The Periodic Table — Brief History
The modern periodic table is the result of more than a century of scientific work. Here are the most important milestones to know:
| Scientist | Year | Contribution |
|---|---|---|
| Döbereiner | 1829 | Law of Triads — middle element's properties average those of the other two in a group of three similar elements |
| Newlands | 1863 | Law of Octaves — every 8th element has similar properties (failed beyond Calcium) |
| Mendeleev | 1869 | Arranged ~66 elements by increasing atomic mass; left gaps for undiscovered elements; predicted their properties |
| Moseley | 1913 | Rearranged elements by atomic number — this is the basis of the modern table |
| Seaborg | 1944 | Discovered 10 new elements; separated the actinide series (placed below main body) |
Locating Elements — Groups, Periods & Blocks
Finding the Period
The period number = the highest principal quantum number (n) occupied by electrons.
Finding the Group
Count valence electrons in the outermost sub-shell(s):
| Last sub-shell | Group |
|---|---|
| ns¹ | 1 |
| ns² | 2 |
| ns²np¹ to ns²np⁶ | 13 to 18 |
| ns² + (n-1)d¹⁻¹⁰ (transition) | 3 to 12 (add both) |
The Four Blocks
Trends in Physical Properties
1. Atomic Radius
2. Electronegativity
3. Ionisation Energy (IE)
IE = energy needed to remove an electron from a neutral gaseous atom. The key rule: 1st IE < 2nd IE < 3rd IE (each successive electron is harder to remove because the remaining electrons are more tightly held).
4. Melting & Boiling Points
In the left (metallic) region of the periodic table: melting points decrease down a group (weaker metallic bonding). In the right (non-metal) region (groups 17, 18): melting points increase down a group (larger molecules → stronger van der Waals forces).
5. Density
Generally increases down a group (mass increases faster than volume). In a period, density peaks in the middle (d-block) where atomic mass is high and atomic radius is relatively small.
6. Metallic Character & Electron Affinity
Going down a group: radius ↑, EN ↓, IE ↓, metallic ↑.
Who Are They?
All Group 1 elements have the valence electronic configuration ns¹. They are soft, silvery metals that are stored in oil to prevent reaction with air and moisture. Hydrogen has the same ns¹ configuration but is NOT a Group 1 metal — it is a gas, non-metal in many compounds, and is studied separately.
Physical Properties Trend (down the group)
| Property | Trend ↓ | Reason |
|---|---|---|
| Atomic radius | Increases | More electron shells added |
| Melting/Boiling point | Decreases | Metallic bond weakens (larger atom, less attractive force between nucleus and electrons) |
| 1st Ionisation Energy | Decreases | Valence electron is farther from nucleus; easier to remove |
| Reactivity | Increases | Easier to lose that one valence electron |
| Density | Generally increases | Mass increases faster than volume; Li, Na, K are less dense than water |
Reactions — the Core Chemistry
With Oxygen (air)
Reactivity increases down the group. The type of oxide formed changes:
All oxides dissolve in water to form strongly alkaline (basic) solutions: M₂O + H₂O → 2MOH
With Water
This is why they are called "alkali metals" — they produce alkaline hydroxide solutions with water. Reactivity: Li (slow fizz) < Na (vigorous, hissing) < K (ignites) < Rb < Cs (explosive).
With Halogens
All Group 1 metals react vigorously with halogens to form ionic halides. Reactivity increases down the group.
Anomalous Behaviour of Lithium
Lithium is the smallest Group 1 atom. Its high charge density (charge/size ratio) gives it more covalent character in its compounds. Key differences:
- Li₂O (normal oxide) — not peroxide like Na
- Li₂CO₃ decomposes on heating (unlike Na₂CO₃ which is stable)
- LiNO₃ decomposes like Mg(NO₃)₂ (gives Li₂O + NO₂ + O₂), not LiNO₂ + O₂
- LiOH and Li₂CO₃ are less soluble than other Group 1 counterparts
- Li resembles Mg — this is the diagonal relationship
Oxides & Hydroxides
Oxides: All Group 1 oxides are basic. They dissolve in water to form the corresponding hydroxide (MOH).
