Chemical reactions
Chemical reactions
Assumed perquisite knowledge:
Students understand the difference between a physical change and a chemical change
Introduction
1.A chemical reaction occurs when substances combine or breakdown to produce something NEW.
2.A chemical reaction is accompanied by a chemical change. eg. permanment change in color or the
production of a gas.
Chemical changes indicate that something new has been produced. eg. the permanent color change
produced as bread is toasted in a toaster or the smoke from a burning match.
3.Chemical reactions are generally difficult to reverse. eg. turning a slice of toast into a slice of fresh
bread.
The chemicals that react are called the reactants. The new materials that are produced are called the
products.
A chemical reaction can be shown by the use of a word equation. The ==> means 'give or produce'
Reactants
==>
Products
chemicals that react
==>
chemicals that are produced
hydrogen + oxygen
==>
water
hydrogen peroxide
==>
water + oxygen
magnesium + oxygen
==>
magnesium oxide
sodium chloride
==>
sodium + chlorine
Types of chemical reactions.
•There are many different types of reactions which include
•Corrosion of metals. Eg. The rusting of iron
•Combustion: Eg. The burning of fossil fuels
•Acid and base : Eg. Hydrochloric acid and sodium hydroxide.
•Acid and carbonate: Eg. Vinegar and bicarb of soda
•Acid and active metal: This type of reaction is also called a single displacement reaction. Eg. The
dissolving of magnesium ribbon in hydrochloric acid.
•Precipitation: This type of reaction is also called a double displacement reaction. Eg. Mixing
colorless lead nitrate and potassium iodide to yellow produce lead iodide.
•Decomposition: Eg. Hydrogen peroxide being broken down into water and oxygen
•Synthesis: Eg. The burning of magnesium in air to produce magnesium oxide.
•Displacement:. Eg. The production of copper wool by from steel wool being placed into a solution
of copper sulphate.
•Redox (Reduction/ Oxidation): Eg. Making a battery.
•Condensation: Eg. Glucose and Fructose producing sucrose
•Dehydration: Eg. Adding concentrated sulphuric acid to sugar
•Substituion: Eg. Bromine reacting with octane in the presence of UV light.
•Addition: Eg. Adding bromine to cyclohexene
•Polymerisation: Eg. making Nylon, polyethylene
•Fermentation: Eg. The making of wine
•Exothermic: Eg. The thermite reaction is used to weld rail tracks together. The reaction is so hot
molten iron is produced!
•Endothermic: Eg. An ice pack. Mix the contents inside the packet and the mixture becomes cold
Acid and bases
Acids and Base reactions
Acids are substances that taste sour. eg. citric acid. Bases are substances that taste bitter. eg. Bi-carb
of soda. There are many different types of acids and bases and most are too corrosive to taste. The
word corrosive means it will eat away. There must be another way to determine whether a chemical
substance is acidic or basic.
Indicators are chemicals that change color in the presence of an acid or a base. Pink litmus paper will
change to blue in the presence of a base. Blue litmus paper will change to pink in the presence of an
acid.
Acids and bases react with each other to produce water and a salt. This type of reaction is called a
neutralization reaction. Acid + Base => Salt + Water
An alkali is a strong soluble base. All of the Group I metals or alkali metals produce alkali solutions on
reaction with water. eg. lithium hydroixde, sodium hydroxide and potassium hydroxide.
A common neutralization reaction carried out in high school is the neutralization of hydrochloric acid
with sodium hydroxide.
Hydrochloric acid + sodium hydroxide => Sodium chloride + Water
2HCl + NaOH => NaCl + H2O
Bromothymol blue is an excellent indicator to add as it is yellow in acid, blue in base and green when
neutral. The solution can be evaporated in a watch glass which is placed on top of a water bath. Salt
crystals appear as the water evaporates.
All neutralization reactions are also exothermic meaning heat is given off. Hydrogen ions produced
from the acid react with hydroxide ions produced from the base. Energy is released as new chemical
bonds are formed as water is produced.
Hydrogen ion + hydroxide ion => water
H+(aq)+ OH-(aq) => H2O(l)
Acids and carbonates
What is a carbonate?
