SCIENCE FORM 4

CHAPTER 5: ENERGY AND CHEMICAL CHANGES5.1 PHYSICAL AND CHEMICAL CHANGES Many things undergo changes in our lives. There of two types of changes: (i) Physical change - Affects the physical properties of a substance, such as its size, shape and state. - Usually reversible. - No new substance is formed. Examples involving physical changes: (a) Melting of ice (b) Evaporation of water (c) Dissolving sugar in water (d) Crystallisation of sodium chloride from its saturated solution (e) Heating of iodine crystals (f) Heating of wax

(ii) Chemical change - Produces new substances that have properties different from those of the original substance. - Usually irreversible. - Examples involving chemical changes: (a) Burning of a paper (b) Color of a peeled apple changes when it is exposed to air (c) Frying an egg (d) Combustion of fuel (e) When iron filings are heated with sulphur powder, the mixture glows brightly. A black solid, iron sulphide is formed. Iron + Sulphur Iron sulphide

(f) When magnesium ribbon is heated, it burns with a bright flame. A white solid, magnesium oxide is formed. Magnesium + Oxygen Magnesium oxide

(g) Zinc reacts with blue copper sulphate solution to form a colorless zinc sulphate solution and brown copper. Zinc + Copper sulphate Zinc sulphate + Copper(h) Iron nail rust when water and oxygen are present. Rust is brown. Iron + Oxygen + Water Rust

(i) Heating green copper carbonate produces black copper oxide and releases carbon dioxide. Copper carbonate Copper oxide + Carbon dioxide(j) Mixing potassium iodide and lead nitrate solution forms lead iodide, a yellow precipitate. Potassium + Lead Potassium + Lead iodide nitrate nitrate iodide

Heating a mixture of iron filings and sulphur powder Reaction of zinc with copper sulphate solution

Aspect

Physical changes

Chemical changes

Formation of new substance

No

Reversibility

Reversible

Energy needed

Less energy needed

More energy needed

5.2 HEAT CHANGE IN CHEMICAL REACTIONS

There are two types of chemical reactions: (i) Exothermic reactions - Reactions which release heat to the surroundings are called exothermic reactions. - Temperature of the surroundings increases.

(ii) Endothermic reactions - Reactions which absorb heat from the surroundings. - Temperature of the surroundings decreases. During chemical reactions, old bonds in the reactants are broken and new bonds in the products are formed. The breaking down of old bonds absorbs heat energy whereas the formation of new bonds releases heat energy.

Breaking an old bond absorbs heat energy

Forming a new bond releases heat energy In exothermic reactions, the heat energy absorbed to break the old bonds in the reactants is less than the heat energy released when new bonds are formed in the products. There is a net of heat energy loss. The temperature of the surroundings increases due to the heat energy that is released.

In endothermic reactions, the heat energy absorbed to break the old bonds in the reactants is greater than the heat energy released when the new bonds are formed in the products. There is a net of heat energy gain. The temperature of the surroundings decreases due to heat energy is absorbed from the surroundings.

In exothermic reactions, the heat energy absorbed to break the old bonds in the reactants is less than the heat energy released when the new bonds are formed in the products

In endothermic reactions, the heat energy absorbed to break the old bonds in the reactants is greater than the heat energy released when the new bonds are formed in the products

Haber process is used to manufacture ammonia in the industry. Ammonia is an important material for making nitrogenous fertilisers. In the Haber process, nitrogen and hydrogen gases are mixed together. The mixture is passed over an iron catalyst.

Iron catalyst Nitrogen + Hydrogen Ammonia 450 C 200 atm

Haber process

Contact process is used to manufacture sulphuric acid in the industry. Step 1: Sulphur is burnt in air to produce sulphur dioxide. Sulphur + Oxygen Sulphur dioxide Step 2: A mixture of sulphur dioxide and air are passed over vanadium(V) oxide catalyst at 450 C to produce sulphur trioxide.

Vanadium(V) oxide catalystSulphur dioxide + Oxygen Sulphur trioxide 450 C

Step 3: Sulphur trioxide is dissolved in concentrated sulphuric acid to produce oleum.

