ENERGY AND REACTIONS

ENERGY AND REACTIONS

ENDOTHERMIC AND EXOTHERMIC REACTIONS

ENDOTHERMIC AND EXOTHERMIC REACTIONS

HEAT OF REACTION

HEAT OF REACTION

BOND ENERGY

BOND ENERGY

Heat of combustionHeat of combustion

Heat of neutralizationHeat of neutralization

HEAT OF FORMATION

HEAT OF FORMATION

Hess’s lawHess’s law

OUR PLANET IS WARMING. WHY?

OUR PLANET IS WARMING. WHY?

The study of the heat changes that accompany chemical reactions is called thermochemistry

The study of the heat changes that accompany chemical reactions is called thermochemistry

Burning fuels in air always produces heat, “burning” food inside our bodies supplies us with the energy we need

Burning fuels in air always produces heat, “burning” food inside our bodies supplies us with the energy we need

Any chemical reaction that produces heat is called an exothermic reaction

Any chemical reaction that produces heat is called an exothermic reaction

EXOTHERMIC AND ENDOTHERMIC REACTIONS

EXOTHERMIC AND ENDOTHERMIC REACTIONS

Combustion of methane is one example of a reaction that produces heat

Combustion of methane is one example of a reaction that produces heat

NH4Cl

EXOTHERMIC AND ENDOTHERMIC REACTIONS

EXOTHERMIC AND ENDOTHERMIC REACTIONS

We note a drop in temperature if we hold a beaker of water at room temperature and add some ammonium chloride (NH4Cl) to it: in this case heat is taken in from the hand.

We note a drop in temperature if we hold a beaker of water at room temperature and add some ammonium chloride (NH4Cl) to it: in this case heat is taken in from the hand.

Any chemical reaction that takes in heat from the surroundings is called an endothermic reaction

Any chemical reaction that takes in heat from the surroundings is called an endothermic reaction

Heat of reaction is the heat change when the moles of reactants, indicated in the balanced equation, react completely.

Heat of reaction is the heat change when the moles of reactants, indicated in the balanced equation, react completely.

HEAT OF REACTIONHEAT OF REACTION

Hydrogen is a fuel which burns in oxygen to form water:

H2(g) + ½ O2(g) H2O(g) H = - 242 kJ

The symbol H (delta H) indicates the variation of hentalpy: the heat change taking place when a chemical reaction occurs at a constant pressure and a constant temperature. This equation states that when one mole of hydrogen reacts with half a mole of oxygen to form one mole of steam 242 kJ of heat are released.

A negative value of H indicates an exothermic reaction

A positive value of H indicates an endothermic reaction

The H value above is referred to as the molar heat of reaction. If the number of moles in the balanced equation are changed, then the heat of reaction also changes:

2 H2(g) + O2(g) 2H2O(g) H = - 484 kJ

Hydrogen is a fuel which burns in oxygen to form water:

H2(g) + ½ O2(g) H2O(g) H = - 242 kJ

The symbol H (delta H) indicates the variation of hentalpy: the heat change taking place when a chemical reaction occurs at a constant pressure and a constant temperature. This equation states that when one mole of hydrogen reacts with half a mole of oxygen to form one mole of steam 242 kJ of heat are released.

A negative value of H indicates an exothermic reaction

A positive value of H indicates an endothermic reaction

The H value above is referred to as the molar heat of reaction. If the number of moles in the balanced equation are changed, then the heat of reaction also changes:

2 H2(g) + O2(g) 2H2O(g) H = - 484 kJ

HEAT OF COMBUSTIONHEAT OF COMBUSTION

A reaction in which a substance is burned is called a combustion reaction.

Consider the reaction in which carbon is burned in oxygen to form carbon dioxide : C(s) + O2(g) CO2(g) H = - 393 kJ

The heat given out in the reaction is called the heat of combustion.

A reaction in which a substance is burned is called a combustion reaction.

