Learning Objective

In this lesson we will learn about the chemical reactivity of metals and some of the different reactions involving metals.

Learning Outcomes

By the end of this lesson you will be able to:

  • Describe how metals form ions when reacting.
  • Write the products of:
  • Metal-oxygen reactions.
  • Metal-water reactions.
  • Metal-acid reactions.
  • Metal displacement reactions.
  • Metal-halogen reactions.
  • Use the reactivity series to predict whether a particular chemical reaction involving a metal will take place.

Metal Reactions

1 | Chemical Reactions of Metals

2 | Metal-Oxygen Reactions

3 | Metal-Water Reactions

4 | Metal-Acid Reactions

5 | Metal-Salt Displacement Reactions

6 | Metal-Halogen Reactions

7 | Summary

 

 metal reactions worksheet  year 10 chemistry pdf workbook  Year 10 Chemistry Print Workbook Australian Curriculum

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Chemical Reactions of Metals

  • The chemical reactivity of metals varies widely.
  • Whether a metal will react depends not only on the other substance, but also the type of metal.
  • For example:
  • All metals react with halogens.
  • Almost all metals react with oxygen.
  • Most metals react with dilute acids.
  • Only some metals react with water.
  • Since metals react by losing electrons, the easier it is for a metal to lose electrons, the more reactive it is.
  • The reactivity of main group metals (those that aren’t transition metals) can be predicted from their location on the periodic table.
  • Reactivity decreases across periods: group 1 metals are the most reactive, followed by group 2, then group 13.
  • Reactivity increases down groups: for example, potassium is more reactive than sodium, and sodium is more reactive than lithium.
  • The relative reactivities of all common metals are displayed in the reactivity series of metals.
  • This lists metals from most reactive to least reactive.
  • The reactivity series can be used to predict whether a particular chemical reaction will take place.

 
reactivity series of metals

The reactivity of metals varies widely.

 


Metal-Oxygen Reactions

  • Chemical reactions involving oxygen are known as oxidation reactions.
  • Corrosion and combustion are two examples of oxidation reactions.
  • Corrosion reactions are generally very slow, whereas combustion reactions are very fast and produce large amounts of heat and light.
  • Metals can undergo both corrosion and combustion reactions, depending on the type of metal and the reaction conditions.
  • * Corrosion reactions are specific to metals, but don’t always involve oxygen; they may also involve other oxidants, such as sulfur.
  • * Combustion reactions are not specific to metals – they also involve a wide variety of fuels, usually carbon compounds.

 

Corrosion of Metals

  • Almost all metals react with oxygen to form metal oxides.
  • These reactions have the general formula:

metal oxygen reaction formula

  • Example
  • iron  +  oxygen    iron (III) oxide
  • 4 Fe (s)  +  3 O2 (g)    2 Fe2O3 (s)
  • The corrosion of iron is known as rusting.
  • Less reactive metals, such as copper and mercury, require heat for the reaction to proceed.
  • The least reactive metals, such as gold and platinum, do not react with oxygen under most conditions.

 
iron metal rust corrosion

Iron reacts slowly with oxygen in the atmosphere to form iron oxide (rust).

(Image: terimakasih0, Pixabay)

 

Combustion of Metals

  • More reactive metals, such as magnesium and aluminium, can combust (burn) under certain conditions.
  • Example
  • Magnesium ribbon readily burns when placed over a Bunsen flame.
  • 2 Mg (s)  +  O2 (g)    2 MgO (s)
  • Whether a metal reacts with oxygen by corrosion or combustion, the products are the same – metal oxides.

 
magnesium ribbon burning oxygen

When ignited over a flame, magnesium burns very bright and hot.

(Image: Capt. John Yossarian, Wikimedia Commons)

 


Metal-Water Reactions

  • Metals in groups 1 and 2 react with water to form a metal hydroxide and hydrogen.
  • These reactions have the general formula:

metal water reaction formula

  • Example
  • sodium  +  water    sodium hydroxide  +  hydrogen
  • 2 Na (s)  +  2 H2O (l)    2 NaOH (s)  +  H2 (g)
  • The reactions involving group 1 metals and water range from vigorous to explosive.
  • Magnesium requires hot water or steam and beryllium requires very hot steam.

 
sodium water reaction

Group 1 metals, such as sodium, react violently with water.

(Image: Tavoromann, Wikimedia Commons)

 


Metal-Acid Reactions

  • Most metals react with dilute acids to form a salt and hydrogen.
  • These reactions have the general formula:

metal acid reaction formula

  • Example
  • zinc  +  hydrochloric acid    zinc chloride  +  hydrogen
  • Zn (s)  +  2 HCl (aq)    ZnCl2 (aq)  +  H2 (g)
  • More reactive metals, such as group 1 metals, are too dangerous to mix with acids, due to the explosive reaction.
  • Metals with an intermediate reactivity, such as magnesium and iron, can be safely reacted with dilute acids.
  • Less reactive metals, such as copper and gold, do not react with dilute acids.
  • The reactivity series of metals can be used to predict whether a metal will react with dilute acids.
  • If hydrogen is added to the reactivity series, it lies between lead and copper.
  • Metals below hydrogen on the reactivity series do not react with dilute acids.
  • This is because metal-acid reactions are a type of single displacement reaction – only metals that are more reactive than hydrogen will displace it in solution.

 
zinc hydrochloric acid reaction

Zinc metal reacts readily with hydrochloric acid.

