Chemistry Unit 11: Air and Water

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Unit 11 Air and Water

1. Describe a chemical test for water.

There are two ways in which you can go about. The first is using anhydrous copper sulphate; in the presence of water, the anhydrous copper sulphate turns from white to blue. Alternatively, you could use anhydrous cobalt chloride; in the presence of water, the anhydrous cobalt chloride turns from blue to pink. 

2. Describe and explain, in outline, the purification of the water supply by filtration and chlorination.

Filtration is used to separate insoluble substances from the water. Afterwards, the water is chlorinated to get rid of bacteria that couldn't be removed through filtration.

3. State some of the uses of water in industry and in the home.

Home: cooking, cleaning, washing

Industry: cooling, planting in agriculture

4. Describe the separation of oxygen and nitrogen from liquid air by fractional distillation.

Like crude oil, liquid air can be separated in a fractionating column. Liquid air is pumped into the bottom of the column and it is heated. Nitrogen gas evaporates from the air and is collected at the top of the column, while liquid oxygen is collected from the bottom. 

5. Describe the composition of clean air as being a mixture of 78% nitrogen, 21% oxygen and small quantities of noble gases, water vapour and carbon dioxide. 

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6. State the common air pollutants as carbon monoxide, sulphur dioxide and oxides of nitrogen, and describe their sources.

Carbon monoxide, sulphur dioxide and nitrogen oxides are what make air dirty. Carbon monoxide comes from the incomplete burning of fuels, so any machinery that burns a fuel could release carbon monoxide. Additionally, burning cigarettes also produce carbon monoxide. Sulphur dioxide and nitrogen oxides both come from burning fossil fuels.

7. Explain the presence of oxides of nitrogen in car exhausts and their catalytic removal.

Oxides of nitrogen are formed in car engines due to heat and pressure. They are poisonous and harmful to humans, so catalytic converters are put into engines. Nitrogen oxide is reduced, meaning that it loses oxygen, so it breaks down into separate nitrogen and oxygen gases. In addition, carbon monoxide is oxidised, meaning it gains oxygen, so it becomes carbon dioxide. 

8. Explain why the proportion of carbon dioxide in the atmosphere is increasing, and why this is important.

Carbon dioxide levels are rising due to excessive human activity, such as factory production and deforestation. This is important because CO2 is one of the gases that contribute to the greenhouse effect, therefore resulting in global warming. Global warming is the general increase of temperature, and this has a great effect on our environment. 

9. State the adverse effect of common air pollutants on buildings and on health.

Air pollutants affect the materials used to build buildings; they cause the materials to become discoloured, brittle, and prone to deterioration. Pollutants also have negative effects on health, including eye, throat and nose irritation, breathing problems, and cancer. 

10. Describe the formation of carbon dioxide: 

• as a product of complete combustion of carbon-containing substances, 

• as a product of respiration, 

• as a product of the reaction between an acid and a carbonate, 

• as a product of thermal decomposition.

When anything containing carbon burns, the carbon reacts with oxygen, gets oxidised, and becomes carbon dioxide. Carbon dioxide is a product of respiration, from the equation glucose + oxygen --> carbon dioxide and water. Similarly, when an acid and carbonate react together, a product of the reaction is carbon dioxide (the other is water). Carbon dioxide is also produced from thermal decomposition. Take the decomposition of calcium carbonate for instance: when calcium carbonate is heated, it breaks down into calcium oxide and carbon dioxide. 

11. Describe the essential conditions for the manufacture of ammonia by the Haber process including the sources of the hydrogen and nitrogen, i.e. hydrocarbons or steam and air.

The Haber process is a process that manufacture ammonia. Nitrogen from the atmosphere and hydrogen, formed through the reacted of methane and steam, are reacted together to form ammonia; this is a reversible reaction.

nitrogen + hydrogen <--> ammonia

N + 3H <--> NH3

The amount of ammonia produced at a time is quite low - about 15%. To produce a higher yield, the following conditions are needed:

• low temperature

• high pressure

• catalyst of iron to speed up the reaction 

This is an important process because ammonia is used in the manufacture of fertilisers used in agriculture. Fertilisers speed up the growth of plants and gives them additional nutrients, so that they grow quickly and healthily. 

12. Describe the rusting of iron in terms of a reaction involving air and water, and simple methods of rust prevention, including paint and other coatings to exclude oxygen.

Iron rusts when it comes into contact with both air and water. Water reacts with the iron over time and the oxygen in the air oxidises the iron. Combined, this forms a layer of rust. 

