Metals and non-metals — AQA GCSE Combined Science
Test yourself on Metals and non-metals with AQA GCSE practice questions.
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Metals and non-metals explained
This statement links an element's position in the periodic table to the type of ion it forms.
Read the full explanation
Metals occupy the left and centre of the periodic table, and their atoms lose electrons from the outer shell so that the number of protons exceeds the number of electrons, giving a positive ion. For example, a sodium atom (2,8,1) loses its single outer electron to become Na⁺, and a magnesium atom (2,8,2) loses two electrons to become Mg²⁺. Non-metals, found towards the upper right, gain or share electrons and typically form negative ions such as O²⁻ and Cl⁻. The statement is a classification rule: if an element reacts by forming positive ions, it is a metal. It does not claim that every metal reaction produces ions, nor that metals never form covalent bonds.
Elements that do not form positive ions are non-metals.
This statement gives a practical way to classify an element as a non-metal: if it does not form positive ions, it is a non-metal. Metals, on the left and centre of the periodic table, lose outer-shell electrons to form positive ions such as Na⁺, Mg²⁺ and Al³⁺. Non-metals, mostly on the right and top, gain or share electrons; they form negative ions such as O²⁻ and Cl⁻, or covalent bonds, and do not form positive ions. For example, oxygen gains two electrons to become O²⁻, and chlorine gains one electron to become Cl⁻. Hydrogen is a non-metal that forms H⁺ in acids, so the statement is a general rule with that familiar exception. Use electron configuration and position in the periodic table to decide whether an element tends to lose or gain electrons.
The majority of elements are metals. Metals are found to the left and towards the bottom of the periodic table. Non-metals are found towards the right and top of the periodic table.
The periodic table separates elements into metals and non-metals, and most known elements are metals. Metals occupy the left side and lower regions, while non-metals cluster towards the right and top. A stepped line running below boron and down through silicon, arsenic and tellurium roughly divides the two groups. Elements touching this line, such as silicon, are metalloids with intermediate properties. Position links to structure: metals lose electrons to form positive ions, whereas non-metals gain or share electrons. For example, sodium on the far left forms Na⁺, while chlorine on the right forms Cl⁻. Hydrogen is a non-metal placed at the top left because of its single electron shell. Use the table's position to predict whether an element is a metal or non-metal.
explain the differences between metals and non-metals on the basis of their characteristic physical and chemical properties. This links to Group 0 (page 74), Group 1 (page 74), Group 7 (page 75) and Bonding, structure and the properties of matter (page 75)
Metals and non-metals differ in physical properties like conductivity and melting point, and in chemical behaviour. Metals have giant structures with delocalised electrons, meaning they conduct electricity, are malleable, and usually have high melting points. Non-metals typically form simple molecules or giant covalent structures; they are often brittle, poor conductors (with exceptions like graphite, which conducts via delocalised electrons), and simple molecular non-metals have lower melting points. Chemically, metals form positive ions by losing electrons. For example, sodium (Group 1) reacts vigorously with water. Non-metals gain or share electrons. Chlorine (Group 7) forms negative ions, while Group 0 noble gases are unreactive non-metals with full outer shells.
explain how the atomic structure of metals and non-metals relates to their position in the periodic table
An element's position in the periodic table is determined by its atomic structure, particularly its atomic number and electron arrangement. Metals occupy the left and centre of the periodic table because their atoms have relatively few outer electrons and lose them easily to form positive ions. Non-metals sit on the right and top, with atoms that have more outer electrons and tend to gain or share electrons. For Groups 1 to 7, the group number equals the number of outer electrons. Group 0 elements, however, have full outer shells (2 for helium, 8 for the rest). The period number equals the number of occupied electron shells. For example, sodium (Group 1, Period 3) has one outer electron and three shells; chlorine (Group 7, Period 3) has seven outer electrons and three shells.
explain how the reactions of elements are related to the arrangement of electrons in their atoms and hence to their atomic number.
An element's reactions depend on how easily its atoms gain, lose or share electrons, governed by its electron arrangement. Atoms react to reach a stable outer shell, usually eight electrons (two for the first shell). Metals, such as sodium (2,8,1), lose their single outer electron to form Na⁺ ions. Non-metals, such as chlorine (2,8,7), gain or share electrons to complete an outer shell. For Groups 1 to 7, the group number equals the number of outer electrons. Since atomic number determines the total number of electrons and their arrangement, it dictates the pattern of reactivity. For example, Group 1 elements all have one outer electron and react similarly, with reactivity increasing down the group as the electron is lost more easily.
Your focus
- State that metals are elements that react to form positive ions.
- Describe electron loss from a metal atom using a named example and its ion symbol.
