A group is a family
When Mendeleev built his periodic table, his strongest evidence was that elements in the same group behave alike: lithium, sodium, and potassium are all reactive metals soft enough to cut with a knife, and chlorine, bromine, and iodine all form salts with metals. With the electron configurations and periodic trends of lessons 2.3.1 to 2.3.5, you can now answer two questions: why are the members of a group alike, and why does each one get more reactive (or less reactive) going down the group?
The alkali metals
The group 1 elements other than hydrogen (lithium, sodium, potassium, rubidium, cesium, and radioactive francium) are the alkali metals碱金属. All have an ns¹ valence configuration and a very low first ionization energy, so they readily lose that one electron to form 1+ ions.
- Physical properties: all are soft, silvery metals (cesium is slightly golden) with low densities and low melting points. From lithium to cesium, the melting point falls steadily; density generally rises, but potassium is less dense than sodium. Sodium and potassium are stored under oil to keep out air and water; lithium is so light it floats on oil, so it’s often sealed in paraffin wax or kept under argon.
- Reaction with water: every alkali metal forms a hydroxide and hydrogen gas. For sodium:
Drop a small piece of sodium into water containing phenolphthalein, and it floats (it’s less dense than water), melts into a shiny ball (the reaction releases heat and sodium melts easily), darts around (pushed by the hydrogen it releases), hisses, and turns the water pink (a base forms). Chinese students memorize these five observations as five characters: 浮、熔、游、响、红. Lithium reacts gently; potassium is more violent, and the heat ignites the hydrogen with a lilac flame; rubidium and cesium explode.
- Reaction with oxygen: the reaction gets more vigorous down the group, and the products get more complex. Lithium forms only Li₂O; sodium burns in air to form pale yellow Na₂O₂ (sodium peroxide); potassium, rubidium, and cesium burn to form even more oxygen-rich compounds.
- Why: down the group, each atom adds an energy level, so the atoms get larger and the first ionization energy drops (Li 520 → Cs 376 kJ/mol). The single outer electron gets easier and easier to remove, so the metals get more reactive. That’s also why potassium comes before sodium in the activity series of lesson 1.8.3.
The halogens
The group 17 elements fluorine, chlorine, bromine, and iodine (plus radioactive astatine) are the halogens卤素. All have an ns²np⁵ valence configuration, so they readily gain one electron to form 1− ions.
- Physical properties: from fluorine to iodine, the elements get darker in color, change from gas to liquid to solid, and have higher melting and boiling points. Bromine is the only nonmetal that is liquid at room temperature, and iodine sublimes easily when heated. Melting and boiling points rise because the molecules get larger and the forces between them get stronger (lesson 2.4.8).
- Reaction with hydrogen: all form hydrogen halides, but less readily down the group. Fluorine and hydrogen combine explosively even in the dark; chlorine and hydrogen react only when exposed to light or a flame; bromine needs heating; and iodine reacts slowly only with continuous heating, while the HI it forms breaks down again, so the reaction is reversible. The stability of the hydrogen halides follows HF > HCl > HBr > HI.
- Displacement reactions: a more reactive halogen displaces a less reactive one from a solution of its salt, just as a more active metal displaces a less active one:
So the halogens’ oxidizing strength (their ability to gain electrons) is F₂ > Cl₂ > Br₂ > I₂. Fluorine is too reactive for these experiments: it reacts violently with the water itself.
- Why: down the group, each atom adds an energy level, so the atoms get larger and the nucleus attracts an incoming electron less strongly. The halogens get worse at gaining electrons and less nonmetallic. The alkali metals react by losing electrons and get more reactive down the group; the halogens react by gaining electrons and get less reactive.
The noble gases
The group 18 elements helium, neon, argon, krypton, xenon, and radon are the noble gases稀有气体. Their outer level is ns²np⁶ (helium is 1s²), already full, so they have the highest first ionization energy in each period and react with almost nothing. They exist as single atoms, and from helium to radon the atoms get larger and the boiling points rise.
The noble gases were once called “inert gases”, on the assumption that they never react. In 1962, chemists made the first noble-gas compound, and xenon compounds such as XeF₂ and XeF₄ soon followed, so the name “noble gases” is now preferred.
Drop each alkali metal into water and compare what happens, then look at how their physical properties change. Next, switch to the halogens and mix halogen solutions with potassium halides in the table. Finally, look at the noble gases’ boiling points and uses:
Drop a small piece of alkali metal into water containing phenolphthalein:
Choose an alkali metal. Rubidium and cesium explode on contact with water, so they're never tried in a school lab; this is only a simulation.
Compare the physical properties of the alkali metals:
Add a halogen solution to a potassium halide solution. Click a cell to see what happens:
| KCl | KBr | KI | |
|---|---|---|---|
| Chlorine water | |||
| Bromine water | |||
| Iodine solution |
Click a cell in the table.
Physical properties of the halogens (melting and boiling points in °C):
| Element | Color and state at room temperature | Melting point | Boiling point |
|---|---|---|---|
| Fluorine (F₂) | pale yellow gas | −219.6 | −188.1 |
| Chlorine (Cl₂) | yellow-green gas | −101.5 | −34.0 |
| Bromine (Br₂) | dark red-brown liquid | −7.2 | 58.8 |
| Iodine (I₂) | purple-black solid | 113.7 | 184.3 |
Boiling points of the noble gases, rising from helium to radon (bar length is measured from absolute zero):
- Helium: light and nonflammable, it fills balloons and airships; liquid helium (−269 °C) cools the superconducting magnets in hospital MRI scanners.
