Predicting chemical reactions

Chemistry is not just about memorising what happens in a test tube; it is about finding the patterns that let you predict what will happen before you even mix the chemicals. By looking at how the Periodic Table is built, we can unlock these patterns.

In these study notes, we will explore how the electronic structures of atoms dictate their chemical behaviour, how to predict reactivity down key groups, and how to use experimental data to build a reactivity series of metals.

What you'll learn:

How the outer-shell electrons of Groups 1, 7, and 0 govern their physical and chemical properties.
How to predict trends in chemical reactivity down these groups.
The distinctive properties of transition metals (separate Chemistry only).
How to construct and use the reactivity series of metals based on reactions with water, acids, and salt solutions.

The Starting Point: Electronic Structure and the Periodic Table

To predict how an element will react, we must look at its atoms. The Periodic Table is organized so that elements with similar properties are placed in vertical columns called groups.

Definition

Group

A vertical column in the Periodic Table. Elements in the same group have the same number of electrons in their outer shell and share similar chemical properties.

Definition

Period

A horizontal row in the Periodic Table. The period number corresponds to the number of electron shells an element's atoms have.

When elements react, they do so to achieve a stable electronic configuration—typically a full outer shell of electrons (a stable octet). The ease with which an atom can lose or gain these outer-shell electrons determines how reactive it is.

Group 1: The Alkali Metals

Group 1 consists of highly reactive metals: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr).

Physical Properties of Group 1

Unlike typical tough metals, Group 1 metals are:

Soft enough to be cut easily with a knife.
Low in density (lithium, sodium, and potassium will actually float on water!).
Low in melting points, which decrease as you go down the group.

Chemical Properties of Group 1

All Group 1 atoms have one electron in their outer shell. To react and become stable, they must lose this single negative electron to form a single positive ion (a cation) with a $1 +$ charge (for example, $Na+\text{Na}^+$ ).

\text{Na} \to \text{Na}^+ + \text{e}^-

Because they only need to lose one electron, they are highly reactive. They react vigorously with:

Water: producing a metal hydroxide solution (alkali) and hydrogen gas ( $H2\text{H}_2$ ).
Oxygen: forming metal oxides, causing them to tarnish quickly in air.
Chlorine: forming solid white metal chlorides.

Trend in Reactivity Down Group 1

As you go down Group 1, reactivity increases.

To understand why, we look at the size of the atoms. As you go down the group, each element has more electron shells. This means:

The single outer electron is further away from the positive nucleus.
The inner electron shells shield the outer electron from the positive pull of the nucleus.
Therefore, the attraction between the nucleus and the outer electron is weaker, making the outer electron easier to lose.

Group 7: The Halogens

Group 7 elements are non-metals known as the halogens: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At).

Physical Properties of Group 7

The halogens exist as diatomic molecules (pairs of atoms joined by a covalent bond, such as $Cl2\text{Cl}_2$ , $Br2\text{Br}_2$ , $I2\text{I}_2$ ). As you go down the group, the melting and boiling points increase because the molecules get larger, resulting in stronger intermolecular forces that require more energy to break. This is why their states change at room temperature:

Fluorine ( $F2\text{F}_2$ ): pale yellow gas
Chlorine ( $Cl2\text{Cl}_2$ ): pale green gas
Bromine ( $Br2\text{Br}_2$ ): red-brown liquid (which vaporises easily into an orange gas)
Iodine ( $I2\text{I}_2$ ): dark grey solid (which sublimes into a purple vapour)

Chemical Properties of Group 7

All halogens have seven electrons in their outer shell. To achieve a stable full outer shell, they must gain one electron to form a halide ion with a $1 -$ charge (for example, $Cl−\text{Cl}^-$ ).

\text{Cl} + \text{e}^- \to \text{Cl}^-

They react with metals to form ionic salts (e.g., sodium chloride) and with non-metals to form covalent molecules (e.g., hydrogen chloride).

Trend in Reactivity Down Group 7

As you go down Group 7, reactivity decreases.

