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chemistry-0620-core | cambridge-igcse
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1. Acids, Bases, and Salts
1.1.1 Making soluble salts by reaction with excess metal
1.1.2 Making soluble salts by reaction with excess base
1.1.3 Making soluble salts by reaction with excess carbonate
1.1.4 Making soluble salts by titration
1.2.1 Solubility of sodium, potassium, ammonium salts
1.2.2 Solubility of nitrates, chlorides, sulfates, carbonates, and hydroxides
1.3.1 Definition of hydrated and anhydrous substances
1.4.1 Making insoluble salts by precipitation
1.5.1 Definition and examples (CuSO₄·5H₂O, CoCl₂·6H₂O)
1.6.1 Reaction of acids with metals, bases, and carbonates
1.6.2 Effect of acids on litmus, thymolphthalein, and methyl orange
1.7.1 Definition of bases and alkalis
1.7.2 Reactions of bases with acids and ammonium salts
1.7.3 Effect of alkalis on litmus, thymolphthalein, and methyl orange
1.8.1 H⁺ ions in acids, OH⁻ ions in alkalis
1.8.2 Compare acidity and alkalinity using pH and indicators
1.9.1 Reaction of acids with alkalis to form water
1.10.1 Definition of strong and weak acids
1.10.2 Hydrochloric acid as a strong acid (HCl → H⁺ + Cl⁻)
1.10.3 Ethanoic acid as a weak acid (CH₃COOH ⇌ H⁺ + CH₃COO⁻)
1.11.1 Classify oxides as acidic (e.g., SO₂, CO₂) or basic (e.g., CuO, CaO)
1.11.2 Definition and examples of amphoteric oxides (Al₂O₃, ZnO)
2. Experimental Techniques and Chemical Analysis
3. Chemical Reactions
3.1.1 Collision theory: kinetic energy of particles
3.1.2 Collision theory: activation energy
3.1.3 How changes in conditions affect reaction rate
3.1.4 How catalysts lower activation energy
3.1.5 Effect of concentration on reaction rate
3.1.6 Effect of gas pressure on reaction rate
3.1.7 Effect of surface area on reaction rate
3.1.8 Effect of temperature on reaction rate
3.1.9 Effect of catalysts on reaction rate
3.1.10 Definition of catalyst and its role
3.1.11 Methods to measure reaction rate
3.1.12 Interpret graphs of reaction rates
3.1.13 Collision theory: particle concentration
3.1.14 Collision theory: frequency of collisions
3.2.1 Definition of reversible reactions
3.2.2 Effect of conditions on reaction direction
3.2.3 Definition of dynamic equilibrium
3.2.4 Effect of temperature on equilibrium
3.2.5 Effect of pressure on equilibrium
3.2.6 Effect of concentration on equilibrium
3.2.7 Effect of catalyst on equilibrium
3.2.8 Haber process equation (N₂ + 3H₂ ⇌ 2NH₃)
3.2.9 Raw materials for Haber process
3.2.10 Conditions used in Haber process
3.2.11 Contact process equation (2SO₂ + O₂ ⇌ 2SO₃)
3.2.12 Raw materials for Contact process
3.2.13 Conditions used in Contact process
3.2.14 Why contact process conditions are used
3.3.1 Definition of redox reactions
3.3.2 Oxidation as gain of oxygen, reduction as loss of oxygen
3.3.3 Identify redox reactions using oxygen gain/loss
3.3.4 Definition of oxidation and reduction in terms of electrons
3.3.5 Oxidation and reduction in terms of oxidation numbers
3.3.6 Redox reactions using color changes (MnO₄⁻, I₂)
3.3.7 Definition of oxidizing and reducing agents
3.3.8 Identify oxidizing and reducing agents in reactions
3.3.9 Use of oxidation numbers to identify redox reactions
3.4.1 Difference between physical and chemical changes
4. Metals
4.1.1 Barrier methods (painting, greasing, plastic coating)
4.1.2 How barriers prevent rust (oxygen/water exclusion)
4.1.3 Galvanizing with zinc as protection
4.1.4 Sacrificial protection using reactive metals
4.1.5 Conditions needed for rusting of iron
4.2.1 Ease of metal extraction depends on reactivity
4.2.2 Blast furnace: iron extracted from hematite
4.2.3 Role of carbon and carbon monoxide in reduction
4.2.4 Thermal decomposition of limestone in furnace
4.2.5 Formation of slag (removes impurities)
4.2.6 Aluminium extraction from bauxite by electrolysis
4.2.7 Blast furnace equations: C + O₂ → CO₂, C + CO₂ → 2CO, Fe₂O₃ + 3CO → 2Fe + 3CO₂
4.2.8 Aluminium extraction: role of cryolite
4.2.9 Why carbon anodes must be replaced in aluminium extraction
4.2.10 Reactions at electrodes in aluminium extraction
4.3.1 Compare physical properties of metals and non-metals
4.3.2 Reactions of metals with acids, water, and oxygen
4.4.1 Uses of aluminium (aircraft, electrical cables, food containers)
4.4.2 Uses of copper (electrical wiring, ductility)
4.5.1 Definition of alloys (mixture of metal with other elements)
4.5.2 Examples: brass (Cu + Zn), stainless steel (Fe + Cr + Ni + C)
4.5.3 Alloys are harder and stronger than pure metals
4.5.4 Why alloys are harder (atoms disrupt layers)
4.6.1 Order of reactivity: K, Na, Ca, Mg, Al, C, Zn, Fe, H, Cu, Ag, Au
4.6.2 Reactions of K, Na, Ca with water
4.6.3 Reaction of Mg with steam
4.6.4 Reactions of Mg, Zn, Fe, Cu, Ag, Au with acids
