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Chemistry Form Five: Chapter 1 - Rate of Reaction
Rate of reaction = change of quantity in reactant or product per unit time.
We usually use water displacement method to collect gas in school laboratory as shown below:
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The reaction is fastest at the start when the reactants are at a maximum
(steepest gradient) The gradient becomes progressively less as reactants are used up and the
reaction slows down.
Finally the graph levels out when one of the reactants is used up and thereaction stops.
The amount of product depends on the amount of reactants used.
The initial rate of reaction is obtained by measuring the gradient at the start
of the reaction. A tangent line is drawn to measure rate of reaction atinstataneous time
hemistry Form 5: Chapter 1 - Collision Theory
According to the collision theory, particles of reactant that achieve activation energy and collide with
correct orientation will result in reaction.
1. Correct Orientation
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Click on the diagram below to play!
2. Activation Energy
Activation energy is the minimum amount of energy that must be overcome by the colliding
particles so that the reaction can occur
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Chemistry Form 5: Chapter 1 - Effect of Concentration on Rate of
Reaction
Experiment to show the effect of concentration on reaction rate
Sodium thiosulphate solution react with dilute sulphuric acid to form a yellow precipitate
of sulphur. In this experiment, the time taken for the formation of sulphur to cover themark 'X' until it disappears from sight can be used to measure rate of reaction.
As the concentration of sulphuric acid is increased, the rate of reaction between sulphuric
acid and sodium thiosulphate increases.
Explanation using collision theory
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When the concentration of the solution of a reactant increases, the number ofparticles per unit volume of the solution also increases.
With more particles per unit volume of the solution of the reactant, the
frequency of collision increases. This causes the frequency of effective collision to increase. Hence, the rate of
reaction increases.
Chemistry Form 5: Chapter 1 - Catalyst Affects the Rate of Reaction
Catalyst is a chemical substance that change the rate of chemical reaction.
Characteristics of catalyst:
Catalyst remains chemically unchanged during reaction. Its chemical composition still
the same before and after reaction.
Catalyst only change the rate of reaction.
Catalyst does not change the quantity of the product formed.
Catalyst is specific in its action.
Only a small amount of catalyst is needed to achieve a big increase in rate of reaction.
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How catalyst increase the rate of reaction:
When a positive catalyst is used in a chemical reaction, it enables the reaction to occur
through an alternative path which requires lower activation energy.
As a result, more colliding particles are able to overcome the lower activation energy.
This causes the frequency of effective collision to increase.
Hence, the rate of reaction increases.
Decomposition of hydrogen peroxide by catalyst of manganese (IV) oxide
hemistry Form 5: Chapter 2 - Hydrogenation
Hydrogenation process is addition reaction to convert alkene becomes alkane. It converts
unsaturated compound to saturated compound.
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Ethene reacts with H2 at 180 C in the presence ofnickel or platinum catalyst to produce ethane
Click on the diagram below to play!
Application of hydrogenation: Making Margarine
Vegetable oils often contain high proportions of polyunsaturated and mono-unsaturatedfats (oils), and as a result are liquids at room temperature. That makes them messy tospread on your bread or toast, and inconvenient for some baking purposes.
You can "harden" (raise the melting point of) the oil by hydrogenating it in the presenceof a nickel catalyst. Conditions (like the precise temperature, or the length of time thehydrogen is passed through the oil) are carefully controlled so that some, but notnecessarily all, of the carbon-carbon double bonds are hydrogenated. This produces a"partially hydrogenated oil" or "partially hydrogenated fat".
hemistry Form 5: Chapter 2 - Manufacture of Ethanol (Hydration)
Ethanol is manufactured by reacting ethene with steam. The reaction is reversible, and
the formation of the ethanol is exothermic.
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Chemistry Form 5: Chapter 2 - Dehydration of Alcohol
ALCOHOL -------> ALKENE
In the dehydration of alcohols, a molecule of water is eliminated from each alcohol
molecule to produce alkene.
There are two methods of dehydration:
a) Ethanol vapour is passed over a heated unglazed porcelain chips, porous pot, pumice stone or
alumina (aluminium oxide).
b) Ethanol is heated under reflux at 170 C with excess concentrated sulphuric acid.
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School Laboratory Experiment
Alkene can be tested by decolourising brown bromine water or decolourising purple
acidified potassium manganate (VII) solution.
hemistry Form 5: Chapter 2 - Carboxylic Acid
Carboxylic acids are organic compounds which form an homologous series with
the general formula of CnH2n+1COOH. Carboxylic acids are compounds which contain a -COOH functional group.
Carboxylic acids are weak acid which ionize partially in water to produce lower
concentration of hydrogen ions compare to strong acid.
