8.1 Modern Periodic Table
⚛️ Basis of Modern Periodic Table
- Arranged by increasing atomic number (Moseley, 1913)
- Properties repeat at regular intervals → Periodicity
- Horizontal rows: Periods (7 total)
- Vertical columns: Groups (18 total)
- Elements in same group = similar chemical properties
🔄 Electronic Configuration & Periodicity
- Across period: Electronic configuration changes gradually
- Down group: Same outermost shell configuration
- Periodicity = recurrence of properties at regular intervals
Mendeleev: Arranged by atomic mass
Modern: Arranged by atomic number (protons in nucleus)
8.2 Salient Features: Periods
📏 Periods in Periodic Table
| Period No. | Name | Elements | Shell Filled |
|---|---|---|---|
| 1st | Very Short | 2 | 1st (K) |
| 2nd | Short | 8 | 2nd (L) |
| 3rd | Short | 8 | 3rd (M) |
| 4th | Long | 18 | 4th (N) |
| 5th | Long | 18 | 5th (O) |
| 6th | Very Long | 32 | 6th (P) |
| 7th | Very Long | 32 | 7th (Q) |
🌟 Special Series
- Lanthanides: 14 elements after La (57) – 4f series
- Actinides: 14 elements after Ac (89) – 5f series
- Also called f-block elements or rare earths
- Placed separately to keep table manageable
8.2 Salient Features: Groups
📊 Groups & Family Names
| Group | Family Name | Outer Config | Type |
|---|---|---|---|
| 1 | Alkali Metals | ns¹ | s-block |
| 2 | Alkaline Earth | ns² | s-block |
| 3-12 | Transition Metals | (n-1)d¹⁻¹⁰ ns² | d-block |
| 13 | Boron Family | ns² np¹ | p-block |
| 14 | Carbon Family | ns² np² | p-block |
| 15 | Nitrogen Family | ns² np³ | p-block |
| 16 | Oxygen Family | ns² np⁴ | p-block |
| 17 | Halogens | ns² np⁵ | p-block |
| 18 | Noble Gases | ns² np⁶ | p-block |
🧩 Blocks Classification
- s-block: Groups 1-2 (alkali & alkaline earth metals)
- p-block: Groups 13-18 (includes all non-metals, metalloids)
- d-block: Groups 3-12 (transition metals)
- f-block: Lanthanides & Actinides (inner transition)
Normal: Groups 1-2, 13-17 – inner shells complete, outer incomplete
Transition: Groups 3-12 – inner d-subshells being filled
8.3 Similarities in Chemical Properties
⚡ Group 1: Alkali Metals
- 1 electron in outermost shell → strong tendency to lose electron
- Electropositive metals
- Reactivity increases down the group
- Oxides are basic: Na₂O + H₂O → 2NaOH
🧪 Group 17: Halogens
- 7 electrons in outermost shell → strong tendency to gain electron
- Electronegative non-metals
- Reactivity decreases down the group
- “Halogen” = salt forming
🔄 Ion Formation Trends
| Direction | Charge on Ion | Example |
|---|---|---|
| Left to Right | +Group Number | Group 1 → 1+, Group 2 → 2+ |
| Right to Left | -(18-Group No.) | Group 17 → 1-, Group 16 → 2- |
8.4 Periodic Properties
📐 Atomic Radius
- Definition: Half distance between nuclei of bonded identical atoms
- Unit: pm (1 pm = 10⁻¹² m)
- Period trend: Decreases left to right (increased nuclear charge)
- Group trend: Increases top to bottom (new shells added)
⚡ Ionization Energy
- Definition: Energy to remove outermost electron
- Unit: kJ mol⁻¹
- Period trend: Increases left to right (smaller size)
- Group trend: Decreases top to bottom (larger size)
- 1st IE < 2nd IE < 3rd IE (increasing difficulty)
🧲 Electron Affinity & Electronegativity
| Property | Definition | Trend Period | Trend Group |
|---|---|---|---|
| Electron Affinity | Energy change adding electron | Increases | Decreases |
| Electronegativity | Ability to attract electrons | Increases | Decreases |
Trend Data Tables
📊 Atomic Radii Trends
| 2nd Period Elements (pm) | |||||||
|---|---|---|---|---|---|---|---|
| Li | Be | B | C | N | O | F | Ne |
| 152 | 113 | 88 | 77 | 75 | 73 | 71 | 69 |
| Group 1 Elements | Inner Electrons | Atomic Radius (pm) |
|---|---|---|
| Li | 2 | 152 |
| Na | 10 | 186 |
| K | 18 | 227 |
| Rb | 36 | 248 |
| Cs | 54 | 265 |
📈 Ionization Energy Data
| 2nd Period IE (kJ mol⁻¹) | |||||||
|---|---|---|---|---|---|---|---|
| Li | Be | B | C | N | O | F | Ne |
| 520 | 899 | 801 | 1086 | 1402 | 1314 | 1681 | 2081 |
F = 4.0 (highest), O = 3.4, N = 3.0, Cl = 3.2
Most electronegative: Top-right corner (F)
Most electropositive: Bottom-left corner (Fr)
Exercise Questions
Multiple Choice Questions:
i) In which period and group will you place the element which is an important part of the solar cell?