4.6.5 Explain reactivity based on ion formation
4.6.6 Displacement reactions of metals
4.6.7 Why aluminium appears unreactive (oxide layer)
5. Stoichiometry
6. Organic Chemistry
6.1.1 Definition of saturated and unsaturated compounds
6.1.2 Draw and interpret displayed formulae
6.1.3 General formulae for alkanes, alkenes, alcohols, carboxylic acids
6.1.4 Definition of functional groups
6.1.5 Definition of a homologous series
6.3.1 Names of fossil fuels: coal, natural gas, petroleum
6.3.2 Methane as the main component of natural gas
6.3.3 Definition of hydrocarbons
6.3.4 Petroleum as a mixture of hydrocarbons
6.3.5 Fractional distillation of petroleum
6.3.6 Properties of fractions (volatility, boiling points, viscosity)
6.3.7 Uses of petroleum fractions (refinery gas, gasoline, naphtha, kerosene, diesel, fuel oil, lubricatin
6.4.1 Single covalent bonds in alkanes (saturated hydrocarbons)
6.4.2 Alkanes are generally unreactive except for combustion and substitution
6.5.1 Double carbon-carbon bonds in alkenes (unsaturated hydrocarbons)
6.5.2 Cracking of large alkanes to produce alkenes and hydrogen
6.5.3 Test for unsaturation using bromine water
6.6.1 Ethanol production by fermentation of glucose
6.6.2 Ethanol production by catalytic hydration of ethene
6.6.3 Combustion of ethanol
6.6.4 Uses of ethanol (solvent, fuel)
6.7.1 Reactions of ethanoic acid with metals, bases, and carbonates
6.7.2 Formation of ethanoic acid by oxidation of ethanol
6.7.3 Ester formation from carboxylic acids and alcohols
6.8.1 Definition of polymers and monomers
6.8.2 Formation of poly(ethene) by addition polymerization
6.8.3 Plastics are made from polymers
6.8.4 Environmental issues of plastics (landfills, ocean waste, toxic gases)
7. States of Matter
8. The Periodic Table
9. Atoms, Elements, and Compounds
9.1.1 Difference between elements, compounds, and mixtures
9.2.1 Structure of an atom (nucleus, electrons, shells)
9.2.2 Relative charges and masses of subatomic particles
9.2.3 Definition of atomic number (proton number)
9.2.4 Definition of mass number (nucleon number)
9.2.5 Electronic configuration for elements (1-20)
9.2.6 Full outer shell in noble gases (Group VIII)
9.2.7 Relation of outer electrons to group number
9.2.8 Relation of electron shells to period number
9.3.1 Definition of isotopes
9.3.2 Symbols of isotopes (e.g., ₆¹²C, ₁₇³⁵Cl⁻)
9.3.3 Why isotopes have the same chemical properties
9.3.4 Calculating relative atomic mass from isotopes
9.4.1 Formation of cations and anions
9.4.2 Definition of ionic bonding
9.4.3 Formation of ionic bonds (dot-and-cross diagrams)
9.4.4 Properties of ionic compounds
9.4.5 Giant lattice structure in ionic compounds
9.5.1 Definition of covalent bonding
9.5.2 Formation of simple covalent molecules (H₂, Cl₂, H₂O, CH₄, NH₃, HCl)
9.5.3 Properties of simple molecular compounds
9.5.4 Dot-and-cross diagrams for CH₃OH, C₂H₄, O₂, CO₂, N₂
9.6.1 Structure of graphite and diamond
9.6.2 Uses of graphite and diamond
9.6.3 Structure of silicon(IV) oxide (SiO₂)
9.6.4 Similarities between diamond and SiO₂
9.7.1 Definition of metallic bonding
9.7.2 Properties of metals (conductivity, malleability, ductility)
10. Chemistry of the Environment
10.1.1 Test for water using cobalt(II) chloride and copper(II) sulfate
10.1.2 Test for purity of water using melting and boiling points
10.1.3 Why distilled water is used in chemistry
10.1.4 Substances in natural water (oxygen, metals, plastics, sewage, microbes, fertilizers)
10.1.5 Harmful substances in water (toxins, nitrates, phosphates, microbes)
10.1.6 Water treatment: sedimentation and filtration
10.1.7 Water treatment: activated carbon for odor removal
10.1.8 Water treatment: chlorination to kill microbes
10.2.1 Use of ammonium salts and nitrates as fertilizers
10.2.2 NPK fertilizers for nitrogen, phosphorus, potassium supply
10.3.1 Composition of clean, dry air (78% N₂, 21% O₂, trace gases)
10.3.2 Sources of CO₂ (fuel combustion), CO (incomplete combustion), CH₄ (decomposition, digestion)
10.3.3 Sources of NOx (car engines), SO₂ (fossil fuels)
10.3.4 Effects of CO₂ and CH₄ (global warming, climate change)
10.3.5 Effects of CO (toxic), particulates (respiratory issues)
10.3.6 Effects of NOx and SO₂ (acid rain, respiratory problems)
10.3.7 How CO₂ and CH₄ cause global warming (heat absorption)
10.3.8 Ways to reduce climate change (trees, hydrogen, renewables)
10.3.9 Reducing acid rain (catalytic converters, low-sulfur fuels, flue gas desulfurization)
10.3.10 Photosynthesis reaction: CO₂ + H₂O → glucose + O₂
10.3.11 Photosynthesis equation: 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂
10.3.12 How NOx forms in car engines
10.3.13 How catalytic converters remove NOx (2CO + 2NO → 2CO₂ + N₂)
11. Electrochemistry
Naming and drawing methane, ethane, ethene, ethanol, ethanoic acid

