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- SO4

2-

• 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:

• 107Ag 106.90509 amu 51.84%

• 109Ag 108.90476 amu 48.16%

• 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.