Consider the reaction in which carbon is burned in oxygen to form carbon dioxide : C(s) + O2(g) CO2(g) H = - 393 kJ

The heat given out in the reaction is called the heat of combustion.

Heat of combustion of a substance is the heat change when one mole of the substance is completely burned in excess oxygen.

Heat of combustion of a substance is the heat change when one mole of the substance is completely burned in excess oxygen.

• The phrase “is completely burned in excess oxygen” means that the product of reaction must be carbon dioxide and not carbon monoxide as it occurs in an incomplete combustion:

C(s) + ½ O2(g) CO(g) H = - 111 kJ

The value of H is not the heat of combustion of carbon.

• Another important phrase is “one mole”.

Consider the thermochemical reaction in which butane is burned:

2 C4H10(g) + 13 O2(g) 8CO2(g) + 10H2O(g) H = - 5720 kJ

The heat of combustion of butane is not - 5720 kJ,

In fact the heat of combustion involves one mole of butane being burned; therefore the correct value is

–5720/2 = -2860 kJ mol-1

HEAT OF COMBUSTIONHEAT OF COMBUSTION

HEAT OF COMBUSTIONHEAT OF COMBUSTION

Heat of combustion is accurately measured using a bomb calorimeter.

This instrument consists of a steel container (the bomb), with a screw-on cap. The sample whose heat of combustion is to be measured is placed in a crucible inside the bomb. The bomb is placed in a container of water (the calorimeter), oxygen is pumped inside and the sample is ignited by electric wires. By measuring the increase in temperature of the water it is possible to measure the heat of combustion of the fuel.

Heat of combustion is accurately measured using a bomb calorimeter.

This instrument consists of a steel container (the bomb), with a screw-on cap. The sample whose heat of combustion is to be measured is placed in a crucible inside the bomb. The bomb is placed in a container of water (the calorimeter), oxygen is pumped inside and the sample is ignited by electric wires. By measuring the increase in temperature of the water it is possible to measure the heat of combustion of the fuel.

HEAT OF COMBUSTIONHEAT OF COMBUSTION

In some cases, in industry, chemists measure the heat given out by a fuel (or food) in terms of the mass of fuel rather than one mole of substance, in order to compare the efficiency of various fuels and to state if the fuel is of good quality.

For example, if you read the side of a box of cornflakes, you will find that 100g of cornflakes provides about 1550 kJ of energy.

Chemists use the term: kilogram calorific value.

In some cases, in industry, chemists measure the heat given out by a fuel (or food) in terms of the mass of fuel rather than one mole of substance, in order to compare the efficiency of various fuels and to state if the fuel is of good quality.

For example, if you read the side of a box of cornflakes, you will find that 100g of cornflakes provides about 1550 kJ of energy.

Chemists use the term: kilogram calorific value.

The kilogram calorific value of a fuel is the heat, the energy, produced when 1 kg of the fuel is completely burned in oxygen.

The kilogram calorific value of a fuel is the heat, the energy, produced when 1 kg of the fuel is completely burned in oxygen.

Some examples of the heat of combustion of various fuelsSome examples of the heat of combustion of various fuels

143.000-286H2(g)Hydrogen

32.750-393C (g)Carbon

29.804-1371C2H5OH(l)Ethanol

48.767-4194C6H14(l)Hexane

49.603-2877C4H10(g)Butane

50.454-2230C3H8(g)Propane

52.000-1560C2H6(g)Ethane

55.625-890CH4(g)Methane

Kilogram caloric value (kJ kg-1)

Heat of combustion H (kJ mol-1) at 25 °C

FormulaFuel

HEAT OF COMBUSTIONHEAT OF COMBUSTION

HEAT OF NEUTRALISATIONHEAT OF NEUTRALISATION

Neutralisation: a reaction between an acid and a base to form a salt and water:

HCl + NaOH NaCl + H2O

Strong acids, strong bases and salts are fully dissociated in water, thus the essential reaction taking place ingnores spectator ions Na+ and Cl-:

H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O

H+ + OH- H2O H = - 57.1 kJ

Neutralisation reactions are nearly always exothermic.

