6.1
Occurrence & Physical Properties of Group 1 Elements
Occurrence
Group 1 elements (alkali metals: Li, Na, K, Rb, Cs, Fr) never occur free in nature due to their extreme reactivity. They are found as ionic compounds in minerals and seawater:
- Lithium: spodumene (LiAlSi₂O⁶), lepidolite
- Sodium: rock salt (NaCl, halite), trona (Na₂CO₃), seawater
- Potassium: sylvite (KCl), carnallite, seawater
- Rubidium & Caesium: trace amounts in minerals
- Francium: radioactive, extremely rare (produced by nuclear reactions)
| Element | Symbol | Z | Config | M.p. (°C) | Density (g/cm³) | Flame colour |
|---|---|---|---|---|---|---|
| Lithium | Li | 3 | [He]2s¹ | 181 | 0.53 | Crimson red |
| Sodium | Na | 11 | [Ne]3s¹ | 98 | 0.97 | Yellow/orange |
| Potassium | K | 19 | [Ar]4s¹ | 64 | 0.86 | Lilac/violet |
| Rubidium | Rb | 37 | [Kr]5s¹ | 39 | 1.53 | Red-violet |
| Caesium | Cs | 55 | [Xe]6s¹ | 28 | 1.87 | Blue-violet |
Trends in Physical Properties Down Group 1
- Melting/boiling point: decreases — weakening metallic bond (larger atomic radius, cations further from electron sea)
- Density: generally increases — increasing atomic mass outweighs increase in atomic volume (though Li, Na, K are all less dense than water, density < 1 g/cm³)
- Atomic radius: increases — new shell added each period
- Ionic radius (M⁺): increases — same trend as atomic radius
- First ionisation energy: decreases — outer electron further away and more shielded
- Electronegativity: decreases
- Reactivity: increases — easier to lose outer electron
All alkali metals are softThey can all be cut with a knife and show a shiny metallic surface when freshly cut (which quickly tarnishes due to reaction with air). All are stored under oil or in inert gas to prevent oxidation.
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Section 6.1 — Group 1 Properties
10 QuestionsQ1 of 10
Group 1 metals are stored under oil because:
Group 1 metals react with water vapour and O₂ in air. Oil prevents contact. Li is less reactive than Na or K.
Q2 of 10
Down Group 1, reactivity with water:
Li: slow fizzing. Na: rapid, melts, moves. K: ignites (purple flame). Rb/Cs: explosive. Reactivity increases down group.
Q3 of 10
The equation for Na reacting with water is:
2Na + 2H₂O → 2NaOH + H₂↑. Product is strongly alkaline NaOH (pH 13-14) + H₂ gas (burns with pop).
Q4 of 10
Why does reactivity increase down Group 1?
Down group: extra shells → outer electron further → weaker attraction → lower IE → easier to lose → more reactive.
Q5 of 10
When Li burns in excess O₂, the product is:
Li forms normal oxide Li₂O. Na forms peroxide Na₂O₂. K, Rb, Cs form superoxides (MO₂). Li⁺ too small to stabilise larger peroxide/superoxide anions.
Q6 of 10
The hydroxides of Group 1 are:
MOH → M⁺ + OH⁻ (fully ionised). Strong bases. Solubility and basicity increase down group.
Q7 of 10
Group 1 nitrates decompose on heating to give:
Na/K/Rb/CsNO₃ → metal nitrite + O₂. LiNO₃ → Li₂O + NO₂ + O₂ (like Group 2 nitrates — Li⁺ behaves like Mg²⁺: diagonal relationship).
Q8 of 10
Which Group 1 flame colour is tested through cobalt blue glass?
K gives lilac flame but Na contamination (yellow) masks it. Cobalt blue glass absorbs yellow sodium light → K's lilac visible.
Q9 of 10
The trend in melting points down Group 1:
Down Group 1: metallic bond weakens (fewer valence electrons per larger atom, larger spacing) → lower melting points. Li (180°C) > Na (98°C) > K (63°C) > Rb (39°C) > Cs (28°C).
Q10 of 10
Group 1 carbonates: which is the most thermally stable?
Thermal stability of carbonates increases down Group 1. Larger cations (lower charge density) polarise CO₃²⁻ less → more stable. Cs₂CO₃ most stable. Li₂CO₃ decomposes most easily.
6.2
Reactivity of Group 1 Elements with O₂, H₂O & Halogens
Reactions with Oxygen
All alkali metals react vigorously with oxygen when exposed to air. The products depend on the metal:
| Metal | Product | Type | Equation |
|---|---|---|---|
| Lithium | Li₂O (lithium oxide) | Normal oxide | 4Li + O₂ → 2Li₂O |
| Sodium | Na₂O₂ (sodium peroxide) | Peroxide | 2Na + O₂ → Na₂O₂ |
| Potassium | KO₂ (potassium superoxide) | Superoxide | K + O₂ → KO₂ |
| Rubidium, Caesium | Superoxides (MO₂) | Superoxide | M + O₂ → MO₂ |
The trend from oxide → peroxide → superoxide reflects increasing ionic radius: larger M⁺ ions stabilise larger anions (O²⁻ < O₂²⁻ < O₂⁻).
Reactions with Water
All Group 1 metals react vigorously with water to produce metal hydroxide + hydrogen gas. Reactivity increases down the group.
General equation:
2M(s) + 2H₂O(l) → 2MOH(aq) + H₂(g)
Li: 2Li + 2H₂O → 2LiOH + H₂ (fizzes steadily, sinks)
Na: 2Na + 2H₂O → 2NaOH + H₂ (melts to a ball, moves rapidly)
K: 2K + 2H₂O → 2KOH + H₂ (ignites H₂, lilac flame)
Rb: 2Rb + 2H₂O → 2RbOH + H₂ (explodes)
Cs: 2Cs + 2H₂O → 2CsOH + H₂ (explosive)
The reaction is exothermic. Down the group: lower IE → easier to lose e⁻ → more vigorous/faster reaction.
