Exothermic reactions include combustion of fuels, many oxidation reactions, acid-alkali neutralisation reactions, reactive metals with water,

moderately reactive metals with strong acids.

Exothermic reactions are used in self-heating cans and hand warmers.

An endothermic chemical reaction absorbs energy from the surrounding, usually in the form of heat energy, so cooling the surroundings, but sometimes the system is heated to provide the heat energy and a high enough temperature to promote the reaction. This means the products have more energy than the reactants and the surroundings have less energy.

Endothermic reactions include thermal decomposition of compounds e.g. carbonates, the reaction between citric acid and sodium hydrogencarbonate, sports injury packs to produce cooling effects

Why is it important to know about energy changes in chemical reactions?

Its important to know how much energy fuels release on combustion i.e. their calorific value.

Its important to know the energy released on burning petrol. diesel, coal or any other fossil fuel and alternative fuels like hydrogen or biofuels (biomass fuels).

The same sort of data is important in knowing how much energy is released on metabolising foods such as fats and carbohydrates.

Accurate energy change data is important in managing chemical processes in industry.

Exothermic reactions may provide their own heat if the process is carried out at high temperatures, energy transfer data provides some of the information needed.

Conversely, excess heat from an exothermic reaction may have to be removed using heat exchangers to avoid ‘overheating’ and excessive reaction rates that could be dangerous. If gases are involved, lack of control could lead to a build up of pressure resulting in an explosion.

Endothermic chemical processes often need a high temperature to promote

the absorption of heat energy, otherwise the reaction rate might economically far too slow. The amount of energy needed can be calculated from energy transfer data.

Heat Changes in Chemical Reactions

When chemical reactions occur, as well as the formation of the products – the chemical change, there is also a heat energy change which can often be detected as a temperature change.

This means the products have a different energy content than the original reactants (see the reaction profile diagrams below).

If the products contain less energy than the reactants, heat is released or given out to the surroundings and the change is called an exothermic reaction (exothermic energy transfer, exothermic energy change of the system).

This is illustrated by the simple energy level diagram above for an exothermic reaction.

The products have less energy than the original reactants (lower energy level) and the difference comes out as heat energy released to the surroundings.

The difference in heights of the energy levels tells you how much energy is released in an exothermic reaction.

The temperature of the system will be observed to rise in an exothermic change.

transfer,  So an exothermic reaction is one which gives out energy to the surroundings, usually in the form of heat energy, hence the rise in temperature.

Examples of exothermic reactions:

The burning or combustion of hydrocarbon fuels e.g. petrol or candle wax, these are very exothermic reactions.

1.The exothermic burning-combustion of fossil fuels is very important source of energy.

2.methane (natural gas) + oxygen ==> carbon dioxide + water (+ heat energy)

CH₄ + 2O₂ ==> CO₂ + 2H₂O

3.The burning of magnesium, reaction of magnesium with acids, or the reaction of sodium with water

2Mg + O₂ ==> 2MgO (+ heat energy)

4.The neutralisation of acids with alkalis

NaOH + HCl ==>NaCl + H₂O (+ heat energy)

Using hydrogen as a fuel in hydrogen-oxygen fuel cells xplosions are caused by VERY fast exothermic reactions producing very fast large expanding volumes of gases.

Other uses of exothermic reactions:

Hand warmers contain chemicals that when mixed together give out heat.

Self-heating cans of coffee, soup or hot chocolate have chemicals contained in the base of the container that when mixed generate enough energy to heat the contents of the can.

If the products contain more energy than the reactants, heat is taken in or absorbed from the surroundings and the change is called an endothermic reaction (endothermic energy endothermic energy change of the system).

This is illustrated by the simple energy level diagram above for an endothermic reaction.

The products have more energy than the original reactants (higher energy level) and the difference comes out as heat energy absorbed from the surroundings.

The difference in heights of the energy levels tells you how much energy is absorbed in an endothermic reaction.

If the change can take place spontaneously, the temperature of the reacting system will fall but, as is more likely, the reactants must be heated to speed up the reaction and provide the absorbed heat.

So an endothermic reaction is one which absorbs energy from the surroundings, usually in the form of heat, hence the observed fall in temperature in some reaction.

