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2.2.6: Molecular and Ionic Compounds

  • Page ID
    210634
  • Learning Objectives

    • Define ionic and molecular (covalent) compounds
    • Predict the type of compound formed from elements based on their location within the periodic table
    • Determine formulas for simple ionic compounds
    • Derive names for common types of inorganic compounds using a systematic approach

    Atomic Bonding

    Video \(\PageIndex{1}\): A brief overview of the bonds we will discuss below.

    In ordinary chemical reactions, the nucleus of each atom (and thus the identity of the element) remains unchanged. Electrons, however, can be added to atoms by transfer from other atoms, lost by transfer to other atoms, or shared with other atoms. The transfer and sharing of electrons among atoms govern the chemistry of the elements. During the formation of some compounds, atoms gain or lose electrons, and form electrically charged particles called ions (Figure \(\PageIndex{1}\); Video \(\PageIndex{2}\)).

    Video \(\PageIndex{2}\): What are ions?

    Figure A shows a sodium atom, N a, which has a nucleus containing 11 protons and 12 neutrons. The atom’s surrounding electron cloud contains 11 electrons. Figure B shows a sodium ion, N a superscript plus sign. Its nucleus contains 11 protons and 12 neutrons. The ion’s electron cloud contains 10 electrons and is smaller than that of the sodium atom in figure A.

    Figure \(\PageIndex{1}\): (a) A sodium atom (Na) has equal numbers of protons and electrons (11) and is uncharged. (b) A sodium cation (Na+) has lost an electron, so it has one more proton (11) than electrons (10), giving it an overall positive charge, signified by a superscripted plus sign.

    You can use the periodic table to predict whether an atom will form an anion or a cation, and you can often predict the charge of the resulting ion. Atoms of many main-group metals lose enough electrons to leave them with the same number of electrons as an atom of the preceding noble gas. To illustrate, an atom of an alkali metal (group 1) loses one electron and forms a cation with a 1+ charge; an alkaline earth metal (group 2) loses two electrons and forms a cation with a 2+ charge, and so on. For example, a neutral calcium atom, with 20 protons and 20 electrons, readily loses two electrons. This results in a cation with 20 protons, 18 electrons, and a 2+ charge. It has the same number of electrons as atoms of the preceding noble gas, argon, and is symbolized Ca2+. The name of a metal ion is the same as the name of the metal atom from which it forms, so Ca2+ is called a calcium ion.

    When atoms of nonmetal elements form ions, they generally gain enough electrons to give them the same number of electrons as an atom of the next noble gas in the periodic table. Atoms of group 17 gain one electron and form anions with a 1− charge; atoms of group 16 gain two electrons and form ions with a 2− charge, and so on. For example, the neutral bromine atom, with 35 protons and 35 electrons, can gain one electron to provide it with 36 electrons. This results in an anion with 35 protons, 36 electrons, and a 1− charge. It has the same number of electrons as atoms of the next noble gas, krypton, and is symbolized Br. (A discussion of the theory supporting the favored status of noble gas electron numbers reflected in these predictive rules for ion formation is provided in a later chapter of this text.)

    Note the usefulness of the periodic table in predicting likely ion formation and charge (Figure \(\PageIndex{2}\)). Moving from the far left to the right on the periodic table, main-group elements tend to form cations with a charge equal to the group number. That is, group 1 elements form 1+ ions; group 2 elements form 2+ ions, and so on. Moving from the far right to the left on the periodic table, elements often form anions with a negative charge equal to the number of groups moved left from the noble gases. For example, group 17 elements (one group left of the noble gases) form 1− ions; group 16 elements (two groups left) form 2− ions, and so on. This trend can be used as a guide in many cases, but its predictive value decreases when moving toward the center of the periodic table. In fact, transition metals and some other metals often exhibit variable charges that are not predictable by their location in the table. For example, copper can form ions with a 1+ or 2+ charge, and iron can form ions with a 2+ or 3+ charge.

    Group one of the periodic table contains L i superscript plus sign in period 2, N a superscript plus sign in period 3, K superscript plus sign in period 4, R b superscript plus sign in period 5, C s superscript plus sign in period 6, and F r superscript plus sign in period 7. Group two contains B e superscript 2 plus sign in period 2, M g superscript 2 plus sign in period 3, C a superscript 2 plus sign in period 4, S r superscript 2 plus sign in period 5, B a superscript 2 plus sign in period 6, and R a superscript 2 plus sign in period 7. Group six contains C r superscript 3 plus sign and C r superscript 6 plus sign in period 4. Group seven contains M n superscript 2 plus sign in period 4. Group eight contains F e superscript 2 plus sign and F e superscript 3 plus sign in period 4. Group nine contains C o superscript 2 plus sign in period 4. Group ten contains N i superscript 2 plus sign in period 4, and P t superscript 2 plus sign in period 6. Group 11 contains C U superscript plus sign and C U superscript 2 plus sign in period 4, A g superscript plus sign in period 5, and A u superscript plus sign and A u superscript 3 plus sign in period 6. Group 12 contains Z n superscript 2 plus sign in period 4, C d superscript 2 plus sign in period 5, and H g subscript 2 superscript 2 plus sign and H g superscript 2 plus sign in period 6. Group 13 contains A l superscript 3 plus sign in period 3. Group 14 contains C superscript 4 negative sign in period 2. Group 15 contains N superscript 3 negative sign in period 2, P superscript 3 negative sign in period 3, and A s superscript 3 negative sign in period 4. Group 16 contains O superscript 2 negative sign in period 2, S superscript 2 negative sign in period 3, S e superscript 2 negative sign in period 4 and T e superscript 2 negative sign in period 5. Group 17 contains F superscript negative sign in period 2, C l superscript negative sign in period 3, B r superscript negative sign in period 4, I superscript negative sign in period 5, and A t superscript negative sign in period 6. Group 18 contains H e in period 1, N e in period 2, A r in period 3, K r in period 4, X e in period 5 and R n in period 6.

