-energy electron that is ejected from the nucleus as a β particle.
positron emission A nuclear decay reaction in which a proton is transformed into a neutron, and a high-energy positron is emitted.
electron capture (EC) A nuclear decay reaction in which an electron in an inner shell reacts with a proton to produce a neutron.
Gamma (γ) emission A nuclear decay reaction that results when a nucleus in an excited state releases energy in the form of a high-energy photon (a γ ray) when it returns to the ground state.
spontaneous fission A nuclear decay reaction in which the nucleus breaks into two pieces with different atomic numbers and atomic masses.
radioactive decay series A series of sequential alpha- and beta-decay reactions that occur until a stable nucleus is finally obtained.
transuranium elements An artificial element that has been prepared by bombarding suitable target nuclei with smaller particles.
Nonionizing radiation Radiation that is relatively low in energy. When it collides with an atom in a molecule or ion, most or all of its energy can be absorbed without causing a structural or a chemical change.
ionizing radiation Radiation of a high enough energy to transfer some as it passes through matter to one or more atoms with which it collides. If enough energy is transferred, electrons can be excited to very high energy levels, resulting in the formation of positively charged ions.
rad (radiation absorbed dose) A unit used to measure the effects of radiation on biological tissues; the amount of radiation that causes 0.01 J of energy to be absorbed by 1 kg of matter.
rem (roentgen equivalent in man) A unit that describes the actual amount of tissue damage caused by a given amount of radiation and equal to the number of rads multiplied by the RBE.
mass defect The difference between the sum of the masses of the components of an atom (protons, neutrons, and electrons) and the measured atomic mass.
nuclear binding energy The amount of energy released when a nucleus forms from its component nucleons, which corresponds to the mass defect of the nucleus.
nuclear fission The splitting of a heavy nucleus into two lighter ones.
critical mass The minimum mass of a fissile isotope capable of supporting sustained fission.
Nuclear fusion The combining of two light nuclei to produce a heavier, more stable nucleus.
thermonuclear reactions A nuclear reaction that requires a great deal of thermal energy to initiate the reaction.
nuclide An atom with a particular number of nucleons.
strong nuclear force An extremely powerful but very short-range attractive force between nucleons that keeps the nucleus of an atom from flying apart (due to electrostatic repulsions between protons).
superheavy elements An element with an atomic number near the magic number of 126.
nuclear transmutation reaction A nuclear reaction in which a nucleus reacts with a subatomic particle or another nuleus to give a product nucleus that is more massive than the starting material.
alpha (α) particle A helium nucleus: $4He.$
beta (β) decay A nuclear decay reaction in which a neutron is converted to a proton and a high-energy electron that is ejected from the nucleus as a β particle.
positron emission A nuclear decay reaction in which a proton is transformed into a neutron, and a high-energy positron is emitted.
electron capture (EC) A nuclear decay reaction in which an electron in an inner shell reacts with a proton to produce a neutron.
Gamma (γ) emission A nuclear decay reaction that results when a nucleus in an excited state releases energy in the form of a high-energy photon (a γ ray) when it returns to the ground state.
spontaneous fission A nuclear decay reaction in which the nucleus breaks into two pieces with different atomic numbers and atomic masses.
radioactive decay series A series of sequential alpha- and beta-decay reactions that occur until a stable nucleus is finally obtained.
transuranium elements An artificial element that has been prepared by bombarding suitable target nuclei with smaller particles.
Nonionizing radiation Radiation that is relatively low in energy. When it collides with an atom in a molecule or ion, most or all of its energy can be absorbed without causing a structural or a chemical change.
ionizing radiation Radiation of a high enough energy to transfer some as it passes through matter to one or more atoms with which it collides. If enough energy is transferred, electrons can be excited to very high energy levels, resulting in the formation of positively charged ions.
rad (radiation absorbed dose) A unit used to measure the effects of radiation on biological tissues; the amount of radiation that causes 0.01 J of energy to be absorbed by 1 kg of matter.
rem (roentgen equivalent in man) A unit that describes the actual amount of tissue damage caused by a given amount of radiation and equal to the number of rads multiplied by the RBE.
mass defect The difference between the sum of the masses of the components of an atom (protons, neutrons, and electrons) and the measured atomic mass.
nuclear binding energy The amount of energy released when a nucleus forms from its component nucleons, which corresponds to the mass defect of the nucleus.
nuclear fission The splitting of a heavy nucleus into two lighter ones.
critical mass The minimum mass of a fissile isotope capable of supporting sustained fission.
Nuclear fusion The combining of two light nuclei to produce a heavier, more stable nucleus.
thermonuclear reactions A nuclear reaction that requires a great deal of thermal energy to initiate the reaction.
effective nuclear charge The nuclear charge an electron actually experiences because of shielding from other electrons closer to the nucleus.
inert-pair effect The empirical observation that the heavier elements of groups 13–17 often have oxidation states that are lower by 2 than the maximum predicted for their group.
tritium A rare isotope of hydrogen that consists of one proton, two neutrons, and one electron.
three-center bond (or electron-deficient bond) A bond in which a hydride ion bridges two electropositive atoms.
graphite intercalation compounds A compound that forms when heavier alkali metals react with carbon in the form of graphite and insert themselves between the sheets of carbon atoms.
crown ethers A cyclic polyether that has four or more oxygen atoms separated by two or three carbon atoms. A central cavity can accommodate a metal ion coordinated to the ring of oxygen atoms.
