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Chapter 3: Problem 48

Which two of the following reactions are oxidationreduction reactions? Explain your answer briefly. Classify the remaining reaction. (a) \(\mathrm{CdCl}_{2}(\mathrm{aq})+\mathrm{Na}_{2} \mathrm{S}(\mathrm{aq}) \rightarrow \mathrm{CdS}(\mathrm{s})+2 \mathrm{NaCl}(\mathrm{aq})\) (b) \(2 \mathrm{Ca}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{CaO}(\mathrm{s})\) (c) \(4 \mathrm{Fe}(\mathrm{OH})_{2}(\mathrm{s})+2 \mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{O}_{2}(\mathrm{g}) \rightarrow 4 \mathrm{Fe}(\mathrm{OH})_{3}(\mathrm{s})\)

Short Answer

Expert verified
Reactions (b) and (c) are redox. Reaction (a) is double displacement.

Step by step solution

01

Understanding Oxidation-Reduction

In an oxidation-reduction (redox) reaction, there is a transfer of electrons between two species. This involves changes in oxidation states of the elements involved. To identify redox reactions, we need to find reactions where oxidation states change.
02

Evaluate Reaction (a)

The reaction is: \( \mathrm{CdCl}_{2}(\mathrm{aq})+\mathrm{Na}_{2} \mathrm{S}(\mathrm{aq}) \rightarrow \mathrm{CdS}(\mathrm{s})+2 \mathrm{NaCl}(\mathrm{aq}) \).- Cadmium (Cd) starts in the +2 state in \( \mathrm{CdCl}_{2} \) and remains +2 in \( \mathrm{CdS} \).- Sulfur (S) starts in the -2 state in \( \mathrm{Na}_{2} \mathrm{S} \) and remains -2 in \( \mathrm{CdS} \).- Sodium (Na) remains +1 throughout.- Chlorine (Cl) remains -1 throughout.Since there are no changes in oxidation states, this is not a redox reaction. It is a double displacement reaction.
03

Evaluate Reaction (b)

The reaction is: \( 2 \mathrm{Ca}(\mathrm{s})+\mathrm{O}_{2}(\mathrm{g}) \rightarrow 2 \mathrm{CaO}(\mathrm{s}) \).- Calcium (Ca) starts at 0 in its elemental form and changes to +2 in \( \mathrm{CaO} \).- Oxygen (O) starts at 0 in \( \mathrm{O}_{2} \) and changes to -2 in \( \mathrm{CaO} \).This reaction involves changes in oxidation states, making it a redox reaction.
04

Evaluate Reaction (c)

The reaction is: \( 4 \mathrm{Fe}(\mathrm{OH})_{2}(\mathrm{s})+2 \mathrm{H}_{2} \mathrm{O}(\ell)+\mathrm{O}_{2}(\mathrm{g}) \rightarrow 4 \mathrm{Fe}(\mathrm{OH})_{3}(\mathrm{s}) \).- Iron (Fe) changes from +2 in \( \mathrm{Fe}(\mathrm{OH})_{2} \) to +3 in \( \mathrm{Fe}(\mathrm{OH})_{3} \).- Oxygen (O) changes from 0 in \( \mathrm{O}_{2} \) to a -2 state in water and the hydroxide.This reaction involves changes in oxidation states, indicating it is a redox reaction.
05

Answer Summary

Reactions \((b)\) and \((c)\) are oxidation-reduction reactions due to changes in oxidation states. Reaction \((a)\) is a double displacement reaction with no changes in oxidation states.

