Electrolysis – Tyler de Witt video
Although quite long, this Tyler de Witt video is a good summary of Electrolysis.
1:59 (Triple only) write ionic half-equations representing the reactions at the electrodes during electrolysis and understand why these reactions are classified as oxidation or reduction
Oxidation: the loss of electrons or the gain of oxygen
Reduction: the gain of electrons or the loss of oxygen
Example: The electrolysis of lead (II) bromide, PbBr2
At the cathode (negative electrode): Pb2+ (l) + 2e– → Pb (l) (reduction)
At the anode (positive electrode): 2Br– (l) → Br2 (g) + 2e– (oxidation)
Example: The electrolysis of aluminium oxide, Al2O3
At the cathode: Al3+ + 3e– → Al (reduction)
At the anode: 2O2- → O2 + 4e– (oxidation)
Example: The electrolysis of sodium chloride solution (NaCl (aq))
At the cathode: 2H+ (aq) + 2e– → H2 (g) (reduction)
At the anode: 2Cl– (aq) → Cl2 (g) + 2e– (oxidation)
Example: The electrolysis of copper sulfate solution (CuSO4 (aq))
At the cathode: Cu2+ (aq) + 2e– → Cu (s) (reduction)
At the anode: 4OH– (aq) → O2 (g) + 2H2O (l) + 4e– (oxidation)
1:60 (Triple only) practical: investigate the electrolysis of aqueous solutions
The diagram shows an electrolytic cell.
The electrolyte is an aqueous solution. For example it might be concentrated sodium chloride, NaCl (aq).
The test tubes over the electrodes must not completely cover them to make sure the ions are free to move throughout the solution.
In the case of NaCl (aq) bubbles of gas will be seen forming at the electrodes. These float up and collect in the test tubes when each gas can be tested to assess its identity.
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) | ? | ? | ? | ? |
Alkali metals and water – videos
The awesome Professor Sir Martyn Poliakoff shows us the reaction of some alkali metals and water.
2:04 (Triple only) explain the trend in reactivity in Group 1 in terms of electronic configurations
As you go down the group the outer electron lost from the group 1 metal is further from the nucleus therefore the electron is less attracted by the nucleus and therefore more easily lost.
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 |
Fluorine – video
Here is a great video on Fluorine, which is the most reactive halogen (group 7):
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.
Halogen reactions with iron wool – video
The reactions of various halogens with iron wool is a great way to see how reactive one halogen is compared to another:
Halogen displacement practical – video
Special thanks to Charlie R for suggesting I add a video on halogen displacement.
2:08 (Triple only) explain the trend in reactivity in Group 7 in terms of electronic configurations
The higher up we go in group 7 (halogens) of the periodic table, the more reactive the element. The explanation concerns how readily these elements form ions, by attracting a passing electron to fill the outer shell.
In fluorine the outer electron shell is very close to the positively charged nucleus, so the attraction between this nucleus and the negatively charged electrons is very strong. This means fluorine is very reactive indeed.
However, for iodine the outer electron shell is much further from the nucleus so the attraction is weaker. This means iodine is less reactive.
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)
Reactions with oxygen – video
This video shows the reaction with oxygen of various elements:
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.
Thermite displacement demonstration goes wrong – video
The thermite reaction is used to repair railways: the reaction between aluminium and iron oxide is highly exothermic and produces molten iron.
It should be properly controlled if demonstrated in the lab: watch this video and see how many mistakes in basic lab safety made by this teacher.
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.
Galvanising – video
This video shows the galvanising of pieces of steel to protect them from rusting:
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.
2:22 (Triple only) know that most metals are extracted from ores found in the Earth’s crust and that unreactive metals are often found as the uncombined element
Most metals are found in the Earth’s crust combined with other elements. Such compounds are found in rocks called ore, rocks from which it is worthwhile to extract a metal.
A few very unreactive metals, such as gold, are found native which means they are found in the Earth’s crust as the uncombined element.
2:23 (Triple only) explain how the method of extraction of a metal is related to its position in the reactivity series, illustrated by carbon extraction for iron and electrolysis for aluminium
Extraction of a metal from its ore typically involves removing oxygen from metal oxides.
