Table of Contents

Chapter 7: Acids, Bases, and Salts

The Characteristics of Acids & Bases

The characteristic properties of acids in reactions with:

Metal:

Only metals above hydrogen in the reactivity series will react with dilute acids. They form a salt and hydrogen gas. The name of the salt depends on the anion in the acid.

Carbonates:

Corresponding metal salt, CO₂, and water will be formed.

Bases:

Metal oxides and metal hydroxides are bases. When they react, a neutralization reaction occurs. A salt and water will be produced.

Acids & base effect on:

Litmus paper

Thymolphthalein

Methyl orange

Indicator	Color in acid	Color in alkali
Litmus	Red	Blue
Phenolphthalein	Colorless	Pink
Methyl orange	Red	Yellow

Bases are oxides or hydroxides of metals, and alkalis are soluble bases.

Base reaction with:

Ammonium salts:

The ammonium salt decomposes when warmed with a base. Ammonia is volatile and easily displaced from the salt by another alkali, so a salt, water, and ammonia are formed.

Aqueous solutions of acid contain H⁺ ions, and aqueous solutions of alkalis contain OH⁻ ions.

Acids are proton donors, bases are proton acceptors.

A strong acid is an acid that is completely dissociated in aqueous solution, and a weak acid is an acid that is partially dissociated in aqueous solution.

Hydrochloric acid is a strong acid:

$\text{HCl (aq)} \rightarrow \text{H}^+ \text{(aq)} + \text{Cl}^- \text{(aq)}$

Ethanoic acid is a weak acid:

$\text{CH}_3\text{COOH (aq)} \leftrightarrow \text{H}^+ \text{(aq)} + \text{CH}_3\text{COO}^- \text{(aq)}$

Neutralisation Reactions

When acids are added to water, they form positively charged hydrogen ions (H⁺).

When alkalis are added to water, they form negatively charged hydroxide ions (OH⁻).

Neutralisation occurs when acid reacts with an alkali; the H⁺ and OH⁻ ions react to produce water H₂O.

$\text{H}^+ \text{(aq)} + \text{OH}^- \text{(aq)} \rightarrow \text{H}_2\text{O (l)}$

Color & pH with Universal Indicator Paper

The pH scale is a numerical scale which is used to show how acidic or alkaline a solution is.

The scale goes from most acidic (1) to most alkaline (14). All pH below 7 is acidic, above 7 is alkaline, and 7 is neutral.

The more H⁺ ions, the stronger the acid, lower pH.

The more OH⁻ ions, the stronger the alkali, higher pH.

The pH scale is logarithmic, meaning each change of 1 on the scale represents a change by a factor of 10.

The universal indicator is a mixture of different indicators to measure pH.

Preparation of Salts

The preparation, separation, and purification of soluble salts by reaction of an acid with:

An Alkali by titration:

$\text{acid} + \text{alkali} \rightarrow \text{salt} + \text{water}$

Carry out titration to work out the precise volume of acid and alkali for neutralization to occur.

Mix the appropriate volumes of acid and alkali so we are left with a solution containing salt and water.

Evaporate the water by gently heating the solution, allow to cool, and the soluble salt will crystallize.

Filter the solid salt, and dry.

Excess metal:

$\text{acid} + \text{metal} \rightarrow \text{salt} + \text{hydrogen}$

Any metal more reactive than hydrogen (e.g., K, Na, Ca, Mg, Al, Zn, etc.) reacts with acid to form a salt and hydrogen gas.

Pour acid into a beaker, add a metal strip, and keep adding until the metal is in excess. Remove excess by filtration, evaporate the filtrate until salt is left behind, and dry.

Insoluble bases:

$\text{acid} + \text{base} \rightarrow \text{salt} + \text{water}$

Add excess metal oxide to acid in a beaker, warm the beaker, and filter out the excess metal oxide. Evaporate the water and dry the salt.

Insoluble carbonate:

$\text{Acid} + \text{Carbonate} \rightarrow \text{Salt} + \text{water} + \text{carbon dioxide}$

Add acid to beaker, warm it, add carbonate until it is in excess, filter the excess, evaporate the filtrate until crystallization point, dry the salt.

Preparation of insoluble salt with the precipitation method:

Place a known volume of soluble salt in a beaker. Keep adding another soluble salt and stir to mix until no more precipitate forms. Filter the precipitate; the residue is the insoluble salt. Wash the residue with water and dry the salt.

