Characteristics of Metals PDF

## Section C ### C17 #### Objectives - describe the physical properties of metals - relate the physical properties of metals to the way in which metal atoms are bonded in the metallic lattice. ### Characteristics of metals Most elements in the periodic table are metals. In fact, metals account for...

## Section C ### C17 #### Objectives - describe the physical properties of metals - relate the physical properties of metals to the way in which metal atoms are bonded in the metallic lattice. ### Characteristics of metals Most elements in the periodic table are metals. In fact, metals account for about 80% of all the elements in the table. Metallurgy, i.e. the study of metals, is one of the oldest applied sciences, dating back to 6000BC with the discovery of gold. Metals have a variety of physical and chemical properties which enable them to play an extremely important part in our daily lives. ### C17.1 Physical properties of metals - Metals are elements whose atoms have a small number of valence electrons, usually 1, 2 or 3. - Most metals are found in Groups I, II and III of the periodic table and between Groups II and III as the transition metals. - In Unit 5.4 you learnt that the type of bonding between metal atoms in a metal is known as metallic bonding and that metals have very distinct physical properties as a result of this bonding. - Recall that a metal lattice is composed of metal cations surrounded by a 'sea' of delocalised, mobile electrons and that the lattice is held together by the strong electrostatic force of attraction between the positive cations and the delocalised, negative electrons known as the metallic bond. - The way in which the metal atoms are bonded in metal lattices helps us to understand the physical properties of metals. #### Metals have high melting points and boiling points - The electrostatic forces of attraction between the positive cations and negative, delocalised electrons are relatively strong. - As a result, large amounts of heat energy are required to separate the atoms in order to melt or boil the metals. - All metals are found in the solid state at room temperature, with the exception of mercury which is a liquid. #### Metals are good conductors of electricity and heat - The delocalised electrons in the metal lattice are not held in any fixed position and are free to move. - These moving electrons act as charge carriers allowing an electric current to be carried through the metal. - They also act as heat carriers allowing heat to be carried through the metal when, for example, one end of a piece of metal is placed in a hot environment, such as a flame, the other end of the metal quickly heats up. #### Metals have a shiny lustre - Lustre is the shininess of a metal. Metals generally have a very shiny, metallic lustre. - This is because the mobile electrons reflect the photons of light back from the metal surface before they are allowed to penetrate far into the surface. #### Metals are hard, malleable and ductile - Metals are considered to be hard because they are not easily damaged when a force is applied to them. - They are, however, malleable meaning that they can be bent and hammered into different shapes. - They are also ductile, meaning that they can be drawn out into wires. - This is because the atoms in the metal are all of the same type, therefore they are all the same size. - If force is applied to the metal, the atoms can roll over each other into new positions without breaking the metallic bond. #### Metals have high densities - The density of any substance is calculated by dividing the mass of the substance by its volume. - Metals have high densities because their atoms are packed very closely together. - This means that there are as many metal atoms as possible in a given volume of metal. #### Summary questions 1. Explain why metals have high melting points and boiling points. 2. Why are metals good conductors of electricity and heat? 3. Explain the terms 'malleable' and 'ductile'. Explain why metals are malleable and ductile. 4. Why do metals have a high density? ### C17.2 Chemical properties and reactions of metals #### Objectives - describe the reactions of various metals with oxygen, water and dilute acids - name the products formed by the reactions - write balanced chemical equations for the reactions of various metals with oxygen, water and dilute acids. - Because metals are elements whose atoms have a small number of valence electrons, usually 1, 2 or 3, they form positive cations in chemical reactions by losing these valence electrons. - This can be summarised as follows: $M \rightarrow M^{n+} + ne^-$ - As a result, metals form ionic compounds when they react and the metal behaves as a reducing agent since it gives electrons to the other reactant, i.e. it causes the other reactant to gain electrons (RIG). - Some metals are relatively