Topic: Key Reactions

Key Reactions quiz

1. State the formula for the iron (II) ion

Question 1 of 14

2. Can a base neutralise an acid?

Question 2 of 14

3. What is the universal indicator colour of a pH 8-10 solution?

Question 3 of 14

4. hydrochloric acid + sodium hydroxide →

Question 4 of 14

5. Write the word and chemical equation for the thermal decompostion of calcium carbonate, CaCO₃

Question 5 of 14

6. What is the pH of a weak alkali?

Question 6 of 14

7. What is the formula for calcium nitrate?

Question 7 of 14

8. If an element conducts electricity, is it typically a metal or a non-metal?

Question 8 of 14

9. State the colour of phenolphthalein in acidic and alkaline solutions

Question 9 of 14

10. Is ammonia acidic or basic?

Question 10 of 14

11. When magnesium burns in air, is the oxide acidic or basic?

Question 11 of 14

12. Describe the test for hydrogen gas

Question 12 of 14

13. Balance the equation for the reaction: ___ HCl(aq) + ___ MnO₂(s) → ___MnCl₂(aq) + ___H₂O(l) + ___ Cl₂(g)

Question 13 of 14

14. What is an acid?

Question 14 of 14


2020-02-16T16:28:50+00:00Categories: Uncategorized|Tags: , |

1:20 understand how to use electrical conductivity and the acid-base character of oxides to classify elements as metals or non-metals


  • conduct electricity
  • have oxides which are basic, reacting with acids to give a salt and water


Non – Metals

  • do not conduct electricity (except for graphite)
  • have oxides which are acidic or neutral


1:25 write word equations and balanced chemical equations (including state symbols): for reactions studied in this specification and for unfamiliar reactions where suitable information is provided


Sodium (Na) reacts with water (H2O) to produce a solution of sodium hydroxide (NaOH) and hydrogen gas (H2).

Word equation:

     sodium + water –> sodium hydroxide + hydrogen

Writing the chemical equation

A chemical equation represents what happens in terms of atoms in a chemical reaction.

Step 1: To write a chemical equation we need to know the chemical formulae of the substances.

     Na + H2O –> NaOH + H2

Step 2: The next step is to balance the equation: write a large number before each compound so the number of atoms of each element on the left hand side (reactants) matches the number on the right (products). This large number is the amount of each compound or element.

During this balancing stage the actual formulas for each compound must not be changed. Only the number of each compound changes.

     2Na + 2H2O –> 2NaOH + H2

If asked for an equation, the chemical equation must be given.


State symbols are used to show what physical state the reactants and products are in.

State symbolsPhysical state
(aq)Aqueous solution (dissolved in water)


A solid piece of sodium (Na) reacts with water (H2O) to produce a solution of sodium hydroxide (NaOH) and hydrogen gas (H2).

     2Na(s) + 2H2O(l) –> 2NaOH(aq) + H2(g)

1:38a understand how to use the charges of these ions in ionic formulae: metals in Groups 1, 2 and 3, non-metals in Groups 5, 6 and 7, Ag⁺, Cu²⁺, Fe²⁺, Fe³⁺, Pb²⁺, Zn²⁺, hydrogen (H⁺), hydroxide (OH⁻), ammonium (NH₄⁺), carbonate (CO₃²⁻), nitrate (NO₃⁻), sulfate (SO₄²⁻)

When given this information of the following ions, it is possible to work out the formulae of ionic compounds which include these ions.

Name of IonFormulaCharge
Silver ionAg++1
Zinc ionZn2++2
Hydrogen ionH++1
Copper (II) ionCu2++2
Iron (II) ionFe2++2
Iron (III) ionFe3++3
Lead (II) ionPb2++2

Ion charges on the periodic table

1:39 write formulae for compounds formed between the ions listed in 1:38

Writing the electron configuration of an atom allows you to work out the electron configuration of the ion and therefore the charge on the ion.



Atom = Mg

Electron configuration = 2,8,2

remove the two electrons from the outer shell to achieve the same electron configuration as the nearest noble gas, Neon (Ne 2,8)

Ion = Mg2+ 



Atom = O

Electron configuration = 2,6

add two electrons to the outer shell to achieve the same electron configuration as the nearest noble gas, Neon (Ne 2,8)

Ion = O2-  [2,8]2-


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)

2:28 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).

IndicatorColour in acidic solutionColour in alkaline solution
Methyl orangeRedYellow

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 acidicweakly acidicneutralweakly alkalinestrongly alkaline

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

2:37 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:



         CAWS CoD


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 acidFormulaSaltFormula ion
sulfuric acidH2SO4sulfateSO42-
hydrochloric acidHClchlorideCl-
nitric acidHNO3nitrateNO3-


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)

3 reactions of acid – video

Underneath are 3 acid reactions :

   alkali      +      acid      →      water      +      salt

   base      +      acid      →      water      +      salt

   carbonate      +      acid      →      water      +      salt      +      carbodioxide

2022-09-15T07:55:36+00:00Categories: Uncategorized|Tags: , , , |

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.



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Section 1: Principles of chemistry

      a) States of matter

      b) Atoms

      c) Atomic structure

     d) Relative formula masses and molar volumes of gases

     e) Chemical formulae and chemical equations

     f) Ionic compounds

     g) Covalent substances

     h) Metallic crystals

     i) Electrolysis

 Section 2: Chemistry of the elements

     a) The Periodic Table

     b) Group 1 elements: lithium, sodium and potassium

     c) Group 7 elements: chlorine, bromine and iodine

     d) Oxygen and oxides

     e) Hydrogen and water

     f) Reactivity series

     g) Tests for ions and gases

Section 3: Organic chemistry

     a) Introduction

     b) Alkanes

     c) Alkenes

     d) Ethanol

Section 4: Physical chemistry

     a) Acids, alkalis and salts

     b) Energetics

     c) Rates of reaction

     d) Equilibria

Section 5: Chemistry in industry

     a) Extraction and uses of metals

     b) Crude oil

     c) Synthetic polymers

     d) The industrial manufacture of chemicals

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