Periodic Table Trends
Discover the remarkable patterns in atomic radius, electronegativity and ionisation energy as you move across periods and down groups in the periodic table.
What Are Periodic Trends?
The periodic table is arranged so that elements with similar properties appear in the same group (vertical column). As you move across a period (left to right) or down a group, atomic properties change in predictable ways. Understanding these trends allows chemists to predict how elements will behave.
The Two Key Factors Driving Trends
Nuclear charge (atomic number)
As atomic number increases across a period, there are more protons attracting electrons more strongly towards the nucleus.
Electron shielding
Inner electron shells shield outer electrons from the full pull of the nucleus. More shells = more shielding = outer electrons held less tightly.
Trend 1: Atomic Radius
Atomic radius is the distance from the nucleus to the outermost electron shell. It determines the physical size of an atom.
Across a Period (left → right)
Atomic radius decreases.
Across a period, the number of protons increases (greater nuclear charge) but electrons are added to the same shell. The increased pull on the electrons draws the electron cloud closer to the nucleus, making the atom smaller.
Down a Group (top → bottom)
Atomic radius increases.
Down a group, electrons are added to new (higher energy) shells that are further from the nucleus. Even though nuclear charge also increases, the shielding effect of inner shells means outer electrons are held less tightly and further out.
Relative Atomic Sizes Across Period 3
Trend 2: Electronegativity
Electronegativity measures an atom’s ability to attract shared electrons in a covalent bond towards itself. Fluorine (F) is the most electronegative element (Pauling scale value 4.0).
Across a Period (left → right)
Electronegativity increases. More protons = stronger nuclear charge = greater attraction for bonding electrons. Non-metals on the right are highly electronegative.
Down a Group (top → bottom)
Electronegativity decreases. More electron shells = greater shielding = outer electrons further from nucleus = weaker attraction for bonding electrons.
| Element | Li | Be | B | C | N | O | F |
|---|---|---|---|---|---|---|---|
| Electronegativity | 1.0 | 1.5 | 2.0 | 2.5 | 3.0 | 3.5 | 4.0 |
Electronegativity increases left to right across Period 2 (Pauling scale).
Trend 3: Ionisation Energy
First ionisation energy is the energy required to remove the first electron from a gaseous atom:
X(g) → X+(g) + e−
Measured in kilojoules per mole (kJ/mol)
Across a Period (left → right)
Ionisation energy generally increases. Greater nuclear charge holds electrons more tightly, so more energy is needed to remove one. Noble gases have the highest ionisation energies in their period.
Down a Group (top → bottom)
Ionisation energy decreases. Outer electrons are further from the nucleus and shielded by more inner shells, so they require less energy to remove. This is why Group 1 metals (like caesium) are so reactive — their outer electron is very easily lost.
Key Vocabulary
| Term | Definition |
|---|---|
| Atomic radius | The size of an atom, measured as half the distance between the nuclei of two bonded identical atoms. |
| Electronegativity | A measure of an atom’s ability to attract bonding electrons in a covalent bond; highest for fluorine (4.0 on Pauling scale). |
| Ionisation energy | The energy required to completely remove the outermost electron from a gaseous atom; measured in kJ/mol. |
| Shielding effect | The reduction in the attractive force of the nucleus on outer electrons, caused by the presence of inner electron shells between them. |
Worked Examples
Predicting which element has a larger atomic radius: Na or Cl.
Given: Na (sodium) is in Group 1, Period 3. Cl (chlorine) is in Group 17, Period 3. Both are in Period 3.
Step 1: Both atoms have their outermost electrons in shell 3 (same period).
Step 2: Na has 11 protons; Cl has 17 protons. Moving across Period 3, nuclear charge increases.
Step 3: Greater nuclear charge pulls electrons closer to the nucleus.
Answer: Na has a larger atomic radius than Cl. As you move left to right across a period, atomic radius decreases due to increasing nuclear charge.
Explaining why fluorine has the highest electronegativity.
Given: Fluorine (F) is in Group 17, Period 2. It has atomic number 9.
Reasoning: F has a very high nuclear charge (9 protons) and is a small atom (Period 2, so only 2 electron shells and minimal shielding). This means F’s nucleus exerts a very strong pull on the shared electrons in any bond it forms.
Answer: Fluorine has the highest electronegativity (4.0) of all elements because its high nuclear charge and minimal shielding combine to give it the strongest attraction for bonding electrons.
Ranking elements by ionisation energy.
Question: Rank Li, Na, and K in order of increasing first ionisation energy (lowest to highest).
Given: All are in Group 1 (alkali metals). Li is in Period 2; Na in Period 3; K in Period 4.
Step 1: Down a group, ionisation energy decreases (outer electron further from nucleus, more shielding).
Step 2: K has the most electron shells (4) → easiest to remove electron → lowest IE. Li has fewest shells → highest IE in this group.
Answer: K < Na < Li (increasing ionisation energy). Actual values: K = 419 kJ/mol, Na = 496 kJ/mol, Li = 520 kJ/mol.
Knowledge Check
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Key Concepts Summary
- •Atomic radius: decreases across a period (increasing nuclear charge); increases down a group (more electron shells).
- •Electronegativity: increases across a period; decreases down a group. Fluorine has the highest value (4.0).
- •Ionisation energy: generally increases across a period; decreases down a group (outer electrons further from nucleus and more shielded).
- •These trends are caused by two competing factors: nuclear charge (increases with atomic number) and electron shielding (increases with number of inner shells).
- •Elements in the same group have similar chemical properties because they have the same number of valence electrons.