USING AND CONTROLLING REACTIONS

Measuring energy changes

Chemical reactions are usually accompanied by either the absorption of energy (heat) or the release of energy (heat, light, electricity, sound). Reactions that produce heat energy include acid – base neutralisation, combustion reactions, the dissolving of some substances and respiration. Combustion reactions usually produce light and heat. Changes that absorb energy include melting, evaporation, and the dissolving of some substances. Reactions that absorb electromagnetic radiation are called photochemical reactions. Photochemical reactions include the reduction of silver ions in black and white photography. Photosynthesis, smog formation and ozone destruction are other examples. Exothermic reactions release heat to the surrounding, therefore the temperature of the surroundings increases. Endothermic reactions absorb heat from their surroundings, therefore the temperature of the surroundings decreases.

Enthalpy Changes for Reactions

The quantity of heat energy released or absorbed when specified amounts of substances react is called the heat of reaction. For reactions carried out at constant pressure, the heat of reaction is called the enthalpy change

(H). The enthalpy change is measured in joules (J) or Kilojoules (kJ) Commonly used is the Molar Enthalpy Change, when it is the energy released or absorbed when 1 mole of a substance fully reacts in a chemical reaction under constant pressure. (units = kJ. Mo1-1)

For Endothermic Reactions H > 0 (+ve values)

For Exothermic Reactions H < 0 ( -ve values)

Thermochemical Equations

Thermochemical equations contain; The mole ratio of reactants The state of each reactant and product The quantity of heat energy released or absorbed per mole

e.g. C2H5OH(l) + O2(g) CO2(g) + H2O(g) ΔH = -1366.8 kJ/mol

MOLAR ENTHALPY OF COMBUSTION

- Is the quantity of heat energy released when 1 mole of a pure element or compound is burnt completely in oxygen under constant pressure. For the complete combustion of hydrocarbons and organic compounds consisting of C, H and O only. The products are gaseous CO2 and water only.

e.g.

Molar Enthalpy of Solution - Quantity of heat energy released or absorbed when 1 mol. Of the substance dissolves in sufficient solvent so that further dilution causes no further release or absorption of heat energy e.g.

Molar Enthalpy of Neutralisation

- Quantity of hear energy released when 1 mol. of hydrogen ions is transferred from an acid to a base in an acid – base reaction occurring in aqueous solution.

e.g.

CALCULATING H VALUES FROM EXPERIMENTAL DATA

The measurement of heat energy changes during a chemical reaction is called calorimetry. The apparatus used for the measurement is called a calorimeter, which is an insulated vessel. The simplest is a polystyrene cup. In a calorimetric determination, the heat absorbed or released by a reaction is absorbed from or by a measured mass of water and can be calculated from the following

H = Specific heat capacity (Cp) of H2O x Temp x volume of water (in mL)

energy released from a chemical reaction (units = Joules)

Or

molar enthalpy (units = kJ mol-1)

Cp of water is 4.18 J.g-1 oC-1 It is the heat required to raise 1 mL of water by 1oC

Assumptions applied to Calorimetric calculations.

1) All heat energy produced or absorbed by the reaction is transferred to the known mass of water. There is no exchange of heat to the calorimeter

2) The reaction occurs quickly so the maximum or minimum temperature of the liquid in the calorimeter to be reached before the liquid begins to return to room temperature.

Approximations

1) The Cp of any aqueous solution is taken to be the same as water. Vis 4.18 J.g-1 oC-1

2) The density of any liquid is the same as the H2O, 1.0g. mL-1, therefore

Mass of solution = volume in mL.

Experimental results

When trying to obtain molar enthalpy of combustion values using calorimetry, the accuracy is very low, when the results are compared with literature values. The low accuracy can be attributed to the following Systematic Errors;

1) all the heat energy released is not absorbed by the water in the calorimeter. It is absorbed by air, the water container.

2) The combustion of alcohols is done in air, therefore combustion is not complete 3) The combustion is not done at constant pressure, which is a requirement 4) Some alcohol evaporates from the wick of a spirit burner.

