Chapter 9Group Properties and Elements

Increase in the atomic size down the first group weakens the interatomic attraction of the metals. This makes them softer down the group and their melting points decrease. Each atom gives only one electron to the electron sea, and in bigger atoms that electron is further from the nucleus, so the metallic bond is weak. Lithium is hard to cut; sodium cuts like cheese; potassium like butter.

Potassium is below sodium in group 1, so it loses its electron more easily and is more reactive: it reacts violently with halogens to form ionic potassium halides.

With fluorine and chlorine the reaction is explosive; with iodine it is slower, because iodine is the least reactive halogen.

Sodium bromide (the bromide ion, Br⁻). The substance that loses electrons is the reducing agent: 2Br⁻ → Br₂ + 2e⁻. These electrons are accepted by chlorine (Cl₂ + 2e⁻ → 2Cl⁻), which is reduced. Na⁺ takes no part (a spectator ion).

The atoms get larger down the group as they have more electrons, so the molecules formed attract each other with a stronger force. I₂ is the biggest halogen molecule (106 electrons), so the forces between its molecules (London forces) are strong enough to hold them in a solid at room temperature (m.p. 114 °C). F₂ and Cl₂, with small molecules, are gases.

Metals often absorb other substances on their surface and activate them. Hydrogen molecules are adsorbed on the surface of finely divided nickel, where the H–H bond is weakened; the oil molecules are adsorbed too. The activated hydrogen atoms add across the C=C double bonds of the oil, turning the liquid oil into a solid fat (margarine, banaspati ghee). The product leaves the surface and the nickel is unchanged.

They burn in oxygen to give basic oxides, more vigorously down the group. Lithium burns to Li₂O, sodium burns brightly to give Na₂O (and Na₂O₂ in excess oxygen), and potassium burns most vigorously (forming KO₂). They tarnish quickly in air, so they are stored under oil.

Fluorine reacts violently: 2F₂ + 2H₂O → 4HF + O₂. Chlorine partly reacts: Cl₂ + H₂O ⇌ HCl + HOCl; HOCl is a bleach and kills germs. Bromine reacts less, and iodine hardly dissolves or reacts. The trend follows the falling reactivity down the group.

Chlorine with water forms hypochlorous acid (HOCl), a strong oxidising agent that kills bacteria and other germs. Only small amounts are needed, so the water stays safe to swim in.

Tensile strength is the ability of a material to resist being pulled apart (the maximum stretching force it can bear before breaking). Iron (as steel), titanium and copper have high tensile strength, so they are used for cables, girders, bridges and wires.

They have ns¹ configuration. This single electron can be removed easily, which makes these metals very reactive.

Lithium reacts with water steadily, sodium vigorously and potassium violently, giving hydrogen and soluble hydroxides.

Moving down, atomic size increases, so it becomes easier for the elements to lose their electron, which is reflected in the increased reactivity of the lower members.

Increase in atomic size weakens the interatomic attraction, so the metals become softer and their melting points decrease: Li 180, Na 98, K 64, Rb 39, Cs 28 °C.

Both size and volume increase, but the increase in mass is greater than the increase in volume, so density (mass per unit volume) increases. Exception: K (0.86) is lighter than Na (0.97).

They react quickly with oxygen and moisture of the air (and violently with water). Kerosene keeps air and water away.

Halogens react with alkali and alkaline earth metals to give salts. These elements are thus named halogens, which means salt-forming. 2Na + Cl₂ → 2NaCl.

Fluorine: pale yellow gas; chlorine: yellowish-green gas; bromine: fuming red-brown liquid; iodine: shiny grey crystals that turn into dark purple vapour when warmed.

The atoms get larger as they have more electrons, so the molecules formed from them attract each other with a stronger force.

Atomic size increases down the group and the tendency to accept an electron from other atoms decreases, making them less reactive.

Each halogen atom has seven outer electrons and needs one more; two atoms share one pair of electrons so that both complete their octets: F–F, Cl–Cl. So they exist as X₂ molecules.

Alkali and alkaline earth metals lose electrons easily and halogens gain them easily, with a large difference in electronegativity, so electrons are transferred: Na⁺Cl⁻, Mg²⁺(Br⁻)₂.

Oxidation is a process in which an electron is lost; reduction is a process in which an electron is gained. Na → Na⁺ + e⁻ (oxidation); Cl₂ + 2e⁻ → 2Cl⁻ (reduction).

The substance which loses an electron is called a reducing agent; the substance which accepts the electron is called an oxidising agent. In 2Na + Cl₂ → 2NaCl, Na is the reducing agent and Cl₂ the oxidising agent.