(Image: Chemicalinterest, Wikimedia Commons)

 


Metal Displacement Reactions

  • Metals can react with salts in solution to form a metal and a salt.
  • These reactions are called metal displacement reactions – a type of single displacement reaction.
  • These reactions have the general formula:

metal salt displacement reaction formula

  • Since metal displacement reactions are a type of single displacement reaction, they will only proceed if the metal that exists as an element is more reactive than the metal that exists as a salt.
  • Example 1
  • Mixing calcium and iron (II) nitrate will result in a chemical reaction as calcium is more reactive than iron:
  • calcium  +  iron (II) nitrate    calcium nitrate  +  iron
  • Ca (s)  +  Fe(NO3)2 (aq)    Ca(NO3)2 (aq)  +  Fe (s)
  • Example 2
  • Mixing copper and iron (II) nitrate will not result in a chemical reaction as copper is less reactive than iron:
  • copper  +  iron (II) nitrate    no reaction
  • Cu (s)  +  Fe(NO3)2 (aq)    no reaction
  • The reactivity series of metals can be used to predict whether a displacement reaction will take place.
  • If the elemental metal is more reactive than the ionic metal, it will displace the ionic metal from the solution – that is, a reaction will occur.
  • If the elemental metal is less reactive than the ionic metal, it will not displace the ionic metal from the solution – that is, a reaction will not occur.

 

Metal Displacement Reactions as Redox Reactions

  • Metal displacement reactions are redox reactions as they involve the transfer of electrons from the more reactive metal to the less reactive metal.
  • The more reactive metal is oxidised (forms ions from atoms), while the less reactive metal is reduced (forms atoms from ions):
  • metal-1  +  metal-2 ion    metal-1 ion  +  metal-2
  • The negative ion (anion) remains unchanged during a metal displacement reaction. It does not participate in the reaction and has no effect on whether the reaction will take place or not.
  • It is therefore referred to as a ‘spectator ion‘ and can be omitted from the formula equation.
  • For example, in the above reaction between calcium and iron (II) nitrate, the nitrate ion is a spectator ion. The formula equation can therefore be simplified as:
  • Ca (s)  +  Fe2+ (aq)    Ca2+ (aq)  +  Fe (s)
  • This is known as a net ionic equation.

 
displacement reaction silver copper metal solution

Copper is more reactive than silver so it will displace silver ions from solution – here, silver metal is forming a precipitate on the wire, while copper is forming ions in solution.

(Image: Toby Hudson, Wikimedia Commons)

 


Metal-Halogen Reactions

  • All metals react with halogens (group 17) to form metal halides.
  • These reactions have the general formula:

metal halogen reaction formula

  • Example
  • lithium  +  chlorine    lithium chloride
  • 2 Li (s)  +  Cl2 (g)    2 LiCl2 (s)

 
lithium chlorine metal halogen reaction

Halogens, such as chlorine, react vigorously with alkali metals, such as lithium.

(Image: Kei Yuen College)


 


Summary

  • Metals react by losing electrons to form positive ions as part of ionic compounds.
  • The reactivity series of metals lists metals from most reactive to least reactive.
  • This can be used to predict whether a chemical reaction will take place or not.
  • Almost all metals react with oxygen to produce metal oxides.
  • metal  +  oxygen    metal oxide
  • Most metals react with dilute acids to produce a salt and hydrogen.
  • metal  +  acid    salt  +  hydrogen
  • Some metals react with water to produce metal hydroxides and hydrogen.
  • metal  +  water    metal hydroxide  +  hydrogen
  • Metals can react with ionic salts if the metal is more reactive than the metal in solution.
  • metal-1  +  metal-2 salt    metal-2  +  metal-1 salt
  • These are known as metal displacement reactions.
  • All metals react with halogens to produce metal halides.
  • metal  +  halogen    metal halide
  • All metal reactions are redox reactions, as they involve the transfer of electrons.

 
reactivity of metals in dilute sulfuric acid

Reactivity of metals in dilute sulfuric acid.

(Image: Capaccio, Wikimedia Commons)

(Header image: 3dman_eu, Pixabay)

 

 metal reactions worksheet  year 10 chemistry pdf workbook  Year 10 Chemistry Print Workbook Australian Curriculum

Click images to preview the worksheet for this lesson and the Year 10 Chemistry Workbook (PDF and print versions)