chemistry, Energetics

• Exothermic reaction


  Chemical reaction which produces heat energy, causing the temperature of the surroundings to increase
• Endothermic reaction


  Chemical reaction which takes in heat energy, causing the temperature of the surroundings to decrease
• No chemical reaction can produce energy from nowhere. Energy is always conserved - it cannot be created or destroyed, but only changed from one form to another.
• Enthalpy change (ΔH)


  The change in heat energy at constant pressure
• Exothermic reactions have negative enthalpy changes (ΔH<0)
• Endothermic reactions have positive enthalpy changes (ΔH>0)
• Measuring energy changes


  1. Energy supplied to water (q = m c ΔT)

  2. Calculating enthalpy change (ΔH = - q / n)
• Energy change of combustion

  The energy change when one mole of a particular fuel is burned fully in oxygen, under standard conditions
• Energy change of neutralisation


  The energy change when one mole of an acid is fully neutralised in solution by an alkali
• Energy change of solution


  The energy change when one mole of a particular substance is dissolved in water (to give a solution of concentration 1 mol/dm3)
• Fossil fuels


  • Coal
  • Crude oil
  • Natural gas
• Fossil fuels were formed over millions of years by the action of high temperature and extreme pressure, deep underground, on animal and plant remains
• Fossil fuels are non-renewable energy sources, since no significant deposits have been formed for millions of years
• The burning of any fossil fuel has considerable environmental impact, because of the pollutants produced
• Incomplete combustion of hydrocarbons

  1. Forming carbon monoxide

  2. Forming carbon (soot)
• Carbon monoxide is a very poisonous pollutant that attaches strongly to haemoglobin in the body's red blood cells, preventing them from carrying oxygen
• Smaller hydrocarbon molecules with shorter carbon chains are less prone to incomplete combustion
• Sulfur dioxide dissolves in water vapour to make sulfurous acid, which may later oxidise in the air to form sulfuric acid
• Carbon monoxide


  Colourless and odourless gas, very dangerous as it is difficult to detect. Many people die each year from carbon monoxide poisoning, usually from running car engines in enclosed spaces with poor ventilation, or from badly maintained domestic heaters, with inadequate air supply.
• Removing carbon monoxide from car engines


  Oxidation to carbon dioxide in the catalytic converter
• Hydrocarbon molecules


  Smaller molecules with shorter carbon chains, such as methane, are less prone to incomplete combustion, as less oxygen is required per mole of fuel for complete combustion. Longer hydrocarbon fuels, such as paraffin, require a greater oxygen supply, so are likely to burn incompletely, giving a sooty, yellow flame.
• Sulfur dioxide


  Dissolves in water vapour to make sulfurous (H2SO3) acid. This may later oxidise in the air to form sulfuric (H2SO4) acid. These fall as acid rain, damaging trees, crops and stone buildings, and killing plants and animals in lakes and rivers. It forms as a result of burning fuels containing sulfur compound impurities, although it can be removed by bubbling the waste gases through an alkaline solution.
• Nitrogen oxides (NO, NO2)


  Form when nitrogen and oxygen in the air react together at the very high temperatures created in a car engine. They dissolve in water vapour to form nitrous (HNO2) and nitric (HNO3) acids, again contributing to acid rain. Nitrogen oxides, along with carbon monoxide, can largely be removed by a car's catalytic converter, turning them to nitrogen and carbon dioxide.
• Unburned hydrocarbons


  Released by car engines when they are unable to burn completely the fuel supplied to them. They can irritate the lungs, and can cause harmful smogs to form, especially on still, sunny days. They can be removed by oxidation in a car's catalytic converter.
• Carbon dioxide


  Produced by the combustion of any fossil fuel, as the carbon in the fuel is oxidised. It is a 'greenhouse gas' that absorbs infrared radiation from the earth's warm surface, trapping heat energy in the atmosphere. Rising CO2 levels are thought to be causing an increase in this 'greenhouse effect', leading to an increase in average temperatures around the world – global warming. Global warming may be causing climate change – alterations in long-term weather patterns, giving rise to the melting of polar ice caps, rising sea levels, and generally more extreme weather.
• Hydrogen


  Can be used as an alternative to fossil fuels. It produces no pollution when burned – only water vapour. However, it is difficult to store and transport, since it is a gas, and forms a very explosive mixture with oxygen. It has to be produced from water, and this requires energy, which must be produced from another source. Even if cheap, reliable hydrogen-powered cars are produced, the production of the hydrogen to power them would require at least as much energy as would be released by burning the fuel.
• Calculating enthalpy changes


  Write a balanced equation for the reaction.

  2. Find the total bond energy of every bond in the reactant molecules.

  3. Find the total bond energy of every bond in the product molecules.

  4. The overall enthalpy change is given by: ΔH = bond energy of reactants - bond energy of products.
• Breaking bonds requires energy (endothermic), and forming bonds releases energy (exothermic).
• The overall enthalpy change for a chemical reaction is the result of two changes in energy: energy must be supplied to the reactants to break all of the bonds in the molecules, and energy is then released when these atoms form bonds with each other to give the molecules of the products.
• In practice, the values obtained experimentally for enthalpy changes can be very different from those obtained theoretically using bond dissociation energies, due to heat loss and incomplete combustion.
• Carbon dioxide


  A relatively unreactive gas, slightly denser than air. It will dissolve sparingly in water, forming a slightly acidic solution (carbonic acid – H2CO3). It can be formed from the thermal decomposition of metal carbonates or the reaction of calcium carbonate with dilute hydrochloric acid. It has many uses, including as a fire extinguisher and in carbonated drinks.