7.0 PROPERTIES AND TRENDS ACROSS A PERIOD
The Periodic Table is a chemist's roadmap. While elements in the same group have similar properties, elements across a period (horizontal row) show a gradual transition from metallic to non-metallic character. This chapter delves into the fundamental reasons behind these trends by examining how atomic structure – specifically the number of protons and the arrangement of electrons – affects key properties like atomic radius, ionic radius, ionization energy, and electronegativity. Understanding these trends allows you to predict and explain the behaviour of elements without memorizing every fact.
7.1 TRENDS IN ATOMIC AND IONIC RADII
7.1.1 Atomic Radius
The atomic radius is the distance from the nucleus to the outermost electron (the valence shell). It is usually measured in picometers (pm) or angstroms (Å).
Trend across a period (left to right): DECREASES
Explanation: As we move from left to right across a period (e.g., from sodium to argon in Period 3), the following happens:
- The number of protons in the nucleus increases (increasing nuclear charge).
- Electrons are added to the same outer shell. There is no significant increase in shielding because the new electrons join the same principal energy level.
- The increasing positive charge pulls the electron cloud inwards more strongly, making the atomic radius smaller.
Example (Period 3): Sodium (Na) has a radius of about 186 pm, while chlorine (Cl) has a radius of about 99 pm. The atoms get smaller across the period.
7.1.2 Ionic Radius
The ionic radius is the size of an ion formed when an atom loses or gains electrons. The trends are slightly different because electron loss and gain change the structure.
Comparing cations (positive ions) with their parent atoms:
- Cations are smaller than their parent atoms. When a metal atom loses electrons to form a positive ion, it loses its entire outer shell. The remaining electrons are pulled in even more tightly by the same nuclear charge. For example, a sodium atom (Na) has a radius of 186 pm, but a sodium ion (Na⁺) has a radius of only about 102 pm.
Comparing anions (negative ions) with their parent atoms:
- Anions are larger than their parent atoms. When a non-metal gains electrons, the increased electron-electron repulsion in the outer shell causes the electron cloud to expand. For example, a chlorine atom (Cl) has a radius of 99 pm, but a chloride ion (Cl⁻) has a radius of about 181 pm.
Trend across a period for ions: This is more complex because we are comparing ions of different charges. Generally, moving from left to right across a period, the positive ions (cations) have a decreasing radius due to increasing nuclear charge. Then, after the noble gas, the negative ions (anions) are much larger, and their size also decreases across the period.
- Example (Period 3 ions): Na⁺ (102 pm) > Mg²⁺ (72 pm) > Al³⁺ (53.5 pm). Then we jump to the much larger Si⁴⁻ (theoretical, not common), P³⁻ (>200 pm), S²⁻ (184 pm), Cl⁻ (181 pm). The anions (S²⁻, Cl⁻) are much larger than the cations, but their size also decreases from S²⁻ to Cl⁻ due to increasing nuclear charge.
7.2 IONIZATION ENERGY
Ionization energy (IE) is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous positive ions. It is a measure of how tightly an atom holds onto its electrons.
First ionization energy refers to the removal of the first electron:
X(g) → X⁺(g) + e⁻
Trend across a period (left to right): INCREASES (generally)
Explanation:
- As we move across a period, the nuclear charge increases (more protons).
- The electrons are being added to the same shell, so the shielding remains roughly constant.
- The increasing nuclear charge holds the outer electrons more tightly, making them harder to remove. Therefore, more energy is required.
Example (Period 3): Sodium (Na) has a low first ionization energy (496 kJ/mol) because it wants to lose its one outer electron. Magnesium (Mg) is higher (738 kJ/mol) because it has a greater nuclear charge. Aluminium (Al) is slightly lower than magnesium (578 kJ/mol) because the electron removed is from a higher-energy p-orbital (it's easier to remove than an s-electron). Silicon (Si) is higher, Phosphorus (P) higher still, Sulfur (S) has a slight drop (because the electron removed is from a doubly occupied p-orbital, causing slight repulsion), then Chlorine (Cl) and Argon (Ar) are very high. Argon has the highest in the period because it has a full outer shell and does not want to lose an electron.
General trend: Across a period, ionization energy increases, with small drops at Group II (s² configuration) and Group V (p³ half-filled stability).
Trend down a group: Ionization energy decreases. As atomic radius increases and shielding increases, the outer electron is further from the nucleus and easier to remove.
7.3 ELECTRONEGATIVITY
Electronegativity is a measure of the tendency of an atom to attract a shared pair of electrons towards itself in a covalent bond. It is a dimensionless number, usually given on the Pauling scale.
Trend across a period (left to right): INCREASES
Explanation:
- As we move across a period, nuclear charge increases.
