XV. Common types of reactions

A. Acids and bases – behavioral patterns, not an absolute (H₂O goes both ways; amphoteric)

high chemical activity makes them “corrosive”

1. acids

sour, tangy taste

give up a H⁺ hydrogen ion (proton)

in aqueous solutions it forms hydronium ion (H₃O⁺)

donates a proton

strong acids break down quickly in water donating their H⁺ (residual anions make good electrical conductors)

2. bases

bitter taste and greasy feel

alkaline

absorb a hydrogen ion (proton)

in aqueous solution it usually forms hydroxide ion (OH^-)

strong bases react quickly in water accepting their H⁺ (residual hydroxide ions OH^- make good electrical conductors)

3. pH scale (-1 to 14)

describes acidity of a solution (how readily it donates a H⁺)

pH = -log [H₃O⁺]

related to concentration of hydronium ions

high concentration = high acidity = low number

pH = 7 = 10^-7M solution = 0.0000001

pH = 6 = 10^-6M solution = 0.000001

pH = 5 = 10^-5M solution = 0.00001

4. reactions form salts – (table 25.1)

ionic compound formed from the reaction between an acid and a base.

cation from base and anion from acid

acid and base neutralize each other so salts not very corrosive

2HCl + Ca(OH)₂ = CaCl₂ + 2H₂0

HF + NaOH = NaF + H₂O

H₂CO₃ + Ca(OH)₂ = CaCO₃ + 2H₂O

B. Oxidation and Reduction

1. processes –
2. in its simplest form it is the transfer of electrons from metallic element to non-metallic (closely related to formation of ionic bonds)
• oxidation – reactant loses one or more electrons to form cation
• reduction – reactant gains one or more electrons to form anion

3. always must occur in pairs. Reactants don’t spontaneously lose or gain electrons must have a source or destination.
• oxidizing agent causes oxidation so it is reduced
• reducing agent causes reduction so it is oxidized

4. half reactions – since each reactant behaves differently, we must show each half of the reaction separately:

oxidation

Mg = Mg²⁺ +2e^-

Li = Li⁺ +e^-

reduction

N + 3e^- = N^3-

Cl + e^- = Cl^-

Al = Al ³⁺ + 3e^- (oxidation)

F + e^- = F^- (reduction)

whole reaction = Al + 3F = AlF₃

must determine the oxidation state of each element first:

5. What about transition elements that have weird valence shell activity?

can have multiple oxidation states depending on what it is paired with

must know the final molecular formula to know what its oxidation state is

6. applications

· photography

· batteries

· combustion

· antioxidants/free radicals

C. Reaction Rates:

molecules must collide (often break bonds)

1. controls:

concentration

orientation of molecules

kinetic energy of molecules (temperature)

2. Activation energy

energy required to break and reform bonds

exothermic – net release H₂ + O = H₂0 + energy (often spontaneous)

endothermic – net absorption N₂ + O₂ + energy = 2NO (not spontaneous, but more easily reversed)

catalyst – provides an alternative reaction pathway involving intermediate reactions, each having a lower activation energy than the uncatalyzed reaction

2O₃ to 3O₂ much faster if Cl is present (breakdown ozone layer)

Cl + O₃ à ClO + O₂

ClO + O₃ à Cl + 2O₂