Question

In each of the following reactions, identify which element is being oxidized and which is being reduced by assigning oxidation numbers. a. \(4 \mathrm{KClO}_{3}(s)+\mathrm{C}_{6} \mathrm{H}_{12} \mathrm{O}_{6}(s) \rightarrow 4 \mathrm{KCl}(s)+6 \mathrm{H}_{2} \mathrm{O}(l)+6 \mathrm{CO}_{2}(g)\) b. \(2 \mathrm{C}_{8} \mathrm{H}_{18}(l)+25 \mathrm{O}_{2}(g) \rightarrow 16 \mathrm{CO}_{2}(g)+18 \mathrm{H}_{2} \mathrm{O}(l)\) c. \(\mathrm{PCl}_{3}(g)+\mathrm{Cl}_{2}(g) \rightarrow \mathrm{PCl}_{5}(g)\) d. \(\mathrm{Ca}(s)+\mathrm{H}_{2}(g) \rightarrow \mathrm{CaH}_{2}(g)\)

Short answer

In the given reactions, a. C is oxidized (0 to +4) and Cl is reduced (+5 to -1). b. C is oxidized (-4 to +4) and O is reduced (0 to -2). c. P is oxidized (+3 to +5) and Cl is reduced (0 to -1). d. Ca is oxidized (0 to +2) and H is reduced (0 to -1).

Step-by-step solution

Assign oxidation numbers for reactants and products

Assign oxidation numbers to each element in the reactants and products of the reaction: 4 KClO3(s) + C6H12O6(s) → 4 KCl(s) + 6 H2O(l) + 6 CO2(g) Reactant side: K: +1 Cl: +5 O: -2 C: 0 H: +1 O: -2 Product side: K: +1 Cl: -1 H: +1 O: -2 C: +4 O: -2

Identify the oxidized and reduced elements

Compare the oxidation numbers and determine which element undergoes oxidation and which element undergoes reduction: - C is oxidized (oxidation number increased from 0 to +4) - Cl is reduced (oxidation number decreased from +5 to -1) #b. Reaction 2: Identifying the oxidized and reduced elements #

Assign oxidation numbers for reactants and products

Assign oxidation numbers to each element in the reaction: 2 C8H18(l) + 25 O2(g) → 16 CO2(g) + 18 H2O(l) Reactants side: C: -4 H: +1 O: 0 Product side: C: +4 O: -2 H: +1 O: -2

Identify the oxidized and reduced elements

Identify which element undergoes oxidation and which undergoes reduction: - C is oxidized (oxidation number increased from -4 to +4) - O is reduced (oxidation number decreased from 0 to -2) #c. Reaction 3: Identifying the oxidized and reduced elements #

Assign oxidation numbers for reactants and products

Assign oxidation numbers to each element in the reaction: PCl3(g) + Cl2(g) → PCl5(g) Reactants side: P: +3 Cl: -1 Cl: 0 Product side: P: +5 Cl: -1

Identify the oxidized and reduced elements

Identify which elements undergo oxidation and which undergo reduction: - P is oxidized (oxidation number increased from +3 to +5) - Cl is reduced (oxidation number decreased from 0 to -1) #d. Reaction 4: Identifying the oxidized and reduced elements #

Assign oxidation numbers for reactants and products

Assign oxidation numbers to each element in the reaction: Ca(s) + H2(g) → CaH2(g) Reactants side: Ca: 0 H: 0 Product side: Ca: +2 H: -1

Identify the oxidized and reduced elements

Identify which elements undergo oxidation and which undergo reduction: - Ca is oxidized (oxidation number increased from 0 to +2) - H is reduced (oxidation number decreased from 0 to -1)

Key concepts

Assigning Oxidation Numbers

In chemical reactions, atoms can gain or lose electrons, altering their oxidation state. Assigning oxidation numbers (ON) helps chemists track these electron transfers. Oxidation numbers are hypothetical charges an atom would have if all bonds were purely ionic. Here are some rules to follow:

• For pure elements, the oxidation number is zero.
• For ions, it is the charge of the ion.
• Oxygen is usually -2, except in peroxides where it's -1.
• Hydrogen is typically +1 when bonded to nonmetals and -1 with metals.
• The sum of oxidation numbers in a neutral compound is zero, while for polyatomic ions, it equals the ion's charge.

Understanding how to assign oxidation numbers is crucial for identifying redox reactions, which is the next key concept.

Redox Reactions

Redox reactions are chemical processes where one substance gives away electrons (oxidizes) and another gains electrons (reduces). The substance losing electrons is the reducing agent, whereas the one gaining electrons is the oxidizing agent.
These reactions are fundamental to a variety of chemical processes, including combustion, respiration, and metallurgy. Let’s break this down using the concept of oxidation states. In essence, if an element’s oxidation number increases during a reaction, it is oxidized; if it decreases, it’s reduced. Take, for example, the oxidation of carbon in a combustion reaction. Carbon starts with an oxidation number of zero and ends with +4, indicating it has lost electrons, underscoring the importance of assigning oxidation numbers.

Chemical Reaction Analysis

Analyzing chemical reactions involves breaking down the processes to understand which reactants are transformed into which products, and how. The analysis goes beyond observing physical changes to interpreting the changes happening on a molecular level, such as bond formation or breakage and electron transfers.
For instance, when assessing the reaction between hydrogen and oxygen to form water, it's vital to note not just that water is produced, but also that oxygen has been reduced and hydrogen oxidized. This analytical approach provides a deeper understanding of the reaction, enabling scientists to predict product formation, gauge reaction spontaneity, and facilitate further applications like energy harnessing or synthesis of new materials.

Identifying Oxidizing and Reducing Agents

In redox chemistry, the substances that accept electrons (and get reduced) are called oxidizing agents, whereas those that donate electrons (and get oxidized) are known as reducing agents. An oxidizing agent gains electrons and its oxidation number decreases, while a reducing agent loses electrons and its oxidation number increases.

• To determine the oxidizing agent, look for the species being reduced.
• To find the reducing agent, identify the species being oxidized.

For instance, in the reaction between magnesium and hydrochloric acid, magnesium acts as the reducing agent because it loses electrons, whereas the hydrogen ions in the acid are the oxidizing agents as they gain electrons. Thus, knowing how to spot these agents is fundamental in understanding the driving forces behind redox reactions.