Chemistry · Lesson 3

Periodic table trends

JAMBWAECNECO

What the syllabus asks

JAMB: periodic table and periodicity, trends in groups and periods, reactivity, properties of first 20 elements, transition metals

WAEC: periodic law, historical development (Döbereiner, Newlands, Mendeleev, Moseley), electronic configuration and position in the table, s-, p- and d-blocks, periodicity of atomic radius, ionisation energy, electron affinity, electronegativity, metallic and non-metallic character, group trends (alkali metals, alkaline earth metals, halogens, noble gases), period 3 trends, transition elements

NECO: periodic law, historical development (Döbereiner, Newlands, Mendeleev, Moseley), electronic configuration and position in the table, s-, p- and d-blocks, periodicity of atomic radius, ionisation energy, electron affinity, electronegativity, metallic and non-metallic character, group trends (alkali metals, alkaline earth metals, halogens, noble gases), period 3 trends, transition elements

What you will learn

Outline the historical development of the periodic table.
State the periodic law and relate electron configuration to period, group and block.
Explain and predict trends in atomic radius, ionisation energy, electron affinity and electronegativity.
Describe metallic and non-metallic character across period 3, including the nature of oxides and chlorides.
Describe properties and reactions of groups 1, 2, 7 and 0.
State the characteristic properties of transition elements.
Use the table to predict properties of unfamiliar elements.

Key points

Periodic law: properties repeat periodically with atomic number (Moseley); Mendeleev left gaps and predicted new elements.
7 periods, 18 groups; s-, p-, d- and f-blocks; metals left of the staircase, non-metals right.
Period = number of shells; main group = number of valence electrons; 4s fills before 3d.
Trends are explained by nuclear charge, number of shells and shielding.
Across a period: radius ↓, ionisation energy ↑ (dips at Al and S), electronegativity ↑, metallic character ↓.
Down a group: radius ↑, ionisation energy ↓, electronegativity ↓; positive ions smaller and negative ions larger than atoms.
Period 3 oxides: basic (Na, Mg) → amphoteric (Al) → acidic (Si to Cl).
Group 1 and 2 metals: reactivity increases down the group; group 2 hydroxides more soluble, sulphates less soluble down.
Halogens: diatomic, darker and higher boiling down the group; reactivity decreases; a higher halogen displaces a lower one.
Noble gases: full shells, monatomic, unreactive; transition metals: variable valency, coloured ions, catalysts, complexes.

Start here: Why we need a periodic table

Imagine a market in Onitsha where traders sold everything in one heap: shoes beside tomatoes beside phone chargers. Finding anything would be a nightmare. By the middle of the 1800s chemists faced the same mess: over 60 elements had been discovered, each with its own properties, and no one knew how they were related. They needed to arrange them so that similar elements sat together.

Early attempts

Döbereiner's triads (1829): J. W. Döbereiner noticed groups of three similar elements, such as lithium, sodium and potassium, or chlorine, bromine and iodine, in which the relative atomic mass of the middle element was roughly the average of the other two. For Cl (35.5) and I (127), the average is 81.25, close to Br (80). But only a few triads could be found.
Newlands' law of octaves (1864): John Newlands arranged the known elements by increasing atomic mass and saw that every eighth element resembled the first, like notes in a musical scale. It worked for light elements but broke down after calcium, and other scientists laughed at the musical comparison.
Mendeleev's periodic table (1869): Dmitri Mendeleev arranged elements by increasing atomic mass but put chemical properties first. Where an element did not fit, he left a gap and boldly predicted that an undiscovered element would fill it. He even swapped some elements (such as tellurium and iodine) out of mass order to keep families together.

Worked example 3.1

Mendeleev predicted "eka-silicon" with relative atomic mass about 72 and density about 5.5 g/cm³. Germanium, found in 1886, has A_r 72.6 and density 5.35 g/cm³. Find the percentage error in his density prediction.

Error = 5.5 − 5.35 = 0.15 g/cm³.
Percentage error = 0.15/5.35 × 100% ≈ 2.8%. This accuracy is why chemists came to trust his table.

≈ 2.8%

The modern periodic law

Mendeleev's few out-of-order pairs were explained in 1913 when Henry Moseley, using X-ray spectra, showed that each element has a unique atomic number (number of protons). Arranged by atomic number, every element falls into the right place with no swapping.

Modern periodic law: the physical and chemical properties of the elements are a periodic function of their atomic numbers.

"Periodic" means repeating at regular intervals, like the days of the week. Properties such as atomic size and reactivity rise and fall in a repeating pattern as Z increases.

⚠️ Common mistake: stating that the modern table is arranged by atomic mass. It is arranged by atomic number.

Try it yourself. Li = 7, Na = 23, K = 39. Check Döbereiner's rule.

Show solution

(7 + 39)/2 = 23 = Na ✓.

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Practice questions

1. Elements in the same group have the same number of

valence electrons
shells
neutrons
protons

Show answer

A: valence electrons
Group = number of outer (valence) electrons for main groups.

2. Which is the most electronegative element?

Chlorine
Oxygen
Caesium
Fluorine

Show answer

D: Fluorine
Fluorine has the highest electronegativity (4.0).

3. Across a period from left to right, atomic radius generally

stays the same
doubles
decreases
increases

Show answer

C: decreases
Nuclear charge increases and pulls electrons closer.

4. Which of these is a noble gas?

Hydrogen
Argon
Nitrogen
Chlorine

Show answer

B: Argon
Group 0/18: He, Ne, Ar, Kr, Xe, Rn.

5. Potassium reacts more vigorously with water than sodium because potassium

loses its outer electron more easily
has fewer electrons
is a non-metal
has a smaller atom

Show answer

A: loses its outer electron more easily
K has a larger atom; its outer electron is further from the nucleus, so it is lost more easily.

Key words

Periodic law: The properties of the elements are a periodic function of their atomic numbers.
Period: A horizontal row of the periodic table.
Group: A vertical column of elements with the same number of valence electrons.
Valence electrons: Electrons in the outermost shell of an atom.
Atomic radius: Half the distance between the nuclei of two bonded atoms of the same element.
First ionisation energy: Energy needed to remove one mole of electrons from one mole of gaseous atoms to form gaseous 1+ ions.
Electron affinity: Energy released when one mole of gaseous atoms each gain an electron to form 1− ions.
Electronegativity: The tendency of an atom in a bond to attract the shared pair of electrons to itself.
Effective nuclear charge: The net positive charge felt by the outer electrons after shielding by inner electrons.
Shielding (screening): Reduction of nuclear attraction on outer electrons by inner electron shells.
Alkali metals: Group 1 elements: Li, Na, K, Rb, Cs, Fr.
Alkaline earth metals: Group 2 elements: Be, Mg, Ca, Sr, Ba, Ra.

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