Chemical Reactivity | General Chemistry 2

Chemical reactivity is studied in this chapter: precipitation, acid-base reactions, oxidation-reduction reactions, combination and decomposition, replacement reactions, combustion

Chemical Reactions

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)

Balancing Chemical Equations

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

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)

Acid - Base

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 - Reduction Reactions

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 et Decomposition Reactions

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)

Replacement Reactions

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
 

Br(l) + CaI(aq) → CaBr(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 Reactions

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)