FUELS

Carbon based fuels

Common carbon based fuels are crude oil, natural gas, coal and oil shale. They have been formed over millions of years by anaerobic breakdown of the remains of plants and animals buried in the ground. They are mainly organic compounds with small amounts of inorganic compounds. In industry, heat energy is obtained from coal, natural gas and oil by burning them in air. Crude oil must be refined into petrol and diesel fuels by chemical processes and physical separation processes in a refinery Carbon based fuels are also a feedstock for the chemical industry to produce products such as synthetic fibres, plastics, paints, solvents, detergents, dyes and pesticides. When the fuels are used for burning they cannot be used to make other products.

Advantages and Disadvantages of Carbon based fuels The advantages of carbon based fuels are: - High enthalpy of combustion, high energy density - They can be burnt where they are needed (eg internal combustion engine) - They are cheap and have many uses - Provide considerable revenue for governments through tax Advantages of specific carbon based fuels include:

Coal – large reserves, cheap to mine. Crude oil – large reserves, easily extracted and transported. Produces a variety of

fuels. Natural Gas – easily extracted, purified and distributed, burns cleanly, Produces lots of energy

The disadvantages of carbon based fuels are; - They are non-renewable, all oil and gas will be depleted before 2100 all coal by 2500. - The pollution created by their extraction, refinement and use is damaging to the our health

and that of the planet. The burning of carbon based fuels also produces huge quantities of carbon dioxide, which contributes to the enhanced greenhouse effect.

- They add to the atmosphere that was not originally there unlike alcohol and wood fuels which put back the CO2 that they took out.

- We depend on these fuels for our lifestyle, life would be made more difficult without them. Disadvantages of specific carbon based fuels include:

Coal - non-renewable, can contain sulphur compounds – combustion produces SO2, NOx, soot, fine ash particles, water.

Crude Oil – non-renewable, contains sulphur compounds – combustion products include SO2, NOx, soot, unburnt hydrocarbons, water.

Natural Gas – non-renewable, combustion products, NOx and water.

Complete Combustion of Fuels

Complete combustion of fuels can only occur when there is unlimited quantities of oxygen. The products of complete combustion are water and carbon dioxide. e.g .

The products of incomplete combustion include carbon (soot), carbon monoxide and unburnt hydro-carbons e.g.

Note that less energy is released when incomplete combustion is allowed to occur.

Harmful Effects of the Products of Incomplete Combustion

Carbon Monoxide

Incomplete combustion of petrol in cars is the major source of carbon monoxide (CO) pollution in the atmosphere of cities CO levels can get as high as 50 to 100ppm at peak traffic times. In the body, CO displaces oxygen from oxyhaemoglobin in the blood to form carboxyhaemoglobin. As oxyhaemoglobin transfers oxygen in the blood, the displacement deprives the body of oxygen. This can lead to impairment of judgement, headaches, dizziness, fatigue, loss of consciousness and even death.

Soot Soot is microcrystalline carbon. It looks like black smoke, with particles being 20 to 300um. Soot can cause a visual pollution problem as it blackens objects. It can also clog up air and fuel inlets in burners, which can lower their efficiency. On plants it will block out light. For humans, soot can be harmful to the respiratory system. While large soot particles can be trapped by cilia in the respiratory tract, smaller particles accumulate in the lung sacs where they can remain for several years. This can inhibit mucous flow, which removes harmful substances. Soot particles may also carry toxic substances on their surface when breathed in. bronchitis, asthma and emphysema are conditions that can be caused or aggravated by long term exposure to soot. Unburnt hydrocarbons have already been discussed.

Heat energy from the Combustion of fuels

The quantity of heat produced by combustion is usually quoted as kJ.mol-1. It can also be quoted as energy density (kJ.g-1), or for liquid fuels kJ.L-1. For fuel mixes it is not possible to get molar enthalpy of combustion values. These different units can be interconverted by using the following expressions

Note that energy density values are similar for fuels. Energy density values are better for comparing fuel efficiencies. Example: Comparing two liquid fuels ethanol and octane, the heat given out per mole on

complete combustion is

Octane has the higher energy density per mole, but liquid fuels are purchased by the litre. So how does the energy density compare litre for litre? First convert the energy to energy per gram by dividing by molar mass

The density of ethanol is 0.785 g mL-1 at 25oC. Energy per litre will be

Can you calculate the energy per litre for octane given its density at 25oC is 0.698 g mL-1?

ELECTROCHEMISTRY

Electrochemical cells

An electrochemical cell comprises the following components:

- two electrodes, one an anode (+ve), the other the cathode (-ve). They are made of either a reactive material or inert material such as platinum or graphite.