stable while others react violently when exposed to other substances, e.g. oxygen, water and acids. - The way in which a particular metal reacts gives an indication of its reactivity. - The reactivity of a metal is related to the metal's position in the periodic table. #### The reactions of metals with oxygen - Some metals can be exposed to air without reacting to any great extent with the oxygen in the air. - These metals include the ones that we use regularly and see around us. - The surface of other metals oxidises immediately on exposure to air and some metals react so vigorously with oxygen they have to be stored under paraffin, e.g. potassium and sodium. - The paraffin acts as a barrier, preventing the metals from coming into contact with the air. - When metals react with oxygen, they form ionic compounds known as metal oxides. - This can be summarised by the following general word equation: > metal + oxygen $\rightarrow$ metal oxide **Table 17.2.1** shows what happens when certain metals react with oxygen in the air. | Metal | Description of the reaction when exposed to dry air | Description of the reaction when heated in air | Products formed when heated and equation | |---|---|---|---| | Potassium | Reacts very readily forming potassium oxide | Burns vigorously with an orange flame | Forms a white powdery solid, potassium oxide: $4K(s) + O_2(g) \rightarrow 2K_2O(s)$ | | Sodium | Reacts very readily forming sodium oxide | Burns vigorously with a lilac flame | Forms a white powdery solid, sodium oxide: $4Na(s) + O_2(g) \rightarrow 2Na_2O(s)$ | | Calcium | Reacts readily to form coating of calcium oxide | Burns very easily with a brick red flame | Forms a white powdery solid called calcium oxide: $2Ca(s) + O_2(g) \rightarrow 2CaO(s)$ | | Magnesium | Reacts slowly to form a coating of magnesium oxide | Burns easily with a bright white flame | Forms a white powdery solid, magnesium oxide: $2Mg(s) + O_2(g) \rightarrow 2MgO(s)$ | | Aluminium | Reacts slowly to form a thin coating of aluminium oxide | Burns when heated strongly, especially if powdered | Forms a white powdery solid, aluminium oxide: $4Al(s) + 3O_2(g) \rightarrow 2Al_2O_3(S)$ | | Zinc | Reacts very slowly to form a thin coating of zinc oxide | Burns when heated strongly, especially if powdered | Forms a white powdery solid, zinc oxide, which is yellow when hot: $2Zn(s) + O_2(g) \rightarrow 2ZnO(s)$ | | Iron | Does not react with dry air | Burns when heated strongly, especially if powdered | Forms a black powdery solid, iron(, ) oxide: $3Fe(s) + 2O_2(g) \rightarrow Fe_3O_4(s)$ | | Copper | Does not react with dry air | Does not burn when heated, but does from an oxide coating if heated very strongly | Forms a black solid, copper(ii) oxide: $2Cu(s) + O_2(g) \rightarrow 2CuO(s)$ | | Silver | Does not react with dry air | Does not react, even when heated very strongly | | - If the metal oxides produced in these reactions can dissolve in water, the resulting solution is alkaline. - For example, potassium oxide reacts with water to form soluble potassium hydroxide, sodium oxide reacts with water to form soluble sodium hydroxide and calcium oxide reacts with water to form calcium hydroxide which is slightly soluble. - Potassium hydroxide, sodium hydroxide and calcium hydroxide are alkalis. - All the other oxides in **Table 17.2.1** are insoluble. - Looking at the reactions of metals with oxygen in the air we see that potassium, sodium, calcium and magnesium are the most reactive metals. - These metals are found in Groups I and II of the periodic table, i.e. to the far left of the table. - In general, the reactivity of metals decreases from left to right across the periodic table. #### The reactions of metals with water - Metals can also be added to water to determine their reactivity. - In general, metals that are very reactive with oxygen are highly reactive when exposed to water. - When a metal reacts with water, the metal hydroxide and hydrogen are produced. - The reaction can be summarised by the following general word equation: > metal + water $\rightarrow$ metal hydroxide + hydrogen - When a metal reacts with steam, the metal oxide and hydrogen are produced. - The reaction can be summarised by the following general word equation: > metal + steam $\rightarrow$ metal oxide + hydrogen **Table 17.2.2** shows what happens when certain metals are allowed to come into contact with water or steam. | Metal | Description of the reaction | Products formed and equation | |---|---|---| | Potassium | Reacts very vigorously with cold water producing a lilac flame | Potassium hydroxide and hydrogen: $2K(s) + 2H_2O(l) \rightarrow 2KOH(aq) + H_2(g)$ | | Sodium | Reacts vigorously with cold water, producing a lilac flame | Sodium hydroxide and hydrogen: $2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)$ | | Calcium | Reacts moderately with cold water | Slightly soluble