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Key Concepts

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

Electron Transfer
In oxidation-reduction (redox) reactions, electron transfer plays a key role. Electrons move from one reactant to another, and this transfer is what drives the change in chemical properties.
For instance, consider the reaction involving calcium and oxygen:
  • Calcium loses electrons and becomes oxidized, transforming from its elemental form (neutral, no charge) to an ion with a charge of +2.
  • Oxygen gains electrons, moving from a neutral state to a -2 charge as it forms the oxide ion.
This transfer of electrons is vital in determining how elements interact during a reaction. By examining which elements lose or gain electrons, you can identify a redox reaction efficiently.
Oxidation States
Oxidation states help us track electron transfer in a reaction. They are like a bookkeeping system for electron charges, showing how many electrons an atom has gained or lost in a compound.
To recognize changes in oxidation states:
  • Evaluate each element in both reactants and products.
  • Check if these numbers change during the reaction.
For example, sulfur in reaction (a) started and ended with the same oxidation state (-2), indicating no electron transfer. However, in reactions (b) and (c), elements did change their oxidation states, confirming the presence of redox reactions. Understanding oxidation states simplifies the process of identifying electron transfer.
Redox Reactions
Redox reactions are chemical processes where oxidation and reduction occur simultaneously. These reactions involve changes in oxidation states because of electron transfer.
  • In oxidation, an element loses electrons, increasing its oxidation state.
  • In reduction, an element gains electrons, decreasing its oxidation state.
Reaction (b) with calcium and oxygen is a classic redox reaction. Calcium's oxidation state increases as it gives electrons to oxygen, which reduces its oxidation state. This interplay defines redox reactions, making them an essential topic in chemistry as electrons are redistributed between different atoms.
Double Displacement Reaction
A double displacement reaction involves the exchange of ions between two compounds, without any change in oxidation states. This means no electron transfer occurs.
In reaction (a), cadmium and sodium swap partners with sulfide and chloride ions:
  • Cadmium pairs with sulfide to form a precipitate.
  • Sodium pairs with chloride to form sodium chloride in solution.
This type of reaction is driven by the formation of a solid precipitate or the creation of water and is distinct from redox reactions. Understanding this distinction helps in quickly identifying whether a reaction involves electron transfer or not.

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Most popular questions from this chapter

Balance the following equations, and then write the net ionic equation. Show states for all reactants and products (s, \(\ell, \mathrm{g},\) aq). (a) the reaction of silver nitrate and potassium iodide to give silver iodide and potassium nitrate (b) the reaction of barium hydroxide and nitric acid to give barium nitrate and water (c) the reaction of sodium phosphate and nickel(II) nitrate to give nickel(II) phosphate and sodium nitrate

Bromine is obtained from sea water by the following redox reaction: $$\mathrm{Cl}_{2}(\mathrm{g})+2 \mathrm{NaBr}(\mathrm{aq}) \rightarrow 2 \mathrm{NaCl}(\mathrm{aq})+\mathrm{Br}_{2}(\ell)$$ (a) What has been oxidized? What has been reduced? (b) Identify the oxidizing and reducing agents.

Write an equation that describes the equilibrium that exists when the weak acid benzoic acid \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{CO}_{2} \mathrm{H}\right)\) dissolves in water. Identify each of the four species in solution as either Bronsted acids or Bronsted bases. Does the equilibrium favor the products or the reactants? (In acting as an acid, the \(-\mathrm{CO}_{2} \mathrm{H}\) group supplies \(\left.\mathrm{H}^{+} \text {to } \mathrm{form} \mathrm{H}_{3} \mathrm{O}^{+} .\right)\)

Write two chemical equations, one in which \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) is a Bronsted acid (in reaction with the carbonate ion, \(\left.\mathrm{CO}_{3}^{2-}\right),\) and a second in which \(\mathrm{HPO}_{4}^{2-}\) is a Bronsted base (in reaction with acetic acid, \(\mathrm{CH}_{3} \mathrm{CO}_{2} \mathrm{H}\) ).

Which compound or compounds in each of the following groups is (are) expected to be soluble in water? (a) \(\mathrm{CuO}, \mathrm{CuCl}_{2}, \mathrm{FeCO}_{3}\) (b) \(\mathrm{AgI}, \mathrm{Ag}_{3} \mathrm{PO}_{4}, \mathrm{AgNO}_{3}\) (c) \(\mathrm{K}_{2} \mathrm{CO}_{3}, \mathrm{KI}, \mathrm{KMnO}_{4}\)
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