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Naming and Drawing Methane, Ethane, Ethene, Ethanol, Ethanoic Acid

Introduction

Understanding the nomenclature and structural representation of organic compounds is fundamental in the study of organic chemistry. This article delves into the naming and drawing of five essential organic molecules: methane, ethane, ethene, ethanol, and ethanoic acid. Tailored for Cambridge IGCSE Chemistry (0620 Core), this comprehensive guide equips students with the necessary skills to accurately identify and depict these compounds, forming a solid foundation for further exploration in organic chemistry.

Key Concepts

Methane (CH₄)

Methane is the simplest alkane, consisting of one carbon atom bonded to four hydrogen atoms. Its molecular formula is CH4, and it is the primary component of natural gas. Methane is a colorless, odorless gas under standard conditions and plays a significant role in energy production and environmental discussions regarding greenhouse gases.

Structural Representation:

The structural formula of methane is depicted as:

Methane Structural Formula

Alternatively, the Lewis structure shows the carbon atom bonded to four hydrogen atoms with single covalent bonds:

Methane Lewis Structure

Ethane (C₂H₆)

Ethane is the simplest alkane with two carbon atoms connected by a single bond, each carbon atom bonded to three hydrogen atoms. Its molecular formula is C2H6. Ethane is a vital component of natural gas and serves as a precursor in the production of other chemicals like ethylene (ethene) through cracking processes.

Structural Representation:

The structural formula of ethane is:

Ethane Structural Formula

The Lewis structure illustrates two carbon atoms connected by a single bond, each bonded to three hydrogen atoms:

Ethane Lewis Structure

Ethene (C₂H₄)

Ethene, also known as ethylene, is the simplest alkene with the molecular formula C2H4. It contains a carbon-carbon double bond, making it unsaturated. Ethene is a crucial industrial chemical used in the production of polymers such as polyethylene, which is widely used in manufacturing plastics.