- an electrolyte, either a molten ionic compound or an aqueous solution containing free

ions.

- a metal wire connecting the two electrodes

- an oxidising agent in contact with the cathode and a reducing agent in contact with the anode

There are two types of Electrochemical Cells. One Galvanic Cells

- Galvanic cells use redox reactions to produce a direct electric current, eg batteries

Electrolytic Cells - Use and external electrical source (D.C.) to cause an non-spontaneous chemical

reaction (electrolysis - redox reactions)

Redox Reactions

A redox reaction involves an oxidising agent reacting with a reducing agent. In the reaction, electrons are transferred from the reducing agent to the oxidising agent Redox reactions can be considered as two half-reactions, one involving oxidation and the other reduction. These half reactions occur simultaneously. It is like lending and borrowing, you can’t have one without the other and they happen at the same time.

Writing Redox Equations for Reactions in Neutral of Acidic Solutions

The procedure for writing redox reactions involves finding an oxidation half reaction, a reduction half equation, summing these and balancing. The steps are shown below: Sulphur dioxide gas is bubbled through an acidified solution containing dichromate ions; the latter is reduced to chromium (III) ions and the sulphur dioxide is oxidised to sulphate ions. The equation is developed as follows: 1) Identification of redox pairs.

2) Balance any element other than hydrogen or oxygen.

3) Balance the oxygen atoms by adding an appropriate number of water molecules.

4) Balance the hydrogen atoms by adding an appropriate number of hydrogen ions to the

left or right.

5) Balance the electrical charge by adding electrons to the left or right.

6) Combine the two half reaction equations to produce an overall:

a) If the electron numbers are not equal then half equations must be multiplied to make them equal.

b) two half equations are then added. Electrons, water and hydrogen ions on opposite sides are cancelled out.

7) Check that all atoms and charges balance.

Galvanic Cells

Galvanic cells are electrochemical cells which produce electrical energy from spontaneous redox reactions. In these cells the reducing agent and oxidising agent are not in contact with each other, they are connected via a metal conducting wire and a salt bridge. A galvanic cell is made of two half cells 1) An oxidation half cell, where the electrode is the anode. 2) A reduction half cell, where the electrode is the cathode.

The two half cells are connected by: 1) An external metal wire where electrons travel from the anode to the cathode 2) A salt bridge connects the solutions. Positive ions move toward the reduction half cell,

negative ions move toward the oxidation half cell.

Half cells can be classified as metal half cells or solution half cells.

Metal Half Cells

A metal half cell consists of a solid metal electrode in contact with a solution containing ions of the metal. Eg. Cu rod standing in a solution of copper sulphate. A rod of zinc standing in zinc sulphate solution. These two half cells can be joined together to make a galvanic cell The shorthand representation of this cell is

When a cell composed of two metal half cells is operating, the more reactive of the two metals is oxidised. In this case it is zinc

the electrons flow through the wire to the cathode where they are accepted by copper ions in the solution. To make Cu(s) on the surface of the electrode

the overall reaction

The salt bridge consists of a concentrated solution of a salt, the ions of which should not react with the ions present in the solution in the half cells. The ions should not be easily oxidised or reduced. NaCl is a good example as Na+ cannot be reduced and Cl- can only be oxidised in high concentrations. Neither ion forms a precipitate in solution. Using your knowledge of the reactivity series it is possible to predict reactions in metal half cell galvanic cells the more reactive metal will be oxidised to its ions the ions of the less reactive metal will be reduced to a metal the more reactive metal will be the anode (less = cathode) electrons flow from anode to cathode +ve ions toward cathode, -ve ions toward anode

Solution Half-Cells

Galvanic cells constructed with inert electrodes in contact with oxidising and reducing agents as solutions can make half cells.