calcium hydroxide and hydrogen: $Ca(s) + 2H_2O(l) \rightarrow Ca(OH)_2(aq) + H_2(g)$ | | Magnesium | Reacts very slowly with cold water and slowly with hot water | With water: very slightly soluble magnesium hydroxide and hydrogen: $Mg(s) + 2H_2O(l) \rightarrow Mg(OH)_2(aq) + H_2(g)$ With steam: magnesium oxide and hydrogen: $Mg(s) + H_2O(g) \rightarrow MgO(s) + 2H_2(g)$ | | Aluminium | Reacts vigorously with steam | Aluminium oxide and hydrogen: $2Al(s) + 3H_2O(g) \rightarrow Al_2O_3(s) + 3H_2(g)$ | | Zinc | Does not react with cold or hot water. Reacts with steam | Zinc oxide and hydrogen: $Zn(s) + H_2O(g) \rightarrow ZnO(s) + H_2(g)$ | | Iron | Does not react with cold or hot water. Reacts with steam | Iron(,) oxide and hydrogen: $3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)$ | | Copper | Does not react with water or steam | | | Silver | Does not react with water or steam | | - It can be seen from **Table 17.2.2** that some of the metals do not react with water. - They do, however, react with steam. - The fact that they require steam to react means that they are less reactive than the ones that react with cold water. - Once again, potassium, sodium and calcium are metals that react most vigorously with water. - When potassium and sodium react with cold water, they both roll into a ball and move rapidly around on the surface of the water as they decrease in size. - As they move, a lilac flame forms around the ball of potassium and an orange flame forms around the ball of sodium. #### The reactions of metals with dilute acids - In general, reactive metals react with acids, except nitric acid, to form a salt and hydrogen. - We can again see that the metals which are very reactive with oxygen and water are also very reactive with acids. - The reaction can be summarised by the following general word equation: > reactive metal + acid $\rightarrow$ salt + hydrogen - The salt that is produced depends on the metal used and the acid. **Table 17.2.3** shows what happens when certain metals react with dilute hydrochloric acid. - When a metal reacts with hydrochloric acid the salt formed is a chloride. | Metal | Description of the reaction | Products formed and equation | |---|---|---| | Potassium | An extremely violent reaction occurs | Potassium chloride and hydrogen: $2K(s) + 2HCl(aq) \rightarrow 2KCl(aq) + H_2(g)$ | | Sodium | A violent reaction occurs | Sodium chloride and hydrogen: $2Na(s) + 2HCl(aq) \rightarrow 2NaCl(aq) + H_2(g)$ | | Calcium | A fairly violent reaction occurs | Calcium chloride and hydrogen: $Ca(s) + 2HCl(aq) \rightarrow CaCl_2(aq) + H_2(g)$ | | Magnesium | A very vigorous reaction occurs | Magnesium chloride and hydrogen: $Mg(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + H_2(g)$ | | Aluminium | A vigorous reaction occurs | Aluminium chloride and hydrogen: $2Al(s) + 6HCl(aq) \rightarrow 2AlCl_3(aq) + 3H_2(g)$ | | Zinc | A fairly vigorous reaction occurs | Zinc chloride and hydrogen: $Zn(s) + 2HCl(aq) \rightarrow ZnCl_2(aq) + H_2(g)$ | | Iron | A very slow reaction occurs | Iron(ii) chloride and hydrogen: $Fe(s) + 2HCl(aq) \rightarrow FeCl_2(aq) + H_2(g)$ | | Copper | Does not react with dilute acids | | | Silver | Does not react with dilute acids | | - Looking at the reactions of metals with hydrochloric acid we can again see that potassium, sodium, calcium and magnesium are the most reactive metals. - The reactions between the metals and dilute sulfuric acid are very similar to the reactions which occur with dilute hydrochloric acid. - However, when a metal reacts with sulfuric acid, the salt formed is a sulfate. ### C17.3 Reactions of metal compounds - You have already come across some of the reactions of metal compounds. - We will review these and look at some new reactions in this unit. - The compounds we will be looking at include metal oxides, metal hydroxides, metal carbonates and metal nitrates. #### Reactions of metal oxides - Metal oxides react with acids to form a salt and water. - The reaction can be summarised by the following general word equation: > metal oxide + acid $\rightarrow$ salt + water - This reaction is known as a neutralisation reaction since the metal oxide is a base that neutralises the acid to form a salt and water. - When a metal oxide reacts with an acid the reaction becomes warmer because it is exothermic. #### Reactions of metal hydroxides - Metal hydroxides also react with acids to form a salt and water. - The reaction can be summarised by the following general word equation: > metal hydroxide + acid $\rightarrow$ salt + water - This reaction is also a neutralisation reaction since the metal hydroxide is a base and it neutralises the acid to form a salt and water. - The reaction is exothermic, therefore it becomes warmer. #### Reactions of metal carbonates - Metal carbonates react with acids to form a salt, carbon dioxide and water. - The reaction can