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Unit 8: Periodicity

8.1 General structure of the Periodic Table. The Periodic Table has changed over the centuries before it came to be what it is today. Chemists had always attempted to group elements in some way or other. The first attempts to classify elements were forwarded by a French nobleman, Antoine Lavoisier. Today’s Periodic Table is based on the work of the Russian Dmitri Mendeleev.

When one looks at a modern Periodic Table, one observes that chemicals with similar chemical properties are found in columns. These columns are called groups (or families). There are 8 groups of elements - Groups 1 - 7 and the last group, Group 0 or 8.

The rows in the Periodic Table are also significant. These rows are called periods. There are 7 periods going down the Periodic Table.

Electronic structure and the Periodic Table Look at the following table:

Number of outer shell electrons 1 2 3 4 5 6 7 lithium beryllium boron carbon nitrogen oxygen fluorine sodium magnesium aluminium silicon phosphorus sulfur chlorine

One should note that the elements in Group 1 all have one outer electron, the elements in Group 2 have 2 outer electrons, and that in Group 7 elements have 7 outer electrons.

For all groups, except Group 0, all the elements in that group have the same number of outer electrons.

Hence if you know that an element, like oxygen, has 6 outer electrons, then it belongs to Group 6.

Reactivity and electron structure The chemical properties of an element are linked to the number of outer electrons.

The nucleus contains protons which are positively charged. Electrons are negatively charged. Hence one would expect that electrons are attracted to the nucleus by attractive forces. This force is called electrostatic force.

When one looks at a group of elements in the Periodic Table, one notes that going down a group, the atoms of the elements get larger, e.g. in Group 1, the K atom is bigger than the Li atom.

We have already discussed that atoms try to gain a noble gas configuration so that they become stable. Let’s look at some groups more closely.

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In Group 1, the elements have one outer electron, and they need to lose this electron to achieve a noble gas configuration.

Since electrons are held to the nucleus by a strong electrostatic force, it is not easy for the atom to lose an electron. The smaller the atom will be, the closer the electrons are to the nucleus and hence the more difficult it will be for the atom to lose its outermost electron.

Following this reasoning, a larger atom will lose its electron more easily than a smaller atom. So a larger atom would be more reactive than a smaller atom. Therefore going down Group 1, the elements are more reactive.

In Group 2, elements need to lose 2 electrons to gain a noble gas configuration. As in Group 1, the reactivity of the elements increases down the group, as the atoms get bigger.

How do the elements in Group 1 compare with those in Group 2 with respect to reactivity?

In the same period, Group 2 elements have more protons in their nuclei. Therefore they have a higher positive charge, and the electrons are held more strongly. Also, Group 2 elements have to lose 2 electrons and not 1. This requires more energy.

These two factors make Group 2 elements less reactive than Group 1 elements in the same period. For example, Li is more reactive than Be; Na is more reactive than Mg.

In other groups where elements have to gain electrons to obtain a noble gas configuration such as in Group 7, the opposite for what we said about Groups 1 and 2 is true.

In group 7, since electrons have to be added on, the smaller the atom, the greater the electrostatic force will be and the easier it will be for the electron to be added on.

Hence in such groups reactivity decreases as the atom gets larger, and hence reactivity decreases down these groups.

Finally note these points:

• Across a period valency increases half-way and then decreases again. • Across a period metallic properties change into non-metallic properties. • Across a period reactivity decreases until half-way through the period, and then increases again. • Across a period elements change from existing as ions (charged particles) (e.g. Na+ ions in

metallic sodium) to covalent molecules (e.g. Cl2).

Metalloids Metalloids are elements that are found near the border that separates the metals from the non- metals. One should note that on the left hand side, the elements have a very strong metallic character while the elements to the right hand side have a strong non-metallic character. The elements that are close to the above mentioned border have both metallic and non-metallic character and so are called metalloids. The following Periodic Table shows which elements are metalloids:

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8.2 Trends in properties across a typical period (Na to Ar).

Valency Note that as one goes from left to right in any period, the valency of the elements increases up to group 4 and then decreases as shown in the following table:

Group 1 Group 2 Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Na Mg Al Si P S Cl Ar 1 2 3 4 3 2 1 0

Oxides One should also note that the oxides of elements across a period follow a particular pattern. Metallic elements form basic oxides; non-metallic elements form acidic oxides while metalloids form amphoteric oxides.

