Exothermic & Endothermic Reactions

• Chemical reactions involve energy changes, either taking in or giving out energy, commonly as heat
• Reactions are classified into two groups: exothermic and endothermic

Exothermic Reactions

• Exothermic reactions release heat energy to the surroundings
• Example: iron and sulphur reacting and glowing, mixing silver nitrate and sodium chloride resulting a temperature rise, adding water to quicklime
• Exo means out
• Reactants become products and release heat energy
• In an energy level diagram, products have less energy than reactants

Other Examples of Exothermic Reactions

• Neutralization of an acid by an alkali
• Combustion of fuels
• Respiration in body cells
• The unit for energy is the joule (J)
• 1 kilojoule (kJ) = 1000 joules

Endothermic Reactions

• Endothermic reactions absorb heat energy from the surroundings
• Endo means in
• Example: barium hydroxide reacting with ammonium chloride and temperature falls, neutralizing ethanoic acid with sodium carbonate also produces a fall in temperature, mixing citric acid and sodium hydrogencarbonate
• Reactants plus heat energy become products
• The energy level diagram shows products having more energy than reactants
• Reactants need to overcome an energy gap, drawing energy from surroundings

Other Examples of Endothermic Reactions

• Cooking food
• Polymerization of ethene to polythene
• Photosynthesis
• Electrolysis

Explaining Energy Changes

• Chemical reactions involve the breaking and forming of bonds
• Breaking bonds requires energy (endothermic)
• Forming bonds releases energy (exothermic)
• In an exothermic reaction, the energy released during bond formation is greater than the energy required for bond breaking
• In an endothermic reaction, the energy released during bond formation is less than the energy required for bond breaking
• Bond energy refers to the energy needed to break a bond (given in kilojoules kJ), or released when bonds form

Calculating energy changes in reactions

• Energy in to break bonds and energy out from bonds forming
• Overall the reaction gives out or takes energy
• For every reaction, even an explosive one, some energy must be supplied to break some bonds to start it off. This is called the activation energy
• Fuels give out energy when they burn - combustion is exothermic.

Reversible Reactions

• Reversible reactions can proceed in both directions
• Shown with a symbol
• An example is heating copper(II) sulphate crystals, they break down into anhydrous copper(II) sulphate, a white powder
• Forward reaction: reactants to products
• Back reaction: products to reactants
• If a reversible reaction is endothermic in one direction, it is exothermic in the other
• The water in copper(II) sulphate crystals is called water of crystallization
• A blue compound is hydrated
• A white compound is anhydrous
• Thermal decomposition is breaking a substance down into simpler substances by heating it
• An example is ammonium chloride is heated, it decomposes

Reversible reactions and dynamic equilibrium

• In a closed system, a reversible reaction eventually reaches a state of dynamic equilibrium
• The forward and back reactions take place at the same rate, so no overall change occurs
• Dynamic equilibrium means the changes are still happening
• Equilibrium means there is no overall change.

Dynamic Equilibrium in Industry

• Many industrial reactions are reversible, like the formation of ammonia
• Equilibrium is reached when ammonia is produced and broken down at the same rate, preventing the reaction from completion

Shifting the equilibrium - Le Chatelier's Principle

• Le Chatelier's principle: When a reversible reaction is in equilibrium and you make a change, it will do what it can to oppose that change
• Changes in conditions can shift equilibrium to increase product yield

Increasing the temperature

• The forward reaction is exothermic and gives out heat
• The back reaction is endothermic and take it in
• Heating shifts equilibrium to the left, favoring reactants and reducing ammonia production

Increasing the pressure

• Pressure results from gas molecule collisions
• Increasing pressure favors the side with fewer gas molecules
• In ammonia production, more ammonia (fewer molecules) is formed to reduce pressure
• Removing ammonia condenses it out in liquid, and nitrogen and hydrogen react to restore equilibrium

Adding a catalyst

• Iron catalysts speeds up the forward and back reactions equally so equilibrium is reached quickly
• Choosing optimum conditions require considering high pressure and catalyst. moderate amount of heat to find a compromise for shift equilibrium and reaction rate
• If the forward reaction is exothermic, temperature increase means yield decrease
• If the forward reaction is endothermic, temperature increase means yield increase
• For an equation, if there are fewer molecules of product than reactant, pressure increase means yield increase

Making ammonia in industry

• The Haber process makes ammonia in industrial, and it is a key chemical for fertilizers
• Raw materials are hydrogen and nitrogen that are mixed and scrubbed (cleaned)
• Gas mixture is compressed to high pressure and heated so that iron catalyst convert the mixture to ammonia, by catalysis
• Cooled and ammonia separates out as liquid
• Unreacted reagents recycled
• The catalyst slowly degrade and needs to be replaced

The economics of making ammonia

• Industrial ammonia production reduces costs
• High pressure increases yield but requires expensive equipment.
• High temperature increases the reaction rate but decreases the equilibrium yield as the catalyst get poisoned, and it's unreacted
• The catalysts get poisoned by impurities and get replaced, so recycle the unreacted gases
• Plants nearby natural gas and water use the chemicals for profit