Common Oxidisers Common Reducing Agents Acidified permanganate ions sulfite ions Iron(lll) ions iodide ions Hydrogen peroxide The inert electrodes are commonly platinum or graphite

By convention, when these cells are constructed, the oxidiser is dissolved in solution with its reduced form and acid if required for the half reaction. Eg. Permanganate ions (oxidiser) would be present in solution with manganese (ll) ions (reduced form) with sulphuric acid (provides H+ ions)

the reducer is placed in solution with its oxidised from. Eg. Fe(ll) ions would be placed in Fe(lll) ions

these cells would function without the oxidised and reduced forms in solution with the reducer and oxidiser respectively. It is included as a matter of convention

FUEL CELLS

Fuel cells are special galvanic cells In a typical fuel cell, gaseous fuels are fed continuously to an anode (negative electrode) compartment and an oxidant (such as oxygen from the air) is fed continuously to the cathode (positive electrode) compartment. While a fuel cell has components and characteristics similar to normal galvanic cells, a fuel cell differs in several respects. The fuel cell is an energy conversion device that theoretically has the capacity to produce an electric current for as long as the fuel and oxidant are supplied to the electrodes. In real life however, corrosion or malfunction of components limits the operating life of fuel cells. Fuel cell modifications include porous electrodes, Impregnated with catalysts to increase the rate of reaction. Porosity

increases surface area. the electrolyte (salt bridge) liquid is usually absorbed into a porous solis matrix that is

between the electrodes. Electrons can’t go through the electrolyte, ions like OH-, H+, CO3

-2 can.

Fuel is continuously supplied to the anode (usually H2) where it is oxidised by losing electrons. At the same time, an oxidant (usually oxygen) is continually supplied to the cathode where it is reduced

The anode (negative) is connected externally to the cathode (+ve) by means of a metal wire

Operating temperatures range from 50oC to 1000oC depending on the electrolyte. The higher temperatures increase the rate of the electrode reactions.

Anode

Cathode

therefore Overall cell reaction is

The electrolyte is KOH(aq), operating temperature is 50o – 200oC Total output of the cell is electrical energy, heat energy and water. In theory any fuel-oxidant combination could be used. In practice only a few combinations have been successful.

ADVANTAGES AND DISADVANTAGES OF FUEL CELLS

Fuel cells have several advantages over normal galvanic cells and over electricity generators that use steam driven turbines.

They continuously produce electric current so long as fuel is continuously supplied

unlike other galvanic cells where the amount of chemicals in the cell determine the lifetime

They have high operating efficiency (70 – 80%) They offer a better mass to power output ratio than normal galvanic cells They use readily available fuels and oxidants They don’t produce pollutant gases such as SO2 and NOx Electrodes and electrolytes are not consumed Electrode reaction products are removed as they are formed and don’t remain in the

cell They require little maintenance They are silent during operation (no moving parts) In different versions they can be used for small to large scale applications, as well as

making mobile power generators Disadvantages include Impurities in the fuel or oxidant can poison the electrode catalysts or contaminate the

electrolyte. Eg. CO2 in the oxidant can react with KOH electrolyte forming K2CO3(s) that can clog the porous electrodes. This restricts air being used as an oxidant

High purity fuels and oxidants are expensive

Medium to high temperatures are needed for cells to function most efficiently

Metal electrodes catalysts such as platinum are expensive

Some of the electrolytes are very corrosive

Rechargeable Galvanic Cells

Rechargeable galvanic cells are also called storage cells. During discharge of the cells, electrode reactions produce an electric current in an external circuit. At the same time the oxidiser and the reducer inside the cell are both used up. During recharge, an external power source provides an electric current to the electrodes. The discharge reactions are reversed, and the original oxidiser and reducer are regenerated. The cell can now be discharged again. Well known examples include the lead-acid cell and the nickel-cadmium (NiCad)

The Lead-Acid Cell

When discharging the overall reaction

The charging reaction is

Note these reactions are complimentary and the materials are regenerated The Nickel-Cadmium Cell

Discharge This is reversed for charging

Electrolytic cells

An electrolytic cell is a electrochemical cell where electric current is used to drive a non-spontaneous redox reaction. This is called electrolysis. Electrolysis reactions are redox reactions and the exchange of electrons takes place at the electrodes (oxidation at the anode, reduction at the cathode). Important Industrial and Domestic Applications are Producing pure metals from their compounds (Zn, Al,Gpl & ll )

Recharging a lead-acid or nicad battery

Electroplating

Refining metals such as copper

Producing non-metal elements like chlorine and hydrogen and compounds like NaOH,

An electrolytic cell

During the operation of an electrolytic cell The power supply pumps electrons through the external circuit from the anode (+ve) to

the cathode (-ve)

In the electrolyte, the positive ions (cations) are attracted to the negatively charged cathode and the –ve ions to the positively charged anode

If the electrolyte is an aqueous solution, water will be in contact with the electrodes.