Structural Representation:

The structural formula of ethene is:

Ethene Structural Formula

The Lewis structure highlights the double bond between the two carbon atoms:

Ethene Lewis Structure

Ethanol (C₂H₆O)

Ethanol, commonly known as alcohol, has the molecular formula C2H6O. It is the simplest alcohol, characterized by a hydroxyl (-OH) functional group attached to a two-carbon chain. Ethanol is widely used as a solvent, in beverages, and as a fuel additive.

Structural Representation:

The structural formula of ethanol is:

Ethanol Structural Formula

The Lewis structure emphasizes the hydroxyl group bonded to the second carbon:

Ethanol Lewis Structure

Ethanoic Acid (CH₃COOH)

Ethanoic acid, also known as acetic acid, has the molecular formula CH3COOH. It is the simplest carboxylic acid, featuring a carboxyl (-COOH) functional group. Ethanoic acid is prevalent in vinegar and serves as a chemical reagent in various industrial processes.

Structural Representation:

The structural formula of ethanoic acid is:

Ethanoic Acid Structural Formula

The Lewis structure highlights the carboxyl group attached to the methyl group:

Ethanoic Acid Lewis Structure

Advanced Concepts

IUPAC Nomenclature Rules

The International Union of Pure and Applied Chemistry (IUPAC) provides a systematic method for naming organic compounds. Understanding these rules is essential for accurately naming molecules like methane, ethane, ethene, ethanol, and ethanoic acid.

Alkanes: Alkanes are saturated hydrocarbons with single bonds. The general formula is CnH2n+2. Naming involves identifying the longest carbon chain and assigning numbers to substituents. For example, ethane (C2H6) indicates two carbon atoms with single bonds.

Alkenes: Alkenes are unsaturated hydrocarbons with one or more double bonds. The general formula is CnH2n. In ethene (C2H4), the double bond is indicated by the suffix '-ene'.

Alcohols: Alcohols contain a hydroxyl (-OH) group. The suffix '-ol' is added to the parent hydrocarbon name. Ethanol (C2H6O) reflects the presence of the hydroxyl group on the second carbon atom.

Carboxylic Acids: Carboxylic acids feature a carboxyl (-COOH) group. The suffix '-oic acid' is used. Ethanoic acid (CH3COOH) is derived from ethane, indicating a two-carbon structure with a carboxylic acid functional group.

Structural Isomerism

Structural isomers are compounds with the same molecular formula but different structural arrangements. For example, ethanol and dimethyl ether both have the molecular formula C2H6O but differ in the arrangement of their atoms.

In ethanol, the hydroxyl group is bonded to the first carbon, whereas in dimethyl ether, an oxygen atom bridges two methyl groups. Understanding structural isomerism is crucial for predicting the physical and chemical properties of organic compounds.

Functional Groups and Their Influence

Functional groups are specific groups of atoms within molecules that determine the characteristic chemical reactions of those molecules. The presence of different functional groups, such as hydroxyl in ethanol and carboxyl in ethanoic acid, imparts distinct properties and reactivities to the compounds.

For instance, alcohols like ethanol can engage in hydrogen bonding, making them more soluble in water compared to alkanes like methane. Carboxylic acids exhibit acidic properties due to the ability to donate a proton from the hydroxyl group.

Stereochemistry in Simple Organic Molecules

While methane and ethane lack stereoisomerism due to their symmetrical structures, ethene exhibits cis-trans isomerism when substituted with different groups. However, in its simplest form (C2H4), ethene is planar and lacks stereoisomers.

Understanding stereochemistry is vital for analyzing more complex organic molecules where spatial arrangements affect the molecules' physical and chemical properties.

Physical and Chemical Properties

The physical and chemical properties of methane, ethane, ethene, ethanol, and ethanoic acid vary significantly due to differences in their molecular structures and functional groups.

  • Methane: Nonpolar, low boiling point, highly flammable.
  • Ethane: Nonpolar, slightly higher boiling point than methane, used as a fuel.
  • Ethene: Unsaturated, participates in addition reactions, used in polymerization.
  • Ethanol: Polar, capable of hydrogen bonding, miscible with water, used as a solvent and in beverages.
  • Ethanoic Acid: Polar, acidic, capable of hydrogen bonding, used in food preservation and as a reagent.