Solubility rules for salts:

Sodium, potassium, & ammonium salts are soluble.

Nitrates are soluble.

Chlorides are soluble, except lead & silver.

Sulfates are soluble, except barium, calcium, & lead.

Carbonates are insoluble except sodium, potassium, & ammonium.

Hydroxides are insoluble except sodium, potassium, ammonium, and calcium (partially).

A hydrated substance is a substance that is chemically combined with water.

An anhydrous substance is a substance that contains no water.

The term water of crystallization refers to the water molecules present in hydrated crystals, including

$\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$ and $\text{CoCl}_2 \cdot 6\text{H}_2\text{O}$
indicating 6 water molecules
are attached/bound to the
cobalt chloride.

Type of Salt	Soluble Salts	Insoluble Salts
Carbonates	Sodium carbonate, Potassium carbonate, Ammonium carbonate	All other carbonates
Nitrates	All nitrates	None
Chlorides	All chlorides except	Lead(II) chloride, Silver chloride
Sulfates	All sulfates except	Calcium sulfate (sparingly soluble), Lead(II) sulfate, Barium sulfate
Ammonium salts & Group I metal salts	All	None

Oxides

Oxides are compounds made from one or more atoms of oxygen combined with one other element. E.g., SO₂, CO₂, CuO, MgO.

Oxides can be classified as acidic or basic based on whether they are bonded to a metal or nonmetal. Some are neutral oxides like N₂O, NO, and CO.

Acidic Oxides:

Formed when a non-metal combines with oxygen.

When dissolved in water, they produce an acidic solution, e.g., CO₂, SO₂, NO₂.

Acidic Oxide	Formula	Acid Produced with Water
Sulphur trioxide	SO₃	Sulphuric acid, H₂SO₄
Sulphur dioxide	SO₂	Sulphurous acid, H₂SO₃
Carbon dioxide	CO₂	Carbonic acid, H₂CO₃
Phosphorus(V) oxide	P₄O₁₀	Phosphoric acid, H₃PO₄

Basic Oxides:

Formed when a metal combines with oxygen.

E.g., CuO, CaO.

Basic Oxide	Formula
Magnesium oxide	MgO
Sodium oxide	Na₂O
Calcium oxide	CaO
Copper(II) oxide	CuO

Amphoteric Oxides:

React with acids & bases to produce salt & water.

E.g., Al₂O₃, ZnO.

Chapter 8: The periodic table

Arrangement of elements

The periodic table is an arrangement of elements and groups in order of increasing proton number (atomic number).

Vertical columns are called groups: I-VII and Group 0 (not VIII).

Horizontal rows are called periods: 1-7.

Metallic characteristics decrease across a period, left to right, and increase down a period.

Metals occur on the left-hand side of the periodic table and non-metals on the right.

Between them lie the metalloids/semi-metals.

Electronic configuration will show the number of occupied shells, which is the same as the period, and the number of electrons on the last shell is the group number.

Group VII Properties

Halogens, the diatomic non-metals, with general trends down the group such as:

Increasing density

Decreasing reactivity

The appearance of halogens at r.t.p:

Chlorine (Cl₂): Pale yellow-green gas

Bromine (Br₂): Red-brown liquid

Iodine (I₂): Grey-black solid

In displacement reactions, the more reactive halogen atoms oxidize the less reactive halide ions, causing halide ions to lose electrons and form halogen atoms.

The halogen atoms then gain electrons to form halide ions, which are reduced. This is a redox reaction.

Transition elements

Transition elements are metals that:

Have high density

High melting point

Form colored compounds

Often act as catalysts as elements and in compounds.

Transition elements have ions with variable oxidation numbers, including iron(II) & iron(III).

Noble Gases

The noble gases are the unreactive, monatomic gases.

They all have a full outer shell of 8 electrons, so they are very stable and unreactive.

Group I Properties

Alkali metals, relatively soft metals with trends down the group:

Decreasing melting point

Increasing density

Increasing reactivity

Chapter 9: Metals

Properties of Metals:

Conduct heat & electricity

Malleable & ductile

High density & high melting point

Properties of Non-Metals:

Don’t conduct heat & electricity

Brittle & easily break

Low density & low melting point

Metals reaction to:

Dilute acids:

Salt + hydrogen formed.

Cold water & steam:

Metal hydroxide + hydrogen.

Oxygen:

Some metals don’t react, e.g., gold.

Others form a metal oxide

Uses of Metals

Aluminium in the manufacture of aircraft because of low density.

