The Periodic Table · Grade 10 Chemistry

A Brief History of the Periodic Table

In 1869, Dmitri Mendeleev arranged the 63 known elements by increasing atomic mass and noticed repeating (periodic) patterns in their properties. He left gaps for undiscovered elements and predicted their properties with remarkable accuracy. Henry Moseley (1913) showed that atomic NUMBER — not mass — is the true organising principle, resolving anomalies in Mendeleev's arrangement.

Key insight: The modern periodic table arranges 118 elements in order of INCREASING ATOMIC NUMBER. The repeating pattern of properties is called periodicity.

Structure of the Periodic Table

The table is organised into PERIODS (horizontal rows) and GROUPS (vertical columns).

Periods: There are 7 periods. The period number tells you how many electron shells (energy levels) an atom of that element has.
Groups: There are 18 groups. Elements in the same group have the same number of valence electrons, giving them similar chemical properties.
The table contains 118 confirmed elements.

Group	Name	Examples	Key Property
Group 1	Alkali Metals	Li, Na, K	Very reactive; form 1+ ions; react vigorously with water
Group 2	Alkaline Earth Metals	Be, Mg, Ca	Reactive; form 2+ ions; less reactive than Group 1
Groups 3–12	Transition Metals	Fe, Cu, Zn, Cr	Variable valency; coloured compounds; good conductors
Group 17	Halogens	F, Cl, Br, I	Very reactive non-metals; form 1− ions; diatomic molecules
Group 18	Noble Gases	He, Ne, Ar	Unreactive; full outer shell; exist as single atoms

Metals, Non-metals, and Metalloids

A bold staircase line separates metals (left/centre) from non-metals (right). Metalloids (semi-metals) sit along the staircase.

Category	Location	Properties	Examples
Metals	Left and centre	Lustrous, malleable, ductile, good conductors, high MP/BP	Na, Fe, Cu, Al
Non-metals	Top right	Dull, brittle (solid), poor conductors, low MP/BP	C, N, O, S, Cl
Metalloids	Along staircase	Intermediate properties; semiconductors	B, Si, As, Ge, Te

Periodic Trends Across a Period (left → right)

As you move LEFT to RIGHT across a period, the number of protons increases while electrons are added to the SAME shell. This means the nuclear charge increases but shielding stays roughly constant — so the nuclear pull on electrons increases.

Atomic radius DECREASES → more protons pull electrons closer to nucleus

Ionisation energy INCREASES → electron is harder to remove from a smaller, higher-Z atom

Electronegativity INCREASES → atom attracts shared electrons more strongly

Metallic character DECREASES → left side is metallic, right side non-metallic

Periodic Trends Down a Group (top → bottom)

As you move DOWN a group, each successive element has one more electron shell. The outer electrons are further from the nucleus and are more shielded by inner electrons — so nuclear pull on outer electrons decreases.

Atomic radius INCREASES → more electron shells

Ionisation energy DECREASES → outer electron further away and more shielded

Electronegativity DECREASES → atom attracts bonding electrons less strongly

Metallic character INCREASES → elements become more metallic going down

Ion Formation

The position of an element in the periodic table predicts the charge of the ion it forms:

Metals (left side) LOSE electrons to form positive ions (cations). Group 1 → 1+ charge; Group 2 → 2+ charge; Group 13 (Al) → 3+ charge.
Non-metals (right side) GAIN electrons to form negative ions (anions). Group 17 → 1− charge; Group 16 → 2− charge; Group 15 → 3− charge.
Noble gases (Group 18) do not typically form ions — their outer shell is already full.

Na → Na⁺ + e⁻ Cl + e⁻ → Cl⁻ Mg → Mg²⁺ + 2e⁻ O + 2e⁻ → O²⁻

Electron Configuration and the Periodic Table

The period number = number of electron shells. The group number (for main group elements) = number of valence electrons. For example, sodium (Na, Z=11) is in Period 3 → 3 shells; Group 1 → 1 valence electron. Configuration: 2,8,1.

Links forward: Chemical bonding (Lesson 4) uses valence electrons directly. Electronegativity values from this lesson determine bond type.

IEB Extension — Advanced Periodicity

Successive ionisation energies provide direct evidence for electron shells. For sodium (Na), the first ionisation energy is low (one valence electron, easy to remove). The second ionisation energy is dramatically higher — the electron is now being removed from a complete inner shell. This large jump in IE values proves the shell structure.