An oxidation half reaction occurs at the anode. For a molten electrolyte, the anions are oxidised. For an aqueous electrolyte, either the anions or the water molecules are oxidised, depending on their relative ease of oxidisation.

Reduction occurs at the cathode. For a molten electrolyte the cations are reduced. For

an aqueous electrolyte either the cations or the water molecules are reduced, depending on the relative ease of reduction

When using an aqueous electrolyte water will be;

Reduced when in preference to Al, Na, K

Oxidised when with SO42- and NO3

-

PRODUCTION OF HIGHLY REACTIVE METALS

The highly reactive metals from Gp l and ll can be done from the electrolysis of molten chlorides of the metals. To do this, inert graphite electrodes are used. This is used to make lithium, Na, Mg, Ca only

RATES OF REACTION

The rate of reaction is the speed at which a reaction occurs.

The rate of a chemical reaction is expressed either as the rate of formation of a product or the rate of consumption of a reactant. The rate of consumption or formation is measured as the change per unit time. The rate of reaction may describe the rate over a period of time (called the average rate of reaction), or the rate at a particular instant of time (called the instantaneous rate of reaction).

Average rate of reaction = change in quantity of reactant product Period of time elapsed Instantaneous rate of Reaction = gradient of the quantity of reactant or product Vs the graph at that particular time. Quantity – time graphs for chemical reactions Plots of concentration, mass or number of moles of a reactant or product against time produce curves of characteristic shape called exponential curves.

FACTORS AFFECTING THE RATES OF REACTIONS Factors can be explained in terms of the Collision theory of chemical reactions. The theory is based on the premise that for reactions to occur - The particles of the reactants must collide - the colliding particles must have sufficient energy to form an activated complex, from

which the final product is formed

Activation Energy of a reaction is the minimum energy required for reactant particles to form the activated complex Collisions of Reactant Particles can be; Productive Collisions – results in the formation of product

Non-Productive Collisions – not resulting in the formation of product

In terms of collision theory, reaction rate is dependent on - The magnitude of the activation energy of the reaction

- the frequency of collisions between reactant particles (temperature)

- the energy of the particle collisions relative to the activation energy

- the orientation of the colliding particles

Reaction rate is increased when AE is lowered, or if collision rate increases or the energy of colliding particles is increased.

FACTORS THAT AFFECT REACTION RATE

1. Concentration of Reactants Rate increases as concentration of reactants increase. This occurs because there are more particles of the reactants therefore more frequent the collisions.

2. Surface area of solid Reactants.

As the surface area of a solid reactant is increased, more particles of the reactant are exposed to collisions with other reactant particles, therefore more collisions, therefore greater rate of formation

3. Pressure of a Gaseous Reaction Mixture As the pressure of a reaction mixture increases, the concentration of reactants increase, therefore more collisions, therefore faster formation rate.

4. Intensity of incident light for photochemical reactions. Absorption of more intense radiation by reactant particles increases their energy. This increases the frequency of productive collisions, therefore faster rate of reaction.

5. Temperature of the Reaction Mixtures Rate increases as the temperature of the reaction mixture is increased For reactions to be productive (forms a product), the reactants must have enough Activation Energy. At a fixed temperature the reactant particles in a reaction mixture have a wide range of Kinetic Energy values

The distribution of collision energies changes as the temperature of the reaction mixture increases, however as the temperature increases, the value of the Activation Energy (EA) remains unchanged, this is because activation energy is independent of temperature.

therefore as the temperature increases, the number of collisions that have a sum of kinetic energies in excess of the activation energy increases. Lowering temperatures can also reduce reaction rates, this is why we cool foods, it reduces oxidisation. If a material has a higher energy the particles will also have higher kinetic energies and hence speed. This will also increase the rate of collisions and hence the reaction rate.

6. Catalysts

A catalyst is not consumed in the reaction. A catalyst provides an alternative pathway from reactant to products. The new pathway has a lower activation energy, therefore more collisions and faster reaction rate.

One of the effects of a catalyst is to absorb reactant particles onto its surface and hold them there in an orientation that allows for more productive collisions. This is especially important for large organic molecules where often there is only one part of the molecule in which a reaction is possible.