    Figure \(\PageIndex{2}\): Some elements exhibit a regular pattern of ionic charge when they form ions.

    Example \(\PageIndex{1}\): Composition of Ions

    An ion found in some compounds used as antiperspirants contains 13 protons and 10 electrons. What is its symbol?

    Solution

    Because the number of protons remains unchanged when an atom forms an ion, the atomic number of the element must be 13. Knowing this lets us use the periodic table to identify the element as Al (aluminum). The Al atom has lost three electrons and thus has three more positive charges (13) than it has electrons (10). This is the aluminum cation, Al3+.

    Exercise \(\PageIndex{1}\)

    Give the symbol and name for the ion with 34 protons and 36 electrons.

    Answer

    Se2−, the selenide ion

    Example \(\PageIndex{2}\): Formation of Ions

    Magnesium and nitrogen react to form an ionic compound. Predict which forms an anion, which forms a cation, and the charges of each ion. Write the symbol for each ion and name them.

    Solution

    Magnesium’s position in the periodic table (group 2) tells us that it is a metal. Metals form positive ions (cations). A magnesium atom must lose two electrons to have the same number electrons as an atom of the previous noble gas, neon. Thus, a magnesium atom will form a cation with two fewer electrons than protons and a charge of 2+. The symbol for the ion is Mg2+, and it is called a magnesium ion.

    Nitrogen’s position in the periodic table (group 15) reveals that it is a nonmetal. Nonmetals form negative ions (anions). A nitrogen atom must gain three electrons to have the same number of electrons as an atom of the following noble gas, neon. Thus, a nitrogen atom will form an anion with three more electrons than protons and a charge of 3−. The symbol for the ion is N3−, and it is called a nitride ion.

    Exercise \(\PageIndex{2}\)

    Aluminum and carbon react to form an ionic compound. Predict which forms an anion, which forms a cation, and the charges of each ion. Write the symbol for each ion and name them.

    Answer

    Al will form a cation with a charge of 3+: Al3+, an aluminum ion. Carbon will form an anion with a charge of 4−: C4−, a carbide ion.

    The ions that we have discussed so far are called monatomic ions, that is, they are ions formed from only one atom. We also find many polyatomic ions. These ions, which act as discrete units, are electrically charged molecules (a group of bonded atoms with an overall charge). Some of the more important polyatomic ions are listed in Table \(\PageIndex{1}\). Oxyanions are polyatomic ions that contain one or more oxygen atoms. At this point in your study of chemistry, you should memorize the names, formulas, and charges of the most common polyatomic ions. Because you will use them repeatedly, they will soon become familiar.

    Table \(\PageIndex{1}\): Common Polyatomic Ions
    Name Formula Related Acid Formula
    ammonium \(\ce{NH4+}\)
    hydronium \(\ce{H_3O^+}\)
    oxide \(\ce{O^{2-}}\)
    peroxide \(\ce{O_2^{2-}}\)
    hydroxide \(\ce{OH^-}\)
    acetate \(\ce{CH_3COO^-}\) acetic acid \(\ce{CH_3COOH}\)
    cyanide \(\ce{CN^-}\) hydrocyanic acid \(\ce{HCN}\)
    azide \(\ce{N_3^-}\) hydrazoic acid \(\ce{HN_3}\)
    carbonate \(\ce{CO_3^{2-}}\) carbonic acid \(\ce{H_2CO_3}\)
    bicarbonate \(\ce{HCO_3^-}\)
    nitrate \(\ce{NO_3^-}\) nitric acid \(\ce{HNO_3}\)
    nitrite \(\ce{NO_2^-}\) nitrous acid \(\ce{HNO_2}\)
    sulfate \(\ce{SO_4^{2-}}\) sulfuric acid \(\ce{H_2SO_4}\)
    hydrogen sulfate \(\ce{HSO_4^-}\)
    sulfite \(\ce{SO_3^{2-}}\) sulfurous acid \(\ce{H_2SO_3}\)
    hydrogen sulfite \(\ce{HSO_3^-}\)
    phosphate \(\ce{PO_4^{3-}}\) phosphoric acid \(\ce{H_3PO_4}\)
    hydrogen phosphate \(\ce{HPO_4^{2-}}\)
    dihydrogen phosphate \(\ce{H_2PO_4^-}\)
    perchlorate \(\ce{ClO_4^-}\) perchloric acid \(\ce{HClO_4}\)
    chlorate \(\ce{ClO_3^-}\) chloric acid \(\ce{HClO_3}\)
    chlorite \(\ce{ClO_2^-}\) chlorous acid \(\ce{HClO_2}\)
    hypochlorite \(\ce{ClO^-}\) hypochlorous acid \(\ce{HClO}\)
    chromate \(\ce{CrO_4^{2-}}\) chromic acid \(\ce{H_2CrO_4}\)
    dichromate \(\ce{Cr_2O_7^{2-}}\) dichromic acid \(\ce{H_2Cr_2O7}\)
    permanganate \(\ce{MnO_4^-}\) permanganic acid \(\ce{HMnO_4}\)