The heat of neutralisation is the heat change when one mole of H+ ions from an acid reacts with one mole of OH- ions from a base.

The heat of neutralisation is the heat change when one mole of H+ ions from an acid reacts with one mole of OH- ions from a base.

The heat of reaction in a reaction of neutralisation depends on the numbers of moles of water formed, therefore the heat of reaction between sulphuric acid and sodium hydroxide is double since two moles of water are formed:

H2SO4 + 2NaOH Na2SO4 + 2H2O H = - 2 · 57 kJ = - 114 kJ

BOND ENERGYBOND ENERGY

What determines the value of H for a particular reaction? Consider the combustion of methane:CH4(s) + 2O2(g) CO2(g) + 2H2O(g) H = - 890 kJ

What determines the value of H for a particular reaction? Consider the combustion of methane:CH4(s) + 2O2(g) CO2(g) + 2H2O(g) H = - 890 kJ

The bond energy is the energy required to break one mole of covalent bonds and to separate the neutral atom completely from each other (the same amount of energy is released when one mole of bonds is formed).

The bond energy is the energy required to break one mole of covalent bonds and to separate the neutral atom completely from each other (the same amount of energy is released when one mole of bonds is formed).

Reactants Products

Energy absorbed to break bonds of reactants.

Energy released when bonds of products are formed.

As known, when a chemical reaction occurs, the bonds of reactants are broken and the bonds of products are formed. Energy is required to break bonds and energy is released when bonds are formed.

As known, when a chemical reaction occurs, the bonds of reactants are broken and the bonds of products are formed. Energy is required to break bonds and energy is released when bonds are formed.

BOND ENERGYBOND ENERGY

Let us consider the energy required to break 4 C – H bonds and the energy required to break the two double bonds in the oxygen molecules; these values (table to the right) are considered positives:

Let us consider the energy required to break 4 C – H bonds and the energy required to break the two double bonds in the oxygen molecules; these values (table to the right) are considered positives:

+ 4 · 412 kJ = + 1648 kJ required

+ 2 · 496 kJ =+ 992 kJ required

743C = O

496O = O

463O – H

612C = C

348C - C

412C - H

Bond energy

( kJ mol-1)Bond

Below are the values of the energy released (negative values) when the bonds of products are formed:

Below are the values of the energy released (negative values) when the bonds of products are formed:

- 2 · 743 kJ = - 1486 kJ released

- 4 · 463 kJ = - 1852kJ released

BOND ENERGYBOND ENERGY

The H for the combustion of methane is obtained making the sum of energy required to break bonds of reactants and the energy released when bonds of products are formed:

The H for the combustion of methane is obtained making the sum of energy required to break bonds of reactants and the energy released when bonds of products are formed:

+ H = - 698kJ

+ 4 · 412 kJ = + 1648 kJ required+ 4 · 412 kJ = + 1648 kJ required

+ 2 · 496 kJ = + 992 kJ required+ 2 · 496 kJ = + 992 kJ required++ENERGYENERGY

REACTANTSREACTANTS

- 2 · 743 kJ = - 1486 kJ released- 2 · 743 kJ = - 1486 kJ released

- 4 · 463 kJ = - 1852kJ released- 4 · 463 kJ = - 1852kJ released++ENERGYENERGY

PRODUCTSPRODUCTS

HEAT OF FORMATIONHEAT OF FORMATION

The heat of formation of a compound is the heat change that takes place when one mole of a compound, in its standard state, is formed from its elements in their standard states.

The heat of formation of a compound is the heat change that takes place when one mole of a compound, in its standard state, is formed from its elements in their standard states.

The standard state of an element or compound is its normal form at 25 °C at one atmosphere pressure (101 kPa)

The heat of formation of water is – 285.8 kJ / mol and may be represented as:

H2(g) + ½ O2(g) H2O(l) H = - 285.8 kJ / mol

The following equation does not represent the heat of formation of water:H2(g) + ½ O2(g) H2O(g) H = - 241.8 kJ / mol

The reason is because the water formed in the above reaction is in the gaseous state. At 25 °C the normal state of water is a liquid.