Reactions with Halogens
All Group 1 metals react directly with halogens to form ionic metal halides (MX):
General: 2M + X₂ → 2MX
2Na + Cl₂ → 2NaCl (sodium chloride)
2K + Br₂ → 2KBr (potassium bromide)
2Li + F₂ → 2LiF (lithium fluoride)
2Na + I₂ → 2NaI (sodium iodide)
Reactivity of alkali metal increases down the group. Reactivity of halogen decreases down the group (F₂ > Cl₂ > Br₂ > I₂).
6.3
Properties of Group 1 Oxides and Hydroxides
Oxides
Group 1 oxides, peroxides, and superoxides all react with water to form alkaline solutions:
Normal oxide + water:
Li₂O + H₂O → 2LiOH
Peroxide + water (produces H₂O₂):
Na₂O₂ + H₂O → 2NaOH + ½O₂
(or: Na₂O₂ + H₂O → 2NaOH + H₂O₂ which decomposes)
Superoxide + water:
4KO₂ + 2H₂O → 4KOH + 3O₂
Na₂O₂ is used in self-contained breathing apparatus (rebreathers) because it both absorbs CO₂ and releases O₂: 2Na₂O₂ + 2CO₂ → 2Na₂CO₃ + O₂.
Hydroxides
All Group 1 hydroxides (MOH) are strong alkalis — they fully dissociate in water:
NaOH(s) → Na⁺(aq) + OH⁻(aq)
KOH(s) → K⁺(aq) + OH⁻(aq)
Basicity increases down the group: LiOH is the least basic (slightly less soluble, weakest alkali of the group); CsOH is the strongest alkali of all known hydroxides. Basicity increases because M–OH bond weakens (M⁺ is larger → less polarising → less tendency to polarise OH⁻ → OH⁻ released more easily).
Hydroxides react with acids:
NaOH + HCl → NaCl + H₂O
KOH + H₂SO₄ → K₂SO₄ + H₂O (after balancing: 2KOH + H₂SO₄ → K₂SO₄ + 2H₂O)
Hydroxides react with CO₂:
2NaOH + CO₂ → Na₂CO₃ + H₂O (excess NaOH)
NaOH + CO₂ → NaHCO₃ (limited NaOH)
6.4
Effect of Heat on Group 1 Carbonates and Nitrates
Thermal Decomposition of Carbonates
Group 1 carbonates are generally thermally stable (unlike Group 2 carbonates which easily decompose). This is because large M⁺ ions do not strongly polarise the CO₃²⁻ ion.
Li₂CO₃ decomposes on heating (Li⁺ is small, more polarising):
Li₂CO₃ → Li₂O + CO₂
Na₂CO₃, K₂CO₃, Rb₂CO₃, Cs₂CO₃:
Stable — do NOT decompose on heating with a Bunsen burner.
(They require extremely high temperatures.)
Li₂CO₃ is anomalous (decomposes easily) because Li⁺ is very small and highly polarising, distorting the CO₃²⁻ ion. This is similar to the diagonal relationship (Li resembles Mg).
Thermal Decomposition of Nitrates
Group 1 nitrates decompose on heating, but the products differ:
LiNO₃ decomposes like Group 2 nitrates (Li⁺ is small, polarising):
4LiNO₃ → 2Li₂O + 4NO₂ + O₂
All other Group 1 nitrates (NaNO₃, KNO₃, etc.) give nitrite + oxygen:
2NaNO₃ → 2NaNO₂ + O₂
2KNO₃ → 2KNO₂ + O₂
The difference: Li⁺ is small and polarising enough to break the N–O bond further, producing NO₂ and Li₂O. Larger alkali metal cations only remove one O per NO⁻, giving NO₂⁻ (nitrite).
6.5
Solubility of Group 1 Compounds
Most Group 1 compounds are highly soluble in water because:
- Low charge density of M⁺ ions: large radius → low polarising power → predominantly ionic compounds → good solubility
- Hydration energy released on dissolving compensates for (and often exceeds) the lattice energy
| Compound | Solubility trend | Notable exception |
|---|---|---|
| Hydroxides (MOH) | Increases down group (LiOH slightly less soluble; CsOH very soluble) | LiOH is less soluble than other Group 1 hydroxides |
| Carbonates (M₂CO₃) | All soluble; increases down group | Li₂CO₃ is less soluble (anomalous — diagonal relationship with Mg) |
| Halides (MX) | All very soluble | LiF is sparingly soluble (high lattice energy due to small Li⁺ + small F⁻) |
| Sulfates (M₂SO₄) | All soluble | No significant exceptions in Group 1 |
| Nitrates (MNO₃) | All very soluble | — |
Lithium anomalies (diagonal relationship with Mg)Li⁺ resembles Mg²⁺ more than it resembles Na⁺. Both have small ionic radii and high charge density. Similarities: Li₂CO₃ decomposes on heating (like MgCO₃); LiF and Li₂CO₃ are sparingly soluble (like MgF₂ and MgCO₃); LiNO₃ decomposes to give NO₂ (like Mg(NO₃)₂).
6.6
Flame Tests for Li⁺, Na⁺ & K⁺
Procedure and Results
Flame tests detect metal ions in compounds by their characteristic emission spectra in the visible region. The technique:
- Clean a nichrome (or platinum) wire loop with concentrated HCl, then test in flame.
- Dip the wire in the test compound (solid or solution).