There are several ways to prevent rust from forming. One would be to simply use a coat of paint; this acts as a barrier to prevent contact with water and air. Another way would be to galvanise the iron, which is to coat the iron in zinc. Similar to paint, the zinc acts as a barrier. Lastly, sacrificial protection can be used. A block of zinc is placed next to the iron, so that any water and air present reacts with the zinc, rather than the iron. 

13. Describe the need for nitrogen-, phosphorus- and potassium-containing fertilisers.

Plants need nitrogen, phosphorus and potassium to grow healthily. Nitrogen helps with the production of proteins, which is a plant's food. Phosphorus helps balance out the pH of the soil, and potassium helps with the production of chloroplasts, which help the plant photosynthesise. 

14. Describe the displacement of ammonia from its salts by warming with an alkali.

When an ammonium salt is warmed with an alkali, the ammonia in the salt is displaced by ammonia gas.

Chemistry Unit 7.1: Speed of Reaction

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Unit 7.1 Speed of Reaction

1. Describe the effect of concentration, particle size, catalysis and temperature on the speeds of reactions.

Here are some factors that speed up the rate of a reaction:

• surface area: the larger the surface area, the faster the rate; this is because this increases the chance of particles colliding with each other
• concentration: a higher concentration means there are more particles in each cm^3 of space, so there are more collisions
• catalysis: a catalyst speeds up the rate of a reaction; it lowers the amount of energy needed for the reaction to take place, so it occurs more easily
• temperature: a higher temperature provides energy for articles to move more quickly, so they collide more often

2. Describe a practical method for investigating the speed of a reaction involving gas evolution.

An example to investigate this is using magnesium and hydrochloric acid to test the rate of reaction. The equation goes as follows:

magnesium + hydrochloric acid --> magnesium chloride + hydrogen

Equipment:

• conical flask

• gas syringe

• stopwatch

• hydrochloric acid

• magnesium

Using this apparatus, we can design an experiment that, for example, measures how long it takes for 20cm^3 of gas, using the volume of magnesium as our independent variable (the one you change). 

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The diagram above uses magnesium ribbons instead of powder, but that's completely fine.

3. Devise a suitable method for investigating the effect of a given variable on the speed of a reaction.

For the experiment above, time was the variable that we changed. The experiment could also be designed around other variables:

• the concentration of acid

• temperature conditions

• surface area of magnesium (chips? ribbon? powder?)

• the reaction with a catalyst and without

4. Interpret data obtained from experiments concerned with speed of reaction.

Once you have carried out the experiment and obtained data, you'll be able to come to a conclusion on what you were investigating.

5. Describe the application of the above factors to the danger of explosive combustion with fine powders (e.g. flour mills) and gases (e.g. mines).

So you know that variables affecting the rate of a reaction are temperature, concentration, and surface area. These are applicable to real life situations.

In a flour mill, flour is produced. Flour particles are very small, so they have a large surface area. If there is a lot of flour in the air, a small spark can cause an explosion between flour and oxygen.

In a coal mine, the air is filled with flammable gases. If the gases reach a certain concentration they can form an explosion with air. 

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6. Describe and explain the effects of temperature and concentration in terms of collisions between reacting particles (concept of activation energy will not be examined).

As mentioned above, a higher temperature will give particles more energy to move and collide, and a higher concentration means there are more particles in the space that will collide. 

As a reaction occurs and the reactants get used up, the concentration of the substance decreases; this is why reactions slow down as they happen over a span of time.

7. Define catalyst as an agent which increases rate but which remains unchanged.

A catalyst is a substance thats speeds up a reaction, but the catalyst itself will not be used up.

Chemistry Unit 9: The Periodic Table

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As reference, here is a Periodic Table; you can refer to this as you go along the syllabus statements. If you didn't know already, you will be provided with a Periodic Table when you take you exams!

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1. Describe the way the Periodic Table classifies elements in order of proton number.

Elements are ordered in numerical order according to the number of protons it contains. Hydrogen has one proton so it is first, followed by helium, lithium, etc.

2. Use the Periodic Table to predict properties of elements by means of groups and periods.

On the Periodic Table there are groups. The important ones that the syllabus covers are alkali metals (group 1), transition metals (group 3), halogens (group 7) and noble gases (group 0/8).




Unit 9.1 Periodic Trends

1. Describe the change from metallic to non-metallic character across a period.

As you run across the Periodic Table horizontally, the character of the element goes from metallic to non-metallic.

