Chapter 12 Reactivity of Metals 12.1 Different Reactivities of Metals Recall an experiment performed in F.3

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Wilfred Hodges
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1 Chapter 12 Reactivity of Metals 12.1 Different Reactivities of Metals Recall an experiment performed in F.3 p.1/9 When freshly cut, potassium has a shiny surface and it reacts vigorously with water, giving out a lilac flame. Gold, on the other hand, does not react with water at all. Reactivity is the readiness to react. Different metals have different reactivities towards substances Putting Five metals in order of reactivity The metals calcium, magnesium, iron, lead and copper are allowed to react with air, water and dilute hydrochloric acid. For a fair comparison of reactivity, the metals (1) should be allowed to react with the same substance in turn; (2) should be in roughly the same state of subdivision; (3) should react under the same conditions of temperature and pressure 12.3 Comparing reactivity of common metals 7.4 Reactivity of metals Most metals, when heated in air, combine with oxygen to form oxides. On exposure to air, most metals lose their shiny appearance and become dull. They are said to tarnish in air. The general reactivity of some common metals when exposed to air are summarized in the following table Metal K Na Ca Mg Al* (protective oxide film prevent further tarnishing) Zn* (protective oxide film prevent further tarnishing) Fe Pb Cu Ag** (react with hydrogen sulphide to form silver sulphide) Au Pt Reaction on exposure to air Tarnish very quickly Calcium is easiest to tarnish in this group. Ag has the least tendency to tarnish in this group. Do not tarnish

2 p.2/9 However, when metals are heated in air, they tend to react more vigorously. The following table summarises the reactivities of different metals. Metal Conditions for reaction/observation Name of oxide Nature of oxide Potassium Burns with a lilac flame to form a white powder potassium superoxide, KO 2 (white) soluble, alkaline Sodium Burns with a yellow flame sodium peroxide, soluble, alkaline Na 2 O 2 (white) Calcium Strong heating is required to make it burn; burns with a brick-red flame calcium oxide, CaO (white) slightly soluble, alkaline Magnesium Strong heating is required; burns with magnesium oxide, almost a bright white flame MgO (white) Aluminium Strong heating is required to burn aluminium powder aluminium oxide, Al 2 O 3 (white) Zinc Strong heating is required to burn zinc Zinc oxide, ZnO powder (yellow when hot and Iron Lead Copper Mercury Silver Gold Platinum Strong heating is required; iron powder burns with yellowish showery sparks It melts upon strong heating to produce a silvery ball; a powder (orange when hot and yellow when cold) is seen on the surface white when cold) iron(ii) diron(iii) oxide, Fe 3 O 4 (black) lead(ii) oxide, PbO (orange when hot and yellow when cold) Its surface turns black on strong copper(ii) oxide, CuO heating (black) A red powder is formed on the surface mercury(ii) oxide, HgO upon very strong heating (red) No change even after strong heating No change even after strong heating No change even after strong heating Appearance of metals and storage methods When a freshly cut metal is left in air for some time the surface will look dull due to the formation of oxide layer. This reaction is more vigorous with reactive metals like sodium and potassium. To protect oxide formation in air, reactive metals are often stored under oil. Moderately reactive metals like calcium can be stored in airtight container. Gold and other inert metals require no special method of storage. Reaction of Metals with water Action of potassium on water A small piece of freshly cut potassium, when put in water, melts into a silvery ball and moves quickly on the surface of water, bursting into lilac flame. The resulting solution is alkaline. Potassium + water potassium hydroxide solution + hydrogen (alkaline solution) Action of sodium on water Sodium reacts with water in a similar manner as potassium. It melts into a silvery ball and moves quickly on the surface of the water, bursting into a bright yellow flame. An alkaline solution is formed.

