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AP Chemistry. Chapter 2B Mr. Solsman. Types of Formulas: 1. Empirical—shows the relative number of atoms of each element in the compound. hydrogen peroxide HO. 2. Molecular—shows the actual number of atoms of each element in a molecule of a compound. hydrogen peroxide H 2 O 2.

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AP Chemistry

Chapter 2B

Mr. Solsman


  • Types of Formulas:

  • 1. Empirical—shows the relative number of atoms of each element in the compound.

  • hydrogen peroxide HO


  • 2. Molecular—shows the actual number of atoms of each element in a molecule of a compound.

  • hydrogen peroxide H2O2


  • 3. Structural—shows the number of atoms and the bonds between them.

  • hydrogen peroxide H-O-O-H


  • Chemical Names

  • 1. Members of a periodic group have the same ionic charge.

  • Group 1A is 1+

  • Group 2A is 2+

  • Group 3A is 3+


  • For A-group cations, ion charge = group number.

  • Exceptions: Sn2+ and Pb2+


  • For A-group cations, ion charge = group number.

  • Exceptions: Sn2+ and Pb2+

  • For A-group anions: ion charge = group number minus 8.

  • O2-, F-


  • Group B elements can form more than one ion. Unfortunately these have to be memorized.


Naming Simple Compounds

  • Type I Binary Ionic Compounds

    • Consists of a cation and an anion.

    • Rules for Naming:

      • A. The cation is always named first & anion second.

      • B. A monatomic cation takes its name from the name of the element. (Many end in –ium)

      • C. A monatomic anion is named by taking the root of the element name and adding –ide.


Naming Type I

  • Name each binary compound

  • CsF

  • AlCl3

  • LiH


Naming Type I

  • Name each binary compound

  • CsF Cesium fluoride

  • AlCl3 Aluminum chloride

  • LiH Lithium hydride


  • Because ionic compounds are arrays of oppositely charged ions, formula units give the relative number of cations and anions in a compound. Ionic compounds generally have only empirical formulas.


  • Because ionic compounds are arrays of oppositely charged ions, formula units give the relative number of cations and anions in a compound. Ionic compounds generally have only empirical formulas.

  • Exceptions: peroxides such as Na2O2 and mercury(I) compounds (Hg2Cl2) have empirical formulas of NaO and HgCl.


  • Ionic compounds have zero net charge so the cation’s positive charges must balance the negative charges of the anions.

  • The criss-cross method can be used to balance charges:

  • Mg2+ and Cl-

  • Ca2+ and O2- 


  • Reduce the subscripts to the smallest whole number that retains the ration of ions.

  • Thus Ca2O2 becomes CaO


  • Complications: some metals form more than one ion. Particularly group B metals. Cobalt for example forms Co2+ and Co3+.

  • Two naming systems are used to describe which is which—systematic and the common or trivial.


Ionic Binary Type II

  • Roman numerals must be used to indicate the charge on the cation if it can have more than one oxidation state.

  • An old system states that the ion with the higher charge has an ending of –ic and the one with the lower charge has an ending of –ous.


Naming Type II

  • Give the systematic name of the following:

  • CuCl

  • HgO

  • Fe2O3

  • MnO2

  • PbCl2


Naming Type II

  • Give the systematic name of the following:

  • CuCl Copper (I) chloride

  • HgO Mercury (II) oxide

  • Fe2O3 Iron(III) oxide

  • MnO2 Manganese(IV) oxide

  • PbCl2 Lead(II) chloride


Naming

  • Name the following:

  • CoBr2

  • CaCl2

  • Al2O3

  • CrCl3


Naming

  • Name the following:

  • CoBr2 Cobalt(II) bromide

  • CaCl2 Calcium chloride

  • Al2O3 Aluminum oxide

  • CrCl3 Chromium(III) chloride


Polyatomic Ions

  • Polyatomic ions are assigned special names that must be memorized.

  • The following table lists the more common polyatomic ions.


