9.1 Properties of Group 1 Elements
⚡ Alkali Metals (ns¹)
- Outer configuration: ns¹ (one valence electron)
- Trend: Reactivity increases down the group
- Reason: Atomic size increases → easier electron loss
- Exception: Hydrogen (gas, non-metal)
💧 Reaction with Water
Trend: Steady (Li) → Vigorous (Na) → Violent (K) → Explosive (Cs)
🔥 Reaction with Chlorine
Trend: Heat needed → Vigorous → Violent
📊 Physical Properties
| Metal | Melting Point (°C) | Density (g/cm³) | Special Features |
|---|---|---|---|
| Li | 180 | 0.53 | Lightest metal |
| Na | 98 | 0.97 | Reacts vigorously with water |
| K | 64 | 0.86 | Violent reaction with water |
| Rb | 39 | 1.53 | Sinks in water |
| Cs | 28 | 1.87 | Explodes in water |
Key Points to Remember:
1. All alkali metals have ns¹ configuration
2. Reactivity ↑ down group (atomic size ↑)
3. Melting point ↓ down group (weaker metallic bonds)
4. Density ↑ down group (mass ↑ > volume ↑)
5. All react with water producing H₂ and metal hydroxide
9.3 Group Properties of Transition Elements
⚙️ d-Block Elements (Groups 3-12)
- Outer configuration: (n-1)d¹⁻¹⁰ ns²
- Properties: Hard metals with high density
- Melting/Boiling points: High
- Oxidation states: Variable (multiple valencies)
- Compounds: Often colored
- Physical: Malleable and ductile
- Special: High tensile strength
⚗️ Catalytic Properties
| Catalyst | Process | Reaction | Product |
|---|---|---|---|
| Iron (Fe) | Haber Process | N₂ + 3H₂ → 2NH₃ | Ammonia (fertilizer) |
| Vanadium(V) oxide | Contact Process | 2SO₂ + O₂ → 2SO₃ | Sulfuric acid |
| Platinum/Palladium | Catalytic Converter | CO → CO₂, NOx → N₂ | Less polluting gases |
| Nickel (Ni) | Hydrogenation | Oil + H₂ → Margarine | Solid fat |
1. Variable oxidation states allow electron transfer
2. Surface absorption activates reactants
3. d-orbitals participate in bonding
🌈 Color & Oxidation States
Transition Metals Memory Aid:
“CHaTiC VaN” (Common Transition Metals: Cr, Hf, Ti, C, V, N)
Properties: Hard, Heavy, High MP/BP, Have colors, Help as catalysts
9.2 Properties of Group 17 Elements
🧪 Halogens (ns²np⁵)
- Outer configuration: ns²np⁵ (seven valence electrons)
- State: Diatomic molecules (F₂, Cl₂, Br₂, I₂)
- Trend: Reactivity decreases down group
- Reason: Atomic size ↑ → electron affinity ↓
- Name: “Halogen” = salt former (Greek: hals + gen)
🎨 Physical Appearance
| Halogen | State at RT | Color | Special |
|---|---|---|---|
| Fluorine (F₂) | Gas | Pale yellow | Most reactive |
| Chlorine (Cl₂) | Gas | Yellowish-green | Swimming pool sterilizer |
| Bromine (Br₂) | Liquid | Red-brown fuming | Only liquid non-metal |
| Iodine (I₂) | Solid | Shiny grey → purple vapor | Sublimes on heating |
⚡ Redox Chemistry
Oxidation: Loss of electron (Na → Na⁺ + e⁻) → Reducing agent
Reduction: Gain of electron (Cl₂ + 2e⁻ → 2Cl⁻) → Oxidizing agent
Halogens: Strong oxidizing agents (F₂ > Cl₂ > Br₂ > I₂)
🔄 Displacement Reactions
Rule: More reactive halogen displaces less reactive halide
Oxidizing power: F₂ > Cl₂ > Br₂ > I₂
💥 Hydrogen Halides (HX)
Thermal stability: HF > HCl > HBr > HI (bond strength decreases)
Bond length: Increases down group → weaker bond
Halogen Facts:
1. All are diatomic (F₂, Cl₂, Br₂, I₂)
2. MP/BP ↑ down group (larger molecules, stronger van der Waals)
3. Reactivity ↓ down group (harder to gain electrons)
4. Color darkens down group
5. Form salts with metals (NaCl, KBr, etc.)
9.4 Properties of Noble Gases
👑 Group 18 Elements
- Elements: He, Ne, Ar, Kr, Xe, Rn
- Outer configuration: ns²np⁶ (except He: 1s²)
- State: Monoatomic gases
- Boiling points: Very low (increase down group)
- Reactivity: Very low (inert gases)
- Reason: Complete octet (stable configuration)
1. Complete valence shell (stable octet)
2. High ionization energies
3. No tendency to gain/lose electrons
4. Xenon forms compounds only with strong oxidizers (XeF₂, XeF₄, XeF₆)
📊 Boiling Point Trend
| Gas | Boiling Point (K) | Atomic Radius (pm) | Uses |
|---|---|---|---|
| He | 4.2 | 31 | Balloons, coolant |
| Ne | 27.1 | 38 | Neon signs |
| Ar | 87.3 | 71 | Welding, light bulbs |
| Kr | 119.8 | 88 | Photography flashes |
| Xe | 165.0 | 108 | Xenon lamps |
| Rn | 211 | 120 | Radioactive, medical |
Noble Gases Mnemonic:
“He Never Argues; Kr Xe Rarely”
He – Helium, Ne – Neon, Ar – Argon, Kr – Krypton, Xe – Xenon, Rn – Radon
Remember: BP increases down group due to increasing van der Waals forces
Exercise Questions – Complete Solutions
Multiple Choice Questions:
i) Which halogen will have the least reactivity with alkaline earth metals?
Answer: (b) Iodine
Detailed Explanation: Halogen reactivity decreases down group 17: F₂ > Cl₂ > Br₂ > I₂. Iodine is least reactive because:
1. Largest atomic size (down group)
2. Lowest electron affinity (hardest to gain electrons)
3. Weakest oxidizing power
4. Poorest ability to accept electron from metals
ii) Which compound do you expect to be coloured?
Answer: (d) NiCl₂
Detailed Explanation: Transition metal compounds are often colored due to d-d electron transitions.
NiCl₂: Nickel(II) chloride – Green color (transition metal compound)
Others: KCI (colorless), BaCl₂ (colorless), AlCl₃ (colorless) – All are normal elements without d-electrons
Reason: Partially filled d-orbitals allow electron transitions that absorb visible light
iii) In which element there exists the strongest forces of attraction between atoms?
Answer: (a) Mg
Detailed Explanation: Metallic bond strength depends on:
1. Number of delocalized electrons
2. Size of cation (smaller = stronger)
3. Charge density
Group 2 trend: Mg > Ca > Sr > Ba (decreasing down group)
Why Mg strongest: Smallest atomic radius (160 pm), highest charge/size ratio, strongest metallic bonding
iv) Elements of which group are all coloured?
Answer: Transition metals (not listed, but implied: d-block elements)
Detailed Explanation: The question has unclear options, but correct answer is transition elements (Groups 3-12).
Reason for color: Partially filled d-orbitals allow d-d transitions that absorb specific wavelengths of light
Examples: Cu²⁺ (blue), Fe³⁺ (yellow/brown), Cr³⁺ (green), MnO₄⁻ (purple)
v) Which halogen acid is unstable at room temperature?
Answer: (b) HI
Detailed Explanation: Thermal stability of HX decreases: HF > HCl > HBr > HI
Why HI unstable:
1. Longest H-I bond (weakest)
2. Iodine has lowest electronegativity
3. HI decomposes: 2HI → H₂ + I₂ at room temperature
4. Bond dissociation energy: HF (565) > HCl (431) > HBr (366) > HI (299 kJ/mol)
vi) Which oxide is the most basic oxide?
Answer: (a) Na₂O
Detailed Explanation: Basicity increases with metallic character and down groups.
Order: Na₂O > Li₂O > MgO > CO (CO is acidic)
Why Na₂O most basic:
1. Na is more metallic than Li (lower ionization energy)
2. Forms strong base NaOH in water: Na₂O + H₂O → 2NaOH
3. Oxide ion (O²⁻) strongly basic
4. CO is acidic (carbon dioxide forms carbonic acid)
vii) Which group elements are the most reactive elements?
