Warm-Up
•Tell me everything you
know about Ionic
Bonding (3 minutes)
Ionic
Game
Warm-Up
1. If element X has an electronegativity of 2.7 and element Y has an
electronegativity of 1.6 the resulting bond is characteristic of
a. an ionic bond because the difference in electronegativity is less than 1.7
b. an ionic bond because the difference in electronegativity is greater than1.7
c. a covalent bond because the difference in electronegativity is less than 1.7
d. a covalent bond because the difference in electronegativity is greater than
1.7
2. Which term describes a substance that has a low melting point and does not
conduct electricity in either the solid or the liquid form?
a. Ionic and molecular
b. Ionic and metallic
c. Covalent and molecular
d. Covalent and metallic
Nomenclature &
Compositional
Stoichiometry
J. McLeod
What will we learn?
• How to use different naming conventions based
upon the type of compound
– Standard system for main group ionic cmpds
– Stock system for transition metal ionic cmpds
– Prefix system for molecules (non-metal/nonmetal)
• How to write formulas from names
• How to calculate percent mass of an element in a
compound
Oxidation Numbers
• The oxidation number is the charge an atom
takes when it loses or gains electrons and
becomes an ion.
• Oxidation Numbers are based on the
number of valence electrons
Chemical Compounds
• Atoms are not chemically stable until they
have 8 valence electrons (octet rule).
• Atoms gain, lose, or share electrons with
other atoms to become chemically stable.
Ionic
Compounds
Crystalline Lattice
Ionic Compounds
• Contains a metal and nonmetal
• The valence electrons from the metal are
transferred to the nonmetal.
• Metal is listed first, followed by the nonmetal.
• Change the ending of the nonmetal to ~ide.
– Examples: sulfide, oxide, phosphide.
• Ionic Examples:
Example: NaCl
Sodium Chloride
Example: BaF2
Barium Fluoride
Example: AlP
Aluminum Phosphide
Standard System Naming
Representative Metals:
• s block & p block metals
• Zn(+2), Cd(+2), Ag(+1)
Representative Metals only have 1 oxidation
number
Ionic Binary Compound
Naming Practice (Standard Naming)
1.Al2S3
1. Aluminum Sulfide
2.Ba3P2
2. Barium Phosphide
3.ZnO
3. Zinc Oxide
4.Ca2C
4. Calcium Carbide
5.Cs3As
5. Cesium Arsenide
6.Fr2O
6. Francium Oxide
More Ionic Binary Compounds
Naming Practice
1.FrF
1. Francium Fluoride
2.GaCl3
2. Gallium Chloride
3.RaBr2
3. Radium Bromide
4.Sr3N2
4. Strontium Nitride
5.Ag2O
5. Silver Oxide
6.Mg3P2
6. Magnesium Phosphide
The 5 steps for writing a binary ionic
compound formula
1. Write the symbols of the two elements
2. Write the oxidation number of each as
superscripts.
3. Drop the positive and negative signs
4. Crisscross the superscripts so they
become subscripts.
5. Reduce when possible
Formula for Aluminum Oxide
1. Write the symbols of the elements
Al
Formula for Aluminum Oxide
2. Write the oxidation numbers for each
element
Al
+3
-2
Formula for Aluminum Oxide
3. Drop the positive & negative sign.
Al
3
2
Formula for Aluminum Oxide
