What this chapter is about
This chapter introduces how elements are organised in the periodic table based on their atomic number and electronic configuration. You will learn about the historical attempts to classify elements, starting from early groupings to Mendeleev's periodic table and finally the modern periodic table based on atomic number.
The chapter explains periodic properties — characteristics that show regular patterns as you move across a period or down a group. These include atomic radius, ionisation enthalpy, electron gain enthalpy, electronegativity and valence. Understanding these trends helps predict chemical behaviour of elements without memorising individual properties.
After studying this chapter, you should be able to locate any element in the periodic table using its electronic configuration, explain why elements in the same group show similar chemistry, and predict relative values of periodic properties for given elements.
Key ideas
- Modern periodic law: Properties of elements are a periodic function of their atomic number, not atomic mass as Mendeleev originally proposed.
- Electronic configuration determines position: The period number equals the highest principal quantum number (n) of occupied orbitals; the block (s, p, d, f) depends on which subshell receives the last electron.
- Atomic radius decreases across a period because increasing nuclear charge pulls electrons closer, even though electrons enter the same shell.
- Atomic radius increases down a group because each successive element has an additional electron shell, placing valence electrons farther from the nucleus.
- Ionisation enthalpy increases across a period (more energy needed to remove an electron) due to stronger nuclear attraction; it decreases down a group as the outermost electron is farther from the nucleus.
- Electron gain enthalpy generally becomes more negative across a period (energy released on gaining an electron), reflecting greater tendency to accept electrons; halogens have the most negative values.
- Electronegativity increases across a period and decreases down a group; fluorine is the most electronegative element.
- Metallic character decreases across a period (elements lose electrons less readily) and increases down a group.
Formulas and facts to remember
1. Effective nuclear charge (Z_eff): Z_eff = Z − σ, where Z is atomic number and σ is shielding constant. It measures the net positive charge felt by a valence electron.
2. Ionisation enthalpy (Δ_i H): Minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state. Unit: kJ mol⁻¹.
3. Electron gain enthalpy (Δ_eg H): Enthalpy change when an electron is added to an isolated gaseous atom. A negative value means energy is released.
4. Electronegativity: A measure of an atom's ability to attract shared electrons in a covalent bond. Pauling scale is commonly used; fluorine = 4.0.
5. Covalent radius: Half the distance between nuclei of two identical atoms bonded covalently.
6. Periods and blocks: Period 1 has 2 elements (1s); periods 2 and 3 have 8 elements each (s and p blocks); period 4 has 18 elements (s, p and d blocks).
7. Diagonal relationship: Certain pairs (Li–Mg, Be–Al, B–Si) show similar properties because the increase in nuclear charge across a period offsets the increase in size down a group.
Worked examples
### Example 1: Locating an element in the periodic table
Problem: An element has atomic number 19. Find its period, group and block.
Solution:
Electronic configuration of Z = 19: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹.
Highest principal quantum number of occupied orbital = 4, so period = 4.
Last electron enters the 4s subshell, so block = s.
For s-block elements, group number = number of electrons in outermost s subshell = 1.
Answer: Period 4, Group 1, s-block. The element is potassium (K).
### Example 2: Comparing ionisation enthalpies
Problem: Arrange the following in order of increasing first ionisation enthalpy: Na, Mg, Al.
Solution:
All three are in period 3. General trend: ionisation enthalpy increases across a period.
Expected order: Na < Mg < Al.
However, Mg (1s² 2s² 2p⁶ 3s²) has a completely filled 3s subshell, which is more stable than Al (3s² 3p¹) where the single 3p electron is easier to remove.
Correct order: Na < Al < Mg.
Answer: Na has the lowest; Mg has the highest first ionisation enthalpy among the three.
### Example 3: Predicting electron gain enthalpy trend
Problem: Why does chlorine have a more negative electron gain enthalpy than fluorine, even though fluorine is more electronegative?
Solution:
Fluorine is very small; its 2p subshell is compact. An incoming electron experiences strong repulsion from electrons already present in this small volume.
Chlorine's 3p subshell is larger, so electron–electron repulsion is less. The additional electron is accommodated more easily, releasing more energy.
Answer: Chlorine releases more energy (Δ_eg H ≈ −349 kJ mol⁻¹) than fluorine (Δ_eg H ≈ −328 kJ mol⁻¹) because of lower inter-electronic repulsion in the larger chlorine atom.
Common mistakes
- Assuming atomic radius always increases with atomic number → it decreases across a period because effective nuclear charge rises while the shell number stays the same.
- Thinking higher electronegativity means more negative electron gain enthalpy in every case → noble gases have positive or near-zero electron gain enthalpy despite being after halogens.
- Confusing ionisation enthalpy with electron gain enthalpy → ionisation removes an electron (energy absorbed); electron gain adds an electron (usually energy released for non-metals).
- Believing Mendeleev arranged elements by atomic number → he used atomic mass; Moseley later established atomic number as the basis.
- Forgetting anomalies due to half-filled and fully filled subshell stability → Mg has higher ionisation enthalpy than Al; N has higher ionisation enthalpy than O.
Quick revision
- Period = highest principal quantum number; group = total valence electrons for main-group elements.
- Across a period: size decreases, ionisation enthalpy increases, electronegativity increases, metallic character decreases.
- Down a group: size increases, ionisation enthalpy decreases, electronegativity decreases, metallic character increases.
- s-block: groups 1–2; p-block: groups 13–18; d-block: groups 3–12; f-block: lanthanoids and actinoids.
- Electron gain enthalpy is most negative for halogens; Cl > F in magnitude due to size effect.