Carbonates are compounds that contain the CO3 group.
Sodium hydrogen carbonate also called bi-carb of soda or baking soda is a common household
carbonate.
It's chemical formula is NaHCO3
Identify which chemical compounds contain the carbonate group.
Click on the links for the answers.
No.
Chemical
Carbonate? Yes/No
1.
Na2SO4
Is this a carbonate?
2.
Na2CO3
Is this a carbonate?
3.
NaNO3
Is this a carbonate?
4.
CuSO4
Is this a carbonate?
5.
CuCO3
Is this a carbonate?
6.
CaCO3
Is this a carbonate?
7.
Ca(NO3)2
Is this a carbonate?
8.
PbSO4
Is this a carbonate?
9.
MgCO3
Is this a carbonate?
10.
SiC
Is this a carbonate?
Show all | Hide all
Acids and Carbonate Reactions
Key point:
Compounds with a carbonate or CO3 group react with acids to produce carbon dioxide, CO2 gas.
General word equation
Acid + Metal carbonate Salt + Water + Carbon dioxide
All carbonates react with acids to produce salt, water and carbon dioxide gas.
The fizz produced in sherbet is a reaction between a food acid and a carbonate.
The ingredients dissolve in your mouth and react with each other producing the carbon dioxide fizz.
Examples of acid and carbonate reactions
1. Hydrochloric acid + sodium carbonate sodium chloride + water + carbon dioxide
2HCl + Na2CO3 2NaCl + H2O + CO2
2. Sulfuric acid + copper carbonate copper sulfate + water + carbon dioxide
H2SO4 + CuCO3 CuSO4 + H2O + CO2
3. Nitric acid + magnesium carbonate magnesium nitrate + water + carbon dioxide
2HNO3 + MgCO3 Mg(NO3)2 + H2O + CO2
4. Acetic acid + sodium hydrogen carbonate sodium acetate + water + carbon dioxide
CH3COOH + NaHCO3 CH3COONa + H2O + CO2
Acids and metals
Acid on metal
Acids react with active metals like zinc, magnesium and iron to produce hydrogen gas and salt.
Chemical change
Chemical change
A chemical change is a sign that indicates that a chemical reaction has occurred.
Chemical changes include
•The production of a gas. Eg. Fizz from sherbet
•A permanent color change. Eg. A red solution changing to a colorless solution
•The production of light. Eg. Glow in the dark light sticks which is an example of chemiluminesce
•The production of a flame. Eg. A Bunsen burner
•The dissolving of an insoluble solid. Eg. Magnesium dissolving in acid.
•The production of an insoluble solid. Eg. The production of a precipitate
These changes indicate that something new has been produced.
Physical changes on the other hand are easy to reverse and don't produce anything new. They only
involve a change in the physical states (solid. liquid and gas). eg. the melting of an ice cube and
steam being produced from a kettle.
Combustion
Combustion is the burning of a fuel with the production of energy or heat
Examples of combustion reactions
1. Cellular respiration
Cellular respiration is the process whereby cells in our bodies obtain energy. Glucose reacts with
oxygen producing carbon dioxide and water with the release of energy.
Glucose + oxygen => carbon dioxide + water + energy
C6H12O6 + 6O2 => 6CO2 + 6H2O + energy
2. The Combustion of fossil fuels
Fossil fuels are made up of hydrocarbons. Hydrocarbons are compounds that contain the elements
hydrogen and carbon. Examples are methane, CH4, propane, C3H8, butane, C4H10 and octane,
C8H18
Hydrocarbons readily burn or undergo combustion reactions.Combustion may be complete or
incomplete.
Complete combustion of fossil fuels results the production carbon dioxide and water. The
incomplete combustion of fossil fuels results in the formation of carbon monoxide and soot
(carbon).
1. Complete combustion of methane: Burns with a blue flame
Methane + oxygen => carbon dioxide and water
CH4 + 2O2 => CO2 + 2H2O
2. Incomplete combustion of methane producing carbon monoxide
Methane + oxygen => carbon monoxide and water
CH4 + 1.5O2 => CO + 2H2O
3. Incomplete combustion of methane producing soot: Burns with a yellow flame
Methane + oxygen => carbon and water
CH4 + O2 => C + 2H2O
As the amount of oxygen is decreased from 2 to 1.5 to 1 in the above three equations the
combustion products respectively move from carbon dioxide to carbon monoxide to carbon or soot.
