Spec point which is NOT in Triple Science
1:51 know that covalent compounds do not usually conduct electricity
Electric current is a flow of charged particles that can move.
Covalent compounds do not conduct electricity.
2:01 understand how the similarities in the reactions of lithium, sodium and potassium with water provide evidence for their recognition as a family of elements
Group 1 metals such as potassium, sodium and lithium, react with water to produce a metal hydroxide and hydrogen. For example:
lithium + water → lithium hydroxide + hydrogen
2Li (s) + 2H₂O (l) → 2LiOH (aq) + H₂ (g)
The observations for the reaction of water with either potassium or sodium or lithium have the following similarities:
- fizzing (hydrogen is produced)
- metal floats and moves around on the water
- metal disappears
In each case a metal hydroxide solution is produced.
These similarities in the reactions provide evidence that the 3 metals are in the same group of the Periodic Table (i.e. have the same number of electrons in their outer shell).
2:02 understand how the differences between the reactions of lithium, sodium and potassium with air and water provide evidence for the trend in reactivity in Group 1
Lithium is the first element in group 1 of the Periodic Table. The observations for the reaction of lithium and water are:
- fizzing (hydrogen gas is released)
- lithium floats and moves around on the water
- lithium disappears
Sodium is the second alkali metal in the group. The reaction of sodium and water is more vigorous than lithium’s:
- fizzing (hydrogen gas is released)
- sodium floats and moves around on the water
- sodium melts into a silver-coloured ball
- sodium disappears
Potassium is the third alkali metal in the group. The reaction of potassium and water is more vigorous than sodium’s:
- fizzing (hydrogen gas is released)
- potassium floats and moves around on the water
- catches fire with a LILAC flame
- potassium disappears
When the group 1 metals react with air they oxidise, showing a similar trend in reactivity as we go down the group of the Periodic Table.
Therefore, as we go down group 1 (increasing atomic number), the elements become more reactive: Li<Na<K<Rb<Cs<Fr
2:03 use knowledge of trends in Group 1 to predict the properties of other alkali metals
From the data in the table, it is possible to deduce the properties of francium from the trends in the other group 1 metals.
For example, we can predict that francium will have a melting point around 20⁰C and a density of just over 2g/cm³.
We can also predict that francium will react violently with water, producing francium hydroxide and hydrogen.
| Alkali metal | Melting point (⁰C) | Density (g/cm³) | Reaction with water | Products |
|---|---|---|---|---|
| lithium (Li) | 181 | 0.53 | fizzing | lithium hydroxide + hydrogen |
| sodium (Na) | 98 | 0.97 | rapid fizzing | sodium hydroxide + hydrogen |
| potassium (K) | 63 | 0.86 | vigorous fizzing and lilac flame | potassium hydroxide + hydrogen |
| rubidium (Rb) | 39 | 1.53 | ? | rubidium hydroxide + hydrogen |
| caesium(Cs) | 29 | 1.88 | ? | caesium hydroxide + hydrogen |
| francium (Fr) | ? | ? | ? | ? |
2:05 know the colours, physical states (at room temperature) and trends in physical properties of chlorine, bromine and iodine
| Element | Colour | State at room temp |
|---|---|---|
| Chlorine (Cl2) | Green | Gas |
| Bromine (Br2) | Red-brown | Liquid |
| Iodine (l2) | Grey | Solid |
Chlorine is a toxic gas, so should be handled in a fume cupboard.
2:06 use knowledge of trends in Group 7 to predict the properties of other halogens
If you look at the trends in the physical properties of the halogens, Cl2, Br2, I2 you can make predictions about the properties of the other halogens.
| Element | Colour | State at room temp |
|---|---|---|
| Fluorine (F2) | Yellow | Gas |
| Astatine (At2) | Black | Solid |
2:07 understand how displacement reactions involving halogens and halides provide evidence for the trend in reactivity in Group 7
Group 7 elements are called the Halogens. As you go up group 7 (decreasing atomic number), the elements become more reactive. For example, fluorine is the most reactive and astatine is the least reactive.
A more reactive halogen will displace a less reactive halogen, e.g. chlorine will displace bromine:
By reacting a halogen solution with a potassium halide solution and making observations, the order of their reactivity can be deduced:
| Potassium chloride, KCl(aq) | Potassium bromide, KBr(aq) | Potassium iodide, KI(aq) | |
|---|---|---|---|
| Chlorine, Cl2(aq) | No change | Colourless to orange | Colourless to brown |
| Bromine, Br2(aq) | No change | No change | Colourless to brown |
| Iodine, I2(aq) | No change | No change | No change |
From the above results, chlorine displaces both bromine and iodine, and bromine displaces iodine. Therefore the order of reactivity is: chlorine is more reactive than bromine, which in turn is more reactive than iodine.
