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
  • analyse patterns in metal activity on the periodic table and explain why they correlate with, for example:
– ionisation energy (ACSCH045)
– atomic radius (ACSCH007)
– electronegativity (ACSCH057)

        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 

        Recap from Module 1: Properties and Structure of Matter

        • The reactivity of metals increases across a period from the right to the left, and increases down a group. This corresponds to a similar trend in the metal's 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.

        Review reactivity with water here

        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. 

        Metal Activity, Atomic Radius and Ionisation Energy

        Atomic radius refers to the size of an atom. 

        Trend of atomic radius on the Periodic Table:

        • Across a Period: Atomic radius decreases as you move from left to right. This is due to increased nuclear charge, which pulls the electrons closer to the nucleus, reducing the size of the atom.
        • Down a Group: Atomic radius increases down a group due to the addition of more electron shells, which places the outermost electrons further from the nucleus.

          

        Metal activity is related to a metal's atomic radius because atomic radius dictates ionisation energy. Ionisation energy is defined as the energy required to remove the outermost electron(s) from an atom to form a cation. It is a critical factor in determining a metal's activity, which refers to its tendency to undergo chemical reactions. 

        Trend of ionisation energy on the Periodic Table:

        • Across a Period: Ionisation energy generally increases from left to right across a period. This increase is due to a greater nuclear charge (more protons) causing a stronger attraction between the nucleus and the electrons, making it harder to remove an electron.
        • Down a Group: Ionisation energy decreases down a group. This decrease can be attributed to the greater distance between the outer electrons and the nucleus, along with increased electron shielding, which reduces the effect of the nuclear charge on the outer electrons.

         

        Generally, metals with larger atomic radii have lower ionisation energies, making them more reactive. The increased distance of the outer electrons from the nucleus makes them easier to remove, enhancing the metal’s activity.

        Metal reactions involve transformation of metal atoms into metal cations. Metals with low ionisation energies tend to be more reactive because they can lose electrons more readily. For example, alkali metals, which have some of the lowest ionisation energies on the periodic table, are also among the most reactive.

        Metal Activity and Electronegativity

        Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. While often discussed in the context of nonmetals, its trends can also help explain the reactivity of metals.

        Trend of electronegativity on the Periodic Table:

        • Across a Period: Electronegativity increases across a period. This is due to smaller atomic radii and a greater effective nuclear charge.
        • Down a Group: Electronegativity decreases down a group as the atomic radius increases and the effective nuclear charge felt by bonding electrons decreases.

         

        Metals with lower electronegativity are typically more active. This lower electronegativity indicates a weaker attraction for electrons, thus facilitating the loss of electrons and enhancing metal activity.

        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

         

        Metals

        K, Na, Ca

        Mg

        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:

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

         

        In metal and water reactions, metal atoms become metal ions in the form of metal hydroxides. Thus, the reactivity of metal with water is greater if the ionisation energy of the metal is lower.

        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. 

          

          
         Table 2: Reactivity of various metals with dilute acid.

        Metals

        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

         

        Example: Reaction between sodium and hydrochloric acid

        $$2Na(s) + 2HCl(aq) \rightarrow 2NaCl(aq) + H_2(g)$$

         

        Example: Reaction between iron and hydrochloric acid

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

         

        In metal and acid reactions, metal atoms become metal ions in the form of salts. Thus, the reactivity of metal with acid is greater if the ionisation energy of the metal is lower.

        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. 

         

         
        Table 3: reactivity of varying metals with oxygen.

        Metals

        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

         

        Example: Reaction between potassium and oxygen:

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

         

        Example: Reaction between magnesium and oxygen:

        $$2Mg(s) + O_2(g) \rightarrow  2MgO(s)$$

         

        In metal and oxygen reactions, metal atoms become metal ions in the form of metal oxides. Thus, the reactivity of metal with oxygen is greater if the ionisation energy of the metal is lower.

        Metal Displacement Reaction

         

        An interesting observation which arises when a metal is submerged into a solution of a less reactive metal ion, is the less reactive metal ion becomes displaced and deposits 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 sulfate solution to lead to metal displacement, causing copper metal to deposit on the surface of the more reactive metal. However silver, being less reactive, does not participate in such reactions.

        Metal displacement reaction is discussed in more detail in redox reactions.

         

         

        RETURN TO MODULE 3: REACTIVE CHEMISTRY