Group 1 Group 2 Group 3 Group 4 Group 5 Group 6 Group 7 Group 8 Na Mg Al Si P S Cl Ar basic basic amphoteric acidic acidic acidic acidic No oxide

Type of bonding The type of bonding that elements form on reacting is also determined by their position in the Periodic Table. One should note that when metals react with non-metals, electrovalent or ionic bonding is formed. On the other hand when non-metals react with each other they undergo covalent bonding.

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Work Out!

By the end of the previous topics you should be able to:

• appreciate that the Periodic Table is a classification of elements in order of increasing atomic number and its use to predict properties of elements • understand the divisions of the Periodic Table; distinguish between a Group (as a vertical column of elements that show similar properties) and a Period (as a row of elements that show a gradual change from metallic to non-metallic character across the period) • describe characteristic properties of metals and non-metals and be aware that some elements exhibit a mixture of properties of metals and non-metals; classify an element as ‘metal’ or ‘non-metal’ on the basis of its properties • identify the following families of elements and recall their positions in the Periodic Table: the alkali metals, alkaline earth metals, transition metals, halogens, noble gases • make the relationship between group number, number of valency electrons and metallic/non-metallic character; understand the correlation between the electron configuration of elements, and charges on ions, to the position of an element in the Periodic Table • utilise the sequence in the valencies of the elements across a period to predict formulae of common compounds of the elements • recall that the oxides of the elements in Period 3 change from basic, through amphoteric to acidic • relate the similarity in chemical properties of elements to the number of electrons in the outer shell • state that the reactivity of the elements in group 1 and 2 increases going down the group; • describe reactions or experiments to support this statement • state that the reactivity of the elements in group 7 decreases going down the group; describe reactions or experiments to support this statement • utilise the trends in the Periodic Table to predict properties of other elements, or their compounds, found in the same group or area of the Periodic Table

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8.3 Trends down typical metallic groups

The alkali metals The most common elements in this group are lithium (Li), sodium (Na) and potassium (K). They are all VERY reactive elements and in fact they have to be stored under oil to prevent them from reacting with air.

There are three other elements in this group: rubidium, caesium and francium, but these metals are too reactive to be used in the lab.

Group 1 metals are soft metals and can be cut with a knife. When they are cut, a shiny surface is revealed.

From the table below, note that they have low densities, as well as low melting and boiling points. Also, they are very good conductors of electricity.

Reaction of alkali metals with water All alkali metals react vigorously with water. If a very small piece of Li, Na or K is dropped in a trough containing water, a very vigorous reaction takes place.

All three metals float on water and react with it to give colourless hydrogen gas. In all three cases enough heat is evolved to melt the metals. In the case of potassium, because of the heat produced, the hydrogen may ignite and burns with a pink flame.

Besides producing hydrogen, these metals also produce an alkaline solution when they react with water (hence their name).

Although all three elements react vigorously with water, K reacts more vigorously than Na which reacts more vigorously than Li.

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To summarise,

alkali metal + water → metal hydroxide (alkali) + hydrogen gas

2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)

Reaction of alkali metals with oxygen Alkali metals burn in oxygen, to form solid white oxides. Each alkali metal burns with a characteristic flame colour. Lithium burns with a red flame, sodium with a yellow flame and potassium with a lilac flame.

Again, K is more reactive with oxygen than Na which is in turn more reactive than Li.

To summarise,

alkali metal + oxygen → metal oxide

4K(s) + O2(g) → 2K2O(s)

Reaction of alkali metals with non-metals Alkali metals react readily with non-metals to form salts. All alkali metal salts are soluble in water. Reactivity depends on both the position of the alkali metal and the position of the non-metal with which the alkali metal is reacting.

For example, sodium reacts with chlorine to form sodium chloride.

To summarise:

alkali metal + non-metal → salt

2Na(s) + Cl2(g) → 2NaCl(s)

This type of reaction is also called a synthesis reaction as a compound is being formed from direct combination of its elements.