2:09 know the approximate percentages by volume of the four most abundant gases in dry air
Air is a mixture of different gases.
The abundance of gases in the air is as follows:
| Gas | % by volume |
|---|---|
| Nitrogen, N2 | 78.1 |
| Oxygen, O2 | 21.0 |
| Argon, Ar | 0.9 |
| Carbon dioxide, CO2 | 0.04 |
2:10 understand how to determine the percentage by volume of oxygen in air using experiments involving the reactions of metals (e.g. iron) and non-metals (e.g. phosphorus) with air
The following 3 experiments can be used to determine that oxygen (O2) makes up approximately 20% by volume of the composition of air.
Copper
The copper is in excess and uses up the oxygen to form copper oxide (CuO).
All the oxygen in the air is therefore used up, and so the volume of the air decreases by about 20% (the percentage of oxygen in air).
Iron
The iron reacts with the oxygen in the air (rusting).
As long as the iron and water are in excess, the total volume of air enclosed by the apparatus decreases by about a fifth (20%) over several days.
Phosphorus
The phosphorus is lit with a hot wire.
It reacts with the oxygen in the air and causes the water level in the bell jar to rise by about 20%.
2:11 describe the combustion of elements in oxygen, including magnesium, hydrogen and sulfur
Magnesium reacts with oxygen producing a bright white flame leaving behind a white ash of magnesium oxide.
magnesium + oxygen → magnesium oxide
2Mg (s) + O₂ (g) → 2MgO
MgO is a base, which can react with an acid to give a salt and water.
Hydrogen reacts with oxygen in an explosive reaction. This is the basis of the ‘squeak pop’ test for hydrogen in test tube. With larger quantities of hydrogen this explosion can be dangerous.
hydrogen + oxygen → water
2H₂ (g) + O₂ (g) → 2H₂O (l)
Sulfur reacts with oxygen producing a blue flame.
sulfur + oxygen → sulfur dioxide
S (s) + O₂ (g) → SO₂ (g)
When sulfur dioxide (SO₂) dissolves in water it forms an acidic solution of sulfurous acid:
SO₂ (g) + H₂O (l) → H₂SO₃ (aq)
2:12 describe the formation of carbon dioxide from the thermal decomposition of metal carbonates, including copper(II) carbonate
thermal decomposition is the process of breaking down by heating.
On heating metal carbonates thermal decompose into metal oxides and carbon dioxide.
Observation: green powder (CuCO3) changes to a black powder (CuO)
2:13 know that carbon dioxide is a greenhouse gas and that increasing amounts in the atmosphere may contribute to climate change
Carbon dioxide (CO2) is a greenhouse gas.
It absorbs infra-red radiation and therefore warms the atmosphere. This leads to global warming.
This may cause climate change.
2:14 Practical: determine the approximate percentage by volume of oxygen in air using a metal or a non-metal
The following 3 experiments can be used to determine that oxygen (O2) makes up approximately 20% by volume of air.
Copper
The copper is in excess and uses up the oxygen to form copper oxide (CuO).
All the oxygen in the air is therefore used up, and so the volume of the air decreases by about 20% (the percentage of oxygen in air).
Iron
The iron reacts with the oxygen in the air (rusting).
As long as the iron, oxygen and water are all in excess, the total volume of air enclosed by the apparatus decreases by about a fifth (20%) over several days.
Phosphorus
The phosphorus is lit with a hot wire.
It reacts with the oxygen in the air and causes the water level in the bell jar to rise by about 20%.
2:15 understand how metals can be arranged in a reactivity series based on their reactions with: water and dilute hydrochloric or sulfuric acid
Some metals are more reactive than others.
The order of reactivity can be determined by adding acid to different metals and observing the rate of reaction.
For example, when hydrochloric acid is added to iron (Fe) then bubbles of hydrogen are produced slowly. However, if the same acid is added to zinc (Zn) then bubbles will be produced more quickly. This tells us that zinc is more reactive than iron.
Instead of using acid, water can be used to test the relative reactivity of metals. However, many metals are too low in the reactivity series to react with water
2:16 understand how metals can be arranged in a reactivity series based on their displacement reactions between: metals and metal oxides, metals and aqueous solutions of metal salts
A metal will displace another metal from its oxide that is lower in the reactivity series. For example, a reaction with magnesium and copper (II) oxide will result in the magnesium displacing the copper from its oxide:
A metal will also displace another metal from its salt that is lower in the reactivity series. For example, the reaction between zinc and copper (II) sulfate solution will result in zinc displacing the copper from its salt:
The blue colour of the copper (II) sulfate solution fades as colourless zinc sulfate solution is formed.
2:17 know the order of reactivity of these metals: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver, gold
A more reactive metal will displace a less reactive metal.
In addition a more reactive metal will react more vigorously than a less reactive metal.