  • Polyatomic ions are covalently bonded atoms that have a net charge which can behave ionically. The polyatomic unit stays together as a unit.

  • Ca2+ + 2 NO3- Ca(NO3)2

  • 2 H+ + SO42-  H2SO4


Naming

  • Name the following polyatomic ions:

  • Na2SO4

  • KH2PO4

  • Fe(NO3)3

  • Na2SO3

  • Na2CO3

  • NaHCO3


Naming

  • Name the following polyatomic ions:

  • Na2SO4 Sodium sulfate

  • KH2PO4 Potassium dihydrogen phosphate

  • Fe(NO3)3 Iron(III) nitrate

  • Na2SO3 Sodium sulfite

  • Na2CO3 Sodium carbonate

  • NaHCO3 Sodium hydrogen carbonate


  • Oxoanion—Most polyatomic ions are oxoanions. These are ions in which a nonmetal is bonded to one or more O atoms.

  • NO2- NO3-

  • SO32- SO42-


  • A. Two oxoanions in the family:

  • The ion with more O atoms takes the nonmetal root and the suffix –ate.

  • The ion with fewer O atoms takes the nonmetal root and the suffix –ite.


  • NO3- nitrate

  • NO2- nitrite

  • SO42- sulfate

  • SO32- sulfite


  • B. Four oxoanions (usually a halogen bonded to oxygen).

  • ClO- ClO2- ClO3- ClO4-


  • ClO4- perchlorate

  • ClO3- chlorate

  • ClO2- chlorite

  • ClO- hypochlorite


  • Hydrates are compounds that have a specific number of water molecules associated with each formula unit.

  • In their formulas, this number is shown after a centered dot.

  • Cu(NO3)2•3H2O CuSO4•5H2O


Binary Covalent Type III

  • Binary covalent compounds are formed between two nonmetals. BrCl3

  • Rules for naming:

    • A. The first element in the formula is named first, using the full name of the element

    • B. The second element is named as if it were an anion.


  • C. Prefixes are used to denote the number of atoms present. (di, tri, tetra, penta, etc.)

  • D. The prefix mono- is never used for naming the first element.


Naming

  • Name the following compounds:

  • PCl5

  • PCl3

  • SF6

  • SO3

  • SO2

  • CO2


Naming

  • Name the following compounds:

  • PCl5 Phosphorus pentachloride

  • PCl3 Phosphorus trichloride

  • SF6 Sulfur hexafluoride

  • SO3 Sulfur trioxide

  • SO2 Sulfur dioxide

  • CO2 Carbon dioxide


  • P4O10

  • Nb2O3

  • Ti(NO3)4


  • P4O10 Tetraphosphorus decaoxide

  • Nb2O3 Niobium(III) oxide

  • Ti(NO3)4 Titanium(IV) nitrate


  • Molecular Masses

  • The molecular mass, formerly molecular weight, is the sum of the atomic masses of the formula unit or molecular compound.


  • NaCl = 23.9898 + 35.4527 = 59.4425 amu

  • H2O = 2(1.0079) + 15.9994 = 18.0152 amu

  • CuSO4•5H2O = 63.546 + 32.066 + 4(15.999) + 10(1.008) + 5(15.999) =

  • 249.683 amu


  • Recall that 1 amu = 1/12 the mass of a carbon-12 atom.

  • The unit amu has recently been replaced by the Dalton, D.


  • Silver (Z = 47) has 46 known isotopes, but only two occur naturally. 107Ag and 109Ag. Given the following mass spec data, calculate the atomic mass of Ag:

  • 107Ag106.90509 amu51.84%

  • 109Ag108.90476 amu48.16%


  • Boron (Z =5) has two naturally occurring isotopes. Calculate the percent abundances of 10B and 11B from the following: atomic mass of B = 10.81 amu; isotopic mass of 10B = 10.0129 amu; isotopic mass of 11B = 11.0093 amu.


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