2. Describe the relationship between Group number, number of outer-shell (valency) electrons and metallic /non-metallic character.

The valency of the outer shell depends on what group the element is in. Here's a little advice: once you get your Periodic Table in the exam, start from Group 1 and label +1, +2, +3, +4, -3, -2, -1, full, ending on Group 8. This determines the valency of the element, helping you to balance equations.




Unit 9.2 Group Properties

1. Describe lithium, sodium and potassium in Group I as a collection of relatively soft metals showing a trend in melting point and reaction with water.

Metal	Hardness	Reactivity	Melting point	Reaction with water
Lithium	Soft	Reactive	180.5°c	Fizzes on water, releases hydrogen
Sodium	Softer	More reactive	97.7°c	Fizzes on water, burns with a yellow flame, release hydrogen
Potassium	Softest	Most reactive	63.4°c	Fizzes on water, burns with a lilac flame, release hydrogen

From my experience, exam boards like to throw in a question worth 6-9 marks about lithium, sodium and potassium comparisons, so it's worth remembering these points.

2. Predict the properties of other elements in Group I, given data where appropriate.

Properties of alkali metals
• relatively soft metals
• low boiling and melting points --> decrease as you go down the group
• reacts with water to form metal hydroxide + hydrogen
• reacts with oxygen to form metal oxide
• reactivity increases as you go down the group
• densities increase as you go down the group

3. Describe the trends in properties of chlorine, bromine and iodine in Group VII including colour, physical state and reactions with other halide ions.

Halogen	State at room temperature	Colour	Reaction with halide ions
Chlorine	Gas	Yellow-green	Oxidises bromine and below
Bromine	Liquid	Red-brown	Oxidises iodine and below
Iodine	Solid	Purple	Oxidises astatate and below

4. Predict the properties of other elements in Group VII, given data where appropriate.

Properties of halogens

• melting and boiling points increase as you go down the group

• reactivity decreases as you go down the group

• kills bacteria

• react with metals to form metal halides

Unit 9.3 Transition Elements

1. Describe the transition elements as a collection of metals having high densities, high melting points and forming coloured compounds, and which, as elements and compounds, often act as catalysts.

Properties of transition metals

• high densities

• high melting points

• forms coloured compounds

• can act as catalysts

Unit 9.4 Noble Gases

1. Describe the noble gases as being unreactive.

Noble gases are unreactive because their outer shells are full. This means that they are stable without having to gain or lose electrons. They are also described as inert. 

2. Describe the uses of the noble gases in providing an inert atmosphere, i.e. argon in lamps, helium for filling balloons.

Under normal conditions, noble gases are odourless, colourless and nonflammable. 

• argon is used in lightbulbs

• helium is used to inflate tires and balloons

• neon is used in neon lights because it glows when electricity is passed through it

• xenon is used in photocopying 

Chemistry Unit 5: Electricity and Chemistry

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Unit 5. Electricity and Chemistry

1. State that electrolysis is the chemical effect of electricity on ionic compounds, causing them to break up into simpler substances, usually elements.

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From the diagram, you can see that the electrolyte (solution) is made up of negative and positive ions; these are called anions and cations, respectively. Through electrolysis, we can use electricity to separate and obtain the anions and cations. The electrolyte is a compound that can conduct electricity for the process to happen. 

2. Use the terms electrode, electrolyte, anode and cathode.

Electrode: the conducting material which passes an electrical current (carbon is commonly used)

Electrolyte: the solution being broken down

Anode: the positive electrode, attracts anions

Cathode: the negative electrode, attracts cations

3. Describe electrolysis in terms of the ions present and the reactions at the electrodes.

The ionic compound will only be able to be broken down if they are in liquid form, as in a liquid the ions can move freely. To obtain a liquid, water can be added to the compound to dissolve it. This way, the anions are free to be attracted to the anode and the cations to the cathode. 

4. Describe the electrode products, using inert electrodes, in the electrolysis of: 

• molten lead(II) bromide, 

• aqueous copper chloride, 

• dilute sulphuric acid.

Molten lead (II) bromide --> lead 2+ ion (to cathode) + bromine 1- ion (to anode)

Aqueous copper chloride --> copper 1+ ion (to cathode) + chlorine 2-  ion (to anode)

Dilute sulphuric acid --> hydrogen (to cathode) + sulphur (to anode)

5. State and use the general principle that metals or hydrogen are formed at the negative electrode (cathode), and that non-metals (other than hydrogen) are formed at the positive electrode (anode).

