Chemistry chapter Periodic Table
Chapter 3 — Periodic Table
Periodicity • Trends • Groups & Periods • Properties of Elements
AFNS Initial Test Preparation | servicesinpak.com
TOPIC 1: History and Development of the Periodic Table
Historical Development
| Scientist | Contribution | Key Limitation |
| Döbereiner (1829) | Law of Triads: groups of 3 elements where middle element properties are average of other two (e.g., Li, Na, K) | Only applied to some elements — not a complete system |
| Newlands (1864) | Law of Octaves: when elements arranged by atomic mass, every 8th element has similar properties (like musical octaves) | Failed after calcium — didn’t work for heavier elements |
| Mendeleev (1869) | Arranged 63 known elements by ATOMIC MASS in rows and columns. Elements with similar properties placed in same column. Left GAPS for undiscovered elements. PREDICTED properties of missing elements. | Used atomic mass not atomic number — some elements were misplaced (e.g., Ar and K, Co and Ni, Te and I) |
| Moseley (1913) | Used X-ray spectroscopy to determine ATOMIC NUMBERS. Rearranged periodic table by ATOMIC NUMBER — resolved all anomalies of Mendeleev’s table. | — |
| KEY FACT | Modern Periodic Law (based on Moseley): ‘The physical and chemical properties of elements are periodic functions of their ATOMIC NUMBERS’ Mendeleev’s original law used ATOMIC MASS — now replaced by ATOMIC NUMBER Key difference: Mendeleev predicted gallium (eka-aluminium), germanium (eka-silicon), scandium Modern periodic table has 118 elements confirmed (as of 2023) |
TOPIC 2: Structure of the Modern Periodic Table
Periods and Groups
The modern periodic table has 7 PERIODS (horizontal rows) and 18 GROUPS (vertical columns). Elements in the same GROUP have similar chemical properties because they have the same number of VALENCE ELECTRONS.
| Period | Elements | Electron Shells Filled | Key Information |
| Period 1 | H, He | n=1 shell only | 2 elements — only 1s subshell |
| Period 2 | Li to Ne | n=1, n=2 | 8 elements — fills 2s and 2p |
| Period 3 | Na to Ar | n=1, 2, 3 | 8 elements — fills 3s and 3p (SHORT period) |
| Period 4 | K to Kr | n=1, 2, 3, 4 | 18 elements — first LONG period (includes 3d transition metals) |
| Period 5 | Rb to Xe | n=1–5 | 18 elements — 4d transition metals |
| Period 6 | Cs to Rn | n=1–6 | 32 elements — 4f lanthanides included |
| Period 7 | Fr to Og | n=1–7 | 32 elements — 5f actinides included (radioactive) |
Groups — Classification
| Group Number | Old Notation | Name | Key Elements | Valence Electrons |
| Group 1 | IA | Alkali Metals (except H) | Li, Na, K, Rb, Cs, Fr | 1 (ns¹) |
| Group 2 | IIA | Alkaline Earth Metals | Be, Mg, Ca, Sr, Ba, Ra | 2 (ns²) |
| Groups 3–12 | IIIB–IIB | Transition Metals (d-block) | Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn etc. | Variable (d electrons) |
| Group 13 | IIIA | Boron Family | B, Al, Ga, In, Tl | 3 (ns² np¹) |
| Group 14 | IVA | Carbon Family | C, Si, Ge, Sn, Pb | 4 (ns² np²) |
| Group 15 | VA | Nitrogen Family (Pnictogens) | N, P, As, Sb, Bi | 5 (ns² np³) |
| Group 16 | VIA | Oxygen Family (Chalcogens) | O, S, Se, Te, Po | 6 (ns² np⁴) |
| Group 17 | VIIA | Halogens | F, Cl, Br, I, At | 7 (ns² np⁵) — most reactive non-metals |
| Group 18 | VIII A (0) | Noble Gases | He, Ne, Ar, Kr, Xe, Rn | 8 (ns² np⁶) — except He=2 |
| KEY FACT | Block classification of the periodic table: • s-block: Groups 1-2 (fills s orbitals) — includes H and He • p-block: Groups 13-18 (fills p orbitals) — includes nonmetals, metalloids, noble gases • d-block: Groups 3-12 (fills d orbitals) — transition metals • f-block: Lanthanides (Ce-Lu) and Actinides (Th-Lr) — rare earth elements Metals: left and centre of table | Non-metals: top right | Metalloids/Semimetals: staircase boundary |
TOPIC 3: Periodic Trends — Most Tested in AFNS
- Atomic Radius
Atomic radius = distance from nucleus to outermost electron shell. Cannot be measured directly — defined as half the distance between two identical bonded atoms (covalent radius) or between adjacent atoms in a metal (metallic radius).
