Periodic table & atomic properties

What you will learn

read the periodic table: groups (columns), periods (rows), metals, non-metals, metalloids,
write electron configurations for the first $20$ elements,
predict group number from the number of valence (outer-shell) electrons,
describe trends in atomic size and reactivity down a group and across a period,
distinguish ionic bonding (metal + non-metal) from covalent bonding (non-metal + non-metal).

Worked example 0 Real-world example: why fluorine is used to harden tooth enamel

Fluoride ion ( $F^-$ ) is added to toothpaste. Fluorine is in Group 17 (halogens). Why does fluorine form an ion with a $-1$ charge, and why is it so reactive?

Fluorine atoms have $9$ protons and $9$ electrons. Configuration: $2, 7$ . (Two in the inner shell, seven in the outer shell.)
Its outer shell is one electron short of a full $8$ (the stable “noble gas” configuration).
Fluorine gains $1$ electron easily, becoming $F^-$ with configuration $2, 8$ .
The strong pull on that single extra electron makes fluorine the most reactive non-metal — it bonds readily with metals (calcium in enamel) to form hard, protective compounds.

Key idea: reactivity is driven by how badly an atom “wants” a full outer shell of $8$ electrons (the octet rule).

1. Structure of the periodic table

Simplified layout of the periodic table. Group number equals the number of valence electrons (for s- and p-block elements).

Groups (columns, numbered 1-18): elements in the same group share the same number of valence electrons, so they behave similarly.
Periods (rows): filling a new period means starting a new electron shell.
Metals are on the left and centre. They lose electrons and form cations (positive ions).
Non-metals are on the right. They gain electrons and form anions (negative ions).
Metalloids (B, Si, Ge, As, Sb, Te) sit on the “staircase” between the two. They show mixed properties.
Noble gases (Group 18) have full outer shells — they are unreactive.

2. Electron configuration (first 20 elements)

Electrons occupy shells around the nucleus. Shells fill in this order: $2, 8, 8, 2$ for the first $20$ elements. Write configurations by filling the innermost shell first.

Element	Z	Configuration	Group
Hydrogen (H)	1	$1$	1
Helium (He)	2	$2$	18
Carbon (C)	6	$2, 4$	14
Nitrogen (N)	7	$2, 5$	15
Oxygen (O)	8	$2, 6$	16
Fluorine (F)	9	$2, 7$	17
Neon (Ne)	10	$2, 8$	18
Sodium (Na)	11	$2, 8, 1$	1
Magnesium (Mg)	12	$2, 8, 2$	2
Sulfur (S)	16	$2, 8, 6$	16
Chlorine (Cl)	17	$2, 8, 7$	17
Argon (Ar)	18	$2, 8, 8$	18
Potassium (K)	19	$2, 8, 8, 1$	1
Calcium (Ca)	20	$2, 8, 8, 2$	2

Worked example 1 Predicting group from configuration

An atom has electron configuration $2, 8, 5$ . What is its atomic number, which group is it in, and name the element.

Total electrons $= 2 + 8 + 5 = 15$ .
Atomic number $Z = 15$ .
Valence (outer shell) electrons $= 5$ , so it is in Group 15.
Element: phosphorus (P).

3. Reactivity trends

Down a group (e.g. Group 1: Li, Na, K, Rb, Cs):

Atomic radius increases (more shells).
Outer electron is held less tightly.
For metals (Groups 1, 2): reactivity increases down the group.
For non-metals (Group 17, halogens): reactivity decreases down the group.

Across a period (left to right):

Atomic radius decreases (more protons pulling the same shell inward).
Metallic character decreases; non-metallic character increases.
Group 1 (very reactive metal) $\to$ Group 18 (unreactive noble gas).

Worked example 2 Which alkali metal is most reactive?

Rank lithium, sodium, potassium and caesium by reactivity in water. Justify.

All are Group 1 metals, each with $1$ valence electron.
Going down the group, the outer electron is in a shell further from the nucleus.
The further away the electron, the more easily it is lost in a reaction.
Reactivity order (least to most): Li $<$ Na $<$ K $<$ Cs.

Caesium reacts explosively even with cold water, whereas lithium merely fizzes.