- Classify an element as a metal or non-metal from its position in the periodic table or the charge on its ion.
Show all 18 objectives
- State that elements which do not form positive ions are non-metals.
- Describe how metals form positive ions by losing electrons and how non-metals form negative ions by gaining electrons.
- Classify an element as a metal or non-metal using its ion formation and position in the periodic table.
- Label the metal and non-metal regions on a blank periodic table.
- Classify given elements as metals or non-metals using their position.
- Explain how the stepped line separates metals from non-metals.
- Explain the physical differences between metals and non-metals using structure and bonding.
- Explain how electron loss or gain relates to the chemical properties of metals and non-metals.
- Use examples from Groups 0, 1 and 7 to illustrate differences in reactivity.
- Identify the group and period of an element from its electron configuration.
- Explain how the number of outer electrons determines whether an element behaves as a metal or a non-metal.
- Relate the atomic structure of Group 0 elements to their position and unreactive nature.
- Describe how the outer-shell electron arrangement of an element determines whether it gains, loses or shares electrons.
- Explain how atomic number determines electron arrangement and therefore the position of an element in the periodic table.
- Relate the reactions of elements in the same group to their similar outer-shell electron arrangements.
Metals and non-metals exam tips
Marking Points
- Metals are elements that lose electrons from their outer shell when they react, producing positively charged ions.
- A positive ion forms because the number of protons in the nucleus is greater than the number of electrons remaining.
- Sodium loses one electron to form Na⁺; magnesium loses two electrons to form Mg²⁺; aluminium loses three to form Al³⁺.
- Non-metals are located towards the upper right of the periodic table and typically gain electrons to form negative ions such as O²⁻ and Cl⁻.
- The rule is used to classify an element from its reaction behaviour or its position in the periodic table, not to predict every possible compound.
- Metals lose electrons from their outer shell to form positive ions, for example Na → Na⁺ + e⁻ and Mg → Mg²⁺ + 2e⁻.
- Non-metals gain electrons to form negative ions, for example O + 2e⁻ → O²⁻ and Cl + e⁻ → Cl⁻.
- Elements that do not form positive ions are classified as non-metals, so the absence of positive-ion formation is the key test in this statement.
- Non-metals can also bond covalently by sharing electrons, so they do not need to form ions at all in many compounds.
- Position in the periodic table supports the classification: metals are mainly on the left and centre, non-metals mainly on the right and top.
- Hydrogen is a non-metal that can form H⁺ in acidic solution, so it is a recognised exception to the simple rule.
- States that most elements in the periodic table are metals, with non-metals forming a minority.
- Identifies metals as occupying the left-hand side and lower portions of the periodic table.
- Identifies non-metals as occupying the right-hand side and upper portions of the periodic table.
- Uses the stepped dividing line to classify an unfamiliar element as metal, non-metal or metalloid.
- Links position to typical behaviour, such as metals forming positive ions and non-metals forming negative ions or sharing electrons.
- Recognises hydrogen as a non-metal positioned at the top left of the periodic table.
- Metals are generally good conductors of electricity because they contain delocalised electrons that are free to move throughout the structure.
- Metals are malleable because layers of positive ions can slide over one another; solid non-metals are often brittle because their atoms are held in fixed covalent lattices or simple molecular structures.
- Metals tend to lose electrons to form positive ions, whereas non-metals tend to gain or share electrons, forming negative ions or covalent bonds.
- Metals usually have high melting points due to strong metallic bonding; many non-metals have lower melting points because they form simple molecular structures with weak intermolecular forces.
- Group 0 noble gases are non-metals with full outer shells, making them unreactive; Group 1 metals are highly reactive because they lose one electron easily.
- For elements in Groups 1 to 7, the group number equals the number of electrons in the outer shell, which controls chemical behaviour.
- Group 0 noble gases have full outer shells (two electrons for helium, eight for the others), which explains their position on the far right and lack of reactivity.
- The period number equals the number of occupied electron shells in the atom.
- Metals are found on the left and in the centre because they have few outer electrons and lose them to form positive ions.
- Non-metals are found on the right and towards the top because they have more outer electrons and gain or share electrons to achieve a full shell.
- Reactions of elements are determined by the number and arrangement of electrons, especially the outer-shell electrons.
- Atoms react to achieve a stable outer electron arrangement, typically eight electrons (two for the first shell).
- Metals lose outer electrons to form positive ions, for example Na → Na⁺ + e⁻.
- Non-metals gain or share electrons to complete their outer shell, for example chlorine gaining one electron to form Cl⁻.
- For Groups 1 to 7, the group number equals the number of outer electrons, so elements in the same group react similarly.