Formula Common Name Source IUPAC Name
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HCO2H formic acid ants (L. formica) methanoic acid
CH3CO2H acetic acid vinegar (L. acetum) ethanoic acid
CH3CH2CO2H propionic acid milk (Gk. protus prion) propanoic acid
CH3(CH2)2CO2H butyric acid butter (L. butyrum) butanoic acid
CH3(CH2)3CO2H valeric acid valerian root pentanoic acid
CH3(CH2)4CO2H caproic acid goats (L. caper) hexanoic acid
CH3(CH2)5CO2H enanthic acid vines (Gk. oenanthe) heptanoic acid
CH3(CH2)6CO2H caprylic acid goats (L. caper) octanoic acid
CH3(CH2)7CO2H pelargonic acid pelargonium (an herb) nonanoic acid
CH3(CH2)8CO2H capric acid goats (L. caper) decanoic acid
hemistry Form 5: Chapter 2 - Esterification
Esterification is a chemical reaction between carboxylic acid combines with an
alcohol in the presence of a catalyst (commonly concentrated sulphuric acid) to
form an ester.
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Examples of Esters methyl butanoate (apple) :
methan
ol
+
butanoic
acid
methylbutanoate
(ester)
+
wate
r
CH3OH+
C3H7COOH C3H7COOCH3 (ester)
+
H2O
ethyl methanoate (rum essence) :
ethan
ol
+
methanoic
acid
ethylmethanoate
(ester)
+
wate
r
C2H5O
H
+
HCOOH HCOOC2H5 (ester)
+
H2O
ethyl butanoate (pineapple) :
ethan
ol
+
butanoic
acid
ethylbutanoate
(ester)
+
wate
r
C2H5O
H
+
C3H7COOH
C3H7COOC2H5
(ester)
+
H2O
pentyl ethanoate (banana) :
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pentan
ol
+
ethanoic
acid
pentylethanoate
(ester)
+
wate
r
C5H11O
H
+
CH3COOH
CH3COOC5H11
(ester)
+
H2O
pentyl butanoate (apricot) :
pentan
ol
+
butanoic
acid
pentylbutanoate
(ester)
+
wate
r
C5H11O
H
+
C3H7COOH
C3H7COOC5H11
(ester)
+
H2O
octyl butanoate (orange) :
octano
l
+
butanoic
acid
octylbutanoate
(ester)
+
wate
r
C8H17O
H
+
C3H7COOH
C3H7COOC8H17
(ester)
+
H2O
methyl ethanoate (solvent) :
methan
ol
+
ethanoic
acid
methylethanoate
(ester)
+
wate
r
CH3OH+
CH3COOH CH3COOCH3 (ester)
+
H2O
ethyl ethanoate (solvent) :
ethan
ol
+
ethanoic
acid
ethylethanoate
(ester)
+
wate
r
C2H5O
H
+
CH3COOH
CH3COOC2H5
(ester)
+
H2O
Esterification by refluxing
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Chemistry Form 5: Chapter 3 - Redox Reaction
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Redox reactions are chemical reactions involving oxidation and reduction
occurring simultaneously. Oxidising agent is the substance that causes oxidation.
Reducing agent is the substance that causes reduction.
Oxidation involves loss of electrons and increase in oxidation number.Reduction involves gain of electrons and decrease in oxidation number.
Example:
The magnesium's oxidation state has increased from 0 to +2 , it has been oxidised.
Magnesium acts as reducing agent. The hydrogen's oxidation state has decreased from+1 to 0 , it has been reduced. The hydrogen ion acts as a oxidising agent.
Click on the diagram below to play!
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There is no change in oxidation number. Therefore, this is not a redox reaction.
Chemistry Form 5: Chapter 3 - Rules of Oxidation Number
There are several rules for assigning the oxidation number to an element. Learning these rules
will simplify the task of determining the oxidation state of an element, and thus, whether it has
undergone oxidation or reduction.
1. The oxidation number of an atom in the elemental state is zero.Example: Cl2 and Al both are 0
2. The oxidation number of a monatomic ion is equal to its charge.
Example: In the compound NaCl, the sodium has an oxidation number of 1+ and
the chlorine is 1-.
3. The algebraic sum of the oxidation numbers in the formula of a compound is zero.
Example: the oxidation numbers in the NaCl above add up to 0
4. The oxidation number of hydrogen in a compound is 1+, except when hydrogen forms
compounds called hydrides with active metals, and then it is 1-.
Examples: H is 1+ in H2O, but 1- in NaH (sodium hydride).
5. The oxidation number of oxygen in a compound is 2-, except in peroxides when it is 1-,
and when combined with fluorine. Then it is 2+.
Example: In H2O the oxygen is 2-, in H2O2 it is 1-.
6. The algebraic sum of the oxidation numbers in the formula for a polyatomic ion is equal
to the charge on that ion.
Example: in the sulfate ion, SO42-, the oxidation numbers of the sulfur and the
oxygens add up to 2-. The oxygens are 2- each, and the sulfur is 6+.
Chemistry Form 5: Chapter 3 - Conversion of Iron (II) to Iron (III) and Iron
(III) to Iron (II)
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Conversion ofFe2+ to Fe3+
Oxidising agent: bromine water
Reducing agent: Fe2+ ions
Oxidation half equation:
Fe2+ ions lose electrons and are oxidized to Fe3+.
