(d) The Periodic Table
- Syllabus
- 2024
- Topic
- —
- Level
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Learning objectives
1.18Periodic Table arrangementUnderstand how elements are arranged in the Periodic Table:• in order of atomic number• in groups and periods.1.19Electronic configurationsUnderstand how to deduce the electronic configurations of the first 20 elements from their positions in the Periodic Table1.20Metals and non-metalsUnderstand how to use electrical conductivity and the acid-base character of oxides to classify elements as metals or non-metals1.21Element position and typeIdentify an element as a metal or a non-metal according to its position in the Periodic Table1.22Group position and electron configurationUnderstand how the electronic configuration of a main group element is related to its position in the Periodic Table1.23Group propertiesUnderstand why elements in the same group of the Periodic Table have similar chemical properties1.24Noble gasesUnderstand why the noble gases (Group 0) do not readily react
Navigate the Periodic Table
Elements are arranged in increasing atomic number, so each step to the next element adds one proton to the nucleus. Atomic number—not mass number or relative atomic mass—controls the order.
| Feature | Direction | Meaning |
|---|---|---|
| Period | Horizontal row | Elements occupy the same number of electron shells. |
| Group | Vertical column | Main-group elements have the same number of outer-shell electrons and related chemistry. |
An element is located by the intersection of its group and period. Magnesium is in Group 2, Period 3; boron is in Group 3, Period 2. Use the element's atomic number to identify it before reading its coordinates.
Moving across a period increases atomic number by one, but relative atomic mass is not directly proportional to atomic number. Do not read groups as rows or periods as columns.
Deduce electron configurations for the first 20 elements
For a neutral atom, the number of electrons equals the atomic number. For the first 20 elements, place electrons into shells in the order 2 in the first shell, then 8 in the second, 8 in the third, and then the fourth shell.
| Element | Atomic number | Electron configuration | Check from position |
|---|---|---|---|
| Boron | 5 | 2,3 | Period 2; Group 3 |
| Sodium | 11 | 2,8,1 | Period 3; Group 1 |
| Silicon | 14 | 2,8,4 | Period 3; Group 4 |
| Calcium | 20 | 2,8,8,2 | Period 4; Group 2 |
Start with the atomic number, fill each inner shell before the next, and confirm that the shell totals add back to the atomic number. The number of occupied shells should match the period.
Do not put more than 2 electrons in the first shell or more than 8 in the second shell. This 2,8,8,2 model is the required deduction pattern for the first 20 elements; it is not a universal filling rule for all later elements.
Classify metals using conductivity and their oxides
Classify an element by combining physical evidence from electrical conductivity with chemical evidence from the acid–base character of its oxide.
| Evidence | Typical metal | Typical non-metal |
|---|---|---|
| Electrical conductivity of the element | Good conductor | Poor conductor |
| Character of the oxide | Basic; reacts with or neutralises acids | Acidic; reacts with bases |
Copper conducts electricity and copper oxide is basic, supporting classification as a metal. Sulfur and chlorine form acidic oxides, supporting classification as non-metals. Silicon dioxide is acidic because it reacts with basic calcium oxide.
Use both requested tests when evidence is available. 'Basic oxide' is not the same as 'alkali': a base need not dissolve in water. Conductivity is evidence about the element, while acid–base character is evidence about its oxide.
Identify metals and non-metals by position
The Periodic Table has a broad metal region on the left and centre and a non-metal region on the upper right. The stepped boundary between these regions is the quickest positional guide.
| Position | Classification examples |
|---|---|
| Left and centre of the table | Metals such as sodium, magnesium, aluminium and potassium |
| Upper-right region, including Group 7 and Group 0 | Non-metals such as chlorine, iodine, oxygen and argon |
Locate the element first, then compare its position with the stepped divide. Period 3 begins with metallic sodium, magnesium and aluminium and continues into non-metals such as sulfur and chlorine.
Hydrogen is a non-metal even though it is placed above Group 1 on the left. Position provides the classification required here; conductivity and oxide tests belong to the separate evidence-based classification objective.
Link electron configuration to group and period
For a main-group element, its electron configuration encodes its Periodic Table position: occupied shells give the period, and outer-shell electrons give the group for Groups 1–7.
| Configuration feature | Position rule | Example |
|---|---|---|
| Number of occupied shells | Period number | 2,8,2 has three shells → Period 3 |
| Number of outer-shell electrons | Group number for Groups 1–7 | 2,8,2 has two outer electrons → Group 2 |
| Full outer shell | Group 0 | 2,8,8 → Group 0, Period 3 |
The relationship works in reverse for the first 20 elements. Group 5, Period 3 means three occupied shells and five outer electrons, so the configuration is 2,8,5.
Do not use total electron number as the group number. Helium is in Group 0 with a full first shell of 2 electrons; the other first-20 noble gases have 8 outer electrons.
Explain why a group shares chemical properties
Elements in the same group have similar chemical properties because their atoms have the same number of electrons in the outer shell.
Chemical reactions involve outer-shell electrons. Atoms in one group therefore tend to lose, gain or share the same number of electrons and form similar types of ions or bonds. Group 1 atoms each have one outer electron and tend to lose that one electron in reactions.
| Same group | What stays the same | What changes down the group |
|---|---|---|
| Example: chlorine, bromine, iodine | Seven outer-shell electrons; similar reaction pattern | More occupied shells and a different period |
Being in the same group explains similarity, not identical reaction speed or every physical property. The causal reason is the same outer-shell electron count, not merely that the elements appear in one column.
Explain the low reactivity of noble gases
Noble gases are in Group 0 and do not readily react because their atoms already have a full outer electron shell.
A full outer shell is a stable arrangement, so a noble-gas atom has little tendency to gain, lose or share electrons. Helium has a full first shell of 2; neon, argon and the later noble gases have a full outer shell of 8.
| Situation | Why a noble gas is suitable |
|---|---|
| Argon atmosphere around reactive magnesium or titanium | Argon does not react with or oxidise the hot materials. |
| Helium in airships | Helium is unreactive and non-flammable, unlike hydrogen. |
Say 'do not readily react' rather than 'can never react'. Their low reactivity follows from electron configuration; it is not simply because they are gases.