- Hold in a roaring blue Bunsen flame and observe the colour.
| Ion | Flame Colour | Notes |
|---|---|---|
| Li⁺ | Crimson red | Distinctive deep red; persistent |
| Na⁺ | Persistent yellow/orange | Very intense; can mask other ions |
| K⁺ | Lilac/violet | Best observed through blue cobalt glass (to filter out yellow from Na⁺ traces) |
| Rb⁺ | Red-violet | — |
| Cs⁺ | Blue-violet | — |
Why Flame Tests Work
The heat of the flame provides energy to promote electrons in the metal ion to higher energy levels (excited states). As the electrons return to the ground state, they emit photons of specific energies corresponding to specific visible wavelengths. Each element has a unique set of energy levels → unique emission spectrum → characteristic colour. This is the atomic emission spectrum applied to qualitative analysis.
6.7
Uses of Group 1 Elements & Their Compounds
| Substance | Uses |
|---|---|
| Lithium (Li) | Lithium-ion batteries (mobile phones, EVs); alloys (lightweight aerospace materials with Al); lithium grease (lubricant) |
| LiCl / Li₂CO₃ | Treatment of bipolar disorder (mood stabiliser); ceramics and glass manufacture; air conditioning (hygroscopic) |
| Sodium (Na) | Coolant in fast breeder nuclear reactors; street lighting (sodium vapour lamps); manufacture of titanium (Na reduces TiCl₄) |
| NaCl | Table salt; food preservation; manufacture of Cl₂, NaOH, Na₂CO₃ (chlor-alkali industry); de-icing roads |
| NaOH | Manufacture of soap, paper, textiles (rayon); drain cleaner; aluminium extraction; food processing |
| Na₂CO₃ | Glass manufacture; water softening; washing soda; paper industry |
| NaHCO₃ | Baking powder; antacid (stomach acid neutralisation); fire extinguishers |
| KNO₃ | Fertiliser; gunpowder (with C and S); food preservation (curing meat) |
| KOH | Manufacture of soft soap; alkaline batteries; CO₂ absorber in laboratory |
| Potassium (K) | K⁺ essential for nerve function and plant growth; sylvite (KCl) mined for fertiliser |
| Na₂O₂ | Bleaching agent; oxygen source in rebreathers (submarines, mines) |
6.8
Hydrogen
Position of Hydrogen in the Periodic Table
Hydrogen (H, Z=1, configuration 1s¹) is usually placed in Group 1 but does not truly belong there. It has properties that make it unique:
| Property | Like Group 1? | Unlike Group 1? |
|---|---|---|
| Electronic config | ns¹ (1s¹) — one outer electron | No inner shells; very small atom |
| Forms +1 ions (H⁺) | Yes — like alkali metal cations | H⁺ is a bare proton (no electrons); unique in chemistry |
| Forms –1 ions (H⁻, hydride) | No — alkali metals do not form anions | H can gain one electron like a halogen (resembles Group 17) |
| Diatomic gas | No — alkali metals are solid metals | H₂ is a gas; forms covalent bonds (like non-metals) |
| Electronegativity | Low (2.2) compared to EN of F | Higher than alkali metals; forms polar covalent bonds |
| Reactivity | Reacts with non-metals, some metals | Does not react with water at room temperature (unlike alkali metals) |
Chemical Properties of Hydrogen
Reaction with metals (reducing agent):
H₂ + 2Na → 2NaH (sodium hydride; H is −1, ionic)
H₂ + Ca → CaH₂ (calcium hydride; ionic)
Reaction with non-metals:
H₂ + F₂ → 2HF (explosive at room temp)
H₂ + Cl₂ → 2HCl (in UV light)
H₂ + O₂ → 2H₂O (combustion)
3H₂ + N₂ → 2NH₃ (Haber process, catalyst, high T and P)
Hydrogen as a fuel: Burns cleanly producing only water. Potential "hydrogen economy" fuel for vehicles. Currently mostly produced from natural gas (steam reforming), but research into electrolysis of water using renewable energy.
Hydrogen isotopes¹H (protium, 99.985%), ²H (deuterium, heavy hydrogen, D), ³H (tritium, T, radioactive). D₂O is heavy water used as a moderator in some nuclear reactors. D has 1 proton + 1 neutron; T has 1 proton + 2 neutrons.
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Section 6.2 — Group 1 Compounds
10 QuestionsQ1 of 10
NaCl is used in which industrial process?
Electrolysis of brine (NaCl(aq)) → Cl₂ (anode), H₂ (cathode), NaOH (solution). Chlor-alkali industry: major source of Cl₂ and NaOH.
Q2 of 10
Na₂CO₃ (washing soda) is used for:
Na₂CO₃ (washing soda): water softening (removes Ca²⁺/Mg²⁺ as insoluble carbonates), glass manufacture, cleaning.
Q3 of 10
NaHCO₃ (baking soda) decomposes when heated to give:
2NaHCO₃ → Na₂CO₃ + H₂O + CO₂. CO₂ causes baked goods to rise. Used in baking powder and as antacid.
Q4 of 10
Which Group 1 ion gives a yellow/orange flame test?
Na⁺: characteristic bright yellow-orange flame. Electron excitation in Na → emission at 589 nm (yellow). Very sensitive — even trace amounts visible.
Q5 of 10
KNO₃ is used in:
KNO₃ (saltpetre/nitre): food preservative (prevents Clostridium botulinum), gunpowder component (oxidising agent), fertiliser.
Q6 of 10
LiOH is used in:
LiOH: lightweight, absorbs CO₂ efficiently. Used in life-support systems. Li is lightest Group 1 metal — LiOH is lightest hydroxide.