As the soot is heated in the flame it produces a yellow flame. In real life there is a mixture of carbon
monoxide and soot from the incomplete combustion of hydrocarbons as well as unburnt
hydrocarbons.
Corrosion
Corrosion reactions
Corrosion is the loss of metallic properties of a metal as the metal reacts with the atmosphere or
water. e.g. strength, lustre or shine and electrical conductivity.
Rust is brownish red in color and is formed from the corrosion of iron. Other metals like copper and
aluminium also corrode or weaken.
For corrosion to occur both water and oxygen need to be present.
Examining the corrosion of iron using nails:
Oxygenated water (shaken)
Humid air
Dry air (anhydrous CaCl2)
Deoxygenated water
(Boiled water/ oil layer)
Salt water
Rust
Rust
No rust
No rust
Extensive rust
If a nail is placed in deoxygenated water it will not rust. If a nail is placed in dry air it will not rust. For
rust to occur both oxygen and water need to be present. The presence of salt accelerates or speeds
up the rate of corrosion.
The chemical formula of rust is Fe2O3.nH2O
The process of rust formation is due to a series of chemical reactions. Points of stress acts in the nail
allow the iron to be easily oxidised. Carbon impurities in the nail allow water to be reduced. An
internal battery or galvanic cell is set up.
Oxidation of iron at points of stress in the crystal lattice: 2Fe(s) ==> 2Fe2+(aq) + 4eReduction of water at the site of carbon impurities: O2(g) + H2O(l) + 4e- ==>4OH-(aq)
Overall equation: 2Fe(s) + O2(g) + H2O(l) ==>Fe(OH)2
The iron(II) hydroxide is converted to rust through a serious of reactions.
The ion(II) hydroixde firstly oxides to iron(III) oxide.
1. Fe(OH)2(s) =oxidation-=> Fe(OH)3
The iron(III) oxide then changes to rust through a dehydration reaction.
2. Fe(OH)3(s) =dehydration=> Fe2O3.nH2O(s) or rust
Rust does adheres loosely to the surface of the metal. This exposes the metal to more and more
water and oxygen allowing rust to continue to proceed.
Decomposition
Decomposition reactions
In a decomposition reaction a chemical substance is broken down into simpler substances.
The general equation for the reaction is A => B + C
Substance A decomposes or breaks apart producing substances B and C.
Decomposition can be achieved by 1. heat, 2. electricity and 3. a catalyst.
Note: Some extremely unstable chemicals can be set of by sound waves!
Examples of decomposition reactions
1. Thermal decomposition of copper carbonate.
If a test tube containing green copper carbonate is heated in a Bunsen burner the powder turns
black and occasionally ‘jumps’ as carbon dioxide gas is given off.
Copper carbonate =heat=> Copper oxide + carbon dioxide
CuCO3 =heat=> CuO + CO2
2. The electrolytic decomposition of water
Electrolysis is the process using electrical energy to speed up or cause a chemical reaction.
Pure water is a poor electrical conductor. Sodium sulfate salt or sulfuric acid needs to be added to
the water first. Note: DO NOT use sodium chloride salt or hydrochloric acid as toxic chlorine gas is
produced!
Two inert electrodes (carbon or platinum) are placed into the water and are connected to an
external D.C power pack or series of batteries. Hydrogen gas is produced at the negative electrode
and oxygen gas at the positive electrode in a ratio of 2:1.
Water =electrolysis=> Hydrogen and Oxygen
2H2O =electrolysis=> 2H2 + O2
3. The catalytic decomposition of hydrogen peroxide
A catalyst is a substance that speeds up a chemical reaction without being used up in the reaction.
Hydrogen peroxide is broken done quite rapidly into oxygen gas and water when a catalyst like
potassium iodide or manganese dioxide is added. Potassium iodide salt is preferred if available as it
makes less mess. The black manganese dioxide powder is difficult to clean up.
