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Chemistry LibreTexts

Chapter 20 Introduction to Organic Chemistry

  • Page ID
    209350
    • Section 12.1: Introduction
      One of the most striking characteristics of transition-metal complexes is the wide range of colors they exhibit.
    • Section 12.3: Coordination Compounds
      The transition elements and main group elements can form coordination compounds, or complexes, in which a central metal atom or ion is bonded to one or more ligands by coordinate covalent bonds. Ligands with more than one donor atom are called polydentate ligands and form chelates. The common geometries found in complexes are tetrahedral and square planar (both with a coordination number of four) and octahedral (with a coordination number of six).
    • 13.1: Introduction
    • 13.2: Rate of a Chemical Reaction
      Reaction rates are reported as either the average rate over a period of time or as the instantaneous rate at a single time. Reaction rates can be determined over particular time intervals or at a given point in time.
    • 13.3: The Rate Law- The Effect of Concentration on Reaction Rate
      The rate law for a reaction is a mathematical relationship between the reaction rate and the concentrations of species in solution. Rate laws can be expressed either as a differential rate law, describing the change in reactant or product concentrations as a function of time, or as an integrated rate law, describing the actual concentrations of reactants or products as a function of time. The rate constant and reaction order are extracted directly from the rate law.
    • 13.4: The Integrated Rate Law- The Dependence of Concentration on Time
      The reaction rate of a zeroth-order reaction is independent of the concentration of the reactants. The reaction rate of a first-order reaction is directly proportional to the concentration of one reactant. The reaction rate of a simple second-order reaction is proportional to the square of the concentration of one reactant. Knowing the rate law of a reaction gives clues to the reaction mechanism.
    • 13.5: The Effect of Temperature on Reaction Rate
      A minimum energy (activation energy,Ea) is required for a collision between molecules to result in a chemical reaction. Plots of potential energy for a system versus the reaction coordinate show an energy barrier that must be overcome for the reaction to occur. The arrangement of atoms at the highest point of this barrier is the activated complex, or transition state, of the reaction. At a given temperature, the higher the Ea, the slower the reaction.
    • 13.6: Reaction Mechanisms
      A balanced chemical reaction does not necessarily reveal either the individual elementary reactions by which a reaction occurs or its rate law. A reaction mechanism is the microscopic path by which reactants are transformed into products. Each step is an elementary reaction. Species that are formed in one step and consumed in another are intermediates. Each elementary reaction can be described in terms of its molecularity. The slowest step in a reaction mechanism is the rate-determining step.
    • 13.7: Catalysis
      Catalysts participate in a chemical reaction and increase its rate. They do not appear in the reaction’s net equation and are not consumed during the reaction. Catalysts allow a reaction to proceed via a pathway that has a lower activation energy than the uncatalyzed reaction. In heterogeneous catalysis, catalysts provide a surface to which reactants bind in a process of adsorption. In homogeneous catalysis, catalysts are in the same phase as the reactants. Enzymes are biological catalysts.
    • Section 12.4: Structure and Isomerization
      Two compounds that have the same formula and the same connectivity do not always have the same shape. There are two reasons why this may happen. In one case, the molecule may be flexible, so that it can twist into different shapes via rotation around individual sigma bonds. This phenomenon is called conformation, and it is covered in a different chapter. The second case occurs when two molecules appear to be connected the same way on paper, but are connected in two different ways in three dimens
    • 14.1: The Concept of Dynamic Equilibrium
      At equilibrium, the forward and reverse reactions of a system proceed at equal rates. Chemical equilibrium is a dynamic process consisting of forward and reverse reactions that proceed at equal rates. At equilibrium, the composition of the system no longer changes with time. The composition of an equilibrium mixture is independent of the direction from which equilibrium is approached.
    • 14.2: The Equilibrium Constant (K)
      The law of mass action describes a system at equilibrium in terms of the concentrations of the products and the reactants. For a system involving one or more gases, either the molar concentrations of the gases or their partial pressures can be used.
