Unit 1

Energy Matters

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• Reaction Rates• Enthalpy changes• Patterns in the Periodic Table• Bonding, Structure and Properties• The Mole• Click here to end

Reaction Rates

Energy changes

• Exothermic changes cause heat to be released to the surroundings.

• Endothermic changes cause absorption of heat from the surroundings.

• A potential energy diagram can be used to show the energy pathway for a reaction.

Energy Diagrams

• We can represent what happens in a chemical reaction using an energy diagram.

• The energy of the reactants and products is shown.

Energy

Reactants

Products

Activated complex

• The enthalpy (energy change) for the reaction can be calculated

H = HR – HP

• For an exothermic reaction H is negative.

Reactants

Products

Energy

Activated complex

H

• The enthalpy (energy change) for the reaction can be calculated

H = HR – HP

• For an endothermic reaction H is positive.

Energy

Reactants

Products

Activated complex

H

• An activated complex is formed.

• This is an unstable collection of atoms, intermediate between reactants and products.

Energy

Reactants

Products

Activated complex

• The energy which is needed to produce the activated complex is called the activation energy (EA).

Energy

EA

• A similar energy diagram can be drawn for an endothermic reaction.

Energy

EA

H

• A similar energy diagram can be drawn for an endothermic reaction.

Energy

EA

H

Following the course of a reaction

• Reactions can be followed by measuring how some quantity we can measure changes with time.

• Reactions can be followed by measuring changes in concentration, mass or volume of either the reactants or products

• The average rate of a reaction, or stage in a reaction, can be calculated by dividing the difference between the initial and final quantities by the time interval.Rate = change

time

Volumeof gas released(ml)

Time (s)

Volumeof gas released(ml)

Time (s)

V1

t1

Volumeof gas released(ml)

Time (s)

V1

t1

V2

t2

Volumeof gas released(ml)

Time (s)

V1

t1

V2

t2

Volume change (V)

V = V2 – V1

Volumeof gas released(ml)

Time (s)

V1

t1

V2

t2

Time change (t)

t = t2 – t1

Volumeof gas released(ml)

Time (s)

V1

t1

V2

t2

Average reaction ratebetween t1 and t2.

Rate = V/t

• The rate of a reaction, or stage in a reaction, is proportional to the reciprocal of the time taken.

• Rate proportional to /t

Factors affecting rate

• The rates of reactions are affected by changes in

• Concentration• Particle size• Temperature.

Collision Theory

• Reactions will only take place when the reacting particles collide.

• Reactions will only take place when the reacting particles collide.

• The particles need to collide at the correct angle.

• The particles need to collide at the correct angle.

• The particles need to collide at the correct angle.

• Collision theory explains the effect of concentration on reaction rates.

• The more particles there are in a given volume, the greater the chance of collision.

Concentration

Collision Theory

• Collision theory explains the effect surface area on reaction rates.

• Collisions can only take place on the surface.

• The larger the surface the more collisions.

Surface Area

Activation energy

Two molecules approach each other

If they don’t have the required activation energy nothing happens.

Two molecules approach each other

If they have the required activation energy the molecules form the

Activated complex

If they have the required activation energy the molecules form the

Activated complex

The activated complex splits apart To form the products.

Temperature

• Each molecule has a kinetic energy.

• Not all molecules in a material have the same kinetic energy.

• Temperature is a measure of the average kinetic energy of the molecules.

Energy Distribution

• Energy distribution diagrams show the numbers of molecules with each kinetic energy.

energy of molecules

number of molecules

• As the temperature increases the kinetic energy distribution changes.

energy of molecules

number of molecules

• At higher temperatures the average kinetic energy of the molecules is greater.

• More molecules will be able to provide the activation energy.

• The reaction will go faster.

Activation energy.

• In some chemical reactions light can be used to increase the number of particles with energy greater than the activation energy (e.g. photography)

• In other reactions shock can increase the number of particles with energy greater than the activation energy.

Excess

• When we carry out a chemical reaction we will usually use more of one of the reagents than is needed.

• That reagent is said to be in excess.

• We can calculate the reagent in excess using a mole equation.

• 1 g of carbon reacts with 10 g of copper(II) oxide. Show by calculation which reagent is in excess.C + 2CuO 2Cu + CO21 mole 2 moles 2 moles 1 mole12g 2x(63.5 + 16)12g 159g

1g 159/12=13.3g

• 1 g of carbon reacts with 10 g of copper(II) oxide. Show by calculation which reagent is in excess.C + 2CuO 2Cu + CO2

1g 13.3gSince 13.3g of copper(II) oxide are needed to react with 1g carbon, then carbon is the reagent in excess in the example given.