    Note that there is a system for naming some polyatomic ions; -ate and -ite are suffixes designating polyatomic ions containing more or fewer oxygen atoms. Per- (short for “hyper”) and hypo- (meaning “under”) are prefixes meaning more oxygen atoms than -ate and fewer oxygen atoms than -ite, respectively. For example, perchlorate is \(\ce{ClO4-}\), chlorate is \(\ce{ClO3-}\), chlorite is \(\ce{ClO2-}\) and hypochlorite is ClO. Unfortunately, the number of oxygen atoms corresponding to a given suffix or prefix is not consistent; for example, nitrate is \(\ce{NO3-}\) while sulfate is \(\ce{SO4^{2-}}\). This will be covered in more detail in the next module on nomenclature.

    The nature of the attractive forces that hold atoms or ions together within a compound is the basis for classifying chemical bonding. When electrons are transferred and ions form, ionic bonds result. Ionic bonds are electrostatic forces of attraction, that is, the attractive forces experienced between objects of opposite electrical charge (in this case, cations and anions). When electrons are “shared” and molecules form, covalent bonds result. Covalent bonds are the attractive forces between the positively charged nuclei of the bonded atoms and one or more pairs of electrons that are located between the atoms. Compounds are classified as ionic or molecular (covalent) on the basis of the bonds present in them.

    Ionic Compounds

    Video \(\PageIndex{3}\): What is an ionic bond?

    When an element composed of atoms that readily lose electrons (a metal) reacts with an element composed of atoms that readily gain electrons (a nonmetal), a transfer of electrons usually occurs, producing ions. The compound formed by this transfer is stabilized by the electrostatic attractions (ionic bonds) between the ions of opposite charge present in the compound. For example, when each sodium atom in a sample of sodium metal (group 1) gives up one electron to form a sodium cation, Na+, and each chlorine atom in a sample of chlorine gas (group 17) accepts one electron to form a chloride anion, Cl, the resulting compound, NaCl, is composed of sodium ions and chloride ions in the ratio of one Na+ ion for each Cl ion. Similarly, each calcium atom (group 2) can give up two electrons and transfer one to each of two chlorine atoms to form CaCl2, which is composed of Ca2+ and Cl ions in the ratio of one Ca2+ ion to two Cl ions.

    A compound that contains ions and is held together by ionic bonds is called an ionic compound. The periodic table can help us recognize many of the compounds that are ionic: When a metal is combined with one or more nonmetals, the compound is usually ionic. This guideline works well for predicting ionic compound formation for most of the compounds typically encountered in an introductory chemistry course. However, it is not always true (for example, aluminum chloride, AlCl3, is not ionic).

    You can often recognize ionic compounds because of their properties. Ionic compounds are solids that typically melt at high temperatures and boil at even higher temperatures. For example, sodium chloride melts at 801 °C and boils at 1413 °C. (As a comparison, the molecular compound water melts at 0 °C and boils at 100 °C.) In solid form, an ionic compound is not electrically conductive because its ions are unable to flow (“electricity” is the flow of charged particles). When molten, however, it can conduct electricity because its ions are able to move freely through the liquid (Figure \(\PageIndex{3}\); Video \(\PageIndex{4}\)).

    This figure shows three photos connected by right-facing arrows. The first shows a light bulb as part of a complex lab equipment setup. The light bulb is not lit. The second photo shows a substances being heated or set on fire. The third shows the light bulb again which is lit.

    Figure \(\PageIndex{3}\): Sodium chloride melts at 801 °C and conducts electricity when molten. (credit: modification of work by Mark Blaser and Matt Evans)

    Video \(\PageIndex{4}\): Watch this video to see a mixture of salts melt and conduct electricity.

    In every ionic compound, the total number of positive charges of the cations equals the total number of negative charges of the anions. Thus, ionic compounds are electrically neutral overall, even though they contain positive and negative ions. We can use this observation to help us write the formula of an ionic compound. The formula of an ionic compound must have a ratio of ions such that the numbers of positive and negative charges are equal.

