AQA Chemistry C3 Part 2

Exothermic - gives out energy, less energy required to break bonds than released when forming them

Endothermic - takes in energy, more energy required to break bonds than is released when forming them

Energy is needed to break bonds & Energy is released when bonds are formed

You can measure energy transfer by taking the temp of the reactants before and the products after

BUT lose energy to surroundings

Fuels release energy - exothermic

Calorimetry - measures amount of energy produced when a fuel's burnt

Working out energy per gram of fuel:

• Weigh the mass of spirit burner + lid before and after & find the difference
• Measure the temp of the water before and after heating & find the difference
• Takes 4.2J of energy to heat 1g of water by 1`C
• 4.2 x amount of water in grams x difference of temp = y
• y/difference of mass = energy per gram of fuel

Exothermic reactions - change in enthalpy(heat) is -ve

Endothermic reactions - change in enthalpy is +ve

Activation energy - min. energy needed to break the bonds & start the reaction

Catalysts lower activation energy BUT don't effect the overall energy change

Bond energy calculations, e.g. H2 + Cl2 --> 2HCl, H-H: +436kJ/mol, Cl-Cl: +242,       H-Cl: +431

• 436 + 242 = 678kJ/mol needed to break the bonds
• 431 x 2 = 862kJ/mol released when forming the bonds
• 862 - 678 = 184kJ/mol: energy change
• Because it's exothermic change in enthalpy is -184kJ/mol

1 calorie = the amount of energy needed to raise 1g of water by 1`C = 4.2joules

1 Calorie = the amount of energy needed to raise 1kg of water by 1`C = 4200joules

• Fats and oils - high amount of energy
• Carbohydrates - some energy but less than fats and oils
• Proteins - same as carbohydrates but used for different things in body

Chemical reactions in your cells going on all the time need energy

BUT taking in more energy than you're using/ your body needs leads to obesity

Flame tests:

• Li+ - red
• Na+ - yellowy/orange
• K+ - lilac
• Ca2+ - brick red
• Ba2+ - green

Add NaOH:

• Ca2+ & Mg2+ - white precipitate
• Al2+ - white precipitate then dissolves in excess
• Fe2+ - sludgy green precipitate
• Fe3+ - red/brown precipitate
• Cu2+ - blue precipitate

Ammonia - smells & turns red litmus paper blue

Carbonates + acid:

• CO2 - turns limewater cloudy
• Copper carbonate - green > black
• Zinc carbonate - white > yellow > white (when cooled)

Sulfates - add barium chloride & HCl, if a white precipitate then it's a sulfate

Halides - add nitric acid & silver nitrate:

• Chloride - white precipitate
• Bromide - cream precipitate
• Iodide - yellow precipitate

Nitrates - add aluminium powder & NaOH, if ammonia then it's a nitrate 

Burn with a yellowy/orange and/or blue flame & solid organic compounds will char

Unsaturated (C=C bonds) - decolourise bromine water

Emperical formula - e.g. 0.4g of hydrocarbon burnt. 1.1g of CO2 & 0.9g of H2O:

• 1) Mass of CO2 - 1.1 x (12/44) = 0.3
• Mass of H2O - 0.9 x (2/18) = 0.1
• 2) Divide by relative atomic mass - 0.3/12 = 0.025moles
• Divide by relative atomic mass - 0.1/1 = 0.1moles
• 3) Divide both by smallest one - 0.025/0.025 = 1
• Divide both by smallest one - 0.1/0.025 = 4
• Ratio - 1:4
• CH4

Advantages - by technicians not chemists, more accurate, faster

Disadvantages - expensive

Types:

• Atomic Absorption Spectroscopy - identifies metals, like flame test, analyses patterns of light absorbed by metal
• Infrared/UV Spectroscopy - identifies compounds, fingerprint of frequencies of infra-red/UV absorbed
• NMR Spectroscopy - for organic compounds, shows what the hydrogen atoms are connected to
• Gas-Liquid Chromatography - for gases & liquids, similar to paper chromatography
• Mass Spectroscopy - for compounds & elements, finds mass - work out what's in there