ENERGY PROFILE DIAGRAMS

In energy terms, the course of a chemical reaction can be represented as an energy profile diagram An energy profile diagram shows the relative potential energy values for the reactants, the products and the activation complex. In order to form the activation complex, the kinetic energy of the colliding particles must exceed the activation energy for the reaction. When the colliding particles have energy less than the activation energy they simply bounce off each other

Exothermic Reaction – reaction releases energy

Endothermic Reaction – reaction absorbs energy

CHEMICAL EQUILIBRIUM

Most chemical reactions are reversible. By convention, in a chemical equation, the reactants are on the left and the products on the right Reactants products This is called the forward reaction. When the products of a reaction react to form the original reactants, the following is used. Reactants products This is the back reaction When the forward and back reactions occur at the same time, the following conversion is used it is known as a reversible reaction Reactants products If a reversible chemical reaction is carried out

In a closed system (no loss or gain of reactants and products and no mass change)

at a constant temperature

then the concentrations of reactants and products eventually reach constant values or a balance. When this occurs the system is said to have reached a position of equilibrium. At equilibrium, the measurable quantities of the system are constant. Colour intensity, pressure and pH. Chemical Equilibrium is described as dynamic, as opposed to static, because, at the equilibrium position, the forward and back reactions are still occurring, but at equal rates. Chemical Equilibrium graphs are drawn as changes in concentration / time curves

THE EQUILIBRIUM CONSTANT

Consider the generalised reaction between A, B etc. to form products C, D etc.

aA + bB cC + dD at the Equilibrium position, the expression

Kc =

c d

a b

C D

A B (Equilibrium Law)

has a constant value at constant temperature. In this expression [C] represents the concentration at equilibrium in mol.L-1. This constant is called the equilibrium constant Kc. It has no units.

e.g.

The magnitude of the equilibrium constant, Kc, provides an indication of the yield of product relative to reactant, present at the equilibrium position. It indicates the extent of the reaction.

Large Kc values (>10) indicate

A high yield of products at the equilibrium position

That the reaction has proceeded to a large extent to the right by the time the equilibrium position is reached

Small Kc values (<0.1) indicates a low yield of products at the equilibrium position

Reaction has proceeded to a small extent to the right by the time equilibrium is

reached.

CALCULATIONS INVOLVING THE EQUILIBRIUM CONSTANT, Kc

A sample calculation involving Kc. To make H2SO4 in the contact process, SO2 needs to be converted to SO3.

2SO2(g) + O2 ⇌ 2SO3(g) for one mixture the concentrations at 7270 were determined by experiment to be

for this the Kc value can be found

given that there was no SO3 initially present, the concentration of O2 and SO2 can be calculated.

The final equilibrium concentrations for a given reaction depend on:

- Initial concentrations of the reactants and products - Temperature - Value of Kc - Pressure (for systems involving gases)

CHANGING THE EXTENT OF A REACTION - LE CHATELIER’S PRINCIPLE

Le Chatelier’s Principle is used to predict how changes to the reaction conditions affect the equilibrium position for a reaction.

Le Chatelier’s Principle is:

If an external change is made to the reaction conditions of a system at equilibrium so that it is no longer at equilibrium, a net reaction will occur (if possible) in the direction that counteracts the change; the internal change will oppose the external change.

Or simply if the equilibrium conditions are changed externally, then the system will attempt to counteract the change if possible.

Concentration Changes External Change System Response to the change Increasing the concentration Equilibrium position is shifted in of a reactant, at constant volume favour of consuming the added and temperature. reactant. Produces product. Equilibrium to the right. Decreasing the concentration Replaces the reactant, equilibrium of a reactant to the left Increase concentration of a Reaction moves the to consume product added product therefore equilibrium to the left Decrease concentration of a Reaction moves to replace product product therefore equilibrium to the right

Pressure Changes

External change Internal change Increasing the pressure of a Equilibrium position is shifted system at equilibrium by to produce a smaller number decreasing the volume of the of molecules in the gas phase. reaction vessel while temperature is constant. Decreasing the pressure of a Equilibrium position is shifted system, by increasing volume, to produce larger numbers of while temperature is constant molecules.

Pressure can also be increased by adding an inert gas.

Eg. N2O4(g) 2NO2(g) increased pressure decreased pressure

Temperature Changes

External change Internal change Increase Equilibrium shifts to endothermic direction Decrease Equilibrium shifts to exothermic direction

CATALYSTS AND EXTENT OF REACTION

While catalysts increase the rate of attainment of equilibria, they do not affect the equilibrium concentration, therefore no change to the Kc value.

