Chapter 8: Reactivity Trends

Question 1

Describe the reactivity of Group 2 metals and their consequent abundance in elemental form.

EXTENSION: Why are they called alkaline earth metals?

Answer 1

The elements in Group 2 of the periodic table are reactive metals and do not occur in their elemental form naturally.

EXTENSION: The name comes from the alkaline properties of the metal hydroxides.

Question 2

• Compare the electron configuration of Group 2 elements to noble gases.
• Describe the general redox reaction of Group 2 elements.
• What is the reducing agent in this reaction?

Answer 2

• Each Group 2 element has two outer shell electrons in its outer s-sub-shell, two more than the electron configuration of a noble gas.
• In a redox reaction, each metal atom is oxidised, losing two electrons to form a 2+ ion with the electron configuration of a noble gas.
• Another species will receive these two electrons and be reduced. The Group 2 element is called a reducing agent because it has reduced another species.

Question 3

Describe the redox reaction of Group 2 elements with oxygen.
Give the general formula and comment on which reactant is oxidised and which is reduced.

Answer 3

• The Group 2 elements all react with oxygen to form a metal oxide, with the general formula MO, made up of M²⁺ and O^2– ions.
  
  • The metal is oxidised, with a positive change in oxidation number. It is the reducing agent.
  • The oxygen is reduced, with a negative change in oxidation number.

Question 4

1. Describe the redox reaction of Group 2 elements with water.
2. Comment on the trend in reactivity.
3. State the balanced equation for the reaction between Sr and H₂O. Describe and explain the changes in oxidation number.

Answer 4

1. The Group 2 elements react with water to form an alkaline metal hydroxide [M(OH)₂] and hydrogen gas [H₂].
2. Magnesium and water react very slowly, but the reacion becomes more and more vigorous with metals further down the group – reactivity increases down the group.
3. Image attached.

Question 5

1. Describe the redox reaction of Group 2 metals with dilute acids.
2. Describe the reactivity trend in Group 2 for this reaction.
3. State the balanced equation for the reaction between Mg(s) and HCl(aq). Describe and explain the changes in oxidation number.
4. What observations could be made in a reaction of magneisum ribbon with dilute hydrochloric acid in a test tube?

Answer 5

1. The Group 2 elements react with dilute acids to form a salt and hydrogen.
2. The reactivity increases down the group.
3. Image attached.
4. The magnesium would react with the dilute hydrochloric acid to give off tiny bubble (of hydrogen).

Question 6

Why does the reactivity of Group 2 elements increase down the group? Explain.

Comment on how they react and the reactivity’s relationship to ionisation energy.

Answer 6

The atoms of Group 2 elements react by losing electrons to form 2+ ions. Therefore the formation of 2+ ions from gaseous atoms requires the input of two ionisation energies.

The ionisation energies decrease down the group. This is because the attraction between the nucleus and the outer electrons decreases as a result of increasing atomic radius and increasing shielding.

The total energy input to form 2+ ions decreases down the group. Hence, the Group 2 elements become more reactive and stronger reducing agents down the group.

Question 7

Describe the reaction of water with oxides of Group 2 elements. Use the example of calcium oxide reacting with water.

Answer 7

The oxides of Group 2 elements react with water, releasing hydroxide ions [OH^–] and forming alkaline solutions of the metal hydroxide.

CaO(s) + H₂O(l) → Ca²⁺(aq) + 2OH^–(aq)

The Group 2 hydroxides are only slightly soluble in water. When the solution becomes saturated, any further metal and hydroxide ions will form a solid precipitate.

Ca²⁺(aq) + 2OH^–(aq) → Ca(OH)₂(s)

Question 8

Describe the trends in solubility of Group 2 hydroxides. Compare Mg(OH)₂(s) to Ba(OH)₂(s).

Answer 8

The solubility of water increases down the group, so the resulting solutions contain more OH^–(aq) ions and are more alkaline.

• Mg(OH)₂(s) is only very slightly soluble in water. The solution has a low OH^–(aq) concentration and a pH≈10.
• Ba(OH)₂(s) is much more soluble in water. The solution has a high OH^–(aq) concentration and a pH≈13.

Question 9

Describe an experiment to show the trend in solubility of Group 2 oxides.

Answer 9

1. Add a spatula of each Group 2 oxide to water in a test tube.
2. Shake the mixture. On this scale, there is insufficient water to dissolve all of the metal hydroxide that forms. You will have a saturated solution of each metal hydroxide with some white solid undissolved at the bottom of the test-tube.
3. Measure the pH of each solution. The alkalinity wil be seen to increase down the group.

Question 10

Describe the uses of Group 2 oxides, hydroxides and carbonates.

Answer 10

• Agriculture. Calcium hydroxide is added to fields as lime by farmers to increase the pH of acidic soils. This neutralises acid in the soil, forming neutral water.
  