The term 'opposites attract' applies here: cathodes (-) attract cations (+), while anodes (+) attract anions (-). To help you with this, think about this trick:

Cations = pussytive --> positive

Even for me, this tip is weird, but in it's own weird way it works!

Metals and hydrogen are positive, so they are the cations in the equation. If the metal is more reactive than hydrogen on the reactivity scale, it will get replaced by hydrogen and hydrogen will be formed at the cathode. The less reactive substance will be formed at the cathode.

For example, in aqueous sodium chloride there is sodium (Na+) and hydrogen (H+). Hydrogen takes sodium's place in the reaction --> 2H+ + 2e- = H2 

This is called a reduction as the negatively charged ions in the circuit are attracted to the positive cations, therefore reducing the cation to neutral. 

Non-metals are negative, so they are the equation's anions. They attract the positive ions from the circuit and also become neutral; this is called oxidisation. 

6. Relate the products of electrolysis to the electrolyte and electrodes used, exemplified by the specific examples in the Core together with aqueous copper(II) sulphate using carbon electrodes and using copper electrodes (as used in the refining of copper).

Electrolysis can also be used to purify metals. When doing so, instead of using carbon rods as electrodes we use the metal we want to purify. 

Let's say we want to purify copper. We take a rod of impure copper as our anode, and pure copper as our cathode. The solution used is aqueous copper sulphate. When the circuit is switched on, copper ions are attracted to the anode as they are cations (all metals are cations!). The impurities of the anode become sludge and the anode decreases in size, while the cathode gains copper ions and increases in size. 

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7. Describe the electroplating of metals, using laboratory apparatus.

Electroplating a metal is to coat a metal with another metal. We do this because maybe the metal being coated is highly reactive and we want to protect them. Metals commonly used to for electroplating are tin, silver, copper and chromium. 

To electroplate, the cathode is the object to be plated and the anode is the metal you want to plate with. The electrolyte is a salt solution of the same metal as the anode. When the circuit is switched on, the anode ions are attracted to the cathode, wearing away the anode. The concentration of the solution stays the same because no ions are being taken from it. 

Say you want to coat a brass key in copper. With alligator clips that are connected to wires, you attach a rod or wire of copper as your anode and the key as your cathode. The wires will be connected to the power supply accordingly. In a beaker of copper sulphate, the anode and cathode are placed into the electrolyte and the circuit is switched on. With this lab equipment, this is a simple process of electroplating. 

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8. Predict the products of the electrolysis of a specified binary compound in the molten state, or in aqueous solution.

Let's use sodium chloride.

At the anode: 2Cl --> Cl2 + 2e-

At the cathode: Na+ + e- --> Na

9. Describe, in outline, the chemistry of the manufacture of 

• aluminium from pure aluminium oxide in molten cryolite, 

• chlorine, hydrogen and sodium hydroxide from concentrated aqueous sodium chloride.

Aluminium: aluminium ore, bauxite, is purified to obtain a white powder; this white powder is aluminium oxide. Aluminium can be extracted from aluminium oxide through electrolysis.

Chlorine, hydrogen and sodium: sodium chloride (brine) is used as a electrolyte for electrolysis, and chlorine, hydrogen and sodium can be obtained from there. 

Chemistry Unit 2: Experimental Techniques

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Unit 2.1 Methods of Separation and Purification

1. Describe paper chromatography.

Chromatography is used to separate two or more dissolved solids in a solution. Paper chromatography uses water to separate the solids. A line is draw on a piece of paper and the solution is dotted on that line. After labelling the dot, the piece of paper is placed into water (or another solvent), the water level meeting a place a little below the pencil line. As the water travels up the paper (this is called capillary action), it separates the solids in the solution. 

We can use chromatography to evaluate the contamination of food or water, which is part of health and safety procedures. 

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2. Interpret simple chromatograms.

The further a substance moves up the paper, the more soluble it is. 

3. Describe methods of separation and purification: filtration, crystallisation, distillation, fractional distillation.

Filtration: when you have a solid and solution, filtering the mixture can help you separate each from the other e.g. sand and water

Crystallisation: crystallisation allows you to separate a solute from a solvent. The solution is left to dry or heated up, evaporating the solvent and leaving the solute in the form of crystals e.g. salt and water

Distillation: distillation helps you obtain a pure liquid from a mixture of liquids. The mixture is heated up in a flask, and the liquid with the lower boiling point will evaporate first. The vapour is cooled in the Liebig condenser and collected in the form of a liquid e.g. ethanol and water

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Fractional distillation: this process is similar to distillation, but it is used to separate more parts in a mixture. The mixture is heated up in a flask, and attached above the flask is a column with beads. This helps to separate the different parts as they evaporate. The vapours are then cooled through the Liebig condenser and collected as liquids e.g. crude oil

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4. Understand the importance of purity in substances in everyday life, e.g. foodstuffs and drugs.