    Example \(\PageIndex{3}\): Predicting the Formula of an Ionic Compound

    The gemstone sapphire (Figure \(\PageIndex{4}\)) is mostly a compound of aluminum and oxygen that contains aluminum cations, Al3+, and oxygen anions, O2−. What is the formula of this compound?

    This is a photograph of a ring with a sapphire set in it.

    Figure \(\PageIndex{4}\): Although pure aluminum oxide is colorless, trace amounts of iron and titanium give blue sapphire its characteristic color. (credit: modification of work by Stanislav Doronenko)

    Solution Because the ionic compound must be electrically neutral, it must have the same number of positive and negative charges. Two aluminum ions, each with a charge of 3+, would give us six positive charges, and three oxide ions, each with a charge of 2−, would give us six negative charges. The formula would be Al2O3.

    Exercise \(\PageIndex{3}\)

    Predict the formula of the ionic compound formed between the sodium cation, Na+, and the sulfide anion, S2−.

    Answer

    Na2S

    Many ionic compounds contain polyatomic ions (Table \(\PageIndex{1}\)) as the cation, the anion, or both. As with simple ionic compounds, these compounds must also be electrically neutral, so their formulas can be predicted by treating the polyatomic ions as discrete units. We use parentheses in a formula to indicate a group of atoms that behave as a unit. For example, the formula for calcium phosphate, one of the minerals in our bones, is Ca3(PO4)2. This formula indicates that there are three calcium ions (Ca2+) for every two phosphate \(\left(\ce{PO4^{3-}}\right)\) groups. The \(\ce{PO4^{3-}}\) groups are discrete units, each consisting of one phosphorus atom and four oxygen atoms, and having an overall charge of 3−. The compound is electrically neutral, and its formula shows a total count of three Ca, two P, and eight O atoms.

    Example \(\PageIndex{4}\): Predicting the Formula of a Compound with a Polyatomic Anion

    Baking powder contains calcium dihydrogen phosphate, an ionic compound composed of the ions Ca2+ and \(\ce{H2PO4-}\). What is the formula of this compound?

    Solution

    The positive and negative charges must balance, and this ionic compound must be electrically neutral. Thus, we must have two negative charges to balance the 2+ charge of the calcium ion. This requires a ratio of one Ca2+ ion to two \(\ce{H2PO4-}\) ions. We designate this by enclosing the formula for the dihydrogen phosphate ion in parentheses and adding a subscript 2. The formula is Ca(H2PO4)2.

    Exercise \(\PageIndex{4}\)

    Predict the formula of the ionic compound formed between the lithium ion and the peroxide ion, \(\ce{O2^2-}\) (Hint: Use the periodic table to predict the sign and the charge on the lithium ion.)

    Answer

    Li2O2

    Because an ionic compound is not made up of single, discrete molecules, it may not be properly symbolized using a molecular formula. Instead, ionic compounds must be symbolized by a formula indicating the relative numbers of its constituent ions. For compounds containing only monatomic ions (such as NaCl) and for many compounds containing polyatomic ions (such as CaSO4), these formulas are just the empirical formulas introduced earlier in this chapter. However, the formulas for some ionic compounds containing polyatomic ions are not empirical formulas. For example, the ionic compound sodium oxalate is comprised of Na+ and \(\ce{C2O4^2-}\) ions combined in a 2:1 ratio, and its formula is written as Na2C2O4. The subscripts in this formula are not the smallest-possible whole numbers, as each can be divided by 2 to yield the empirical formula, NaCO2. This is not the accepted formula for sodium oxalate, however, as it does not accurately represent the compound’s polyatomic anion, \(\ce{C2O4^2-}\).

    Molecular Compounds

    Video \(\PageIndex{5}\): What are covalent bonds?

    Many compounds do not contain ions but instead consist solely of discrete, neutral molecules. These molecular compounds (covalent compounds) result when atoms share, rather than transfer (gain or lose), electrons. Covalent bonding is an important and extensive concept in chemistry, and it will be treated in considerable detail in a later chapter of this text. We can often identify molecular compounds on the basis of their physical properties. Under normal conditions, molecular compounds often exist as gases, low-boiling liquids, and low-melting solids, although many important exceptions exist.

    Whereas ionic compounds are usually formed when a metal and a nonmetal combine, covalent compounds are usually formed by a combination of nonmetals. Thus, the periodic table can help us recognize many of the compounds that are covalent. While we can use the positions of a compound’s elements in the periodic table to predict whether it is ionic or covalent at this point in our study of chemistry, you should be aware that this is a very simplistic approach that does not account for a number of interesting exceptions. Shades of gray exist between ionic and molecular compounds, and you’ll learn more about those later.

    Example \(\PageIndex{5}\): Predicting the Type of Bonding in Compounds

    Predict whether the following compounds are ionic or molecular:

    1. KI, the compound used as a source of iodine in table salt
    2. H2O2, the bleach and disinfectant hydrogen peroxide
    3. CHCl3, the anesthetic chloroform
    4. Li2CO3, a source of lithium in antidepressants

    Solution

    1. Potassium (group 1) is a metal, and iodine (group 17) is a nonmetal; KI is predicted to be ionic.
    2. Hydrogen (group 1) is a nonmetal, and oxygen (group 16) is a nonmetal; H2O2 is predicted to be molecular.
    3. Carbon (group 14) is a nonmetal, hydrogen (group 1) is a nonmetal, and chlorine (group 17) is a nonmetal; CHCl3 is predicted to be molecular.
    4. Lithium (group 1) is a metal, and carbonate is a polyatomic ion; Li2CO3 is predicted to be ionic.

