1. Draw the Lewis dot structures for the following molecular compounds:

a. CCl4

b. SF6

c. O2

d. CN-

e. SO3

2. Decide if the bonds are polar or nonpolar:

a. H-Br b. N-N c. P-Cl

Unit 4: Chemical Formulas and Reactions - Chapter 9 & 11

Part One: Chemical Formulas

1. Distinguish among atoms, a molecular formula for covalent molecules, and a formula unit for ionic compounds.

2. Distinguish between ionic compounds and covalent molecules. 3. Memorize the charges of common monoatomic ions. 4. Know that most transition metals that have multiple charges, and use the Stock Naming System

(Roman numerals) naming these metals in an ionic compound. 5. Memorize the names, formulas and charges of the common polyatomic ions. 6. Write chemical formulas for binary ionic compounds (only 2 elements present). 7. Name binary ionic compounds when given the chemical formula. 8. Identify by name and write the chemical formulas for ternary ionic compounds (with polyatomic

ions).

4

Tetrahedral; 109˚: Usually

nonpolar CH4

CF4

Trigonal Pyrimidal: 107˚

Usually polar: NH3, PCl3

Bent: 104.5˚

Usually polar: H2O, OF2

9. Identify the name of a covalent molecule from the formula. Write the chemical formula for a covalent molecule given the name. 13. Identify by name and write formulas for common acids.

Naming Acids:

Naming AcidsNormal ending

____-ide

____-ate

____-ite

Acid name is…

hydro-___-ic acid

_____-ic acid

_____-ous acid

1. If the anion attached to the

hydrogen ion ends in -ide, the prefix

hydro is used and the ending

changes to -ic.

Example : HCl - hydrochloric acid

2. If the anion ends in -ate

(polyatomic) then the ending

changes to -ic.

Example: HClO3 - chloric acid

3. If the anion ends in -ite the

ending changes to -ous.

Example: HClO2 - chlorous acid

Practice Questions:

1. Write the correct chemical formulas for the compounds a-i and write the names for j-o.

a. Sodium chloride j. BaSO4

b. Iron (II) bromide k. NH4NO3

c. Potassium oxide l. Mn2O3

d. Aluminum sulfide m. AgC2H3O2

e. Calcium hypochlorite n. PbCl2

f. Magnesium nitrate o. Zn(ClO3)2

g. Carbon dioxide p. HCl

h. Phosphorus pentachloride q. HNO3

i. Sulfur trioxide r. H2SO4

Part Two: Chemical Reactions

A. Equations show:

1. the reactants which enter into a reaction.

2. the products which are formed by the reaction.

3. the amounts of each substance used and each substance produced.

B. Two important principles to remember:

1. Every chemical compound has a formula which cannot be altered.

2. A chemical reaction must account for every atom that is used. This is an application of the Law of

Conservation of Matter which states that in a chemical reaction atoms are neither created nor

destroyed.

C. Some things to remember about writing equations:

1. The diatomic elements when they stand alone are always written H2, N2, O2, F2, Cl2, Br2, I2

2. The sign, → , means "yields" and shows the direction of the action.

3. A small delta, (), above the arrow shows that heat has been added.

4. A double arrow, ↔ , shows that the reaction is reversible and can go in both directions.

5. Before beginning to balance an equation, check each formula to see that it is correct. NEVER

change a formula during the balancing of an equation.

6. Balancing is done by placing coefficients in front of the formulas to insure the same number of

atoms of each element on both sides of the arrow.

Practice Balancing Equations

a. _____C6H12O6 + _____O2 → ____CO2 + _____H2O Type: _____________________

b. _____Fe2S3 + _____HCl → _____FeCl3 + _____H2S Type: _____________________

7. Always consult the Activity Series of metals and nonmetals before attempting to write equations

for single replacement reactions.

D. Four basic types of chemical reactions:

1. Combination (synthesis):

two or more elements or compounds may combine to form a more complex compound.