Work Out

1. Which elements are known as the alkali metals? Why are the chemical properties of these elements, all very similar? 2. Sodium chloride is an ionic compound. What does this mean? 3. Complete the following table:

Name Formula Colour in flame test Effect of heat Sodium chloride Yellow No reaction

Na2CO3

Potassium carbonate K2CO3 → LiOH

NaHCO3

4. Complete the following chemical equations, make them balance, and name the products:

a. K(s) + H2O(l) →

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b. Li(s) + Cl2(g) →

c. KOH(aq) + CuSO4(aq) →

d. Na2CO3(s) + H2O(l) + CO2(g) → 5. How could you tell the difference between sodium hydrogencarbonate and anhydrous sodium carbonate? 6. Write a balanced chemical equation for the reaction and name the products: e. Lithium + water f. Lithium burning in oxygen g. Potassium burning in oxygen h. Sodium carbonate + hydrochloric acid

Alkaline earth metals The alkaline earth metals are found in Group 2 in the Periodic Table and include beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba) and radium (Ra); the most common are Mg and Ca.

All the metals in the group have two outer electrons and need to lose these two electrons to achieve a noble gas configuration (so they have a valency of 2).

They are harder than the Group 1 metals and although they are silver-grey in colour, they tend to tarnish quickly when left exposed to air, because they are so reactive that they quickly form a layer of oxide on their surface.

Like the alkali metals they are good conductors of heat and electricity.

Reaction of the alkaline earth metals with oxygen Like alkali metals, alkaline earth metals react with oxygen but at a slower rate. For this reason, it is not necessary to store them under oil but one must note that they do tend to form a layer of oxide on standing in air.

For example, calcium reacts with the oxygen in air to form calcium oxide.

To summarise:

alkaline earth metal + oxygen → alkaline earth metal oxide

2Ca(s) + O2(g) → 2CaO(s)

Reaction of alkaline earth metals with water Both calcium and magnesium react with water but less vigorously than the alkali metals.

Before reacting magnesium with water, magnesium must be cleaned with emery paper or dilute acid to remove the oxide layer that forms upon it.

When clean magnesium is placed in cold water, a few bubbles of hydrogen are formed slowly on its surface.

Mg(s) + 2H2O(l) → Mg(OH)2(s) + H2(g)

Calcium reacts in a similar manner but more vigorously since reactivity increases down the group.

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Ca(s) + 2H2O(l) → Ca(OH)2(s) + H2(g)

Since the hydroxides formed are slightly soluble in water, they both give an alkaline solution (hence their name) but as the concentration of hydroxide increases in the solution, a white precipitate of the hydroxide is formed.

Reaction of alkaline earth metals with acids Both calcium and magnesium react with dilute acids to give the corresponding salt and hydrogen (in other words they displace hydrogen from the acid).

Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)

Mg(s) + H2SO4(aq) → MgSO4(aq) + H2(g)

Like in Group 1, one observes that as you go down the group the reactivity increases (i.e. Ca is more reactive than Mg).

Alkaline earth metal carbonates 1. Magnesium carbonate occurs as magnesite and in association with calcium carbonate as dolomite. Calcium carbonate exists in two crystalline forms: calcite, (e.g. limestone, chalk, marble); and aragonite, found in coral shells. Limestone and dolomite are both very important industrially as they are used in various industrial processes, e.g. extraction of iron and the Solvay process (the industrial process used to make sodium carbonate).

2. Both carbonates are sparingly soluble in water and on strong heating they decompose to give the oxide and carbon dioxide. Like all carbonates, they also react readily with dilute acids to give the salt, carbon dioxide and water.

3. Magnesium carbonate is used to produce magnesium oxide which is used in steel furnaces.

4. Calcium carbonate which is called limestone has many uses, such as: • it is a basic raw material in the Solvay Process. • it is used to make cement. Cement is made by heating limestone and clay. Cement is then usually mixed with rubble, sand and water to form concrete. • it is also used in the glass industry, as well as to make quicklime (“ġir” in Maltese) and slaked lime.

When heated CaCO3 decomposes to form the calcium oxide (CaO) which is called quicklime and carbon dioxide:

ℎ푒푎푡 CaCO3(s) → CaO(s) + CO2(g)

When water is added to quicklime, a chemical reaction happens and slaked lime is formed. Slaked lime is what whitewashers use to whitewash walls. It is slightly soluble in water so this solution would be alkaline.

