admin

SPM Form 5 Chemistry Chapter 3 – Oxidation and Reduction

by

Revision Notes

  1. Redox Reaction
    1. Redox Reaction in Terms of Gain and Loss of Oxygen
    2. Redox Reaction in Terms of Gain and Loss of Hydrogen
    3. Oxidation Number
    4. Oxidation Number and IUPAC Nomenclature
    5. Redox Reaction in Terms of Change of Oxidation Number
    6. Redox Reaction in Terms of Transfer of Electron
  2. Types of Redox Reaction
    1. Oxidising and Reducing Agents
    2. Change of Iron(II) to Iron(III) and Vice Versa
    3. Displacement of Metals from Solution of Their Salts
    4. Displacement of Halogens
    5. Transfer of Electrons at A Distance
  3. Corrosion as a Redox Reaction
    1. Prevention of Rusting
  4. Series of Reactivities of Metals
    1. Position of Carbon in The Series of Reactivity of Metals
    2. Position of Hydrogen in The Series of Reactivity of Metals
    3. Extraction of Metals from Their Ores
  5. Electrolytic and Chemical Cells
    1. Redox Reaction in Chemical Cells
    2. Dry Cell
    3. Alkaline Cell
    4. Lead Acid Accumulator

Videos

  1. Oxidation and Reduction – Loss or Gain of Oxygen or Hydrogen | Loss or Gain Electron | Change of Oxidation States
  2. Oxidation and Reduction in Term of Changes of Oxidation State
  3. Oxidising Agent & the Half Equation
  4. Redox Reaction – Oxidation of Iron (II) to Iron (III)
  5. Redox Reaction – Reduction of Iron (III) to Iron (II)
  6. Reaction between Chlorine Water and Potassium Bromide Solution | Oxidation and Reduction
  7. The Oxidising Agent and Reducing Agent
  8. Reaction between Potassium Dichromate (VI) and Potassium Iodide  – Part 1 | Part 2 | Part 3

Graph of Product/Reactant Change Against Time

by

Graph of Product/Reactant Change Against Time

  1. In a chemical reaction,
    1. the reactants will decrease over time
    2. the product will increase over time.
  2. the rate of reaction will decrease over time owing to the decrease in concentration and total surface area of reactants.
  3. In a graph of quantity of product/reactant over time, the rate of reaction is equal to the gradient of the graph.

Example
The reaction between dilute hydrochloric acid and excess marble will produce calcium chloride and gas of carbon dioxide. Sketch the graph of

  1. the mass of the marble against time.
  2. the volume of carbon dioxide against time.
  3. the concentration of hydrochloric acid against time.
  4. the concentration of calcium chloride against time.

Answer:
a.

b.

c.

d.

Finding Average Rate of Reaction

by

Finding Average Rate of Reaction

Rate of reaction is a measure of how fast a reaction occur, or how much the reactant/product change in a period of time.
Rates of reaction = Quantity change of reactants/products Total time for the reaction

Example
In a chemical reaction, 2.5g of calcium carbonate react completely with excess hydrochloric acid to produce 600cm³ of carbon dioxide gas in 1.5 minutes. Find the rate of reaction in term of
a. decreasing mass of calcium carbonate
b. increasing volume of carbon dioxide gas produced

Answer:
a.
Change of the amount of reactant=-2.5g Tima taken for the change=1.5minute=90s Rate of Reaction = -2.5g / 90s = 0.027gs-1

b.
Change of the amount of product
=600cm3 Tima taken for the change
= 1.5minute = 90s