CHEMICAL INDUSTRY

The Chemical Industry uses chemical reactions and physical processes on a large scale to convert raw materials into useful products. For a particular production process, the conditions of the chemical reactions are controlled to produce the best yield of product possible at an economic rate of formation. Conditions that give the highest yield may not give a fast rate of reaction and there may be a need for compromise between conflicting factors. Yield of product The yield of a product from a chemical reaction is the quantity of product obtained. the theoretical yield is the quantity of product predicted by the stoichiometry of the

chemical reaction.

the actual yield is the quantity of product actually obtained. Chemical companies want to maximise the actual yield. Ammonia Production by the Haber Process The Haber Process is the industrial process used throughout the world to produce ammonia. It involves a reaction between nitrogen and hydrogen gases at 400O C and 250 atmospheres of pressure, with an iron catalyst. The yield for this is 45% of the theoretical yield.

H = -46 kJ . mol-1

Using Le Chatelier’s Principle you would predict using high pressure, less molecules on the product side

low temperature, because the forward reaction is exothermic

The actual pressure used is high. The actual temperature used is moderately high. This is a compromise between yield of ammonia (greater at lower temperatures) and rate of formation (faster at higher temperatures) The iron catalyst increases the rate of formation of ammonia.

SULPHURIC ACID PRODUCTION BY THE CONTACT PROCESS

Sulphuric acid is produced via the Contact Process with several steps

or

(oleum)

The key step in this series of reactions is the production of SO3. Using Le Chatelier’s Principle a high yield can be predicted using high pressure and low temperature. The actual pressure is atmospheric pressure. High pressure is not used because the resulting increase in yield is not worth it. The actual temperature is moderately high. This is a compromise between yield of SO3 and rate of formation of SO3 (faster at higher temperatures). A vanadium pentoxide catalyst increases the rate of formation of sulphur trioxide. The conditions used in industry are those which provide the most economical yield of sulphur trioxide (in time and cost).

FLOW CHARTS

A flow diagram is often used to represent the movement of materials through the various components of the plant. The materials are classified as follows: Raw materials are substances converted by chemical or physical means into useful products. Eg. Coal, oil, natural gas, air, limestone, water. Waste products of chemical conversion processes are those for which there is no use or market. They can be dangerous

By-products are products of the chemical conversion process that are not the main product, but do have a commercial value or a use within the plant itself.

ENERGY COSTS

Energy use is a major cost item in most industrial processes, often outweighing the costs of raw materials and labour. Energy use must be kept to a minimum, as fuels are expensive and non-renewable. Energy costs can be minimised by using the following energy conversion measures. Transferring heat energy – heat energy released by an exothermic reaction can be

transferred to another part of the process by using heat exchangers. Heat can also be taken from hot waste gases before they are expelled.

Operating chemical processes at lower temperatures – using a catalyst allows a reaction to occur at lower temperature with the same rate.

Running a process continuously – if a process has to be regularly stopped and started, it can be more costly in terms of energy and time than a process which can be run continuously.

METAL PRODUCTION

THE REACTIVITY OF METALS

When metals react, they undergo oxidation – the metal atoms lose electrons to form metal ions. The more easily they lose electrons, the more reactive the metal. M → Mn+ + ne-

Conversely, the more easily a metal is oxidised, the less easily its ions are reduced to the metal.

THE METAL REACTIVITY SERIES

A more reactive metal will displace the ions of a less reactive metal from solution

THE NATURAL OCCURRENCE OF METALS

Most metals occur in the earth’s crust as compounds called minerals. The most unreactive metals such as gold, silver and platinum occur in the uncombined or native state. Less reactive metals are often found combined with other elements

METAL ORES

The production of metal ores may involve one or more of the following stages: Concentration of the mineral. This involves removal of most of the metal gangue, therefore

concentration of the metal mineral. This product is called a mineral concentrate. This stage is not required if the percentage of gangue in the ore is low.

Conversion of the mineral in the concentrate into a substance suitable for reduction. This is a chemical conversion process. One of the most common of these processes is the roasting of sulphide minerals to produce the metal oxide.

Reduction of metal compounds to the metal – can be done by using chemical reducing agents such as carbon monoxide or by electrolysis

Refining of the metal - involves the removal of trace quantities of impurities from the metal formed in the reduction stage.

