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C.7.A name ionic compounds containing main group transition metals, covalent compounds, acids, and bases, using International Union of
Pure and Applied Chemistry (IUPAC) nomenclature rules
C.7.B write the chemical formulas of common polyatomic ions, ionic compounds containing main group or transition metals, covalent compounds, acids
C.7.C construct electron dot formulas to illustrate ionic and covalent bonds
C.7.D describe the nature of metallic bonding and apply the theory to explain metallic properties such as thermal and electrical conductivity, malleability, and ductility
C.7.E predict molecular structure for molecules with linear, trigonal planar, or tetrahedral electron pair geometries using Valence Shell
Electron Pair Repulsion (VSEPR) theory

Naming Chemical Compounds… slide 4
Naming Acids and Bases… slide 11
Writing Chemical Formulas… slide 16
Chemical Bonding… slide 22
Molecular Geometry, VSEPR Theory… slide 31

C.7.A name ionic compounds containing main group transition metals, covalent compounds, using
International Union of Pure and Applied Chemistry
(IUPAC) nomenclature rules.

 There are three main types of compounds when working on Naming Compounds.
 Metal Binary Compounds (Ionic) – Contain a metal and a non-metal. They form an ionic bond.
 Non-Metal Binary Compounds (Molecular)–
Contain two non-metals. They form a covalent bond.
 Ternary Compounds (Polyatomic)– Contain polyatomic ions. The formula will have three or more elements in it.

 Name the first element. (This will always be the metal.)
 Replace the ending on the second element with an
“ide” ending. ( This element will be the non-metal)
 Example:
 NaCl sodium and chlorine becomes sodium chloride
 MgS magnesium and sulfur becomes magnesium sulfide

 When some atoms can have more than one possible charge, you name the charge on the atom.
 The following elements must have a roman numeral:
Cr-Cu, Au, Hg, Sn, & Pb
 Copper +1 and +2 Iron +2 and +3
Cu +1 is copper (I)
Cu +2 is copper (II)
Fe +2 is iron (II)
Fe +3 is iron (III)
CuCl is copper (I) chloride FeCl2 is iron (II) chloride
CuCl2 is copper (II) chloride FeCl3 is iron (III) chloride

 Name the first element
 Replace the ending on the second element with “ide”
 Use Prefixes to indicate the number of atoms in the formula.
 *Exception: A prefix is not required when the first element only has 1 atom.
 Ex:
 CO2 carbon and oxygen is carbon dioxide
 N2O nitrogen and oxygen is dinitrogen monoxide

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1 atom = mono
2 atoms = di
3 atoms = tri
4 atoms = tetra
5 atoms = penta
6 atoms = hexa
7 atoms = hepta
8 atoms = octa
9 atoms = nona
10 atoms = deca

 Name the first part of the compound. Element or polyatomic ion.
 Name the second part of the compound. Element or polyatomic ion.
 Example:
 MgSO4 NH4OH ammonium hydroxide magnesium sulfate
 K3PO4 potassium phosphate

 Acids without Oxygen are named with the prefix
“hydro” and end in “ic”
 Examples:
 HCl hydrochloric acid
 HF hydrofluoric acid
 HBr hydrobromic acid

 Acids with oxygen have several forms.
 The “ic” or regular ending for an acid comes from the polyatomic ion with the “ate” ending. This gives the regular count for the oxygen for this type of acid.
 Example:
 H2SO4
 SO4 is sulfate so this acid is called sulfuric acid

 Once you know the “ic” ending you count the number of oxygen in the other forms to find the name for the acid. (REMEMBER: The regular “ic” form comes from the polyatomic ion that ends with “ate”)
 Two less oxygen hypo ________ ous acid
 One less oxygen ________ ous acid
 Regular “ic” form ________ ic acid
 One more oxygen per ________ ic acid

 The other names for the acids will come from the count based from the “regular acid name”
 H2SO4 -ate ending so it is sulfuric acid
 H2SO3 -ite ending so it is sulfurous acid
 H2SO2 two less oxygen will have the prefix hypo and the –ous ending. hyposulfurous acid.
 H2SO5 one more oxygen will have a prefix per and the regular -ic ending. persulfuric acid

C.7.B write the chemical formulas of common polyatomic ions, ionic compounds containing main group or transition metals, covalent compounds, acids

 Write chemical symbol for each part of the compound.
 Write the charge (oxidation #) for the element.
Do the charges add together and equal zero?
 Yes, Stop this is the formula. The number of electrons given away is the same as what is being taken by the second atom.
 No, Cross the absolute value of the charge to the opposite element as a subscript. Multiply the new subscript by the charge and see if the new values will add together and equal zero. If yes, Stop you have the formula

 potassium bromide Formula
K +1 Br -1 +1 + -1 = 0 Yes KBr
 magnesium chloride
Mg +2 Cl -1 +2 + -1 = +1 No
Mg 1 Cl 2 Mg (1 x +2)= +2 Cl (2 x -1)= -2
Yes MgCl2