- Atomic number gives the number of protons and hence the number of electrons, so it determines the electron arrangement and the reactivity pattern.
Examiner Tips
- 💡When asked to classify an element, quote both its position in the periodic table and the charge on the ion it forms.
- 💡Show the electron loss explicitly, for example sodium 2,8,1 becomes Na⁺ with the configuration 2,8.
- 💡Use the words positive ion and electron loss together so the link between structure and charge is clear.
- 💡Link each classification to electron transfer: metals lose electrons to form positive ions, non-metals gain or share electrons.
- 💡When explaining an example, write the ion with a superscript charge, such as Mg²⁺ or O²⁻, and state the number of electrons gained or lost.
- 💡If a question asks you to identify a non-metal, look for an element that forms negative ions or covalent bonds rather than positive ions.
- 💡Use the periodic table position as supporting evidence, but base your final answer on ion formation or electron transfer.
- 💡Sketch a simple periodic table outline and shade the metal and non-metal regions to anchor the pattern visually.
- 💡When asked to classify an element, quote its position relative to the stepped line before naming it as metal or non-metal.
- 💡Use the terms left, right, top and bottom precisely, since vague answers such as 'one side' lose clarity.
- 💡Link each physical property to its structural explanation, for example 'metals conduct electricity because delocalised electrons can move'.
- 💡When comparing, give a paired statement for a metal and a non-metal rather than describing only one side.
- 💡When explaining reactivity, connect the number of outer electrons to the ease of losing or gaining electrons to achieve a full shell.
- 💡Remember to explicitly mention helium's two outer electrons if discussing Group 0, as it is an exception to the octet rule.
- 💡Link each statement to a specific group, such as Group 1 metals losing one electron or Group 7 non-metals gaining one electron.
- 💡When explaining a group trend, do not compare outer-shell electron numbers as they are identical; instead, compare the number of shells, distance from the nucleus, shielding, and ease of electron loss or gain.
Common Mistakes
- Thinking that all metals form positive ions in every reaction; correction: the statement describes the characteristic behaviour of metals when they do form ions, and some metal compounds are covalent.
- Writing ion charges without superscript signs, for example using a plain plus sign after the symbol; correction: use Unicode superscript notation such as Na⁺ and Mg²⁺.
- Confusing the number of electrons lost with the group number for every element; correction: for main-group metals the group number usually equals the number of outer electrons lost, but transition metals can form more than one ion, such as Fe²⁺ and Fe³⁺.
- Assuming every non-metal forms negative ions in all compounds; correction: many non-metals, such as carbon and hydrogen in covalent compounds, share electrons instead of forming ions.
- Treating hydrogen as a metal because it can form a positive ion; correction: hydrogen is a non-metal, and its formation of H⁺ in acids is a specific exception to the general rule.
- Confusing the charge number with the number of electrons gained or lost; correction: O²⁻ shows a 2⁻ charge because oxygen gains two electrons, not because it gains one.
- Thinking that all elements on the right of the periodic table are non-metals without checking; correction: use electron loss or gain and ion formation as the deciding evidence.
- Assuming every element on the left is a metal and every element on the right is a non-metal without checking the stepped line; correct by noting metalloids such as silicon sit on the boundary.
- Placing hydrogen with the metals because it is at the top left; correct by recalling hydrogen is a non-metal with one electron.
- Thinking the periodic table is mostly non-metals; correct by stating that the majority of elements are metals.
- Saying all non-metals are gases: correction — many non-metals are solids at room temperature, such as sulfur and carbon, and bromine is a liquid.
- Stating that no non-metals conduct electricity: correction — graphite is a form of carbon (a non-metal) that conducts electricity due to delocalised electrons.
- Claiming elemental non-metals have ionic bonds: correction — elemental non-metals are bonded covalently or exist as single atoms; ionic bonds only form between metals and non-metals.
- Stating that the group number always equals the number of outer electrons: correction — this is true for Groups 1 to 7, but Group 0 elements have full outer shells (8 electrons, except helium which has 2).
- Confusing group number with period number: correction — group number relates to outer electrons, while period number gives the number of occupied shells.
- Thinking atomic number equals the number of outer electrons: correction — atomic number is the number of protons, which equals the total electrons in a neutral atom, not just the outer ones.
- Saying reactivity depends only on the number of shells rather than the outer-shell arrangement; correct by linking reactivity to outer electrons and the drive for a stable shell.
- Writing that atoms 'want' or 'need' eight electrons without explaining electron transfer or sharing; correct by describing loss, gain or sharing of electrons to reach a stable arrangement.
- Confusing atomic number with mass number when explaining electron arrangement; correct by stating atomic number equals the number of protons and electrons in a neutral atom.