For some formation of metals some stages may not be necessary.

Iron ore deposits mined in Australia have a high concentration (>65%) and the ore does not need further concentration.

Aluminium metal produced in electrolytic cells is of such purity

(>99%) that further refining is usually not necessary. Similarly zinc metal produced by electrolysis does not usually require refining. The percentage of zinc in the ore is usually so low that the ore needs concentrating before further processing.

THE ELECTROLYTIC PRODUCTION OF ZINC

Crushing and Grinding The ore is first crushed and ground. The main aim of this is to produce separate fine particles of the zinc sulphide mineral and the gangue. The ore is crushed in two stages to produce pieces about 2cm in size. Grinding reduces the zinc and gangue to a powder with particles 200 microns. The end result is a very large surface area. Mineral concentration by Froth Floatation The fine powder from the grinding mill is subjected to froth floatation. The powder is mixed with water in large floatation tanks or cells to form a slurry. Frothing agents, eg pine oil, are added with water soluble substances called “collectors”. The most common collectors are sodium or potassium xanthates such as sodium ethylxanathe and potassium isopropylxanthate.

Xanthate ions have similar structures to soap and detergent anions, with an ionic head and a non-polar tail represented as

Xanthate ions become attached to the grains containing zinc and sulphide ions, they are attracted by electrostatic force between the negatively charged heads and the Zn+2 ions. The grain acquires a non-polar (hydrophobic) surface as the tails point outwards.

Air is blown into the slurry to form hydrophobic air bubbles that are stabilised by the frothing agent. The hydrophobic mineral grains attach to the air bubbles, and to the surface as a froth. The froth is skimmed off and collected for subsequent processing. The other grains of gangue are not carried to the surface, but remain as a sludge at the bottom of the tank

The zinc concentrate obtained from floatation contains about 50% zinc sulfide. Roasting of the Zinc Sulphide Concentrate is roasted in air to get ZnO + SO2

Leaching of the Zinc Oxide Sulphur dioxide from the roasting stage is used to make H2SO4, which is then used to convert ZnO to ZnSO4(aq)

Zinc powder is added to the solution to displace any less reactive metal ion impurities such as Ag+, Cd+2, Cu+2.

The metals formed here create a residue that is collected and processed. Electrolysis of the Zinc Sulphate Solution - To do this, the reactions are as follows Anode – Water is oxidised to make O2 and H+ ions

Cathode – Zinc ions to zinc

Overall

PRODUCTION OF THE MORE REACTIVE METALS

Metals above zinc on the reactivity series can not be produced by electrolysis of aqueous solutions. The metals from aluminium and above on the metal reactivity series are produced industrially, either by the electrolysis of suitable compounds in the molten state or by using a reactive metal reducing agent at high temperatures.

DETAILS OF THE PRODUCTION OF ALUMINIUM BY ELECTROLYSIS

The reduction of aluminium ions to aluminium is not easy. The final stage in the industrial production of aluminium involves the electrolysis of molten alumina, Al2O3. This needs to be done as it is too electropositive to be done by other methods. It is too stable to be done by carbon reduction as it would require very high temperatures (around 2000oC) As the melting point of Alumina is 20300C, it is mixed with cryolite Na3AlF6 and other ionic compounds to form a mix with a melting point of less than 10000C. Even at this temperature it consumes a huge amount of energy. Electrolysis is carried out at 940 – 9800C in a carbon lined steel cathode vessel using carbon anodes. Fresh alumina is added continuously to replace the alumina consumed. Anode – Oxide ions → oxygen At this temperature, carbon anodes burn in the oxygen and are eaten away therefore must be replaced

Cathode - Aluminium ions → Aluminium

Overall

CHEMICAL REDUCTION OF METAL OXIDES TO METALS

Metals below Al on the reactivity series can be produced industrially by the reduction of the metal oxides with carbon. This is possible because they are more easily reduced. e.g.

ENERGY CONSUMED IN THE PRODUCTION OF METALS

The reduction stage usually consumes the greatest amount of energy and therefore is the most costly. For stages prior to and following the reduction, energy use is low. The most energy intensive reduction processes are those involving electrolysis of a molten, non aqueous electrolyte as used for Sodium and Aluminium production. Electrolysis of an aqueous solution uses considerably less energy.