Examples of endothermic reactions

The thermal decomposition of limestone

calcium carbonate (limestone) ==> calcium oxide (lime) + carbon dioxide

CaCO₃ (+ heat energy) ==>CaO + CO₂

This only happens at temperatures above 900^oC.

the cracking of oil fractions

e.g. octane (+ heat energy) ==> hexane + ethene

C₈H₁₈ ==> C₆H₁₄ + C₂H₄

Again this needs a very high temperature AND a catalyst too.

These are two very important endothermic reactions used in the chemical industry.

Dissolving ammonium nitrate in water doesn’t need heating, the salt spontaneously dissolves and the temperature of the water/solution immediately falls as energy from the surroundings is absorbed, in fact from the water itself.

Other uses of endothermic reactions:

Some sports injury packs have a mixture of chemicals (or maybe a salt and water) that when mixed undergo an endothermic energy change, thereby absorbing heat from the surroundings, cooling some poor bruised limb! Rather more convenient and less messy than packs of ice!

One of the most important endothermic reactions, for which most of animal life depends is photosynthesis.

The energy from sunlight is absorbed as water and carbon dioxide are converted to glucose and oxygen.

6H₂O + 6CO₂ + sunlight energy ==> C₆H₁₂O₆ + 6O₂

However, on ‘burning’ the glucose/carbohydrates in our bodies, the ‘stored’ sunlight energy is released to keep us warm and drive all the chemical processes in our cells, so the opposite reaction is exothermic!

C₆H₁₂O₆ + 6O₂  ==>  6H₂O_(l) + 6CO₂  + heat/chemical energy

The difference between the energy levels of the reactants and products gives the overall energy change for the reaction

At a more advanced level the heat change is called the enthalpy change is denoted by delta H, ΔH.

ΔH is negative (-ve) for exothermic reactions i.e. heat energy is given out and lost from the system to the surroundings which warm up.

ΔH is positive (+ve) for endothermic reactions i.e. heat energy is gained by the system and taken in from the surroundings which cool down OR, as is more likely, the system is heated to provide the energy needed to effect the change.

Extra NOTE: In some exothermic changes, no heat is not released e.g. in batteries and fuel cells, where the energy is released as electrical energy.

Reversible Reactions and energy changes

If the direction of a reversible reaction is changed, the energy change is also reversed.

For a reversible reaction, the energy released in the exothermic reaction is numerically equal to the heat absorbed in the reverse reaction.

For example: the thermal decomposition of hydrated copper(II) sulphate is a very good example to observe in the school laboratory, even though it is not practical to measure the actual energy changes involved.

blue hydrated copper(II) sulphate + heat    white anhydrous copper(II) sulphate + water

CuSO₄.5H₂O_(s)     CuSO_4(s) + 5H₂O_(g)

On heating the blue solid, hydrated copper(II) sulphate, steam is given off and the white solid of anhydrous copper(II) sulphate is formed and left as the residue.

This is a thermal decomposition and is endothermic as heat is absorbed (taken in) from the surroundings.

The energy is needed to break down the crystal structure and drive off the water.

When the white solid is cooled and a few drops of water added, blue hydrated copper(II) sulphate is reformed and heat energy is given out to the surroundings, the mixture hots up!

The reverse reaction is exothermic as heat is given out.

i.e. on adding water to white anhydrous copper(II) sulphate the mixture heats up as the blue crystals reform.

Water molecules recombine with the copper ion releasing energy when the new bonds are formed.

Types of Chemical Reactions

Chemical reactions are processes in which substances change into other substances.

You know a chemical reaction takes place if one or more of these occur:

• Color changes – Different combinations of molecules reflect light differently. A color change indicates a change in molecules.
• Heat content changes – In all chemical reactions, the heat content of the reactants and the heat content of the products is never the

same. Sometimes the difference is great and can be easily detected. At other times, the difference is slight and more difficult to detect.

• Gas produced – Whenever a gaseous product forms in a liquid solution, bubbles can be seen. A colorless gas produced in a

reaction of solids is much harder to detect.

• Precipitate forms – Precipitates are insoluble products formed by a reaction taking place in a liquid solution. This insoluble product will eventually settle to the bottom, but might immediately appear by turning the clear solution cloudy.

Most chemical reactions can be placed into one of five basic types:

1. Decomposition Reactions

• A compound breaks into parts.
• compound → element + element
• 2H₂O → 2H₂ + O₂

Some decomposition complications with heat:

• Some acids, when heated, decompose into an acidic oxide and H₂O.

H₂SO₃ → SO₂ + H₂O

• Metallic hydroxides, when heated, decompose into a metallic oxide and H₂O.