| Direction | Trend | Reason | Example |
| Across a period (left → right) | DECREASES ⭐ (atomic radius gets SMALLER) | Nuclear charge (Z) increases → more protons attract electrons more strongly → electrons pulled closer → radius decreases. Same number of shells but more nuclear pull. | Na (186 pm) > Mg (160 pm) > Al (143 pm) > Si (118 pm) > P (110 pm) > S (104 pm) > Cl (99 pm) |
| Down a group (top → bottom) | INCREASES ⭐ (atomic radius gets LARGER) | New electron shell added at each period → outermost electrons further from nucleus. Also: shielding effect increases (inner electrons shield outer electrons from nuclear charge). | Li(152) < Na(186) < K(227) < Rb(248) < Cs(265 pm) |
| QUICK TIP | Atomic radius trend summary: • SMALLEST atom = Helium (He) — noble gas, smallest atomic radius • Actually, among elements with more than 2 electrons: FLUORINE has smallest atomic radius ⭐ • LARGEST atomic radius = Francium (Fr) — bottom-left of table • Going LEFT → radius increases | Going DOWN → radius increases • CATION smaller than parent atom (lost electrons → less repulsion → shells contract) • ANION larger than parent atom (gained electrons → more repulsion → shells expand) |
- Ionisation Energy (IE)
Ionisation Energy = energy required to REMOVE one electron from a gaseous atom (or ion). The energy for the first removal is First Ionisation Energy (IE₁):
X(g) → X⁺(g) + e⁻ ΔH = +IE₁ (always positive — endothermic)
| Direction | Trend | Reason |
| Across a period (left → right) | INCREASES ⭐ | Atomic radius decreases → electrons held more tightly → harder to remove. Also: nuclear charge increases. |
| Down a group (top → bottom) | DECREASES ⭐ | Atomic radius increases → outermost electrons further from nucleus → held less tightly → easier to remove. Shielding increases. |
| Exceptions | NOT perfectly smooth | N > O (N has stable half-filled 2p³ — harder to remove than O’s 2p⁴ which has one paired electron) Mg > Al (Mg has stable 2s² full subshell) |
| KEY FACT | Successive ionisation energies: IE₁ < IE₂ < IE₃ < IE₄ … (each successive IE is larger — harder to remove from increasingly positive ion) Sudden large jump in IE indicates crossing a shell boundary: Na: IE₁=496, IE₂=4562 kJ/mol → HUGE jump after 1st (2nd electron from inner shell) Mg: HUGE jump after 2nd (3rd from inner shell) → confirms Mg in Group 2 (2 valence electrons) |
- Electron Affinity (EA)
Electron Affinity = energy change when a neutral gaseous atom GAINS one electron to form a negative ion:
X(g) + e⁻ → X⁻(g) ΔH = EA (usually negative — exothermic = energy released)
- Across a period (left → right): EA generally BECOMES MORE NEGATIVE (more exothermic) — atoms more eager to gain electrons as they approach noble gas configuration
- Down a group: EA BECOMES LESS NEGATIVE — atom larger, electron added further from nucleus, less attraction
- EXCEPTION: F has lower EA than Cl ⭐ — F is so small that adding an electron causes significant electron-electron repulsion in compact 2p orbital. Cl has higher EA than F!
- Group 2 (filled s²) and Group 5 (half-filled p³) have lower EA due to extra stability of their configurations
- Electronegativity
Electronegativity = ability of an atom to ATTRACT bonding electrons toward itself in a covalent bond. Pauling scale (dimensionless, F=4.0).