Q7 of 10
The 'diagonal relationship': Li resembles Mg because:
Diagonal relationship: elements diagonally adjacent in periods 2-3 often have similar properties. Li/Mg: both form normal oxides, both carbonates decompose on heating, LiNO₃ decomposes like Mg(NO₃)₂.
Q8 of 10
Group 1 metals all have oxidation state +1 in compounds because:
Group 1: 1 valence electron. Very low IE₁. Lose 1 electron readily → M⁺. IE₂ is enormously high (inner shell) → +2 not formed.
Q9 of 10
Which compound dissolves in water to give strongly alkaline solution?
KOH → K⁺ + OH⁻. Strong alkali. Metal hydroxides of Group 1 are strong bases. Other salts (LiCl, Na₂SO₄, NaNO₃) give neutral solutions.
Q10 of 10
The reaction of Na with ethanol gives:
Na + C₂H₅OH → C₂H₅ONa + ½H₂. Similar to Na + H₂O but slower (ethanol less polar, OH less acidic than water). Product is sodium ethoxide — a strong base.
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Exercises
- State and explain the trend in reactivity of Group 1 elements with water as you go down the group from Li to Cs. Write equations for Li and K with water.
Reactivity increases down the group: Li fizzes gently; Na moves rapidly and melts; K ignites the hydrogen; Rb and Cs explode. This is because ionisation energy decreases down the group (outer electron is in a higher shell, further from nucleus and more shielded) → easier to lose one electron → more reactive. Equations:
2Li + 2H₂O → 2LiOH + H₂ (steady effervescence)
2K + 2H₂O → 2KOH + H₂ (ignites H₂, lilac flame) - Compare the products of heating NaNO₃ and LiNO₃. Write balanced equations and explain why the products differ.
NaNO₃ (and other Group 1 nitrates except LiNO₃): gives nitrite + oxygen.
2NaNO₃ → 2NaNO₂ + O₂
LiNO₃: behaves like a Group 2 nitrate (diagonal relationship with Mg).
4LiNO₃ → 2Li₂O + 4NO₂ + O₂
Reason: Li⁺ is very small and has a high charge density → highly polarising → distorts the NO⁻ ion more strongly → breaks the N–O bond further, producing NO₂ gas and Li₂O. Larger Group 1 cations have lower polarising power → only remove one oxygen → give stable NO₂⁻ (nitrite). - Explain why Li₂CO₃ decomposes on heating but Na₂CO₃ does not. What is this similarity to Mg called, and why does it occur?
Li₂CO₃ → Li₂O + CO₂ (on heating). Li⁺ is very small (ionic radius 76 pm) and has a high charge density → polarises the CO₃²⁻ ion strongly → distorts the ion enough to release CO₂. Na⁺ (ionic radius 102 pm) is much larger with lower charge density → does not polarise CO₃²⁻ sufficiently → Na₂CO₃ is thermally stable. This resemblance between Li and Mg (both decompose their carbonates, both have less soluble hydroxides and fluorides) is called the diagonal relationship. It occurs because Li⁺ and Mg²⁺ have similar charge density and ionic radius ratios, giving similar polarising power despite being in different groups.
- Describe the flame test procedure for distinguishing between solutions containing Na⁺ and K⁺. What practical precaution is needed and why?
Procedure: clean a nichrome wire with conc. HCl; test by holding in flame (no colour). Dip wire in test solution; hold in roaring flame; observe colour. Na⁺ gives a persistent yellow/orange flame; K⁺ gives a lilac/violet flame.
Precaution: observe the K⁺ flame through a blue cobalt glass filter. Na⁺ produces an extremely intense yellow colour that can mask the lilac of K⁺ even if only traces of sodium are present. The cobalt glass absorbs/transmits yellow and blue wavelengths, filtering out the yellow of Na so the lilac of K becomes visible. Also, the nichrome wire must be thoroughly cleaned with conc. HCl between tests to avoid cross-contamination. - Explain the anomalous position of hydrogen in the periodic table. In what ways does H resemble Group 1, and in what ways does it differ?
Resemblances to Group 1: (1) Electronic config 1s¹ — one outer electron like alkali metals. (2) Forms H⁺ ions (proton) by losing its electron, similar to M⁺ cations. (3) Reacts with non-metals to form binary compounds (HCl, HF, H₂O etc.). (4) EN~2.2, lower than most non-metals.
Differences from Group 1: (1) H⁺ is a bare proton (no electrons, no core), completely unlike Na⁺ or K⁺ which have electron cores. (2) H can form H⁻ (hydride ion) by gaining an electron — alkali metals never form anions. (3) H₂ is a gas (diatomic covalent molecule); alkali metals are solid metals. (4) H does not react with water at room temperature (alkali metals react vigorously). (5) H is a non-metal; alkali metals are reactive metals. - Write balanced equations for the reactions of sodium with: (a) oxygen (excess), (b) water, (c) chlorine, (d) HCl(aq).
(a) 2Na + O₂ → Na₂O₂ (sodium peroxide, in excess O₂)
(b) 2Na + 2H₂O → 2NaOH + H₂
(c) 2Na + Cl₂ → 2NaCl
(d) Na + HCl → NaCl + ½H₂ (or: 2Na + 2HCl → 2NaCl + H₂)
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Multiple Choice Quiz — 25 Questions
Unit 6 Quiz
Select one answer per questionQ1
Reactivity of Group 1 metals increases down the group because:
Down Group 1: IE₁ decreases because outer electron is in a successively higher shell, farther from nucleus and more shielded → easier to lose → greater reactivity.
Q2
The product when sodium burns in excess oxygen is:
Na burns in excess O₂ to give sodium peroxide Na₂O₂. Li gives normal oxide Li₂O. K and heavier metals give superoxides MO₂.