An electron is lost only when there is a substance to accept it. Every oxidation (loss) must be matched by a reduction (gain) of the same number of electrons.

A halogen having more oxidising power displaces an ion of another halogen from its compound.

No. Iodine is a weaker oxidising agent than bromine (oxidising power Cl₂ > Br₂ > I₂), so it cannot take the electron from Br⁻. Only a halogen above another in the group can displace it.

The colourless KI solution turns brown (iodine in solution); with a few drops of starch it turns blue-black. Cl₂ + 2KI → 2KCl + I₂.

Halogens react with hydrogen to give hydrogen halides: H₂ + X₂ → 2HX. Hydrogen halides behave as strong acids in water (HF is weak).

As the halogen atom gets bigger, the bonding pair gets further away from the halogen nucleus, so the bond length increases and the bond gets weaker. The weaker the bond, the less heat energy is needed to break it. HF > HCl > HBr > HI.

Fluorine is very small and the most electronegative element, so HF molecules are held together by strong hydrogen bonds. It boils at 19.5 °C; HCl (−85 °C), HBr and HI are gases.

Elements from group 3 to group 12 are d-block or transition elements. All are metals, often hard with higher densities and high melting and boiling points. They show variable oxidation states, form coloured compounds, are malleable and ductile, and act as catalysts.

Metals often absorb (adsorb) other substances on their surface and activate them in this process. Bonds in the adsorbed molecules are weakened, so they react faster; the products leave and the metal is unchanged.

Iron. The ammonia is used to prepare urea fertilizer.

It is the manufacture of sulphuric acid. Platinum was originally used, but vanadium pentoxide (V₂O₅) is now preferred.

A device used in the exhaust of an automobile which converts more harmful gases produced in the engine into gases that do not pollute the atmosphere. Platinum, palladium and rhodium are the catalysts used.

A solid fat made by the hydrogenation of vegetable oils using a nickel catalyst. Margarine is less likely to spoil than butter.

A transition metal can lose different numbers of electrons and so form ions with different charges: iron forms Fe²⁺ (iron(II), FeCl₂) and Fe³⁺ (iron(III), FeCl₃); copper forms Cu⁺ and Cu²⁺.

All have eight electrons (s²p⁶) in their outermost shells except He, which has s². Since their outer shells are complete, they show very little chemical reactivity.

Their outer shells are already complete, so their atoms have no tendency to share electrons with each other: they stay as monoatomic gases (He, Ne, Ar), unlike O₂ or Cl₂.

Helium fills weather balloons and airships (light, does not burn); argon fills electric bulbs so the hot filament does not burn; neon glows red in advertising signs.

Metals can be hammered into thin sheets: malleability. Metals can be drawn into wires: ductility.

Metallic bonding (a sea of free electrons) allows metals to be the best conductors of heat and electricity. The mobile electrons carry charge and pass on heat energy quickly.

Non-metals do not have free electrons, so they cannot conduct heat and electricity. Graphite is the only exception: the electrons between its layers are loosely held and become mobile.

Non-metals do not have free electrons, so the bonds between their particles are weak; they break into pieces when hammered or stretched. Sulphur and phosphorus exist as powders.

Mercury is a liquid metal; iodine is a lustrous (shiny) non-metal.

Metals are held by strong metallic bonds throughout the solid, so much energy is needed to melt them. Most non-metals are small molecules held by weak forces, which break easily (sulphur melts at 115 °C). Graphite and diamond are exceptions: giant covalent networks with very high melting points.

Helium (b.p. −269 °C, only 4 °C above absolute zero). The atoms of noble gases are held together only by very weak attractive forces (London forces), which grow with the size and number of electrons of the atom. Helium has the smallest atom with only two electrons, so these forces are weakest of all; very little energy separates its atoms. Boiling points rise down the group: He −269, Ne −246, Ar −186, Kr −153, Xe −108, Rn −62 °C.

All alkali metals react with chlorine to form white, ionic metal chlorides, 2M + Cl₂ → 2MCl. The reaction becomes more vigorous as we go down the group.

The reason: atoms get bigger down the group, so the outer electron is lost more easily to chlorine.

Metals are held together by strong metallic bonds: every positive ion is attracted to the sea of mobile electrons in all directions. A lot of energy is needed to break these bonds, so metals have high melting points and are solids (mercury is the only liquid metal).