    Exercise \(\PageIndex{5}\)

    Using the periodic table, predict whether the following compounds are ionic or covalent:

    1. SO2
    2. CaF2
    3. N2H4
    4. Al2(SO4)3
    Answer a

    molecular

    Answer b

    ionic

    Answer c

    molecular

    Answer d

    ionic

    Naming Chemicals Compounds

    Nomenclature, a collection of rules for naming things, is important in science and in many other situations. This module describes an approach that is used to name simple ionic and molecular compounds, such as NaCl, CaCO3, and N2O4. The simplest of these are binary compounds, those containing only two elements, but we will also consider how to name ionic compounds containing polyatomic ions, and one specific, very important class of compounds known as acids (subsequent chapters in this text will focus on these compounds in great detail). We will limit our attention here to inorganic compounds, compounds that are composed principally of elements other than carbon, and will follow the nomenclature guidelines proposed by IUPAC. The rules for organic compounds, in which carbon is the principle element, will be treated in a later chapter on organic chemistry.

    Ionic Compounds

    To name an inorganic compound, we need to consider the answers to several questions. First, is the compound ionic or molecular? If the compound is ionic, does the metal form ions of only one type (fixed charge) or more than one type (variable charge)? Are the ions monatomic or polyatomic? If the compound is molecular, does it contain hydrogen? If so, does it also contain oxygen? From the answers we derive, we place the compound in an appropriate category and then name it accordingly. We will begin with the nomenclature rules for ionic compounds.

    Compounds Containing Only Monatomic Ions

    The name of a binary compound containing monatomic ions consists of the name of the cation (the name of the metal) followed by the name of the anion (the name of the nonmetallic element with its ending replaced by the suffix –ide). Some examples are given in Table \(\PageIndex{2}\).

    Table \(\PageIndex{2}\): Names of Some Ionic Compounds
    NaCl, sodium chloride Na2O, sodium oxide
    KBr, potassium bromide CdS, cadmium sulfide
    CaI2, calcium iodide Mg3N2, magnesium nitride
    CsF, cesium fluoride Ca3P2, calcium phosphide
    LiCl, lithium chloride Al4C3, aluminum carbide

    Compounds Containing Polyatomic Ions

    Compounds containing polyatomic ions are named similarly to those containing only monatomic ions, except there is no need to change to an –ide ending, since the suffix is already present in the name of the anion. Examples are shown in Table \(\PageIndex{3}\).

    Table \(\PageIndex{3}\): Names of Some Polyatomic Ionic Compounds
    KC2H3O2, potassium acetate (NH4)Cl, ammonium chloride
    NaHCO3, sodium bicarbonate CaSO4, calcium sulfate
    Al2(CO3)3, aluminum carbonate Mg3(PO4)2, magnesium phosphate

    Ionic Compounds in Your Cabinets

    Every day you encounter and use a large number of ionic compounds. Some of these compounds, where they are found, and what they are used for are listed in Table \(\PageIndex{4}\). Look at the label or ingredients list on the various products that you use during the next few days, and see if you run into any of those in this table, or find other ionic compounds that you could now name or write as a formula.

    Table \(\PageIndex{4}\): Everyday Ionic Compounds
    Ionic Compound Name Use
    NaCl sodium chloride ordinary table salt
    KI potassium iodide added to “iodized” salt for thyroid health
    NaF sodium fluoride ingredient in toothpaste
    NaHCO3 sodium bicarbonate baking soda; used in cooking (and in antacids)
    Na2CO3 sodium carbonate washing soda; used in cleaning agents
    NaOCl sodium hypochlorite active ingredient in household bleach
    CaCO3 calcium carbonate ingredient in antacids
    Mg(OH)2 magnesium hydroxide ingredient in antacids
    Al(OH)3 aluminum hydroxide ingredient in antacids
    NaOH sodium hydroxide lye; used as drain cleaner
    K3PO4 potassium phosphate food additive (many purposes)
    MgSO4 magnesium sulfate added to purified water
    Na2HPO4 sodium hydrogen phosphate anti-caking agent; used in powdered products
    Na2SO3 sodium sulfite preservative

    Compounds Containing a Metal Ion with a Variable Charge

    Most of the transition metals can form two or more cations with different charges. Compounds of these metals with nonmetals are named with the same method as compounds in the first category, except the charge of the metal ion is specified by a Roman numeral in parentheses after the name of the metal. The charge of the metal ion is determined from the formula of the compound and the charge of the anion. For example, consider binary ionic compounds of iron and chlorine. Iron typically exhibits a charge of either 2+ or 3+, and the two corresponding compound formulas are FeCl2 and FeCl3. The simplest name, “iron chloride,” will, in this case, be ambiguous, as it does not distinguish between these two compounds. In cases like this, the charge of the metal ion is included as a Roman numeral in parentheses immediately following the metal name. These two compounds are then unambiguously named iron(II) chloride and iron(III) chloride, respectively. Other examples are provided in Table \(\PageIndex{5}\).