CaO(s) + H2O(l) → Ca(OH)2(s)

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As the whitewash stands on the wall, in time it reacts with carbon dioxide in the air to form calcium carbonate.

Ca(OH)2(s) + CO2(g) → CaCO3(s) + H2O(l)

Commercially produced compounds of alkali and alkaline earth metals Washing soda is used to soften hard water by removing the calcium and magnesium soluble ions from water.

Epsom salts consist of hydrated magnesium sulfate (MgSO4.7H2O). They are used for purgative purposes and for other medical uses.

Plaster of Paris is hydrated calcium sulfate that has lost three-quarters of its water of crystallisation. When water is added to Plaster of Paris, it heats up, expands somewhat and quickly sets. It is used to make casts and to cover wall surfaces and ceilings.

Milk of magnesia contains magnesium hydroxide, Mg(OH)2; this is basic in nature and is used to counterattack excess acidity in the stomach.

Work Out

1. Make a list of 6 calcium and magnesium compounds, giving their chemical name and formulae and also where possible, their common names, occurrence and uses. 2. Describe with aid of labelled diagrams, an experiment demonstrating the reaction between a. Magnesium and water b. Magnesium and steam Give the chemical equation for each reaction. 3. Outline the uses of: a. Limestone b. Quicklime c. Lime water d. Slaked Lime 4. Why is slaked lime sometimes known as hydrated lime? 5. Finely divided limestone is known as agricultural lime since it is used to reduce the acidity in the soil. Explain how the acidity is reduced. 6. Draw the electron structure of a) calcium, b) magnesium. 7. Write an ionic equation for the ionisation of calcium. 8. Write a balanced chemical equation for each reaction and name the products:

a. MgCO3(aq) + H2SO4(aq) →

b. Ca(s) + O2(g) → c. Mg(s) + ZnO(s) → d. Ca(s) + HCl(aq) → 9. Alkaline earth metals are reducing agents. Explain why and give two examples.

By the end of this topic you should be able to:

Group 1 ~ the Alkali metals; Li, Na and K

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• recall the typical ‘non-metallic’ physical properties (soft, low density, etc.) of these metals • utilise the results of reacting lithium, potassium and sodium with water to show their similarity as a ‘family’ of elements and confirm their order of reactivity • recall that the oxides of the alkali metals are basic, and the aqueous hydroxides are alkaline; that these are neutralised by acids to form salts • predict the properties and reactivity of other metals in this group • predict the general patterns of behaviour of the common compounds of these elements by making links to other related topics in the syllabus

Group 2 ~ the alkaline earth metals; Mg and Ca

• describe and understand the reactions of magnesium and calcium with oxygen, with water and • dilute hydrochloric acid; magnesium with steam; utilise the results of these reactions to show the similarity of the elements and confirm their order of reactivity • recall that limestone is a naturally-occurring form of calcium carbonate and explain the link between limestone districts and hardness of water • describe the thermal decomposition of calcium carbonate to produce calcium oxide • describe the effect of water on calcium oxide and recall that the solution produced is alkaline • understand that oxides and hydroxides of calcium, that appear to be insoluble, are in fact ‘sparingly soluble’ • explain why calcium oxide and calcium hydroxide are used to neutralise soil acidity • describe the reactions by which the common salts of magnesium and calcium can be prepared by making links to the general methods of preparing salts covered in Unit 5, Topic 5.4 • make predictions about the properties and reactivity of other metals in this group • predict the general patterns of behaviour of other common compounds of these elements by making links to related topics in the syllabus

8.4 The transition metals Iron (Fe) and copper (Cu) are two examples of a group of metals called the transition metals.

Transition metals, which are found in between Groups 2 and 3 of the Periodic Table, have properties that differentiate them from other metals. These include:

a) they have multiple valencies b) they have high densities, melting and boiling points c) transition metals undergo oxidation and reduction very readily, and so they are good catalysts d) they also adsorb (accumulate on a surface) gases onto their surface very easily and this makes them good catalysts in gaseous reactions e) they form coloured complex ions

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In this topic we will be looking more closely at two examples of these transition metals - iron and copper.