Sulphur + Concentrated Oleum trioxide sulphuric acid

Step 4: The oleum is diluted with water to produce concentrated sulphuric acid.

Oleum + Water Concentrated sulphuric acid

Contact process

5.3 THE REACTIVITY SERIES OF METALS Why do gold, silver and platinum exist as elements in the Earths crust and whereas metals such as sodium and potassium do not exist as elements in the Earths crust? Some metals are more reactive than others. Different metals show different reactivity with water, acids and oxygen.

The reactions involving metals: (i) Reaction of reactive metal with water to produce alkali and hydrogen. Metal + Water Alkali + Hydrogen (ii) Reaction of reactive metal with dilute acid to produce a salt and hydrogen. Metal + Dilute acid Salt + Hydrogen (iii) Reaction of metal with oxygen to form metal oxide Metal + Oxygen Metal oxide

Based on their reactivity with oxygen, metals and carbon (non-metal) can be arranged in a reactivity series as follows: Potassium Sodium Calcium Magnesium Aluminium Carbon Zinc Iron Tin Lead Copper Silver Gold

The position of carbon in the reactivity series is determined by comparing its attraction for oxygen with other metals. When carbon is heated with a metal oxide, carbon can remove oxygen from the metal oxide if carbon is more reactive than the metal. When zinc oxide and iron oxide are heated with carbon, the oxygen in the metal oxides are removed by carbon. Therefore, carbon is more reactive than zinc and iron. When carbon is heated with aluminium oxide, the carbon cannot remove oxygen from aluminium oxide. This means that carbon is more reactive than zinc and iron but less reactive than aluminium. Therefore, carbon is positioned between aluminium and zinc in the reactivity series.

5.4 APPLICATION OF REACTIVITY SERIES OF METALS Most metals in the Earths crust are reactive. They react readily with other elements to form compounds such as oxides, sulphides and carbonates. These compounds are called ores.

Ore

Composition

Bauxite

Aluminium oxide

Cassiterite

Tin(IV) oxide

Haematite

Iron(III) oxide

Sphaletite

Zinc sulphide

Sedimentary rock

Calcium carbonate

The method of extracting a metal from its ore depends on its position in the reactivity series. Metals below carbon in the reactivity series can be extracted from their oxides with using carbon because they are less reactive than carbon. Carbon is used for extraction because it is cheap and readily available. Metals above carbon in the reactivity series cannot be extracted using carbon because they are more reactive than carbon. They are extracted using electrolysis. Tin ore or cassiterite is tin(IV) oxide. Tin is extracted by heating cassiterite with carbon (coke) and limestone at high temperature in a blast furnace. Carbon removes the oxygen from tin(IV) oxide because it is more reactive than tin. Tin(IV) oxide + Carbon Tin + Carbon dioxide Limestone is added to react with the impurities in the tin ore. Slag is formed. Two products are collected at the bottom of the blast furnace. Molten slag floats on the molten tin. This allow the two products to flow separately.

Extraction of tin in a blast furnace

5.5 ELECTROLYSIS Electrolysis is the decomposition of an electrolyte by electricity. An electrolyte is a liquid or solution which contains free-moving ions that can conduct electricity. The electrodes are conductors which carry electricity into or out of an electrolyte. The electrode joined to the positive terminal of the dry cell is called the anode whereas the electrode joined to the negative terminal is called the cathode. During electrolysis, the positively-charged ions (cations) are attracted to the cathode and receive electrons. The negatively-charged ions (anions) are attracted to the anode and release electrons.

Electrolysis

Refer to the diagram above: In the complete circuit, the bulb lights up. This is because molten lead bromide conducts electricity. At the anode, the bromide ions release electrons to become bromine atoms. The bromine atoms combine together to form bromine gas (brown colour). At the cathode, lead ions receive electrons to become lead atoms. Lead metal is formed (grey solid).