Answer: (b) First group
Detailed Explanation: Alkali metals (Group 1) are most reactive metals.
Reason:
1. Lowest ionization energies (easy to lose electron)
2. Largest atomic sizes down group (Cs most reactive)
3. React violently with water, air, halogens
Note: Halogens (Group 17) are most reactive non-metals (F₂ most reactive)
x) Which property is correct for group 1 elements?
Answer: (a) Low catalytic activity
Detailed Explanation: Alkali metals have low catalytic activity because:
1. Too reactive – react with reactants
2. Form compounds easily rather than facilitate reactions
3. Not transition metals (no d-orbitals for catalysis)
Other options wrong:
(b) High density – False, alkali metals are light (Li floats on water)
(c) Low electrical conductivity – False, they are good conductors
(d) High melting point – False, they have low melting points
Short Answer Questions:
i. Why does it become easier to cut an alkali metal when we move from top to bottom in group 1?
Answer: Metallic bonding weakens down the group, making metals softer and easier to cut.
Detailed Explanation:
1. Atomic size increases down group (more electron shells)
2. Metallic bond strength decreases (valence electrons farther from nucleus)
3. Interatomic forces weaken
4. Result: Metals become softer (Li hard, Na soft, K very soft, Rb/Cs can be cut with knife)
5. Also reflected in decreasing melting points: Li(180°C) → Cs(28°C)
ii. Predict the reactivity of potassium towards halogens.
Answer: Potassium reacts violently with all halogens, forming potassium halides.
Detailed Explanation:
1. K is very reactive alkali metal (loses 1 electron easily)
2. Reacts violently with F₂ (most dangerous), vigorously with Cl₂, Br₂, I₂
3. Reactions: 2K + X₂ → 2KX (X = F, Cl, Br, I)
4. Heat produced can cause fire/explosion
5. Reactivity order with halogens: F₂ > Cl₂ > Br₂ > I₂
iii. In reaction Cl₂ + 2NaBr → 2NaCl + Br₂, chlorine acts as oxidizing agent. Which is the reducing agent?
Answer: Br⁻ (from NaBr) is the reducing agent.
Detailed Explanation:
Oxidation: 2Br⁻ → Br₂ + 2e⁻ (Br⁻ loses electrons – reducing agent)
Reduction: Cl₂ + 2e⁻ → 2Cl⁻ (Cl₂ gains electrons – oxidizing agent)
Why Br⁻ reduces Cl₂: Cl₂ has higher oxidizing power than Br₂
Standard electrode potentials: Cl₂/Cl⁻ = +1.36V, Br₂/Br⁻ = +1.09V
iv. Why does iodine exist in solid state at room temperature?
Answer: Iodine has largest molecular size among halogens, leading to strong intermolecular forces.
Detailed Explanation:
1. I₂ molecules are largest (atomic radius 140 pm)
2. Strong London dispersion forces between molecules
3. Higher molecular weight (254 g/mol)
4. Requires more energy to overcome intermolecular forces
5. MP/BP trend: F₂(-220°C) < Cl₂(-101°C) < Br₂(-7°C) < I₂(114°C)
6. I₂ sublimes (solid → gas directly) on gentle heating
Constructed Response Questions:
i. Which noble gas should have the lowest boiling point and why?
Answer: Helium (He) has lowest boiling point (4.2 K).
Detailed Explanation:
1. Smallest atomic size (31 pm)
2. Weakest van der Waals forces (London dispersion forces)
3. Only 2 electrons (monoatomic, spherical)
4. Least polarizable electron cloud
5. Boiling point trend: He(4.2K) < Ne(27K) < Ar(87K) < Kr(120K) < Xe(165K) < Rn(211K)
6. Larger atoms have stronger intermolecular forces → higher BP
ii. Compare the reactions of alkali metals with chlorine.
Answer: Reactivity increases down group: Li (slow with heat) → Na (vigorous) → K (violent) → Rb/Cs (explosive).
