Ap chemistry
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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

AP Chemistry

Chapter 2B

Mr. Solsman


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  • Types of Formulas:

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

  • hydrogen peroxide HO


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

  • hydrogen peroxide H2O2


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

  • hydrogen peroxide H-O-O-H


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  • 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+


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

  • Exceptions: Sn2+ and Pb2+


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

  • Exceptions: Sn2+ and Pb2+

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

  • O2-, F-


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


Naming simple compounds

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

Naming Type I

  • Name each binary compound

  • CsF

  • AlCl3

  • LiH


Naming type i1

Naming Type I

  • Name each binary compound

  • CsF Cesium fluoride

  • AlCl3 Aluminum chloride

  • LiH Lithium hydride


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  • 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.


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  • 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.


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  • 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- 


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

  • Thus Ca2O2 becomes CaO


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  • 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

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

Naming Type II

  • Give the systematic name of the following:

  • CuCl

  • HgO

  • Fe2O3

  • MnO2

  • PbCl2


Naming type ii1

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

Naming

  • Name the following:

  • CoBr2

  • CaCl2

  • Al2O3

  • CrCl3


Naming1

Naming

  • Name the following:

  • CoBr2 Cobalt(II) bromide

  • CaCl2 Calcium chloride

  • Al2O3 Aluminum oxide

  • CrCl3 Chromium(III) chloride


Polyatomic ions

Polyatomic Ions

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

  • The following table lists the more common polyatomic ions.


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  • 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


Naming2

Naming

  • Name the following polyatomic ions:

  • Na2SO4

  • KH2PO4

  • Fe(NO3)3

  • Na2SO3

  • Na2CO3

  • NaHCO3


Naming3

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


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  • 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-


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  • 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.


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  • NO3- nitrate

  • NO2- nitrite

  • SO42- sulfate

  • SO32- sulfite


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

  • ClO- ClO2- ClO3- ClO4-


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  • ClO4- perchlorate

  • ClO3- chlorate

  • ClO2- chlorite

  • ClO- hypochlorite


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  • 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 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.


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  • 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.


Naming4

Naming

  • Name the following compounds:

  • PCl5

  • PCl3

  • SF6

  • SO3

  • SO2

  • CO2


Naming5

Naming

  • Name the following compounds:

  • PCl5 Phosphorus pentachloride

  • PCl3 Phosphorus trichloride

  • SF6 Sulfur hexafluoride

  • SO3 Sulfur trioxide

  • SO2 Sulfur dioxide

  • CO2 Carbon dioxide


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  • P4O10

  • Nb2O3

  • Ti(NO3)4


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  • P4O10 Tetraphosphorus decaoxide

  • Nb2O3 Niobium(III) oxide

  • Ti(NO3)4 Titanium(IV) nitrate


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  • Molecular Masses

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


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  • 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


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

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


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  • 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%


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  • 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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