SYLLABUS OVERVIEW:
In order to study elements in an organised manner, they need to be classified. Dmitri Mendeleev gave the first major periodic table based on atomic masses, but it could not explain the positions of isotopes and rare earth metals.
These defects were removed when Henry Moseley put forward the modern periodic table.
Modern Periodic Law: "The physical and chemical properties of elements are the periodic functions of their atomic number."
Periodic Table: A tabular arrangement of elements in groups (vertical columns) and periods (horizontal rows) highlighting the regular trends in properties.
The modern periodic table has 18 vertical columns known as groups. Important groups:
There are 7 horizontal rows known as periods. The number of shells present in an atom determines its period.
Q: Given the atomic number of Sodium is 11. Find its Group and Period.
Solution:Electronic configuration of Na (11) = 2, 8, 1.
It has 3 shells $\rightarrow$ Period 3.
It has 1 valence electron $\rightarrow$ Group 1.
Periodicity: The reappearance of properties at regular intervals or gradual variation in properties at regular intervals is called periodicity.
Cause of periodicity: The recurrence of similar electronic configuration, i.e. having the same number of electrons in the outermost orbit.
Example: All Halogens (Group 17) have seven valence electrons, hence they show similar properties like being coloured non-metals and strong oxidising agents.
Valency is the combining capacity of an element.
Properties which show a regular gradation as we move across a period or down a group are Periodic Properties.
It is the distance between the centre of the nucleus of an atom and its outermost shell.
Depends on two factors: (i) Number of shells (increases size) and (ii) Nuclear charge (increases inward pull, decreases size).
Key Exceptions:
1. Inert Gases: Size of inert gases is larger than the halogens before them due to complete octets causing maximum inter-electronic repulsions.
2. Cation is always smaller than the parent atom (higher effective nuclear charge). Anion is always larger than the parent atom (lower effective nuclear charge).
Q: Arrange Mg, Cl, Na, S, Si in increasing order of atomic size.
Solution:They belong to Period 3. Atomic size decreases from left to right. Left to right order: Na, Mg, Si, S, Cl.
Increasing size: Cl < S < Si < Mg < Na.
Nature of Oxides: Across a period, the basic nature of oxides decreases and acidic nature increases (e.g., $Na_2O$ is strongly basic, $SO_3$ is acidic). Down a group, basic nature increases.
The energy required to remove an electron from a neutral isolated gaseous atom and convert it into a positively charged gaseous ion. ($M + \text{I.E.} \rightarrow M^+ + e^-$)
Unit: $\text{kJ mol}^{-1}$ or eV/atom.
Note: Helium has the highest ionisation energy; Caesium has the lowest.
The amount of energy released while converting a neutral gaseous isolated atom into a negatively charged gaseous ion by the addition of an electron. ($X + e^- \rightarrow X^- + \text{E.A.}$)
Crucial Exception: Fluorine has a lower E.A. than Chlorine. Reason: Fluorine's atom is extremely small, causing strong inter-electronic repulsions for any incoming electron.
The tendency of an atom in a molecule to attract the shared pair of electrons towards itself.
Fluorine has the highest electronegativity (4.0).
| Property | Alkali metals [Group 1] | Halogens [Group 17] |
|---|---|---|
| Elements | Li, Na, K, Rb, Cs | F, Cl, Br, I, At |
| Valence Electrons | 1 (Highly electropositive) | 7 (Highly electronegative) |
| Atomic Size | Largest in their period. | Smallest in their period. |
| Ionisation Energy | Lowest I.E. in their period. | High I.E. in their period. |
| Chemical Action | Strong reducing agents. | Strong oxidising agents. |
| Properties | Across the period (Left to right) | Down the group (Top to bottom) |
|---|---|---|
| No. of valence electrons | Increases | Remains same |
| Atomic size (radius) | Decreases (↓) | Increases (↑) |
| Ionisation energy | Increases (↑) | Decreases (↓) |
| Metallic character | Decreases (↓) | Increases (↑) |
| Non-metallic character | Increases (↑) | Decreases (↓) |
| Electron affinity | Increases (↑) | Decreases (↓) |
| Electronegativity | Increases (↑) | Decreases (↓) |
Core application questions to master the class lecture.
Give reasons: The size of a $Cl^-$ ion is greater than the size of a Cl atom.
Solution:A Chlorine atom (Cl) has 17 protons and 17 electrons. When it gains an electron to form a chloride ion ($Cl^-$), it has 17 protons but 18 electrons. The effective positive charge per electron becomes less, so the inward pull decreases and the size expands.
A group of elements: Boron, Aluminium, Gallium, Indium, Thallium.
(a) Which element has the most metallic character?
(b) Which element would be expected to have the highest electronegativity?
(a) Thallium. (Because metallic character increases down the group due to increasing atomic size).
(b) Boron. (Electronegativity decreases down the group. Boron is at the top).
Give reasons: The oxidising power of elements increases from left to right along a period.
Solution:Oxidising power depends on the tendency to gain electrons (non-metallic character). Across a period from left to right, atomic size decreases and nuclear charge increases, which strongly increases the atom's ability to pull and gain incoming electrons.