Chapter 4

Ionic Compounds

Chapter 4 suggested problems: 31, 32, 36, 38, 40, 48, 50, 60, 62, 64, 66, 68, 70, 94

  1. Ions

    1. Substances that have gained or lost one or more valence electrons

    2. Metals tend to form cations, nonmetals tend to form anions

    3. Ion formation, the Periodic Table, and the Octet rule

    4. Ionization energy - energy is required to add or lose electrons but if the resulting compound is more stable (i.e. lower in energy) it can occur

  2. Chemical bonds

    1. Molecules - two or more atoms held tightly together by chemical bonds

    2. A chemical bond occurs between two atoms

    3. Each atom donates one electron to the bond, and the way the electron is donated is a function of the type of bond

    4. Bonds are based on the electrostatic attraction of the positively-charged nuclei of the bonding atoms for the negatively-charged electrons

      1. F = (kq1q2)/r2

    5. Three types of chemical bonds and three types of compounds

      1. Metallic bonds - between M-M, "electron sea" model

      2. Ionic bonds - between cations and anions, usually between M-NM, or between monatomic and polyatomic ions, bonding electrons not shared equally (ion formation or electron transfer)

      3. Covalent bonds - between NM-NM, bonding electrons shared more or less equally

    6. Metallic compounds

      1. Have metallic bonds

      2. Includes pure metals and alloys

      3. "Electron sea" model

    7. Ionic compounds

      1. Held together by ionic bonds

      2. Includes salts and most rocks and minerals

      3. Ionic compounds dissociate in water if they dissolve and are capable of carrying charge (conducting electricity)

      4. High MP and BP and the relationship between MP/BP and the strength of bonding forces

    8. Molecular compounds (covalent compounds)

      1. Held together by covalent bonds

      2. Includes organic compounds and polymers

      3. Many molecular compounds dissolve but most of them do not dissociate

      4. Polymer - a molecular chain of many smaller repeating units called monomers

    9. The behavior of the atoms in forming bonds or ions can be predicted by their position in the Periodic Table

      1. Is the element a metal or a nonmetal?

      2. Of which group is the element a member?

      3. The Octet Rule: atoms are more stable when they are isoelectronic with the nearest noble gas

    10. In forming a chemical bond, both bonding atoms become isoelectronic with the nearest noble gas (or progress in that direction)

      1. Ionic compounds through forming ions

      2. Covalent compounds through sharing electrons

  3. Lewis electron dot structures (Lewis structures) for atoms

    1. Octet rule - generally governs observed behavior

    2. Lewis structures generally consist of the elemental symbol surrounded by one dot for each valence electron of the substance

      1. Valence electrons are the outer shell s and p electrons

      2. Electrons in filled d shells behave as inner core electrons

      3. In partially filled d shells the d electrons are valence electrons (transition metals), but the electrons in filled d subshells are unreactive

    3. Examples - 2nd and 3rd period elements

  4. Nomenclature of ionic compounds

    1. Ionic compounds are those that contain ionic bonds; between metals and nonmetals

    2. We will learn the IUPAC rules of nomenclature (systematic nomenclature) for ionic and covalent compounds

    3. Do not worry about the Stock (old) nomenclature system (e.g. ous, ic, ate, ite, hypo, per)

    4. The nomenclature of organic compounds has its own set of rules that will not be discussed in this class (you'll have to wait until Chem 1120)

    5. Naming monatomic cations

      1. Group 1 and 2 metals have only one cation: element name + "ion"

      2. Transition metals and p-block metals generally have more than one cation: element name + (charge in Roman numerals) + "ion" - (see Table 4.2, p. 83, McMurry & Castellion)

