Chemical Formulas and Chemical Compounds
A correctly written chemical formula must represent the known facts about the analytically
determined composition of the compound
Used to represent amounts of substance
-H2O can represent: 1 mole of water molecules
1 molecule of water
1 molar mass of water molecules
Monatomic Ions: Ions formed from a single atom
-Main group elements form ions with noble gas configurations
-Group 1 metals:
-Group 2 metals:
-Group 13 metals:
-Group 14: becomes difficult
-Tin and lead: Lose 2 e-, become +2
-Group 15,16,17 (non-metals)
-Form ions by gaining e-d block elements form 2+ and 3+ ions
-Monatomic ions are followed by the name “ion”
-Negative ions combined with a cation: Drop ending of element name and add “ide”
Binary Compounds: Compounds composed of 2 different elements
Binary Ionic Compound: Compound composed of a metal and a non-metal
1. Write cation first
2. Subscripts represent number of atoms present
a. If no subscript: One is understood
b. Charges are not included in chemical formulas
1. Compound is neutral
c. Charges can be crossed over to get subscripts
d. Find number of ions needed to equalize + and – charges
Steps:
1. Write symbols for ions side by side, positive ion first
2. Cross over charge values to give subscripts
3. Check subscripts and write formula
Na + Cl
Mg + Br
Al + O
1
Li2O
CaCl2
Different names are needed for positive ions of 2 different charges formed by the same metal (dblock)
2 Different naming systems
1. Old: Ion with lower charge - ous
Ion with higher charge – ic
Ex: Cu+
Cu2+
2. New: Charge is indicated by Roman numeral
(Stock system)
Ex: Cu+
Cu+2
Nomenclature: Term that refers to methods of naming chemical compound
Compounds Containing Polyatomic Ions
-All but ammonium are – charged
-Oxyanions: Polyatomic ions that contain O2
-In several cases 2 different oxyanions are formed from the same 2 elements
-NO2and
NO3(ite: less oxygen)
(ate: more oxygen)
-Ex:
Examples:
Sodium Carbonate
Aluminum Sulfate
Ammonium sulfate
Copper (II) nitrate
2
Fe(NO3)2
Cu2SO4
Binary Molecular Compounds
-Compounds between nonmetals: Molecular
-Old System (Greek prefixes)
-Least electronegative element is given 1st
-Oxygen and halogens are 2nd
-Prefix Rules
1. Used with name of 1st element
a. Only if more than one atom of that element is present
2. Second element
a. Prefix if more than 1 compound can be formed by the 2 elements
b. The root of the name of the second element
c. Ending: “ide” (indicates compound containing 2 elements)
Examples:
N2O5
P8Cl7
S3F9
Petacarbon hexachloride
Tetranitrogen octaphosphide
Trifluoride hexaiodide
Acids: Molecular compound that contains one or more H
Binary acid: H and a halogen (a more electronegative non-metal)
-Occurs in a water solution
-Ex: HCl
Oxyacid: Acids containing H, O, and a 3rd element
3
-Many polyatomic ions are produced by loss of H ions from oxyacids
Examples:
Sulfuric acid
Phosphoric acid
Nitric acid
Salt: Ionic compound composed of cation and anion of an acid
Ex: CaSO4
Hydrates: Compounds that attract and hold water molecules in their crystal structure
Water of Hydration: Water locked in compound, can be removed by heating
Anhydrous: Solid residue remaining after water has been removed
Formula of hydrated compound: Place raised dot after anhydrous formula followed by the
number of water molecules per formula unit of compound
Examples:
Copper II sulfate pentahydrate (hydrate: water)
Sodium carbonate hepta hydrate
Oxidation Numbers or oxidation states: Assigned to atoms in molecules, including molecular ions,
to show general distribution of electrons among bonded atoms
-Ownership is given to the more electronegative atom in bond
Rules for assigning oxidation Numbers
1. Uncombined elements: 0
2. Monatomic ion = to charge
3. Fluorine = -1
4. Oxygen = -2
-Except in: Peroxides H2O2: O = -1
Superoxides KO2: O = -1/2
Compounds with Fluorine O = +2
4
H = +1
-Except in compounds with metals (NaH) –1
6. The more electronegative element in a binary compound
-Oxidation # = to charge of ion
7. Algebraic sum of oxidation #s in neutral compound = 0
8. Algebraic sum of oxidation #s in polyatomic ions = the charge
5.
Practice: Assign the oxidation # for the purple element
1. CF4 ____
2. PCl3 _____
3. SO2 ______
4. HNO3 _____
5. SiO2 _______
6. KH ________
7. P2O5 _______
8. HClO2 _______
9. PCl3 _______
10. PCl5 _______
11. POCl3 _______
12. SO3 2-_________
Empirical Formula: Simplest ratio of atoms in a compound
Ex: B2H6 : Empirical Formula BH3
Steps:
1. Convert each amount to moles
2. Divide each mole amount by the smallest mole amount
a. Results in your subscripts
b. Must be close to a whole number; if not multiply by 2
3. Percents can be interchanged for grams
5
Ex:
If you have 2.128 g Cl and 1.203 g Ca what is the empirical formula?
If you have 3.50% Fe and 1.50% O what is the empirical formula?
Molecular Formula: Gives the actual # of atoms of each element in a molecular compound
(empirical formula)X = molecular formula
-X is the integral factor
-Multiply subscripts in empirical formula by the integral factor
X = Molecular formula mass/empirical formula mass
Ex:
If the molecular formula mass is 283.889g and the empirical formula is P2O5 what is the molecular
formula?
6