Uses of electrolysis in industry: (i) Electroplating - To prevent iron objects from corrosion, they are electroplated with a thin layer of unreactive metals such as copper, silver and chromium. - Electroplating make these objects resistant to corrosion and more attractive. - At the anode, the copper dissolves to form positively-charged copper ions. - At the cathode, the copper ions receive electrons to form a coat of copper on the iron spoon.

A rheostat is used to control the current flow in the circuit so that a small current is used and the object to be electroplated must be cleaned with a sandpaper before electrolysis. These steps are to obtain good results.

Iron spoon electroplated with copper (ii) Extraction of reactive metals - Metals which are more reactive than carbon are extracted from their ores by electrolysis. - For example, aluminium can be extracted from its ore, bauxite (aluminium oxide). - Bauxite is first purified and then dissolved in cryolite. This is to lower the melting point of aluminium oxide.

- When aluminium oxide melts, aluminium ions and oxide ions are free to move. - When electricity is passed through the electrolyte, the positively-charged aluminium ions are attracted to the cathode. They receive electrons and become aluminium atoms. - The molten aluminium formed is channelled into moulds. - At the anode, the oxide ions lose electrons to become oxygen atoms. The oxygen atoms combine together to form oxygen gas.

Aluminium is extracted from bauxite by electrolysis

(iii) Purification of metals - When electricity is passed through the electrolyte, the copper anode dissolves to form copper ions. At the same, the impurities settle to the bottom. - These positively-charged ions are attracted to the cathode. They receive electrons and form copper atoms and causes copper to be deposited on the pure copper. Purifying copper by electrolysis 5.6 THE PRODUCTION OF ELECTRICAL ENERGY FROM CHEMICAL REACTIONS A simple cell consists of two electrodes and an electrolyte. The more reactive metal is the negative terminal which releases electrons whereas the less reactive metal acts as a positive terminal which receives electrons.

A simple cell Refer to the diagram above: Magnesium and copper strips are used as electrodes with copper sulphate solution as the electrolyte. Magnesium (negative terminal) is more reactive than copper (positive terminal). Therefore, magnesium atoms release electrons to form magnesium ions which move into the electrolyte. As a result, the magnesium strip becomes thinner. The electrons flow to the copper strip through the wire and this produces electricity. Positively-charged copper ions from the electrolyte receive the electrons and form copper atoms. This causes copper to be deposited on the copper strip. Therefore, the copper strip becomes thicker. At the same time, the blue colour of the copper sulphate solution fades.

Type of cell

Uses

Advantages

Disadvantages

Dry cells

Used in portable devices such as radios, cassette players, cameras and toys

Light Small Portable Cheap

Non-rechargeable Not long-lasting Leakage may happen when the zinc case becomes thinner

Lead-acid accumulators

Used in vehicles

High voltage Rechargeable Long-lasting if well taken care of

Heavy Expensive Acid might spill Distilled water needs to be added

Alkaline batteries

Used in radios, torch lights and toys which need large electric current for long-lasting periods

Long-lasting Large current Constant voltage

Non-rechargeable

Mercury cell

Used in watches, calculators, hearing aides and measuring instruments

Small Portable Steady voltage Constant current Long-lasting

Expensive Non-rechargeable

Nickel-cadmium batteries

Used in electronic devices such as digital cameras

Rechargeable Long-lasting

Expensive

5.7 CHEMICAL REACTIONS THAT OCCUR IN THE PRESENCE OF LIGHT During photosynthesis, chlorophyll absorbs light energy to split water molecules into hydrogen and oxygen. The oxygen is released into the atmosphere whereas the hydrogen atoms react with carbon dioxide to form glucose. lightCarbon + Water Glucose + Oxygen dioxide chlorophyll

Photosynthesis requires light

5.8 INNOVATIVE EFFORTS IN THE DESIGN OF EQUIPMENT USING CHEMICAL REACTIONS AS SOURCES OF ENERGY We should use electric cells efficiently to prevent wastage. Turn off all the electrical devices when they are not in use. After using, we should dispose them wisely to reduce environmental pollution. Fuel cells are electric cells which does not run down or require recharging. A hydrogen fuel cell uses hydrogen as fuel to react with oxygen to produce energy.The product is water which will not pollute the environment.