    Table \(\PageIndex{5}\): Names of Some Transition Metal Ionic Compounds
    Transition Metal Ionic Compound Name
    FeCl3 iron(III) chloride
    Hg2O mercury(I) oxide
    HgO mercury(II) oxide
    Cu3(PO4)2 copper(II) phosphate

    Out-of-date nomenclature used the suffixes –ic and –ous to designate metals with higher and lower charges, respectively: Iron(III) chloride, FeCl3, was previously called ferric chloride, and iron(II) chloride, FeCl2, was known as ferrous chloride. Though this naming convention has been largely abandoned by the scientific community, it remains in use by some segments of industry. For example, you may see the words stannous fluoride on a tube of toothpaste. This represents the formula SnF2, which is more properly named tin(II) fluoride. The other fluoride of tin is SnF4, which was previously called stannic fluoride but is now named tin(IV) fluoride.

    Example \(\PageIndex{6}\): Naming Ionic Compounds

    Name the following ionic compounds, which contain a metal that can have more than one ionic charge:

    1. Fe2S3
    2. CuSe
    3. GaN
    4. CrCl3
    5. Ti2(SO4)3

    Solution

    The anions in these compounds have a fixed negative charge (S2−, Se2− , N3−, Cl, and \(\ce{SO4^2-}\)), and the compounds must be neutral. Because the total number of positive charges in each compound must equal the total number of negative charges, the positive ions must be Fe3+, Cu2+, Ga3+, Cr3+, and Ti3+. These charges are used in the names of the metal ions:

    1. iron(III) sulfide
    2. copper(II) selenide
    3. gallium(III) nitride
    4. chromium(III) chloride
    5. titanium(III) sulfate

    Exercise \(\PageIndex{6}\)

    Write the formulas of the following ionic compounds:

    1. chromium(III) phosphide
    2. mercury(II) sulfide
    3. manganese(II) phosphate
    4. copper(I) oxide
    5. chromium(VI) fluoride
    Answer a

    CrP

    Answer b

    HgS

    Answer c

    Mn3(PO4)2

    Answer d

    Cu2O

    Answer e

    CrF6

    Erin Brokovich and Chromium Contamination

    In the early 1990s, legal file clerk Erin Brockovich (Figure \(\PageIndex{2}\)) discovered a high rate of serious illnesses in the small town of Hinckley, California. Her investigation eventually linked the illnesses to groundwater contaminated by Cr(VI) used by Pacific Gas & Electric (PG&E) to fight corrosion in a nearby natural gas pipeline. As dramatized in the film Erin Brokovich (for which Julia Roberts won an Oscar), Erin and lawyer Edward Masry sued PG&E for contaminating the water near Hinckley in 1993. The settlement they won in 1996—$333 million—was the largest amount ever awarded for a direct-action lawsuit in the US at that time.

    Figure \(\PageIndex{5}\): (a) Erin Brockovich found that Cr(VI), used by PG&E, had contaminated the Hinckley, California, water supply. (b) The Cr(VI) ion is often present in water as the polyatomic ions chromate, \(\ce{CrO4^2-}\) (left), and dichromate, \(\ce{Cr2O7^2-}\) (right).

    Chromium compounds are widely used in industry, such as for chrome plating, in dye-making, as preservatives, and to prevent corrosion in cooling tower water, as occurred near Hinckley. In the environment, chromium exists primarily in either the Cr(III) or Cr(VI) forms. Cr(III), an ingredient of many vitamin and nutritional supplements, forms compounds that are not very soluble in water, and it has low toxicity. Cr(VI), on the other hand, is much more toxic and forms compounds that are reasonably soluble in water. Exposure to small amounts of Cr(VI) can lead to damage of the respiratory, gastrointestinal, and immune systems, as well as the kidneys, liver, blood, and skin.

    Despite cleanup efforts, Cr(VI) groundwater contamination remains a problem in Hinckley and other locations across the globe. A 2010 study by the Environmental Working Group found that of 35 US cities tested, 31 had higher levels of Cr(VI) in their tap water than the public health goal of 0.02 parts per billion set by the California Environmental Protection Agency.

    Molecular (Covalent) Compounds

    The bonding characteristics of inorganic molecular compounds are different from ionic compounds, and they are named using a different system as well. The charges of cations and anions dictate their ratios in ionic compounds, so specifying the names of the ions provides sufficient information to determine chemical formulas. However, because covalent bonding allows for significant variation in the combination ratios of the atoms in a molecule, the names for molecular compounds must explicitly identify these ratios.