Rate of Reaction
= 600cm3/90s
=6.7cm3s-1

Understanding Rate of Reaction

by
  1. Rate of reaction is defined as the change in the amount of reactants or products per unit time.
  2. It is a measure of how fast a reaction occurs.
    1. Fast Reaction = Rate of reaction is high.
    2. Slow Reaction = Rate of reaction is low
  3. A fast reaction taken shorter time for the reaction to complete.
Example of fast reaction
Type of Reaction Example
Combustion Combustion of magnesium in oxygen \[2Mg + {O_2} \to 2MgO\] Combustion of ethane (C2H6) \[{C_2}{H_6} + \frac{7}{2}{O_2} \to 2C{O_2} + 3{H_2}O\]
Reaction between reactive metal and water Reaction between potassium and water \[2K + 2{H_2}O \to 2KOH + {H_2}\]
Reaction between metal carbonate and acid Reaction between limestone/ marble and sulphuric acid \[\begin{gathered} CaC{O_3} + {H_2}S{O_4} \hfill \\ \to CaS{O_4} + C{O_2} + {H_2}O \hfill \\ \end{gathered} \]
Ionic precipitation (Double decomposition) Precipitation of silver(I) chloride \[AgN{O_3} + HCl \to AgCl + HN{O_3}\]
Example of slow reaction
Type of Reaction Example
Photosynthesis \[6C{O_2} + 6{H_2}O \to {C_6}{H_{12}}{O_6} + 6{O_2}\]
Rusting \[4Fe + 3{O_2} + 2{H_2}O \to 2F{e_2}{O_3} \bullet 2{H_2}O\]
Fermentation \[{C_6}{H_{12}}{O_6} \to 2{C_2}{H_5}OH + 2C{O_2}\]

SPM Form 5 Chemistry Chapter 1 – Rate of Reaction

by

Revision Notes

  1. Rate of Reaction
    1. Average Rate of Reaction
    2. Average Rate of Reaction – Measurable Quantities
    3. Average Rate of Reaction – Immeasurable Quantities
  2. Analysing Rate of Reaction from Graph
    1. Average Rate of Reaction from a Graph
    2. Instantaneous Rate of Reaction from a Graph
  3. Factors Affecting the Rate of Reaction
    1. Total Surface Area
    2. Concentration of Reactants
    3. Temperature of the Reactant
    4. Presence of Catalyst
    5. Pressure of Gas
  4. Collision Theory
  5. Application of Rate of Reaction

Videos

  1. Understanding Rate of Reaction
  2. Finding Rate of Reaction (Example)
  3. Finding Average Rate of Reaction from Graph
  4. Finding Instantaneous Rate
  5. Finding Instantaneous Rate (Example)
  6. Factors Affecting the Rate of Reaction
  7. Factors Affecting the Rate of Reaction – Size of Particle
  8. Factors Affecting the Rate of Reaction – Concentration
  9. Factors Affecting the Rate of Reaction – Temperature
  10. Factors Affecting the Rate of Reaction – Catalyst

Ostwald Process

by

Ostwald Process

Introduction

  1. Industrially, nitric acid is made by the catalytic oxidation of ammonia over heated platinum. 
  2. Oxidising ammonia produces oxides of nitrogen which can then be dissolved in water to produce nitric acid.

Reaction

  1. Initially, nitrogen(II) oxide will be formed from the catalytic oxidation of ammonia using the transition metal platinum.
    Ammonia + Oxygen → Nitrogen(II) Oxide + Steam
    4NH3 (g) + 5O2 (g) → 4NO (g) + 6H2O (g)
  2. The nitrogen(II) oxide is rapidly cooled before combining with oxygen (from the excess air) to form nitrogen(IV) oxide.
    2NO (g) + O2 (g) → 2NO2 (g)
  3. The nitrogen(IV) oxide, mixed with excess air, is then allowed to react with water to form nitric acid.
    Nitrogen(IV) Oxide + Oxygen (air) + Water → Nitric acid
    4NO2 (g) + O2 (g) + 2H2O (1) → HNO3 (aq)

Uses of Nitric Acid

  1. Most of the nitric acid made is used to make the all-important fertilisers, such as ammonium nitrate.
  2. Other uses of the nitric acid include making explosive, like nitroglycerine, or TNT (trinitrotoluene), and making dyes. Modern dyes are azo dyes, which can be formed by the reduction of various nitro-compounds.

