Periodic Trends

Elements in the same group have the same valence electron configuration (same number and type of valence electrons). Because chemical reactivity depends primarily on valence electrons, these elements undergo similar types of reactions and form compounds with similar formulas, even though they differ in atomic size and total electron count.

An element in Period 4, Group 15 is in the p-block, so its highest-energy electrons occupy the 4p subshell. The periodic table's column placement within the p-block indicates three p electrons (p³), giving a valence configuration of 4s²4p³. This is arsenic (As).

Noble gases have completely filled s and p subshells in their outermost shell (e.g., Ne is 2s²2p⁶). This completely filled valence shell creates an exceptionally stable electron configuration with no tendency to gain, lose, or share electrons, resulting in extremely low reactivity.

Across a period, each successive element has one additional proton and one additional electron, but the added electron enters the same principal energy level. Electrons in the same shell provide poor shielding for one another, so the increase in nuclear charge is not offset by an equal increase in shielding. As a result, $Z_{\mathrm{eff}}$ increases, and valence electrons are held more tightly.

Phosphorus has a greater $Z_{\mathrm{eff}}$ experienced by its valence electron. Both Na (Z = 11) and P (Z = 15) have 10 core electrons shielding the valence shell, so $Z_{\mathrm{eff}}$ ≈ +1 for Na and $Z_{\mathrm{eff}}$ ≈ +5 for P. Because P has more protons but the same number of core electrons, its valence electrons experience a significantly stronger net attraction to the nucleus.

Moving down a group, the valence electrons occupy a higher principal energy level, meaning they are significantly farther from the nucleus. Additionally, each new shell of core electrons provides substantial shielding, greatly reducing the effective nuclear charge experienced by the outermost electrons. Although the total nuclear charge Z increases, the combined effect of greater distance and increased shielding dominates, resulting in a weaker net electrostatic attraction according to Coulomb's law.

Across a period, the number of protons increases while electrons are added to the same principal energy level. Electrons in the same shell shield each other poorly, so the effective nuclear charge increases from left to right. The increasing $Z_{\mathrm{eff}}$ means valence electrons are held more tightly, requiring more energy to remove, so first ionization energy generally increases.

Nitrogen has an electron configuration of 1s²2s²2p³ with three unpaired p electrons, while oxygen has a configuration of 1s²2s²2p⁴ with one pair of electrons sharing an orbital. The electron-electron repulsion between the two paired electrons in oxygen's 2p orbital destabilizes that electron, making it easier to remove. This additional repulsion more than offsets the slight increase in $Z_{\mathrm{eff}}$ from nitrogen to oxygen, resulting in a lower IE₁ for oxygen.

Calcium has a higher first ionization energy than potassium. Both are in Period 4, but calcium has one additional proton (Z = 20 vs. Z = 19) while both have 18 core electrons. Calcium's valence electrons experience a greater $Z_{\mathrm{eff}}$ (≈ +2 vs. ≈ +1), so more energy is required to remove an electron from calcium than from potassium.

From sodium to chlorine, the number of protons increases from 11 to 17 while all valence electrons are added to the n = 3 shell. The core electron count remains at 10, so $Z_{\mathrm{eff}}$ increases from approximately +1 to +7. This increasing effective nuclear charge pulls the valence electrons closer to the nucleus, progressively contracting the atomic radius across the period.

All five ions have 10 electrons. The species with the most protons exerts the greatest $Z_{\mathrm{eff}}$ on those electrons, pulling them inward and producing the smallest radius. The order from smallest to largest is Mg²⁺ (12 protons) < Na⁺ (11) < F⁻ (9) < O²⁻ (8) < N³⁻ (7). Greater nuclear charge among isoelectronic species results in a smaller ionic radius.

When chlorine gains one electron to form Cl⁻, the number of electrons increases from 17 to 18, but the nuclear charge remains at 17 protons. The additional electron increases electron-electron repulsion in the valence shell without any increase in nuclear attraction. This excess repulsion causes the electron cloud to expand, making Cl⁻ significantly larger than the neutral Cl atom.

Across a period, $Z_{\mathrm{eff}}$ increases because protons are added while core electron shielding remains nearly constant. The incoming electron is therefore attracted more strongly into the valence shell, making the process more exothermic. As a result, electron affinity becomes more negative (more favorable) from left to right across a period.

Fluorine has a higher $Z_{\mathrm{eff}}$ relative to its atomic radius than iodine because it is higher in Group 17 with fewer shielding shells. The incoming electron in fluorine enters a smaller, closer n = 2 shell where it experiences a stronger effective nuclear charge. Therefore, fluorine is predicted to have a more exothermic electron affinity than iodine. (In practice, chlorine has a slightly more exothermic EA than fluorine due to extreme electron-electron repulsion in fluorine's small 2p subshell, but the general group trend predicts fluorine > iodine.)

Noble gases have completely filled valence shells (ns²np⁶). An additional electron would need to enter the next higher principal energy level (n+1), which is much farther from the nucleus and heavily shielded. The weak attraction between the nucleus and this distant incoming electron means the process releases little or no energy and may require energy input, resulting in a near-zero or positive electron affinity.

Across a period, the number of protons increases while electrons are added to the same principal energy level. Core electron shielding remains essentially constant, so $Z_{\mathrm{eff}}$ increases. The higher $Z_{\mathrm{eff}}$ gives each atom a greater ability to attract shared bonding electrons toward itself, causing electronegativity to increase from left to right.

Nitrogen has a higher electronegativity than phosphorus. Both are in Group 15, but nitrogen's valence electrons are in the n = 2 shell while phosphorus's are in the n = 3 shell. Nitrogen has a smaller atomic radius and fewer shielding shells, so its nucleus exerts a stronger electrostatic pull on bonding electrons according to Coulomb's law. The shorter distance between the bonding electrons and nitrogen's nucleus results in a greater attractive force.

When two bonded atoms have equal or nearly equal electronegativities (small ΔEN), they share electrons evenly, forming a nonpolar covalent bond. When the electronegativity difference is moderate, electrons are shared unequally, producing a polar covalent bond with partial charges (δ⁺ and δ⁻). When the electronegativity difference is very large, the more electronegative atom effectively takes the electron from the less electronegative atom, resulting in electron transfer and an ionic bond.

Germanium has a lower first ionization energy than silicon. Both elements are in Group 14, but germanium is one period below silicon. Moving down a group, the valence electrons occupy a higher principal energy level with greater shielding and a larger distance from the nucleus, weakening the nuclear attraction. Therefore, less energy is required to remove a valence electron from germanium.

Aluminum is between magnesium and silicon in Period 3. The general trend predicts an IE₁ between 738 and 786 kJ/mol. However, aluminum's outermost electron is in a 3p subshell (higher energy than Mg's 3s), which makes it easier to remove. This suggests Al's IE₁ is actually slightly lower than Mg's, closer to approximately 575–580 kJ/mol (actual: 577 kJ/mol). The Group 2 → Group 13 exception must be considered when interpolating.