If the ore contains a metal which is below carbon in the reactivity series then the metal is extracted by reaction with carbon in a displacement reaction.
If the ore contains a metal which is above carbon in the reactivity series then electrolysis (or reaction with a more reactive metal) is used to extract the metal.
2:24 (Triple only) be able to comment on a metal extraction process, given appropriate information
Extraction of a metal from its ore typically involves removing oxygen from metal oxides.
If the ore contains a metal which is below carbon in the reactivity series then the metal is extracted by reaction with carbon in a displacement reaction.
If the ore contains a metal which is above carbon in the reactivity series then electrolysis (or reaction with a more reactive metal) is used to extract the metal.
2:25 (Triple only) explain the uses of aluminium, copper, iron and steel in terms of their properties the types of steel will be limited to low-carbon (mild), high-carbon and stainless
| Aluminium | |
|---|---|
| Use | Property |
| Aircrafts and cans | Low density / resists corrosion |
| Power cables | Conducts electricity / ductile |
| Pots and pans | Low density / strong (when alloyed) / good conductor of electricity and heat |
Aluminium resists corrosion because it has a very thin, but very strong, layer of aluminium oxide on the surface.
| Copper | |
|---|---|
| Use | Property |
| Electrical wires | very good conductor of electricity and ductile |
| Pots and pans | very good conductor of heat / very unreactive / malleable |
| Water pipes | unreactive / malleable |
| Surfaces in hospitals | antimicrobial properties / malleable |
| Iron | |
|---|---|
| Use | Property |
| Buildings | Strong |
| Saucepans | Conducts heat / high melting point / malleable |
| Steel | ||
|---|---|---|
| Type of steel | Iron mixed with | Some uses |
| Mild steel | up to 0.25% carbon | nails, car bodies, ship building, girders |
| High-carbon steel | 0.6%-1.2% carbon | cutting tools, masonry nails |
| Stainless steel | Chromium (and nickel) | cutlery, cooking utensils, kitchen sinks |
Mild steel is a strong material that can easily be hammered into various shapes (malleable). It rusts easily.
High-carbon steel is harder than mild steel but more brittle (not as malleable).
Stainless steel forms a strong, protective oxide layer so is very resistant to corrosion.
2:26 (Triple only) know that an alloy is a mixture of a metal and one or more elements, usually other metals or carbon
An alloy is a mixture of a metal with, usually, other metals or carbon.
For example, brass is a alloy of copper and zinc, and steel is an alloy of iron and carbon.
2:27 (Triple only) explain why alloys are harder than pure metals
Alloys are harder than the individual pure metals from which they are made.
In an alloy, the different elements have slightly different sized atoms. This breaks up the regular lattice arrangement and makes it more difficult for layers of ions to slide over each other.
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
2:33 (Triple only) describe how to carry out an acid-alkali titration
Titration is used to find out precisely how much acid neutralises a certain volume of alkali (or vice versa).
The diagram shows the titration method for a neutralisation reaction between hydrochloric acid and sodium hydroxide, using phenolphthalein as an indicator. The indicator changes colour when neutralisation occurs.
The conical flask is swirled to mix the solutions each time alkali is added. When reading the burette it is important to be aware that the numbers on the scale increase from top to bottom. Readings are usually recorded to the nearest 0.05cm³ so all readings should be written down with 2 decimal places. The second decimal place is given as a ‘0’ if the level of the solution is on a line, or ‘5’ if it is between the lines. The volume of alkali added is calculated by subtracting the final reading from the initial reading. Various indicators can be used such as phenolphthalein or methyl orange. However universal indicator should not be used since it has a wide range of colours rather than one specific colour change so it would be unclear when the precise endpoint of titration was achieved.
This process is repeated a number of times. The first time it is done roughly to get a good approximation of how much alkali needs to be added. On subsequent attempts, the alkali is added very slowly when approaching the correct volume.
(Triple only) Titration calculations – videos
Here are a couple of videos explaining how to perform titration calculations:
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)