4. Crisscross the superscripts so they
become subscripts
Al
3
2
Formula for Aluminum Oxide
5. Reduce subscripts when possible. (not
possible here)
Al
2
3
Examples of Reduction
of Subscripts
Sr2O2
Al3P3
Pb2O4
Ba3N2
SrO
AlP
PbO2
Ba3N2
Ionic Binary Compound
Formula Practice
1.Potassium iodide
1. KI
2.Lithium bromide
2. LiBr
3.Magnesium oxide
3. MgO
4.Sodium nitride
4. Na3N
5. Radium phosphide
5. Ra3P2
6.Strontium chloride
6. SrCl2
More Ionic Binary Compound
Formula Practice
1. Rubidium Sulfide
1. Rb2S
2. Gallium Fluoride
2. GaF3
3. Thallium Iodide
3. TlI3
4. Cesium sulfide
4. Cs2S
5. Aluminum bromide 5. AlBr3
6. Potassium Nitride
6. K3N
Ionic Binary Compound
(Stock Naming)
Transition Metals:
• d & f blocks
• Sn+2 & Sn+4
• Pb +2 & Pb+4
Transition Metals
Roman numerals are used in the name of the transition
metal in the compound to show the oxidation number
of the cation.
Sc
+3
Ti
+3
+4
V
+5
+4
Cr
+3
+6
Mn
+4
+6
Fe
+2
+3
Co
+2
+3
Ni
+2
+3
Cu
+2
+1
Examples
Mn+4 Manganese (IV)
Mn+6 Manganese (VI)
Fe+2 Iron (II)
Fe+3 Iron (III)
Cu+1 Copper(I)
Cu+2 Copper (II)
Stock System Naming
Compounds are named by the name of
the transition metal cation followed by the
oxidation number as the roman numeral
and then proper name of the anion.
Cu2O
Copper (I) oxide
CuO
Copper (II) Oxide
Stock System Nomenclature
How does one determine the charge of the
transition metal?
1. Find the charge of the anion
2. Calculate total negative charges (multiple by
number of anions)
3. Cation must be of equal value but opposite charge
4. Divide by the number of available cations
5. Write the roman numeral corresponding to the
charge on the transition metal
Stock System Nomenclature
Examples:
+3 total -3 total
+3
-1
CoCl3
Cobalt (III) chloride
Stock System Nomenclature
Examples:
+6 total -6 total
+2
-3
Ni3P2
Nickel (II) Phosphide
Stock System Nomenclature
Examples:
+14 total -14 total
+7
-2
Mn2O7
Manganese (VII) Oxide
Naming binary compounds
containing a transition metal
1. Fe2O3
2. NiBr2
3. V2O5
4. Cu3P2
5. PbS2
6. MnO2
1. Iron (III) Oxide
2. Nickel (II) Bromide
3. Vanadium (V) Oxide
4. Copper (II) Phosphide
5. Lead (IV) Sulfide
6. Manganese( IV) Oxide
Naming binary compounds
containing a transition metal
1. Co2S3
2. CrI2
3. VO2
4. SnI4
5. Au2S
6. TiF3
1. Cobalt (III) Sulfide
2. Chromium (II) Iodide
3. Vanadium (IV) Oxide
4. Tin (IV) Iodide
5. Gold (I) Sulfide
6. Titantium (III) Fluoride
Formulas for Transition Metals
Binary Ionic Compounds
• Remember: The Roman Numeral IS the
oxidation number for the element.
Example: Gold (III) Oxide
Au2O3
Formulas for Transition Metals
1. Iron (II) Oxide
1. FeO
2. Lead(IV) Sulfide
2. PbS2
3. Vanadium (IV) Fluoride 3. VF4
4. Palladium (II) nitride
4. Pd3N2
5. Ruthenium (III) iodide
5. RuI3
Polyatomic Ions
Polyatomic (many atoms) ions are covalent
molecules with a charge. They behave as if
they were one-atom ion.
1
Polyatomic Ions
Look at pg 7 of your reference sheets.
What is the difference between ammonium and
the rest of the polyatomics?
• Treat polyatomic ions as you would a normal
ion – crisscross oxidation numbers to