Detailed Comparison:
2Li + Cl₂ → 2LiCl
Requires heating
Slow reaction
Forms white LiCl
2Na + Cl₂ → 2NaCl
Vigorous reaction
Yellow flame
Forms white NaCl
2K + Cl₂ → 2KCl
Violent reaction
Lilac flame
Forms white KCl
2Rb + Cl₂ → 2RbCl
Explosive reaction
Requires caution
Forms white halides
All reactions produce: White crystalline salts (metal chlorides) soluble in water
iv. Name any three elements in periodic table which exist as liquids.
Answer: Bromine (Br), Mercury (Hg), Gallium (Ga), Cesium (Cs), Francium (Fr) – near room temp.
Detailed List:
1. Bromine (Br): Only liquid non-metal at RT (melting point -7°C)
2. Mercury (Hg): Liquid metal at RT (melting point -39°C)
3. Gallium (Ga): Melts at 30°C (melts in hand)
Others: Cesium (28°C), Francium (27°C) – but radioactive
Note: Room temperature typically 20-25°C
v. Why are transition elements different from normal elements?
Answer: Transition elements have partially filled d-orbitals giving unique properties.
Detailed Differences:
• Variable oxidation states
• Colored compounds
• Paramagnetic
• Catalytic properties
• Form complexes
• High MP/BP
• Hard, dense metals
• Fixed oxidation states
• Colorless compounds
• Diamagnetic
• No catalytic activity
• Simple ions
• Variable MP/BP
• Variable hardness
Key reason: Partially filled d-orbitals in transition elements allow unique electron configurations and bonding
Descriptive Questions:
i. Explain role of catalytic converter in automobile.
Answer: Converts harmful exhaust gases to less harmful gases using transition metal catalysts.
Detailed Explanation:
Location: In automobile exhaust system
Catalysts: Platinum (Pt), Palladium (Pd), Rhodium (Rh) – transition metals
Reactions catalyzed:
1. Oxidation: 2CO + O₂ → 2CO₂ (carbon monoxide to carbon dioxide)
2. Oxidation: CₓHᵧ + (x + y/4)O₂ → xCO₂ + (y/2)H₂O (hydrocarbons to CO₂ + H₂O)
3. Reduction: 2NO → N₂ + O₂ (nitric oxide to nitrogen + oxygen)
4. Reduction: 2NO₂ → N₂ + 2O₂ (nitrogen dioxide to nitrogen + oxygen)
Structure: Ceramic honeycomb coated with catalyst metals
Benefit: Reduces air pollution (CO, NOₓ, unburnt hydrocarbons)
ii. Why do chemical reactivities of alkali metals increase down group whereas decrease for halogens?
Answer: Opposite trends due to opposite electron behaviors – metals lose electrons, halogens gain electrons.
Detailed Comparison:
Reactivity ↑ down group
Reason: Atomic size ↑
• Outer electron farther from nucleus
• Easier to lose electron (↓ IE)
• More reactive down group
• Cs most reactive metal
Reactivity ↓ down group
Reason: Atomic size ↑
• Outer shell farther from nucleus
• Harder to gain electron (↓ EA)
• Less reactive down group
• F₂ most reactive non-metal
Key Concept: Metals react by losing electrons (easier as size increases), non-metals react by gaining electrons (harder as size increases)
iv. Both alkali metals and halogens are very reactive elements with roles opposite to each other. Explain.
Answer: They are complementary reactive elements – metals lose electrons, halogens gain electrons.
Detailed Explanation:
Alkali Metals (Electron Donors):
1. Have 1 valence electron (ns¹)
2. Low ionization energy (easy to lose electron)
3. Strong reducing agents
4. Form cations (M⁺)
5. Reactivity order: Cs > Rb > K > Na > Li
Halogens (Electron Acceptors):
1. Need 1 electron (ns²np⁵)
2. High electron affinity (easy to gain electron)
3. Strong oxidizing agents
4. Form anions (X⁻)
5. Reactivity order: F₂ > Cl₂ > Br₂ > I₂
Perfect Combination: React violently together to form ionic salts (MX)
v. Why hydrogen bromide is thermally unstable compared to hydrogen chloride?
Answer: H-Br bond is longer and weaker than H-Cl bond due to larger bromine atom.