This is the exact opposite of Group 1! Let's explain why:

As you go down the group, the atoms get larger because they have more electron shells.
The outer shell is further away from the positive nucleus, and there is more shielding from inner shells.
This makes it harder for the positive nucleus to attract and capture the incoming electron needed to fill its outer shell.

Common Mistake

Confusing Group 1 and Group 7 trends

A very common exam mistake is to assume reactivity increases down all groups. Remember: Group 1 metals want to lose an electron (easier when big), while Group 7 non-metals want to gain an electron (easier when small).

Group 0: The Noble Gases

Group 0 consists of helium (He), neon (Ne), argon (Ar), krypton (Kr), and xenon (Xe).

Properties of Group 0

Unreactive (Inert): They do not easily react to form molecules. This is because they already have a full outer shell of electrons (helium has 2; the others have 8). They do not need to lose, gain, or share electrons.
Monatomic: They exist as single, unbonded atoms.
Low boiling points: Their boiling points increase as you go down the group because the atoms get heavier, increasing the weak attractive forces between them.

Transition Metals (Separate Chemistry Only)

Note: This section is only for students taking separate GCSE Chemistry (Triple Science); it does not appear in the Combined Science papers.

Between Groups 2 and 3 sits a large block of metallic elements known as the transition metals (such as iron, copper, chromium, and manganese).

Transition Metals vs. Group 1

While Group 1 elements are soft, highly reactive, and have low densities, transition metals fit our classic picture of metals. They have:

High melting points (except mercury, which is liquid at room temperature).
High densities (they feel heavy for their size).
Low reactivity (they do not react vigorously with water or oxygen).
High strength and hardness.

Unique Chemical Features

Transition metals have three main chemical tricks that set them apart:

They form coloured compounds: For example, copper(II) sulfate is bright blue, while iron(II) hydroxide is green.
They can form ions with different charges: Iron can form $Fe2+\text{Fe}^{2+}$ or $Fe3+\text{Fe}^{3+}$ ions. Copper can form $Cu+\text{Cu}^+$ or $Cu2+\text{Cu}^{2+}$ .
They make excellent catalysts: A catalyst is a substance that speeds up a chemical reaction without being used up itself. For example, iron is the catalyst used in the industrial Haber Process to make ammonia.

The Reactivity Series of Metals

By comparing how different metals react with water, dilute acids, and other metal solutions, we can arrange them into a list called the reactivity series.

Reactivity Series Ladder

Reacting Metals with Water

Highly reactive metals react with cold water to produce a metal hydroxide and hydrogen gas:

\text{Metal(s)} + \text{Water(l)} \to \text{Metal Hydroxide(aq)} + \text{Hydrogen(g)}

Potassium, Sodium, Calcium: React vigorously with cold water.
Magnesium: Reacts very slowly with cold water, but reacts rapidly with steam to form magnesium oxide ( $MgO\text{MgO}$ ) and hydrogen.
Zinc, Iron, Copper: Do not react with cold water.

Reacting Metals with Dilute Acid

Metals above hydrogen in the reactivity series react with dilute acids to produce a salt and hydrogen gas:

\text{Metal(s)} + \text{Acid(aq)} \to \text{Salt(aq)} + \text{Hydrogen(g)}

Potassium and Sodium: React too violently to be safely tested with acid in a school lab.
Calcium, Magnesium, Aluminium, Zinc, Iron: React with a steady stream of bubbles (hydrogen gas). The rate of bubbling indicates their relative reactivity.
Copper, Silver, Gold: Do not react with dilute acids at all.

Reactivity and Positive Ions (Higher Tier Only)

For Higher Tier, you must be able to explain metal reactivity in terms of ion formation.

Key Idea

Metal Reactivity = Ease of Cation Formation

The reactivity of a metal is directly related to its tendency to lose its outer-shell electrons to form its positive ion (cation).

When a metal reacts, it undergoes oxidation (loss of electrons). A metal that is high up in the reactivity series (like potassium) loses its outer electron very easily. A metal lower down (like copper) holds onto its electrons tightly, making it much harder to form a positive ion.