  • Ca(OH)₂(s) + 2H⁺(aq) →→ Ca²⁺(aq) + 2H₂O(l)
• Medicine. Group 2 bases are often used as antacids for treating acid indigestion. Many indigestion tablets use magnesium and calcium carbonates as the main ingredients. Milk of magnesia – a suspension of white magnesium hydroxide in water – can also be used.
  
  • Mg(OH)₂(s) + 2HCl(aq) →→ MgCl₂(aq) + 2H2O(l)
  • CaCO₃(s) + 2HCl(aq) →→ CaCl₂(aq) + H₂O(l) + CO₂(g)

Question 11

Describe the characteristic physical properties of halogens.

Answer 11

The halogens, Group 17 (7) of the periodic table, are the most reactive non-metallic group. The elements do not occur in their elemental form in nature. On Earth, the halogens occur as stable halide ions dissolved in sea water or combined with sodium or potassium as solid deposits.

At RTP, all the halogens exist as diatomic molecules. The group contains elements in all three physical states at RTP, changing from gas to liquid to solid down the group. In their solid states, the halogens form lattices with simple molecular structures.

Question 12

Explain the trend in the boiling point of halogens.

Answer 12

The boiling point of halogens increases down the group. This is because, as you go down Group 7

• there are more electrons
• there are stronger London forces
• more energy is required to break the intermolecular forces
• the boiling point increases.

Question 13

Explain the redox reactions of halogens.

Answer 13

Redox reactions are the most common type of reaction of the halogens. Each halogen atom is reduced, gaining one electron to form a 1– halide ion with the electron configuration of the nearest noble gas.

Another species loses electrons to halogen atoms – it is oxidised. The halogen is called an oxidising agent because it has oxidised another species.

Question 14

Describe the reactivity of halogens. Refer to halogen-halide displacement reactions.

Answer 14

The results of halogen-halide displacement reactions show that the reactivity of the halogens decreases down the group.

Question 15

How is a halogen-halide displacement reaction carried out?

Answer 15

A solution of each halogen is added to aqueous solutions of the other halides. If the halogen added is more reactive the the halide present

• a reaction takes place, the halogen displacing the halide from solution
• the solution changes colour.

Question 16

1. Why is cyclohexane added to solutions of halogens in water?
2. Describe the colours of halogen solutions in water and cyclohexane.

Answer 16

1. Solutions of iodine and bromine in water can appear a similar orange-brown colour, depending on the concentration. To tell them apart, an organic non-polar solvent such as cyclohexane can be added and the mixture shaken as the non-polar halogens dissolve more readily in cyclohexane than in water.
2. Image attached

Question 17

Give the full and ionic equation of the redox reaction between aquesous solutions of chlorine and sodium bromide. What is the change in oxidation number.

Answer 17

Question 18

Why are fluorine and astatine not considered in halogen-halide displacement reactions?

Answer 18

• Fluorine is a pale yellow gas, reacting with almost any substance that it comes in contact with.
• Astatine is extremely rare because it is radioactive and decays rapidly.

Question 19

Explain the reactivity trend of Group 7 elements.

Answer 19

In redox reactions, halogens react by gaining electrons. As you go down the group

• the atomic radius increases
• there are more inner shells so shielding increases
• there is less nuclear attraction to capture an electron from another species
• reactivity decreases.

In the halogens, fluorine is the strongest oxidising agent, gaining electrons from other species more rapidly than the other halogens. The halogens become weaker oxidising agents down the group.

Question 20

What is disproportionation? Give two examples.

Answer 20

Disproportionation is a redox reaction in which the same element is both oxidised and reduced. The reaction of chlorine with water and with cold, dilute sodium hydroxide are to examples of disproportionation reactions.

Question 21

Describe and explain the disproportionation reaction between chlrine and water.

Answer 21

When small amounts of chlorine are added to water, a disproportionation reaction takes place. For each chlorine molecule, one chlorine atom is oxidised and the other chlorine atom is reduced.

Question 22

1. Describe how bacteria are killed by the products from the disproportionation reaction between chlorine and water.
2. How can the bleachening effect of chloric(I) acid be demonstrated?

Answer 22

1. The two products of the reaction are chloric(I) acid [HClO) and hydrochloric acid [HCl]. The bacteria are killed by chloric(I) acid and chlorate(I) ions [ClO^–].
2. By adding some indicator to a solution of chlorine and water. The indicator first turns red, from the presence of two acids. The colour then disppears as the bleaching action of chloric(I) acid takes effect.

Question 23

Describe and explain the disproportionation reaction between chlorine and cold, dilute aqueous sodium hydroxide.

Answer 23

The reaction of chlorine and water is limited by the low solubility of chlorine in water. If the water contains dissolved sodium hydroxide, much more chlorine dissolves and another disporportionation reaction takes place.

The resulting solution contains a large concentration of chlorate(I) [ClO^–] ions from the sodium chlorate [NaClO] that is formed.

Question 24

What are the benefits and risks associated with chlorine use?

Answer 24

Although chlorine is beneficial in ensuring that our water is fit to drink and that bacteria are killed, chlorine is also an extremely toxic gas. Chlorine is a respiratory irritant in small concentrations, and large concentrations can be fatal.