As mentioned above, checking the purity of things is important as contamination can result in health and safety issues, especially food or drugs. They can cause you to become very sick.

5. Identify substances and assess their purity from melting point and boiling point information.

Take water for example. We all know that the boiling point of water is 100°c. If the boiling point has changed, this indicates that the water is not pure. This is the same with the melting point.

6. Suggest suitable purification techniques, given information about the substances involved.

With the above methods of separation and purification, I have given some examples of what each method can be used for. 

Chemistry Unit 12: Sulphur

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Unit 12: Sulphur

1. Describe the manufacture of sulphuric acid by the Contact process, including essential conditions.

The manufacture of sulphuric acid is called the Contact Process and this is the process: firstly, you need to burn sulphur in oxygen to create sulphur dioxide. The sulphur dioxide is then passed over a vanadium pentoxide catalyst to form sulphur trioxide. Essential conditions for maximum yield of sulphur trioxide is 450°c and a moderate amount of atmospheric pressure. The sulphuric trioxide is then dissolved in water to produce sulphuric acid. 

Here's how it goes with symbol equations:

1. sulphur + oxygen --> sulphur dioxide

S + O2 --> SO2

2. sulphur dioxide + oxygen --> sulphur trioxide (vanadium pentoxide catalyst, 450°c, atmospheric pressure)

SO2 + O2 --> SO3

3. sulphur trioxide + water --> sulphuric acid

SO3 + H2O --> H2SO4

2. Describe the properties of dilute sulphuric acid as a typical acid.

● eats away and dissolve objects

● readily absorbs moisture

● corrosive

● reacts with bases

● doesn't conduct electricity

Chemistry Unit 8: Acids, Bases and Salts

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Unit 8.1 The characteristic properties of acids and bases

Before I start going into the syllabus, let's define acids and bases clearly!

Acids: a substance that dissolves in water to produce hydrogen

Base: a substance that neutralises an acid

A lot of the time, you'll hear the word alkali. An alkali is a type of base that can dissolve in water. An alkali is a base, but a base is not an alkali.

1. Describe neutrality and relative acidity and alkalinity in terms of pH (whole numbers only) measured using full-range indicator and litmus.

When talking about pH, you've got the whole pH scale in mind, where you identify the pH according to a certain colour. You can use litmus paper or universal indicator (UI) to check a substance's pH. 

Neutral: pH 7, turns UI green, turns litmus paper purple

Acid: low pH (less than 7), turns UI red-yellow, turns litmus paper red

Alkali: high pH (greater than 7), turns UI blue-purple, turns litmus paper blue

Below is an image displaying the different colours universal indicator turns according to pH levels, and also what substances are of that pH level.

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2. Describe the characteristic reactions between acids and metals, bases (including alkalis) and carbonates.

Here is a handy-dandy four-step system to this:

● acid + alkali --> salt + water

● acid + metal --> salt + hydrogen

● acid + carbonate --> salt + water + carbon dioxide

● acid + oxides (bases) --> salt + water

3. Describe and explain the importance of controlling acidity in the environment (air, water and soil)

Soil: Soil is used to grow crops, so it is important for it to be neutral. If it happens to be too alkaline or acidic, the crops tend to grow poorly. Acidity is usually the problem for soil, so a base can help neutralise it. Bases include limestone, slaked lime, or quick lime (more on that later!).

Water: Factory waste is often acidic, and it can leak into water. To prevent this from happening, it needs to be neutralised. Again, slaked lime is used for this. 

Air: Burning fossil fuels releases gases into the air, such as nitrogen oxides and sulphur dioxide. They react with water and air, leading to acid rain. As you can tell by the name, it's not that great; it causes buildings to erode and will negatively affect soil and water. 

Unit 8.2 Types of oxides

1. Classify oxides as either acidic or basic, related to metallic and non-metallic character of the other element.

First, let's define oxide: it is a compound made up of oxygen and another element. Next, acidic oxide: a compound of oxygen and a non-metal, most commonly a gas. Lastly, basic oxide: a compound of oxygen and a metal.