    • 14.3: Expressing the Equilibrium Constant in Terms of Pressure
      An equilibrated system that contains products and reactants in a single phase is a homogeneous equilibrium; a system whose reactants, products, or both are in more than one phase is a heterogeneous equilibrium.
    • 14.4: Calculating the Equilibrium Constant From Measured Equilibrium Concentrations
      Various methods can be used to solve the two fundamental types of equilibrium problems: (1) those in which we calculate the concentrations of reactants and products at equilibrium and (2) those in which we use the equilibrium constant and the initial concentrations of reactants to determine the composition of the equilibrium mixture. When an equilibrium constant is calculated from equilibrium concentrations, concentrations or partial pressures are use into the equilibrium constant expression.
    • 14.5: Heterogenous Equilibria - Reactions Involving Solids and Liquids
      When the products and reactants of an equilibrium reaction form a single phase, whether gas or liquid, the system is a homogeneous equilibrium. In such situations, the concentrations of the reactants and products can vary over a wide range. In contrast, a system whose reactants, products, or both are in more than one phase is a heterogeneous equilibrium, such as the reaction of a gas with a solid or liquid.
    • 14.6: The Reaction Quotient- Predicting the Direction of Change
      The reaction Quotient has the same form as the equilibrium constant expression, but it is derived from concentrations obtained at any time. When a reaction system is at equilibrium, Q=K . Graphs derived by plotting a few equilibrium concentrations for a system at a given temperature and pressure can be used to predict the direction in which a reaction will proceed. Points that do not lie on the line or curve are nonequilibrium states.
    • 14.7: Finding Equilibrium Concentrations
      Various methods can be used to solve the two fundamental types of equilibrium problems: (1) those in which we calculate the concentrations of reactants and products at equilibrium and (2) those in which we use the equilibrium constant and the initial concentrations of reactants to determine the composition of the equilibrium mixture. When an equilibrium constant is calculated from equilibrium concentrations, concentrations or partial pressures are use into the equilibrium constant expression.
    • 14.8: Le Châtelier’s Principle- How a System at Equilibrium Responds to Disturbances
      Systems at equilibrium can be disturbed by changes to temperature, concentration, and, in some cases, volume and pressure; volume and pressure changes will disturb equilibrium if the number of moles of gas is different on the reactant and product sides of the reaction. The system's response to these disturbances is described by Le Châtelier's principle: The system will respond in a way that counteracts the disturbance. Adding a catalyst affects the reaction rates but does not alter equilibrium.
    • Section 12.5: Bonding in Coordinate Compounds
      Crystal field theory treats interactions between the electrons on the metal and the ligands as a simple electrostatic effect. The presence of the ligands near the metal ion changes the energies of the metal d orbitals relative to their energies in the free ion. Both the color and the magnetic properties of a complex can be attributed to this crystal field splitting. The magnitude of the splitting depends on the nature of the ligands bonded to the metal.
    • 15.10: Acid Strength and Molecular Structure
      Inductive effects and charge delocalization significantly influence the acidity or basicity of a compound. The acid–base strength of a molecule depends strongly on its structure. The weaker the A–H or B–H+ bond, the more likely it is to dissociate to form an H+H+ ion. In addition, any factor that stabilizes the lone pair on the conjugate base favors the dissociation of H+H+ , making the conjugate acid a stronger acid.
    • 15.11: Lewis Acids and Bases
      Lewis proposed that the electron pair is the dominant actor in acid-base chemistry. An Lewis acid is a substance that accepts a pair of electrons, and in doing so, forms a covalent bond with the entity that supplies the electrons. A Lewis base is a substance that donates an unshared pair of electrons to a recipient species with which the electrons can be shared. Lewis acis/base theory is a powerful tool for describing many chemical reactions used in organic and inorganic chemistry.
    • 15.12: Acid rain
      The damaging effects of acid rain have led to strong pressure on industry to minimize the release of harmful reactants. Acid rain is rainfall whose pH is less than 5.6, the value typically observed, due to the presence of dissolved carbon dioxide. Acid rain is caused by nitrogen oxides and sulfur dioxide produced by both natural processes and the combustion of fossil fuels. Eventually, these oxides react with oxygen and water to give nitric acid and sulfuric acid.
