Unit 1 Energy Matters

Reaction rates 1. When following the course of a reaction

rate = change

time

2. When timing a reaction (how long before colour change):

rate = 1/time (s-1) time = 1/rate

Reaction rates

Increased by

concentration or temperature

particle size

catalyst

Catalysts

• Homogeneous

same state as reactants e.g. enzymes• Heterogeneous

different state from reactants e.g. in catalytic converters; iron in Haber process.

Catalysts can be poisoned.

Excess reactante.g. 15 g of calcium carbonate were reacted with 50 cm3 of 4 mol l-1 hydrochloric acid. Calculate the mass of carbon dioxide produced.

i. Write balanced equation CaCO3 + 2HCl CaCl2 + H2O + CO2

ii. Calculate moles of reactants usually n = m/gfm for one and n = cv for other e.g. n CaCO3

= 15/100 = 0.15; nHCl = 4 x 0.05 = 0.2

iii. Compare moles of reactants to equation e.g. need 1 CaCO3 :2 HCl; have 0.15:0.2 so excess CaCO3;

reaction stops when HCl runs outiv. Calculate moles of product from moles of chemical used up

e.g. 2 moles HCl: 1 mole CO2 so 0.2 gives 0.1 mole CO2

v. Use m = n x gfm to calculate mass of product e.g. 0.1 x 44 = 4.4 g CO2

EnthalpyExothermic reactions release energy to surroundings, which heat up e.g. reaction mixture gets hotter.

Endothermic reactions take in energy from surroundings, which cool down

Enthalpy definitions

Enthalpy of combustion:

Energy change when one mole of a substance burns completely

e.g. CH4 (g) + 2 O2 (g) CO2 (g) + 2 H2O (l)

Enthalpy of solution:Energy change when one mole of substance dissolves completely in watere.g. NaCl (s) Na+ (aq) + Cl – (aq)

Enthalpy of neutralisation:Energy change when one mole of water forms from the neutralisation of an acidH+ (aq) + OH- (aq) H2O (l)

Calculating enthalpy changesEh = cmT – all to do with water!

Eh = heat energy (kJ)

m = mass of water (Kg; 1 litre = 1 Kg; 100 ml = 0.1 Kg)

c = specific heat capacity of water (4.18 kJ kg-1 oC-1)

T = temperature change

NB neutralisations: add volume of acid to alkali; average starting temperature

H = Eh /n - to do with stuff burned/dissolve/neutralised

H = enthalpy change (kJ mol-1)

n = moles of stuff dissolved/burned/acid neutralised

Insert negative if exothermic

Patterns in the Periodic TableUse data book to describe patterns in, e.g., atomic

size (covalent radius), boiling points, etc.

Atomic size, first ionisation energy, electronegativity

Trends explained by

• increasing nuclear charge (protons) from left to right in period.

• increasing electron shells (and shielding of outer electrons) going down a group.

No values for Noble gases for atomic size or electronegativity

since they don’t bond..

Melting and boiling points and density

Trends explained by differences in bonding.

Bonding, structure and properties of elements

Metallic bonding – attraction between positive ions and outer electrons

Covalent networks – strong covalent bonds (shared pairs of e-) between atoms

Covalent molecules – strong covalent bonds within molecules; weak van der Waals forces between molecules

Monatomic elements – weak van der Waals forces between atoms

PropertiesMetals and graphite conduct electricity.

Covalent networks have high melting and boiling points as change of state involves breaking covalent bonds.

Covalent molecules have lower melting and boiling points as change of state involves breaking weak van der Waals forces.

Metals have reasonably high melting and boiling points; increases as strength of metallic bond increases.

Bonding, structure and properties of compounds

Intramolecular bonding determined by

difference in electronegativity of 2 atoms • Ionic – big difference

• Pure covalent – no difference

• Polar covalent – in between difference!!

Polar molecules

Polar molecules

Intermolecular bonding between molecules• Van der waals – temporary dipoles; weak

• Permanent dipole – permanent dipole interactions; stronger

• Hydrogen bonding – when H attached to O, N, F; strongest

PropertiesIonic compounds conduct electricity in solution and as melts.

High boiling and melting points indicate breaking of strong bonds when compound changes state; either ionic or covalent network.

Low melting point and boiling points indicate molecular covalent bonding, with weak van der waals interactions between molecules.

Polar interactions and hydrogen bonding elevate boiling points of molecules.

Hydrogen bonding increases viscosity; ice less dense than water.

Polar liquids deflected by a charged rod.

Like dissolves like: polar & ionic substances dissolve in water; non-polar substances dissolve in nonpolar solvents.

The mole

For solutionsn = c V

For solidsm = n x gfm

For gasesV = n x Vmol

The moleExamples1. Which has the greatest number of atoms, 3 g of ethane or

1.6 g of methane?

2. Calculate the molar volume of oxygen from the following data. Mass of empty 1 litre flask = 205.42 gMass of flask + oxygen = 206.77 g

3. Calculate the volume of hydrogen gas released when 2 g of calcium reacts with excess hydrochloric acid. Take the molar volume of hydrogen to be 24 litres.

4. 100 cm3 of propane gas was ignited with 750 cm3 of oxygen. What was the composition of the resultant mixture?

Ethane, with 0.8 moles of atoms, methane only has 0.5 moles

Since 1.35 g has a volume of 1 litre, the molar volume of oxygen is 23.7 litres (32/1.35 x 1 litre)

40 g calcium gives 24 litres so 2 g gives 1.2 litres

250 cm3 oxygen, 300 cm3 carbon dioxide as C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l)