| Direction | Trend | Reason | Most/Least Electronegative |
| Across a period (left → right) | INCREASES ⭐ | Smaller atomic radius + more nuclear charge → stronger pull on shared electrons | Most electronegative in period = rightmost nonmetal (F in period 2) |
| Down a group (top → bottom) | DECREASES ⭐ | Larger atomic radius + more shielding → weaker pull on shared electrons | Least electronegative in group = bottommost element |
| Overall extremes | — | — | MOST electronegative = Fluorine (F=4.0) ⭐ LEAST electronegative = Caesium/Francium (~0.7) ⭐ |
| ⭐ REMEMBER FOR AFNS TEST
Electronegativity values (Pauling scale) to memorise: F=4.0 ⭐ (most electronegative element — ALWAYS) O=3.5 | N=3.0 | Cl=3.2 | Br=2.8 | C=2.5 | H=2.1 Na=0.9 | K=0.8 | Cs=0.7 Electronegativity difference determines bond type: <0.4 = Non-polar covalent | 0.4–1.7 = Polar covalent | >1.7 = Ionic bond H₂O: O-H bond = 3.5−2.1 = 1.4 (polar covalent) | NaCl: 3.2−0.9 = 2.3 (ionic) |
TOPIC 4: More Periodic Trends
- Metallic Character
- Across a period (left → right): DECREASES — metals on left, metalloids in middle, non-metals on right
- Down a group: INCREASES — elements become more metallic going down
- Most metallic element = Francium (Fr) — bottom left ⭐
- Least metallic = Fluorine — top right (most non-metallic)
- Metalloids (semimetals): B, Si, Ge, As, Sb, Te, Po, At — borderline properties, semiconductors
- Melting and Boiling Points
No single simple trend — depends on bonding type and structure:
- Period 3 trend: Na, Mg, Al (metallic bonds — increase with more delocalized electrons), Si (giant covalent — very high MP), P, S, Cl, Ar (molecular — weak van der Waals, low MP)
- Within a group of metals: generally decreases going down (weaker metallic bonds)
- Within noble gases (Group 18): increases going down (more electrons → stronger London dispersion forces) He < Ne < Ar < Kr < Xe
- S has higher BP than O despite same group (S has more electrons, stronger van der Waals)
- Density
- Generally increases across a period (more protons/electrons, similar atomic radius → higher mass per volume)
- Down Group 1 (alkali metals): Li < Na < K < Rb < Cs (increases, except K slightly anomalous)
- Osmium (Os) and Iridium (Ir) are the densest elements (~22.6 g/cm³)
- Lithium (Li) is the least dense metal (0.53 g/cm³) — floats on water
TOPIC 5: Periodicity of Properties — Groups in Detail
Group 1 — Alkali Metals (Li, Na, K, Rb, Cs, Fr)
Configuration: ns¹. Lose 1 electron → form M⁺ ions. Soft, low-density, low-melting-point metals.
- React vigorously with water: 2M + 2H₂O → 2MOH + H₂ ↑ (Li slow, Na vigorous, K catches fire) ⭐
- React with oxygen: Li → Li₂O (oxide) | Na → Na₂O₂ (peroxide) | K,Rb,Cs → MO₂ (superoxide) ⭐
- React with halogens: 2M + X₂ → 2MX (ionic halides)
- Reactivity INCREASES down the group (IE decreases → easier to lose electron)
- All form strongly ALKALINE hydroxide solutions (MOH)
Group 2 — Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra)
Configuration: ns². Lose 2 electrons → form M²⁺. Harder, higher-melting than Group 1.
- React with water (less vigorous than Group 1): M + 2H₂O → M(OH)₂ + H₂ (Be=no reaction, Mg=slow, Ca=steady, Sr,Ba=vigorous)
- Flame colours: Li=crimson red, Na=yellow/orange, K=lilac/violet, Ca=brick red, Sr=crimson, Ba=apple green, Cu=blue-green ⭐
- Reactivity increases down the group. Beryllium is anomalous — forms covalent compounds.
Group 17 — Halogens (F, Cl, Br, I, At)
Configuration: ns² np⁵. Need 1 electron → form X⁻ ions (halide ions). Most reactive non-metals.
- Physical state at room temperature: F₂=yellow gas, Cl₂=yellow-green gas, Br₂=red-brown liquid, I₂=purple-grey solid ⭐
- Reactivity DECREASES down group (EA decreases, harder to attract electron)
- Fluorine = most electronegative, most reactive halogen ⭐
- Displacement reactions: more reactive halogen displaces less reactive from salt solution Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq) ⭐
- Hydrogen halides: HF (very weak acid, strong H-bond), HCl > HBr > HI (acid strength increases down group)
Group 18 — Noble Gases (He, Ne, Ar, Kr, Xe, Rn)
Configuration: ns² np⁶ (He=1s²). Full outer shell → chemically inert (unreactive). All gases at room temperature.