Aluminium in the manufacture of overhead electrical cables because of low density and good electrical conductivity.

Aluminium in food containers because of resistance to corrosion.

Copper in electrical wiring due to good conduction and ductility.

Extraction of Metals

The ease of obtaining metals from ore depends on their position of the metals in the reactivity series.

The extraction of iron from hematite ore in the blast furnaces takes place by:

The burning of carbon (coke) to provide heat & produce carbon dioxide:

$\text{C} + \text{O}_2 \rightarrow \text{CO}_2$

The reduction of carbon dioxide to carbon monoxide:

$\text{C} + \text{CO}_2 \rightarrow 2\text{CO}$

The reduction of iron(III) oxide by carbon monoxide:

$\text{Fe}_2\text{O}_3 + 3\text{CO} \rightarrow 2\text{Fe} + 3\text{CO}_2$

Thermal decomposition of calcium carbonate (limestone) to produce calcium oxide:

$\text{CaCO}_3 \rightarrow \text{CaO} + \text{CO}_2$

The formation of slag:

$\text{CaO} + \text{SiO}_2 \rightarrow \text{CaSiO}_3$

The main ore of aluminium is bauxite and aluminium is extracted by electrolysis.

Bauxite is first purified to produce Al₂O₃ (aluminium oxide).

Aluminium oxide is then dissolved in molten cryolite.

This is because Al₂O₃ has a high melting point which needs a lot of energy, but the mixture has a lower melting point without interfering with the reaction.

The mixture is placed in the electrolysis cell, and the carbon anodes are regularly replaced.

This is because the carbon reacts with oxygen in Al₂O₃ & are burnt away over time. The electrode is made from steel & lined with graphite.

At the cathode:

$\text{Al}^{3+} + 3\text{e}^- \rightarrow \text{Al}$

At the anode:

$2\text{O}^{2-} \rightarrow \text{O}_2 + 4\text{e}^-$

Alloys and Their Properties

An alloy is a mixture of metal with other elements including:

Brass – Copper & zinc

Stainless Steel – Iron, and other elements like chromium, nickel, and carbon.

Alloys can be harder, stronger, than pure metals and more useful. This is because the different sized atoms from different elements mean the layers can’t slide over each other.

Brass – Made of Cu & Zn but stronger than both, used for musical instruments, ornaments & door knobs.

Stainless Steel – Used in cutlery as it is hard & is resistant to rusting

Corrosion of Metals

The conditions required for rusting of iron and steel to form hydrated iron(III) oxide are iron, water, and oxygen.

Some common barrier methods to prevent rusting include painting, greasing, and coating with plastic.

Sacrificial Protection is the process of using a more reactive metal attached to a less reactive metal so the more reactive metal will corrode first. Zinc is used to galvanize.

All the barrier methods prevent the iron from being exposed to air and oxygen.

Galvanising is the process of using zinc to coat the iron. Zinc is more reactive, so it will lose electrons more easily and oxidize/rust faster, so the zinc is sacrificed.

Reactivity Series

The order of reactivity is:

Potassium, Sodium, Calcium, Magnesium, Aluminium, Carbon, Zinc, Iron, Hydrogen, Copper, Silver, Gold.

Metal atoms form positive ions by loss of electrons when they react with other substances. The tendency to lose electrons is a measure of how reactive the metal is. A metal high up in the series loses electrons more easily than the ones below it.

Displacement reactions occur when a more reactive metal displaces a less reactive metal from a compound. The more reactive metals lose electrons easier, making them better reducing agents.

Example:

Aluminium is high in the reactivity series, but in reality, it doesn’t react with water and is slow with dilute acids. This is because it reacts readily with oxygen to form a protective layer of aluminium oxide; this layer prevents reaction with water and acids.

Chapter 10: Chemistry of the environment

Water

Testing for the Presence of Water:

Anhydrous cobalt(II) chloride:

Cobalt(II) chloride turns blue to pink in contact with water.

Equation:

Anhydrous cobalt(II) chloride + water ⇌ Hydrated cobalt(II) chloride

$\text{CoCl}_2 (\text{s}) + 6\text{H}_2\text{O (l)} \rightleftharpoons \text{CoCl}_2 \cdot 6\text{H}_2\text{O (s)}$

Anhydrous copper(II) sulfate:

Copper(II) sulfate turns white to blue in contact with water.