Q3
The equation for potassium reacting with water is:
2K + 2H₂O → 2KOH + H₂. All Group 1 metals react with water to give MOH + H₂.
Q4
The flame test colour for sodium is:
Na⁺ gives a very intense, persistent yellow/orange flame. Li = crimson; K = lilac; Cu = blue-green.
Q5
Why is blue cobalt glass used when performing a flame test for potassium?
Sodium's intense yellow easily masks potassium's lilac. Cobalt blue glass absorbs yellow wavelengths, allowing the lilac of K⁺ to be seen clearly.
Q6
When heated, NaNO₃ produces:
2NaNO₃ → 2NaNO₂ + O₂. All Group 1 nitrates (except LiNO₃) give nitrite + oxygen. LiNO₃ gives Li₂O + NO₂ + O₂.
Q7
Li₂CO₃ decomposes on heating while Na₂CO₃ does not because:
Li⁺ has a very small ionic radius → high charge density → strongly polarises the CO₃²⁻ anion → distorts it enough to release CO₂. Na⁺ is larger, less polarising → Na₂CO₃ is thermally stable.
Q8
Which Group 1 compound is used in self-contained breathing apparatus?
Na₂O₂ (sodium peroxide) absorbs CO₂ and releases O₂: 2Na₂O₂ + 2CO₂ → 2Na₂CO₃ + O₂. Used in rebreathers for submarines and firefighters.
Q9
Which property of Group 1 metals decreases down the group?
Melting point decreases down Group 1 (Li 181°C → Cs 28°C) because the metallic bond weakens as atomic radius increases. Atomic radius, density, and reactivity all increase down the group.
Q10
The diagonal relationship means Li resembles:
Li (Period 2, Group 1) resembles Mg (Period 3, Group 2) — the diagonal neighbour. Both have small, highly polarising cations. This is the Li–Mg diagonal relationship.
Q11
Hydrogen differs from Group 1 metals because it:
H can gain an electron to form H⁻ (hydride) like a halogen. Alkali metals only lose electrons. H is a diatomic non-metal gas — very different from solid metallic alkali metals.
Q12
The product of reacting potassium with oxygen (in excess) is:
K (and heavier alkali metals Rb, Cs) react with excess O₂ to form superoxides MO₂. K + O₂ → KO₂ (potassium superoxide).
Q13
NaHCO₃ is used as an antacid because:
NaHCO₃ + HCl → NaCl + H₂O + CO₂. The HCO⁻ ion neutralises H⁺ → reduces excess acidity in the stomach.
Q14
Which of the following is sparingly soluble in water (unlike most Group 1 compounds)?
Li₂CO₃ is sparingly soluble (diagonal relationship with MgCO₃). Li⁺ is small with high charge density → high lattice energy → less easily dissolved.
Q15
Which Group 1 element has the lowest first ionisation energy?
IE₁ decreases down Group 1. Cs (Z=55) has the outermost electron in n=6 → furthest from nucleus, most shielded → lowest IE₁ → most reactive alkali metal.
Q16
The reaction of Na₂O₂ with CO₂ is important in rebreathers because it:
2Na₂O₂ + 2CO₂ → 2Na₂CO₃ + O₂. Exhaled CO₂ is absorbed; O₂ is released for breathing. Also: 2Na₂O₂ + 2H₂O → 4NaOH + O₂.
Q17
Flame tests work because:
Flame heat excites electrons to higher energy levels. Returning to ground state, electrons emit photons of specific frequencies → characteristic colours. This is atomic emission spectroscopy in action.
Q18
Which use best describes Na₂CO₃?
Na₂CO₃ (washing soda) is essential in glass manufacture (with SiO₂ and CaCO₃) and as a water softener (precipitates Ca²⁺/Mg²⁺ as carbonates). Na (liquid) is used as a reactor coolant; KNO₃ is in gunpowder.
Q19
The reaction of NaOH with CO₂ (excess NaOH) gives:
Excess NaOH + CO₂: 2NaOH + CO₂ → Na₂CO₃ + H₂O. Limited NaOH: NaOH + CO₂ → NaHCO₃.
Q20
Hydrogen is sometimes placed in Group 17 because it:
H can gain an electron to achieve a full shell (1s², like He). H⁻ (hydride ion) resembles halide ions (F⁻, Cl⁻) in being a −1 anion. Hence H is sometimes placed in Group 17 as well as Group 1.
Q21
All Group 1 elements are stored under oil because:
Alkali metals are extremely reactive with water and oxygen in air → stored under mineral oil or dry inert gas (Ar/N₂) to prevent oxidation and reaction with atmospheric moisture.
Q22
Which alkali metal has the highest melting point?
M.p. decreases down Group 1. Li (181°C) has the highest m.p. in the group because it has the smallest atomic radius → strongest metallic bond (cations closest to electron sea).
Q23
Sodium hydroxide reacts with hydrochloric acid to give:
NaOH + HCl → NaCl + H₂O. This is a simple neutralisation reaction (acid + base → salt + water).
Q24
KNO₃ is used in fertilisers because:
KNO₃ provides both K (for enzyme function, water regulation) and N (for protein/chlorophyll synthesis) — two of the three macronutrients (N, P, K) essential for plant growth.
Q25
Lithium batteries are widely used in portable electronics because lithium:
Li has the lowest atomic mass of all metals (6.94 g/mol) → batteries store more charge per unit mass → high energy density → lightweight batteries for phones, laptops, and electric vehicles.
