Chemical Reactions Involving Metals

This is part of Year 11 HSC Chemistry course under the topic of Prediction Reactions of Metals

HSC Chemistry Syllabus

  • Conduct practical investigations to compare the reactivity of a variety of metals in:
    • Water
    • Dilute Acid
    • Oxygen
    • Other metal ions in solution


      Chemical Reactions involving metals

      This video goes various chemical reactions involving reactions between water of different temperature, acids of different concentration, oxygen, and other metal ions in solution with metals. 

      Predicting Metal Reactivity 


      • The reactivity of metals increases from the right to the left side of the NESA periodic table.
      • Similarly, an upward trend from the bottom left to the top right of the periodic table represents increasing ionisation energy.
      • Considering these two trends, we should expect that metals with a low first ionisation energy such as sodium or potassium are more reactive.
      • Non-metals reside at the top right of the periodic table and are generally less reactive than metals. 

      Transition Metals

      Transition metals, such as elements like nickel, are generally less reactive than group 1 and 2 metals. This reduced reactivity of the metals is attributed to their strong nuclear attractions towards the nucleus which result because of the large number of electrons in their inner shells. 

      Metals Reacting with Water

      When metals come into contact with water, they typically form a metal hydroxide or oxide while releasing hydrogen gas. The reactivity of the metal and the water's temperature, are important points to take into account when determining the nature of the reaction. For instance, metals like potassium, sodium, and calcium, all vigorously react with cold water to release hydrogen gas. This gas can ignite due to the immense energy that these reactions release. 

      On the other hand, less reactive metals like magnesium only react with hot water, while transition metals like aluminium, zinc, and iron, require steam and red-hot conditions to react. Notably, metals like lead, tin, copper, mercury, gold, and silver, remain unreactive with water, attributed to their large ionisation potentials



      K, Na, Ca


      Mg, Al, Zn, Fe

      Pb, Sn, Cu, Hg, Ag, Au

      Reactivity with water

      Reacts with cold water to form hydroxide ions and release hydrogen gas

      Reacts with hot water to form hydroxide ions and release hydrogen gas

      Reacts with steam until red hot to form oxide ions and release hydrogen gas

      Does not react with water

      Table 1: reactivity of various metals with water at varying temperatures 


      Example: Reaction between sodium and water:

      $$2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g)$$


      Example: Reaction between zinc and water:

      $$Zn(s) + H_2O(l) \rightarrow ZnO(aq) + H_2(g)$$


      Metals Reacting with Dilute Acids

      Metals tend to react with dilute acids, leading to the formation of salts and the liberation of hydrogen gas. The intensity of the reaction can be a gauge of the metal's reactivity. For instance, highly reactive metals like potassium, sodium, and calcium, when reacting with acids, can unleash energy potent enough to ignite the hydrogen gas produced. 



      K, Na

      Ca, mg

      Al, Zn, Fe, Sn

      Pb, Cu, Hg, Ag, Au

      Reactivity with dilute acid

      Effervesces very rapidly to produce hydrogen gas which may ignite

      Bubbles rapidly to produce hydrogen gas

      Bubbles slowly to moderately as hydrogen gas is released

      No reaction

       Table 2: Reactivity of various metals with dilute acid


      Example: Reaction between iron and hydrochloric acid

      $$Fe(s) + 2HCl(aq) \rightarrow FeCl_2(aq) + H_2(g)$$


      Metals Reacting with Oxygen

      Metals will also engage in reaction with oxygen to yield metal oxides, a process that is accelerated if the metals are combusted. Reactive metals such as potassium, sodium, and calcium generate oxide layers rapidly. These layers are highly reactive and can spontaneously combust when exposed to air. Conversely, metals like magnesium, aluminium, and zinc may yield less reactive oxide layers which protect the metals from further reaction. Inert metals like lead, tin, and copper, develop protective oxide coatings during heating, which halt further reactions. Notably, gold and silver remain unresponsive to oxygen. 



      K, Na, Ca

      Mg, Al, Zn, Fe

      Pb, Sn, Cu, Hg

      Ag, Au

      Reactivity with Oxygen

      Burn rapidly to form oxides or peroxides

      Burn readily if powdered or as fine fibres to form oxides

      Become coated with oxide layers during heating

      No reaction

      Table 3: reactivity of varying metals with oxygen.


      Example: Reaction between potassium and oxygen:

      $$4K(s) + O_2(g) \rightarrow 2K_2O(s)$$


      Example: Reaction between silver and oxygen

      $$2Ag(s) + O_2(g) \rightarrow  2Ag_2O(s) \text{  (no reaction)} $$


      Metal Displacement

      An interesting observation which arises when a metal is submerged into a solution of a less reactive metal, is the less reactive metal begins to become displaced and starts to deposit out of the solution. This experimental approach to testing metal reactivity is pivotal to creating the metal reactivity series. 

      For example, metals like iron, lead, zinc, and magnesium, which exhibit more reactivity than copper, will react with copper sulphate solution to lead to metal displacement. However silver, being less reactive, does not participate in such reactions. Concepts regarding metal displacement are explored further in the redox reactions video and section of module 3 chemistry.