Transition metals have irregularities in electron configuration. The d-subshell is slightly lower in energy than 4s, causing:

Chromium (Cr, Z=24): expected [Ar]3d⁴4s² but actually [Ar]3d⁵4s¹ — half-filled d-subshell is extra stable
Copper (Cu, Z=29): expected [Ar]3d⁹4s² but actually [Ar]3d¹⁰4s¹ — fully-filled d-subshell is extra stable

Effective nuclear charge (Zeff) = Z − shielding constant (σ). As you move across a period, Z increases by 1 each step but shielding increases by less than 1 (inner electrons shield more than outer electrons), so Zeff rises — this explains why atomic radius decreases and IE increases across a period.

Interactive Periodic Table — Periods 1–4

Mode

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Element Info

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Legend

Alkali metals Alkaline earth Transition metals Other metals Metalloids Non-metals Noble gases

Question 1

What does the period number of an element tell you?

The period number equals the number of electron shells. Sodium is in Period 3, so it has 3 electron shells.

Question 2

Moving left to right across Period 3, the atomic radius:

Across a period, nuclear charge increases but electrons are added to the same shell — stronger pull on electrons makes the atom smaller.

Question 3

Which group contains elements that most readily form 1− ions?

Group 17 (halogens) have 7 valence electrons — they gain one electron to complete their outer shell, forming 1− ions.

Question 4

Which of the following is a metalloid?

Silicon sits along the staircase line separating metals from non-metals — it's a metalloid (semiconductor) used in computer chips.

Question 5

Down Group 1 (alkali metals), which trend is correct?

Going down a group, each element has one more electron shell. The outer electrons are further from the nucleus, so atomic radius increases.

Question 6

An element is in Period 3, Group 2. What is the charge of the ion it most likely forms?

Group 2 elements have 2 valence electrons, which they lose to form 2+ cations. Magnesium (Mg, Period 3, Group 2) forms Mg²⁺.

Question 7 (Analysis)

Consider three elements: sodium (Na, Period 3, Group 1), magnesium (Mg, Period 3, Group 2), and potassium (K, Period 4, Group 1). Ranked from LARGEST to SMALLEST atomic radius, which order is correct?

This requires combining both trends. K is in Period 4 (4 electron shells) while Na and Mg are both in Period 3 (3 shells) — an extra shell outweighs any group difference, so K is largest. Comparing Na and Mg (same period, same number of shells): Mg has one more proton (Z=12 vs Z=11), giving a stronger nuclear pull on the same outer shell, so Mg is smaller than Na. Overall: K > Na > Mg.

IEB Question 1 IEB

The successive ionisation energies (in kJ·mol⁻¹) of an unknown element are: 738, 1451, 7733, 10 543. Between which removals is there a large jump, and what does this indicate?

The large jump occurs between the 2nd and 3rd ionisation energies (1451 → 7733 kJ·mol⁻¹). This means the first 2 electrons were removed from the outer shell easily, but the 3rd comes from an inner shell — confirming the element has 2 valence electrons. The element is magnesium (Mg, Group 2).

IEB Question 2 IEB

Chromium has the electron configuration [Ar]3d⁵4s¹ rather than the expected [Ar]3d⁴4s². The best explanation is:

A half-filled d-subshell (d⁵, one electron in each of the five d-orbitals) has extra stability due to exchange energy. This causes Cr to 'promote' one 4s electron into the 3d level. The same applies to Cu ([Ar]3d¹⁰4s¹) which prefers a fully-filled d-subshell.

Answer in full sentences. Show working where applicable.

Question 1

Explain why the atomic radius decreases across a period from left to right. Use the terms 'nuclear charge', 'electron shells', and 'shielding' in your answer.

Question 2

An element has the electron configuration 2,8,7. (a) In which period and group is it? (b) What type of ion will it form? (c) Name the element and give its symbol.

Question 3

Compare the properties of metals, non-metals, and metalloids. Include at least three properties for each category.

Question 4

Explain the difference between ionisation energy and electronegativity. How do these trends change across Period 2?

Question 5

Using the concept of electron configuration, explain why elements in the same group have similar chemical properties. Give sodium (Na) and potassium (K) as examples.

Question 6 — Reading First Ionisation Energy Data

The table below gives the first ionisation energy of each element in Period 3:

Element	Na	Mg	Al	Si	P	S	Cl	Ar
1st IE (kJ·mol⁻¹)	496	738	577	786	1012	1000	1251	1521

(a) Describe the overall trend in first ionisation energy from Na to Ar, using values from the table.
(b) Using the concepts of nuclear charge and electron shells, explain why this overall trend occurs.
(c) Two elements in the table break the smooth increasing pattern (their IE is lower than the element before them). Identify these two elements from the data.
(d) Potassium (K) is the first element of Period 4. Would you expect its first ionisation energy to be higher or lower than sodium's (496 kJ·mol⁻¹)? Name the trend (across a period or down a group) that applies here, and explain your reasoning.