Iron Iron is the second most abundant metal after aluminium; it forms 4% of Earth’s crust.

Iron is found mostly in the form of ores. The most common ores are the iron oxides (Fe2O3, called haematite and Fe3O4, called magnetite) and iron(II) carbonate (Fe2CO3). The latter is roasted in air to drive off water and the carbonate loses CO2 and forms the oxide.

Extraction of iron A blast furnace is used in the process of extracting iron from its ores. The blast furnace is loaded with limestone (calcium carbonate), coke (nearly pure carbon) and iron ore (usually Fe2O3).

1. Hot air is blasted in through holes at the bottom. This causes the carbon (in the coke) to burn:

C(s) + O2(g) → CO2(g)

2. Also the limestone starts to decompose:

ℎ푒푎푡 CaCO3(s) → CaO(s) + CO2(g)

3. The CO2 reacts with more hot coke to produce carbon monoxide:

CO2(g) + C(s) → 2CO(g)

4. Carbon monoxide rises to the furnace and reduces the iron(III) oxide ore by taking oxygen from it. The temperature required for this to happen is around 700 oC.

Fe2O3(s) + 3CO(g) → 2Fe(l) + 3CO2(g)

5. The molten iron trickles to the bottom of the furnace where it is tapped off.

6. The calcium oxide formed when the limestone decomposes (see step 2), reacts with any earth or sandy materials in the ore to form liquid slag, which is mostly calcium silicate.

CaO(s) + SiO2(s) → CaSiO3(l) (slag)

Calcium silicate also trickles down to the bottom of the furnace but since it is less dense than the molten iron, it floats on top of it. Slag is a waste product and is used in road building.

7. The waste gases which are mainly nitrogen and oxides of carbon, (e.g. CO2 and CO) escape from the top of the furnace and are used to heat incoming air. This helps reduce heating costs in the furnace.

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8. The iron produced in the blast furnace is called “pig” or cast iron and contains impurities (some 4% carbon).

Iron is mostly converted into alloys, e.g. stainless steel.

Uses of iron 1. Pig iron is brittle and hard and does not have many uses. It is sometimes used to make gas cylinders, stoves, hot-water radiators, bases of Bunsen burners and other metals parts that need not be strong.

2. Wrought iron is the purest form of commercial iron. It is strong and malleable (can be worked and hammered into various shapes). It is used to make nails, chains, farm machines, cores of electromagnets because it cannot be permanently magnetised, etc.

3. Iron is used in the formation of alloys. An alloy is a mixture of metals. Steel is one important alloy made up of iron and carbon. There are various kinds of steels (e.g. stainless steel which also contains some chromium and nickel). Many items are made from steel, e.g. trains, ships, kettles etc.

Some reactions of iron

Action of steam on heated iron 1. Some iron powder is placed in the middle part of a combustion tube.

2. Water is boiled gently. Also, the combustion tube is warmed at the end where the steam comes in so that it does not condense.

3. The iron is heated and steam is passed over it slowly. This is controlled by controlling the amount of heating of the water. The iron changes to a blue-black compound, i.e. triiron tetraoxide (or iron (II, III) oxide). This is a mixed oxide which consists of 1 part FeO and 1 part

Fe2O3.

3Fe(s) + 4H2O(g) ⇌ 4H2(g) + Fe3O4(s)

Action of hydrochloric acid on iron 1. A piece of iron wire is placed in a test tube containing dilute hydrochloric acid. Bubbles of gas form on iron.

2. Iron reacts with dilute hydrochloric acid to form iron(II) chloride and hydrogen gas.

Fe(s) + 2HCl(aq) → FeCl2(aq) + H2(g)

Action of chlorine gas on iron 1. Rust-free iron wire is coiled around a pencil and placed in a combustion tube.

2. The reaction should take place in a fume cupboard as chlorine is poisonous.

3. Dry chlorine is passed through the apparatus to drive out any air.

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4. The iron is heated strongly and chlorine is passed over it. As soon as the iron glows red-hot and starts to react, the heating of iron is stopped.