4. Ionic and covalent bonding

When a metal meets a non-metal, the metal loses electrons and the non-metal gains them. The opposite charges attract, forming an ionic bond (e.g. sodium chloride, NaCl).

When two non-metals meet, neither can fully take the other’s electrons. They share pairs of electrons — a covalent bond (e.g. water H $_2$ O, methane CH $_4$ ).

Worked example 3 Why does table salt exist as NaCl?

Write the electron configurations of sodium and chlorine. Explain how they bond.

Na: $2, 8, 1$ . Cl: $2, 8, 7$ .
Na loses $1$ electron to achieve $2, 8$ (the stable neon configuration), becoming Na $^+$ .
Cl gains $1$ electron to achieve $2, 8, 8$ (the stable argon configuration), becoming Cl $^-$ .
Na $^+$ and Cl $^-$ attract electrostatically, forming an ionic compound: NaCl.

Key idea: both atoms end up with full outer shells — the octet rule is satisfied, and the compound is stable.

Worked example 4 Covalent bonding in water

Explain the bonding in H $_2$ O.

Oxygen (Z = 8, configuration $2, 6$ ) needs $2$ more electrons for a full outer shell.
Each hydrogen (Z = 1, configuration $1$ ) needs $1$ more electron.
Oxygen shares a pair of electrons with each of two hydrogen atoms.
Each shared pair is a single covalent bond. All three atoms now have filled outer shells.

Key idea: covalent bonding is sharing, not transferring, electrons. This happens when both atoms need electrons.

Practice: Year 10

Fluency

Table literacy

What is the difference between a group and a period?
Name the elements in Group 1 (first five).
Name the elements in Group 18 (first four).
Where are metals located on the periodic table? Non-metals?
What is a metalloid? Name two.
Why are noble gases unreactive?

Fluency

Electron configuration

Write the electron configuration of (a) carbon, (b) oxygen, (c) neon, (d) sodium.
Which group is an element with configuration $2, 8, 6$ in? Name it.
Which group is an element with configuration $2, 8, 8, 1$ in? Name it.
How many valence electrons does nitrogen (Z = 7) have?
Draw the electron-shell diagram for magnesium (Z = 12).

Fluency

Ions and bonding

What charge does a Group 1 metal ion carry? A Group 2 metal?
What charge does a Group 17 ion carry? A Group 16 non-metal?
Is sodium chloride (NaCl) ionic or covalent? Justify.
Is methane (CH $_4$ ) ionic or covalent? Justify.
Write the formula for the ionic compound formed between magnesium and oxygen.

Reasoning

Trends and explain

Explain why reactivity increases down Group 1 but decreases down Group 17.
Predict which is more reactive: sodium or magnesium. Justify using electron configurations.
Why do elements in the same group have similar chemical properties?
Explain why atoms form ions — what is the “goal” of the process?
A student claims “helium should be in Group 2 because it has 2 electrons.” Explain why helium is in Group 18 instead.

Problem solving

Applications

Lithium-ion batteries rely on Li $^+$ ions moving between electrodes. Why is lithium used rather than sodium or potassium? (Consider mass and reactivity.)
Argon is used inside incandescent light bulbs. Why is it safer than oxygen or nitrogen?
Fluorine forms HF, chlorine forms HCl, bromine forms HBr. Which of these three reactions is most vigorous? Predict using the group trend.
Write the formula and bonding type for (a) potassium chloride, (b) calcium oxide, (c) carbon dioxide, (d) ammonia (NH $_3$ ).

Challenge

Reasoning

Harder reasoning

Mendeleev predicted the existence of “eka-silicon” (later discovered as germanium) with specific properties. What aspects of the periodic table made such predictions possible? Why did the modern table switch from atomic mass to atomic number as the ordering principle?
Transition metals often form ions with more than one possible charge (e.g. Fe $^{2+}$ and Fe $^{3+}$ ). Why does this complication not appear for Group 1 or Group 17 elements?
The radius of Na is about $186$ pm, while Na $^+$ is only about $102$ pm. Explain the dramatic shrinkage. Compare with Cl ( $99$ pm) vs Cl $^-$ ( $181$ pm), where the opposite happens.
Water (H $_2$ O) is a covalent compound, yet it dissolves NaCl (ionic). Explain why, in terms of water molecules’ polarity and the charged ions in the crystal.