Unit 1 Energy Matters

1. Unit 1 Energy Matters

2. Reaction rates 1. When following the course of a reaction rate = change time • When timing a reaction (how long before colour change): • rate = 1/time (s-1) time = 1/rate

3. Reaction rates Increased by concentration or temperature particle size catalyst

4. 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.

5. Excess reactant e.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. • Write balanced equation • CaCO3 + 2HCl CaCl2 + H2O + CO2 • 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 • 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 out • 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 • Use m = n x gfm to calculate mass of product • e.g. 0.1 x 44 = 4.4 g CO2

6. Enthalpy Exothermic reactions release energy to surroundings, which heat up e.g. reaction mixture gets hotter. Endothermic reactions take in energy from surroundings, which cool down

7. 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 water e.g. NaCl (s) Na+ (aq) + Cl – (aq) Enthalpy of neutralisation: Energy change when one mole of water forms from the neutralisation of an acid H+ (aq) + OH- (aq) H2O (l)

8. Calculating enthalpy changes Eh = cmDT – 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-1oC-1) DT = temperature change NB neutralisations: add volume of acid to alkali; average starting temperature DH = Eh /n - to do with stuff burned/dissolve/neutralised DH = enthalpy change (kJ mol-1) n = moles of stuff dissolved/burned/acid neutralised Insert negative if exothermic

9. Patterns in the Periodic Table Use 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.

10. Bonding, structure and properties of elements

11. 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 forcesbetween molecules Monatomic elements – weak van der Waals forces between atoms

13. Properties Metals 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.

14. 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!!

15. Polar molecules

16. Polar molecules

17. 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

18. Properties Ionic 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.

19. The mole For solutions n = c V For solids m = n x gfm For gases V = n x Vmol

20. The mole • Examples • Which has the greatest number of atoms, 3 g of ethane or • 1.6 g of methane? • Calculate the molar volume of oxygen from the following data. • Mass of empty 1 litre flask = 205.42 g • Mass 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)