- Zero valency — do NOT normally form compounds (small exceptions: XeF₂, XeF₄ formed under extreme conditions)
- Boiling points increase down group: He (−269°C) < Ne (−246°C) < Ar (−186°C) < Kr (−152°C) < Xe (−108°C) — more electrons → stronger van der Waals
- Used in: He=filling balloons, cryogenics | Ne=advertising signs (red light) | Ar=welding gas, light bulbs | Kr,Xe=flash photography
- The modern periodic table is arranged in order of increasing:
- A) Atomic mass
- B) Atomic number
- C) Number of neutrons
- D) Valence electrons
✔ Answer: B) Atomic number
- The periodic table has how many periods (rows)?
- A) 5
- B) 6
- C) 7
- D) 8
✔ Answer: C) 7
- The periodic table has how many groups (columns)?
- A) 7
- B) 14
- C) 18
- D) 20
✔ Answer: C) 18
- Elements in the same group have the same number of:
- A) Protons
- B) Neutrons
- C) Valence electrons
- D) Energy levels
✔ Answer: C) Valence electrons
- Elements in the same period have the same number of:
- A) Valence electrons
- B) Electron shells (energy levels)
- C) Neutrons
- D) Protons
✔ Answer: B) Electron shells (energy levels)
- Atomic radius generally increases as you go:
- A) Across a period left to right
- B) Down a group
- C) Right to left in a period
- D) Both B and C
✔ Answer: D) Both B and C
- Ionisation energy generally increases as you go:
- A) Down a group
- B) Right to left in a period
- C) Left to right across a period
- D) From metals to metalloids
✔ Answer: C) Left to right across a period
- Electronegativity generally increases:
- A) Down a group
- B) Left to right across a period
- C) From non-metals to metals
- D) As atomic mass increases
✔ Answer: B) Left to right across a period
- Group 1 elements are called:
- A) Alkaline earth metals
- B) Halogens
- C) Alkali metals
- D) Noble gases
✔ Answer: C) Alkali metals
- Group 2 elements are called:
- A) Alkali metals
- B) Alkaline earth metals
- C) Transition metals
- D) Halogens
✔ Answer: B) Alkaline earth metals
- Group 17 elements are called:
- A) Noble gases
- B) Alkali metals
- C) Halogens
- D) Transition metals
✔ Answer: C) Halogens
- Group 18 elements are called:
- A) Halogens
- B) Alkali metals
- C) Noble gases
- D) Lanthanides
✔ Answer: C) Noble gases
- Noble gases are chemically inert because they have:
- A) High atomic mass
- B) Complete outermost electron shells
- C) No protons
- D) Many neutrons
✔ Answer: B) Complete outermost electron shells
- Who developed the first periodic table arranged by atomic mass?
- A) Moseley
- B) Newlands
- C) Dmitri Mendeleev
- D) Bohr
✔ Answer: C) Dmitri Mendeleev
- Moseley modified the periodic table by arranging elements by:
- A) Atomic mass
- B) Atomic number
- C) Number of neutrons
- D) Electron configuration
✔ Answer: B) Atomic number
- Metallic character increases:
- A) Across a period left to right
- B) Down a group
- C) Right to left only
- D) As ionisation energy increases
✔ Answer: B) Down a group
- Sodium (Na) has atomic number 11. It belongs to group:
- A) 2
- B) 11
- C) 1
- D) 17
✔ Answer: C) 1
- Chlorine has atomic number 17. It belongs to group:
- A) 1
- B) 7
- C) 17
- D) 18
✔ Answer: C) 17
- The element with atomic number 2 is:
- A) Hydrogen
- B) Helium
- C) Lithium
- D) Neon
✔ Answer: B) Helium
- The element with atomic number 8 is:
- A) Carbon
- B) Nitrogen
- C) Oxygen
- D) Fluorine
✔ Answer: C) Oxygen
- Transition metals are found in groups:
- A) 1–2
- B) 3–12
- C) 13–17
- D) 17–18
✔ Answer: B) 3–12
- Electron affinity is the energy change when an atom:
- A) Loses an electron
- B) Gains an electron
- C) Loses a proton
- D) Gains a neutron
✔ Answer: B) Gains an electron
- Which period contains the element Sodium (Na)?
- A) Period 2
- B) Period 3
- C) Period 4
- D) Period 1
✔ Answer: B) Period 3
- Diagonal relationships in the periodic table occur between:
- A) Elements in the same group
- B) Elements in the same period
- C) An element and the one diagonally below-right
- D) Noble gases
✔ Answer: C) An element and the one diagonally below-right
- Which of these is a metalloid (semi-metal)?
- A) Sodium
- B) Chlorine
- C) Silicon
- D) Iron
✔ Answer: C) Silicon