These properties influence their applications and reactivity in various chemical processes.

Reactivity and Chemical Reactions

Each of the compounds discussed exhibits unique reactivity patterns based on their functional groups and bonding.

  • Methane: Undergoes substitution reactions, such as halogenation, typically requiring UV light.
  • Ethane: Similar to methane, participates in substitution reactions with halogens.
  • Ethene: Engages in addition reactions, such as hydrogenation, halogenation, and hydrohalogenation.
  • Ethanol: Undergoes oxidation to form ethanoic acid and can participate in esterification reactions.
  • Ethanoic Acid: Undergoes neutralization reactions with bases to form acetate salts and can participate in esterification.

Understanding these reactions is essential for predicting product formation and manipulating chemical pathways in organic synthesis.

Applications in Industry and Everyday Life

The compounds methane, ethane, ethene, ethanol, and ethanoic acid have widespread applications across various industries and everyday life.

  • Methane: Used as a fuel for heating, electricity generation, and as a feedstock for producing hydrogen and other chemicals.
  • Ethane: Primarily used in the petrochemical industry to produce ethylene through cracking.
  • Ethene: Essential in the production of polyethylene plastics, ethylene oxide, and other polymers.
  • Ethanol: Utilized in alcoholic beverages, as a solvent in pharmaceuticals, and as an additive in fuels (bioethanol).
  • Ethanoic Acid: Employed in the food industry as vinegar, in the production of synthetic fibers, and as a chemical reagent.

These applications highlight the significance of these compounds in both industrial processes and daily products.

Environmental Impact

The environmental implications of these organic compounds vary based on their usage and reactivity.

  • Methane: A potent greenhouse gas contributing to climate change.
  • Ethane: Released during natural gas extraction, contributing to greenhouse gas effects.
  • Ethene: Used in manufacturing plastics, leading to plastic pollution.
  • Ethanol: Considered a renewable fuel but can lead to evaporation losses and air quality issues.
  • Ethanoic Acid: Biodegradable but excessive use can lead to environmental acidity changes.

Awareness of these impacts is crucial for developing sustainable practices in the production and usage of these compounds.

Safety and Handling

Proper safety measures are essential when handling these organic compounds due to their varying degrees of toxicity and flammability.

  • Methane: Highly flammable; risk of explosion in confined spaces.
  • Ethane: Flammable gas; requires proper ventilation to prevent accumulation.
  • Ethene: Flammable; exposure can lead to respiratory issues.
  • Ethanol: Flammable liquid; ingestion can be toxic.
  • Ethanoic Acid: Corrosive; can cause burns upon contact with skin and eyes.

Adherence to safety protocols ensures safe handling and minimizes risks associated with these compounds.

Comparison Table

Compound Structural Formula Functional Group Uses
Methane (CH₄) Methane Structural Formula Alkane Fuel, energy source
Ethane (C₂H₆) Ethane Structural Formula Alkane Production of ethylene
Ethene (C₂H₄) Ethene Structural Formula Alkene Polyethylene production
Ethanol (C₂H₆O) Ethanol Structural Formula Alcohol Solvent, beverages, fuel additive
Ethanoic Acid (CH₃COOH) Ethanoic Acid Structural Formula Carboxylic Acid Food preservation, chemical reagent

This table highlights the structural formulas, functional groups, and primary uses of each compound, providing a clear comparison of their key aspects.

Summary and Key Takeaways

  • Methane, ethane, ethene, ethanol, and ethanoic acid are fundamental organic compounds with distinct structures and properties.
  • Accurate nomenclature is essential for identifying and communicating about these molecules in chemistry.
  • Functional groups play a crucial role in determining the reactivity and applications of organic compounds.
  • Understanding structural isomerism and stereochemistry enhances the comprehension of molecular behavior.
  • These compounds have diverse applications, from energy production to industrial manufacturing, impacting various aspects of daily life and the environment.

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Examiner Tip
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Tips

To remember functional groups, associate the suffix with the group: '-ane' for alkanes (single bonds), '-ene' for alkenes (double bonds), '-ol' for alcohols (OH group), and '-oic acid' for carboxylic acids (COOH group). Use molecular models or drawing software to practice structural and Lewis structures regularly. Additionally, create flashcards for each compound's name, formula, and functional group to reinforce memorization for exams.