- Atomic radius decreases.
- The atoms get smaller and the nucleus has a greater pull on any bonding electrons. Therefore, electronegativity increases.
Example (Period 3): Sodium (Na) has a very low electronegativity (0.93) – it tends to lose electrons, not attract them. Magnesium (1.31), Aluminium (1.61), Silicon (1.90), Phosphorus (2.19), Sulfur (2.58), Chlorine (3.16) – chlorine is one of the most electronegative elements. Argon has no electronegativity value because it does not typically form bonds.
Trend down a group: Electronegativity decreases. Even though nuclear charge increases, atomic radius increases significantly, and shielding increases, making the pull on bonding electrons weaker. For example, fluorine is the most electronegative element (3.98), but chlorine (3.16) is less electronegative, bromine even less, and iodine less still.
7.4 SUMMARY OF TRENDS ACROSS A PERIOD
| Property | Trend across a period (left → right) | Reason |
|---|---|---|
| Atomic radius | Decreases | Increasing nuclear charge pulls electrons in closer; same shell. |
| Ionic radius | Cations decrease, then jump to larger anions, then anions decrease. | Loss of shell for cations; increased repulsion for anions; nuclear charge increases. |
| Ionization energy | Increases (with small drops) | Increasing nuclear charge holds electrons more tightly. |
| Electronegativity | Increases | Smaller atoms with higher nuclear charge attract bonding electrons more. |
| Metallic character | Decreases | Atoms lose electrons less easily; tendency to gain electrons increases. |
✍️ COMPREHENSIVE PRACTICE QUESTIONS
Section A: Short Answer & Definitions
- Define atomic radius. State and explain the trend in atomic radius across Period 3 from sodium to argon.
- Why is a sodium ion (Na⁺) smaller than a sodium atom (Na)?
- Why is a chloride ion (Cl⁻) larger than a chlorine atom (Cl)?
- Define first ionization energy.
- Explain the general trend in first ionization energy across Period 3.
- Define electronegativity. Which element in Period 3 has the highest electronegativity and why?
- Why does electronegativity increase across a period?
- State the trend in metallic character across a period.
Section B: Application & Explanation
- Explain why the atomic radius of magnesium is smaller than that of sodium, even though magnesium has more protons and electrons.
- Arrange the following in order of increasing atomic radius: Al, Si, Mg, Cl. Explain your order.
- Explain why the first ionization energy of aluminium (578 kJ/mol) is slightly lower than that of magnesium (738 kJ/mol), despite aluminium being further to the right in the period.
- Why is there a slight drop in ionization energy between phosphorus and sulfur in Period 3? (Hint: consider electron configuration: P is [Ne]3s²3p³, S is [Ne]3s²3p⁴.)
- Arrange the following ions in order of increasing ionic radius: Mg²⁺, O²⁻, F⁻, Na⁺, Al³⁺. Explain your reasoning.
- Explain why fluorine has a higher electronegativity than chlorine, even though chlorine has more protons.
- Based on periodic trends, predict whether strontium (Sr, below calcium) will have a higher or lower ionization energy than calcium. Explain.
- Why does sodium (a metal) have a low electronegativity, while chlorine (a non-metal) has a high electronegativity?
- An atom X has a smaller atomic radius than atom Y, but both are in the same period. Which atom is likely to have the higher ionization energy? Explain.
- Explain why the ionic radius of Al³⁺ is much smaller than that of Na⁺, even though both are in the same period.
Section C: Fill in the Blanks
- Across a period, atomic radius __________ due to increasing __________ charge.
- A positive ion (cation) is __________ than its parent atom because it has lost its outer shell.
- Ionization energy is the energy required to remove an electron from a __________ atom.
- The general trend for first ionization energy across a period is to __________.
- Electronegativity is a measure of an atom's ability to __________ a shared pair of electrons.
- The element with the highest electronegativity in the Periodic Table is __________.
- Down a group, electronegativity __________ because atomic radius __________.
Section D: Data Analysis & Comparison
- The table below shows the atomic radii of elements in Period 3. Plot a rough graph and explain the trend.
Element: Na, Mg, Al, Si, P, S, Cl
Radius (pm): 186, 160, 143, 117, 110, 104, 99 - The first ionization energies (kJ/mol) for Period 3 are:
Na: 496, Mg: 738, Al: 578, Si: 787, P: 1012, S: 1000, Cl: 1251, Ar: 1520- a) Why is the value for Na so low?
- b) Why is there a drop from Mg to Al?
- c) Why is there a drop from P to S?
- d) Why does Ar have the highest value?
