Topic 3

Compounds:

• a compound is a substance made up of two or more difference chemical elements combined in a fixed ratio

Definition of chemical bonds:

• a strong attractive force that exists between atoms in a molecule

• valence electrons play a fundamental role in chemical bonding

Structural vs molecular vs empirical formula:

• structural - CH3COOH

• molecular - C2H4O2

• empirical - CH20

Structural formula:

• is a formula which shows the arrangement of atoms in the molecule of a compound

Molecular formula:

• is a formula which shows the exact number and type of atoms present in a molecule

Empirical formula:

• is a formula which shows the simplest, whole-number ratio of atoms of each element in a molecule

Representing molecular structure:

• in a chemical formula, a subscript (e.g. the ‘2’ in H2) indicates the number of atoms of an element within a single molecule, while a coefficient (e.g. the ‘2’ in 2H2) represents the total number of molecules of that formula unit involved in a chemical reaction

Stoichiometry:

• stoichiometry is a section of chemistry that involves using relationships between reactants and/or products in a chemical reaction to determine desired quantitative data

The law of conservation of mass (Antoine Lavoisier 1743 - 1794):

• the law of conservation of mass states that the total mass present before a chemical reaction is the same as the total mass present after the chemical reaction

Balancing reactions:

• equations are balanced by including stoichiometric coefficients so that the number of atoms of each type is the same on each side of the equation

• H2(g) + Cl2(g) → HCl(g)

Predicting reactions using equations:

• knowing formulas and stoichiometry allows you to predict things

Types of chemical reactions:

1. addition reactions

2. elimination reactions

3. single replacement reactions

4. double replacement reactions

5. combustion reactions

6. acid-base reactions

Addition/combination reactions:

• a combination reaction is a reaction in which 2 or more simple reactants to form a complex product

• X + Y → XY

• 2Na + Cl2 → 2NaCl

Examples of addition/combination reactions:

• hydrogenation

• halogenation

• hydration

Elimination/decomposition reactions:

• a decomposition reaction is a reaction in which a compound breaks down into two or more simpler substances

• XY → X + Y

• 4HNO3 → 4NO2 + 2H2O + O2

• 2CuO → 2Cu +O2

Examples of elimination/decomposition reactions:

• dehydrogenation

• dihydroxylation

• dehalogenation

Single substitution reactions:

• single displacement is when an element from on reactant switches with an element of the other to form two new reactants

• X + YZ → Y + XZ

• Fe + CuSO4 → Cu + FeSO4

• F2 + 2NaCl → 2NaF

Double substitution reactions:

• double displacement is when two elements from on reactants switched with two elements of the other to form two new reactants

• AX + BY → AY + BX

• AgNO3 + NaCl → AgCl + NaNO3

• 2Kl + Pb(NO3)2 → 2KNO3 + Pbl2

Combustion reactions:

• combustion is the formation of CO2 and H2O from the reaction of a chemical and O2

• C3H8 + 5O2 → 3CO2 + 4H2O

• C2H5OH + 3O2 → 2CO2 + 3H2O

• 2Mg + O2 → 2MgO

Acid-base reactions:

• acid-base reactions are when two reactants form salts and water

• HCl + NaOH → H2O + NaCl

• HCl + CaCO3 → CaCl2 + H2O + CO2

• H2SO4 + CaCO3 → CaSO4 + H2O + CO2

Reactions at equilibrium:

• a stage in a chemical reaction in which the rate at which the reactants are converted to products is equal to the rate at which products are converted back to reactants

• H2 + N2 $\lrArr$ 2NH3

• 2NO2 $\lrArr$ N2O4

• HCl + H2O $\lrArr$ H3O+ + Cl-

Le Chatelier’s Principle:

• when a stress is applied to a system at equilibrium, the equilibrium will shift to minimise the effects of the stress

• A + 2B $\lrArr$ C + D

What would happen if you changed the conditions by increasing the concentration of A?

• the position of equilibrium moves to the right

What would happen if you changed the conditions by decreasing the concentration of A?

• the position of equilibrium moves to the left

What would happen if you changed the conditions by increasing the pressure?

• the position of equilibrium moves right

What would happen if you changed the conditions by decreasing the pressure?

• the position of equilibrium moves to the left

What would happen if you changed the conditions by decreasing the pressure?

• the position of equilibrium is not affected

The first law of thermodynamics:

• the first law of thermodynamics is a formation of the law of conservation of energy in the context of thermodynamic pressure

• system + surroundings = universe

Exothermic reactions:

• an exothermic reaction is a reaction that releases energy to the surroundings

• the energy is released as heat energy, so the surroundings warm up

Endothermic reactions:

• an endothermic reaction is a reaction that absorbs energy from the surroundings

• the energy is absorbed as heat energy, so the surroundings cool down

What would happen if you changed the conditions by increasing the temperature?

• the position of equilibrium moves to the left

What would happen if you changed the conditions by increasing the temperature?

• the position of equilibrium moves to the right

Acidification of sea water:

• CO2(g)$\lrArr$ CO2(aq)

• CO2(aq) + H2O (l) $\lrArr$ H2CO3(aq)

• H2CO3(aq) $\lrArr$ H+(aq) + HCO3-(aq)

• HCO3-(aq) $\lrArr$ H=(aq) + CO3 2-(aq)

The Haber Process:

• high pressure will favour the forward reaction, however high pressure is expensive

• reaction pressure is usually 200 atm

• low temperature will favour the forward reaction, however low temperature means a slow reaction

• reaction temperature is usually 500 degrees

The contact process:

• a three-step process used in the manufacture of sulphuric acid

1. make sulphur dioxide

2. convert sulphur dioxide into sulphur trioxide

3. convert sulphur trioxide into sulphuric acid

Key takeaways:

• a chemical formula represents the composition of a compound using symbols and subscripts to show the elements and their ratios

• subscripts indicate the number of atoms of each element in one molecule or formula unit (e.g. H2O has 2 hydrogen atoms and 1 oxygen atom)

• chemical equations symbolically represent reactions using formulas, showing the reactants, products and their stoichiometric ratios

• the law of conservation of mass states that matter is neither created nor destroyed in a chemical reaction - the total mass of reactants equals the total mass of products

• balancing equations ensures the same number of each type of atom appears on both sides of the equation

Common types of reactions include:

• addition

• elimination

• displacement (single and double)

• combustion

• neutralisation

Dynamic equilibrium:

• a stage where the rate of the forward reaction equals the rate of the reverse reaction

• denoted with a double arrow $\lrArr$

Le Chateliers Principle:

• when a stress is applied to a system at equilibrium, the position of equilibrium shifts to minimise that stress