Basic form: A + X → AX

Examples of combination reactions:

1. Metal + oxygen → metal oxide 2Mg(s) + O2(g) → 2MgO(s)

2. Nonmetal + oxygen → nonmetallic oxide C(s) + O2(g) → CO2(g)

3. Metal oxide + water → metallic hydroxide MgO(s) + H2O(l) → Mg(OH)2(s)

4. Nonmetallic oxide + water → acid CO2(g) + H2O(l) → H2CO3(aq)

5. Metal + nonmetal → salt 2 Na(s) + Cl2(g) → 2NaCl(s)

6. A few nonmetals combine with each other. 2P(s) + 3Cl2(g) → 2PCl3(g)

2. Decomposition:

A single compound breaks down into its component parts or simpler compounds.

Basic form: AX → A + X

Examples of decomposition reactions:

1. CaCO3(s) → CaO(s) + CO2(g)

2. Ca(OH)2(s) → CaO(s) + H2O(g)

3. 2KClO3(s) → 2KCl(s) + 3O2(g)

4. 2 NaHCO3(s) → Na2O(s) + 2 CO2(g) + H2O(g)

5. H2SO4 → H2O(l) + SO3(g)

3. Single Replacement Reactions: a more active element takes the place of another element in a compound and sets the less active

one free.

Basic forms:

o Single Replacement - Metal: A + BX → AX + B

o Single Replacement- Halogen: Y2 + AX → AY + X2

Examples of Single Replacement reactions:

1. Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

2. 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)

3. Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)

4. Cl2(g) + 2NaBr(aq) → 2NaCl(aq) + Br2(l)

NOTE: Refer to the activity series for metals and nonmetals to predict products of replacement

reactions. If the free element is above the element to be replaced in the compound, then the reaction will

occurr. If it is below, then no reaction occurs.

4. Double Replacement Reactions: occurs between ions in aqueous solution. A reaction will occur when a pair of ions come together

to produce at least one of the following:

1. a precipitate

2. a gas

3. water or some other non-ionized substance.

Basic form: AX + BY → AY + BX

Examples of Double Replacement Reactions:

1. Formation of precipitate. NaCl (aq) + AgNO3(aq) → NaNO3(aq) + AgCl(s)

2. Formation of a gas. HCl(aq) + FeS(s) → FeCl2(aq) + H2S(g)

3. Formation of water. (If the reaction is between an acid and a base it is called a neutralization

reaction.) HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

NOTE: Use the solubility rules to decide whether a product of an ionic reaction is insoluble in water and

will thus form a precipitate. If a compound is soluble in water then it should be shown as being in

aqueous solution, or left as separate ions. It is, in fact, often more desirable to show only those ions that

are actually taking part in the actual reaction. Equations of this type are called net ionic equations.

5. Combustion of Hydrocarbons:

Another important type of is the of the combustion of a hydrocarbon. When a hydrocarbon is burned with

sufficient oxygen supply, the products are always carbon dioxide and water vapor. If the supply of

oxygen is low or restricted, then carbon monoxide will be produced. This is why it is so dangerous to have

an automobile engine running inside a closed garage or to use a charcoal grill indoors.

Examples of Combustion Reactions:

1. Hydrocarbon (CxHy) + O2(g) → CO2(g) + H2O(g)

2. CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)

3. 2C4H10(g) + 13O2(g) → 8CO2(g) + 10H2O(g)

Practice Questions:

1. Identify the type of reaction, predict the products and balance:

a. _____C2H5OH + _____O2 →

b. _____Mg + _____HCl →

c. _____Sr(OH)2 + _____HBr →

d. _____SrCO3→

2. Identify the spectator ions in the following complete ionic equations:

a. Al3+ + Cl- + K+ + OH- → Al(OH)3 + K+ + Cl- spectator ions: ______________

b. Pb2+ + NO3- + Na+ + Cl- → PbCl2 + Na+ + Cl- spectator ions: ______________