Hydroxides: All Group 1 hydroxides are strong bases. They dissolve readily in water. Solubility increases down the group (LiOH least soluble → CsOH most soluble).
Effect of Heat on Nitrates & Carbonates
Sodium → Caesium all give nitrite + O₂ on heating.
(Behaves like Group 2 nitrates)
(Behaves like Group 2 carbonates)
Solubility Trends
| Compound type | Trend down Group 1 |
|---|---|
| Hydroxides (MOH) | Solubility increases |
| Carbonates (M₂CO₃) | Solubility increases |
| Sulphates (M₂SO₄) | Solubility decreases |
Flame Tests
| Element | Flame Colour |
|---|---|
| Lithium (Li) | Deep Red / Crimson |
| Sodium (Na) | Bright Yellow / Orange |
| Potassium (K) | Lilac / Violet |
| Rubidium (Rb) | Red |
| Caesium (Cs) | Blue |
Important Uses
Lithium: Batteries (phones, pacemakers), Li₂CO₃ to toughen glass, alloys for aircraft.
Sodium: NaOH (caustic soda) in soap/paper making; NaCl for seasoning and manufacturing; Na vapour lamps for street lighting; molten Na as coolant in nuclear reactors.
Who Are They?
Valence configuration: ns². All are solid metals at room temperature. They lose 2 electrons to form M²⁺ ions. Called "alkaline earth metals" because their oxides dissolve in water to give alkaline solutions, and were historically called "earths" (they were found in Earth's crust).
Physical Properties Trend
| Property | Trend ↓ the group | Reason |
|---|---|---|
| Atomic radius | Increases | Extra electron shells added |
| Melting/Boiling point | Decreases (generally) | Metallic bond weakens with larger atomic radius |
| 1st Ionisation Energy | Decreases | Larger atom; outer electrons farther from nucleus |
| Reactivity | Increases | Easier to lose 2 electrons as radius grows |
| Electronegativity | Decreases | Larger radius → less ability to attract electrons |
Reactions with Common Substances
| Reactant | Product | Notes |
|---|---|---|
| Cold water | M(OH)₂ + H₂ | Be: no reaction. Mg: no reaction. Ca → Ba: increasingly vigorous |
| Steam | MO + H₂ | Mg reacts with steam: Mg + H₂O(g) → MgO + H₂ |
| Oxygen | MO (mostly) | Ba forms BaO₂ (peroxide) in excess O₂ |
| Halogens (e.g. Cl₂) | MCl₂ | All react to form ionic halides except BeCl₂ (covalent) |
| Dilute HCl | MCl₂ + H₂ | All react to form the chloride salt + hydrogen gas |
| Nitrogen | M₃N₂ | Requires extremely high temperature |
Properties of Compounds
Oxides
Group 2 oxides are basic (ionic). They react with water: MO + H₂O → M(OH)₂. Exception: BeO is amphoteric — it reacts with both acids and bases.
Hydroxides — Solubility Trend
Sulphates — Opposite trend!
Carbonates
Sparingly soluble, and their solubility decreases slightly down the group.
Thermal Stability of Carbonates, Hydroxides & Nitrates
Thermal stability increases down the group. Compounds of smaller ions (top of group) have more covalent character and decompose more easily. BeCO₃ decomposes most easily; BaCO₃ is the most stable.
Anomalous Behaviour of Beryllium
Be is unique because its Be²⁺ ion is very small and highly charged → high polarising power → more covalent character than expected.
As base: Be(OH)₂ + 2H⁺ → Be²⁺ + 2H₂O
As acid: Be(OH)₂ + 2OH⁻ → [Be(OH)₄]²⁻
Dissolves in non-polar organic solvents (unlike other Group 2 chlorides which are ionic). Hydrolyses in water to form Be(OH)₂ + HCl.
Flame Tests for Group 2
| Element | Flame Colour |
|---|---|
| Beryllium / Magnesium | Colourless / Bright White |
| Calcium | Brick Red |
| Strontium | Crimson Red |
| Barium | Apple Green |
Test for Ba²⁺ ions
Add SO₄²⁻ ions (e.g. Na₂SO₄ solution): a white precipitate of BaSO₄ forms, insoluble even in acid.