Equation:

Anhydrous copper(II) sulfate + water ⇌ Hydrated copper(II) sulfate

$\text{CuSO}_4 (\text{s}) + 5\text{H}_2\text{O (l)} \rightarrow \text{CuSO}_4 \cdot 5\text{H}_2\text{O (s)}$

Testing the Purity of Water:

Water has a boiling point of 100°C and a melting point of 0°C.

Impure water will have a higher boiling point and will melt below 0°C.

Note: Distilled water is used in practicals rather than tap water because it has a few chemical impurities.

Water from natural sources includes substances like:

Dissolved oxygen, metal compounds, plastics, sewage, harmful microbes, nitrates from fertilizers, phosphates from fertilizers & detergents.

Some of these substances are useful because:

Dissolved oxygen is used by aquatic life for respiration.

Some metal compounds provide essential minerals for life.

Some of these substances are potentially harmful like:

Some metal compounds are toxic.

Some plastics harm aquatic life.

Sewage contains harmful microbes that cause diseases.

Nitrates & phosphates cause deoxygenation of water & harm life.

The treatment of domestic water supply:

Step 1: Sedimentation & filtration to remove solids.

Step 2: Use of carbon to remove tastes & flavors.

Step 3: Chlorination to kill all the microbes.

Fertilisers

Ammonium salts & nitrates are used as fertilizers.

NPK fertilizers provide elements: nitrogen, phosphorus, & potassium for improved plant growth.

Nitrogen – Chlorophyll & protein for healthy leaves.

Phosphorus – Promotes healthy roots.

Potassium – Promotes the growth of healthy fruit & flowers.

Common fertilizer compounds:

Ammonium nitrate (NH₄NO₃)

Potassium sulfate (K₂SO₄)

Ammonium phosphate ((NH₄)₃PO₄)

Air Quality and Climate

The composition of clean dry air has approx: 78% nitrogen, N₂, 21% oxygen O₂; the remainder is a mixture of noble gases & carbon dioxide, CO₂.

The sources of air pollutants:

CO₂ from complete combustion of carbon-containing fuels.

CO & particulates from incomplete combustion of carbon-containing fuels.

Methane from the decomposition of vegetation & waste gas from animal digesting.

Oxides of nitrogen from car engines.

Sulfur dioxide from the combustion of fossil fuels which contain sulfur compounds.

The air pollutants have adverse effects such as:

Carbon dioxide leads to increased global warming & leads to climate change.

Carbon monoxide is a toxic gas.

Particulates cause an increased risk of respiratory problems & cancer.

Methane leads to increased global warming & leads to climate change.

Oxides of nitrogen cause acid rain, photochemical smog & respiratory issues.

Sulfur dioxide causes acid rain.

The Greenhouse Effect:

The greenhouse gases like CO₂, methane, nitrous oxide, etc., absorb heat that is reflected off the ground from the sun, and this absorbed heat is re-emitted into all directions. This reduces thermal energy lost to space & increases temp on Earth, leading to global warming & climate change.

Strategies that reduce the effect of these environmental issues are:

For climate change:

Planting trees, reduction in livestock farming, decreasing the use of fossil fuels & using renewable energy.

For acid rain:

Use catalytic converters in vehicles, reduce emissions of sulfur dioxide by using low-sulfur fuels & flue gas desulfurization with calcium oxide.

Oxides of nitrogen form in car engines & are removed by catalytic converters by mixing with carbon monoxide.

$2\text{CO} + 2\text{NO} \rightarrow 2\text{CO}_2 + \text{N}_2$

Photosynthesis:

The reaction between carbon dioxide & water to make glucose and oxygen in the presence of chlorophyll and using energy from light.

Word Equation:

$\text{Carbon dioxide} + \text{Water} \rightarrow \text{Glucose} + \text{Oxygen}$

Chemical Equation:

$6\text{CO}_2 + 6\text{H}_2\text{O} \rightarrow \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2$

Chapter 11: Organic Chemistry

Formulae, Functional Groups, and Terminology

Drawing & interpreting displayed formula:

Displayed formula shows spatial arrangement of all atoms and bonds in a molecule.
E.g.,

$\text{C}_5\text{H}_{12}$

Homologous series:

A family of similar compounds with similar chemical properties due to the presence of the same functional group.

The general characteristics:

Have the same functional group.

Have the same general formula.

Differ from one member to the next by a -CH₂- unit.

Displaying a trend in physical properties.

Sharing similar chemical properties.

Saturated:

Saturated compound has molecules in which all carbon-carbon bonds are single bonds.