📝
Unit Test — 50 Marks
Section A — Short Answer
20 marksQ1 [4 marks]
State and explain the trend in reactivity of Group 1 metals with water as you descend from Li to Cs. Write equations for the reactions of Li and K with water, and describe the observations for each. [4]
Trend: reactivity increases down the group. [1]
Explanation: IE₁ decreases as atomic radius increases and shielding increases → outer electron is easier to lose → more vigorous electron transfer to water → greater reactivity. [1]
2Li + 2H₂O → 2LiOH + H₂. Observations: fizzes steadily; Li moves on surface; sinks. Solution turns litmus blue (alkaline). [1]
2K + 2H₂O → 2KOH + H₂. Observations: very vigorous; K moves rapidly; H₂ ignites with a lilac flame; may explode. Solution strongly alkaline. [1]
Explanation: IE₁ decreases as atomic radius increases and shielding increases → outer electron is easier to lose → more vigorous electron transfer to water → greater reactivity. [1]
2Li + 2H₂O → 2LiOH + H₂. Observations: fizzes steadily; Li moves on surface; sinks. Solution turns litmus blue (alkaline). [1]
2K + 2H₂O → 2KOH + H₂. Observations: very vigorous; K moves rapidly; H₂ ignites with a lilac flame; may explode. Solution strongly alkaline. [1]
Q2 [4 marks]
Compare the products of thermal decomposition of NaNO₃ and LiNO₃. Write balanced equations for each and explain why the products differ in terms of the polarising power of the cation. [4]
2NaNO₃ → 2NaNO₂ + O₂ [1]
4LiNO₃ → 2Li₂O + 4NO₂ + O₂ [1]
Na⁺ (ionic radius 102 pm) has a moderate charge density → polarises NO⁻ only slightly → one O is removed → stable nitrite NaNO₂ forms. [1]
Li⁺ (ionic radius 76 pm) is very small with high charge density → strongly polarises the NO⁻ ion → distorts it sufficiently to break N–O bonds → produces NO₂ gas and Li₂O. This anomalous behaviour is part of the diagonal relationship between Li and Mg. [1]
4LiNO₃ → 2Li₂O + 4NO₂ + O₂ [1]
Na⁺ (ionic radius 102 pm) has a moderate charge density → polarises NO⁻ only slightly → one O is removed → stable nitrite NaNO₂ forms. [1]
Li⁺ (ionic radius 76 pm) is very small with high charge density → strongly polarises the NO⁻ ion → distorts it sufficiently to break N–O bonds → produces NO₂ gas and Li₂O. This anomalous behaviour is part of the diagonal relationship between Li and Mg. [1]
Q3 [4 marks]
Describe the flame test procedure for identifying Na⁺, K⁺, and Li⁺ ions. State the colour observed for each and explain why each element produces a different colour. [4]
Procedure: clean nichrome wire with conc. HCl, test in flame (no colour). Dip in test solution; hold in roaring blue flame; observe colour. Clean wire between each test. [1]
Li⁺ → crimson red. Na⁺ → persistent yellow/orange. K⁺ → lilac (view through blue cobalt glass to filter out any Na yellow). [1]
Different colours arise because each element has a unique set of electron energy levels. Flame heat promotes electrons to higher levels (excited states). Electrons fall back to ground state emitting photons of specific energy (E=hf). Each element emits at characteristic frequencies corresponding to specific colours — the visible part of its atomic emission spectrum. [2]
Li⁺ → crimson red. Na⁺ → persistent yellow/orange. K⁺ → lilac (view through blue cobalt glass to filter out any Na yellow). [1]
Different colours arise because each element has a unique set of electron energy levels. Flame heat promotes electrons to higher levels (excited states). Electrons fall back to ground state emitting photons of specific energy (E=hf). Each element emits at characteristic frequencies corresponding to specific colours — the visible part of its atomic emission spectrum. [2]
Q4 [4 marks]
Write balanced equations for the reactions of sodium with: (a) oxygen (excess) [1]; (b) water [1]; (c) explain why Na₂O₂ is used in rebreathers, with a relevant equation [2].
(a) 2Na + O₂ → Na₂O₂ (sodium peroxide forms in excess O₂) [1]
(b) 2Na + 2H₂O → 2NaOH + H₂ [1]
(c) In a rebreather (closed-circuit breathing apparatus, e.g. for submariners), the person exhales CO₂. Na₂O₂ simultaneously absorbs the CO₂ and releases O₂ for breathing: 2Na₂O₂ + 2CO₂ → 2Na₂CO₃ + O₂. This makes the system self-sustaining: no external O₂ supply is needed. [1] Also: Na₂O₂ reacts with exhaled moisture: 2Na₂O₂ + 2H₂O → 4NaOH + O₂, providing additional O₂. [1]
(b) 2Na + 2H₂O → 2NaOH + H₂ [1]
(c) In a rebreather (closed-circuit breathing apparatus, e.g. for submariners), the person exhales CO₂. Na₂O₂ simultaneously absorbs the CO₂ and releases O₂ for breathing: 2Na₂O₂ + 2CO₂ → 2Na₂CO₃ + O₂. This makes the system self-sustaining: no external O₂ supply is needed. [1] Also: Na₂O₂ reacts with exhaled moisture: 2Na₂O₂ + 2H₂O → 4NaOH + O₂, providing additional O₂. [1]
Q5 [4 marks]
Explain the anomalous behaviour of lithium in Group 1 with reference to the diagonal relationship with magnesium. Give THREE specific examples where Li behaves more like Mg than like Na. [4]
Li⁺ (ionic radius 76 pm, charge +1) has a similar charge density to Mg²⁺ (ionic radius 72 pm, charge +2). Because charge density = charge/size, both ions have comparable polarising power despite different charges. [1]
Three examples where Li resembles Mg more than Na: [3, 1 each]
(1) Thermal decomposition of carbonate: Li₂CO₃ decomposes on heating (Li₂CO₃ → Li₂O + CO₂), like MgCO₃ → MgO + CO₂. Na₂CO₃ is thermally stable.
