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Chapter 14: Problem 63

Explain why most metals used in catalysis are transition metals.

Short Answer

Expert verified
Transition metals are commonly used in catalysis mainly because of their ability to adopt multiple and stable oxidation states, which enable them to transfer electrons with the reactants, thereby lowering the activation energy of the reaction. In addition, they can form complexes, bringing the reactants closer to each other, which further speeds up the reaction.

Step by step solution

01

Understanding Catalysis

Catalysis is a phenomenon where a substance, called a catalyst, speeds up a chemical reaction without being consumed in the process. They work by providing an alternative reaction pathway with a lower activation energy.
02

Characteristics of Transition Metals

Transition metals are elements whose atom has a partially filled d sub-shell. They exhibit unique features such as the ability to form stable, multiple oxidation states and the capability to form complexes.
03

Linking Transition Metals properties with Catalysis

Owing to their ability to adopt multiple and stable oxidation states, transition metals can transfer electrons with species involved in the reaction, thereby lowering the reaction's activation energy. The formation of complexes is also crucial in bringing the reaction participants close together, further enhancing the rate of reaction.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Catalysis
Catalysis involves using a catalyst to increase the speed of a chemical reaction without being consumed. This means that catalysts participate in the reaction cycle multiple times. They do not change the final products of the reaction. Instead, they alter the pathway the reaction takes.
  • A catalyst can be homogeneous, meaning it exists in the same phase as the reactants, or heterogeneous, existing in a different phase.
  • Catalysts are important in a wide range of industries and processes, from food production to the manufacture of pharmaceuticals.
By lowering the activation energy, catalysts enable reactions to proceed more quickly or under less extreme conditions.
Activation Energy
Activation energy is the minimum energy required for reactants to engage in a chemical reaction. It acts as a barrier, preventing reactions from occurring spontaneously.
  • Environments with lower activation energies allow reactions to happen faster or at lower temperatures.
  • Catalysts are substances that help lower this activation energy, providing an alternate route for the reaction.
By reducing the activation energy, catalysts increase the rate of reaction. This is why they are widely used in chemical manufacturing and other industrial processes.
Oxidation States
Oxidation states, sometimes called oxidation numbers, are indicators of the degree of oxidation or reduction a substance can undergo. They represent the hypothetical charge an atom would have if all bonds were ionic.
  • Transition metals can have several different oxidation states, which means they can lose or gain different numbers of electrons.
  • This versatility is due to their partially filled d sub-shells.
These characteristics make transition metals particularly effective as catalysts, as they can facilitate complex reactions by stabilizing various oxidation states.
Chemical Reaction
A chemical reaction is a process where substances known as reactants are transformed into different substances, known as products. During a reaction, bonds between atoms in the reactants break, and new bonds form to create the products.
  • Chemical reactions can be endothermic, absorbing energy, or exothermic, releasing energy.
  • The rate of a chemical reaction can be influenced by various factors such as temperature, concentration, and the presence of a catalyst.
Understanding chemical reactions is vital for industries that rely on changing raw materials into valuable products, such as the pharmaceutical, petrochemical, and food industries.

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Most popular questions from this chapter

Define activation energy. What role does activation energy play in chemical kinetics?

Many reactions involving heterogeneous catalysts are zero order; that is, rate \(=k\). An example is the decomposition of phosphine \(\left(\mathrm{PH}_{3}\right)\) over tungsten (W): $$ 4 \mathrm{PH}_{3}(g) \longrightarrow \mathrm{P}_{4}(g)+6 \mathrm{H}_{2}(g) $$ It is found that the reaction is independent of \(\left[\mathrm{PH}_{3}\right]\) as long as phosphine's pressure is sufficiently high \((\geq 1 \mathrm{~atm}) .\) Explain.

What do we mean by the mechanism of a reaction?

Variation of the rate constant with temperature for the first-order reaction $$ 2 \mathrm{~N}_{2} \mathrm{O}_{5}(g) \longrightarrow 2 \mathrm{~N}_{2} \mathrm{O}_{4}(g)+\mathrm{O}_{2}(g) $$ is given in the following table. Determine graphically the activation energy for the reaction. $$ \begin{array}{lc} \mathrm{T}(\mathrm{K}) & \mathrm{k}\left(\mathrm{s}^{-1}\right) \\ \hline 273 & 7.87 \times 10^{3} \\ 298 & 3.46 \times 10^{5} \\ 318 & 4.98 \times 10^{6} \\ 338 & 4.87 \times 10^{7} \end{array} $$

As we know, methane burns readily in oxygen in a highly exothermic reaction. Yet a mixture of methane and oxygen gas can be kept indefinitely without any apparent change. Explain.
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