Non-metals mostly exist as small separate molecules (H₂, N₂, O₂, Cl₂). The forces between these molecules are weak, so they are gases. Those with larger molecules or networks (S₈, P₄, I₂, carbon) are solids. Only bromine is a liquid.

Mercury and bromine are the only two elements that are liquid at a standard room temperature of 25 °C.

The differences come from the partly filled d sub-shell, whose electrons can take part in bonding.

Chlorine is a stronger oxidising agent than bromine. Chlorine atoms are smaller, so the nucleus attracts an incoming electron more strongly; oxidising power decreases down the group: Cl₂ > Br₂ > I₂.

Bromine cannot displace chlorine from NaCl: Br₂ + NaCl → no reaction. Both can oxidise iodide to iodine.

Fluorine is the most reactive; iodine (among the common halogens) is the least reactive.

Astatine, below iodine, is radioactive and very rare.

Harmful exhaust gases

Burning petrol or diesel in a car engine produces, besides CO₂ and water, harmful gases: carbon monoxide (poisonous), oxides of nitrogen NOₓ (cause acid rain and smog) and unburnt hydrocarbons.

What a catalytic converter is

A catalytic converter is a device used in the exhaust of an automobile which converts the more harmful gases produced in the engine into gases which do not pollute the atmosphere. Platinum, palladium and rhodium are the catalysts used.

Construction

A steel case holds a ceramic honeycomb coated with a very thin layer of the metals, giving a huge surface area. The hot exhaust gases pass through the channels.

Reactions

How the catalyst works

The gases are adsorbed on the metal surface, where their bonds are weakened and they react quickly at the exhaust temperature; the harmless products leave the surface.

Different ways of reacting

Alkali metals react by losing their one outer electron; halogens react by gaining one electron. Both depend on how strongly the nucleus holds the outer shell, and that grows weaker down a group as the atoms get bigger.

Alkali metals: reactivity increases

When we move from top to bottom, the atomic size increases. Owing to this it becomes easier for elements to lose an electron, which is reflected in the increased reactivities of the lower members. The outer electron is further from the nucleus and better shielded: ionisation energy falls (Li 520 → Cs 377 kJ mol⁻¹). Li reacts steadily with water, Na vigorously, K violently.

Halogens: reactivity decreases

Atomic size increases down the group and the tendency to accept an electron from other atoms decreases, making them less reactive. A bigger atom attracts an incoming electron less strongly: electronegativity falls (F 4.0 → I 2.7) and oxidising power falls (Cl₂ > Br₂ > I₂).

Metals

In metals the positive ions are arranged in rows and held by strong metallic bonds with a sea of mobile electrons. These bonds act in all directions; when force is applied the rows slide over each other without breaking, because the electron sea still holds them. So metals bend rather than break: they are malleable, ductile, tough and have high tensile strength (they can hold heavy weights). They are hard and dense, with high melting points.

Non-metals

Non-metals do not have free electrons, so the bonds between their atoms are weak and they break down when stretched. Most solid non-metals are made of separate molecules (S₈, P₄, I₂) held by weak intermolecular forces; a blow simply separates them, so they are brittle and crumble. Gaseous non-metals have no strength at all.

Exception

Diamond, a giant covalent network, is the hardest natural substance, but it is still brittle and shatters with a sharp blow.

Alkali metals: electron givers

With ns¹ in the outer shell, they have a strong tendency to lose one electron and form cations (M⁺). They are electropositive metals and reducing agents (they are oxidised): Na → Na⁺ + e⁻. Reactivity increases down the group.

Halogens: electron takers

With ns²np⁵, they have a strong tendency to gain one electron and form anions (X⁻). They are electronegative non-metals and oxidising agents (they are reduced): Cl₂ + 2e⁻ → 2Cl⁻. Reactivity decreases down the group.

When they meet

Sodium gives its electron to chlorine; both reach noble-gas configurations, and an ionic salt forms with a large release of energy.

Bond length

The bond length between hydrogen and halogen increases down the group, because as the halogen atom gets bigger, the bonding pair of electrons gets further away from the halogen nucleus. Bromine is bigger than chlorine, so the H–Br bond (141 pm) is longer than H–Cl (127 pm).

Bond strength

The bond between hydrogen and halogen therefore gets weaker. H–Br 366 kJ mol⁻¹ against H–Cl 431 kJ mol⁻¹.

Thermal stability

The weaker the bond, the less heat energy is needed to break it. Hence the thermal stability of hydrogen halides decreases down the group, and HBr decomposes more easily on heating than HCl:

Metals

Metals are elements which can generally form cations easily and tend to form metallic bonds. Examples: copper, silver, iron, aluminium, gold, platinum, zinc.