\text{K(s)} \to \text{K}^+ \text{(aq)} + \text{e}^- \quad (\text{Very easy - highly reactive})

\text{Cu(s)} \to \text{Cu}^{2+} \text{(aq)} + 2\text{e}^- \quad (\text{Difficult - low reactivity})

Displacement Reactions

A displacement reaction occurs when a more reactive metal takes the place of a less reactive metal in its compound.

For example, if you put a piece of reactive magnesium ribbon into a blue solution of copper sulfate, the magnesium will displace the copper:

\text{Magnesium(s)} + \text{Copper Sulfate(aq)} \to \text{Magnesium Sulfate(aq)} + \text{Copper(s)}

\text{Mg(s)} + \text{CuSO}_4\text{(aq)} \to \text{MgSO}_4\text{(aq)} + \text{Cu(s)}

During this reaction, the blue colour of the copper sulfate solution fades because magnesium sulfate is colourless, and red-brown copper metal deposits on the magnesium.

Tip

Displacement is a Competition

Think of a displacement reaction as a competition for a prize (the sulfate or oxide ion). The more reactive metal is stronger and will always win the prize away from a weaker, less reactive metal.

Deducing Reactivity from Experimental Results

You can determine the order of reactivity for unknown metals by reacting them systematically with different metal salt solutions and observing whether a reaction occurs.

Example

Deducing reactivity from displacement observations

A student reacted four unknown metals ( $W$ , $X$ , $Y$ , and $Z$ ) with solutions of their nitrates. The table below displays the results, where a tick ( $✓\checkmark$ ) represents a reaction and a cross ( $×\times$ ) represents no reaction.

Metal	Nitrate of $W$	Nitrate of $X$	Nitrate of $Y$	Nitrate of $Z$
$W$	—	$×\times$	$×\times$	$×\times$
$X$	$✓\checkmark$	—	$✓\checkmark$	$✓\checkmark$
$Y$	$✓\checkmark$	$×\times$	—	$✓\checkmark$
$Z$	$✓\checkmark$	$×\times$	$×\times$	—

Deduce the order of reactivity for these four metals, from most reactive to least reactive.

Count the total number of displacement reactions each metal can achieve. The more reactive a metal is, the more other metals it can displace from their compounds.
- Metal $X$ reacted with 3 solutions ( $W$ , $Y$ , and $Z$ ).
- Metal $Y$ reacted with 2 solutions ( $W$ and $Z$ ).
- Metal $Z$ reacted with 1 solution ( $W$ ).
- Metal $W$ reacted with 0 solutions.
Order the metals based on their reaction counts. Since $X$ displaced the most metals (3), it is the most reactive. $Y$ displaced fewer (2), so it is next. $Z$ displaced only one (1), and $W$ was unable to displace any other metal (0).
Formulate the final reactivity sequence. The reactivity order from most reactive to least reactive is:

X > Y > Z > W

Exam technique

In the exam

Explain the trends using the "three-step electronic rule": Whenever you are asked to explain why reactivity increases or decreases down a group, make sure to state:
- How the atomic radius / number of shells changes.
- How this affects the strength of attraction between the outer electrons and the positive nucleus.
- How easily the outer electrons are lost (Group 1) or gained (Group 7).
Include state symbols: In equation questions, you are often expected to provide state symbols: $(s)\text{(s)}$ for solid, $(l)\text{(l)}$ for liquid, $(g)\text{(g)}$ for gas, and $(aq)\text{(aq)}$ for aqueous (dissolved in water). Remember that acids and metal salt solutions are always $(aq)\text{(aq)}$ .
Link reactivity to ions (Higher Tier): If asked why a metal is more reactive than another, phrase your answer in terms of its "tendency to form positive ions".

Self review

Check yourself

Why does potassium react more vigorously with water than sodium does? Write your answer in terms of atomic structure.
Chlorine is bubbled through a solution of potassium bromide. Write down what you would expect to see, and complete the balanced chemical equation: $Cl2(g)+2KBr(aq)→…\text{Cl}_2\text{(g)} + 2\text{KBr(aq)} \to \dots$
[Higher Tier] Explain why magnesium is able to displace zinc from zinc sulfate solution, making reference to electron transfer and ion formation.