The presence of Fe3+
ions is confirmed by the formation ofbrown precipitate with excess of NaOH solution.
Fe2+ --------> Fe3+ + e
Reduction half equation:
Bromine molecules which give bromine water its brown
colour gain electrons and are reduced to colourless bromide
ions.
Br2 + 2 e -------> 2 Br-
Overall ionic equation:
2 Fe2+ + Br2 2 Fe3+ + 2 Br-
Observation:
Brown bromine water decolourises. The solution changes
colour from pale green (Fe2+) to yellow (Fe3+).
Conversion ofFe3+ to Fe2+
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Oxidising agent: Fe2+ ions
Reducing agent: zinc
Oxidation half equation:
Zinc atoms lose their electrons and are oxidized to zinc ions,
Zn2+.
Zinc powder dissolves in iron (III) chloride.
Zn -------> Zn2+ + 2 e
Reduction half equation:
Fe3+ ions accept electrons and are reduced to Fe2+.
The presence of Fe2+ ions is confirmed by the formation of
green precipitate with excess of NaOH solution.
Fe3+ + e Fe2+
Overall ionic equation:
2 Fe3+ + Zn ---------> 2 Fe2+ + Zn2+
Observation:
Zinc powder dissolves into solution. The solution changes
colour from brown (Fe3+) to pale green (Fe2+).
Chemistry Form 5: Chapter 3 - Redox Reaction in Displacement of Metal
Reactivity Series of metals
Displacement of metals from solution is a redoxreaction whereby a less reactive metal ion is
displaced from its salt solution by a more
reactive metal. As a result, the less reactivemetal ion is deposited as a solid metal while the
more reactive metal dissolves in the solution.
The general formula for a displacement reaction
is:
M (s) + Xn+ (aq) -----> Mn+ (aq) + X (s)
where metal M is the more reactive than metal
X.
Most reactive
K
Na
Ca
Mg
Al
Zn
Fe
Pb
[H]Cu
Ag
Least reactive
Example of displacement reaction:
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Fe (s) + CuSO4 (aq) -----> FeSO4 (aq) + Cu (s)
Iron displaces copper from the solution because it is more reactive than copper metal.Iron, being more reactive, loses its electrons readily. The electrons are transferred fromthe iron atoms to the copper(II) ions in the solution. Copper(II) ions are reduced tocopper metal and iron atoms become oxidised to iron (II) ions. Iron acts as reducingagent whereas copper (II) ions act as a oxidising agent.
Oxidation reaction: Fe (s) -----> Fe2+ (aq) + 2e-
Reduction reaction: Cu2+ (aq) + 2e- -----> Cu (s)Overall Redox reaction: Fe (s) + Cu2+ (aq) -----> Fe2+ (aq) + Cu (s)
hemistry Form 5: Chapter 3 - Rusting of Iron
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Rusting is a corrosion of iron.
For iron to rust, oxygen and water must be present.
In the presence of acids and salts, rusting occurs faster because thesesubstances increase the electrical conductivity of water, making water a betterelectrolyte.
Oxygen acts as the oxidizing agent and iron acts as the reducing agent .
The surface of iron at the middle of the water droplet serves as the anode at which oxidation occurs. The
iron atoms lose electrons to form iron (II) ions.
The electrons flow to the edge of the water droplet where there is plenty of dissolved oxygen. The iron
surface there serves as cathode at which reduction occurs. Oxygen gains the electrons and is reduced to
hydroxide ions.
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The iron (II) ions produced combine with the hydroxide ions to iron (II) hydroxide. The Fe(OH)2 is then
further oxidized by oxygen to form iron (III) oxide, Fe2O3 known as rust
Chemistry Form 5: Chapter 3 - Redox Reaction in Electrochemistry Reaction
More electropositive metal undergoes oxidation reaction by releasing electronsand act as a reducing agent.
Less electropositive metal undergoes reduction reaction by gaining electrons and
act as a oxidising agent.
Electrons flow from more electropositive metal to less electropositive metal.
Chemistry Form 5: Chapter 4 - Exothermic and Endothermic Reaction
Chemical energy is needed to transform a chemical substance into a new product through
chemical reaction. Therefore, breaking or formation of chemical bond involves energy, whichmay be either absorbed or released from a chemical reaction.
To break the chemical bond, energy from surrounding is absorbed resulting decrease of
temperature of surrounding.
To form the chemical bond, energy from reaction is released to the surrounding resulting
increase of temperature of surrounding.
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Soaps are sodium or potassium salts of fatty acids. Soaps are prepared by hydrolyzingfats or oils under alkaline condition. This reaction is called saponification.
The fats or oils are hydrolysed first to form glycerol and fatty acids. The acids then react with an
alkali to form the corresponding sodium or potassium salts.
The soap formed can be precipitated by adding sodium chloride. This is because sodium
chloride lowers the solubility of soap in water.
The glycerol and excess sodium hydroxide solution are removed by rinsing the soap formed
with water.
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