The standard state of an element or compound is its normal form at 25 °C at one atmosphere pressure (101 kPa)

The heat of formation of water is – 285.8 kJ / mol and may be represented as:

H2(g) + ½ O2(g) H2O(l) H = - 285.8 kJ / mol

The following equation does not represent the heat of formation of water:H2(g) + ½ O2(g) H2O(g) H = - 241.8 kJ / mol

The reason is because the water formed in the above reaction is in the gaseous state. At 25 °C the normal state of water is a liquid.

HEAT OF FORMATIONHEAT OF FORMATION

gas0O2Oxygen

solid0FeIron

liquid- 278C2H5OHEthanol

gas- 46.2 NH3Ammonia

gas- 393CO2Carbon dioxide

gas- 111COCarbon monoxide

gas - 84. 7C2H6Ethane

gas- 74.9CH4Methane

Normal state at 25 °C

Heat of formaton

H (kJ mol-1)FormulaSubstance

The heat of formation of any element in its standard state is zero. The heats of formation of most compounds have negative value of H: energy is given out when the compound is formed from its elements. The more negative the value of the heat of formation is, the more stable is the compound.

The heat of formation of any element in its standard state is zero. The heats of formation of most compounds have negative value of H: energy is given out when the compound is formed from its elements. The more negative the value of the heat of formation is, the more stable is the compound.

HESS’S LAWHESS’S LAW

Hess’s law states that if a chemical reaction takes place in a number of stages, the sum of the heat changes in the separate stages, is equal to the heat change when the reaction is carried out in one stage.

Hess’s law states that if a chemical reaction takes place in a number of stages, the sum of the heat changes in the separate stages, is equal to the heat change when the reaction is carried out in one stage.

A B C D

HA D = HA B + HB C + HC D

In other words, the amount of energy liberated or absorbed in a chemical reaction is the same whether the reaction takes place in one o several steps, as shown below:

In other words, the amount of energy liberated or absorbed in a chemical reaction is the same whether the reaction takes place in one o several steps, as shown below:

HESS’S LAWHESS’S LAW

Consider the two routes for converting carbon and oxygen to carbon dioxide

Consider the two routes for converting carbon and oxygen to carbon dioxide

C C O2

C O

+ O2

+1/2 O2 +1/2 O 2

H 1 = - 394 kJ mol -

1

H 3

= - 283 kJ mol -

1

H2 = ?

C C O2

C O

+ O2

+1/2 O2 +1/2 O 2

H 1 = - 394 kJ mol -

1

H 3

= - 283 kJ mol -

1

H2 = ?

HESS’S LAWHESS’S LAW

Hess’s law may be used to calculate the unknown heat of the reaction that converts carbon to carbon monoxide: C + 1/2 O2 CO

Hess’s law may be used to calculate the unknown heat of the reaction that converts carbon to carbon monoxide: C + 1/2 O2 CO

H 1 = H 2 + H 3 - 394 = x - 283 x = - 394 + 283 = -111 kJ mol-1

C(s) + 1/2 O2(g) CO(g) H 2 = - 111 kJ mol-1

HESS’S LAWHESS’S LAW

Hess’s Law is useful because:

• the value of the heat of reaction can be calculated indirectly, if it proves difficult to find it directly in the laboratory. (It is difficult to measure the heat of formation of CO

since some CO2 is also formed during the experiment).

• thermochemical reactions can be added (+), subtracted (-), multiplied (x) or divided (:) to calculate unknown heat of reaction.

Hess’s Law is useful because:

• the value of the heat of reaction can be calculated indirectly, if it proves difficult to find it directly in the laboratory. (It is difficult to measure the heat of formation of CO

since some CO2 is also formed during the experiment).

• thermochemical reactions can be added (+), subtracted (-), multiplied (x) or divided (:) to calculate unknown heat of reaction.

