Why Nitrogen Holds Its Electrons Tighter Than Oxygen

Sodium explodes in water, chlorine is a poisonous gas, and argon does almost nothing — yet they sit side by side in one period. The differences come from two competing effects: the growing pull of the nucleus and the shielding of outer electrons by inner ones.

Once you understand those two effects, size, ionisation enthalpy, electronegativity and reactivity all follow.

This part covers atomic and ionic radii, ionisation and electron gain enthalpy, electronegativity, and trends in valence, metallic character and reactivity.

Atomic radius decreases across a period because nuclear charge rises while electrons enter the same shell, and increases down a group because new shells are added; a cation is smaller than its atom, an anion larger, and among isoelectronic species the one with more protons is smaller.

Across period 2, lithium is the largest atom and fluorine among the smallest; down group 1, lithium, sodium, potassium grow steadily.

**Worked example — isoelectronic ions with $10$ electrons. Compare the ratio of protons to electrons:

- O${}^{2-}$**: $\frac{8}{10}=0.8$
- **F${}^{-}$**: $\frac{9}{10}=0.9$
- **Na${}^{+}$**: $\frac{11}{10}=1.1$
- **Mg${}^{2+}$**: $\frac{12}{10}=1.2$

More protons per electron means a stronger pull, so the size order is

$${\text{O}}^{2-}>{\text{F}}^{-}>{\text{Na}}^{+}>{\text{Mg}}^{2+}$$

An everyday example. Pulling the drawstring of a cloth bag harder squeezes the same contents into a smaller bundle — like a larger nuclear charge acting on the same number of electrons.

The substance. Noble gas radii are van der Waals radii, so they look larger than neighbouring atoms and should not be compared directly.

Ionisation enthalpy — the energy to remove the outermost electron from a gaseous atom — rises across a period and falls down a group as nuclear pull and shielding change; electron gain enthalpy generally becomes more negative across a period; and stable half-filled or filled subshells cause the irregularities.

Irregularities in first ionisation enthalpy (kJ/mol):

- **Be ($899$) > B ($801$) — boron's electron is in a 2p orbital, higher in energy and shielded by the 2s pair
- N ($1402$) > O ($1314$)** — nitrogen's half-filled $2p^3$ is extra stable, while oxygen's paired 2p electrons repel each other

Successive ionisation. Sodium needs $496$ kJ/mol to lose its first electron but $4562$ kJ/mol for its second — the huge jump shows the second electron comes from a stable noble gas core.

Electron gain enthalpy:

- Halogens have the most negative values; **chlorine ($-349$ kJ/mol) is more negative than fluorine ($-328$ kJ/mol) because fluorine's small 2p shell is crowded, so an added electron is repelled
- Noble gases have positive values — an extra electron must enter a new shell

An everyday example. Sodium lamps light up easily because sodium atoms give up and regain their loosely held outer electron with little energy.

The substance. The second ionisation enthalpy is always larger than the first**, since the electron is removed from a positive ion.

Electronegativity is an atom's tendency to attract shared electrons in a bond; a large difference between two atoms gives an ionic bond, a moderate difference a polar covalent bond and a small difference a nearly non-polar bond, and oxides of low-electronegativity metals are basic while those of high-electronegativity non-metals are acidic.

Scales. On the Pauling scale, fluorine has the highest value, $4.0$. The Mulliken scale takes the average of ionisation enthalpy and electron affinity. Electronegativity rises across a period and falls down a group.

Worked example — bond type from differences.

- NaCl: $3.0-0.9=2.1$ — ionic
- HCl: $3.0-2.1=0.9$ — polar covalent
- C–H: $2.5-2.1=0.4$ — nearly non-polar

Nature of oxides.

- Basic: Na${}_2$O + H${}_2$O $\to$ 2NaOH
- Acidic: SO${}_3$ + H${}_2$O $\to$ H${}_2$SO${}_4$
- Amphoteric: Al${}_2$O${}_3$ reacts with both acids and bases
- Neutral: CO, NO, N${}_2$O

Across period 3, oxides change from strongly basic Na${}_2$O to strongly acidic Cl${}_2$O${}_7$.

An everyday example. A tug-of-war between two unequal teams drags the rope towards the stronger side, just as the more electronegative atom pulls the shared electrons.

The substance. Electronegativity is a relative scale, not a directly measured quantity — values differ slightly from one scale to another.

Valence of representative elements equals the number of valence electrons or eight minus that number; metallic character falls across a period and rises down a group; and second-period elements differ from their group because of their small size, high electronegativity and lack of d orbitals, while diagonal pairs such as Li–Mg resemble each other.

Worked example — period 3 formulas.

- Hydrides: NaH, MgH${}_2$, AlH${}_3$, SiH${}_4$, PH${}_3$, H${}_2$S, HCl
- Oxides: Na${}_2$O, MgO, Al${}_2$O${}_3$, SiO${}_2$, P${}_4$O${}_{10}$, SO${}_3$, Cl${}_2$O${}_7$

Reactivity. Alkali metals grow more reactive down the group as they lose electrons more easily; halogens grow less reactive down the group as they gain electrons less readily.

Anomalous second period. Boron can hold at most four bonds, as in $[{\text{BF}}_4{]}^{-}$, because it has no d orbitals, while aluminium forms $[{\text{AlF}}_6{]}^{3-}$.

Diagonal relationship. Li–Mg, Be–Al and B–Si pairs show similar properties because moving right raises electronegativity and moving down lowers it, so the effects balance. Lithium and magnesium both form nitrides when heated in air, and their carbonates both decompose on heating.

An everyday example. Two cousins can look more alike than two siblings — much like lithium resembling magnesium more than sodium in some ways.

The substance. The first element of each group is the least typical member, because its size and electronegativity are extreme.