Results are in... 90% of Mr Gundry's students achieved a grade between A*-C! How'd you do? :)
C3 Energy Calculations
An exothermic reaction between a fuel and oxygen, that produces carbon dioxide (and/or monoxide) and water.
A substance that releases a lot of energy when burned.
Not all fuels release the same amount of energy when burned. The temperature rise of the surroundings when a fuel burns is related to the amount of energy given off. We can calculate the energy using this equation:
- Q = energy released by fuel in joules (J)
- c = specific heat capacity of object heated (usually water), and is the amount of energy needed to heat 1g of the object by 1 °C. Water has a specific heat capacity of 4.2 J/g°C
- m = the mass of the object heated measured in grams (g)
- T = the temperature rise measured in °C (the Δ means 'change in' temperature)
This is a slight topic cross over with physics, and the following video explains what specific heat capcity is, and how you can reverse calculate it. This is not needed for your GCSE in chemistry.
Energy Level Diagrams
During every chemical reaction bonds are broken, and new bonds are formed. Depending on the amount of energy needed to break or make the bonds, will depend on whether, overall, energy is released or absorbed.
Bond breaking is the first step in any chemical reaction. This requires the input of energy, and so initially all reactions will absorb energy from the surroundings. The bonds in the reactants need to be broken before they can react and form new bonds.
Bond making transfers energy into the surrounding environment, and if more energy is needed to break bonds than is released the reaction is endothermic overall. If the amount of energy released is more than the energy needed to break bonds then the reaction is exothermic overall.
Exothermic reactions release energy to the surroundings. This is because more energy is released by making new bonds, than needed to break bonds.
Examples of exothermic reactions include: combustion and neutralisation
Endothermic reactions absorb energy from the surroundings. This is because more energy is needed to break the bonds in the reactants, than is released by forming new bonds in the products.
Examples of endothermic reactions are: thermal decomposition and electrolysis.
Just like in the law of the conservation of mass (which says matter cannot be created or destroyed - the reason why all chemical equations must be balanced), the amount of energy at the start of a reaction, will always equal the energy at the end of the reaction. This might seem confusing as energy level diagrams show either energy being gained or lost, but what they don't show is energy stored in chemical bonds.
Thinking back to our notes on rates, we know that for a successful collision to happen there is a minimum amount of energy needed for a reaction to happen. We called this the activation energy. We can show this on our energy level diagrams as an initial hump before the products are formed.
If we add a catalyst to a reaction we know this lowers the activation energy needed, and that hump will be much smaller.
Using current fuels is bad for the environment as they release pollutants (such as carbon dioxide) that contribute to global warming, global dimming, acid rain etc...
Using hydrogen as a fuel produces only one product: water. Hydrogen is burned in oxygen to only produce water! This is much better than what we currently use. It's most efficiently used in something called a hydrogen fuel cell.
Page last updated: 14/04/2017