Predicting chemical reactions

What you'll learn:

How the outer-shell electrons of Groups 1, 7, and 0 govern their physical and chemical properties.
How to predict trends in chemical reactivity down these groups.
The distinctive properties of transition metals (separate Chemistry only).
How to construct and use the reactivity series of metals based on reactions with water, acids, and salt solutions.

The Starting Point: Electronic Structure and the Periodic Table

To predict how an element will react, we must look at its atoms. The Periodic Table is organized so that elements with similar properties are placed in vertical columns called groups.

Definition

Group

A vertical column in the Periodic Table. Elements in the same group have the same number of electrons in their outer shell and share similar chemical properties.

Definition

Period

A horizontal row in the Periodic Table. The period number corresponds to the number of electron shells an element's atoms have.

Group 1: The Alkali Metals

Group 1 consists of highly reactive metals: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr).

Physical Properties of Group 1

Unlike typical tough metals, Group 1 metals are:

Soft enough to be cut easily with a knife.
Low in density (lithium, sodium, and potassium will actually float on water!).
Low in melting points, which decrease as you go down the group.

Chemical Properties of Group 1

\text{Na} \to \text{Na}^+ + \text{e}^-

Because they only need to lose one electron, they are highly reactive. They react vigorously with:

Water: producing a metal hydroxide solution (alkali) and hydrogen gas ( $H2\text{H}_2$ ).
Oxygen: forming metal oxides, causing them to tarnish quickly in air.
Chlorine: forming solid white metal chlorides.

Trend in Reactivity Down Group 1

As you go down Group 1, reactivity increases.

To understand why, we look at the size of the atoms. As you go down the group, each element has more electron shells. This means:

The single outer electron is further away from the positive nucleus.
The inner electron shells shield the outer electron from the positive pull of the nucleus.
Therefore, the attraction between the nucleus and the outer electron is weaker, making the outer electron easier to lose.

Group 7: The Halogens

Group 7 elements are non-metals known as the halogens: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At).

Physical Properties of Group 7

Fluorine ( $F2\text{F}_2$ ): pale yellow gas
Chlorine ( $Cl2\text{Cl}_2$ ): pale green gas
Bromine ( $Br2\text{Br}_2$ ): red-brown liquid (which vaporises easily into an orange gas)
Iodine ( $I2\text{I}_2$ ): dark grey solid (which sublimes into a purple vapour)

Chemical Properties of Group 7

\text{Cl} + \text{e}^- \to \text{Cl}^-

They react with metals to form ionic salts (e.g., sodium chloride) and with non-metals to form covalent molecules (e.g., hydrogen chloride).

Trend in Reactivity Down Group 7

As you go down Group 7, reactivity decreases.

This is the exact opposite of Group 1! Let's explain why:

As you go down the group, the atoms get larger because they have more electron shells.
The outer shell is further away from the positive nucleus, and there is more shielding from inner shells.
This makes it harder for the positive nucleus to attract and capture the incoming electron needed to fill its outer shell.

Common Mistake

Confusing Group 1 and Group 7 trends

Group 0: The Noble Gases

Group 0 consists of helium (He), neon (Ne), argon (Ar), krypton (Kr), and xenon (Xe).

Properties of Group 0

Unreactive (Inert): They do not easily react to form molecules. This is because they already have a full outer shell of electrons (helium has 2; the others have 8). They do not need to lose, gain, or share electrons.
Monatomic: They exist as single, unbonded atoms.
Low boiling points: Their boiling points increase as you go down the group because the atoms get heavier, increasing the weak attractive forces between them.

Transition Metals (Separate Chemistry Only)

Note: This section is only for students taking separate GCSE Chemistry (Triple Science); it does not appear in the Combined Science papers.

Between Groups 2 and 3 sits a large block of metallic elements known as the transition metals (such as iron, copper, chromium, and manganese).