HESS’S LAWHESS’S LAW

Calculate the heat of formation of CH4 from the following heats of combustion:

a) C(s) + O2(g) CO2 (g) H = - 393 kJ mol -

1

b) H2(g) + 1/2 O2(g) H2O (g) H = - 286 kJ mol -

1

c) CH4(g) + 2 O2(g) CO2 (g) + 2 H2O (g) H = - 879 kJ mol -

1

Calculate the heat of formation of CH4 from the following heats of combustion:

a) C(s) + O2(g) CO2 (g) H = - 393 kJ mol -

1

b) H2(g) + 1/2 O2(g) H2O (g) H = - 286 kJ mol -

1

c) CH4(g) + 2 O2(g) CO2 (g) + 2 H2O (g) H = - 879 kJ mol -

1

We need to write the balanced equation of formation of methane from his elements:

C(s) + 2 H2(g) CH4 (g) H = ?

We need to write the balanced equation of formation of methane from his elements:

C(s) + 2 H2(g) CH4 (g) H = ?

We must build up the reaction of formation of methane using the above mentioned thermochemical reactions. Introduce C

a) C(s) + O2(g) CO2 (g) H = - 393 kJ mol -1

We must build up the reaction of formation of methane using the above mentioned thermochemical reactions. Introduce C

a) C(s) + O2(g) CO2 (g) H = - 393 kJ mol -1

Introduce 2 H2. Multiply equation b) by 2. Do not forget to multiply H by 2 as well.

b) x 2 : 2 H2(g) + O2(g) 2 H2O (g) H = - 572 kJ mol -

1

Introduce 2 H2. Multiply equation b) by 2. Do not forget to multiply H by 2 as well.

b) x 2 : 2 H2(g) + O2(g) 2 H2O (g) H = - 572 kJ mol -

1

HESS’S LAWHESS’S LAW

HESS’S LAWHESS’S LAW

Get the left-hand side of the required equation by adding the following equations:

a) C(s) + O2(g) CO2 (g) H = - 393 kJ mol -1

b) x 2 : 2 H2(g) + O2(g) 2 H2O (g) H = - 572 kJ mol -

1

d) C(s) + 2 H2(g) + 2 O2(g) CO2 (g) + 2 H2O (g) H = - 965 kJ mol -1

Get the left-hand side of the required equation by adding the following equations:

a) C(s) + O2(g) CO2 (g) H = - 393 kJ mol -1

b) x 2 : 2 H2(g) + O2(g) 2 H2O (g) H = - 572 kJ mol -

1

d) C(s) + 2 H2(g) + 2 O2(g) CO2 (g) + 2 H2O (g) H = - 965 kJ mol -1

HESS’S LAWHESS’S LAW

Introduce CH4 We have CH4 on the left-hand side of equation c). We need it on the right-hand side of the equation for the formation of methane. Reverse equation c), remembering to reverse the sign of H too

c) reversed CO2 (g) + 2 H2O (g) CH4(g) + 2 O2(g) H = + 879 kJ mol -

1

Introduce CH4 We have CH4 on the left-hand side of equation c). We need it on the right-hand side of the equation for the formation of methane. Reverse equation c), remembering to reverse the sign of H too

c) reversed CO2 (g) + 2 H2O (g) CH4(g) + 2 O2(g) H = + 879 kJ mol -

1

Now add c) reversed to d) equation:

d) C(s) + 2 H2(g) + 2 O2(g) CO2 (g) + 2 H2O (g) H = - 965 kJ mol -1

c) reversed CO2 (g) + 2 H2O (g) CH4(g) + 2 O2(g) H = +879 kJ mol -1

C(s) + 2 H2(g)+ 2 O2(g)+CO2 (g)+ 2 H2O (g) CO2 (g)+ 2 H2O (g) + CH4(g) + 2 O2(g)

We obtain the required equation of formation of methane:

C(s) + 2 H2(g) CH4 (g)