(2) Decomposition of nitrate: LiNO₃ decomposes to Li₂O + NO₂ + O₂ (like Mg(NO₃)₂); NaNO₃ gives NaNO₂ + O₂.
(3) Solubility of Li₂CO₃ and LiF: Both are sparingly soluble (like MgCO₃ and MgF₂), unlike the highly soluble Na₂CO₃ and NaF.
Three examples where Li resembles Mg more than Na: [3, 1 each]
(1) Thermal decomposition of carbonate: Li₂CO₃ decomposes on heating (Li₂CO₃ → Li₂O + CO₂), like MgCO₃ → MgO + CO₂. Na₂CO₃ is thermally stable.
(2) Decomposition of nitrate: LiNO₃ decomposes to Li₂O + NO₂ + O₂ (like Mg(NO₃)₂); NaNO₃ gives NaNO₂ + O₂.
(3) Solubility of Li₂CO₃ and LiF: Both are sparingly soluble (like MgCO₃ and MgF₂), unlike the highly soluble Na₂CO₃ and NaF.
Section B — Extended Answer
30 marksQ6 [8 marks]
Describe the trends in physical properties of Group 1 elements (Li to Cs): melting point, density, atomic radius, first ionisation energy, and reactivity. For each property, state the trend, explain it using atomic structure, and give supporting data where relevant. [8]
Melting point [1.5]: Decreases Li(181)>Na(98)>K(64)>Rb(39)>Cs(28)°C. Metallic bond weakens: increasing atomic radius → larger M⁺ cations further from electron sea → weaker electrostatic attraction → less energy to break bond → lower m.p. All alkali metals have relatively low m.p. compared to other metals due to only 1 delocalised electron.
Density [1.5]: Generally increases (Li 0.53 → Cs 1.87 g/cm³). Atomic mass increases faster than atomic volume. Notable: Li, Na, K all less dense than water (density <1 g/cm³) → float on water during reaction.
Atomic radius [1.5]: Increases (Li 134 pm → Cs 262 pm). Each successive element has one more full electron shell (n increases) → outer electrons in progressively higher, larger shells. Shielding by inner electrons also increases, reducing effective nuclear charge on outer shell.
First ionisation energy [1.5]: Decreases Li(526)>Na(496)>K(419)>Rb(403)>Cs(376) kJ/mol. Outer electron (always ns¹) is in a higher shell each step → greater distance from nucleus and increased shielding → effective nuclear charge on outer electron decreases → less energy to remove it.
Reactivity [2]: Increases Li<Na<K<Rb<Cs. Directly related to IE₁: lower IE → easier to lose outer electron in reactions. Li barely fizzes with water; Na melts and moves rapidly; K ignites H₂; Rb/Cs explode. The fundamental driver is the decreasing ionisation energy, itself caused by the outer electron being increasingly remote and shielded.
Density [1.5]: Generally increases (Li 0.53 → Cs 1.87 g/cm³). Atomic mass increases faster than atomic volume. Notable: Li, Na, K all less dense than water (density <1 g/cm³) → float on water during reaction.
Atomic radius [1.5]: Increases (Li 134 pm → Cs 262 pm). Each successive element has one more full electron shell (n increases) → outer electrons in progressively higher, larger shells. Shielding by inner electrons also increases, reducing effective nuclear charge on outer shell.
First ionisation energy [1.5]: Decreases Li(526)>Na(496)>K(419)>Rb(403)>Cs(376) kJ/mol. Outer electron (always ns¹) is in a higher shell each step → greater distance from nucleus and increased shielding → effective nuclear charge on outer electron decreases → less energy to remove it.
Reactivity [2]: Increases Li<Na<K<Rb<Cs. Directly related to IE₁: lower IE → easier to lose outer electron in reactions. Li barely fizzes with water; Na melts and moves rapidly; K ignites H₂; Rb/Cs explode. The fundamental driver is the decreasing ionisation energy, itself caused by the outer electron being increasingly remote and shielded.
Q7 [8 marks]
Describe fully the reactions of Group 1 metals with oxygen, water, and halogens. For each reaction type: state the product(s), write balanced equations (using Li, Na, or K as examples), explain the trend in reactivity, and note any differences in product type down the group. [8]
With oxygen [3]: All react vigorously when exposed to air, producing metal oxides (type depends on metal). 4Li + O₂ → 2Li₂O (normal oxide). 2Na + O₂ → Na₂O₂ (peroxide). K + O₂ → KO₂ (superoxide, as do Rb and Cs). Trend: Li gives simple oxide; heavier alkali metals form peroxides/superoxides because their larger M⁺ cations can stabilise the larger anions (O²⁻ < O₂²⁻ < O₂⁻). Reactivity increases down group: Cs reacts explosively. All products are basic — react with water to give alkaline solutions.
With water [3]: General: 2M + 2H₂O → 2MOH + H₂. All reactions are exothermic. 2Li + 2H₂O → 2LiOH + H₂ (steady fizzing). 2Na + 2H₂O → 2NaOH + H₂ (vigorous; Na melts/moves). 2K + 2H₂O → 2KOH + H₂ (H₂ ignites; lilac flame). Rb/Cs: explosive. All produce strongly alkaline solutions (MOH dissociates fully). Reactivity increases down group for the same reasons as with oxygen (decreasing IE₁).