For example, potassium takes a shorter time to react than sodium:
2:18 know the conditions under which iron rusts
Iron rusts when oxygen and water are present.
Reaction:
2:19 understand how the rusting of iron may be prevented by: barrier methods, galvanising and sacrificial protection
Barrier Methods: Rusting may be prevented by stopping the water and oxygen getting to the iron with a barrier of grease, oil, paint or plastic.
Galvanising: (coating in zinc) also prevents water and oxygen getting to the iron, but with galvanising even if the barrier is broken the more reactive zinc corrodes before the less reactive iron. During the process, the zinc loses electrons to form zinc ions.
Sacrificial Protection: Zinc blocks are attached to iron boat hulls and underground pipelines to act as sacrificial anodes. Zinc is more reactive than iron, so oxygen in the air reacts with the zinc to form a layer of zinc oxide instead of the iron.
2:20 in terms of gain or loss of oxygen and loss or gain of electrons, understand the terms: oxidation, reduction, redox, oxidising agent, reducing agent, in terms of gain or loss of oxygen and loss or gain of electrons
Oxidation
- Oxidation is the loss of electrons. For example a sodium atom (Na) loses an electron to become a sodium ion (Na⁺). Another example is a chloride ion (Cl⁻) losing an electron to become a chlorine atom (Cl).
- Another definition of oxidation is the gain of oxygen. For example if carbon combines with oxygen to form carbon dioxide, the carbon is being oxidised.
Reduction
- Reduction is the gain of electrons. For example a sodium ion (Na⁺) gains an electron to become a sodium atom (Na). Another example is a chlorine atom (Cl) gaining an electron to become a chloride ion (Cl⁻).
- Another definition of reduction is the loss of oxygen. For example when aluminium oxide is broken down to produce aluminium and oxygen, the aluminium is being reduced.
Redox: A reaction involving oxidation and reduction.
A good way to remember the definitions of oxidation and reduction in terms of electrons is:
- OILRIG : Oxidation Is the Loss of electrons and Reduction Is the Gain of electrons
Oxidising agent: A substance that gives oxygen or removes electrons (it is itself reduced).
Reducing agent: A substance that takes oxygen or gives electrons (it is itself oxidised).
2:21 practical: investigate reactions between dilute hydrochloric and sulfuric acids and metals (e.g. magnesium, zinc and iron)
Metals which are above hydrogen in the reactivity series will react with dilute hydrochloric or sulfuric acid to produce a salt and hydrogen.
metal + acid → salt + hydrogen
For example:
magnesium + hydrochloric acid → magnesium chloride + hydrogen
Mg (s) + 2HCl (aq) → MgCl₂ (aq) + H₂ (g)
This is a displacement reaction.
There is a rapid fizzing and a colourless gas is produced. This gas pops with a lighted splint, showing the gas is hydrogen.
The reaction mixture becomes warm as heat is produced (exothermic).
The magnesium disappears to leave a colourless solution of magnesium chloride.
If more reactive metals are used instead of magnesium the reaction will be faster so the fizzing will be more vigorous and more heat will be produced.
2:28 describe the use of litmus, phenolphthalein and methyl orange to distinguish between acidic and alkaline solutions
Indicators are substances which change colour in response to a change in pH (acid or alkali).
| Indicator | Colour in acidic solution [H+] | Colour in alkaline solution [OH-] |
|---|---|---|
| Litmus | Red | Blue |
| Methyl orange | Red | Yellow |
| Phenolphthalein | Colourless | Pink |
Methyl orange is orange in a neutral solution
2:29 understand how to use the pH scale, from 0–14, can be used to classify solutions as strongly acidic (0–3), weakly acidic (4–6), neutral (7), weakly alkaline (8–10) and strongly alkaline (11–14)
The pH scale ranges from 0 to 14, and tells you how acidic or how alkaline a solution is.
| strongly acidic | weakly acidic | neutral | weakly alkaline | strongly alkaline | |
|---|---|---|---|---|---|
| pH | 0-3 | 4-6 | 7 | 8-10 | 11-14 |
2:30 describe the use of Universal Indicator to measure the approximate pH value of an aqueous solution
An indicator is a substance that has more than one colour form depending on the pH.
Universal indicator is a mixture of different dyes which change colour in a gradual way over a range of pH.
2:31 know that acids in aqueous solution are a source of hydrogen ions and alkalis in a aqueous solution are a source of hydroxide ions
An acid is source of hydrogen ions (H+).
An alkali is source of hydroxide ions (OH–).
2:32 know that bases can neutralise acids
Metal oxides, metal hydroxides and ammonia (NH₃) are called bases.
Bases neutralise acids by combining with the hydrogen ions in them.