Important Uses & Limestone Chemistry
CaCO₃ (limestone) →heat ~1000°C→ CaO (quicklime) + CO₂
CaO + H₂O → Ca(OH)₂ (slaked lime / ishwagara)
Ca(OH)₂ is used to reduce soil acidity, in cement making, and as a cheap base in industry.
Cement: Made by strongly heating limestone + clay → "clinker" → mixed with gypsum (CaSO₄·2H₂O).
Other uses: Mg in chlorophyll (photosynthesis); Mg(OH)₂ as antacid medicine; Ca in bones/teeth; Ba compounds in X-ray contrast agents.
Who Are They?
Valence configuration: ns²np¹. These are the first p-block elements. Except for Boron, they are all metals. Boron is a non-metal with very high melting point and low density.
Physical Properties Trend
| Property | Trend ↓ | Reason |
|---|---|---|
| Atomic radius | Increases | Extra electron shells |
| Melting/Boiling point | Decreases generally | Weaker metallic bonding as atoms get larger |
| 1st Ionisation Energy | Decreases | Outer electrons farther from nucleus |
Reactions of Aluminium (the most important element)
With Oxygen
This protective oxide layer on the surface prevents further corrosion — that is why aluminium "resists corrosion" even though it is reactive.
With Hydrochloric Acid
With Concentrated H₂SO₄
Note: Al does NOT react with HNO₃ because the Al₂O₃ layer protects it (passivation).
With NaOH (alkali)
This is why aluminium utensils must NOT be washed with strong alkaline solutions — the alkali dissolves the protective oxide layer and then attacks the metal.
With Halogens
Amphoteric Behaviour of Al₂O₃ and Al(OH)₃
Amphoteric = reacts with both acids AND bases. Al₂O₃ is a key example:
Al(OH)₃ (aluminium hydroxide) shows the same behaviour.
Identification of Al³⁺ Ions in Solution
Anomalous Behaviour of Boron
Boron is a non-metal, unlike all other Group 13 members. Reasons: small size + high nuclear charge/size ratio + high electronegativity.
| Property | Boron (anomalous) | Other Group 13 metals |
|---|---|---|
| Electrical conductivity | Poor conductor | Good conductors |
| Melting/Boiling point | Very high (2075°C) | Much lower |
| Oxide nature | Acidic (B₂O₃) | Amphoteric or basic |
| Compounds | Only covalent | Some ionic compounds |
| Reaction with dilute acid | No reaction | Dissolve in acids → H₂ evolved |
Inert Pair Effect (Thallium)
Heavier members (Ga, In, Tl) can exhibit the +1 oxidation state in addition to +3. This is the inert pair effect: as you go down the group, the s electrons become more tightly held by the nucleus and less available for bonding. Only the p electron is used → oxidation state +1.
Uses of Group 13 Elements
Aluminium: Cooking utensils, window frames, overhead power cables, aircraft alloys, food packaging, firefighter suits (reflects radiant heat), clay for cement/pots.
Boron: Nuclear reactor control rods (absorbs neutrons), borosilicate glass (test tubes, vacuum flasks), semiconductor additive.
Gallium: Semiconductors, LED (light-emitting diodes).
Who Are They?
Valence configuration: ns²np² — 4 valence electrons. This group shows the clearest non-metal → metalloid → metal transition. Carbon (non-metal), Si & Ge (metalloids/semiconductors), Sn & Pb (metals).
Physical Properties
| Property | Trend ↓ | Notes |
|---|---|---|
| Atomic radius | Increases | Extra electron shells |
| Melting/Boiling point | Decreases | Inter-atomic bonds weaken as atoms get larger |
| Density | Increases | Mass increases faster than volume |
| Electronegativity | Decreases | Larger radius → less ability to attract |
| Electrical conductivity | Generally increases | Metallic character increases down group |
Carbon — Special Features
Allotropes of Carbon
Catenation
Carbon's unique ability to form long chains and rings bonded to other carbon atoms (C–C, C=C, C≡C). This is the basis of organic chemistry. Silicon forms Si–O–Si bonds instead of long Si–Si chains.