Catalysts

• Catalysts speed up reactions, without being changed by the reaction.

• Catalysts are used in many industrial processes.

• They reduce the temperature needed, so reducing energy costs.

Industrial Process Catalyst used

Haber Process Iron

Ostwald Process Platinum

Contact Process Vanadium(V) oxide

Catalytic Cracking Aluminium oxide

Collision Theory

• Not all collisions are successful, because they need to have the appropriate activation energy.

• Activation energy is the energy needed to start the reaction.

Catalysts

• Catalysts work by providing an alternative pathway for the reaction.

• This pathway has a lower activation energy

Energy

• Catalysts work by providing an alternative pathway for the reaction.

• This pathway has a lower activation energy

Energy

• Heterogeneous catalysts are in a different state from the reactants they catalyse.

• Homogeneous catalysts are in the same state as the reactants they catalyse.

• Heterogeneous catalysts work by the adsorption of reactant molecules.

• The adsorption of the molecules loosens bonds and makes it easier for the substance to react.

• The surface activity of a catalyst can be reduced by poisoning, when surface sites are taken over by other substances, preventing reactants being adsorbed.

• Impurities in the reactants result in the industrial catalysts having to be regenerated or renewed.

Catalytic converters

• Catalytic convertors are fitted to cars to catalyse the conversion of poisonous carbon monoxide and oxides of nitrogen to carbon dioxide and nitrogen.

• Cars with catalytic converters only use ‘lead-free’ petrol to prevent poisoning of the catalyst.

Enzymes

• Enzymes catalyse the chemical reactions which take place in the living cells of plants and animals.

• Enzymes are used in many industrial processes.

Reaction Rates

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Enthalpy changes

Enthalpy changes

• An enthalpy change is the energy produced or released in a chemical reaction.

• Enthalpy change is given the symbol H.

• Enthalpy change is measured in kilojoules per mole (Kj/mol)

Enthalpy of combustion

• The enthalpy of combustion of a substance is the enthalpy change when one mole of the substance burns completely in oxygen.

Enthalpy of solution

• The enthalpy of solution of a substance is the enthalpy change when one mole of the substance dissolves in water.

Enthalpy of neutralisation

• The enthalpy of neutralisation of an acid is the enthalpy change when the acid is neutralised to form one mole of water.

Calculating enthalpy changes

• To calculate enthalpy changes we use the equation:H = cmTH is enthalpy changec is (4.18 Kj/mol)m is number of kilograms (litres) of water

T is the temperature change

Enthalpy changes

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Patterns in the Periodic Table

Patterns in the Periodic Table

• The modern Periodic Table is based on the work of Mendeleev.

• Mendeleev arranged the known elements in order of increasing atomic masses.

• He combined this putting elements with similar chemical properties in the same vertical group.

• He left gaps for elements that had not been discovered at that time.

• Certain physical properties show trends, which repeat from one period of the periodic table to the next.

Density and hardness

• Elements at the left side of the Periodic Table indicate a gradual increase in these properties

• From Group 1 elements to Group 4 elements. Then there is a gradual decrease to low values at Group 7 and Group 0 (mostly gases).

Melting and boiling points

• The melting point and boiling point of an element gives an indication of the size of the forces that hold together the atoms or molecules.

• The higher the melting and boiling point the stronger the forces.

Boiling points

Melting points

Atomic size

• Covalent radius is half the distance between the centres of two covalently bonded atoms.

• We use it as a measure of atomic size.

Covalent radius

• We can see that covalent radius decreases as we move from left to right across a period.

• This is because the nuclear charge is increasing.

• Attraction to outer electrons increases.

• Thus the atom is smaller

• We can see that covalent radius decreases as we move from top to bottom down a period.

• This is because atoms have more shells of electrons.

• Nuclear charge is shielded by the inner electron shells.

• Thus the atom is bigger

Ionisation enthalpy

• First ionisation enthalpy (energy) is the energy to remove one mole of electrons from one mole of free, gaseous atoms.

X(g) X+(g) + e

Ionisation enthalpies of the first 20 elements

Ionisation enthalpy

• If we plot ionisation enthalpy against atomic number we can see two things.

• Ionisation enthalpy increases as we move from left to right across a period.

• Ionisation enthalpy decreases as you move down a group.

• As we move across a period from left to rightAtomic size decreasesCharge on the nucleus increasesThe force of attraction on the outer electrons is greaterMore energy is needed to remove an electron

• As we move down a groupAtomic size increasesExtra layers of electrons help shield the outer electrons from nuclear attraction.The force of attraction on the outer electrons is lessLess energy is needed to remove an electron

• Second ionisation enthalpy (energy) is the energy to remove a second mole of electrons.