be summarised by the following general word equation: > metal carbonate + acid $\rightarrow$ salt + carbon dioxide + water - When a metal carbonate reacts with an acid, effervescence is seen as the carbon dioxide is evolved. #### Decomposition of metal compounds - Many metal compounds decompose when they are heated. - **Table 17.3.1** summarises the effects of heat on metal nitrates, metal carbonates and metal hydroxides. | Metal | Metal compound | Nitrates | Carbonates | Hydroxides | |---|---|---|---|---| | Potassium | Decompose to form the metal nitrite and oxygen: $2NaNO_3(s) \rightarrow 2NaNO_2(s) + O_2(g)$ | Do not decompose. The carbonates are stable. | Do not decompose. The hydroxides are stable. | | Sodium | Decompose to form the metal oxide, nitrogen dioxide and oxygen: $2Mg(NO_3)_2(S) \rightarrow 2MgO(s) + 4NO_2(g) + O_2(g)$ | Decompose to form the metal oxide and carbon dioxide: $MgCO_3(s) \rightarrow MgO(s) + CO_2(g)$ | Decompose to form the metal oxide and water (as steam): $Mg(OH)_2(s) \rightarrow MgO(s) + H_2O(g)$ | | Calcium | | | | | | Magnesium | | | | | | Aluminium | | | | | | Zinc | Decompose to form the metal oxide, nitrogen dioxide and oxygen: $2Cu(NO_3)_2(S) \rightarrow 2CuO(s) + 4NO_2(g) + O_2(g)$ | Decompose to form the metal oxide and carbon dioxide: $CuCO_3(s) \rightarrow CuO(s) + CO_2(g)$ | Decompose to form the metal oxide and water (as steam): $Cu(OH)_2(S) \rightarrow CuO(s) + H_2O(g)$ | | Iron | | | | | | Lead | | | | | | Copper | | | Silver carbonate does not exist since it is too unstable. | Silver hydroxide does not exist since it is too unstable. | | Silver | Decomposes to form the metal, silver, nitrogen dioxide and oxygen: $2AgNO_3(s) \rightarrow 2Ag(s) + 2NO_2(g) + O_2(g)$ | | | | - When some metal compounds are heated and they decompose, colour changes are observed. - For example, when the compounds of copper are heated they change colour from blue or blue-green to black because black copper(II) oxide is formed in each reaction. - When compounds of lead are heated, they change colour from white to yellow because yellow lead(II) oxide is produced. - Tests given in **Table 22.3.1** on page 354 can be used to identify the oxygen, nitrogen dioxide, carbon dioxide and water vapour produced when metal compounds decompose. - Looking at the decomposition of metal compounds we can see that potassium and sodium form stable compounds which do not decompose on heating or, in the case of the nitrates, only decompose slightly. - However, silver forms compounds which are too unstable to exist or, in the case of silver nitrate, decomposes very easily and completely when heated. ### Key concepts - Metals are elements whose atoms have a small number of valence electrons, usually 1, 2 or 3, and they form positive cations in chemical reactions by losing these valence electrons. - A metal lattice is composed of metal cations surrounded by a 'sea' of delocalised electrons. - The electrostatic attraction between positive cations and delocalised negative electrons, known as the metallic bond, holds the metal lattice together. - In general, metals have high melting points and boiling points, conduct electricity and heat, have a shiny lustre, are malleable and ductile and have high densities. - The properties of metals can be explained by relating them to the structure of the metallic lattice. - Metals form positive cations in chemical reactions by losing valence electrons. As a result they form ionic compounds. - Metals behave as reducing agents in reactions. - Metals react with oxygen to form metal oxides. - If metal oxides can dissolve in water they form alkaline solutions. - Potassium, sodium, calcium and magnesium react with water to form the metal hydroxide and hydrogen. - Other metals, except copper, react with steam to form the metal oxide and hydrogen. - Reactive metals react with hydrochloric acid and sulfuric acid to form a salt and hydrogen. - Metal oxides react with acids to form a salt and water. The reaction is a neutralisation reaction. - Metal hydroxides react with acids to form a salt and water. The reaction is also a neutralisation reaction. - Metal carbonates react with acids to form a salt, carbon dioxide and water. - The nitrates of potassium and sodium decompose when heated to form the metal nitrite and oxygen. - The nitrates of other metals, except silver, decompose when heated to form the metal oxide, nitrogen dioxide and oxygen. - Silver nitrate decomposes when heated to form silver, nitrogen dioxide and oxygen. - The carbonates and hydroxides of potassium and sodium are not decomposed when heated. - The carbonates of other metals, except silver, decompose when heated to form the metal oxide and carbon dioxide. Silver carbonate does not exist. - The hydroxides of metals, except silver, decompose when heated to form the metal oxide and water as steam. Silver hydroxide does not exist.