H = (- 965 + 879) kJ mol -1 = - 86 kJ mol -1

Now add c) reversed to d) equation:

d) C(s) + 2 H2(g) + 2 O2(g) CO2 (g) + 2 H2O (g) H = - 965 kJ mol -1

c) reversed CO2 (g) + 2 H2O (g) CH4(g) + 2 O2(g) H = +879 kJ mol -1

C(s) + 2 H2(g)+ 2 O2(g)+CO2 (g)+ 2 H2O (g) CO2 (g)+ 2 H2O (g) + CH4(g) + 2 O2(g)

We obtain the required equation of formation of methane:

C(s) + 2 H2(g) CH4 (g)

H = (- 965 + 879) kJ mol -1 = - 86 kJ mol -1

HESS’S LAWHESS’S LAW

Calculate the heat liberated using the equation:

Heat liberated = (ms + me) · c · (t2 – t1)

ms= mass of solution in kg

me = mass equivalent in water of the calorimeter that absorbs heat (about 30 g)

c= specific heat capacity of solution in J kg-1 K-1 (assume 4060 J kg-1 K-1 for the solution)

t2 – t1 = temperature rise

Note: Assume that the density of the solution is 1 g / cm 3

Calculate the heat liberated using the equation:

Heat liberated = (ms + me) · c · (t2 – t1)

ms= mass of solution in kg

me = mass equivalent in water of the calorimeter that absorbs heat (about 30 g)

c= specific heat capacity of solution in J kg-1 K-1 (assume 4060 J kg-1 K-1 for the solution)

t2 – t1 = temperature rise

Note: Assume that the density of the solution is 1 g / cm 3

STUDENT EXPERIMENTSTUDENT EXPERIMENT

Calculate the heat of neutralisationKnowing the heat liberated when this number of moles of acid was neutralized, you can now calculate the amount of heat liberated when one mole of the acid is neutralized.

Calculate the heat of neutralisationKnowing the heat liberated when this number of moles of acid was neutralized, you can now calculate the amount of heat liberated when one mole of the acid is neutralized.

STUDENT EXPERIMENTSTUDENT EXPERIMENT

Experiment

To determine the heat of reaction (heat of neutralisation) of hydrochloric acid with sodium hydroxide.

Experiment

To determine the heat of reaction (heat of neutralisation) of hydrochloric acid with sodium hydroxide.

Introduction

In this experiment , an equal number of moles of hydrocloric acid are mixed with an equal number of moles of sodium hydroxide in a calorimeter. The rise in temperature is then measured and from this the heat of neutralization is calculated.

Introduction

In this experiment , an equal number of moles of hydrocloric acid are mixed with an equal number of moles of sodium hydroxide in a calorimeter. The rise in temperature is then measured and from this the heat of neutralization is calculated.

Apparatus required

Thermometer (reading to 0.1 °C), two beakers (100 ml), a calorimeter (250 ml), two burettes (50 ml).

Apparatus required

Thermometer (reading to 0.1 °C), two beakers (100 ml), a calorimeter (250 ml), two burettes (50 ml).

Chemicals required

Hydrochloric acid (1.0 M), sodium hydroxide (1.0 M).

Chemicals required

Hydrochloric acid (1.0 M), sodium hydroxide (1.0 M).

Procedure

• Using a burette, place 50 ml of the 1 M hydrochloric acid solution into the first beaker.

• Using the second burette, place 50 ml of the 1 M solution of sodium hydroxide into the second beaker.

• Measure separately the temperature of the two solutions.

• When the two solutions are at the same temperature, place the acid into the calorimeter and quickly add the base to the acid, stirring well.

• Place the lid on the calorimeter, to prevent heat loss, and record the maximum temperature reached.

Procedure

• Using a burette, place 50 ml of the 1 M hydrochloric acid solution into the first beaker.

• Using the second burette, place 50 ml of the 1 M solution of sodium hydroxide into the second beaker.

• Measure separately the temperature of the two solutions.

• When the two solutions are at the same temperature, place the acid into the calorimeter and quickly add the base to the acid, stirring well.