Answer: (a) Third period and 14th group
Explanation: Silicon (Si) is used in solar cells. Atomic number 14: Period 3, Group 14 (Carbon family).
ii) Identify the electronic configuration of the outermost shell of a transition metal.
Answer: (b) nd¹⁰ns²
Explanation: Transition metals have (n-1)d and ns orbitals. General configuration: (n-1)d¹⁻¹⁰ ns².
iii) Which is the softest metal?
Answer: (a) Na
Explanation: Sodium is so soft it can be cut with a knife. Alkali metals (Group 1) are generally soft.
iv) A yellow solid element exists in allotropic forms which is also present in fossil fuels.
Answer: (d) Sulphur
Explanation: Sulfur is yellow, has allotropes (rhombic, monoclinic), and is present in coal and petroleum.
v) How many electrons can nitrogen accept in its outermost shell?
Answer: (b) 3
Explanation: Nitrogen has 5 valence electrons (2s²2p³), needs 3 to complete octet (2s²2p⁶).
vi) Which element is the most reactive element?
Answer: (c) Fluorine
Explanation: Fluorine has highest electronegativity (4.0) and strongest electron affinity, making it most reactive non-metal.
vii) Which element has the highest melting point?
Answer: (a) Na
Explanation: Among alkali metals, melting point decreases down group: Li (180°C) > Na (98°C) > K (63°C) > Rb (39°C) > Cs (28°C).
x) The element having less value of ionization energy and less value of electron affinity is likely to belong to:
Answer: (a) Group 1
Explanation: Group 1 metals have low IE (easy to lose electrons) and low EA (don’t readily accept electrons).
Short Answer Questions:
i. Why was atomic number chosen to arrange the elements in the periodic table?
Answer: Atomic number determines chemical properties. When arranged by atomic number, elements show regular periodicity in properties. Moseley discovered that atomic number (protons) not atomic mass should be the basis.
ii. What is the significance of the word “periodic”?
Answer: “Periodic” means repeating at regular intervals. In periodic table, properties of elements repeat periodically when arranged by increasing atomic number.
iii. Why does the size of a period increase as we move down the periodic table?
Answer: Lower periods accommodate more electrons in new shells and subshells. Period 1: only s-subshell (2 elements), Period 2: s and p (8), Period 4: s, d, p (18), Period 6: s, f, d, p (32).
iv. In a group, the elements have the same number of electrons in the outermost shell. Why is it so?
Answer: Group number indicates number of valence electrons. All elements in same group have identical electronic configuration in outermost shell (ns¹ for Group 1, ns²np⁵ for Group 17, etc.).
v. Do you expect calcium to be more reactive than sodium? Give the reason.
Answer: No, sodium is more reactive than calcium. Sodium (Group 1) has 1 valence electron with lower ionization energy (496 kJ/mol) than calcium (Group 2, 2 valence electrons, 1st IE = 590 kJ/mol). Sodium loses its single electron more easily.
vi. Which element has maximum and minimum atomic radius in third period?
Answer: Maximum: Sodium (Na) – 186 pm; Minimum: Chlorine (Cl) – 99 pm. Atomic radius decreases across period due to increasing nuclear charge.