determine the formula.
• The only difference is that when you have
more than one of a specific polyatomic ion in
a formula you must encase it in parenthesis.
Ternary Ionic Compounds
1. Ca(NO3)2
2. Magnesium Phosphate 2. Mg3 (PO4) 2
3. Ba(OH) 2
3. Barium Hydroxide
4. Na2SO4
4. Sodium Sulfate
5. K2CO3
5. Potassium Carbonate
6. Aluminum Phosphate
6. AlPO4
7. Gallium Sulfate
7. Ga2(SO4) 3
8. Cesium Acetate
8. CsC2H3O2
9. Strontium Permanganate
9. Sr(MnO4)2
10.Rubidium Dichromate
10.Rb2Cr2O7
1. Calcium Nitrate
Ionic Mixed Practice
Write the name
1. NH4Cl
2. CaBr2
3. GaP
4. Rb2CO3
5. Mg(NO2)2
6. CuSe
7. FeTe
Write the formula
8. potassium chlorate
9. gold (I) oxide
10. cadmium iodate
11. cesium fluoride
12. aluminum sulfide
13. nickel (II) Nitrate
14. chromium (VI) Nitride
Covalent
Compounds
Water Molecule
Hydrogen
Atom
Oxygen
Atom
Hydrogen
Atom
Covalent Compounds
• Contains 2 or more nonmetals
• The valence electrons are shared between the
nonmetals.
• Must use prefixes in the name.
• Name tells you the formula.
• Example: N2O4 – dinitrogen tetroxide
• You CANNOT reduce the formulas
• If the first element is singular there is no prefix.
• The second element must ALWAYS have a prefix.
Covalent Prefixes
Mono -1
Di – 2
Tri – 3
Tetra – 4
Penta – 5
Hexa – 6
Hepta – 7
Octa – 8
Non – 9
Deca -10
A prefix tells you the number of atoms of that
element in the compound.
Naming Covalent Compounds
1. N2O3
1. Dinitrogen trioxide
2. CH4
2. Carbon tetrahydride
3. PO5
3. Phosphorus pentoxide
4. S2F4
4. disulfur tetrafluoride
5. P4O10
5. Tetraphosphorus decoxide
Writing Formulas for Covalent
Compounds
1. Carbon Monoxide
1. CO
2. Phosphorus Pentachloride
2. PCl5
3. Sulfur hexafluoride
3. SF6
4. Dinitrogen trioxide
4. N2O3
5. Sulfur dioxide
5. SO2
Ionic & Covalent Structure
Ionic Compounds form a
Water
crystalline lattice –
repeating pattern of
ions.
Covalent compounds
form individual
molecules that are not
connected to each
Na+1 Ions
Cl-1 Ions
Sodium Chloride
other.
How do I know which naming system
to use?
Identify the FIRST element of the compound
as either:
– Representative Metal Standard System
NaCl LiBr Al2O3 Ba(OH)2
– Transition Metal
Stock System
CoCl2 CoCl3 Cr2O3 FeO
(ROMAN NUMERALS)
– Non-metal
Prefix System
CO2 PCl5 As2O5 NO2
Name These Compounds
1. PCl3
1. Phosphorus trichloride
2. Sr3N2
2. Strontium Nitride
3. KOH
3. Potassium Hydroxide
4. NH3
4. Nitrogen trihydride
5. CrCO3
5. Chromium (II) Carbonate
Write formulas for these compounds
1. Calcium Chromate
1. CaCrO4
2. Sodium Bromide
2. NaBr
3. Sulfur hexafluoride
3. SF6
4. Carbon tetrachloride
4. CCl4
5. Potassium Phosphate
5. K3PO4
Warm-Up – Label the metals as TM, RM, or NM
1.Aluminum Sulfide ____ _____________
2.AsCl5
____ ______________
3.Barium nitride
____ ______________
4.FeCl2
____ ______________
5.Calcium hydride
____ ______________
6.CCl4
____ ______________
7.Cd3N2
____ ______________
8.Cesium nitride
____ ______________
9.Chromium (III) oxide ____ ______________
10.Copper (I) sulfate ____ ______________
Acid Naming
Examples
• HI
• ~ide