Transition Metals vs. Group 1

While Group 1 elements are soft, highly reactive, and have low densities, transition metals fit our classic picture of metals. They have:

High melting points (except mercury, which is liquid at room temperature).
High densities (they feel heavy for their size).
Low reactivity (they do not react vigorously with water or oxygen).
High strength and hardness.

Unique Chemical Features

Transition metals have three main chemical tricks that set them apart:

They form coloured compounds: For example, copper(II) sulfate is bright blue, while iron(II) hydroxide is green.
They can form ions with different charges: Iron can form $Fe2+\text{Fe}^{2+}$ or $Fe3+\text{Fe}^{3+}$ ions. Copper can form $Cu+\text{Cu}^+$ or $Cu2+\text{Cu}^{2+}$ .
They make excellent catalysts: A catalyst is a substance that speeds up a chemical reaction without being used up itself. For example, iron is the catalyst used in the industrial Haber Process to make ammonia.

The Reactivity Series of Metals

By comparing how different metals react with water, dilute acids, and other metal solutions, we can arrange them into a list called the reactivity series.

Reactivity Series Ladder

Reacting Metals with Water

Highly reactive metals react with cold water to produce a metal hydroxide and hydrogen gas:

\text{Metal(s)} + \text{Water(l)} \to \text{Metal Hydroxide(aq)} + \text{Hydrogen(g)}

Potassium, Sodium, Calcium: React vigorously with cold water.
Magnesium: Reacts very slowly with cold water, but reacts rapidly with steam to form magnesium oxide ( $MgO\text{MgO}$ ) and hydrogen.
Zinc, Iron, Copper: Do not react with cold water.

Reacting Metals with Dilute Acid

Metals above hydrogen in the reactivity series react with dilute acids to produce a salt and hydrogen gas:

\text{Metal(s)} + \text{Acid(aq)} \to \text{Salt(aq)} + \text{Hydrogen(g)}

Potassium and Sodium: React too violently to be safely tested with acid in a school lab.
Calcium, Magnesium, Aluminium, Zinc, Iron: React with a steady stream of bubbles (hydrogen gas). The rate of bubbling indicates their relative reactivity.
Copper, Silver, Gold: Do not react with dilute acids at all.

Reactivity and Positive Ions (Higher Tier Only)

For Higher Tier, you must be able to explain metal reactivity in terms of ion formation.

Key Idea

Metal Reactivity = Ease of Cation Formation

The reactivity of a metal is directly related to its tendency to lose its outer-shell electrons to form its positive ion (cation).

\text{K(s)} \to \text{K}^+ \text{(aq)} + \text{e}^- \quad (\text{Very easy - highly reactive})

\text{Cu(s)} \to \text{Cu}^{2+} \text{(aq)} + 2\text{e}^- \quad (\text{Difficult - low reactivity})

Displacement Reactions

A displacement reaction occurs when a more reactive metal takes the place of a less reactive metal in its compound.

For example, if you put a piece of reactive magnesium ribbon into a blue solution of copper sulfate, the magnesium will displace the copper:

\text{Magnesium(s)} + \text{Copper Sulfate(aq)} \to \text{Magnesium Sulfate(aq)} + \text{Copper(s)}

\text{Mg(s)} + \text{CuSO}_4\text{(aq)} \to \text{MgSO}_4\text{(aq)} + \text{Cu(s)}

During this reaction, the blue colour of the copper sulfate solution fades because magnesium sulfate is colourless, and red-brown copper metal deposits on the magnesium.

Tip

Displacement is a Competition

Think of a displacement reaction as a competition for a prize (the sulfate or oxide ion). The more reactive metal is stronger and will always win the prize away from a weaker, less reactive metal.

Deducing Reactivity from Experimental Results

You can determine the order of reactivity for unknown metals by reacting them systematically with different metal salt solutions and observing whether a reaction occurs.