Chlorine in drinking water can react with organic hydrocarbons such as methane, formed from decaying vegetation. Chlorinated hydrocarbons are formed, which are suspected of causing cancer.

Question 25

Describe and explain the test for carbonates.

Answer 25

Carbonates react with acids to form carbon dioxide gas. Na₂CO₃(aq) + 2HNO₃(aq) →→ 2NaNO₃(aq) + CO₂(g) + H₂O(l) This reaction forms the basis for a test for the carbonate ion, CO3^2–. 1. In a test tube, add dilute nitric acid to the solid or solution to be tested. 2. If you see bubbles, the unknown compound could be a carbonate. 3. Bubble the gas through lime water – a saturated aqueous solution of calcium hydroxide. 4. Carbon dioxide reacts to form a fine white precipitate of calcium carbonate, which turns the lime water cloudy. CO₂(g) + Ca(OH)₂(aq) →→ CaCO₃(s) + H₂O(l)

Question 26

Describe and explain the test for identifying sulfates.

Answer 26

Most sulfates are soluble in water, but barium sulfate [BaSO₄] is very insoluble. The formation of a white precipitate of barium sulfate is the basis for the sulfate test, in which equous barium ions are added to a solution of an unknown compound. The ionic equation for this is Ba²⁺(aq) + SO₄^2–(aq) →→ BaSO₄(s) The Ba²⁺(aq) ions are usually added as aqueous barium chloride or barium nitrate. Barium nitrate should be used if a halide test may be carried out after the sulfate test, as barium chloride introduces chloride ions to the solution.

Question 27

Describe and explain the test for halides.

Answer 27

Most halides are soluble in water, but silver halides are insoluble. Aqueous silver ions react with aqueous halide ions to form precipitates of silver halides. The ionic equation for this is Ag⁺(aq) + X^–(aq) →→ AgX(s) This reaction forms the basis of a halide test: 1. Add aqueous silver nitrate [AgNO₃] to an aqueous solution of a halide. 2. The silver halide precipitates are different colours – silver chloride is white, silver brominde is cream, and silver iodide is yellow. 3. Add aqueous ammonia to test the solubility of the precipitate. This stage is useful in distincting the different halides from each other.

Question 28

What is the sequence of tests if you are asked to analyse the anion(s) in an unknown inorganic compound?

Answer 28

For anions, the correct order for tests is: (1) carbonate [CO₃^2–] (2) sulfate [SO₄^2–] (3) halides [Cl^–, Br^– and I^–]

Question 29

Why is there a correct order to test for anions?

Answer 29

* **Carbonate test** In the carbonate test, you add a dilute acid and are looking for effervescence from carbon dioxide gas. Neither sulfate nor halide ions produce bubbles with dilute acid. The carbonate test can be carried out without the possibility of an incorrect conclusion. * **Sulfate test** In the sulfate test, you add a solution containing Ba²⁺(aq) ions and are looking for a white precipitate of BaSO₄(s). Barium carbonate [BaCO₃] is also white and insoluble which is why it is important to carry out the carbonate test first and to proceed to the sulfate test when it is known that there is no carbonate present. * **Halide test** In the halide test, you add a solution containing Ag⁺(aq) ions, as AgNO₃(aq), and are looking for a precipitate. Silver carbonate [Ag₂CO₃] and silver sulfate [Ag₂SO₄] are both insoluble in water and will form as precipitates in this test. It is therefore important to carry out the halide test last.

Question 30

Describe and explain the steps to carry out if you are asked to analyse the anions in a mixture of chemicals.

Answer 30

1. Carry out the carbonate test. If you see bubbles, continue adding dilute nitric acid until the bubbling stops. All carbonate ions will then have been removed and there will be none left to react in the next tests. 2. To the solution left from the carbonate test, add an excess of barium nitrate(aq). Any sulfate ions present will precipitate out as barium sulfate. 3. Filter the solution to remove the barium sulfate. 4. To the solution left from the sulfate test, add AgNO₃(aq). 5. Any carbonate or sulfate ions initially present have already been removed so any precipitate formed must involve halide ions. 6. Add NH₃(aq) to confirm which halide you have.

Question 31

Describe and explain the test for the ammonium ion.

Answer 31

When heated together, aqueous ammonium ions [NH₄⁺] and aqueous hydroxide ions [OH^–] react to form ammonia gas [NH₃]. NH₄⁺(aq) + OH^–(aq) →→ NH₃(g) + H₂O(l) This reaction forms the basis for a test of the ammonium ion. 1. Aqueous sodium hydroxide [NaOH] is added to a solution of an ammonium ion. 2. Ammonia gas is produced. You are unlikely to see gas bubbles as ammonia is very soluble in water. 3. The mixture is warmed and ammonia gas is released. 4. You may be able to smell the ammonia, but it is easy to test the gas with moist pH indicator paper. Ammonia is alkaline and its presence will turn the paper blue.