X+(g) X2+(g) + e

• If we look at successive ionisation enthalpies for the same element, we see that there is a point where there is a large increase.

• This comes when the next electron has to be taken from an inside shell, nearer the nucleus where the electron is more strongly held.

Element 1st I.E.(kJmol-1)

2nd I.E.(kJmol-1)

3rd I.E.(kJmol-1)

4th I.E.(kJmol-1)

Na 502 4560 6920 9540

Mg 744 1460 7750 10500

Al 584 1830 2760 11600

Electronegativity

• Atoms of different elements have different attractions for bonding electrons.

• Pauling worked out values for electronegativity.

• Electronegativity is a measure of the attraction an atom involved in a bond has for the electrons of the bond.

• Electronegativity values increase across a period and decrease down a group.

Patterns in the Periodic Table

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Bonding, structure and

properties

Bonding

• Bonds are forces which hold particles together.

• There are two main kinds of forces

• Intramolecular forces

• Intermolecular forces

Intramolecular forces

• Intramolecular forces are the forces which hold atoms together.

• Covalent bonds are typical intramolecular forces.

Intermolecular forces

• Intermolecular forces are the forces which exist between molecules.

• Van der Waal’s forces, dipole-dipole attractions and hydrogen bonds are typical intermolecular forces.

Metallic bonding

• Metals can lose their outer shell electrons to gain a stable electron arrangement.

• Metal atoms are arranged in a lattice and can also delocalise their outer shell electrons, allowing them to move freely between the atoms in the lattice.

• This brings the metals closer to obtaining a stable electron arrangement.

• Since the metal atoms become positively charged ions they attract the free moving electrons in the lattice.

• This attraction forms a metallic bond which is very strong.

Covalent bonding

• As with all bond formations, the atoms must first collide with one another.

• When some atoms collide with each other, the electrons in the outer shell can be shared between the atoms.

• The electrons of the two atoms are both negatively charged and repel each other.

• When a collision takes place with sufficient energy to form a compound, the outer energy levels overlap and the atoms share the electrons.

• The overlap area has an increase in negative charge, which is strongly attracted by the positive nuclei of both atoms.

• This draws the two atoms closely together.

• The electrostatic force of attraction between the nuclei and the shared electrons forms a strong covalent bond.

e-

e- ++

Polar covalent bonding

• Sometimes on atom has a greater force of attraction than the other.

• This leads to polar covalent bonding, where there are slight charges (shown by and ) on the atoms.

e-

e- ++

What kind of bonding?

• The type of bond depends on the difference in electronegativity between the bonded atoms.

Covalent 0

Borderline 1.5

Ionic 3.0

Ionic bonding

• An ionic bond usually occurs between a metal and a non-metal and involves ions, which are charged atoms (or groups of atoms).

• In ionic bonding, electrons are transferred from one atom to another allowing both atoms to achieve a stable electron arrangement.

• For example, sodium and chlorine atoms would form an ionic bond making the compound sodium chloride as shown below :

van der Waal’s forces

• van der Waals’ forces are forces of attraction which can operate between all atoms and molecules.

• van der Waals’ forces are much weaker than all other types of bonding.

• Random movement of electrons in an atom can cause a temporary imbalance in charge.

• The force of attraction between two such atoms is called van der Waal’s force.

• Thus van der Waals’ forces are a result of electrostatic attraction between temporary dipoles, caused by movement of electrons in atoms and molecules.

• The strength of van der Waals’ forces is related to the size of the atoms or molecules.

Polar molecules

• A molecule is described as polar if it has a permanent dipole e.g. hydrogen chloride

H

Cl

• A molecule with polar bonds will not be polar if it is arranged symmetrically e.g.

Cl

Cl C Cl

Cltetrachloromethane non-polar

Cl

Cl C Cl

Htrichloromethane polar

• The dipoles of adjacent molecules will attract each other, as shown opposite.

• Permanent dipole-permanent dipole interactions are additional electrostatic forces of attraction between polar molecules.

• Permanent dipole-permanent dipole interactions are stronger than van der Waals’ forces for molecules of equivalent size.

• The melting and boiling points of polar substances are higher than the melting and boiling points of non-polar substances with similar molecular sizes.

• This is due to the extra forces between molecules.

Solvents

• Ionic compounds and polar molecular compounds tend to be soluble in polar solvents such as water and insoluble in non-polar solvents.

• Non-polar molecular substances tend to be soluble in non-polar solvents and insoluble in polar solvents.

