AP CHEMISTRY NOTES 2-1

ASSIGNING OXIDATION NUMBERS

RULES FOR ASSIGNING OXIDATION NUMBERS:

1. The oxidation number of any free element (including diatomic elements) is always “0”.

2. The sum of the oxidation number of all of the atoms in a compound is always equal to “0”.

3. The sum of the oxidation numbers of all of the atoms in a polyatomic ion is always equal to the charge

of that ion.

4. The oxidation number of all elements in Group IA (except hydrogen) is always +1.

5. The oxidation number of all elements in Group IIA is always +2.

6. The following elements always have the oxidation numbers indicated:

Al = +3

Cd = +2

Zn = +2

Ag = +1

7. The oxidation number of hydrogen atoms in a compound is usually +1. (In hydrides, however, the

oxidation number of hydrogen is -1.)

8. The oxidation number of oxygen atoms in a compound is usually -2. (In peroxides, however, the

oxidation number of oxygen is -1. The peroxide ion is O22-.

9. In a binary ionic compound, the oxidation number of elements in Group VIIA is -1.

Examples:

NaI FeCl3

Br2 NaNO3

KClO4 ZnCr2O7

Na2O2 S8

Cr(MnO4)6 Pb(SO4)2

AP CHEMISTRY NOTES 2-2

WRITING NET IONIC EQUATIONS

A net ionic equation is one in which all species are shown in their ionized form (if appropriate – see below)

and spectator ions (those that appear unchanged on both sides of the equation) are left out.

RULES:

1. Strong acids and bases ionize (break apart to form ions; ie. “split”)

a. Strong acids include HCl hydrochloric acid

HI hydroiodic acid

HBr hydrobromic acid

H2SO4 sulfuric acid

HNO3 nitric acid

HClO4 perchloric acid

H2CrO4 chromic acid

HMnO4 permanganic acid

(Note: Concentrated sulfuric acid does not ionize)

b. Strong bases include all hydroxides of Groups IA and IIA metals except magnesium

2. Soluble ionic compounds ionize (know your solubility rules!!)

3. Molecular compounds, solids, oxides, and pure liquids do not ionize

4. All “spectator ions” are removed from the equation

EXAMPLES: Balance the following equations and write net ionic equations for each.

CaCO3 + HNO3 _______________> Ca(NO3)2 + CO2 + H2O

ICl + H2O _______________> HCl + HIO

H2C2O4 + NaOH _______________> Na2C2O4 + H2O

Fe + HCl _______________> FeCl2 + H2

Na2CO3 + HCl _______________> NaCl + CO2 + H2O

Ba(C2H3O2)2 + K2SO4 _______________> KC2H3O2 + BaSO4

AP CHEMISTRY NOTES 2-3

OXIDATION-REDUCTION REACTIONS

OXIDATION-REDUCTION REACTIONS – the changes that occur when electrons are transferred between

reactants (also known as a redox reaction)

OIL RIG

Oxidation Is Loss of electrons . . . . . . . Reduction Is Gain of electrons

Oxidation and reduction ALWAYS occur simultaneously

Oxidation – the loss of electrons by a substance (the oxidation number rises, becoming more positive)

Reduction – the gain of electrons by a substance (the oxidation number lowers, becoming more negative)

Oxidizing Agent – the substance in a redox reaction which causes another substance to be oxidized

(it itself is reduced)

Reducing Agent – the substance in a redox reaction which causes another substance to be reduced

(it itself is oxidized)

EXAMPLES:

HCl + NaOH _______________> NaCl + H2O

Mg + S _____________> MgS

As + H2 _______________> AsH3

K2Cr2O7 + H2O + S _______________> KOH + Cr2O3 + SO3

AP NOTES 2-4

WRITING REDOX HALF-REACTIONS

RULES: 1. Re-write the reactions as net ionic equations. 2. Find the substance oxidized and substance reduced. 3. Write the oxidation half-reaction and the reduction half-reaction and add electrons to balance. EXAMPLES:

Cu + AgNO3 _______________> Ag + Cu(NO3)2

HNO3 + HI _______________> NO + I2 + H2O

Zn + MnO2 + NH4Cl __________> ZnCl2 + Mn2O3 + NH3 + H2O

K2Cr2O7 + H2O + S _______________> KOH + Cr2O3 + SO3

AP NOTES 2-5

BALANCING REDOX REACTIONS – Acidic Solution

RULES:

1. Write the net ionic equation for the reaction.

2. Write separate half-reactions for the oxidation and reduction processes (balance for substance

changing charge).

3. Balance the half-reaction “by mass”:

a. Use H2O to balance oxygens where needed

b. Add H+ to balance hydrogens on the other side

4. Balance both reactions “by charge” (multiply each half-reaction by an appropriate number to

equalize the number of electrons lost or gained).

5. Add the half-reactions and subtract terms that appear on both sides of the equation.

6. Check the final equation to be sure that atoms are conserved, charge is conserved, and all

electrons have cancelled.

EXAMPLES:

Al(NO3)3 + Fe _______________> Fe(NO3)2 + Al

Co + Br2 _______________> CoBr3

HNO3 + HI _______________> NO + I2 + H2O (Acidic Solution)

S + HNO3 _______________> SO2 + NO + H2O (Acidic Solution)

AP CHEMISTRY NOTES 2-6

BALANCING REDOX REACTIONS – Basic Solution

(UNIT 2)

RULES:

1. Write the net ionic equation for the reaction.

2. Write separate half-reactions for the oxidation and reduction processes (balance for substance

changing charge).

3. Balance the half-reaction “by mass”:

a. Use H2O to balance oxygens where needed

b. Add H+ to balance hydrogens on the other side

c. Add the same number of OH- ions to both sides of the equation to equal the H+ ions

added previously

d. Convert the H+ / OH- ions to water on one side of the equation and leave the lone

OH- ions on the other side alone

4. Continue as for reactions in acidic solution.

EXAMPLES:

Ag + CN- + O2 _______________> Ag(CN)2- + H2O

Cr + CrO42- _______________> Cr(OH)3

AP Chemistry NOTES 2-7

STATISTICAL ANALYSIS:

PERCENT ERROR & STANDARD DEVIATION

PERCENT ERROR

Accuracy – how close a measurement or calculation is to the accepted value

One statistical method used to report / measure accuracy is percent error:

Example: In an experiment, the diameter of a molecule was determined to be 0.267 nm. The

accepted value was found to be 0.290 nm. What was the percent error?

STANDARD DEVIATION

Precision – how close together the values for multiple measurements of the same quantity are

One statistical method used to report / measure precision is standard deviation:

Where Σ = “the sum of”

_

d1 = xi - x

x1 = each measurement value

_

x = average (mean) of values

N = the number of measurements

Example: The following values were determined for the volume of a tin cylinder (in cm3). Determine

the standard deviation of the values.

Value of Measurement

xi

Deviation

_

di = (xi – x)

Squared Deviation

di2

10.11

10.13

10.10

10.12

10.15

10.11

10.12

mean = __________ sum = _____________

Example: The mass of an object was measured six times and the following values were obtained:

13.34 g, 13.08 g, 13.58 g, 13.42 g, 13.29 g, and 13.45 g. Determine the standard deviatio