The key reaction is:
acid + base → salt + water
An example of this is:
sulfuric acid + copper oxide → copper sulfate + water
H₂SO₄ + CuO → CuSO₄ + H₂O
2:34 know the general rules for predicting the solubility of ionic compounds in water: common sodium, potassium and ammonium compounds are soluble, all nitrates are soluble, common chlorides are soluble, except those of silver and lead(II), common sulfates are soluble, except for those of barium, calcium and lead(II), common carbonates are insoluble, except for those of sodium, potassium and ammonium, common hydroxides are insoluble except for those of sodium, potassium and calcium (calcium hydroxide is slightly soluble)
| Salt | Solubility | Exceptions |
|---|---|---|
| sodium (Na+), potassium (K+) and ammonium (NH4+) | soluble | none |
| nitrates (NO3-) | soluble | none |
| chlorides (Cl-) | soluble | silver chloride (AgCl) and lead (II) chloride (PbCl2) |
| sulfates (SO42-) | soluble | barium sulfate (BaSO4), calcium sulfate (CaSO4) and lead (II) sulfate (PbSO4) |
| carbonates (CO32-) | insoluble | sodium carbonate (Na2CO3), potassium carbonate (K2CO3) and ammonium carbonate ((NH4)2CO3) |
| hydroxides (OH-) | insoluble | sodium hydroxide (NaOH), potassium hydroxide (KOH) and calcium hydroxide (Ca(OH)2) (calcium hydroxide is slightly soluble) |
2:35 understand acids and bases in terms of proton transfer
An acid is a proton (H⁺) donor.
A base is a proton (H⁺) acceptor.
A proton is the same as a hydrogen ion. A good way to think about that is to realise that a hydrogen atom is just one proton and zero neutrons surrounded by only one electron. If that atom becomes an ion by the removal of the electron, then only one proton is left.
When sulfuric acid reacts with copper (II) oxide (CuO):
Cu²⁺O²⁻ (s) + H₂SO₄ (aq) → Cu²⁺ (aq) + SO₄²⁻ (aq) + H₂O (l)
H₂SO₄ is an acid. It donates protons (H⁺) to CuO, the base.
2:36 understand that an acid is a proton donor and a base is a proton acceptor
An acid is a proton donor.
A base is a proton acceptor.
A proton is the same as a hydrogen ion. A good way to think about that is to realise that a hydrogen atom is just one proton and zero neutrons surrounded by only one electron. If that atom becomes an ion by the removal of the electron, then only one proton is left.
2:37 describe the reactions of hydrochloric acid, sulfuric acid and nitric acid with metals, bases and metal carbonates (excluding the reactions between nitric acid and metals) to form salts
Acid reactions summary
alkali + acid → water + salt
base + acid → water + salt
carbonate + acid → water + salt + carbon dioxide
metal + acid → salt + hydrogen
To assist remembering this list, many pupils find it useful to remember this horrid looking but very effective mnemonic:
AAWS
BAWS
CAWS CoD
MASH
Acids are a source of hydrogen ions (H⁺) when in solution. When the hydrogen in an acid is replaced by a metal, the compound is called a salt. The name of the salt depends on the acid used. For example if sulfuric acid is used then a sulfate salt will be formed.
| Parent acid | Formula | Salt | Formula ion |
|---|---|---|---|
| sulfuric acid | H2SO4 | sulfate | SO42- |
| hydrochloric acid | HCl | chloride | Cl- |
| nitric acid | HNO3 | nitrate | NO3- |
Acid + Alkali and Acid + Base
A base is a substance that can neutralise an acid, forming a salt and water only.
Alkalis are soluble bases. When they react with acids, a salt and water is formed. The salt formed is often as a colourless solution. Alkalis are a source of hydroxide ions (OH⁻) when in solution.
alkali + acid → water + salt
base + acid → water + salt
Examples of acid + alkali reactions:
- sodium hydroxide + hydrochloric acid → sodium chloride + water
- NaOH (aq) + HCl (aq) → NaCl (aq) + H₂O (l)
- potassium hydroxide + sulfuric acid → potassium sulfate + water
- 2KOH (aq) + H₂SO₄ (aq) → K₂SO₄ (aq) + 2H₂O (l)
Example of an acid + base reaction:
CuO (s) + H₂SO₄ (aq) → CuSO₄ (aq) + H₂O (l)
Acid + Carbonate
carbonate + acid → water + salt + carbon dioxide
A carbonate is a compound made up of metal ions and carbonate ions. Examples of metal carbonates are sodium carbonate, copper carbonate and magnesium carbonate.
When carbonates react with acids, bubbling is observed which is the carbon dioxide being produced. If the acid is in excess the carbonate will disappear.
