CC13 · Groups in the periodic tableTopic 6 — Groups in the periodic table
Alkali metals, halogens, noble gases and group trends.
Revise the key ideas
Group 1: alkali metals
Group 1 alkali metals include lithium, sodium and potassium. They have one outer-shell electron and commonly form +1 ions by losing it.
They are soft metals, can be cut with a knife in a controlled demonstration, and have relatively low melting points compared with most metals. They conduct electricity.
They react readily with non-metals such as oxygen, forming oxygen-containing compounds. Fresh surfaces tarnish in air, so samples are commonly stored under oil.
With water they form a metal hydroxide and hydrogen, making an alkaline solution: metal + water → metal hydroxide + hydrogen.
Lithium floats, moves and fizzes relatively gently. Sodium reacts more vigorously, often melts into a ball and moves across the water. Potassium reacts still more vigorously and can ignite with a lilac flame.
These water reactions release heat. Sodium and potassium can melt because heat exceeds their relatively low melting temperatures; lithium does not normally melt in the same demonstration.
The balanced sodium equation is 2Na + 2H₂O → 2NaOH + H₂. Hydroxide ions in the solution explain why an indicator shows alkalinity.The reaction is exothermic and increasingly vigorous down group 1.
Alkali-metal demonstrations require teacher control, shielding and small samples. Do not attempt them at home or assume oil storage makes the metal harmless.
Explaining group 1 reactivity
Reactivity increases down group 1: lithium < sodium < potassium. Predict rubidium and caesium to be more reactive than potassium, using the trend.
Down group 1, atoms have more electron shells. The outer electron is further from the nucleus. Inner electrons reduce the nucleus’s attraction to it: this is shielding.
Although nuclear charge increases, distance and shielding reduce the effective attraction to the outer electron. It is lost more readily, making the metal more reactive.Group 1 reacts by electron loss, not gain.
Lithium is 2,1; sodium 2,8,1; potassium 2,8,8,1. The same outer count gives similar reactions; the growing shell count helps explain changing reactivity.One outer electron remains; added inner shells increase distance and shielding.
Forming a +1 ion is oxidation: M → M⁺ + e⁻. Do not explain the trend by saying group 1 atoms gain electrons or that more protons alone must increase attraction.
Atomic number rises down the group, but atomic number by itself is not a complete explanation. Connect shell count, shielding, attraction and electron loss.
Group 7: halogen properties
Group 7 halogens have seven outer-shell electrons and commonly gain one electron to form −1 halide ions. They are non-metals. Their molecules contain two atoms (they are diatomic), for example Cl₂, Br₂ and I₂.
At room temperature chlorine is a pale green gas, bromine a red-brown liquid and iodine a grey-black solid. Iodine vapour is purple; do not confuse vapour colour with the bulk solid.
From chlorine to bromine to iodine, melting point, boiling point and density generally increase. Stronger intermolecular attractions between larger molecules explain the boiling-point trend.
Approximate source-table values are chlorine melting −101 °C / boiling −34 °C; bromine −7 °C / 59 °C; iodine 114 °C / 184 °C. Use supplied values for state predictions.At about 20 °C, compare temperature with melting and boiling points.
Halogens react with metals to form metal halides: 2Na + Cl₂ → 2NaCl. Sodium chloride is table salt; fluoride compounds are used in toothpaste, rather than elemental fluorine.
Hydrogen reacts with halogens to form hydrogen halides. Dissolving hydrogen chloride in water gives hydrochloric acid; hydrogen chloride gas and the aqueous acid are different conditions.
Chlorine bleaches damp blue litmus paper, often after it first turns red. Chlorine can kill microorganisms and remove colour, but harmful gas must be handled with proper controls.
Halogen reactivity and displacement
Halogen reactivity decreases down the group: fluorine > chlorine > bromine > iodine. Predict astatine to be less reactive than iodine from this pattern.
Down the group, extra shells increase distance and shielding, weakening the attraction for an incoming electron. The atom gains an electron less readily.Halogens react by gaining electrons; the trend is opposite to group 1.
A more reactive halogen displaces a less reactive halogen from a halide solution. Chlorine displaces bromine from bromide and iodine from iodide; bromine displaces iodine from iodide.Chloride is not displaced by bromine or iodine; a halogen cannot displace itself.
Chlorine gains electrons and is reduced: Cl₂ + 2e⁻ → 2Cl⁻. Bromide ions lose electrons and are oxidised: 2Br⁻ → Br₂ + 2e⁻. The displacement is redox.Chlorine is reduced; bromide, not chlorine, is oxidised.
Iodine cannot displace chlorine from chloride or bromine from bromide. Equal-halogen combinations do not produce a displacement reaction.
Use small controlled volumes and observe solution colour changes; compare with reference solutions. Colours depend on solvent and concentration, so use the given practical information.
Group 0: noble gases
Group 0 noble gases include helium, neon, argon and krypton. At room temperature they are colourless gases with low boiling points. Each particle is a single atom: they are monatomic.
Their atoms have complete outer shells: helium has two electrons in its first shell, while neon and argon have eight in their outer shells.A full first shell contains two; helium does not need eight electrons.
Full shells explain their very low reactivity compared with other groups and why they do not commonly form ions or compounds. Some heavier noble gases can form compounds under special conditions, so avoid “never react”.
Boiling point and density generally increase down the familiar noble-gas group. Larger atoms have stronger intermolecular attractions, requiring more energy to separate them on boiling.
Noble gases are non-flammable. Their low reactivity and occurrence as uncombined atoms made their discovery difficult; several were identified in the late nineteenth century.
Helium's low density and non-flammability make it suitable for balloons and airships; hydrogen is light but flammable.
Argon provides an unreactive atmosphere for welding and other oxidation-sensitive processes, including protective gas applications. Neon emits red-orange light in electrical discharge tubes; krypton is used in some photographic flash lamps.Match the application to the property, rather than simply memorising names.
Although ordinary noble gas samples conduct poorly, an electrical discharge can excite or ionise gas and produce light. Do not equate a glowing sign with normal metal-like conduction.
Watch CC13 · Groups in the periodic table · Topic 6 — Groups in the periodic table
Revise groups in the periodic table with this narrated video. Use the player controls to pause, seek, adjust the volume or mute. Turn English captions on or off using the captions menu.
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