hydroiodic acid

hydrochloric acid

hydrosulfuric acid
• HCl
• ~ide
• H2 S
• ~ide
Examples
• HNO3
• ~ate

nitric acid

carbonic acid

sulfurous acid
• H2CO3
• ~ate
• H2SO3
• ~ite
Examples
• hydroselenic acid
• -ide

H2Se

H2CrO4

HClO
• chromic acid
• -ate
• hypochlorous acid
• -ite
Acid Nomenclature
1. Nitric Acid
1. HNO3
2. Hydrochloric Acid
2. HCl
3. Acetic Acid
3. HC2H3O2
4. Hydrofluoric Acid
4. HF
5. Carbonic Acid
5. H2CO3
6. Nitrous Acid
6. HNO2
7. Phosphoric Acid
7. H3PO4
Acid Nomenclature
1. HBrO3
1. Bromic Acid
2. H2Cr2O7
2. Dichromic Acid
3. HClO4
3. Perchloric Acid
4. H2SO3
4. Sulfurous Acid
5. HIO3
5. Iodic Acid
6. HBr
6. Hydrobromic Acid
7. H2CrO4
7. Chromic Acid
Oxidation Numbers
Rules
1. F is always -1
2. H with nonmetals is +1; H with metals is -1
3. Oxygen = -2
4. Elements = 0
5. Metal ions are their charge
6. The most electronegative atom will take its
charge as if it was alone by itself as an anion
7. The sum of all oxidation numbers are equal to
the charge on the compound or ion
Oxidation Numbers
+1 + ? + -8 = 0
+1 +7 -2
KMnO4
Oxidation Numbers
+2 + ? + -6 = 0
+1 +4 -2
H2SO3
Oxidation Numbers
? + -6 = -1
+5 -2
ClO3
-1
Oxidation Numbers
? + -3 = 0
+3 -1
PCl3
Oxidation Number Practice
1.
2.
3.
4.
5.
6.
7.
SCl2
OF2
H2SO4
AsCl5
N2O4
K2CO3
Na2Cr2O7
1.
2.
3.
4.
5.
6.
7.
S = +2
O = +2
S=+6
As = +5
N = +4
C = +4
Cr=+6
Warm-Up
Formula:
1. Radium Phosphide
2. Iron (II) Hydrogen
Sulfate
3. Dinitrogen monoxide
Name:
4. K2Cr2O7
5. TiO2
6. CO
Acids:
7. Hydrosulfuric Acid
8. Thiocyanic Acid
9. HClO2
10. HMnO4
Warm-Up
Formula:
1. Radium Phosphide
2. Iron (II) Hydrogen
Sulfate
3. Dinitrogen monoxide
4. Hydrophosphoric Acid
Name:
5. K2Cr2O7
6. TiO2
7. CO
8. HClO2
Oxidation Numbers:
9. HNO2
10. KCN
11. Na2CO3
12. AlPO4
Review: Molar Mass
Molar mass is the mass of 1 mole of a compound. It is
calculated by adding each of the atomic masses of the
elements making up the compound
Example 1:
What’s the molar mass of water (H2O)
2H
1O
Water:
H: 2 * 1.008 g/mole = 2.02 g/mol
O: 1 * 16.00 g/mol = 16.00 g/mol
18.02 g/mol
Percent Composition
• Percent Composition is the ratio of the mass
of an element to the total mass of the
compound
• We assume we have one mole of compound
to make the calculations easier – we can use
the periodic table.
Molar Mass of Water
H2 O
18g/mol
Percent Composition
Na2SO4: Find % Composition for Sodium, Sulfur, & Oxygen
Na:
S:
O:
23g/mol x 2 Na = 46 g/mol
32g/mol x 1 S = 32 g/mol
16g/mol x 4 O = 64 g/mol
142 g/mol
46 g / mol
x100%  32.4%
142 g / mol
32 g / mol
x100%  22.5%
%S =
142 g / mol
64 g / mol
%O = 142 g / mol x100%  45.1%
% Na =
Percent Composition
KBrO3: Find % Composition for each element
K:
Br:
O:
39.1g/mol x 1 K = 39.1 g/mol
79.9g/mol x 1 Br= 79.9 g/mol
16g/mol x 3 O = 48 g/mol
167 g/mol
39.1g / mol
%K = 167 g / mol x100%  23.4%
79.9 g / mol
x100%  47.8%
%Br =
167 g / mol
%O =
48.0 g / mol
x100%  28.7%
167 g / mol
Percent Composition of Hydrates
Hydrates: % water in a hydrate
CuSO4·12H2O
Molar Mass of H2O is: 18 g/mol
Cu:
S:
O:
H2O:
63.5 g/mol x 1 Cu
32 g/mol x 1 S
16g/mol x 4 O
18 g/mol x 12 H2O
=
=
=
=
63.5 g/mol
32.0 g/mol
64 g/mol
216 g/mol
375.5 g/mol
216 g / mol
x100%  57.5%
% H2O = 375.5 g / mol
Percent Composition
Na2SO4·5H2O
Na
S
O
H2O
23 g/mol x
32 g/mol x
16 g/mol x
18 g/mol x
2 Na
1S
4H
5 H2O
=
=
=
=
% H2O = 90 g / mol x100%  38.8%
232 g / mol
46 g/mol
32 g/mol
64 g/mol
90 g/mol
232 g/mol
Warm-Up
1.What is the oxidation number for
carbon in H2CO3
2.Name: SnO
3.Write the Formula: dinitrogen
trioxide
4. Find % Comp of Carbon in #1.
Percent Composition
1. Ammonium Sulfite (H)
2. Aluminum Acetate (C)
3. Copper (II) Hydroxide (O)
4. Magnesium Carbonate (Mg)
5. Iron (II) Phosphate (P)
6. Beryllium Nitride (Be)
7. Lithium Phosphide (Li)
Empirical vs. Molecular Formulas
• Empirical Formula: A chemical formula that
indicates the relative proportions of the
elements in a molecule rather than the actual
number of atoms of the elements. (most
reduced)
• Molecular Formula: A chemical formula that
shows the number and kinds of atoms in a
molecule. (can be reduced)
Example
1. Which of the following is an empirical
formula?
A. H2O
B. C2H2
C. H2O2
D. C12H6O12
Example
2. Which of the following is a molecular
