4.1.1 Properties and trends in Groups 1, 7 and 0

Soft, low-density alkali metals react with water

1. Lithium, sodium and potassium are soft enough to cut with a knife, and each freshly cut surface is shiny.
   1. The shiny surface tarnishes quickly as the metal reacts with oxygen in the air, so the metals are stored under oil.
2. The alkali metals have low densities, so lithium, sodium and potassium all float on water.
3. They have low melting points for metals, and the melting points decrease down the group.
4. Every alkali metal reacts with water to form hydrogen gas and a metal hydroxide, for example: $$2Na(s)+2H_{2}O(l)\to 2NaOH(aq)+H_2(g)$$
   1. The metal hydroxide dissolves to give an alkaline solution, which turns universal indicator purple.

Group 1 reactivity increases down the group

1. Every alkali metal atom has one outer electron, which it loses to form an ion with a $1+$ charge, such as ${Na}^{+}$.
2. Going down the group, each atom has one more shell of electrons, so the outer electron is further from the nucleus.
   1. The extra inner shells also shield the outer electron from the attraction of the positive nucleus.
3. The weaker attraction means the outer electron is lost more easily, so reactivity increases down the group.
4. This is why potassium reacts with water more vigorously than sodium, and sodium more vigorously than lithium.
5. The trend predicts that rubidium and caesium, below potassium, react with water even more violently.

Halogens darken and boil higher down Group 7

1. Each halogen exists as diatomic molecules, such as ${Cl}_2$, ${Br}_2$ and $I_2$, in which two atoms share a pair of electrons.
2. At room temperature, chlorine is a pale green gas, bromine is a red-brown liquid and iodine is a grey-black solid.
   1. Bromine gives off an orange-brown vapour, and iodine gives off a purple vapour when it is warmed.
3. The colour gets darker and the melting and boiling points increase down the group.
   1. The molecules get larger down the group, so the intermolecular forces between them get stronger and more energy is needed to separate them.
4. Halogens react with metals to form metal halides, which are ionic salts, for example: $$2Na(s)+{Cl}_2(g)\to 2NaCl(s)$$

Falling Group 7 reactivity explains halogen displacement

1. Every halogen atom has seven outer electrons and gains one electron to form a halide ion with a $1-$ charge, such as ${Cl}^{-}$.
2. Going down the group, the outer shell is further from the nucleus and more shielded by inner shells.
   1. An incoming electron is attracted less strongly, so it is gained less easily and reactivity decreases down the group.
3. Fluorine is therefore the most reactive halogen, and iodine is the least reactive of the common halogens.
4. A halogen displaces a less reactive halogen from a solution of its halide salt, for example: $${Cl}_2(aq)+2KBr(aq)\to 2KCl(aq)+{Br}_2(aq)$$
   1. The colourless potassium bromide solution turns orange as bromine forms.
5. A less reactive halogen cannot displace a more reactive one, so iodine does not react with potassium bromide solution.

Full outer shells make the noble gases unreactive

1. Helium has a full first shell of two electrons, and the other noble gases have eight outer electrons.
   1. A full outer shell gives the atoms no tendency to lose, gain or share electrons, so the noble gases are very unreactive.
2. The noble gases are monatomic, existing as single atoms rather than molecules.
3. They are colourless gases at room temperature with very low boiling points.
4. Boiling points increase down the group, from $-269{}^{\circ }C$ for helium to $-108{}^{\circ }C$ for xenon.
   1. The atoms get larger down the group, so the forces between the atoms get stronger.
5. Density also increases down the group, because the atoms get heavier.

4.1.2 Properties of the transition metals

Transition metals are hard, strong and dense

1. Transition metals such as iron, copper, nickel and chromium sit in the central block, between Groups 2 and 3.
2. Like other metals they are shiny, conduct electricity and heat, and can be hammered into shape.
3. Most have high melting points, so iron melts at $1538{}^{\circ }C$ while sodium melts at $98{}^{\circ }C$.
4. Most have high densities, so copper has a density of $8.96{g/cm}^3$ while potassium has a density of $0.86{g/cm}^3$.
5. They are hard and strong, which is why iron, in the form of steel, is used for bridges, girders and car bodies.
   1. The Group 1 metals, by contrast, are soft enough to cut with a knife.