3 Sodium + water Sodium hydroxide solution + hydrogen (alkaline solution) p.3/9 Action of calcium on water Care must be taken to choose pieces of calcium granules that are still shiny (and not dull) in appearance. Why? In the space below, sketch a diagram showing how the hydrogen could be trapped by reacting calcium with water. Write a word equation describing the reaction between calcium and water. Action of magnesium, aluminium, zinc and iron on steam Sketch a diagram below, showing how aluminium reacts with steam and how hydrogen could be collected, using simple apparatus in the chemistry laboratory. Write word equations for (a) magnesium reacting with steam (b) zinc reacting with steam (c) iron reacting with steam

4 p.4/9 Aluminium does not react with steam, no matter how high the temperature is. Explain why this is so. Reactions of Metals with Dilute Hydrochloric Acid Dilute hydrochloric acid is chosen as it is a typical acid without any oxidizing or dehydrating properties (unlike nitric acid and hydrochloric acid in the concentrated form) Metals above hydrogen in the reactivity series react with dilute hydrochloric acid. Hydrogen gas is given off and the metal chlorides are formed. Group I metals react explosively with dilute acids and should not be attempted in the school laboratory. Some metals resist reaction with dilute acids because of a protective oxide coating. For example, the aluminium oxide coating on the surface of aluminium prevents reaction. The coating can be removed by cleaning with concentrated hydrochloric acid. Write word equations between the following metals (oxide coating removed) with dilute hydrochloric acid. Magnesium, aluminium, lead and copper 12.4 The Metal Reactivity Series By comparing their reactions with air, water and dilute hydrochloric acid, we can arrange common metals in order of reactivity. The order is called the metal reactivity series (MRS). Potassium, K Sodium, Na Calcium, Ca Magnesium, Mg Aluminium, Al Zinc, Zn Iron, Fe Lead, Pb Copper, Cu Mercury, Hg Silver, Ag Gold, Au Most reactive Decreasing reactivity Least reactive

5 Q.12.2 Experiment Strong heating in air Metal A B C D forms a black burns with a lilac powder flame burns with a brickred flame p.5/9 burns with a dazzling white flame Reaction with cold moderate reaction no reaction violent reaction; no reaction water burns by itself Reaction with dilute hydrochloric acid fast reaction No reaction (experiment not performed) fast reaction (a) Arrange the four metals in decreasing order of reactivity. Give reasons for your choice. (b) Give possible names of the four metals Chemical Equations Word equation It is used to show the names of reactants and products: e.g. Chemical equation It is used to show the rearrangement of atoms that occur in the reaction. Symbols and formulae are needed in writing a chemical equation. e.g. State symbols The physical states of reactants and products are shown using state symbols. (s) = (l) = (g) = (aq) = How to write and balance a chemical equation 1. The names of reactants and products are written. 2. The names of reactants and products are translated to their chemical symbols or formulae. 3. Balance the equation and check that it is right by counting the number of each type of atoms for the reactants and products of both the left and right hand sides. 4. Add state symbols to the reactants and products. Examples 1. Write a balanced chemical equation for the reaction between zinc and hydrochloric acid to give hydrogen and zinc chloride.

6 2. When copper(ii) oxide reacts with carbon monoxide, copper and carbon dioxide are formed. Write a balanced chemical equation for this chemical change. p.6/9 Rewrite the following word equations as balanced chemical equations and add appropriate state symbols. (a) methane(ch 4 ) + oxygen(o 2 ) carbon dioxide(co 2 ) + water(h 2 O) (b) calcium oxide + hydrochloric acid calcium chloride + water 12.6 A Simple Explanation for Different Reactivities of Metals When magnesium burns in oxygen, magnesium loses electrons to form magnesium ions. Mg Mg e - Oxygen accepts electrons to become oxide ions. O 2 + 4e - 2O 2- In general, metals react by losing electrons and their reactivity, therefore, depends on how easily a metal loses electrons. A portion of the periodic table is shown. I II III IV V VI VII 0 Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Ga Br Kr Which element loses electron most readily? Explain briefly. Which element gains electron most readily? Explain briefly. For the Groups I, II and III, which metal is the most reactive? Explain in terms of their readiness to lose electrons.