    6. Naming monatomic anions: element name - end + "ide" (see Table 4.3, p. 83, McMurry & Castellion)

    7. Polyatomic ions (see Table 4.4, p. 84, McMurry & Castellion)

      1. Polyatomic cations: Hg22+ mercury (I), NH4+ammonium

      2. Polyatomic anions
        SO42- sulfate
        SO32- sulfite
        NO3- nitrate
        NO2- nitrite
        OH- hydroxide
        CO32- carbonate
        CrO42- chromate
        MnO4- permanganate
        C2H3O2- acetate
        PO43- phosphate
        CN- cyanide
        ClO4- perchlorate
        ClO3- chlorate
        ClO2- chlorite
        ClO- hypochlorite

      3. Note that many of these are oxyanions

    8. Ionic compound names

      1. Cations are named first, anions are named second

      2. When naming the compound, drop the "ion" portion of the cation name when adding the cation name to the anon name

    9. Formulas of ionic compounds

      1. Molecular formula - a shorthand notation indicating the types and numbers of atoms in a compound

        1. By convention the molecular formulae of ionic compounds list cations first, then anions

        2. Shortcomings: 2 or more compounds may have the same molecular formula but with the atoms arranged differently - structural (constitutional) isomers e.g. ethanol vs. dimethyl ether

      2. Compounds are electrically neutral, so there must be a balance between the net positive charge of the cations and the net negative charge of the anions

      3. Note: you must pay attention to charge when writing the formulas of ionic compounds, worry about mass balance later

      4. "Cross-multiplying" rule and warning

      5. Examples: going from formulas to names and from names to formulas

        1. FeCl3, NaNO3, (NH4)2SO4, Ba3P2

        2. Potassium chlorate, aluminum sulfite, sodium acetate, rubidium oxide

  5. Nomenclature of binary molecular compounds

    1. Covalent compounds contain covalent bonds; between nonmetals and nonmetals

    2. Binary compounds consist of two elements

    3. Element order established by convention (i.e., history): right to left and bottom to top

      1. B - Si C - Sb As P N - H - Te Se S - I Br Cl - O - F

        1. Only exceptions to rule are H and O

    4. Naming rules

      1. Compound name has elements in same order as molecular formula

      2. First element: exact name

      3. Second element: "ide" suffix (named as if it is an anion)

      4. Prefixes denote numbers of atoms in compound (subscripts in molecular formulae):
        1
        mono
        2
        di
        3
        tri
        4
        tetra
        5
        penta
        6
        hexa
        7
        hepta
        8
        octa
        9
        nona
        10
        deca

        1. Exception 1: first element by itself is never "mono" e.g. nitrogen dioxide

        2. Exception 2: chop "o" / "a" from prefix is element name begins with a vowel e.g. carbon monoxide

      5. Examples

        1. XeF6, KrF2, ICl5, N4S4, P2O5, NO, N2O

        2. Iodine heptafluoride, dinitrogen pentoxide, tetraphosphorus decoxide

  6. Nomenclature of acids and bases (see Table 4.6, p. 89, McMurry & Castellion)

    1. Acids are substances that can donate a hydrogen ion

    2. Bases are substances that can donate a hydroxide ion

    3. The names of acids and bases are based on common accepted names and not on the systematic IUPAC nomenclature

    4. Some typical acids:
      nitric acid HNO3
      nitrous acid HNO2
      sulfuric acid H2SO4
      sulfurous acid H2SO3
      hydrofluoric acid HF
      hydrochloric acid HCl
      hydrobromic acid HBr
      hydroiodic acid HI
      carbonic acid H2CO3
      phosphoric acid H3PO4
      perchloric acid HClO4
      chloric acid HClO3
      chlorous acid HClO2
      hypochlorous acid HClO

    5. Some typical bases
      lithium hydroxide LiOH
      sodium hydroxide NaOH
      potassium hydroxide KOH
      rubidium hydroxide RbOH
      cesium hydroxide CsOH
      calcium hydroxide Ca(OH)2
      strontium hydroxide Sr(OH)2
      barium hydroxide Ba(OH)2
      ammonium hydroxide NH4OH
      ammonia NH3

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