Examples of acid + carbonate reactions:
- calcium carbonate + hydrochloric acid → calcium chloride + water + carbon dioxide
- CaCO₃ (s) + 2HCl (aq) → CaCl₂ (aq) + H₂O (l) + CO₂ (g)
- potassium carbonate + hydrochloric acid → potassium chloride + water + carbon dioxide
- K₂CO₃ (aq) + 2HCl (aq) → 2KCl (aq) + H₂O (l) + CO₂ (g)
Acid + Metal
metal + acid → salt + hydrogen
Metals will react with an acid if the metal is above hydrogen in the reactivity series.
When metals react with acids, bubbling is observed which is the hydrogen being produced. If the acid is in excess the metal will disappear.
Examples of acid + metal reactions:
- magnesium + sulfuric acid → magnesium sulfate + hydrogen
- Mg (s) + H₂SO₄ (aq) → MgSO₄ (aq) + H₂ (g)
- aluminium + hydrochloric acid → aluminium chloride + hydrogen
- 2Al (s) + 6HCl (aq) → 2AlCl₃ (aq) + 3H₂ (g)
- copper + hydrochloric acid → no reaction (since copper is below hydrogen in the reactivity series)
2:38 know that metal oxides, metal hydroxides and ammonia can act as bases, and that alkalis are bases that are soluble in water
A base is a substance that neutralises an acid by combining with the hydrogen ions in them to produce water.
A base usually means a metal oxide, a metal hydroxide or ammonia.
Alkalis are bases which are soluble in water.
Some metal oxides are soluble in water and react with it to form solutions of metal hydroxides. For example:
Na₂O (s) + H₂O (l) → 2NaOH (aq)
Apart from this and other group 1 oxides (such as potassium oxide) most other metal oxides are not soluble in water.
One exception is calcium oxide which does dissolve slightly in water to form calcium hydroxide (known as limewater):
CaO (s) + H₂O (l) → Ca(OH)₂ (aq)
Ammonia is another base. Ammonia reacts with water to form ammonium ions and hydroxide ions:
NH₃ (aq) + H₂O (l) ⇋ NH₄⁺ (aq) + OH⁻ (aq)
All the solutions produced here contain hydroxide ions (OH⁻) so they are all alkalis.
2:39 describe an experiment to prepare a pure, dry sample of a soluble salt, starting from an insoluble reactant
Excess Solid Method:
Preparing pure dry crystals of copper sulfate (CuSO4) from copper oxide (CuO) and sulfuric acid (H2SO4)
| Step | Explanation |
|---|---|
| Heat acid (H2SO4) in a beaker | Speeds up the rate of reaction |
| Add base (CuO) until in excess (no more copper oxide dissolves) and stir with glass rod | Neutralises all the acid |
| Filter the mixture using filter paper and funnel | Removes any excess copper oxide |
| Gently heat the filtered solution (CuSO4) | To evaporate some of the water |
| until crystals form on a glass rod | Shows a hot saturated solution formed |
| Allow the solution to cool so that hydrated crystals form | Copper sulfate less soluble in cold solution |
| Remove the crystals by filtration | Removes crystals |
| Dry by leaving in a warm place | Evaporates the water |
2:42 practical: prepare a sample of pure, dry hydrated copper(II) sulfate crystals starting from copper(II) oxide
Excess Solid Method:
Preparing pure dry crystals of copper sulfate (CuSO4) from copper oxide (CuO) and sulfuric acid (H2SO4)
| Step | Explanation |
|---|---|
| Heat acid (H2SO4) in a beaker | Speeds up the rate of reaction |
| Add base (CuO) until in excess (no more copper oxide dissolves) and stir with glass rod | Neutralises all the acid |
| Filter the mixture using filter paper and funnel | Removes any excess copper oxide |
| Gently heat the filtered solution (CuSO4) | To evaporate some of the water |
| until crystals form on a glass rod | Shows a hot saturated solution formed |
| Allow the solution to cool so that hydrated crystals form | Copper sulfate less soluble in cold solution |
| Remove the crystals by filtration | Removes crystals |
| Dry by leaving in a warm place | Evaporates the water |
2:44a describe tests for these gases: hydrogen, carbon dioxide
Test for hydrogen gas (H2)
- Test: use a lit splint
- Result: burns with a squeaky pop
Test for carbon dioxide (CO2)
- Test: bubble through limewater
- Result: limewater turns cloudy
2:44 describe tests for these gases: hydrogen, oxygen, carbon dioxide, ammonia, chlorine
Tests for gases
| Gas | Test | Result if gas present |
|---|---|---|
| hydrogen (H2) | Use a lit splint | Gas pops |
| oxygen (O2) | Use a glowing splint | Glowing splint relights |
| carbon dioxide (CO2) | Bubble the gas through limewater | Limewater turns cloudy |
| ammonia (NH3) | Use red litmus paper | Turns damp red litmus paper blue |
| chlorine (Cl2) | Use damp litmus paper | Turns damp litmus paper white (bleaches) |
2:45 describe how to carry out a flame test
A flame test is used to show the presence of certain metal ions (cations) in a compound.