5. The end of the combustion tube near the receiving bottle is warmed and black crystals sublime in the bottle.

2Fe(s) + 3Cl2(g) → 2FeCl3(s)

6. If instead of chlorine, hydrogen chloride gas is used, iron(II) chloride is formed.

Fe(s) + 2HCl(g) → FeCl2(s) + H2(g)

Iron(II) and iron(III) compounds

Hydroxides - Fe(OH)2 and Fe(OH)3 1. Iron(II) hydroxide is a green solid that is formed when sodium hydroxide is added to pure iron(II) sulfate. Both reactants have to be dissolved in freshly boiled water to remove any dissolved oxygen which tends to oxidise the iron(II) salt produced.

FeSO4(aq) + 2NaOH(aq) → Na2SO4(aq) + Fe(OH)2(s)

2. Iron(III) hydroxide is a reddish-brown solid that is formed by precipitation when sodium hydroxide and iron(III) sulfate are mixed together.

Fe2(SO4)3(aq) + 6NaOH(aq) → 3Na2SO4(aq) + 2Fe(OH)3(s)

Since both iron(II) and iron(III) hydroxide are basic, they react with acids to produce a salt and water only.

Fe(OH)2(s) + 2HCl(aq) → FeCl2(aq) + 2H2O(l)

Fe(OH)3(s) + 3HCl(aq) → FeCl3(aq) + 3H2O(l)

Also, Fe(OH)2 oxidises readily to the more stable Fe(OH)3 form in the presence of oxygen or oxidising chemical agents.

Chlorides - FeCl2 and FeCl3 1. Iron(II) chloride is a white solid which is formed by reacting hydrogen chloride gas with iron.

Fe(s) + 2HCl(g) → FeCl2(s) + H2(g)

2. Iron(III) chloride is a black solid that sublimes. It is formed by reacting chlorine with iron.

2Fe(s) + 3Cl2(g) → 2FeCl3(s)

Copper Copper was one of the earliest metals known to humans. The metal occasionally occurs free in nature, but it is also found with other elements in compounds such as copper pyrites (CuFeS2), malachite (CuCO3, Cu(OH)2), and cuprite (Cu2O).

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Copper metal is obtained industrially by heating copper oxide with carbon which gives a fairly pure source of copper metal since carbon reduces copper oxide to the metal.

CuO(s) + C(s) → Cu(s) + CO(g)

Since this form of copper is not pure enough for most purposes, copper is purified using electrolytic purification as detailed in the chapter on electrolysis.

The metal is comparatively soft, malleable, ductile and a good conductor of electricity.

The pure metal is reddish-brown in colour but it tarnishes in air and becomes covered with a greenish patina (i.e. a mixture of copper (II) sulfate and copper (II) hydroxide).

Native Copper Patina- copper pipes covered copper dome

The most familiar compounds are the copper(II) compounds, although copper(I) compounds do exist as well.

Action of nitric acid on copper Since copper is below hydrogen in the ECS (Electrochemical Series), it will not react with dilute acids unless they are also oxidising agents.

Since dilute nitric acid is such an acid, copper does react with it, liberating nitrogen monoxide gas.

3Cu(s) + 8HNO3(aq) → 3Cu(NO3)2(aq) + 2NO(g) + 4H2O(l)

If the nitric acid is concentrated the following reaction is given:

Cu(s) + 4HNO3(aq) → Cu(NO3)2(aq) + 2NO2(g) + 2H2O(l)

This reaction is important because it is a way by which copper(II) ions in solution and from which other copper(II) compounds can be produced.

Action of conc. sulfuric acid on copper Concentrated sulfuric acid is also an oxidizing agent and it oxidizes copper metal to form copper(II) sulfate. The following equation shows the equation for this reaction:

Cu(s) + 2H2SO4(l) → CuSO4(s) + SO2(g) +2H2O(l)

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However, one must note that dilute, non-oxidising acids such as dilute HCl and dilute H2SO4 do not react with copper metal.

Reduction of copper(II) oxide by hydrogen Copper(II) oxide is reduced to the metal by heating it with hydrogen since hydrogen is a reducing agent.

CuO(s) + H2(g) → Cu(s) + H2O(l)

If the hot copper is exposed to air it will immediately react with the oxygen in air to reform black copper(II) oxide. So the copper produced must be allowed to cool.