Did You Know
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Did You Know

Methane, the simplest hydrocarbon, is not only a key energy source but also the primary component of natural gas used worldwide for heating and electricity. Ethanol serves as a renewable fuel alternative, significantly reducing carbon emissions when used in vehicles. Additionally, ethanoic acid is essential in the production of polymers and textiles, playing a crucial role in the manufacturing industry.

Common Mistakes
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Common Mistakes

One common mistake is confusing the structural and Lewis formulas of organic compounds. For example, students often misrepresent ethene's double bond in its Lewis structure. Another frequent error is incorrect nomenclature, such as labeling ethanoic acid as ethanol, leading to confusion between functional groups. Lastly, misidentifying functional groups, like confusing the hydroxyl group in ethanol with the carboxyl group in ethanoic acid, can hinder accurate chemical analysis.

FAQ

What is the molecular formula of ethanoic acid?
The molecular formula of ethanoic acid is C₂H₄O₂.
How does ethene differ from ethane?
Ethene is an alkene with a carbon-carbon double bond (C₂H₄), making it unsaturated, whereas ethane is an alkane with only single bonds (C₂H₆), making it saturated.
Why is ethanol considered a versatile compound?
Ethanol is versatile because it is used in beverages, as a solvent, in hand sanitizers, and as a biofuel, showcasing its wide range of applications across different industries.
What functional group characterizes ethanoic acid?
Ethanoic acid is characterized by the carboxyl group (-COOH).
How is methane produced naturally?
Methane is primarily produced naturally through the decomposition of organic matter in anaerobic environments, such as wetlands and the digestive systems of ruminants.
Can ethene be used to ripen fruits?
Yes, ethene acts as a plant hormone that promotes the ripening of fruits by stimulating the production of enzymes that break down cell walls.
1. Acids, Bases, and Salts
1.1.1 Making soluble salts by reaction with excess metal
1.1.2 Making soluble salts by reaction with excess base
1.1.3 Making soluble salts by reaction with excess carbonate
1.1.4 Making soluble salts by titration
1.2.1 Solubility of sodium, potassium, ammonium salts
1.2.2 Solubility of nitrates, chlorides, sulfates, carbonates, and hydroxides
1.3.1 Definition of hydrated and anhydrous substances
1.4.1 Making insoluble salts by precipitation
1.5.1 Definition and examples (CuSO₄·5H₂O, CoCl₂·6H₂O)
1.6.1 Reaction of acids with metals, bases, and carbonates
1.6.2 Effect of acids on litmus, thymolphthalein, and methyl orange
1.7.1 Definition of bases and alkalis
1.7.2 Reactions of bases with acids and ammonium salts
1.7.3 Effect of alkalis on litmus, thymolphthalein, and methyl orange
1.8.1 H⁺ ions in acids, OH⁻ ions in alkalis
1.8.2 Compare acidity and alkalinity using pH and indicators
1.9.1 Reaction of acids with alkalis to form water
1.10.1 Definition of strong and weak acids
1.10.2 Hydrochloric acid as a strong acid (HCl → H⁺ + Cl⁻)
1.10.3 Ethanoic acid as a weak acid (CH₃COOH ⇌ H⁺ + CH₃COO⁻)
1.11.1 Classify oxides as acidic (e.g., SO₂, CO₂) or basic (e.g., CuO, CaO)
1.11.2 Definition and examples of amphoteric oxides (Al₂O₃, ZnO)
2. Experimental Techniques and Chemical Analysis
3. Chemical Reactions
3.1.1 Collision theory: kinetic energy of particles
3.1.2 Collision theory: activation energy
3.1.3 How changes in conditions affect reaction rate
3.1.4 How catalysts lower activation energy
3.1.5 Effect of concentration on reaction rate
3.1.6 Effect of gas pressure on reaction rate
3.1.7 Effect of surface area on reaction rate
3.1.8 Effect of temperature on reaction rate
3.1.9 Effect of catalysts on reaction rate
