Unit 9: Chemical Properties of d-block Metals
Discover the unique chemistry of transition metals, from their varied reactivity to their vital role as catalysts.
9.11 d-block Metals (Fe, Cu, Au): Reactivity with Air, Water, Acids
The d-block metals, or transition metals, are generally less reactive than s-block and p-block metals. Their reactivity varies considerably across the block, but most are stable enough to be used as structural materials and in coinage.
Reaction with Air (Oxygen)
Most d-block metals are resistant to reaction with air at room temperature, often due to forming a protective oxide layer or simply being unreactive.
- Iron (Fe): Reacts with oxygen only when heated, forming iron(II,III)
oxide, a mixed oxide often written as $Fe_3O_4$. It is famously susceptible to slow
reaction with oxygen in the presence of water (rusting).
$3Fe(s) + 2O_2(g) \xrightarrow{\Delta} Fe_3O_4(s)$ - Copper (Cu): Is quite unreactive with air but will slowly form a black
layer of copper(II) oxide when heated.
$2Cu(s) + O_2(g) \xrightarrow{\Delta} 2CuO(s)$ - Gold (Au): Is extremely unreactive and does not react with oxygen, which is why it remains shiny indefinitely.
Reaction with Water
Most d-block metals do not react with cold water.
- Iron (Fe): Does not react with pure water but will react with steam at
high temperatures to produce iron(II,III) oxide and hydrogen gas.
$3Fe(s) + 4H_2O(g) \rightleftharpoons Fe_3O_4(s) + 4H_2(g)$ - Copper (Cu) and Gold (Au): Do not react with water or steam. This unreactivity is why copper is used for water pipes.
Reaction with Acids
Reactivity with acids depends on the metal's position in the reactivity series relative to hydrogen.
- Iron (Fe): Being more reactive than hydrogen, it reacts with dilute
non-oxidizing acids like HCl and $H_2SO_4$ to produce an iron(II) salt and hydrogen gas.
$Fe(s) + 2HCl(aq) \rightarrow FeCl_2(aq) + H_2(g)$ - Copper (Cu): Being less reactive than hydrogen, it does not react with dilute non-oxidizing acids. It will, however, react with oxidizing acids like nitric acid ($HNO_3$).
- Gold (Au): Is extremely unreactive and does not react with common acids. It will only dissolve in aqua regia, a highly corrosive mixture of nitric acid and hydrochloric acid.
Solved Examples:
-
Why are copper and gold used in coins while iron is not?
Solution: Copper and gold are very unreactive towards air and water. They do not corrode or tarnish easily, so they maintain their appearance over time. Iron, however, rusts easily, making it unsuitable for coinage. -
Write a balanced equation for the reaction of iron with dilute sulfuric
acid.
Solution: $Fe(s) + H_2SO_4(aq) \rightarrow FeSO_4(aq) + H_2(g)$ -
A student places a copper turning in a test tube of dilute hydrochloric
acid. What will they observe?
Solution: They will observe nothing. Copper is less reactive than hydrogen and cannot displace it from a non-oxidizing acid like HCl.
9.12 Transition Metals (Variable Oxidation States, Catalytic Abilities)
A transition metal is formally defined as a d-block element that forms at least one stable ion with a partially filled d-subshell. This definition excludes elements like Scandium (forms only $Sc^{3+}$ with an empty d-subshell) and Zinc (forms only $Zn^{2+}$ with a full d-subshell). These partially filled d-orbitals give rise to their most distinctive chemical properties.
Variable Oxidation States
The energies of the 4s and 3d subshells are very close. This means that transition metals can lose not only their 4s electrons but also a variable number of 3d electrons to form ions with different charges (oxidation states).
- Iron (Fe): Has the electron configuration [Ar] $3d^6 4s^2$. It can lose the two 4s electrons to form the Fe²⁺ (iron(II)) ion. It can also lose one more 3d electron to form the more stable Fe³⁺ (iron(III)) ion.
- Copper (Cu): Has the configuration [Ar] $3d^{10} 4s^1$. It can lose the 4s electron to form Cu⁺ (copper(I)) or lose an additional 3d electron to form the more common Cu²⁺ (copper(II)) ion.
Catalytic Abilities
The ability to exist in multiple oxidation states and the presence of available d-orbitals make transition metals and their compounds excellent catalysts. They can provide an alternative reaction pathway with lower activation energy by temporarily bonding with reactants.
- Iron (Fe) is used as the catalyst in the Haber-Bosch process for manufacturing ammonia: $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$.
- Nickel (Ni) is used to catalyze the hydrogenation of alkenes in margarine production.
Solved Examples:
-
Explain why zinc (Zn) is considered a d-block element but not a transition
metal.
Solution: Zinc's electron configuration is [Ar] $3d^{10} 4s^2$. It only forms one stable ion, $Zn^{2+}$, by losing its two 4s electrons. The configuration of the $Zn^{2+}$ ion is [Ar] $3d^{10}$. Since this ion has a completely full d-subshell, zinc does not meet the definition of a transition metal. -
What are the two common oxidation states of copper? Write the formula for
copper(I) oxide and copper(II) oxide.
Solution: The two common oxidation states are +1 and +2. Copper(I) oxide is $Cu_2O$. Copper(II) oxide is $CuO$. -
How does a catalyst like iron speed up a reaction?
Solution: The iron catalyst provides an active site on its surface where reactant molecules (N₂ and H₂) can adsorb and their bonds can be weakened. This creates a new reaction pathway with a lower activation energy, allowing the reaction to proceed much faster.
9.13 Rusting of Iron
Rusting is the common term for the corrosion of iron and its alloys, such as steel. It is a slow electrochemical process that requires the presence of both oxygen and water. The rust itself is hydrated iron(III) oxide, $Fe_2O_3 \cdot nH_2O$.
Unlike the protective oxide layer on aluminum, rust is flaky, porous, and brittle. It does not protect the underlying iron. Instead, it flakes off, exposing fresh iron to continue the rusting process, which can eventually destroy the entire object.
The overall simplified chemical equation for rusting is:
$$ 4Fe(s) + 3O_2(g) + 2nH_2O(l) \rightarrow 2Fe_2O_3 \cdot nH_2O(s) $$The process can be accelerated by the presence of dissolved salts (like NaCl), which act as an electrolyte, and by acidic conditions. Preventing rust is a major industrial concern and is achieved through methods like painting, coating with oil, or sacrificial protection (using a more reactive metal like zinc).
Solved Examples:
-
What two substances are essential for iron to rust?
Solution: Oxygen (from the air) and water. -
Why does a car rust faster in coastal areas or in regions where salt is used
on icy roads?
Solution: Salt dissolved in water creates an electrolyte solution. This solution speeds up the electrochemical process of rusting by facilitating the movement of ions, making the corrosion of the car's iron body panels happen much more quickly. -
An iron nail is placed in a test tube containing boiled, deionized water,
and a layer of oil is placed on top. Will the nail rust? Explain.
Solution: No, the nail will not rust. Boiling the water removes dissolved oxygen. The layer of oil prevents oxygen from the air from dissolving back into the water. Since oxygen is one of the two essential conditions for rusting, the process cannot occur. -
How is the oxide layer on iron different from the oxide layer on
aluminum?
Solution: The oxide layer on aluminum ($Al_2O_3$) is tough, non-porous, and adheres strongly to the surface, protecting the metal. The oxide layer on iron (rust, $Fe_2O_3 \cdot nH_2O$) is flaky, porous, and does not adhere well, offering no protection to the underlying metal.