Unsaturated:

Unsaturated compound has molecules in which one or more carbon-carbon bonds are not single bonds.

Structural formula:

Unambiguous description of the way the atoms in a molecule are arranged.
E.g.,

$\text{CH}_2 = \text{CH}_2$

$\text{CH}_3\text{CH}_2\text{OH}, \text{CH}_3\text{COOCH}_3$

Structural isomers:

Compounds with the same molecular formula but different structural formulae.

E.g.,

$\text{C}_4\text{H}_{10}$ as $\text{CH}_3\text{CH}_2\text{CH}_2\text{CH}_3$ and $\text{CH}_3 \text{CH}(\text{CH}_3) \text{CH}_3$ $\text{C}_4\text{H}_8$

$\text{CH}_3\text{CH}_2\text{CH}=\text{CH}_2$ and $\text{CH}_3\text{CH}=\text{CHCH}_3$

Naming Organic Compounds

Carboxylic Acids

A carboxylic acid

is an organic compound containing a carboxyl group ( $\text{O=}\underset{\text{OH}}{\text{C}}$ ) bonded to an R group ( $\text{R}$ ).

Reaction of Ethanoic Acid with:

Metals – a salt of ethanoate. e.g., sodium ethanoate
( $\text{CH}_3\text{COO}\text{Na}$

$\text{Acid} + \text{metal} \rightarrow \text{Salt} + \text{hydrogen}$

Bases – a salt of ethanoate.

$\text{Acid} + \text{base} \rightarrow \text{Salt} + \text{water}$

Carbonates – a salt of ethanoate.

$\text{Acid} + \text{Carbonate} \rightarrow \text{Salt} + \text{Water} + \text{Carbon dioxide}$

Formation of Ethanoic Acid:

Ethanol + Oxygen $\rightarrow$ Ethanoic acid + Water

$\text{C}_2\text{H}_5\text{OH} + \text{O}_2 \rightarrow \text{CH}_3\text{COOH} + \text{H}_2\text{O}$

To oxidize ethanol, the oxidizing agent acidified aqueous potassium manganate (VII) $\left[\text{KMnO}_4\right]$ is added to ethanol.

$\text{C}_2\text{H}_5\text{OH} + \text{KMnO}_4 \rightarrow \text{CH}_3\text{COOH} + \text{MnO}_2 + \text{KOH} + \text{H}_2\text{O}$

Alternatively, bacteria can be used to oxidize during vinegar (ethanoic acid) production:
Acetic acid bacteria oxidize the acid and form vinegar.

Carboxylic acid reaction with alcohol using acid catalyst to form an ester:

Fischer esterification process of ester formation:

acid
Carboxylic acid + alcohol $\overset{\text{acid}}{\longrightarrow}$ ester + water

Note: R may be an H atom, alkyl, or aryl group.

Polymers

Polymers are large molecules are built up from many smaller molecules called monomers.

The molecules are covalently bonded. Different linkages exist depending on the monomers & polymerisation. Linkages can be: amide, ester link.

Poly(ethene) is formed by the addition polymerisation of ethene monomers. Addition polymerisation is where a chain adds a new monomer to a growing polymer.

Example:

For other alkenes, to find their single monomer, break their double bonds.

Addition polymerisation is the repeated addition of monomers that have double/triple bonds.

Condensation polymerisation is where there is repeated condensation of two different bi-functional or tri-functional monomers.

Eg: Polyamides from dicarboxylic acid and a diamine

Polyesters from a dicarboxylic acid and a diol

Plastics are made from polymers.
Polymeric substances can’t be digested/decomposed, so they’re non-biodegradable, which causes many environmental concerns.

Example:

Disposal in landfills accumulation in oceans formation of toxic gas when burned

Plastic made out of PET can be converted back to monomers and be re-polymerised, which allows recycling.

Proteins are natural polyamides and they are formed from amino acid monomers with the structure:

The structure of proteins:

Fuels

Fossil fuels:

Coal, natural gas, & petroleum.
Methane is the main component in natural gases.
Hydrocarbons are the compounds that contain hydrogen and carbon only.
E.g.:

$\text{CH}_4$

(methane)

etroleum is a mix of hydrocarbons. The petroleum is separated into useful components by fractional distillation.
The properties of fractions obtained change from the bottom to top of the fractionating column by:

Decreasing chain length, higher volatility, lower boiling points, lower viscosity

FRACTION	NUMBER OF CARBON ATOMS	BOILING POINT RANGE /°C
REFINERY GAS	1 – 4	BELOW 25
GASOLINE / PETROL	4 – 12	40 – 100
NAPHTHA	7 – 14	100 – 150
KEROSENE / PARAFFIN	12 – 16	150 – 240
DIESEL / GAS OIL	15 – 19	220 – 250
FUEL OIL	19 – 25	250 – 320
LUBRICATING OIL	20 – 50	300 – 350
BITUMEN	MORE THAN 70	MORE THAN 350

Alcohols

Alcohol is also called ethanol (C₂H₅OH).