3.1.10 Definition of catalyst and its role
3.1.11 Methods to measure reaction rate
3.1.12 Interpret graphs of reaction rates
3.1.13 Collision theory: particle concentration
3.1.14 Collision theory: frequency of collisions
3.2.1 Definition of reversible reactions
3.2.2 Effect of conditions on reaction direction
3.2.3 Definition of dynamic equilibrium
3.2.4 Effect of temperature on equilibrium
3.2.5 Effect of pressure on equilibrium
3.2.6 Effect of concentration on equilibrium
3.2.7 Effect of catalyst on equilibrium
3.2.8 Haber process equation (N₂ + 3H₂ ⇌ 2NH₃)
3.2.9 Raw materials for Haber process
3.2.10 Conditions used in Haber process
3.2.11 Contact process equation (2SO₂ + O₂ ⇌ 2SO₃)
3.2.12 Raw materials for Contact process
3.2.13 Conditions used in Contact process
3.2.14 Why contact process conditions are used
3.3.1 Definition of redox reactions
3.3.2 Oxidation as gain of oxygen, reduction as loss of oxygen
3.3.3 Identify redox reactions using oxygen gain/loss
3.3.4 Definition of oxidation and reduction in terms of electrons
3.3.5 Oxidation and reduction in terms of oxidation numbers
3.3.6 Redox reactions using color changes (MnO₄⁻, I₂)
3.3.7 Definition of oxidizing and reducing agents
3.3.8 Identify oxidizing and reducing agents in reactions
3.3.9 Use of oxidation numbers to identify redox reactions
3.4.1 Difference between physical and chemical changes
4. Metals
4.1.1 Barrier methods (painting, greasing, plastic coating)
4.1.2 How barriers prevent rust (oxygen/water exclusion)
4.1.3 Galvanizing with zinc as protection
4.1.4 Sacrificial protection using reactive metals
4.1.5 Conditions needed for rusting of iron
4.2.1 Ease of metal extraction depends on reactivity
4.2.2 Blast furnace: iron extracted from hematite
4.2.3 Role of carbon and carbon monoxide in reduction
4.2.4 Thermal decomposition of limestone in furnace
4.2.5 Formation of slag (removes impurities)
4.2.6 Aluminium extraction from bauxite by electrolysis
4.2.7 Blast furnace equations: C + O₂ → CO₂, C + CO₂ → 2CO, Fe₂O₃ + 3CO → 2Fe + 3CO₂
4.2.8 Aluminium extraction: role of cryolite
4.2.9 Why carbon anodes must be replaced in aluminium extraction
4.2.10 Reactions at electrodes in aluminium extraction
4.3.1 Compare physical properties of metals and non-metals
4.3.2 Reactions of metals with acids, water, and oxygen
4.4.1 Uses of aluminium (aircraft, electrical cables, food containers)
4.4.2 Uses of copper (electrical wiring, ductility)
4.5.1 Definition of alloys (mixture of metal with other elements)
4.5.2 Examples: brass (Cu + Zn), stainless steel (Fe + Cr + Ni + C)
4.5.3 Alloys are harder and stronger than pure metals
4.5.4 Why alloys are harder (atoms disrupt layers)
4.6.1 Order of reactivity: K, Na, Ca, Mg, Al, C, Zn, Fe, H, Cu, Ag, Au
4.6.2 Reactions of K, Na, Ca with water
4.6.3 Reaction of Mg with steam
4.6.4 Reactions of Mg, Zn, Fe, Cu, Ag, Au with acids
4.6.5 Explain reactivity based on ion formation
4.6.6 Displacement reactions of metals
4.6.7 Why aluminium appears unreactive (oxide layer)
5. Stoichiometry
6. Organic Chemistry
6.1.1 Definition of saturated and unsaturated compounds
6.1.2 Draw and interpret displayed formulae
6.1.3 General formulae for alkanes, alkenes, alcohols, carboxylic acids
6.1.4 Definition of functional groups
6.1.5 Definition of a homologous series
6.3.1 Names of fossil fuels: coal, natural gas, petroleum
6.3.2 Methane as the main component of natural gas
6.3.3 Definition of hydrocarbons
6.3.4 Petroleum as a mixture of hydrocarbons
6.3.5 Fractional distillation of petroleum
6.3.6 Properties of fractions (volatility, boiling points, viscosity)
6.3.7 Uses of petroleum fractions (refinery gas, gasoline, naphtha, kerosene, diesel, fuel oil, lubricatin
6.4.1 Single covalent bonds in alkanes (saturated hydrocarbons)
6.4.2 Alkanes are generally unreactive except for combustion and substitution