Example

Deducing reactivity from displacement observations

Metal	Nitrate of $W$	Nitrate of $X$	Nitrate of $Y$	Nitrate of $Z$
$W$	—	$×\times$	$×\times$	$×\times$
$X$	$✓\checkmark$	—	$✓\checkmark$	$✓\checkmark$
$Y$	$✓\checkmark$	$×\times$	—	$✓\checkmark$
$Z$	$✓\checkmark$	$×\times$	$×\times$	—

Deduce the order of reactivity for these four metals, from most reactive to least reactive.

Count the total number of displacement reactions each metal can achieve. The more reactive a metal is, the more other metals it can displace from their compounds.
- Metal $X$ reacted with 3 solutions ( $W$ , $Y$ , and $Z$ ).
- Metal $Y$ reacted with 2 solutions ( $W$ and $Z$ ).
- Metal $Z$ reacted with 1 solution ( $W$ ).
- Metal $W$ reacted with 0 solutions.
Order the metals based on their reaction counts. Since $X$ displaced the most metals (3), it is the most reactive. $Y$ displaced fewer (2), so it is next. $Z$ displaced only one (1), and $W$ was unable to displace any other metal (0).
Formulate the final reactivity sequence. The reactivity order from most reactive to least reactive is:

X > Y > Z > W

Exam technique

In the exam

Explain the trends using the "three-step electronic rule": Whenever you are asked to explain why reactivity increases or decreases down a group, make sure to state:
- How the atomic radius / number of shells changes.
- How this affects the strength of attraction between the outer electrons and the positive nucleus.
- How easily the outer electrons are lost (Group 1) or gained (Group 7).
Include state symbols: In equation questions, you are often expected to provide state symbols: $(s)\text{(s)}$ for solid, $(l)\text{(l)}$ for liquid, $(g)\text{(g)}$ for gas, and $(aq)\text{(aq)}$ for aqueous (dissolved in water). Remember that acids and metal salt solutions are always $(aq)\text{(aq)}$ .
Link reactivity to ions (Higher Tier): If asked why a metal is more reactive than another, phrase your answer in terms of its "tendency to form positive ions".

Self review

Check yourself

Why does potassium react more vigorously with water than sodium does? Write your answer in terms of atomic structure.
Chlorine is bubbled through a solution of potassium bromide. Write down what you would expect to see, and complete the balanced chemical equation: $Cl2(g)+2KBr(aq)→…\text{Cl}_2\text{(g)} + 2\text{KBr(aq)} \to \dots$
[Higher Tier] Explain why magnesium is able to displace zinc from zinc sulfate solution, making reference to electron transfer and ion formation.

Predicting chemical reactions

What you'll learn:

The Starting Point: Electronic Structure and the Periodic Table

Group 1: The Alkali Metals

Physical Properties of Group 1

Chemical Properties of Group 1

Trend in Reactivity Down Group 1

Group 7: The Halogens

Physical Properties of Group 7

Chemical Properties of Group 7

Trend in Reactivity Down Group 7

Group 0: The Noble Gases

Properties of Group 0

Transition Metals (Separate Chemistry Only)

Transition Metals vs. Group 1

Unique Chemical Features

The Reactivity Series of Metals

Reacting Metals with Water

Reacting Metals with Dilute Acid

Reactivity and Positive Ions (Higher Tier Only)

Displacement Reactions

Deducing Reactivity from Experimental Results

Recap questions

You've reached the end

Predicting chemical reactions Revision Guide

Predicting chemical reactions

What you'll learn:

The Starting Point: Electronic Structure and the Periodic Table

Group 1: The Alkali Metals

Physical Properties of Group 1

Chemical Properties of Group 1

Trend in Reactivity Down Group 1

Group 7: The Halogens

Physical Properties of Group 7

Chemical Properties of Group 7

Trend in Reactivity Down Group 7

Group 0: The Noble Gases

Properties of Group 0

Transition Metals (Separate Chemistry Only)

Transition Metals vs. Group 1

Unique Chemical Features

The Reactivity Series of Metals

Reacting Metals with Water

Reacting Metals with Dilute Acid

Reactivity and Positive Ions (Higher Tier Only)

Displacement Reactions

Deducing Reactivity from Experimental Results

Recap questions

You've reached the end

Predicting chemical reactions Revision Guide