It is formed by:

Fermentation of aqueous glucose (Plant Sugar) at 25-35°C in the presence of yeast and in the absence of oxygen. Ethanol can also be made by the catalytic addition of steam to ethene at 300°C and 6000 kPa (60 atm) in the presence of an acid catalyst.

Advantages & Disadvantages of:

Fermentation	Catalytic Addition of Steam to Ethene
Less energy	Continuous process
Cheap	Less labour needed
Renewable raw materials	Fast process
Labour intensive	Expensive
Time consuming	High energy needed
	Non-renewable raw materials

Combustion of Ethanol

C₂H₅OH + 3O₂ (g) → 2CO₂ (g) + 3H₂O (l)

They combust completely in the presence of water. When oxygen is scarce, they produce water and carbon monoxide or carbon (soot).

Ethanol can be used as a:

Solvent

Fuel

Alkanes

The bonding in alkanes is single covalent, and alkanes are saturated hydrocarbons.

Alkanes are generally unreactive (with the exception of combustion and chlorine substitution).

Substitution reaction is where one atom or group of atoms is replaced by another atom or a group of atoms.

Examples of Alkanes:

1 carbon atom – Methane (CH₄)

2 carbon atoms – Ethane (C₂H₆)

3 carbon atoms – Propane (C₃H₈)

4 carbon atoms – Butane (C₄H₁₀)

5 carbon atoms – Pentane (C₅H₁₂)

Alkane	Structural Formula	Displayed Formula
Methane	CH₄
Ethane	CH₃CH₃
Propane	CH₃CH₂CH₃
Butane	CH₃CH₂CH₂CH₃

Substitution of Alkanes with Chlorine

This is a photochemical reaction where ultraviolet light provides the activation energy (Ea).

$\text{CH}_4 + \text{Cl}_2 \xrightarrow{\text{UV Light}} \text{CH}_3\text{Cl} + \text{HCl}$

$\text{Methane + Chlorine} \rightarrow \text{Chloromethane + Hydrochloric acid}$

Alkenes

Bonding in alkenes includes a double carbon-carbon covalent bond, and alkenes are unsaturated hydrocarbons.

Examples:

Ethene Propene But-1-ene:

Manufacturing Alkenes

A process of catalytic cracking is used to convert long-chain molecules into short-chain and more useful hydrocarbons.

Short-chain alkenes are formed by cracking long-chained alkenes.

Cracking is done by heating the hydrocarbon molecules to 600-700°C to vaporize them.

The vapors pass over a hot powdered catalyst like aluminum/silica, which breaks their covalent bonds, causing thermal decomposition.

This is done because short-chained alkenes are more useful hydrocarbons.

In an addition reaction, only one product is formed.

Properties of Alkenes:

Alkene + Bromine / Aqueous Bromine:

Ethene + Bromine → 1,2-Dibromoethane

C₂H₄ (g) + Br₂ (aq) → C₂H₄Br₂ (aq)

Alkene + Hydrogen in Presence of Nickel Catalyst:

C₂H₄ (g) + H₂ (g) → C₂H₆ (g)

Alkene + Steam in Presence of Acid Catalyst:

C₂H₄ (g) + H₂O (g) → C₂H₅OH (g)

Test to Distinguish Between Saturated/Unsaturated Hydrocarbons

Bromine Water (brown water) + Ethene (unsaturated) → 1,2-Dibromoethane (colorless)

If the color doesn’t change, it’s Alkane.