Here are some examples of each type of oxide:
acidic oxide: sulphuric acid (H2SO4), nitric acid (HNO3)
basic oxide: sodium hydroxide (NaOH), copper oxide (CuO)

So now that you know what the different oxides are, it's easier to classify them into groups. 

Acidic oxides:
● react with water to give and acid
i.e carbon dioxide + water --> carbonic acid
     CO2 + H2O --> H2CO3
● react with bases to form a salt
i.e. sulphuric acid + copper oxide --> copper sulphate + water
     H2SO4 + CuO --> CuSO4 + H2O

Basic oxides:
● react with acids to form a salt and water
i.e. copper oxide + hydrochloric acid --> copper chloride + water

     CuO + 2HCl --> CuCl2 + H2O


2. Further classify some oxides as neutral, given relevant information.

Some oxides are neither basic or acidic, so they are neutral oxides. They don't react with acids or bases either. Examples of neutral oxides include nitrous oxide (N2O) and carbon monoxide (CO). 




Unit 8.3 Preparation of salts

1. Describe the preparation, separation and purification of salts using techniques selected from section C2.1 and the reactions specified in C8.1.

Preparing the salt:
1. Add an excess of your carbonate to your acid
2. Test the pH with U.I. paper to check if the solution is neutral
3. Filter the solution to get rid of the excess carbonate
4. Heat the liquid until most of it has evaporated and you are left with salt crystals

Separating the salt:
Let's say we have a mixture of salt and pebbles
1. Add water to your mixture so the salt dissolves
2. Filter the mixture
3. Evaporate the liquid

Purifying the salt:
If it's an insoluble salt: filtration
If it's a soluble salt: distillation

2. Suggest a method of making a given salt from suitable starting materials, given appropriate information. 

Remember that handy-dandy four-step system from above? Here, it comes to use.

Acid + alkali --> salt + water
Hydrochloric acid + sodium hydroxide --> sodium chloride + water
HCl + NaOH --> NaCl + H2O
1. Add 30cm2 of sodium hydroxide to a flask
2. Add two drops of the indicator phenolphthalein 
3. Add the hydrochloric acid to the flask
4. The solution will turn colourless when neutralised
5. Heat the solution to evaporate the liquid and obtain the salt

Acid + metal --> salt + hydrogen
Sulphuric acid + zinc --> zinc sulphate + hydrogen
H2SO4 + Zn --> ZnSO4 + H2
1. Add zinc to a flask of the acid
2. When the zinc dissolves, hydrogen bubbles will appear
3. Filter the solution to remove the excess zinc
4. Heat the solution to evaporate the liquid and obtain the salt

Acid + base --> salt + water
Sulphuric acid + iron(II) oxide --> iron sulphate + water
H2SO4 + FeO --> FeSO4 --> H2O
1. Add and excess of iron oxide to the sulphuric acid
2. Filter the solution to remove the excess iron oxide
3. Heat the solution to evaporate the liquid and obtain the salt




Unit 8.4 Identification of ions and gases

1. Use the following tests to identify aqueous cations, anions, and gases.

Aqueous cations:
● ammonium
● copper(II)
● iron(II)
● iron(III)
● zinc 

Cation	Test	What happens if cation is present?
Ammonium	Add dilute sodium hydroxide, warm up	Ammonia gas is released, damp red litmus paper turns blue
Copper(II)	Add dilute sodium hydroxide or ammonia solution	A blue precipitate forms
Iron(II)	Add dilute sodium hydroxide or ammonia solution	A pale green precipitate forms
Iron(III)	Add dilute sodium hydroxide or ammonia solution	A red-brown precipitate forms
Zinc	Add dilute sodium hydroxide or ammonia solution	A white precipitate forms

Anions:

● carbonate

● chloride

● nitrate

● sulphate

Anion	Test	What happens if anion is present?
Carbonate	Add dilute hydrochloric acid	Bubbles that give off gas turn limewater a milky-white
Chloride	Add the same volume of nitric acid as chloride, add aqueous silver nitrate	A white precipitate forms
Nitrate	Add sodium hydroxide, then aluminium	Ammonia gas will be given off
Zinc	Add dilute sodium hydroxide or ammonia solution	A white precipitate forms

Gases:

● ammonia

● carbon dioxide

● chlorine

● hydrogen

● oxygen

Anion	Test	What happens if gas is present?
Ammonia	Damp red litmus paper	Paper turns blue
Carbon dioxide	Limewater	Limewater turns milky-white
Chlorine	Damp blue litmus paper	Paper turns white
Hydrogen	Lighted splint	A loud POP
Oxygen	Glowing splint	Splint relights