Energy Changes

Exothermic - transfers energy to the surroundings (exits)

An exothermic reaction is shown by a rise in temperature

-Examples:

• Combustion - very exothermic as it gives out lots of energy
• Neutralisation reactions - acid + alkali
• Many oxidation reactions - sodium + water releases energy and the sodium is oxidised

-Everyday Uses:

• Hand warmers -  oxidation of iron + air (with a salt solution catalyst) to release energy
• Self heating cans - between the chemicals in their bases

Endothermic - takes in energy from the surroundings

Endothermic reactions are shown by a fall in temperature

-Examples:

• Citric Acid + Sodium Hydrogencarbonate - reaction in Sherbet
• Thermal Decomposition - CaCO3 (+heat) -> CO2 +CaO

-Everyday Uses:

• Sport Injury Packs - allows the injury pack to become instantly cooler without having to use a freezer

-Reaction Profiles:

Reaction Profiles - show the relative energy of the reactants and products in a reaction and how the energy changes over the course of the reaction

Exothermic - products are lower than the reactants

Endothermic - products are higher than the reactants

-Activation Energy:

Activiation Energy is the minimum amount of energy for the reactants to start colliding and react.

The greater the activation energy, the more energy needed to start the reaction. The energy supplied to the reaction could be by heating the reaction mixture

-Required Practical:

• 1) Pour 30cm3 of 2M dilute HCl into a Polystyrene Cup using a 50cm3 measuring beaker
• 2) Stand the cup inside of a beaker 
• 3) Measure the temperature of the acid and record
• 4) Pour 5cm3 of 2M NaOH into a 10cm3 measuring cylinder
• 5) Put the NaOH into the cup, fit the lid and gently stir using the themometer through the hole
• 6) Record the temperature when it stops changing.
• 7) Repeat 4-6 until a total of 40cm3 of NaOH is added
• 8) Repear 1-8 for a second trial and calculate an average
• 9) Plot a graph with Mean Temperature (y-axis) and Volume of NaOH added (x-axis) with a line of best fit

Energy supplied to break bonds - endothermic 

Energy released when forming bonds - exothermic 

-Bond Energy Calculations:

Each chemical bond has a particular bond energy assosiated with it - this can be used to calculate the overall energy change of a reaction

Endothermic - energy chnage is positive as it's taken in energy from its surroundings

Exothermic - energy change is negative as it's released energy to its surroundings

-Example:

• H2 + C2 -> 2HCl
• Energy required to break bonds - (1x H-H) + (1x Cl-Cl) = 678kJ/mol
• Energy released from formation - (2x H-Cl) = 862kJ/mol
• Overall change - 678kJ/mol - 862kJ/mol = -184kJ/mol (exothermic) 

Electochemical cell - 2 different electrodes in contact with an electolyte:

• 2 electrodes must conduct electricity (usually metals)
• Electolyte is a liquid wich reacts with the electrodes as it contains ions

The chemical reaction between the 2 electrodes and the electrolyte creates a charge difference between the electrodes.

The electrodes are connected by a wire so the charge is able to flow - electricity is produced.

A voltmeter could be connected to measure the voltage of the cell.

Voltage of a cell depends on the electrolyte and the electrodes used - different metals react differently with the same electrolyte

• Different electrode = different reactivities
• Different electrolyte = different ions in the liquid

Bigger the difference in reactivity of the electrodes, the bigger the voltage of the cell

-Batteries:

• A battery is formed by connecting 2 or more cells in series.
• Voltages of the cells are combined to give a larger overall voltage

Overtime the reacting particles are used up and are turned into the product of the reaction. The reaction can no longer carry on and no electricity is produces - irreversable (i.e. alkaline batteries)

If the reaction in a cell is reversable, the cell can be recharged by connecting it to an external electric current

Fuel cell - an electrical cell that's supplied with a fuel + oxygen

The fuel that enters the cell would become oxidised and sets up a potential difference in the cell

There are many type of fuel cells which use different types of electrolytes and fuels

• i.e Hydrogen-Oxygen fuel cell

• Electrolyte - usually Potassium Hydroxide solution
• Electrodes - usually porous carbon with a catalyst

• Anode - H2 loses electrons to produce H+ ions (oxidation):
  • 2H+ + 2e -> H2
• Cathode - O2 gains electrons and reacts with H+ ions to form H20 (reduction):
  • 02 + 4H+ + 4e -> 2H20 
• Electrons flow through an external circuit from anode to cathode - this is the electric current
• Overall reaction:
  • 2H2 + O2 -> 2H20

-Advantages in Vehicles:

• Less pollutants released - by-products are water + heat
• Less expensive - than batteries in electric vehicles
• Less time-consuming - doesn't need to be recharged as often