- A platinum or nichrome wire is dipped into concentrated hydrochloric acid to remove any impurities.
- The wire is dipped into the salt being tested so some salt sticks to the end.
- The wire and salt are held in a non-luminous (roaring) bunsen burner flame.
- The colour is observed.
Properties of the platinum or nichrome wire is:
- Inert
- High melting point
2:46 know the colours formed in flame tests for these cations: Li⁺ is red, Na⁺ is yellow, K⁺ is lilac, Ca²⁺ is orange-red, Cu²⁺ is blue-green
When put into a roaring bunsen burner flame on a nichrome wire, compounds containing certain cations will give specific colours as follows.
| Ion | Colour in flame test |
|---|---|
| lithium (Li⁺) | red |
| sodium (Na⁺) | yellow |
| potassium (K⁺) | lilac |
| calcium (Ca²⁺) | orange-red |
| copper (II) (Cu²⁺) | blue-green |
2:47 describe tests for these cations: NH₄⁺ using sodium hydroxide solution and identifying the gas evolved, Cu²⁺, Fe²⁺ and Fe³⁺ using sodium hydroxide solution
Underneath are the tests for : ammonium ion test, copper (II) ion test, iron (II) ion test, iron (III) ion test
Describe tests for the cation NH4+, using sodium hydroxide solution and identifying the ammonia evolved
Describe tests for the cations Cu2+, Fe2+ and Fe3+, using sodium hydroxide solution
First, add sodium hydroxide (NaOH), then observe the colour:
2:48a describe a test for CO₃²⁻ using hydrochloric acid and identifying the gas evolved
Describe tests for anions: Carbonate ions (CO32-)
2:48 describe tests for these anions: Cl⁻, Br⁻ and I⁻ using acidified silver nitrate solution, SO₄²⁻ using acidified barium chloride solution, CO₃²⁻ using hydrochloric acid and identifying the gas evolved
Describe tests for anions: Halide ions (Cl–, Br– and I–)
Underneath are the tests for:
chloride ion test, bromide ion test, iodide ion test, sulfate ion test, carbonate ion test
Describe tests for anions: Sulfate ions (SO42–)
Describe tests for anions: Carbonate ions (CO32-)
2:49 describe a test for the presence of water using anhydrous copper(II) sulfate
Add anhydrous copper (II) sulfate (CuSO4) to a sample.
If water is present the anhydrous copper (II) sulfate will change from white to blue.
2:50 describe a physical test to show whether a sample of water is pure
If the sample is pure water it will boil at 100oC
3:01 know that chemical reactions in which heat energy is given out are described as exothermic, and those in which heat energy is taken in are described as endothermic
Exothermic: chemical reaction in which heat energy is given out.
Endothermic: chemical reaction in which heat energy is taken in.
(So, in an exothermic reaction the heat exits from the chemicals so temperature rises)
3:02 describe simple calorimetry experiments for reactions such as combustion, displacement, dissolving and neutralisation
Calorimetry allows for the measurement of the amount of energy transferred in a chemical reaction to be calculated.
EXPERIMENT1: Displacement, dissolving and neutralisation reactions
Example: magnesium displacing copper from copper(II) sulfate
Method:
- 50 cm3 of copper(II) sulfate is measured and transferred into a polystyrene cup.
- The initial temperature of the copper sulfate solution is measured and recorded.
- Magnesium is added and the maximum temperature is measured and recorded.
- The temperature rise is then calculated. For example:
| Initial temp. of solution (oC) | Maximium temp. of solution (oC) | Temperature rise (oC) |
|---|---|---|
| 24.2 | 56.7 | 32.5 |
Note: mass of 50 cm3 of solution is 50 g
The cup used is polystyrene because:
polystyrene is an insulator which reduces heats loss
EXPERIMENT2: Combustion reactions
To measure the amount of energy produced when a fuel is burnt, the fuel is burnt and the flame is used to heat up some water in a copper container
Example: ethanol is burnt in a small spirit burner
Method:
- The initial mass of the ethanol and spirit burner is measured and recorded.
- 100cm3 of water is transferred into a copper container and the initial temperature is measured and recorded.
- The burner is placed under of copper container and then lit.
- The water is stirred constantly with the thermometer until the temperature rises by, say, 30 oC
- The flame is extinguished and the maximum temperature of the water is measured and recorded.