 Same rules as normal ionic compounds. The charge for the transition metal will come from the name of the compound.
 iron (III) chloride
 Fe +3 Cl -1 +3 + -1 = +2 No
 Fe1 Cl 3 Fe (1 x +3) +3 Cl (3 x -1) -3
Yes FeCl3

 Use the prefix to determine the subscript of each element in the formula.
 NO PREFIX on the first element indicates a subscript of 1
 Write the correct formula using the correct symbol and subscript for each element.
 Ex: carbon dixoide CO
2

 The rules for polyatomic ions will be the same as ionic compounds.
 *Polyatomic ions must be placed in parenthesis if the subscript is larger than 1 when criss-crossing.
 magnesium sulfate
 Mg +2 SO4 -2 MgSO4
 iron (III) phosphate
 Fe +3 PO4 -3 FePO4

 sodium hydroxide
 Na +1 OH -1 NaOH
 calcium hydroxide
 Ca +2 OH -1 Ca(OH)2 aluminum phosphate

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Write the symbol and charge (oxidation #) of each element.
If the charges do not add up to zero, criss-cross the oxidation #.
Ex: hydrosulfuric acid
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H +1 S -2 = H
2
S

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Write the symbol and the charge for the polyatomic ion (oxyanion).
If the charges do not add up to zero, criss-cross the oxidation numbers.
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1 more oxygen
MEMORIZED(-ate)
1 less oxygen (-ite) per ____________ic
____________ic
____________ous
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2 less oxygen hypo___________ous
H
2
SO
3 sulfurous acid

 There are three main types of chemical bonding. ionic, covalent, and metallic.
 Ionic bonding occurs when there is a transfer of electrons.
 Covalent bonding occurs when atoms share electrons.
 Metallic bonding consist of the attraction of free floating valance electrons for positively charged metal ions.

 Electro negativities are used to determine what type of bond is formed when atoms come together in a chemical reaction.
 To find the type of bond find the difference in the electronegativities.
 If the difference is greater than 1.67 an ionic bond is formed.
 If the difference is less than 1.67 a covalent bond is formed.

 All atoms want to obtain eight electrons in the valence energy level. To do so they will give, take, or share electrons.

 The element with the fewest atoms goes in the center.
 The other atoms go around the central atom.
 Show the transfer of the electrons with a positive for the atom that lost the electrons and a negative for the atoms that gain the electrons.

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NaCl sodium chloride sodium: (1.01) chlorine: (2.83)
2
2
6
1 Cl:
2
2
6
2
5
Sodium transfers the 3s 1 to chlorine to complete the 3p energy level.
The electronegativity difference is 1.72
An ionic bond is formed.

 The element with the fewest atoms goes in the center.
 The other elements go around the central atom.
 A bonding pair can only form where there is an unpaired electron.
 Shared pairs or bonding pairs are shown with a dash.
One dash equals two electrons.

 AsI 3 arsenic triiodide arsenic (2.20) iodine (2.21)
As: 1s 2 2s 2 2p 6
3s 2 3p 6 4s 2 3d 10 4p 3
I: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 4d 10 5p 5
The electronegativity difference is .01
A covalent bond is formed. The atoms share the electrons.

C.7.E predict molecular structure for molecules with linear, trigonal planar, or tetrahedral electron pair geometries using Valence Shell Electron Pair Repulsion (VSEPR) theory

 The shape that a covalently bonded substance will take is referred to as its Molecular Geometry.
 The shape is determined by the central atom, and the number of shared and unshared electron pairs around the atom.
 Electron pairs around the central atom will spread out as far as possible to minimize the repulsive forces.
 This gives bond angles depending on the shape.

Total number of electron pairs.
Number of shared pairs
2 2
Number of unshared pairs
Shape
0 Linear
Bond Angle
180 0

Total number of electron pairs.
Number of shared pairs
Number of unshared pairs
Shape
3 3 0
Trigonal
Planar
Bond Angle
120 0

Total number of electron pairs.
Number of shared pairs
Number of unshared pairs
4 4 0
Shape
Tetrahedral
Bond Angle
109.5
0

Total number of electron pairs.
Number of shared pairs
Number of unshared pairs
Shape
4 3 1
Bond Angle
Trigonal
Pyramidal 107.3
0

Total number of electron pairs.
Number of shared pairs
Number of unshared pairs
Shape
4 2 2 Bent
Bond Angle
104.5
0

Linear
Tetrahedral
Trigonal Planar
Trigonal Pyramidal
Bent