Transition metals are less reactive than Group 1

1. Transition metals react very slowly or not at all with cold water, while every Group 1 metal reacts with cold water and releases hydrogen.
   1. Iron rusts only slowly, over days or weeks, when it is in contact with both water and oxygen.
2. Copper does not react with water or with dilute acids, so it is used for water pipes and hot water tanks.
3. Gold and silver are so unreactive that they stay shiny for centuries, which suits jewellery and coins.
4. Transition metal atoms form positive ions less readily than Group 1 atoms, which is why they react less vigorously.

Many transition metals form ions with different charges

1. A Group 1 metal always forms ions with a $1+$ charge, but many transition metals form more than one positive ion.
   1. Iron forms both iron(II) ions, ${Fe}^{2+}$, and iron(III) ions, ${Fe}^{3+}$.
   2. Copper forms both copper(I) ions, ${Cu}^{+}$, and copper(II) ions, ${Cu}^{2+}$.
2. The Roman numeral in the name of a compound gives the charge on the metal ion.
3. Different charges give different formulae, so iron(II) chloride is ${FeCl}_2$ and iron(III) chloride is ${FeCl}_3$.
   1. Compounds of the same metal with different ion charges have different properties, including different colours.

Many transition metal compounds are coloured

1. Compounds of Group 1 metals, such as sodium chloride, are white solids that dissolve to give colourless solutions.
2. Many transition metal compounds are coloured, both as solids and in solution.
   1. Copper(II) sulfate solution is blue, iron(II) sulfate solution is pale green and iron(III) chloride solution is orange-brown.
   2. Potassium manganate(VII) contains manganese and forms a deep purple solution.
3. The colour depends on the ion charge, which is why iron(II) and iron(III) compounds have different colours.
4. The colours help in identifying ions, because each of these metal ions gives a precipitate of its own colour with sodium hydroxide solution.

Transition metals and their compounds are useful catalysts

1. Many transition metals and their compounds speed up reactions without being used up, so a small amount can be used again and again.
2. Iron is the catalyst in the Haber process, which makes ammonia for fertilisers: $$N_2(g)+3H_2(g)\rightleftharpoons 2{NH}_3(g)$$
3. Nickel catalyses the reaction of vegetable oils with hydrogen to make margarine.
4. Manganese(IV) oxide catalyses the decomposition of hydrogen peroxide: $$2H_2{O}_2(aq)\to 2H_{2}O(l)+O_2(g)$$
5. Platinum and rhodium in car catalytic converters speed up reactions that remove carbon monoxide and nitrogen oxides from exhaust gases.

4.1.3 Reactivity predicted from position in the Periodic Table

An element's group predicts its type of reaction

1. Elements in the same group have the same number of outer electrons, so they take part in similar reactions.
2. Metals in Groups 1 and 2 react by losing their outer electrons, forming ions with $1+$ and $2+$ charges.
   1. Sodium reacts with water to form hydrogen and sodium hydroxide, so potassium and lithium, in the same group, are predicted to do the same.
3. Non-metals in Groups 6 and 7 react with metals by gaining electrons, forming ions with $2-$ and $1-$ charges.
   1. Chlorine reacts with sodium to form sodium chloride, so bromine is predicted to form sodium bromide in the same way.
4. Two non-metals, such as carbon and oxygen, react by sharing electrons to form a covalent compound.
5. Group 0 elements have full outer shells, so no reaction is predicted for them.

Using ion charges to predict product formulae

1. The charges on the ions must cancel out, so the formula of an ionic product follows from the two group numbers.
2. A Group 1 metal and a Group 7 non-metal combine in a $1:1$ ratio, as in $KBr$ and $LiCl$.
3. A Group 2 metal and a Group 7 non-metal combine in a $1:2$ ratio, so calcium and chlorine form calcium chloride, ${CaCl}_2$.
4. A Group 1 metal and a Group 6 non-metal combine in a $2:1$ ratio, so lithium and oxygen form lithium oxide, ${Li}_{2}O$.
5. With the formula known, balance the equation using the diatomic formula of the non-metal element: $$4Li(s)+O_2(g)\to 2{Li}_{2}O(s)$$

Position within a group predicts reactivity

1. Metals become more reactive down a group, because the outer electrons are further from the nucleus and are lost more easily.
   1. Caesium, near the bottom of Group 1, is predicted to react with water even more violently than potassium.
2. Non-metals become less reactive down a group, because an extra electron is attracted less strongly by the nucleus.
   1. Astatine, at the bottom of Group 7, is predicted to be the least reactive halogen, so it should not displace iodine from potassium iodide solution.
3. Predicted reactivity also shows whether a displacement will happen, because only a more reactive element can take the place of a less reactive one in its compound.