With halogens [2]: General: 2M + X₂ → 2MX (metal halide salt, ionic). 2Na + Cl₂ → 2NaCl. 2K + Br₂ → 2KBr. All reactions vigorous; all produce white ionic halide salts. Reactivity of alkali metal increases down group. Reactivity of halogen decreases F₂>Cl₂>Br₂>I₂. All halides are soluble except LiF (sparingly soluble — diagonal relationship).
With water [3]: General: 2M + 2H₂O → 2MOH + H₂. All reactions are exothermic. 2Li + 2H₂O → 2LiOH + H₂ (steady fizzing). 2Na + 2H₂O → 2NaOH + H₂ (vigorous; Na melts/moves). 2K + 2H₂O → 2KOH + H₂ (H₂ ignites; lilac flame). Rb/Cs: explosive. All produce strongly alkaline solutions (MOH dissociates fully). Reactivity increases down group for the same reasons as with oxygen (decreasing IE₁).
With halogens [2]: General: 2M + X₂ → 2MX (metal halide salt, ionic). 2Na + Cl₂ → 2NaCl. 2K + Br₂ → 2KBr. All reactions vigorous; all produce white ionic halide salts. Reactivity of alkali metal increases down group. Reactivity of halogen decreases F₂>Cl₂>Br₂>I₂. All halides are soluble except LiF (sparingly soluble — diagonal relationship).
Q8 [6 marks]
Hydrogen is anomalous in the periodic table. (a) State two ways hydrogen resembles Group 1 elements. [2] (b) State two ways hydrogen differs from Group 1 elements. [2] (c) Write equations for hydrogen reacting with sodium and with nitrogen, and explain the oxidation state of H in each product. [2]
(a) Resemblances to Group 1 [2]: (1) Electron configuration 1s¹ — one electron in the outermost (and only) shell, like ns¹ of alkali metals. (2) Forms H⁺ ions by losing its electron, similar to M⁺ cation formation; H⁺ is involved in acid-base chemistry like alkali metal cations. [1 each, any 2]
(b) Differences from Group 1 [2]: (1) H forms H⁻ (hydride ion, −1) by gaining an electron — alkali metals never form anions. This resembles halogen chemistry. (2) H₂ is a gas (diatomic, covalent) at room temperature; alkali metals are solid reactive metals. Also acceptable: H does not react with water at room temperature; H is a non-metal; H⁺ is a bare proton with no electrons. [1 each, any 2]
(c) H₂ + 2Na → 2NaH (sodium hydride). Na is +1; H is −1 (oxidation state −1 in metal hydrides — H has gained an electron). [1] 3H₂ + N₂ → 2NH₃ (Haber process; needs Fe catalyst, high T and P). H is +1 (bonded to more electronegative N; H loses electron density to N). [1]
(b) Differences from Group 1 [2]: (1) H forms H⁻ (hydride ion, −1) by gaining an electron — alkali metals never form anions. This resembles halogen chemistry. (2) H₂ is a gas (diatomic, covalent) at room temperature; alkali metals are solid reactive metals. Also acceptable: H does not react with water at room temperature; H is a non-metal; H⁺ is a bare proton with no electrons. [1 each, any 2]
(c) H₂ + 2Na → 2NaH (sodium hydride). Na is +1; H is −1 (oxidation state −1 in metal hydrides — H has gained an electron). [1] 3H₂ + N₂ → 2NH₃ (Haber process; needs Fe catalyst, high T and P). H is +1 (bonded to more electronegative N; H loses electron density to N). [1]
Q9 [8 marks]
A student dissolves a Group 1 metal compound in water and tests the solution. The flame test gives a lilac colour. On heating the solid compound, a colourless gas is evolved that rekindles a glowing splint. (a) Identify the metal ion present and the compound. [2] (b) Write the equation for the thermal decomposition. [2] (c) Explain what gas rekindled the splint. [1] (d) Describe two further chemical tests you could carry out on the original compound to confirm its identity, with expected results. [3]
(a) Lilac flame → K⁺ ion present. Gas that rekindles a glowing splint = O₂. A Group 1 potassium compound that evolves O₂ on heating = potassium nitrate (KNO₃). [2]
(b) 2KNO₃ → 2KNO₂ + O₂ [2]
(c) The gas is oxygen (O₂). O₂ supports combustion, so a glowing splint reignites in its presence. [1]
(d) Any two from [1.5 each]: (1) Test for NO⁻ (nitrate) ion: Add dilute H₂SO₄ and a small piece of Fe(II) sulfate; carefully layer concentrated H₂SO₄. A brown ring (iron(III) complex) at the interface confirms NO⁻. (2) Add AgNO₃(aq) after acidifying with dilute HNO₃: No precipitate with NO⁻ (confirms not a halide). (3) Dissolve in water; test pH: KNO₃ solution is approximately neutral (K⁺ neutral; NO⁻ very weakly basic) — pH ~7. (4) Test solubility in water: Highly soluble (confirming not Li₂CO₃ or LiF which are sparingly soluble). [3 total]
(b) 2KNO₃ → 2KNO₂ + O₂ [2]
(c) The gas is oxygen (O₂). O₂ supports combustion, so a glowing splint reignites in its presence. [1]
(d) Any two from [1.5 each]: (1) Test for NO⁻ (nitrate) ion: Add dilute H₂SO₄ and a small piece of Fe(II) sulfate; carefully layer concentrated H₂SO₄. A brown ring (iron(III) complex) at the interface confirms NO⁻. (2) Add AgNO₃(aq) after acidifying with dilute HNO₃: No precipitate with NO⁻ (confirms not a halide). (3) Dissolve in water; test pH: KNO₃ solution is approximately neutral (K⁺ neutral; NO⁻ very weakly basic) — pH ~7. (4) Test solubility in water: Highly soluble (confirming not Li₂CO₃ or LiF which are sparingly soluble). [3 total]