- The burner and the remaining ethanol is reweighed. For example:
| Mass of water (g) | Initial temp of water (oC) | Maximum temp of water (oC) | Temperature rise (oC) | Initial mass of spirit burner + ethanol (g) | Final mass of spirit burner + ethanol (g) | Mass of ethanol burnt (g) |
|---|---|---|---|---|---|---|
| 100 | 24.2 | 54.2 | 30.0 | 34.46 | 33.68 | 0.78 |
The amount of energy produced per gram of ethanol burnt can also be calculated:
3:03 calculate the heat energy change from a measured temperature change using the expression Q = mcΔT
Calorimetry allows for the measurement of the amount of energy transferred in a chemical reaction to be calculated.
EXPERIMENT1: Displacement, dissolving and neutralisation reactions
Example: magnesium displacing copper from copper(II) sulfate
Method:
- 50 cm3 of copper(II) sulfate is measured and transferred into a polystyrene cup.
- The initial temperature of the copper sulfate solution is measured and recorded.
- Magnesium is added and the maximum temperature is measured and recorded.
- The temperature rise is then calculated. For example:
| Initial temp. of solution (oC) | Maximium temp. of solution (oC) | Temperature rise (oC) |
|---|---|---|
| 24.2 | 56.7 | 32.5 |
Note: mass of 50 cm3 of solution is 50 g
EXPERIMENT2: Combustion reactions
To measure the amount of energy produced when a fuel is burnt, the fuel is burnt and the flame is used to heat up some water in a copper container
Example: ethanol is burnt in a small spirit burner
Method:
- The initial mass of the ethanol and spirit burner is measured and recorded.
- 100cm3 of water is transferred into a copper container and the initial temperature is measured and recorded.
- The burner is placed under of copper container and then lit.
- The water is stirred constantly with the thermometer until the temperature rises by, say, 30 oC
- The flame is extinguished and the maximum temperature of the water is measured and recorded.
- The burner and the remaining ethanol is reweighed. For example:
| Mass of water (g) | Initial temp of water (oC) | Maximum temp of water (oC) | Temperature rise (oC) | Initial mass of spirit burner + ethanol (g) | Final mass of spirit burner + ethanol (g) | Mass of ethanol burnt (g) |
|---|---|---|---|---|---|---|
| 100 | 24.2 | 54.2 | 30.0 | 34.46 | 33.68 | 0.78 |
The amount of energy produced per gram of ethanol burnt can also be calculated:
3:08 practical: investigate temperature changes accompanying some of the following types of change: salts dissolving in water, neutralisation reactions, displacement reactions and combustion reactions
Calorimetry allows for the measurement of the amount of energy transferred in a chemical reaction to be calculated.
EXPERIMENT1: Displacement, dissolving and neutralisation reactions
Example: magnesium displacing copper from copper(II) sulfate
Method:
- 50 cm3 of copper(II) sulfate is measured and transferred into a polystyrene cup.
- The initial temperature of the copper sulfate solution is measured and recorded.
- Magnesium is added and the maximum temperature is measured and recorded.
- The temperature rise is then calculated. For example:
| Initial temp. of solution (oC) | Maximium temp. of solution (oC) | Temperature rise (oC) |
|---|---|---|
| 24.2 | 56.7 | 32.5 |
Note: mass of 50 cm3 of solution is 50 g
EXPERIMENT2: Combustion reactions
To measure the amount of energy produced when a fuel is burnt, the fuel is burnt and the flame is used to heat up some water in a copper container
Example: ethanol is burnt in a small spirit burner
Method:
- The initial mass of the ethanol and spirit burner is measured and recorded.
- 100cm3 of water is transferred into a copper container and the initial temperature is measured and recorded.
- The burner is placed under of copper container and then lit.
- The water is stirred constantly with the thermometer until the temperature rises by, say, 30 oC
- The flame is extinguished and the maximum temperature of the water is measured and recorded.
- The burner and the remaining ethanol is reweighed. For example:
| Mass of water (g) | Initial temp of water (oC) | Maximum temp of water (oC) | Temperature rise (oC) | Initial mass of spirit burner + ethanol (g) | Final mass of spirit burner + ethanol (g) | Mass of ethanol burnt (g) |
|---|---|---|---|---|---|---|
| 100 | 24.2 | 54.2 | 30.0 | 34.46 | 33.68 | 0.78 |
The amount of energy produced per gram of ethanol burnt can also be calculated:
3:09 describe experiments to investigate the effects of changes in surface area of a solid, concentration of a solution, temperature and the use of a catalyst on the rate of a reaction
The rate of a chemical reaction can be measured either by how quickly reactants are used up or how quickly the products are formed.
The rate of reaction can be calculated using the following equation:
The units for rate of reaction will usually be grams per min (g/min)
An investigation of the reaction between marble chips and hydrochloric acid:
Marble chips, calcium carbonate (CaCO3) react with hydrochloric acid (HCl) to produce carbon dioxide gas. Calcium chloride solution is also formed.
Using the apparatus shown the change in mass of carbon dioxide can be measure with time.