Reactivity trends across a period

1. A Group 1 metal is more reactive than the Group 2 metal in the same period, so sodium reacts vigorously with cold water while magnesium reacts very slowly.
   1. A Group 2 atom must lose two electrons, and its nucleus has more protons attracting them, so it forms ions less readily.
2. The transition metals, in the centre of the table, are much less reactive than the metals in Groups 1 and 2.
3. Among the non-metals, a Group 7 element is usually more reactive than the Group 6 element in the same period, so fluorine is more reactive than oxygen.
4. Combining both trends puts the most reactive metals at the bottom left and the most reactive non-metal, fluorine, at the top right, next to Group 0.
5. Group 0 breaks the pattern, because full outer shells make its elements almost completely unreactive.

4.1.4 Reactivity of metals and the reactivity series

More reactive metals form positive ions more readily

1. When a metal reacts with water or a dilute acid, its atoms lose electrons and become positive ions.
   1. Magnesium atoms become ${Mg}^{2+}$ ions, while the hydrogen given off comes from the water or the acid.
2. A metal whose atoms form positive ions very readily, such as potassium, reacts vigorously.
3. A metal whose atoms form positive ions only with difficulty, such as copper, reacts slowly or not at all.
4. Placing metals in order of this tendency gives the reactivity series: potassium, sodium, lithium, calcium, magnesium, aluminium, zinc, iron, copper, silver and gold.
   1. Hydrogen is often placed between iron and copper as a reference point, because metals above it react with dilute acids to release hydrogen.

Reactions with water rank the most reactive metals

1. A reactive metal reacts with water to form a metal hydroxide and hydrogen: $$2K(s)+2H_{2}O(l)\to 2KOH(aq)+H_2(g)$$
2. Potassium reacts so vigorously that the hydrogen catches fire, and sodium fizzes rapidly as it melts into a ball.
3. Lithium and calcium fizz steadily, and calcium forms a cloudy solution because calcium hydroxide is only slightly soluble.
4. Magnesium reacts only very slowly with cold water, giving a few bubbles over several days.
5. Zinc, iron and copper show no visible reaction with cold water.
6. Water therefore separates the most reactive metals well but cannot rank the less reactive ones, which need another test.

Reactions with dilute acids rank the middle metals

1. A metal above hydrogen reacts with a dilute acid to form a salt and hydrogen: $$Mg(s)+2HCl(aq)\to {MgCl}_2(aq)+H_2(g)$$
2. With the same dilute acid, magnesium fizzes vigorously, zinc fizzes steadily and iron gives slow bubbles.
   1. The faster the bubbles of hydrogen form, the more readily the metal forms positive ions.
3. Zinc reacts with dilute sulfuric acid in the same way: $$Zn(s)+H_2{SO}_4(aq)\to {ZnSO}_4(aq)+H_2(g)$$
4. Copper, below hydrogen, shows no reaction with dilute hydrochloric or sulfuric acid.
5. The gas can be confirmed as hydrogen because it burns with a squeaky pop when a lit splint is held at the mouth of the tube.

More reactive metals displace less reactive metals from salts

1. A metal that forms positive ions more readily takes the place of a less reactive metal in its compound.
2. Magnesium displaces copper from copper(II) sulfate solution: $$Mg(s)+{CuSO}_4(aq)\to {MgSO}_4(aq)+Cu(s)$$
   1. The blue colour fades, a pink-brown coating of copper forms on the magnesium and the solution gets warmer.
3. Copper does not react with magnesium sulfate solution, because copper is less reactive than magnesium.
4. Magnesium also displaces iron from iron(II) sulfate solution, and the iron that forms can be picked up with a magnet wrapped in cling film.
5. The further apart two metals are in the series, the larger the temperature rise when they react.
6. Displacement also ranks metals that do not react with dilute acids, since copper displaces silver from silver nitrate solution.

Deducing a reactivity order from fair comparisons

1. Use the same size and form of each metal, such as equal-mass strips or powders, because surface area also changes how fast a metal reacts.
2. Keep the volume and concentration of the acid or solution, and the starting temperature, the same for every metal.
3. Change only the metal, so that any difference in the results comes from its reactivity.
4. Rank the metals using a measured result, such as the temperature rise, the volume of hydrogen collected in the first minute or the number of displacement reactions each metal gives.
5. The metal with the most reactions, the largest temperature rise or the fastest gas production goes at the top of the order.