Ca(OH)₂ → CaO + H₂O

• Metallic carbonates, when heated, decompose into a metallic oxide and CO₂.

Li₂CO₃ → Li₂O + CO₂

Metallic chlorates, when heated, decompose into metallic chlorides and O₂.

• The products of combustion are always carbon dioxide and water.
  hydrocarbon + oxygen → carbon dioxide + water
• CH₄ + 2O₂ → CO₂ + 2H₂O
• When metallic substances combine with oxygen, the result is an oxidation-reduction reaction.

The rusting of iron – 4Fe + 3O₂ → 2Fe₂O₃

Chemical reactions can be classified in other ways as well:

Neutralization Reactions

• Special types of double displacement reactions that involve the reaction between an acid and base to form a salt and water.
• acid + base → salt + water
• Heat is usually given off in neutralization reactions.
• A suspension of solid magnesium hydroxide in water is widely used as an antacid to neutralize excess stomach acid:

Mg(OH)_{2 (s)} + 2HCl _(aq) → MgCl_{2 (aq)} + 2H₂O _(l)

Oxidation-Reduction Reactions

• Any reaction in which elements experience a change in oxidation number.
• one atom gains e^&minus and another atom loosese^&minus

S + O₂ → SO₂

• In the reaction above, sulfur and oxygen both have an oxidation number of zero before the reaction. After the reaction, sulfur is +4 and oxygen is −2.

Precipitation Reactions

• Aqueous reactions that involve the formation of a precipitate (solid).
• soluble compound + soluble compound → insoluble compound

2KI _(aq) + Pb(NO₃)_{2 (aq)} → 2KNO_{3 (aq)} + PbI_{2 (s)}

The physical state symbol (aq) says the reaction is taking place in a water solution. The physical state symbol (s) says the lead (II) iodide is a solid – therefore insoluble

FIRST AND SECOND LAW OF THERMODYNAMICS.

Thermodynamics is a branch of physics which deals with the energy and work of a system. Thermodynamics deals only with the large scale response of a system which we can observe and measure in experiments. In aerodynamics, the thermodynamics of a gas obviously plays an important role in the analysis of propulsion systems but also in the understanding of high speed flows. The first law of thermodynamics defines the relationship between the various forms of energy present in a system (kinetic and potential), the work which the system performs and the transfer of heat. The first law states that energy is conserved in all thermodynamic processes.

The First Law of Thermodynamics – Energy is conserved.

The internal energy of a system is the sum of all the kinetic and potential energies of all its components.

ΔE = E_final − E_initial

The equation above indicates the change in internal energy, ΔE, is the difference between the final energy, E_final, and the initial energy, E_initial.

Thermodynamic quantities like ΔE have three parts: a number, a unit, and a sign.

• The number and unit give the magnitude of the charge.
• The sign gives the direction.

• A positive ΔE, when E_final>E_initial, indicates the system gained energy from its surroundsings. (diagram B below)
• A negative ΔE, when E_final<E_initial, indicates the system lost energy to its surroundsings. (diagram A below)

In a chemical reaction the initial state of the system refers to the reactants and the final state refers to the products.

2 H_2(g) + O_2(g) → 2 H₂O_(g)

In the reaction above, the system loses energy to the surroundings as heat. Because heat is lost from the system, the internal energy of the products (final state) is less than that of the reactants (initial state), and ΔE is negative.

Another way of saying this – as heat is added to or removed from a system, work is done on or by the system.

Using q to represent the heat added to or removed from the system, and w to represent work, the First Law of Thermodynamics can be represented by the equation:

ΔE = q + w

For q

• +  means system gains heat
• −  means system loses heat

For w

• +  means work is done on system
• −  means work is done by system

For ΔE

• +  means net gain of energy by system
• −  means net loss of energy by system

When a process (such as a chemical reaction) occurs in which the system absorbs heat, the process is called endothermic.

A process in which the system loses heat is called exothermic.

The work involved in the expansion or compression of gases is called pressure-volume work, or P-V work.

When the pressure is constant, the sign and magnitude of the P-V work is given by

w = − P ΔV

where P is pressure and ΔV is the change in volume of the system

(ΔV = V_final − V_initial).

Enthalpy, H is the thermodynamic function that accounts for heat flow in processes occurring at constant pressure when no forms of work are performed other than P-V work.

H = E + PV

At constant pressure, a change in enthalpy equals the change in internal energy plus the product of the constant pressure times the change in volume.