formula?
A. C6H12O11
B. SO2
C. CH4
D. C6H12
Empirical Formula Calculations
1. Treat your percentages as grams (out of
100) Thus, % will become mass in grams. E.g.
69.58 % Ba becomes 69.58 g Ba.
2. Calculate the # of moles (mol = g ÷ g/mol)
3. Divide by the lowest number of moles
4. Write the simplest formula from mol ratios
Determining Empirical Formulas
1. A compound consists of 32.4 % Na, 22.5 % S, and 45.1
% O. Determine the simplest formula.
2.
Na: 32.4 g ÷ 22.99 g/mol = 1.409 mol Na
S: 22.5 g ÷ 32.06 g/mol = 0.702 mol S
O: 45.1 g ÷ 16.00 g/mol = 2.819 mol O
3.
Na
S
O
mol
1.409
0.702
2.819
mol
1.409
0.702
2.819
(reduced)
/ 0.702
/ 0.702
/0.702
=2
=1
=4
4: the simplest formula is Na2SO4
Determining Empirical Formulas
1. A compound consists of 69.58% Ba, 6.090% C, 24.32%
O. Determine the simplest formula.
2: Ba: 69.58 g  137.33 g/mol = 0.50666 mol Ba
C: 6.090 g 12.01 g/mol = 0.50708 mol C
O: 24.32 g  16.00 g/mol = 1.520 mol O
3.
mol
mol
(reduced)
Ba
C
O
0.507
0.507
1.520
0.507/
0.507/
1.520/
0.507 = 0.507 = 0.507 =
1
1.001
3.000
4: the simplest formula is BaCO3
Determining Empirical Formulas
1. A compound consists of 92.3% C, 7.7%H. Determine
the simplest formula.
2: C: 92.3 g  12.01 g/mol = 7.69 mol C
H: 7.7 g 1.00 g/mol = 7.7 mol H
3.
mol
mol
(reduced)
C
H
7.69
7.7
7.69/
7.7/ 7.69
7.69 = 1 = 0.99
4: the simplest formula is CH
Determining Molecular Formulas
Using the previous empirical formula if my
molecular formulas has a mass of 78g, what
would be the molecular formula?
Mass of CH = 13g
78/13= 6
C6H6
Chemical Formulas Vocab
•
•
•
•
•
•
•
•
•
•
•
•
Acids
Anion
Binary Compound
Cation
Chemical Bond
Covalent Compounds
Crystalline Lattice
Diatomic
Empirical Formula
Hydrates
Ionic Compound
Molar Mass
•
•
•
•
•
•
•
•
•
•
•
•
Mole Ratio
Molecular Formula
Molecule
Monoatomic Molecule
Oxidation Number
Percent Composition
Polyatomic Ion
Prefix System
Standard System
Stock System
Ternary Compound
Valence Electrons
Things you should know
The rules for writing the names of salts and molecules.
The rules for naming salts and molecules.
The rules for naming & writing formulas for acids
Determine oxidation numbers.
Calculate the weight percent of any element in a compound from its
formula.
• Derive the empirical formula of a compound or salt from the %
composition information.
• Determine molecular formulas from empirical formulas
•
•
•
•
•
Warm-Up
1. What is the empirical formula for a
compound that has 85.7% C & 14.3% H.
2. What is the molecular formula if the molar
mass is 84g?
3. What is the oxidation number for carbon in
C6H12?
4. What is the name for BaSO4?
5. What is the formula for Iron(III)Oxide?
Warm-Up
1. An example of an empirical formula is
a.
b.
c.
d.
C2 H 2
H 2 O2
CaCl2
C2Cl2
a.
b.
c.
d.
Cannot be determined
Remains the same
Increases
Decreases
2. As the mass number of the isotopes of
hydrogen increase, the number of electrons
WarmUp
1. Which figure correctly represents the Lewis
structure for that atom?
2. Which physical changes are endothermic?
a. Melting and freezing
b. Melting and evaporation
c. Condensation and sublimation
d. Condensation and deposition
WarmUp
1. In which compound do the atoms have the greatest
difference in electronegativity?
a. NaBr
b. AlCl3
c. LiI
d. KF
2. Generally, how many valence electrons are needed for
atoms to be most stable?
a. 6
b. 8
c. 18
d. 32
WarmUp
1. Compared to the atomic radius of a sodium atom, the atomic radius of a
magnesium atom is smaller. The smaller radius is primarily a result of the
magnesium atom having
2.
3.
4.
a.
b.
c.
d.
More principal energy levels
Fewer principal energy levels
A larger nuclear charge
A smaller nuclear charge
a.
b.
c.
d.
A beta particle
A neutron
Deuterium
An alpha particle
a.
b.
c.
d.
Lower electronegativities and higher ionization energies
Lower electronegativities and lower ionization energies
Higher electronegativities and higher ionization energies
Higher electronegativities and lower ionization energies
a.
b.
c.
d.
Al and Ba
Ni and P
Cl and S
Na and K
The particle has a mass of approximately four amu and a positive 2 charge?
Compared to atoms of metals, atoms of nonmetals generally have
Which two elements have the most similar chemical properties?