Chemical Reactivity | General Chemistry 2

Precipitation: formation of a solid, called precipitate.
A solid is formed if ions of an insoluble salt are present.

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

Types of Equations:

- Molecular equation: shows the formulas of the compounds.
     AgNO₃ (aq) + NaCl (aq) → AgCl (s) + NaNO₃ (aq)

- Ionic equation: shows the ions of the compounds.
     Ag⁺ (aq) + NO₃^- (aq) + Na⁺ (aq) + Cl^- (aq) → AgCl (s) + Na⁺ (aq) + NO₃^- (aq)

- Net ionic equation: shows only the ions that form a solid.
     Ag⁺ (aq) + Cl^- (aq) → AgCl (s)

Arrhenius definitions of acids and bases:

- Acids: substances that have H in their formulas and that dissociate in water to give H⁺ ions.
- Bases: substances that have OH in their formulas and that dissociate in water to give HO^- ions.

HCl is an Arrhenius acid: HCl (aq) → H⁺ (aq) + Cl^- (aq)
NaOH is an Arrhenius base: NaOH (aq) → Na⁺ (aq) + HO^- (aq)

Brönsted acid-base theory:

- Acids: substances that can donate a proton (proton donor)
- Bases: substances that can accept a proton (proton acceptor)
The acid molecules HA gives an H+ to the base molecule B: HA + B → A^- + HB⁺

Conjugate acid-base pair (HX / X^-): acid-base pair that differ only in the presence or absence of a proton

Ex: NH₃ (aq) + H₂SO₄ (aq) → NH₄⁺ (aq) + HSO₄^- (aq)
NH₃ is a base, H₂SO₄ is an acid.
NH₄⁺ / NH₃ and H₂SO₄ / HSO₄^- are two conjugate acid-base pairs

 Lewis acid-base theory:

- Acids: electron-pair acceptors ⇒ must be electron deficient or have an orbital that can accept electrons: BF₃, AlCl₃, SO₂
- Bases: electron-pair donors ⇒ must have a lone pair of electrons: NH₃
When a Lewis base gives electrons from a lone pair to a Lewis acid, a covalent bond forms between the molecules and the product is called an adduct.

NH₃ + BF₃ → NH₃BF₃
NH₃BF₃ is an adduct

Oxidation number:

- for an ion: charge of the ion
- for an atom: hypothetical charge if it was an ion
⇒ group IA: + 1; group IIA: + 2; oxygen: - 2; halogens: - 1; H with a nonmetal: + 1; H with a metal: - 1

Oxidation-reduction reactions (Redox reactions):
reactions in which the oxidation number of an atom, molecule or ion changes.
- Oxidation: particle becomes more positively charged (loss of electrons) ⇒ oxidation number increases
- Reduction: particle becomes less positively charged (gain of electrons) ⇒ oxidation number decreases

Oxidizing agent: substance that is reduced in a redox reaction
Reducing agent: substance that is oxidized in a redox reaction

Balancing Redox Reactions:

1) Divide the reaction into half-reactions
2) Balance atoms (other than O and H)
3) Balance O atoms by adding H₂O molecules
4) Balance H atoms by adding H⁺ ions
5) Balance charge by adding electrons
6) Multiply each half-reaction by an integer to have: 
number of electrons lost = number of electrons gained
7) Add the half-reactions together

Under basic conditions:
8) Add HO^- ions to react with all H⁺ ions: HO^- + H⁺ → H₂O

Combination reactions:
2 or more elements (or simple compounds) combine to form 1 product.
 

2 Na (s) + Cl₂ (g) → 2 NaCl (s)
CO₂ (g) + H₂O (l) → H₂CO₃ (aq)

Decomposition reaction:
one substance splits into 2 or more simpler substances.
 

CaCO₃ (s) → CaO (s) + CO₂ (g)
MgSO₄ (s) → MgO (s) + SO₃ (g)

Single replacement:
1 element takes the place of a different element in a reacting compound.
An element will displace another element in a compound if it is more active than the element in the compound.
 

Br₂ (l) + CaI₂ (aq) → CaBr₂ (aq) + I₂ (s) 
[activity of halogen gases decreases down the column]

Fe (s) + H₂SO₄ (aq) → FeSO₄ (aq) + H₂ (g)

Double replacement:
2 elements in the reactants exchange place.

NaCl (aq) + AgNO₃ (aq) → NaNO₃ (aq) + AgCl (s)
BaCl₂ (aq) + Na₂SO₄ (aq) → 2 NaCl (aq) + BaSO₄ (aq)

Combustion reaction:
a compound (generally an alkane) reacts with O₂ and produces CO₂, H₂O and energy (heat and light).
 

CH₄ (g) + 2 O₂ (g) → CO₂ (g) + 2 H₂O (g)
2 C₂H₆ (g) + 7 O₂ (g) →4 CO₂ (g) + 6 H₂O (g)