Characteristics of Ammonia

by

Characteristics of Ammonia

  1. Ammonia gas can turn a moist red litmus paper to blue.
  2. As an alkali, ammonia can react with acid to form salt and water.
    Example
    H2SO4(aq) + 2NH3(aq) → (NH4)2SO4(aq)
    HNO3(aq) + NH3(aq) → NH4NO3(aq)
    H3PO4(aq) + 3NH3(aq) → (NH4)3PO4(aq)
  3. Ammonia dissolve into water to form ammonium and hydroxide ion.
    NH3 + H2O→ NH4+ + OH–
  4. The hydroxide ion can react with many kinds of positive ion to form precipitate.
    Example
    Mg2+ + 2OH– → Mg(OH)2
    Fe2+ + 2OH– → Fe(OH)2
    Al3+ + 3OH– → Al(OH)2

Testing for Ammonia

  1. Ammonia is the only common alkaline gas so it can be identified with moist red litmus paper turning blue.
  2. Concentrated ammonia when reacts with concentrated hydrochloric acid produces white fume.
    Ammonia gas + Hydrogen chloride gas → ammonium chloride
    NH3 (g) + HC1 (g) → NH4C1

Haber Process

by

Haber Process

  1. Ammonia is manufactured in industries through Haber Process.
  2. In the Haber process, nitrogen gas, N2 from the air is mixed with hydrogen gas, H2 derived mainly from natural gas.
  3. The mixture is compressed to a high pressure of 200 atmospheres at a temperature of about 450°C.
  4. Iron is used as a catalyst to speed up the rate of reaction.
  5. Chemical equation below shows the reaction.
    N2 (g) + 3H2 (g) → 2NH3 (g)
  6. About 98% of mixture are converted into ammonia, NH3.
  7. The unreacted nitrogen gas, N2 and hydrogen gas, H2 is recycled and passed back into the reactor together with the new source of nitrogen gas, N, and hydrogen gas, H2.

Summary

Ammonia

by
  1. Ammonia is a compound of nitrogen and hydrogen with the formula NH3.
  2. It is a colourless gas with a characteristic pungent smell.
  3. Ammonia is an essential compound in industry.
  4. Although in extensive use, ammonia is both corrosive and hazardous.

Uses of Ammonia

  1. The uses of ammonia include
    1. manufacturing nitrogenous fertilisers
    2. as a cooling agent in refrigerator
    3. to prevent coagulation of latex
    4. as raw material to manufacture nitric acid (Ostwald process)
    5. to make explosive
    6. as cleaning agent to remove grease

(Click on the image to enlarge)

Environmental and Health Issues of Sulphur Dioxide

by

Environmental and Health Issues of Sulphur Dioxide

Sulphur dioxide, SO2 is one of the by-products of the Contact Process. It is one of the sources of environmental pollution.

Acid Rain

  1. Sulphur dioxide (SO2) is the pollutant primarily associated with acid rain.
  2. Acid rain occurs when pH of the rain is between 2.4 and 5.0. This is due to the reaction of sulphur dioxide, SO2 with rainwater.
    SO2 + H2O→ H2SO3
  3. The negative effect of acid rains includes
    1. corrosion of concrete building and metal structure.
    2. corrosion of monuments and statues made from marble
    3. causes erosion of top soil.
    4. killing aquatic life.

Health Effects

  1. SO2 is an irritant when it is inhaled and at high concentrations may cause severe problems in asthmatics such as narrowing of the airways, known as bronchoconstriction.
  2. Asthmatics are considerably more sensitive to the effects of SO2 than other individuals.

Sources of SO2

  1. The principal source of SO2 is from the combustion of fossil fuels in domestic premises and, more importantly, non-nuclear power stations.
  2. Other industrial processes such as manufacturing of sulphuric acid also contribute to the presence of SO2 in the air.