Cryptands Consisting of three $(−OCH2CH2O−)n$ chains connected by two nitrogen atoms, this compound can completely encapsulate a metal ion of the appropriate size, coordinating to the metal by the lone pairs of electrons on each oxygen and the two nitrogen atoms.
organometallic compounds A compound that contains a metal covalently bonded to a carbon atom of an organic species.
ion pumps A complex assembly of proteins that selectively transport ions across cell membranes by their high affinity for ions of a certain charge and radius.
ionophores A molecule that facilitates the transport of metal ions across membranes.
effective nuclear charge The nuclear charge an electron actually experiences because of shielding from other electrons closer to the nucleus.
inert-pair effect The empirical observation that the heavier elements of groups 13–17 often have oxidation states that are lower by 2 than the maximum predicted for their group.
tritium A rare isotope of hydrogen that consists of one proton, two neutrons, and one electron.
three-center bond (or electron-deficient bond) A bond in which a hydride ion bridges two electropositive atoms.
graphite intercalation compounds A compound that forms when heavier alkali metals react with carbon in the form of graphite and insert themselves between the sheets of carbon atoms.
crown ethers A cyclic polyether that has four or more oxygen atoms separated by two or three carbon atoms. A central cavity can accommodate a metal ion coordinated to the ring of oxygen atoms.
Cryptands Consisting of three $(−OCH2CH2O−)n$ chains connected by two nitrogen atoms, this compound can completely encapsulate a metal ion of the appropriate size, coordinating to the metal by the lone pairs of electrons on each oxygen and the two nitrogen atoms.
organometallic compounds A compound that contains a metal covalently bonded to a carbon atom of an organic species.
ion pumps A complex assembly of proteins that selectively transport ions across cell membranes by their high affinity for ions of a certain charge and radius.
ionophores A molecule that facilitates the transport of metal ions across membranes.
Metallurgy A set of processes by which metals are extracted from their ores and converted to more useful forms.
coordination compound A chemical compound with one or more metal complexes.
soft acids An acid with the highest affinity for soft bases. It tends to be a cation of a less electropositive metal.
isomers Two or more compounds with the same molecular formula but different arrangements of their atoms.
structural isomers Two or more compounds that have the same molecular formula but differ in which atoms are bonded to one another.
geometrical isomers Complexes that differ only in which ligands are adjacent to one another or directly across from one another in the coordination sphere of a metal.
mer An isomer in which three ligands lie on a spherical meridian.
crystal field theory (CFT) A bonding model based on the assumption that metal–ligand interactions are purely electrostatic in nature, which explains many important properties of transition-metal complexes.
crystal field splitting energy The difference in energy between the $eg$ set of $d$ orbitals $(dz2$ and $dx2−y2)$ and the $t2g$ set of $d$ orbitals $(dxy$, $dxz$, $dyz)$ that results when the five $d$ orbitals are placed in an octahedral crystal field.
spin-pairing energy (P) The energy that must be overcome to place an electron in an orbital that already has one electron.
spectrochemical series An ordering of ligands by their crystal field splitting energies.
crystal field stabilization energy (CFSE) The additional stabilization of a metal complex by selective population of the lower-energy $d$ orbitals (the t2g orbitals).
Jahn–Teller theorem A theory that states that a non-linear molecule with a spatially degenerate electronic ground state will undergo a geometrical distortion to remove the degeneracy and lower the overall energy of the system.
siderophores An organic ligand that has a high affinity for Fe(III) and is secreted into the surrounding medium to increase the total concentration of dissolved iron.
metalloenzyme A protein that contains one or more tightly bound metal ions and catalyzes a biochemical reaction.
group transfer reactions Reactions involving the transfer of a group, catalyzed by metal ions that do not change their oxidation states during the reaction.
Metallurgy A set of processes by which metals are extracted from their ores and converted to more useful forms.
coordination compound A chemical compound with one or more metal complexes.
soft acids An acid with the highest affinity for soft bases. It tends to be a cation of a less electropositive metal.
isomers Two or more compounds with the same molecular formula but different arrangements of their atoms.
structural isomers Two or more compounds that have the same molecular formula but differ in which atoms are bonded to one another.
geometrical isomers Complexes that differ only in which ligands are adjacent to one another or directly across from one another in the coordination sphere of a metal.
mer An isomer in which three ligands lie on a spherical meridian.
crystal field theory (CFT) A bonding model based on the assumption that metal–ligand interactions are purely electrostatic in nature, which explains many important properties of transition-metal complexes.
crystal field splitting energy The difference in energy between the $eg$ set of $d$ orbitals $(dz2$ and $dx2−y2)$ and the $t2g$ set of $d$ orbitals $(dxy$, $dxz$, $dyz)$ that results when the five $d$ orbitals are placed in an octahedral crystal field.
spin-pairing energy (P) The energy that must be overcome to place an electron in an orbital that already has one electron.
spectrochemical series An ordering of ligands by their crystal field splitting energies.
crystal field stabilization energy (CFSE) The additional stabilization of a metal complex by selective population of the lower-energy $d$ orbitals (the t2g orbitals).