    Compounds Composed of Two Elements

    When two nonmetallic elements form a molecular compound, several combination ratios are often possible. For example, carbon and oxygen can form the compounds CO and CO2. Since these are different substances with different properties, they cannot both have the same name (they cannot both be called carbon oxide). To deal with this situation, we use a naming method that is somewhat similar to that used for ionic compounds, but with added prefixes to specify the numbers of atoms of each element. The name of the more metallic element (the one farther to the left and/or bottom of the periodic table) is first, followed by the name of the more nonmetallic element (the one farther to the right and/or top) with its ending changed to the suffix –ide. The numbers of atoms of each element are designated by the Greek prefixes shown in Table \(\PageIndex{6}\).

    Table \(\PageIndex{6}\): Nomenclature Prefixes
    Number Prefix Number Prefix
    1 (sometimes omitted) mono- 6 hexa-
    2 di- 7 hepta-
    3 tri- 8 octa-
    4 tetra- 9 nona-
    5 penta- 10 deca-

    When only one atom of the first element is present, the prefix mono- is usually deleted from that part. Thus, \(\ce{CO}\) is named carbon monoxide, and \(\ce{CO2}\) is called carbon dioxide. When two vowels are adjacent, the a in the Greek prefix is usually dropped. Some other examples are shown in Table \(\PageIndex{7}\).

    Table \(\PageIndex{7}\): Names of Some Molecular Compounds Composed of Two Elements
    Compound Name Compound Name
    SO2 sulfur dioxide BCl3 boron trichloride
    SO3 sulfur trioxide SF6 sulfur hexafluoride
    NO2 nitrogen dioxide PF5 phosphorus pentafluoride
    N2O4 dinitrogen tetroxide P4O10 tetraphosphorus decaoxide
    N2O5 dinitrogen pentoxide IF7 iodine heptafluoride

    There are a few common names that you will encounter as you continue your study of chemistry. For example, although NO is often called nitric oxide, its proper name is nitrogen monoxide. Similarly, N2O is known as nitrous oxide even though our rules would specify the name dinitrogen monoxide. (And H2O is usually called water, not dihydrogen monoxide.) You should commit to memory the common names of compounds as you encounter them.

    Example \(\PageIndex{7}\): Naming Covalent Compounds

    Name the following covalent compounds:

    1. SF6
    2. N2O3
    3. Cl2O7
    4. P4O6

    Solution

    Because these compounds consist solely of nonmetals, we use prefixes to designate the number of atoms of each element:

    1. sulfur hexafluoride
    2. dinitrogen trioxide
    3. dichlorine heptoxide
    4. tetraphosphorus hexoxide

    Exercise \(\PageIndex{7}\)

    Write the formulas for the following compounds:

    1. phosphorus pentachloride
    2. dinitrogen monoxide
    3. iodine heptafluoride
    4. carbon tetrachloride
    Answer a

    PCl5

    Answer b

    N2O

    Answer c

    IF7

    Answer d

    CCl4

    Binary Acids

    Some compounds containing hydrogen are members of an important class of substances known as acids. The chemistry of these compounds is explored in more detail in later chapters of this text, but for now, it will suffice to note that many acids release hydrogen ions, H+, when dissolved in water. To denote this distinct chemical property, a mixture of water with an acid is given a name derived from the compound’s name. If the compound is a binary acid (comprised of hydrogen and one other nonmetallic element):

    1. The word “hydrogen” is changed to the prefix hydro-
    2. The other nonmetallic element name is modified by adding the suffix -ic
    3. The word “acid” is added as a second word

    For example, when the gas \(\ce{HCl}\) (hydrogen chloride) is dissolved in water, the solution is called hydrochloric acid. Several other examples of this nomenclature are shown in Table \(\PageIndex{8}\).

    Table \(\PageIndex{8}\): Names of Some Simple Acids
    Name of Gas Name of Acid
    HF(g), hydrogen fluoride HF(aq), hydrofluoric acid
    HCl(g), hydrogen chloride HCl(aq), hydrochloric acid
    HBr(g), hydrogen bromide HBr(aq), hydrobromic acid
    HI(g), hydrogen iodide HI(aq), hydroiodic acid
    H2S(g), hydrogen sulfide H2S(aq), hydrosulfuric acid

    Oxyacids

    Many compounds containing three or more elements (such as organic compounds or coordination compounds) are subject to specialized nomenclature rules that you will learn later. However, we will briefly discuss the important compounds known as oxyacids, compounds that contain hydrogen, oxygen, and at least one other element, and are bonded in such a way as to impart acidic properties to the compound (you will learn the details of this in a later chapter). Typical oxyacids consist of hydrogen combined with a polyatomic, oxygen-containing ion. To name oxyacids:

    1. Omit “hydrogen”
    2. Start with the root name of the anion
    3. Replace –ate with –ic, or –ite with –ous
    4. Add “acid”

    For example, consider H2CO3 (which you might be tempted to call “hydrogen carbonate”). To name this correctly, “hydrogen” is omitted; the –ate of carbonate is replace with –ic; and acid is added—so its name is carbonic acid. Other examples are given in Table \(\PageIndex{9}\). There are some exceptions to the general naming method (e.g., H2SO4 is called sulfuric acid, not sulfic acid, and H2SO3 is sulfurous, not sulfous, acid).