Thermal decomposition of copper(II) nitrate and carbonate These two salts decompose on heating like typical nitrates and carbonates, as shown by the following chemical equations:

ℎ푒푎푡 2Cu(NO3)2(s) → 2CuO(s) + 4NO2(g) + O2(g)

ℎ푒푎푡 CuCO3(s) → CuO(s) + CO2(g)

Copper(I) compounds Copper(I) compounds do exist, but most of these are unstable in the presence of dilute acids and water.

Copper(I) oxide is produced by reacting copper(II) sulfate, glucose and sodium hydroxide. Glucose is a “reducing” sugar and reduces the copper(II) ions to copper(I) ions. This is the chemical reaction that happens during Benedict’s test for reducing sugars.

Some uses of copper 1. Being highly malleable and ductile, it is shaped into wires, rods, etc. 2. It is a very good conductor of heat and electricity (only second to silver) and so it is used in boilers, kettles, electricity wires, dynamos etc. 3. It forms various important alloys such as brass, which is a mixture of copper and zinc (used to make screws and nuts) and bronze, which is a mixture of copper and tin.

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Work Out

By the end of this topic you should be able to:

• appreciate that the study of these two metals and their compounds illustrates typical transition metal properties • recall that transition metals have high melting points, high densities, a variable valency and their compounds are generally coloured as exemplified by iron and copper • recall specific examples of the use of transition metals and their compounds as catalysts • state that iron is obtained from iron ore and describe the extraction of iron by the process of • smelting in the blast furnace; including an outline diagram, and a description of the key reactions occurring in different parts of the blast furnace • recall and explain why iron from the blast furnace has limited uses; relate the uses of steel to its properties • apply the term rusting to the corrosion of iron; explain that rusting is an example of oxidation • describe the action of steam, hydrogen chloride and chlorine on iron • describe the precipitation of iron(ll) hydroxide and iron(lll) hydroxide • explain, (or identify descriptions of), the conversion of iron(ll) to iron(lll) compounds and vice versa. • identify and explain the oxidation/reduction reactions occurring • describe the oxidation of copper by using concentrated nitric acid • discuss the preparation, and simple chemical properties, of the common compounds of copper, limited to items specified in this section and related topics in the syllabus • state that an alloy is a mixture of metals, or of metals with non-metals, giving examples • interpret and explain a given reaction scheme, where the metal or its compounds is/are converted from one to the other

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8.5 Trends down a typical non-metal group.

The halogen group The family of elements in Group 7 of the Periodic Table is known as the halogens. The name halogen means “salt producer” since all of them produce ionic salts of the type M+X- (where M is a metal and X is the halogen, example: Na+Cl-).

Group 7 consists of fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At) in this order.

Each halogen has 7 electrons in its outer shell and it needs either to gain or share 1 electron to achieve a noble gas configuration.

The elements themselves exist as simple covalent molecules (e.g. Br2, Cl2), and may also form inter halogen compounds, e.g. ClBr (FI doesn’t exist).

Trends in physical properties 1. The melting points and boiling points increase down the group. A result of this is that fluorine and chlorine are gases; bromine is a liquid while iodine and astatine are solids at RTP.

2. The size of the atoms going down the group increases. This property is similar to the rest of the groups in the Periodic Table.

3. One also notes that the halogens have distinctive colours:

• Fluorine is pale yellow • Chlorine is pale green • Bromine is orange brown • Iodine is dark violet • Astatine is black and is highly radioactive.

Trends in chemical properties

Reaction of the halogens with water Chlorine and bromine dissolve in water and react with it to form acidic solutions.

Iodine does not react readily with water.

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Fluorine on the other hand is very reactive as a gas but forms a weak acid which is highly corrosive.

This shows that as one goes down the group the reactivity with water decreases.

Reaction of the halogens with hydrogen Fluorine reacts spontaneously with hydrogen without the need of sunlight.

Chlorine reacts spontaneously with hydrogen even in diffused sunlight. In direct bright sunlight the mixture explodes.

Sunlight is not enough to start off the reaction between bromine and hydrogen. The mixture must be heated to 473K (200 oC) before any reaction occurs.