Jahn–Teller theorem A theory that states that a non-linear molecule with a spatially degenerate electronic ground state will undergo a geometrical distortion to remove the degeneracy and lower the overall energy of the system.
siderophores An organic ligand that has a high affinity for Fe(III) and is secreted into the surrounding medium to increase the total concentration of dissolved iron.
metalloenzyme A protein that contains one or more tightly bound metal ions and catalyzes a biochemical reaction.
group transfer reactions Reactions involving the transfer of a group, catalyzed by metal ions that do not change their oxidation states during the reaction.
functional groups The structural units that chemists use to classify organic compounds and predict their reactivities under a given set of conditions.
conformational isomer (or conformer) Isomers whose three-dimensional structures differ because of rotation about a σ bond.
structural isomers Isomers that have the same molecular formula but differ in which atoms are bonded to one another.
stereoisomers Molecules that have the same connectivity but whose component atoms have different orientations in space.
specific rotation The amount (in degrees) by which the plane of polarized light is rotated when the light is passed through a solution that contains 1.0 g of a solute per 1.0 mL of solvent in a tube 10.0 cm long.
electrophile An electron-deficient species that needs electrons to complete its octet.
radical Highly reactive species that have an unpaired valence electron.
nucleophile An electron-rich species that has a pair of electrons available to be shared with another atom.
substitution reaction A chemical reaction in which one atom or a group of atoms in a substance is replaced by another atom or a group of atoms from another substance.
elimination reactions A chemical reaction in which adjacent atoms are removed, or “eliminated,” from a molecule, resulting in the formation of a multiple bond and a small molecule.
addition reaction A chemical reaction in which the components of a species A–B are added to adjacent atoms across a carbon-carbon multiple bond.
pyrolysis reaction A high-temperature decomposition reaction that can be used to form fibers of synthetic polymers.
dehydration reaction A reaction that proceeds by eliminating the elements of water
electrophilic aromatic substitution reactions A reaction in which a −H of an arene is replaced (substituted) by an electrophilic group in a two-step process.
Grignard reagents An organometallic compound that has stabilized carbanions, whose general formula is RMgX, where X is Cl, Br, or I.
quaternary ammonium salts A salt that consist of an anion and a cation in which all four H atoms of the ammonium ion $(NH4+)$ are replaced by alkyl groups.
proteins A biological polymer with more than 50 amino acid residues linked together by amide bonds.
carbohydrates A polyhydroxy aldehyde or a polyhydroxy ketone with the general formula $Cn(H2O)m.$
Lipids A family of compounds that includes fats, waxes, some vitamins, and steroids and characterized by their insolubility in water.
Nucleic acids A linear polymer of nucleotides that is the basic structural component of DNA and RNA.
functional groups The structural units that chemists use to classify organic compounds and predict their reactivities under a given set of conditions.
conformational isomer (or conformer) Isomers whose three-dimensional structures differ because of rotation about a σ bond.
structural isomers Isomers that have the same molecular formula but differ in which atoms are bonded to one another.
stereoisomers Molecules that have the same connectivity but whose component atoms have different orientations in space.
specific rotation The amount (in degrees) by which the plane of polarized light is rotated when the light is passed through a solution that contains 1.0 g of a solute per 1.0 mL of solvent in a tube 10.0 cm long.
electrophile An electron-deficient species that needs electrons to complete its octet.
radical Highly reactive species that have an unpaired valence electron.
nucleophile An electron-rich species that has a pair of electrons available to be shared with another atom.
substitution reaction A chemical reaction in which one atom or a group of atoms in a substance is replaced by another atom or a group of atoms from another substance.
elimination reactions A chemical reaction in which adjacent atoms are removed, or “eliminated,” from a molecule, resulting in the formation of a multiple bond and a small molecule.
addition reaction A chemical reaction in which the components of a species A–B are added to adjacent atoms across a carbon-carbon multiple bond.
pyrolysis reaction A high-temperature decomposition reaction that can be used to form fibers of synthetic polymers.
dehydration reaction A reaction that proceeds by eliminating the elements of water
electrophilic aromatic substitution reactions A reaction in which a −H of an arene is replaced (substituted) by an electrophilic group in a two-step process.
Grignard reagents An organometallic compound that has stabilized carbanions, whose general formula is RMgX, where X is Cl, Br, or I.
quaternary ammonium salts A salt that consist of an anion and a cation in which all four H atoms of the ammonium ion $(NH4+)$ are replaced by alkyl groups.
proteins A biological polymer with more than 50 amino acid residues linked together by amide bonds.
carbohydrates A polyhydroxy aldehyde or a polyhydroxy ketone with the general formula $Cn(H2O)m.$
Lipids A family of compounds that includes fats, waxes, some vitamins, and steroids and characterized by their insolubility in water.
Nucleic acids A linear polymer of nucleotides that is the basic structural component of DNA and RNA.