7 p.7/9 Conclusion A Metal higher in the reactivity series has a higher reactivity, and its atoms would lose outer shell electrons to form cations more easily Applications of the Reactivity Series Reduction of Metal Oxides Oxidation and reduction The addition of oxygen to a substance is called oxidation. The removal of oxygen from a substance is called reduction Reducing metal oxides There are 3 main ways by which metal oxides may be reduced to form the metals. (1) Heating the metal oxide alone. (2) Reducing the metal oxides by another metal. (3) Reducing the metal oxide by carbon. Action of heat on metal oxides Oxides of metals below copper in the reactivity series, e.g. silver oxide, mercury(ii) oxide decompose on heating to give oxygen and the metal. Write equations to represent the thermal decomposition of (1) mercury(ii) oxide (2) silver oxide Reduction with metals A metal oxide could be reduced by heating it with another metal higher in the reactivity series. When copper(ii) oxide is heated with magnesium, copper and magnesium oxide are formed. Write a chemical reaction to represent this reaction. Reduction with carbon Oxides of less reactive metals can be reduced by carbon (heated charcoal) to form the metal and carbon dioxide. The process is very important and is used widely in industry to extract the metal from its ore. For example, write an equation to represent the extraction of copper from its ore of copper(ii) oxide by heating with charcoal (charcoal is a form of carbon). Draw a diagram to show how a sample of copper(ii) oxide can be reduced to copper by heating with town gas in the school laboratory. Write an equation for the process and explain the colour changes of the reduction.

8 Are all metal oxides reduced by heated charcoal? Which metal oxides are able to be reduced by heated charcoal? p.8/9 Oxides of metals from zinc downwards in the reactivity series can be reduced by carbon (heated charcoal). Mercury oxides and silver oxides can be reduced by heating the oxides alone. The following table summarizes the reduction of metal oxides with carbon at 1500 C Oxides Heating with carbon K 2 O Na 2 O CaO no action MgO Al 2 O 3 ZnO Fe 2 O 3 reduced with increasing ease from ZnO to CuO PbO CuO HgO can be decomposed to metals by heating alone Ag 2 O Displacement reactions of metals in aqueous solution A metal higher in the metal reactivity series can displace a metal lower in the series from the solution of the compound. For example, zinc can displace copper from a solution of copper(ii) sulphate. However, copper cannot displace zinc from a solution of zinc sulphate. Zn(s) + CuSO 4 (aq) Cu(s) + ZnSO 4 (aq) Cu(s) + ZnSO 4 (aq) no reaction 12.8 Ionic Equations Ionic equation is a simple representation of a chemical reaction between ions. When silver nitrate and sodium chloride react to give silver chloride and sodium nitrate, write a balanced chemical equation. Convert the chemical equation into an ionic equation, making use of the concept of spectator ions. Rewrite the following equation into an ionic equation : CuSO 4 (aq) + 2NaOH(aq) Cu(OH) 2 (s) + Na 2 SO 4 (aq) Q.12.8 Predict, with reasons, whether a reaction takes place in each of the following. Write an ionic equation for any reaction which occurs. (a) Magnesium is added to lead(ii) nitrate solution.

9 (b) Silver is added to copper(ii) sulphate solution. p.9/9 (c) Magnesium is added to dilute hydrochloric acid. (Hint: Dilute hydrochloric acid contains H + (aq) and Cl - (aq) ions.) 12.9 Extraction of Metals from their Ores In nature, metals usually occurred in combined state in compounds called ores. The process of obtaining a metal from its ore is called extraction. The method of extraction depends on the reactivity of the metals concerned. (1) Reactive metals are obtained from electrolysis of their molten ores, e.g. K, Na, Ca, Mg and Al. (2) Moderate reactive metals are obtained from their ores by reduction with carbon. e.g. 2ZnO + C 2Zn + CO 2 (3) Unreactive metals are obtained from their ores simply by heating, e.g. Cu, Hg and Au. K Na Ca Mg Al Zn Fe Pb Cu Ag Extracted by electrolysis of the molten ores Extracted by heating ores with carbon Extracted by heating ores in air Refer to a summary on p.217 & 218 Book 1A.

1. Which of the given statements about the reaction below are incorrect?

1. Which of the given statements about the reaction below are incorrect? 2PbO(s) + C(s) 2Pb(s) + CO 2 (g) a. Lead is getting reduced b. Carbon dioxide is getting oxidised c. Carbon is getting oxidised