One to two spatulas of potassium iodide salt is added to a conical flask containing 50mL of 5%
hydrogen peroxide solution. The solution is gently mixed and bubbles of gas are observed to form. If
a glowing splint is placed into the conical flask it will re-ignite due to the production of oxygen gas.
Hydrogen peroxide =catalyst=> Oxygen and water
2H2O2 =catalyst=> O2 + 2H2O
Displacement
Displacement reactions
Metals differ in their reactivities. Magnesium is more reactive than zinc and zinc is more reactive
than iron.
Metals can be placed in order of decreasing chemical activity. This is called the activity series of
metals.
Activity series of metals
Metal ion
Metal
Symbol
Mg2+
Magnesium
Mg
Most active metal
Zn2+
Zinc
Na
Fe2+
Iron
Fe
Pb2+
Lead
Pb
Cu2+
Copper
Cu
Ag+
Silver
Ag
Least active metal
An active metal will displace the ion of a less active metal from solution.
A metal higher on the table will react with a metal ion solution lower on the table.
An example of a displacement reaction is placing a piece of steel wool into a solution of copper
sulfate.
Note: Don't use the commercial brands of steel wool that are protected from rusting. They will not
react.
The steel wool is turned into pink copper wool. Over time the solution also looses its blue color and
turns olive green.
Iron (steel wool) + Copper sulfate => Iron sulfate + Copper
Fe +
CuSO4
=>
FeSO4 +
Cu
Blue
Green
Pink
The reaction is also exothermic producing heat.
Precipitation
Precipitation reactions
In a precipitation reaction two solutions are mixed together resulting in the formation of an
insoluble solid which is called the precipitate.
Redox
Reduction oxidation reactions
Reduction oxidation reactions are called redox reactions. They involve the transfer of electrons from
one species to another and are represented by two half equations, the oxidation half equation and
the reduction half equation.
Oxidation half equation: Eg. Mg ==> Mg2+ + 2e- OIL: Oxidation Is Loss of electrons by a chemical
species. The Mg species looses electrons and is oxidised.
Reduction half equation: Eg. Cu2+ + 2e- ==> Cu RIG: Reduction Is Gain of electrons by a species.The
Cu2+ ion gains electrons and is reduced.
Notice the electrons, symbol e-, are on different sides of the two equations in equal numbers.
Oxidation reduction reactions always occur together. One chemical species looses electrons and the
other species gains electrons.
The resulting net reaction is called a redox reaction. Eg. Mg + Cu2+ ==> Mg2+ + Cu
A galvanic cell is a device which separates these two half reactions allowing the electrons to flow
through an external wire. We call this device a battery.
Synthesis
Synthesis reaction
The word synthesis means to put together. In a synthesis reaction two elements combine to produce
a compound. A + B => C
The reaction of a metal with a non-metal to produce a compound is an example of a synthesis
reaction.
Examples of synthesis reactions
1. The burning of magnesium metal
If a piece of magnesium ribbon is ignited a very bright light is given off and a white powder or ash is
produced/ Note: the burning of magnesium produces UV light which may damage the eyes if looked
upon directly. Wear safety goggles and do not directly look at the light.
Magnesium + oxygen => magnesium oxide
2Mg + O2 => 2MgO
2. The reaction of sodium with chlorine
If a piece of sodium metal is melted in a deflagrating spoon and placed into a gas jar containing
chlorine gas a violent reaction occurs. If the jar is observed crystals of sodium chloride salt are seen
which have formed on the sides of the glass.
Sodium + chlorine => Sodium chloride
2Na + Cl2 => 2NaCl
3. The reaction of alumnium with bromine
Aluminum is resistant towards reaction because of a tough impermeable oxide layer which strongly
adheres to the metal. If however, a sheet of aluminum foil is scratched with steel wool the
aluminum is exposed it will react violently with pure liquid bromine
Aluminum + bromine => aluminum bromide
2Al + 3Br2 => 2AlBr3
4. The burning of iron
A nail will not react when heated over a Bunsen burner. However, if the surface area of the iron is
increased by using steel wool a reaction is seen to occur. Increasing the surface area increases the
rate of reaction.
Iron + oxygen => Iron(II) oxide
2Fe + O2 => 2FeO
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