Chemical Reactivity | General Chemistry 2

Chemical reaction: process in which one or more substances (the reactants) are converted to one or more different substaces (the products)

A chemical reaction is represented by a chemical equation:

• chemical formulas of the reactants on the left-hand side 
• chemical formulas of the products on the right-hand side
• reactants and products are separated by an arrow
• each individual substance’s chemical formula is separated from others by a '+' sign.

The states of the substance (solid (s), liquid (l), gas (g) or aqueous (aq)) can be added after each individual substance’s chemical formula.

Mg (s) + Cl₂ (g) → MgCl₂ (s)

Atoms are neither created nor destroyed in a chemical reaction
⇒ chemical equation must be balanced (same number of each type of atom on both sides)
 

Li + Br₂ → LiBr    is not balanced
2 Li + Br₂  → 2 LiBr      is balanced

Easy steps for balancing chemical equations:

1. write the unbalanced chemical equation

2. determine how many atoms of each element are present on each side of the arrow

3. balance atoms present in a single molecule of reactant and product first

4. balance any oxygen or hydrogen atoms last

5. at the end, check your equation to make sure you cannot reduce the balancing coefficients

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 equations: show the formulas of the compounds

     AgNO₃ (aq) + NaCl (aq) → AgCl (s) + NaNO₃ (aq)

• Ionic equations: show the ions of the compounds

     Ag⁺ (aq) + NO₃^- (aq) + Na⁺ (aq) + Cl^- (aq) → AgCl (s) + Na⁺ (aq) + NO₃^- (aq)

• Net ionic equations: show only the ions that form a solid

     Ag⁺ (aq) + Cl^- (aq) → AgCl (s)

Arrhenius acid-base theory:

• 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 donates electrons from its 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 compounds) combine to form 1 product
 

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

Decomposition reaction:
1 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)