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4.4: Oxidation / Reduction Reactions

REDOX REACTIONS

OXIDATION:

Increase in Oxygen numbers

Reduce number of Hydrogens

Increase Oxidation numbers

Remove electrons

REDUCTION:

Decrease Oxygen numbers

Increase Hydrogen numbers

Decrease Oxidation numbers

Add electrons

ASSIGNING OXIDATION NUMBERS

1. Any substance in elemental form = 0
Example:
O2, Na, S8, P4

2. Monoatomic ion oxydation number equal to its charge
Example:

F⁻' = - 1

3. Oxygen and Hydrogen's oxydation numbers are as follows:

O = -2

H = +1 (when it's paired with a non-metal)

H = -1 (when it's paired with a metal)

4. Sum of oxydation numbers:

Compound sum = 0

Polyatomic ion sum = charge of ion

4.3: Neutralization Reactions

NEUTRALIZATION REACTIONS

ACID (H+) + BASE (OH-) -------> SALT (Ionic compound) + WATER

STRONG ACIDS: Completely ionize, that is they are soluble in water.

HCl, HBr, HI, HClO3, HClO4, HNO3, H2SO4

All other acids are weak and therefore they stay together in H2O (do not separate as ions)

Example:

MOLECULAR EQUATION:

$\large \fn_cm 2CH_{3}COOH\;\;+\;\;Ba(OH)_{2}\;\;\rightarrow \;\;(CH_{3}COO)_{2}Ba\;\;+\;\;2H_{2}O$

IONIC EQUATION:

$\dpi{120} \fn_cm 2CH_{3}COOH\;+\;\underline{Ba^{2+}}\;+\;2OH^{-}\;\rightarrow \;(2CH_{3}COO)^{-}\;+\;\underline{Ba^{2+}}\;+\;2H_{2}O$

NET IONIC EQUATION:

$\dpi{120} \fn_cm 2CH_{3}COOH\;\;+\;\;2OH^{-}\;\rightarrow \;(2CH_{3}COO)^{-}\;\;+\;\;2H_{2}O$

20.1: Electrochemistry

ELECTROCHEMISTRY

Oxidation Numbers: Assign and determine if a reaction is going under oxidation or reduction

a) Oxidation / Reducing Agent: Loses electrons

b) Reduction / Oxidizing Agent: Gains electrons

Rules:

1. Pure, most stable form = 0

2. Ion in solution (Pb+2) = Charge of the ion

3. For the most part: Oxygen = -2, Fluorine = -1, Hydrogen = +1

Example:

$\large {\color{Blue}\overset{0}{Cd(s)} }\: +\: {\color{Red}\overset{+4}{NiO_{2}} }\: +\: 2H_{2}O\: \rightarrow \: {\color{Blue}\overset{+2}{Cd(OH)_{2}} }\: +\: {\color{Red}\overset{+2}{Ni(OH)_{2}} }$