Bromine needs to be heated substantially for it to react with hydrogen to form hydrogen bromide.

Iodine and hydrogen even when heated to 773K (500 oC) and in the presence of a catalyst, do not react completely.

This again shows that reactivity with hydrogen decreases down the group, i.e. as the halogen atoms become larger.

Reaction of the halogens with iron All halogens react with iron albeit at different rates as shown by experiment. As one goes down the group the order of reactivity decreases. This is evident during experiment since the reactions start out as being spontaneous with fluorine to ending with considerable amount of heat required for the reaction to happen with iodine.

Explanations and conclusions As one can see from the above reactions, halogens all react similarly and their reactivity decreases down the group. As one goes down the group, the atom is bigger, (i.e. the atomic number is bigger), and therefore there must be some link between the size of the atom and reactivity.

First of all, the halogens react similarly because they all have 7 outer electrons and need 1 electron to achieve a noble gas configuration.

Remember that as the atomic size increases:

a. the outer electron progressively gets further away from the attractive effect of the nucleus. b. there are an increasing number of completed electron shells, each of which “shields” the outer electron from the attractive effect of the nucleus. c. the positive charge of the nucleus increases.

The net effect of these three factors is that outer electrons are more easily lost and so less easily gained.

The easier halogens gain electrons, the more reactive they are, it is harder for large atoms to gain electrons, therefore the larger the halogen atom is, the less reactive it will be.

Unit 8: Periodicity Page 19 of 20

This explains why the order of reactivity for the halogens is: fluorine, then chlorine, then bromine and finally iodine.

Displacement of one halogen by another As stated earlier the reactivity of a halogen depends on the ability to gain an electron.

So, when chlorine gas is bubbled through a solution containing bromide ions, the chlorine “grabs” the electrons of the bromide ions and these are converted back into bromine gas molecules. In turn, the chlorine gas molecules become ions.

However, if iodine is mixed in a solution containing bromide ions, no reaction occurs because iodine is less reactive than bromine and cannot take away its ions.

Summarising:

- - Cl2(g) + 2Br (aq) → Br2(g) + 2Cl (aq)

- I2(s) + 2Br (aq) → no reaction

Fluorine displaces all the other halogens. Chlorine displaces bromine and iodine. Bromine displaces iodine. Iodine displaces none of the common halogens.

Note that oxidising ability decreases down the halogen group. Chlorine is a better oxidising agent than iodine. It is an oxidising agent because it causes iodide ions (say) to lose their electrons and form molecules (REMEMBER: Oxidation is LOSS of electrons).

Work Out

By the end of this topic you should be able to:

• recall the physical characteristics such as state at room temperature and variation in colour • show that they are a ‘family’ of elements by discussing similarities in their chemical properties • describe and explain the relative reactivity of the halogens as illustrated by halogen-halide displacement reactions; explain this relative reactivity of the elements as oxidising agents • predict the properties and relative reactivity of other halogens • state or predict typical properties of halides of common metals

Unit 8: Periodicity Page 20 of 20

8.6 The noble gases The noble gases are a group of chemical elements with very similar properties: under standard conditions, they are all odourless, colourless, monatomic gases, with a very low chemical reactivity. The six noble gases that occur naturally are helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and the radioactive radon (Rn).

The properties of the noble gases can be well explained by modern theories of atomic structure: their outer shell of valence electrons is considered to be "full", giving them little tendency to participate in chemical reactions.

Neon, argon, krypton, and xenon are obtained from air using the method of fractional distillation of liquid air. Noble gases have several important applications in industries such as lighting, welding, and space exploration. A helium-oxygen breathing gas is often used by deep-sea divers at depths of seawater over 180 feet (55 m) to keep the diver from experiencing oxygen toxaemia and nitrogen narcosis. After the risks caused by the flammability of hydrogen became apparent, it was replaced with helium in balloons.

Colours produced by different noble gases when used in neon tubes

(Taken from http://en.wikipedia.org/wiki/Noble_gas)

By the end of this topic you should be able to:

• state that the noble gases are a family of very unreactive elements • relate this lack of reactivity to their having full outer shells of electrons • recall the monoatomic nature of noble gases • relate some uses of noble gases to their physical properties and lack of chemical reactivity