Constructed Response Questions:
i. Suppose a new element is discovered. Where would you like to accommodate this element in the periodic table?
Answer: Determine its atomic number. Place it in appropriate period (based on electron shells) and group (based on valence electrons). If it has atomic number > 118, it would extend period 7 or start period 8.
ii. What is the first element of the periodic table? Will it lose an electron or gain it?
Answer: Hydrogen (H). It can both lose electron (forming H⁺) like alkali metals or gain electron (forming H⁻) like halogens, but typically loses electron in reactions with non-metals.
iii. Atomic radii of boron and aluminum are 88 pm and 125 pm respectively. Which element is expected to lose electron or electrons easily?
Answer: Aluminum (Al) will lose electrons more easily. Larger atomic radius means outer electrons are farther from nucleus, experiencing less attraction. Lower ionization energy for Al (577 kJ/mol) vs B (801 kJ/mol).
iv. How would you find the atomic radius of an atom?
Answer: Measure distance between nuclei of two identical bonded atoms (covalent radius for non-metals, metallic radius for metals). Atomic radius = half of this distance. For example, Cl-Cl bond = 198 pm, so Cl radius = 99 pm.
Descriptive Questions:
i. Which information is needed to locate elements if you don’t know atomic number? Is atomic mass helpful?
Answer: Need atomic number (protons) for exact placement. Atomic mass is not reliable because of isotopes and doesn’t always increase with atomic number. Mendeleev used atomic mass but left gaps for undiscovered elements; modern table uses atomic number.
ii. How many blocks of elements are present? Are blocks helpful in studying properties?
Answer: Four blocks: s, p, d, f. Yes, very helpful: s-block = reactive metals, p-block = non-metals/metalloids, d-block = transition metals with variable oxidation states, f-block = inner transition metals with similar properties.
iii. Explain variation in atomic radius and ionization energy across periods.
Answer:
Atomic radius: Decreases left to right. Increasing nuclear charge pulls electrons closer, no new shells added.
Ionization energy: Increases left to right. Smaller atoms hold electrons more tightly, harder to remove.
Exceptions: Be vs B (p-orbital easier to remove than s), N vs O (half-filled p³ more stable).
v. Write names of four non-metals which exist in solid state at normal temperature.
Answer: Carbon (C), Sulfur (S), Phosphorus (P), Iodine (I). Also: Selenium (Se), Tellurium (Te). Most non-metals are gases (N, O, F, Cl, noble gases) or liquids (Br).
Investigative Questions:
i. Arrangement of elements in periodic table is a remarkable achievement. Comment citing benefits.
Answer: Benefits: 1) Systematic organization of all known elements, 2) Prediction of properties of undiscovered elements, 3) Understanding periodicity and trends, 4) Classification into metals/non-metals/metalloids, 5) Basis for chemical bonding theories, 6) Prediction of compound formation, 7) Foundation for quantum mechanical understanding of atoms.
iii. Modern periodic table is amended form of Mendeleev’s table. Elaborate differences.
Answer:
Mendeleev (1869): Based on atomic mass, left gaps for undiscovered elements, 8 groups, no concept of atomic number.
Modern (Moseley 1913): Based on atomic number (protons), explains position anomalies (Ar-K, Co-Ni), 18 groups with subgroups, includes noble gases, separates lanthanides/actinides, based on electronic configuration.
Key difference: Atomic number (modern) vs atomic mass (Mendeleev).
Quick Exercises from Text:
1. Element with configuration sp³. Find period, group, block.
Answer: Period: 2 (n=2), Group: 14 (2+2=4 valence electrons), Block: p-block (p-orbital being filled). Example: Carbon (C) – 1s²2s²2p².
2. Element with eight electrons in outermost shell. Group? Physical state?
Answer: Group 18 (Noble gases: ns²np⁶). Physical state: Gas at room temperature (all noble gases are gases).
3. Element belongs to 16th group and is a gas. Which period?
Answer: Period 2 (Oxygen, O) or Period 3 (Sulfur, S) but sulfur is solid. Only oxygen in Group 16 is gas at room temp, so Period 2.