6.5.1 Double carbon-carbon bonds in alkenes (unsaturated hydrocarbons)
6.5.2 Cracking of large alkanes to produce alkenes and hydrogen
6.5.3 Test for unsaturation using bromine water
6.6.1 Ethanol production by fermentation of glucose
6.6.2 Ethanol production by catalytic hydration of ethene
6.6.3 Combustion of ethanol
6.6.4 Uses of ethanol (solvent, fuel)
6.7.1 Reactions of ethanoic acid with metals, bases, and carbonates
6.7.2 Formation of ethanoic acid by oxidation of ethanol
6.7.3 Ester formation from carboxylic acids and alcohols
6.8.1 Definition of polymers and monomers
6.8.2 Formation of poly(ethene) by addition polymerization
6.8.3 Plastics are made from polymers
6.8.4 Environmental issues of plastics (landfills, ocean waste, toxic gases)
7. States of Matter
8. The Periodic Table
9. Atoms, Elements, and Compounds
9.1.1 Difference between elements, compounds, and mixtures
9.2.1 Structure of an atom (nucleus, electrons, shells)
9.2.2 Relative charges and masses of subatomic particles
9.2.3 Definition of atomic number (proton number)
9.2.4 Definition of mass number (nucleon number)
9.2.5 Electronic configuration for elements (1-20)
9.2.6 Full outer shell in noble gases (Group VIII)
9.2.7 Relation of outer electrons to group number
9.2.8 Relation of electron shells to period number
9.3.1 Definition of isotopes
9.3.2 Symbols of isotopes (e.g., ₆¹²C, ₁₇³⁵Cl⁻)
9.3.3 Why isotopes have the same chemical properties
9.3.4 Calculating relative atomic mass from isotopes
9.4.1 Formation of cations and anions
9.4.2 Definition of ionic bonding
9.4.3 Formation of ionic bonds (dot-and-cross diagrams)
9.4.4 Properties of ionic compounds
9.4.5 Giant lattice structure in ionic compounds
9.5.1 Definition of covalent bonding
9.5.2 Formation of simple covalent molecules (H₂, Cl₂, H₂O, CH₄, NH₃, HCl)
9.5.3 Properties of simple molecular compounds
9.5.4 Dot-and-cross diagrams for CH₃OH, C₂H₄, O₂, CO₂, N₂
9.6.1 Structure of graphite and diamond
9.6.2 Uses of graphite and diamond
9.6.3 Structure of silicon(IV) oxide (SiO₂)
9.6.4 Similarities between diamond and SiO₂
9.7.1 Definition of metallic bonding
9.7.2 Properties of metals (conductivity, malleability, ductility)
10. Chemistry of the Environment
10.1.1 Test for water using cobalt(II) chloride and copper(II) sulfate
10.1.2 Test for purity of water using melting and boiling points
10.1.3 Why distilled water is used in chemistry
10.1.4 Substances in natural water (oxygen, metals, plastics, sewage, microbes, fertilizers)
10.1.5 Harmful substances in water (toxins, nitrates, phosphates, microbes)
10.1.6 Water treatment: sedimentation and filtration
10.1.7 Water treatment: activated carbon for odor removal
10.1.8 Water treatment: chlorination to kill microbes
10.2.1 Use of ammonium salts and nitrates as fertilizers
10.2.2 NPK fertilizers for nitrogen, phosphorus, potassium supply
10.3.1 Composition of clean, dry air (78% N₂, 21% O₂, trace gases)
10.3.2 Sources of CO₂ (fuel combustion), CO (incomplete combustion), CH₄ (decomposition, digestion)
10.3.3 Sources of NOx (car engines), SO₂ (fossil fuels)
10.3.4 Effects of CO₂ and CH₄ (global warming, climate change)
10.3.5 Effects of CO (toxic), particulates (respiratory issues)
10.3.6 Effects of NOx and SO₂ (acid rain, respiratory problems)
10.3.7 How CO₂ and CH₄ cause global warming (heat absorption)
10.3.8 Ways to reduce climate change (trees, hydrogen, renewables)
10.3.9 Reducing acid rain (catalytic converters, low-sulfur fuels, flue gas desulfurization)
10.3.10 Photosynthesis reaction: CO₂ + H₂O → glucose + O₂
10.3.11 Photosynthesis equation: 6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂
10.3.12 How NOx forms in car engines
10.3.13 How catalytic converters remove NOx (2CO + 2NO → 2CO₂ + N₂)
11. Electrochemistry
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