    • 15.1: Introduction
      Heartburn is caused by a buildup of excessive amounts of stomach acid, particularly HCl. This acid is used to digest the food we eat, but it can often back up into the esophagus causing that burning sensation many of us are familiar with.
    • 15.2: The Nature of Acids and Bases
      In chemistry, acids and bases have been defined differently by three sets of theories: One is the Arrhenius definition defined above, which revolves around the idea that acids are substances that ionize (break off) in an aqueous solution to produce hydrogen (H+) ions while bases produce hydroxide (OH-) ions in solution. The other two definitions are discussed in detail alter in the chapter and include the Brønsted-Lowry definition and the Lewis theory.
    • 15.3: Definitions of Acids and Bases
      A compound that can donate a proton (a hydrogen ion) to another compound is called a Brønsted-Lowry acid. The compound that accepts the proton is called a Brønsted-Lowry base. The species remaining after a Brønsted-Lowry acid has lost a proton is the conjugate base of the acid. The species formed when a Brønsted-Lowry base gains a proton is the conjugate acid of the base. Thus, an acid-base reaction occurs when a proton is transferred from an acid to a base.
    • 15.4: Acid Strength and the Acid Dissociation Constant (Ka)
      Acid–base reactions always contain two conjugate acid–base pairs. Each acid and each base has an associated ionization constant that corresponds to its acid or base strength. Two species that differ by only a proton constitute a conjugate acid–base pair. The magnitude of the equilibrium constant for an ionization reaction can be used to determine the relative strengths of acids and bases.
    • 15.5: Autoionization of Water and pH
      Water is amphiprotic: it can act as an acid by donating a proton to a base to form the hydroxide ion, or as a base by accepting a proton from an acid to form the hydronium ion ( H3O+ ). The autoionization of liquid water produces OH− and H3O+ ions. The equilibrium constant for this reaction is called the ion-product constant of liquid water (Kw) and is defined as Kw=[H3O+][OH−] . At 25°C, Kw is 1.01×10−14 ; hence pH+pOH=pKw=14.00 .
    • 15.6: Finding the [H3O+] and pH of Strong and Weak Acid Solutions
      Acid–base reactions always contain two conjugate acid–base pairs. Each acid and each base has an associated ionization constant that corresponds to its acid or base strength. Two species that differ by only a proton constitute a conjugate acid–base pair. The magnitude of the equilibrium constant for an ionization reaction can be used to determine the relative strengths of acids and bases.
    • 15.7: Base Solutions
    • 15.8: The Acid-Base Properties of Ions and Salts
      A salt can dissolve in water to produce a neutral, a basic, or an acidic solution, depending on whether it contains the conjugate base of a weak acid as the anion ( A−A− ), the conjugate acid of a weak base as the cation ( BH+ ), or both. Salts that contain small, highly charged metal ions produce acidic solutions in water. The reaction of a salt with water to produce an acidic or a basic solution is called a hydrolysis reaction.
    • 15.9: Polyprotic Acids
      An acid that contains more than one ionizable proton is a polyprotic acid. The protons of these acids ionize in steps. The differences in the acid ionization constants for the successive ionizations of the protons in a polyprotic acid usually vary by roughly five orders of magnitude. As long as the difference between the successive values of Ka of the acid is greater than about a factor of 20, it is appropriate to break down the calculations of the concentrations sequentially.
    • Section 12.56: Applications of Coordination Compounds
      In this section, we describe several systems that illustrate the roles transition metals play in biological systems. Our goal is for you to understand why the chemical properties of these elements make them essential for life. We begin with a discussion of the strategies organisms use to extract transition metals from their environment.
    • Section 12.7: Transition Metals and Coordination Compounds (Exercises)
    • Section 12.2: Properties of Transition Metals
      The transition metals are elements with partially filled d orbitals, located in the d-block of the periodic table. The reactivity of the transition elements varies widely from very active metals such as scandium and iron to almost inert elements, such as the platinum metals. The type of chemistry used in the isolation of the elements from their ores depends upon the concentration of the element in its ore and the difficulty of reducing ions of the elements to the metals.