Detailed Explanation:
Bond Parameters:
H-Cl: Bond length = 127 pm, Bond energy = 431 kJ/mol
H-Br: Bond length = 141 pm, Bond energy = 366 kJ/mol
Reasons for HBr instability:
1. Br atom larger than Cl → longer H-Br bond
2. Weaker bond (lower bond dissociation energy)
3. Br less electronegative than Cl (2.96 vs 3.16)
4. HBr decomposes at lower temperature: 2HBr ⇌ H₂ + Br₂
5. Thermal stability trend: HF > HCl > HBr > HI
Practical effect: HBr stored in dark, sometimes with stabilizer
vii. V₂O₅ catalyst preferred over platinum in oxidation of SO₂. Give reasons.
Answer: Cheaper, resistant to poisoning, efficient at moderate temperatures.
Detailed Reasons:
1. Cost: V₂O₅ much cheaper than Pt (precious metal)
2. Poison Resistance: V₂O₅ not poisoned by arsenic impurities in SO₂
3. Temperature: Works optimally at 400-450°C (industrial conditions)
4. Efficiency: High conversion rate (>99%)
5. Durability: Longer lifespan, less sensitive to impurities
6. Availability: Vanadium more abundant than platinum
Reaction: 2SO₂ + O₂ ⇌ 2SO₃ (Contact Process for H₂SO₄)
Historical: Pt was originally used but failed due to As₂O₃ poisoning
Investigative Questions:
i. Explain role of sodium as heat transfer agent in nuclear power plant. Which property utilized?
Answer: Sodium transfers heat from reactor core to steam generators using its high thermal conductivity and low melting point.
Detailed Explanation:
Role: Liquid sodium coolant in fast breeder reactors
Properties utilized:
1. High thermal conductivity: Efficient heat transfer
2. Low melting point (98°C): Liquid at operating temperatures
3. High boiling point (883°C): Doesn’t vaporize easily
4. Low viscosity: Easy to pump
5. Low neutron absorption: Doesn’t slow neutrons (important for fast reactors)
Advantages over water:
• Higher temperature operation
• No pressure buildup (liquid metal)
• Better heat transfer
Safety concern: Reacts violently with water/air (requires careful handling)
ii. Why and how does lithium behave differently from rest of alkali metals?
Answer: Lithium shows diagonal relationship with magnesium (Group 2) due to small size and high charge density.
Detailed Differences:
• Harder metal
• Higher melting point (180°C)
• Reacts slowly with water
• Forms Li₃N with nitrogen
• Carbonate/bicarbonate unstable
• Forms covalent compounds
• Smallest alkali metal
• Soft metals
• Low melting points
• React vigorously with water
• Don’t form nitrides
• Stable carbonates
• Form ionic compounds
• Larger atomic sizes
Reasons for differences:
1. Smallest atomic radius (152 pm)
2. Highest charge density (charge/size ratio)
3. Strongest polarizing power (distorts electron clouds)
4. Forms more covalent bonds due to small size
5. Similar to magnesium (diagonal relationship in periodic table)
From Chapter Text:
Exercise: How do halogens react with water?
Answer: Halogens react with water in two ways: disproportionation (Cl₂, Br₂) or oxidation (F₂).
Detailed Reactions:
1. Fluorine (F₂): Violent oxidation: 2F₂ + 2H₂O → 4HF + O₂
2. Chlorine (Cl₂): Disproportionation: Cl₂ + H₂O ⇌ HCl + HOCl (hypochlorous acid)
3. Bromine (Br₂): Similar to chlorine but slower: Br₂ + H₂O ⇌ HBr + HOBr
4. Iodine (I₂): Very slight reaction with water
Chlorine water uses: Bleaching, disinfecting (swimming pools)
Reason: Disproportionation – same element oxidized and reduced
Exercise: Keeping in view trends of reactivity in first group, how would they react with oxygen?
Answer: Reactivity with oxygen increases down group: Li (oxide), Na (peroxide), K/Rb/Cs (superoxide).
Detailed Reactions:
Lithium: 4Li + O₂ → 2Li₂O (normal oxide)
Sodium: 2Na + O₂ → Na₂O₂ (sodium peroxide)
Potassium: K + O₂ → KO₂ (potassium superoxide)
Rubidium/Cesium: Similar superoxides
Trend: As size increases, larger anions stabilized
Products: Oxides (O²⁻), Peroxides (O₂²⁻), Superoxides (O₂⁻)
Reactivity: Increases down group (Cs most vigorous)