Oxidation States: +2 and +4
Group 14 elements can form compounds in both +2 and +4 oxidation states. The trend in stability is crucial:
The Inert Pair Effect (Ge, Sn, Pb)
As you go down Group 14, the s electrons become more tightly held by the nucleus and less willing to participate in bonding. Only the 2 p electrons are used → oxidation state +2. With a strong oxidising agent, all 4 electrons can be removed → oxidation state +4.
Reactions of Oxides and Chlorides with Water
| Compound | Reaction with Water | Why? |
|---|---|---|
| CCl₄ | Does NOT hydrolyse | Carbon has no vacant d-orbitals to accommodate the O atom during reaction |
| SiCl₄ | Hydrolyses: SiCl₄ + 2H₂O → SiO₂ + 4HCl | Silicon has vacant d-orbitals → can form 5-coordinate transition state |
| SnCl₄ | Hydrolyses: SnCl₄ + 2H₂O → SnO₂ + 4HCl | Same reason as Si |
| CO₂ | CO₂ + H₂O → H₂CO₃ (weak acid) | Acidic oxide; H₂CO₃ does NOT react with acids |
| SiO₂ | Does NOT react with water | Giant covalent structure |
| SiO₂ | Reacts with HF only: SiO₂ + 4HF → SiF₄ + 2H₂O | HF forms stable Si–F bonds |
| PbO and PbO₂ | Amphoteric — react with both acids and bases | Pb is metallic; its oxide shows intermediate character |
Thermal Stability of Halides
Thermal stability of tetrahalides decreases down the group: CCl₄ > SiCl₄ > GeCl₄ > SnCl₄ > PbCl₄.
PbCl₄ decomposes readily at room temperature: PbCl₄ → PbCl₂ + Cl₂.
PbBr₄ and PbI₄ do not exist because Br and I are not strong enough oxidising agents to oxidise Pb to +4.
Chemical Tests for Group 14 Ions
| Ion | Reagent | Observation |
|---|---|---|
| Pb²⁺ | HCl | White precipitate of PbCl₂ (dissolves on heating) |
| Pb²⁺ | H₂S | Black precipitate of PbS |
| Pb²⁺ | KI | Yellow precipitate of PbI₂ |
| Pb²⁺ | OH⁻ (NaOH) | White precipitate of Pb(OH)₂, dissolves in excess |
| Sn²⁺ | H₂S | Brown precipitate of SnS |
| CO₃²⁻ | Dilute HCl | Colourless CO₂ gas → turns limewater milky |
| HCO₃⁻ vs CO₃²⁻ | MgSO₄ solution | CO₃²⁻ → white precipitate; HCO₃⁻ → no change |
Diagonal Relationship (Li/Mg, Be/Al, B/Si)
Pairs of elements diagonally adjacent in periods 2 and 3 show similar properties because the decrease in size going right across a period and increase in size going down a group cancel out, giving similar charge/size ratios.
| Diagonal Pair | Key Similarities |
|---|---|
| Li / Mg | Both form only normal oxides (not peroxides). Li₂CO₃ & MgCO₃ both decompose on heating. Hydroxides decompose on heating. Hydroxides are weak alkalis. |
| Be / Al | Both oxides (BeO, Al₂O₃) are amphoteric. Both chlorides are covalent. Both rendered passive by conc. HNO₃. Similar atomic radii. |
| B / Si | Both non-metals with giant covalent structures. Both form weak acids (H₃BO₃, H₂SiO₃). Both have acidic oxides soluble in alkali. Both form hydrides that ignite in air. |
Uses of Group 14 Elements
Carbon: Fuel (charcoal/coal), reducing agent in metallurgy, lubricant (graphite), electrodes, pencil leads, diamond in cutting tools & jewellery, radiocarbon dating (C-14).
Silicon: Semiconductors in electronics (transistors, chips), silicates in glass & ceramics, sand in glass production, silicone implants.
Tin: Tin-plating steel cans to prevent corrosion (tinned food cans).
Lead: Lead-acid batteries, radiation shielding (X-rays), lining of sulphuric acid vessels.
Cassiterite (SnO₂) is the tin ore mined in Rwanda.
Who Are They?
Valence configuration: ns²np³ — 5 valence electrons (3 unpaired p electrons). Common oxidation states: +3 and +5. Nitrogen and phosphorus are non-metals; arsenic and antimony are metalloids; bismuth is a metal.