Alkene	Structural Formula	Displayed Formula
Ethene	CH₂=CH₂
Propene	CH₂=CHCH₃
But-1-ene	CH₂=CHCH₂CH₃
But-2-ene	CH₃CH=CHCH₃

Chapter 12: Experimental techniques & Chemical analysis

Experimental Design

Stopwatches – Measure time

Thermometers – Measure temperature

Balances – Measure Mass

Burettes – Measure volume of liquid or a gas

Volumetric Pipettes – To transfer a specific amount of liquid

Measuring Cylinder – Measure liquid volume (less accurate than burette)

Gas Syringe – Insert or withdraw a specific volume of gas

Advantages and Disadvantages of Methods & Apparatus

Apparatus Choice	Advantage	Disadvantage
Temperature probe versus thermometer	– More precise readings – Can easily take multiple repeat readings – Can use automatic sampling over longer periods of time	– Probe can be corroded by some reagents – Probes more expensive to replace
Volumetric pipette versus a measuring cylinder	– Pipette measures very accurately	– Harder to use and only measures one fixed volume
Gas syringe versus inverted cylinder for collecting gases	– Gas syringes are easy to set up and keep the gas dry	– The pistons can stick – Limited volume can be collected – Delicate and expensive
Microscale versus normal scale quantities	– Less wasteful of reagents – Saves energy – Safer	– Hard to see what’s happening – Lose a lot of material in separating and purifying products

Solvent:

A substance that dissolves a solute.

Solute:

A substance that is dissolved in a solvent.

Solution:

A mixture of one or more solutes dissolved in a solvent.

Saturated Solution:

A solution containing the maximum concentration of a solute dissolved in the solvent at a specified temperature.

Residue:

A substance that remains after any evaporation, distillation, filtration, or any similar process.

Filtrate:

A liquid/solution that has passed through a filter.

Acid-Base Titrations

Titration is a method of analyzing the concentrations of acids and bases. Titrations are also used to prepare salts.

The typical materials needed are:

50 cm³ burette

25 cm³ volumetric pipette

A suitable indicator

The indicator is required to show the end-point of titration.

Chromatography

A technique to separate substances that have different solubilities in a solvent.

Pencil line drawn on paper; sample placed on the line.

Paper put into solvent container above the line.

Solvent travels up the paper.

Different substances travel to different lengths.

Locating agents are substances that react with the sample to produce a colored product. The agent is used on chromatograms to locate things like amino acids/sugars.

Chromatograms

Pure:

All substance in the same place.

Mixture:

Separate to show all differences.

Impure:

Shows more than one spot.

$\text{R}_f = \frac{\text{distance travelled by substance}}{\text{distance travelled by solvent}}$

Retention Factor (R $_f$ ):

The R $_f$ value of a compound is always the same.

Separation & Purification

Suitable Solvent:

For a mixture of solids, can dissolve some.

Filtration:

For undissolved solids from solid & liquid.

Crystallisation:

Separate dissolved solid by evaporation.

Simple Distillation:

Separate liquid & soluble solid.

Fractional Distillation:

Separate two or more liquids.

Assessing Purity

Pure substances boil/melt at specific temperatures.

Mixtures have a range of temperatures.

Melting & boiling points can be used to assess purity.

The closer measured M.P (Melting Point) & B.P (Boiling Point) is to the actual point, the purer the substance is.

Identification of Ions and Gases

Tests to identify Anions:

Anion	Test	Result
Carbonate (CO₃²⁻)	Add dilute acid and test the gas released	Effervescence, gas produced is CO₂, which turns limewater milky
Chloride (Cl⁻)	Acidify with dilute nitric acid and add aqueous silver nitrate	White precipitate formed
Bromide (Br⁻)	Acidify with dilute nitric acid and add aqueous silver nitrate	Cream precipitate formed
Iodide (I⁻)	Acidify with dilute nitric acid and add aqueous silver nitrate	Yellow precipitate formed
Nitrate (NO₃⁻)	Add aqueous NaOH and aluminium foil, warm gently and test the gas released	Gas given off is ammonia, has a pungent smell, and turns moist red litmus paper blue
Sulfate (SO₄²⁻)	Acidify with dilute nitric acid and add aqueous barium nitrate	White precipitate formed
Sulphite (SO₃²⁻)	Add dilute acid, warm gently, and test the gas released	Gas decolorizes purple acidified aqueous potassium manganate(VII) solution

Using NaOH & NH₃ to identify the cations:

Metal Cation	Effect of Adding NaOH	Effect of Adding Ammonia Solution
Aluminium (Al³⁺)	White precipitate, dissolves in excess NaOH to form a colorless solution	White precipitate, insoluble in excess ammonia, white precipitate remains
Ammonium (NH₄⁺)	Ammonia produced if warmed	–
Calcium (Ca²⁺)	White precipitate, insoluble, so remains in excess NaOH	Very faintly visible white precipitate
Chromium (III) (Cr³⁺)	Green precipitate which forms a green solution in excess	Grey-green precipitate, insoluble in excess
Copper (II) (Cu²⁺)	Light blue precipitate, insoluble in excess	Light blue precipitate, soluble in excess to form dark blue color
Iron (II) (Fe²⁺)	Green precipitate, insoluble in excess	Green precipitate, insoluble in excess
Iron (III) (Fe³⁺)	Red-brown precipitate, insoluble in excess	Red-brown precipitate, insoluble in excess
Zinc (Zn²⁺)	White precipitate, dissolves in excess to form colorless solution	White precipitate, dissolves in excess to form colorless solution

Tests to identify gases:

Gas	Appearance of Gas	Test	Result
Ammonia (NH₃)	Colorless, pungent smell	Damp red litmus paper	Turns blue
Carbon dioxide (CO₂)	Colorless and odorless	Bubble through limewater	Limewater turns milky/cloudy
Chlorine (Cl₂)	Pale green, choking smell	Damp blue litmus paper	Turns red and quickly is bleached
Hydrogen (H₂)	Colorless and odorless	Hold a lighted splint in mouth of test tube	Burns with a ‘squeaky pop’ sound
Oxygen (O₂)	Colorless and odorless	Hold a glowing splint	Splint relights
Sulfur dioxide (SO₂)	Colorless, pungent choking smell	Add to acidified aqueous potassium manganate(VII)	Turns from purple to colorless

Flame test for cations:

Cation	Colour of Flame
Li⁺	Red
Na⁺	Yellow
K⁺	Lilac
Ca²⁺	Orange-red
Cu²⁺	Blue-green
Ba²⁺	Apple green

Chapter 7: Acids, Bases, and Salts

The Characteristics of Acids & Bases

The characteristic properties of acids in reactions with:

Metal:

Carbonates:

Bases:

Acids & base effect on:

Bases are oxides or hydroxides of metals, and alkalis are soluble bases.

Base reaction with:

Ammonium salts:

Neutralisation Reactions

Color & pH with Universal Indicator Paper

Preparation of Salts

An Alkali by titration:

acid+alkali→salt+water\text{acid} + \text{alkali} \rightarrow \text{salt} + \text{water}

Excess metal:

Insoluble bases:

Insoluble carbonate:

Preparation of insoluble salt with the precipitation method:

Solubility rules for salts:

Oxides

Acidic Oxides:

Basic Oxides:

Amphoteric Oxides:

Chapter 8: The periodic table

Arrangement of elements

Group VII Properties

The appearance of halogens at r.t.p:

Transition elements

Noble Gases

Group I Properties

Chapter 9: Metals

Properties of Metals:

Properties of Non-Metals:

Metals reaction to:

Dilute acids:

Cold water & steam:

Oxygen:

Uses of Metals

Extraction of Metals

The main ore of aluminium is bauxite and aluminium is extracted by electrolysis.

Alloys and Their Properties

Corrosion of Metals

Reactivity Series

The order of reactivity is:

Chapter 10: Chemistry of the environment

Water

Testing for the Presence of Water:

Anhydrous cobalt(II) chloride:

Anhydrous copper(II) sulfate:

Testing the Purity of Water:

Water from natural sources includes substances like:

The treatment of domestic water supply:

Fertilisers

Common fertilizer compounds:

Air Quality and Climate

The sources of air pollutants:

The air pollutants have adverse effects such as:

The Greenhouse Effect:

Strategies that reduce the effect of these environmental issues are:

For climate change:

For acid rain:

Photosynthesis:

Word Equation:

Carbon dioxide+Water→Glucose+Oxygen\text{Carbon dioxide} + \text{Water} \rightarrow \text{Glucose} + \text{Oxygen}

Chemical Equation:

Chapter 11: Organic Chemistry

Formulae, Functional Groups, and Terminology

Drawing & interpreting displayed formula:

Homologous series:

The general characteristics:

Saturated:

Unsaturated:

Structural formula:

Structural isomers:

Naming Organic Compounds

Carboxylic Acids

Reaction of Ethanoic Acid with:

Carboxylic acid reaction with alcohol using acid catalyst to form an ester:

Polymers

$\text{acid} + \text{alkali} \rightarrow \text{salt} + \text{water}$

$\text{Carbon dioxide} + \text{Water} \rightarrow \text{Glucose} + \text{Oxygen}$