As the marble chips react with the acid, carbon dioxide is given off.
The purpose of the cotton wool is to allow carbon dioxide to escape, but to stop any acid from spraying out.
The mass of carbon dioxide lost is measured at intervals, and a graph is plotted:
Experiment to investigate the effects of changes in surface area of solid on the rate of a reaction:
The experiment is repeated using the same mass of chips, but this time the chips are larger, i.e. have a smaller surface area.
Since the surface area is smaller, the rate of reaction is less.
Both sets of results are plotted on the same graph.
If instead the chips were smashed into powder (and again same mass of chips used) the surface area would be much larger and so the rate of reaction higher (steeper line on graph).
Experiment to investigate the effects of changes in concentration of solutions on the rate of a reaction:
The experiment is again repeated using the exact same quantities of everything but this time with half the concentration of acid. The marble chips must however be in excess. The reaction with the half the concentration of acid happens slower and produces half the amount of carbon dioxide.
Experiment to investigate the effects of changes in temperature on the rate of a reaction:
The experiment is once again repeated using the exact same quantities of everything but this time at a higher temperature. The reaction with the higher temperature happens faster.
Experiment to investigate the effects of the use of a catalyst on the rate of a reaction:
Hydrogen peroxide naturally decomposes slowly producing water and oxygen gas.
Manganese (IV) oxide can be used as a catalyst to speed up the rate of reaction.
The rate of reaction can be measured by measuring the volume of oxygen produced at regular intervals using a gas syringe.
Both sets of results are plotted on the same graph.
Experiment to investigate the reaction between varying concentrations of sodium thiosulfate and hydrochloric acid
Sodium thiosulfate (Na2S2O3) and hydrochloric acid (HCl) are both colourless solutions. They react to form a yellow precipitate of sulfur.
sodium thiosulfate + hydrochloric acid → sodium chloride + sulfur dioxide + sulfur + water
Na2S2O3(aq) + 2HCl(aq) → 2NaCl(aq) + SO2(g) + S(s) + H2O(l)
To investigate the effects of changes in concentration of sodium thiosulfate on the rate of a reaction, the conical flask is placed above a cross. The reaction mixture is observed from directly above and the time for a cross to disappear is measured. The cross disappears because a precipitate of sulfur is formed.
In order to change the concentration of sodium thiosulfate, the volumes of sodium thiosulfate and water are varied (see results table). However the total volume of solution must always be kept the same as to ensure that the depth of the solution remains constant.
In this reaction, sulfur dioxide gas (SO2), which is poisonous is produced therefore the experiment must be carried out in a well ventilated room.
The results are recorded in the table below and then plotted onto a graph.
| Volume of Na2S2O3(aq) (cm3) | Volume of water (cm3) | Concentration of Na2S2O3(aq) (mol/dm3) | Time taken for cross to disappear (s) | Rate of reaction (s-1) (1/time) |
|---|---|---|---|---|
| 50 | 0 | 0.10 | 45 | 0.0222 |
| 40 | 10 | 0.08 | 60 | 0.0167 |
| 30 | 20 | 0.06 | 80 | 0.0125 |
| 20 | 30 | 0.04 | 13 | 0.0769 |
| 10 | 40 | 0.02 | 255 | 0.0039 |
The graph shows that the rate of reaction is directly proportional to the concentration.
The experiment can also be repeated to show how temperature affects the rate of reaction.
In this experiment the concentration of sodium thiosulfate is kept constant but heated to range of different temperatures.
As a rough approximation, the rate of reaction doubles for every 10oC temperature rise.
3:10 describe the effects of changes in surface area of a solid, concentration of a solution, pressure of a gas, temperature and the use of a catalyst on the rate of a reaction
Increasing the surface area of a solid increases the rate of a reaction.
Increasing the concentration of a solution increases the rate of a reaction.
Increasing the pressure of a gas increases the rate of a reaction.
Increasing the temperature increases the rate of a reaction.
Using a catalyst increases the rate of a reaction.
3:11 explain the effects of changes in surface area of a solid, concentration of a solution, pressure of a gas and temperature on the rate of a reaction in terms of particle collision theory
Increasing the surface area of a solid:
- more particles exposed
- more frequent collisions
- increase the rate of a reaction
Increasing the concentration of a solution or pressure of a gas:
- more particles in same space
- more frequent collisions
- increase rate of reaction
Increasing the temperature:
- particles have more kinetic energy
- more frequent collisions
- and a higher proportion of those collisions are successful because the collision energy is greater or equal to the activation energy
- increase rate of reaction
3:12 know that a catalyst is a substance that increases the rate of a reaction, but is chemically unchanged at the end of the reaction
A catalyst is a substance that increases the rate of a reaction, but is chemically unchanged at the end of the reaction.