- Explain the difference in ionic radii between:
- a) Na⁺ (102 pm) and Mg²⁺ (72 pm)
- b) F⁻ (133 pm) and O²⁻ (140 pm)
Section E: Challenge / Multi-Concept
- Predict and explain the trend in electronegativity and ionization energy for the elements in Group I (Li to Fr) and across Period 2 (Li to Ne).
- An element X has a high first ionization energy, a high electronegativity, and a small atomic radius. Is X likely to be a metal or a non-metal? Suggest which group it might belong to.
- Explain why the second ionization energy of an element is always higher than the first ionization energy. Use sodium as an example.
- Compare the trends in atomic radius and ionic radius for the elements across Period 3. Why is the change in ionic radius more dramatic when moving from cations to anions?
- Using your knowledge of periodic trends, predict the following for bromine (Br, Period 4, Group VII) relative to chlorine (Cl, Period 3, Group VII):
- a) Atomic radius (larger or smaller?)
- b) Ionization energy (higher or lower?)
- c) Electronegativity (higher or lower?)
- Explain why the ionization energy of oxygen is lower than that of nitrogen, despite oxygen having a higher nuclear charge. (Hint: consider electron pairing in the p-orbitals.)
- Connect the concepts: How do the trends in ionization energy and electronegativity explain why metals form cations and non-metals form anions?
📝 ANSWERS TO SELECTED QUESTIONS
1. Atomic radius decreases across a period due to increasing nuclear charge pulling electrons inwards.
2. Na⁺ has lost its outer shell, and the remaining electrons are pulled in more tightly.
3. Cl⁻ has gained an electron, increasing repulsion and expanding the electron cloud.
4. Energy to remove 1 mole of electrons from 1 mole of gaseous atoms.
5. Ionization energy increases across a period due to higher nuclear charge holding electrons tighter.
6. Electronegativity: attraction for bonding electrons. Highest in Period 3 is Cl (small, high nuclear charge).
7. Atoms get smaller, nuclear charge increases, so attraction for bonding electrons increases.
8. Metallic character decreases (elements become less likely to lose electrons).
9. Mg has higher nuclear charge than Na, same shell, so greater pull on electrons.
10. Mg (160) > Al (143) > Si (117) > Cl (99). Atomic radius decreases left to right.
11. Al's outer electron is in a 3p orbital (higher energy, easier to remove) compared to Mg's 3s orbital.
12. P has half-filled 3p³ subshell (extra stability). S has a paired electron in one of the 3p orbitals, creating repulsion, making it easier to remove.
13. Increasing ionic radius: Al³⁺ (smallest, high charge) < Mg²⁺ < Na⁺ < F⁻ < O²⁻ (largest, extra repulsion). Order based on nuclear charge and number of electrons.
14. F is smaller than Cl, less shielding, so its nucleus attracts bonding electrons more strongly, despite fewer protons.
15. Sr is below Ca, so it has larger atomic radius and more shielding, thus lower ionization energy.
16. Na has low nuclear attraction for electrons, tends to lose them; Cl has high nuclear attraction, tends to gain/share.
17. X with smaller radius has higher ionization energy (electrons held tighter).
18. Al³⁺ has much higher nuclear charge (13+) than Na⁺ (11+), and has lost all outer electrons, so it's tiny.
19. decreases, nuclear
20. smaller
21. gaseous
22. increase
23. attract
24. fluorine
25. decreases, increases
26. a) Na has one outer electron, easy to lose. b) Al's electron from p-orbital (higher energy). c) S has electron pairing repulsion. d) Ar has full octet, very stable.
27. a) Mg²⁺ has higher nuclear charge pulling on fewer electrons. b) O²⁻ has more electrons (10 vs 10? Actually F⁻ has 10 e⁻, O²⁻ has 10 e⁻ as well? O²⁻ has 10 e⁻, F⁻ has 10 e⁻ – O²⁻ has lower nuclear charge (8+) than F⁻ (9+), so O²⁻ is larger.
28. Down Group I: IE decreases, electronegativity decreases (atomic radius increases). Across Period 2: IE increases, electronegativity increases.
29. Non-metal, likely Group VII or VI.
30. After losing one electron, the atom becomes a positive ion with a higher effective nuclear charge on remaining electrons, so more energy needed to remove another.
32. a) Larger (down group). b) Lower (easier to lose electron). c) Lower (larger atom, less attraction).
33. N has half-filled 2p³ subshell (stable), O has 2p⁴ with one paired electron (repulsion), so easier to remove.
34. Metals have low IE (lose e⁻ easily to form cations). Non-metals have high IE and high electronegativity (gain/share e⁻ to form anions).
These notes are your complete guide to the trends across a period. Visualize the atoms shrinking as nuclear charge increases. Understand why energy is needed to remove electrons and why some atoms are electron "greedy." Master these trends, and the Periodic Table becomes a powerful tool for prediction. Keep learning.