• Place the lid on the calorimeter, to prevent heat loss, and record the maximum temperature reached.

STUDENT EXPERIMENTSTUDENT EXPERIMENT

STUDENT EXPERIMENTSTUDENT EXPERIMENT

Summarize results as follows:

Temperature of HCl before mixing = °C

Temperature of NaOH before mixing = °C

Highest temperature reached after mixing = °C

Temperature rise = °C

Number of moles of acid used = mol

Number of moles of base used = mol

Mass of solution = Kg

Summarize results as follows:

Temperature of HCl before mixing = °C

Temperature of NaOH before mixing = °C

Highest temperature reached after mixing = °C

Temperature rise = °C

Number of moles of acid used = mol

Number of moles of base used = mol

Mass of solution = Kg

• Huge amounts of energy are needed for heating, generating electricity and for transport.

• Fuels give off numerous emissions that many say are bringing about great climate changes, because of the greenhouse effect of global warming.

• All economic activity that requires energy consumption contributes to these emissions producing greenhouse gases.

• Huge amounts of energy are needed for heating, generating electricity and for transport.

• Fuels give off numerous emissions that many say are bringing about great climate changes, because of the greenhouse effect of global warming.

• All economic activity that requires energy consumption contributes to these emissions producing greenhouse gases.

CLIMATE CHANGECLIMATE CHANGE

Why is the planet warming?

Energy: is the largest source of CO2 – mainly emissions

from fossil fuel combustion to produce electricity.

Energy: is the largest source of CO2 – mainly emissions

from fossil fuel combustion to produce electricity.

CLIMATE CHANGECLIMATE CHANGE

Who and what contributes to greenhouse gases?Who and what contributes to greenhouse gases?

Industry: mainly contributes to greenhouse gas emissions through energy use, including direct consumption of fossil fuels and use of electricity (e.g. CO2 is a by-product of cement manufacture).

Industry also produces some CFC’s although they, along with substances that were attacking the ozone layer, were banned by the Montreal Protocol in 1989.

Industry: mainly contributes to greenhouse gas emissions through energy use, including direct consumption of fossil fuels and use of electricity (e.g. CO2 is a by-product of cement manufacture).

Industry also produces some CFC’s although they, along with substances that were attacking the ozone layer, were banned by the Montreal Protocol in 1989.

Agriculture: is the largest producer of CH4.Agriculture: is the largest producer of CH4.

Transport: CO2 emissions come from fuel combustion.Transport: CO2 emissions come from fuel combustion.

… is a gas in the atmosphere that freely allows radiation from the sun to reach the earth’s surface, but traps the heat radiated back from the earth’s surface towards space.

The heating effect is analogous to the manner in which the glass of a greenhouse traps the sun’s radiation to warm the air inside the greenhouse.

… is a gas in the atmosphere that freely allows radiation from the sun to reach the earth’s surface, but traps the heat radiated back from the earth’s surface towards space.

The heating effect is analogous to the manner in which the glass of a greenhouse traps the sun’s radiation to warm the air inside the greenhouse.

CLIMATE CHANGECLIMATE CHANGE

Greenhouse gas …Greenhouse gas …

Why are greenhouse gases dangerous for the environment?Why are greenhouse gases dangerous for the environment?

Heat

Sun’s radiation

CO2 freely allows radiation from the sun, but traps heat radiated back from the earth’s surface.

Vegetation subtracts CO2 from the atmosphere, reducing damage by greenhouse gases.

CLIMATE CHANGECLIMATE CHANGE

Greenhouse gasGreenhouse gas

Solutions?

Hydrogen has the best potential of becoming the fuel of the future.

CLIMATE CHANGE

Hydrogen is a carbon-free energy carrier. When used in fuel cells, there are no harmful emissions.

The use of energy may lead to climate changes. It is thus necessary to make the transition to cleaner and environmentally favourable energy carriers.

Hydrogen can be produced from sustainable, renewable sourcesand may contribute to meeting the growth in world energy demand.