Hydrogen bonds

• Hydrogen bonds are electrostatic forces of attraction between molecules which contain these highly polar bonds.

• A hydrogen bond is stronger than other forms of permanent dipole-permanent dipole interaction but weaker than a covalent bond.

• Bonds consisting of a hydrogen atom bonded to an atom of a strongly electronegative element such as fluorine, oxygen or nitrogen are highly polar.

• The anomalous boiling points of ammonia, water and hydrogen fluoride are a result of hydrogen bonding.

• Boiling points, melting points, viscosity and miscibility in water are properties of substances which are affected by hydrogen bonding.

• Hydrogen bonding between molecules in ice results in an expanded structure which causes the density of ice to be less than that of water at low temperatures.

Structure

• Different materials have different structures

• Metallic• Ionic• Covalent molecular• Covalent network• Monatomic

Metallic Structure

• A metallic structure consists of a giant lattice of positively charged ions and delocalised outer electrons.

Ionic Structure

• An ionic structure consists of a giant lattice of oppositely charged ions.

• This structure only is found in compounds.

Covalent Molecular Structure

• A covalent molecular structure consists of discrete molecules held together by weak intermolecular forces.

• This can be found in elements (Cl2) or compounds (CH4)

Covalent Network Structure

• A covalent network structure consists of a giant lattice of covalently bonded atoms.

• This structure is found in elements (C(diamond)) or compounds (SiO2 or SiC)

Monatomic Structure

• A monatomic structure consists of discrete atoms held together by van der Waals’ forces.

• This is found in the Noble Gases.

Structure

• The first 20 elements in the Periodic Table can be categorised according to bonding and structure.

Li, Be, Na, Mg, Al, K, Ca have metallic structure

 H2, N2, O2, F2, Cl2, P4, S8 and C (fullerenes) have covalent molecular structure

B, C (diamond, graphite), Si) have covalent network structure

He, Ne and Ar have monatomic structure

Bonding, structure and properties

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The mole

The Mole

• One mole of any substance contains the formula mass in grams.

massformulagram

gmassmoles

)(

• A 1 mole per litre (mol/l) solution contains one mole in each litre of solution.

)(lvolume

molesofnumberionconcentrat

• Each material is made up of different types of particles: atoms, molecules, ions.

• The basic blocks of each substance are called formula units.

Formula UnitsSubstance Example Formula

Units

Atomic element

copper Cu atoms

Molecular element

chlorine Cl2 molecules

Covalent compound

methane CH4 molecules

Ionic compound

calcium chloride

1xCa2+ ions 2xCl- ions

The Avogadro Constant

• One mole of any substance contains 6.02×1023 formula units.

• This number is called the Avogadro Constant

• Equimolar amounts of substances contain equal numbers of formula units.

• How many atoms in 2g of calcium?

• Calcium is Ca, formula mass 40.• 2g is 2/40 = 0.05 moles• Number of atoms = 0.05 x 6 x

1023

= 3 x 1022

• How many ions in 10 of calcium carbonate?

• Calcium carbonate is CaCO3, formula mass 100.

• 10g is 10/100 = 0.1 moles• Number of formula units

= 0.1x6x1023 = 6x1022

• Each formula unit has 2 ions.Number of ions = 2 x 6x1022 = 1.2x1023

Molar Volume

• The molar volume (in units of l mol-1) is the same for all gases at the same temperature and pressure.

• This is because the sizes of molecules is insignificant compared to the distances between them.

• The volume of a gas can be calculated from the number of moles and vice versa.

• How many atoms are in 2.5 litres of hydrogen gas? (molar volume = 25l)

• 2.5 litres =2.5/25 = 0.1 moles• Number of H2 molecules =

0.1x6x1023 =6x1022

• Number of atoms =2x6x1022 =1.2x1023

Reacting Volumes

• Since equal volumes of any gas contain equal numbers of molecules then we can relate mole numbers to molecules.CH4 + 2O2 CO2 + 2H20

1 mole 2 moles 1mole 2 moles1 vol 2 vols 1 vol 2 vols

Reacting Volumes

• What volume of oxygen will be needed to burn 20 ml of propane? C3H8 + 5O2 3CO2 + 4H20

1 mole 5 moles 3 moles 4 moles1 vol 5 vols 3 vol 4 vols 20 ml 100 ml

The Mole

• The mole is the key to all our calculations.

• The next slide shows how the mole is related to all the other quantities we need.

• It also includes a quantity we will not meet until later.

1 mole

1 gramformulamass

1 litre of1 mol/l solution

6 x 1023

formula units

1 molarvolume

n x 96500 coulombs

The Mole

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The End

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