- Neon: glows red when electricity passes through it, which is how neon signs work.
- Argon: shields hot metal from air during welding, and fills light bulbs.
- Xenon: reacts with fluorine to form compounds such as XeF₂ and XeF₄, so noble gases aren't completely unreactive.
Example 1: Predicting rubidium’s properties
Rubidium (Rb) is in period 5, group 1. Predict how its reaction with water compares with potassium’s, write the equation, and give the formula of its chloride.
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Structure. Rubidium’s valence configuration is 5s¹; it has 5 shells, one more than potassium.
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From structure to properties. The atom is larger, its outer electron is farther out, and its first ionization energy is lower (K 419, Rb 403 kJ/mol), so it loses that electron more easily and reacts with water even more violently than potassium, sometimes explosively.
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Write the equation. Alkali metals react with water to form a strong base and hydrogen:
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Compounds. With one valence electron, rubidium forms Rb⁺, which combines 1 : 1 with Cl⁻ as RbCl; RbOH is an even stronger base than KOH.
Example 2: What can bromine displace?
Bromine water is added to a KCl solution and to a KI solution. What do you see in each? Write equations for any reactions that occur.
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The activity order. As oxidizing agents, Cl₂ > Br₂ > I₂. A more active halogen displaces a less active one from its salt solution, not the other way round.
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Into KCl solution. Bromine is less active than chlorine and can’t displace it, so nothing reacts; the solution just shows the orange-yellow of the bromine water.
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Into KI solution. Bromine is more active than iodine and displaces it, turning the solution brown:
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How can you show iodine formed? Add starch solution: it turns blue, showing elemental iodine is present.
- Na or K: which reacts more violently with water?
Potassium has one more energy level than sodium, so its outer electron is easier to remove.
- Li or Na: which has the higher melting point?
Alkali-metal melting points fall down the group: Li 180.5 °C, Na 97.8 °C.
- F₂ or Cl₂: which combines with hydrogen more easily?
Fluorine reacts explosively with hydrogen even in the dark; chlorine needs light or a flame.
- I₂ or Br₂: which combines with hydrogen more easily?
The halogens combine with hydrogen less readily down the group, and bromine is above iodine.
- HCl or HBr: which is more stable?
Hydrogen halide stability follows HF > HCl > HBr > HI.
- KOH or NaOH: which is the stronger base?
Potassium is more metallic than sodium, and KOH is the stronger base.
- Ar or Kr: which has the higher boiling point?
Noble-gas boiling points rise down the group: Ar −185.8 °C, Kr −153.4 °C.
- Na or K: which is denser?
Alkali-metal density generally rises down the group, but potassium is an exception: Na 0.968 g/cm³, K 0.862 g/cm³.
- Br₂ or Cl₂: which is the stronger oxidizing agent?
Chlorine displaces bromine from KBr solution, so Cl₂ > Br₂.
Common mistakes
Mistake 1: “Alkali-metal density increases steadily down the group.” It rises overall, but potassium (0.862 g/cm³) is less dense than sodium (0.968 g/cm³). Lithium, sodium, and potassium are all less dense than water, so they float.
Mistake 2: “Elements always get more reactive down a group.” That’s true for the alkali metals, but the halogens are the reverse. Alkali metals react by losing electrons, and larger atoms lose them more easily. Halogens react by gaining electrons, and the larger the atom, the weaker the nucleus’s pull on an incoming electron, so the less reactive the halogen.
Mistake 3: “Noble gases are inert and never react with anything.” They are very unreactive, but not completely. Since 1962, chemists have made xenon compounds such as XeF₂ and XeF₄.
Quick check
Which statement about the alkali metals is correct?
From lithium to cesium, first ionization energy drops, so the outer electron is lost more easily and the reaction with water gets more violent. Melting point falls; potassium is less dense than sodium; and every alkali metal reacts with water, which is why sodium and potassium are stored under oil.
Chlorine water is added to a KI solution, and the solution turns yellow-brown. What does this show?
The reaction is Cl₂ + 2KI → 2KCl + I₂: chlorine displaced iodine, so chlorine gains electrons more readily than iodine. The I⁻ ions lost electrons to become I₂.
Summary
- Alkali metals: ns¹ valence configuration, readily lose one electron to form 1+ ions. From lithium to cesium, melting point falls, density generally rises (potassium is the exception), and reactions with water and oxygen get more violent as the metals get more reactive.
- Halogens: ns²np⁵ valence configuration, readily gain one electron to form 1− ions. From fluorine to iodine, color darkens, melting and boiling points rise, reaction with hydrogen gets harder, the hydrogen halides get less stable, and oxidizing strength follows F₂ > Cl₂ > Br₂ > I₂.
- Noble gases: full outer level, very unreactive, made of single atoms; boiling points rise from helium to radon. Xenon forms compounds such as XeF₂.
On to AP (for later)
AP doesn’t ask you to memorize each group’s reactions, but it uses them to test other ideas. Alkali metals reacting with water and halogen displacement are both oxidation–reduction reactions (U4.9). The rising boiling points of the halogens and noble gases come from stronger London dispersion forces between larger atoms and molecules (U3.1).