Non-metals

Non-metals generally gain electrons easily and show a greater variety of colours and physical states. Examples: oxygen, nitrogen, chlorine, sulphur, carbon, bromine.

Physical properties

Chemical properties

The reaction

1. Platinum is poisoned

Platinum is rendered inactive due to the presence of arsenic as impurity in sulphur dioxide. Arsenic compounds stick to the platinum surface and block it, so the catalyst stops working; V₂O₅ is not poisoned so easily.

2. Cost

Platinum is an expensive catalyst; vanadium pentoxide is far cheaper, and large amounts are needed in a sulphuric acid plant.

3. Durability

V₂O₅ lasts for years and works well at the operating temperature of about 450 °C.

The job

In some nuclear reactors (fast breeder reactors) a huge amount of heat is produced in the core and must be carried away quickly to a heat exchanger, where it boils water into steam that drives the turbines.

Why liquid sodium

Property used

Mainly its high thermal conductivity as a liquid metal, together with its low melting point.

Why

How it is different

Diagonal relationship

In many of these ways lithium resembles magnesium, the element diagonally below it in group 2.

Aluminium

Magnesium

So aluminium is safe and durable for cooking, while magnesium would corrode, contaminate food and be a fire risk.

Electronic configuration

All group 1 elements have ns¹ configuration in their outermost shells. This single electron can be removed easily, which makes these metals very reactive, except hydrogen, which is a gas and a non-metal.

Reaction with water

Lithium reacts with water steadily, sodium vigorously and potassium violently, giving water-soluble hydroxides and hydrogen.

Reaction with chlorine

The reaction with chlorine becomes more vigorous down the group. 2Na + Cl₂ → 2NaCl.

Reaction with oxygen

They burn to form basic oxides, which dissolve to give alkalis: K₂O + H₂O → 2KOH.

Softness and melting point

Increase in atomic size weakens the interatomic attraction, so they become softer and their melting points decrease.

Density

The increase in mass is greater than the increase in volume, so density increases down the group (K is an exception).

General

Group 17 elements have seven electrons in their outermost shells. They are electronegative non-metals with a strong tendency to accept one electron. They exist as diatomic molecules and are very reactive.

Physical properties

Fluorine is a pale yellow gas, chlorine yellowish green, bromine a fuming red-brown liquid and iodine shiny grey crystals giving purple vapour. Atomic radii, melting and boiling points increase down the group.

Reactivity

They react with alkali and alkaline earth metals to give salts (metal halides, usually ionic). Unlike metals, reactivity decreases down the group.

Oxidising agents

Halogens are oxidising agents (they accept electrons) and their oxidising power decreases down the group: Cl₂ > Br₂ > I₂.

Displacement

A halogen with more oxidising power displaces an ion of another halogen from its compound:

Hydrogen halides

H₂ + X₂ → 2HX; strong acids in water; thermal stability HF > HCl > HBr > HI.

Definition

Elements at the centre of the modern periodic table from group 3 to group 12 are called d-block or transition elements. All transition elements are metals with similar properties.

Properties

How they catalyse

Metals often absorb other substances on their surface and activate them.

Industrial catalysts

Members

Elements in group 18 are called noble gases: helium, neon, argon, krypton, xenon and radon.

Why unreactive

All have eight electrons (s²p⁶) in their outermost shells except helium, which has s². Since their outer shells are complete, they show very little chemical reactivity. They have no tendency to lose, gain or share electrons, and they have the highest ionisation energies of their periods.

Properties

Uses

1. Thermal conductivity

Metals are the best conductors of heat: their free electrons pass energy on quickly (cooking pots are made of aluminium, copper and steel). Non-metals are poor conductors of heat (insulators): sulphur, wood and plastics.

2. Electrical conductivity

Metallic bonding allows metals to be the best conductors of electricity (copper wires). Non-metals have no free electrons and cannot conduct, except graphite, which has mobile electrons between its layers.

3. Malleability and ductility

Metals can be hammered into thin sheets (malleability) and drawn into wires (ductility), because rows of ions slide past each other while the electron sea keeps holding them. Non-metals are brittle: they break into pieces when hammered and cannot be drawn into wires.

4. Melting and boiling points

Strong metallic bonds give metals generally high melting and boiling points (iron 1538 °C); alkali metals are soft exceptions. Non-metals have low melting and boiling points (sulphur 115 °C; oxygen boils at −183 °C), except graphite and diamond (above 3500 °C).