Physical Properties Trend
| Property | Trend ↓ | Notes |
|---|---|---|
| Atomic radius | Increases | Extra shells; larger atoms |
| Metallic character | Increases | N & P: non-metals → As & Sb: metalloids → Bi: metal |
| 1st Ionisation Energy | Decreases | Outer electrons farther from nucleus; easier to remove |
| Electronegativity | Decreases | Larger radius → less ability to attract electrons |
| Melting point | Increases generally (in solids) | N is a gas; solid members have higher MP going down |
Why Is Nitrogen So Unreactive? (Inertness of N₂)
Nitrogen gas exists as N≡N, a triple bond. The bond energy of N≡N is 942 kJ/mol — one of the strongest bonds in chemistry. Compared to: N–N single bond (247 kJ/mol), O=O double bond (498 kJ/mol).
Valency of 3 and 5 in Group 15
All Group 15 elements can form 3 bonds (using the 3 unpaired p electrons).
With the exception of nitrogen, they can also form 5 bonds by using vacant d-orbitals (sp³d hybridisation).
Reactions of Group 15 Elements
With Oxygen
With Hydrogen
Stability of hydrides decreases down the group: NH₃ most stable → BiH₃ least stable. Basicity also decreases: NH₃ > PH₃ > AsH₃ > SbH₃ > BiH₃.
With Chlorine
With Metals
Ammonia (NH₃) — Laboratory and Industrial
Laboratory Preparation
NH₃ is less dense than air → collected by upward delivery. Dried over quicklime (CaO), NOT concentrated H₂SO₄ (it reacts with NH₃).
Industrial Production: Haber Process
| Condition | Value | Reason |
|---|---|---|
| Temperature | 450–500°C | Compromise: lower T gives better yield but too slow; higher T is faster but very low yield |
| Pressure | 200–1000 atm | High pressure favours the product side (fewer moles of gas) |
| Catalyst | Finely divided iron (Fe) + Al₂O₃ promoter | Increases reaction rate without affecting equilibrium position |
Chemical Properties of Ammonia
Uses of Ammonia
Fertilisers (main use — e.g. ammonium nitrate NH₄NO₃), livestock feed supplement, paper manufacturing, food fermentation, manufacturing of nitric acid.
Nitric Acid (HNO₃) — Laboratory and Industrial
Laboratory Preparation
Industrial Production: Ostwald Process
Three steps, starting from ammonia:
Reactions of Nitric Acid with Metals
| Metal | Conc. HNO₃ | Dilute HNO₃ |
|---|---|---|
| Copper (Cu) | Cu + 4HNO₃(conc.) → Cu(NO₃)₂ + 2NO₂ + 2H₂O (brown gas) | 3Cu + 8HNO₃(dil.) → 3Cu(NO₃)₂ + 2NO + 4H₂O (colourless) |
| Lead (Pb) | Pb + 4HNO₃(conc.) → Pb(NO₃)₂ + 2NO₂ + 2H₂O | — |
| Magnesium (Mg) | — | With very dilute (1%): Mg + 2HNO₃ → Mg(NO₃)₂ + H₂ |
| Zinc (Zn) | 4Zn + 10HNO₃(conc.) → 4Zn(NO₃)₂ + NH₄NO₃ + 3H₂O | — |
Reactions with Non-metals (concentrated HNO₃)
Uses of Nitric Acid
Production of fertilisers (ammonium nitrate), manufacturing dyes, artificial fibres, drugs, explosives (TNT).
Phosphorus — Key Points
Allotropes of Phosphorus
Phosphoric Acid (H₃PO₄)
Produced when P₂O₅ reacts with water: P₄O₁₀ + 6H₂O → 4H₃PO₄. A triprotic weak acid with three ionisable H atoms. Used in fertilisers, detergents, rust removal, and food flavouring (cola drinks).
Environmental Impact
• Contribute to photochemical smog and acid rain (HNO₃ in rain).
• N₂O (laughing gas) is a potent greenhouse gas.
• Haber process burns fossil fuels → CO₂ → global warming.
• Ostwald process emits NO and NO₂ → must be carefully controlled.