    Table \(\PageIndex{9}\): Names of Common Oxyacids
    Formula Anion Name Acid Name
    HC2H3O2 acetate acetic acid
    HNO3 nitrate nitric acid
    HNO2 nitrite nitrous acid
    HClO4 perchlorate perchloric acid
    H2CO3 carbonate carbonic acid
    H2SO4 sulfate sulfuric acid
    H2SO3 sulfite sulfurous acid
    H3PO4 phosphate phosphoric acid

    Summaries

    How Atoms Bond

    Video \(\PageIndex{6}\): A summary of chemical bonds.

    Metals (particularly those in groups 1 and 2) tend to lose the number of electrons that would leave them with the same number of electrons as in the preceding noble gas in the periodic table. By this means, a positively charged ion is formed. Similarly, nonmetals (especially those in groups 16 and 17, and, to a lesser extent, those in Group 15) can gain the number of electrons needed to provide atoms with the same number of electrons as in the next noble gas in the periodic table. Thus, nonmetals tend to form negative ions. Positively charged ions are called cations, and negatively charged ions are called anions. Ions can be either monatomic (containing only one atom) or polyatomic (containing more than one atom).

    Compounds that contain ions are called ionic compounds. Ionic compounds generally form from metals and nonmetals. Compounds that do not contain ions, but instead consist of atoms bonded tightly together in molecules (uncharged groups of atoms that behave as a single unit), are called covalent compounds. Covalent compounds usually form from two or more nonmetals.

    Naming Chemical Compounds

    Video \(\PageIndex{7}\): "You know what's an awful lot like unexpectedly waking up in Belguim with a bunch of people speaking several different languages you don't understand? Chemistry."

    Chemists use nomenclature rules to clearly name compounds. Ionic and molecular compounds are named using somewhat-different methods. Binary ionic compounds typically consist of a metal and a nonmetal. The name of the metal is written first, followed by the name of the nonmetal with its ending changed to –ide. For example, K2O is called potassium oxide. If the metal can form ions with different charges, a Roman numeral in parentheses follows the name of the metal to specify its charge. Thus, FeCl2 is iron(II) chloride and FeCl3 is iron(III) chloride. Some compounds contain polyatomic ions; the names of common polyatomic ions should be memorized. Molecular compounds can form compounds with different ratios of their elements, so prefixes are used to specify the numbers of atoms of each element in a molecule of the compound. Examples include SF6, sulfur hexafluoride, and N2O4, dinitrogen tetroxide. Acids are an important class of compounds containing hydrogen and having special nomenclature rules. Binary acids are named using the prefix hydro-, changing the –ide suffix to –ic, and adding “acid;” HCl is hydrochloric acid. Oxyacids are named by changing the ending of the anion (-ate to –ic, and -ite to -ous) and adding “acid;” H2CO3 is carbonic acid.

    Glossary

    binary acid
    compound that contains hydrogen and one other element, bonded in a way that imparts acidic properties to the compound (ability to release H+ ions when dissolved in water)
    binary compound
    compound containing two different elements.
    covalent bond
    attractive force between the nuclei of a molecule’s atoms and pairs of electrons between the atoms
    covalent compound
    (also, molecular compound) composed of molecules formed by atoms of two or more different elements
    ionic bond
    electrostatic forces of attraction between the oppositely charged ions of an ionic compound
    ionic compound
    compound composed of cations and anions combined in ratios, yielding an electrically neutral substance
    molecular compound
    (also, covalent compound) composed of molecules formed by atoms of two or more different elements
    monatomic ion
    ion composed of a single atom
    nomenclature
    system of rules for naming objects of interest
    oxyacid
    compound that contains hydrogen, oxygen, and one other element, bonded in a way that imparts acidic properties to the compound (ability to release H+ ions when dissolved in water)
    oxyanion
    polyatomic anion composed of a central atom bonded to oxygen atoms
    polyatomic ion
    ion composed of more than one atom

    Contributors and Attributions

    • Paul Flowers (University of North Carolina - Pembroke), Klaus Theopold (University of Delaware) and Richard Langley (Stephen F. Austin State University) with contributing authors. Textbook content produced by OpenStax College is licensed under a Creative Commons Attribution License 4.0 license. Download for free at http://cnx.org/contents/85abf193-2bd...a7ac8df6@9.110).

    • Adelaide Clark, Oregon Institute of Technology
    • Crash Course Chemistry: Crash Course is a division of Complexly and videos are free to stream for educational purposes.
    • TED-Ed’s commitment to creating lessons worth sharing is an extension of TED’s mission of spreading great ideas. Within TED-Ed’s growing library of TED-Ed animations, you will find carefully curated educational videos, many of which represent collaborations between talented educators and animators nominated through the TED-Ed website.
    • Fuse School, Open Educational Resource free of charge, under a Creative Commons License: Attribution-NonCommercial CC BY-NC (View License Deed: https://creativecommons.org/licenses/by-nc/4.0/)

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