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Chapter 7: Problem 104

(a) One of the alkali metals reacts with oxygen to form a solid white substance. When this substance is dissolved in water, the solution gives a positive test for hydrogen peroxide, \(\mathrm{H}_{2} \mathrm{O}_{2}\). When the solution is tested in a burner flame, a lilac-purple flame is produced. What is the likely identity of the metal? (b) Write a balanced chemical equation for the reaction of the white substance with water.

Short Answer

Expert verified
The likely identity of the metal is potassium (K), as the lilac-purple flame is associated with potassium. The balanced chemical equation for the reaction of the white substance (potassium superoxide) with water is: \[ 2 \mathrm{KO}_{2} (s) + 2 \mathrm{H}_{2}\mathrm{O} (l) \longrightarrow 2 \mathrm{KOH} (aq) + \mathrm{H}_{2}\mathrm{O}_{2} (aq) \]

Step by step solution

01

(a) Identify the alkali metal

We are given that the alkali metal reacts with oxygen to form a solid white substance. The presence of a lilac-purple flame is indicative of the alkali metal being potassium (K). This is because each alkali metal produces a characteristic flame color due to its atomic emission spectrum, and the lilac-purple flame is associated with potassium. Thus, the likely identity of the metal is potassium (K).
02

(b) Write balanced chemical equation for the reaction of the white substance with water

We have established that the metal in question is potassium (K). When potassium reacts with oxygen, it forms potassium superoxide (KO鈧�), which is a solid white substance. The reaction of potassium superoxide with water produces potassium hydroxide (KOH) and hydrogen peroxide (H鈧侽鈧�). The balanced chemical equation for this reaction is: \[ 2 \mathrm{KO}_{2} (s) + 2 \mathrm{H}_{2}\mathrm{O} (l) \longrightarrow 2 \mathrm{KOH} (aq) + \mathrm{H}_{2}\mathrm{O}_{2} (aq) \] This equation represents the reaction between the potassium superoxide and water, resulting in the formation of potassium hydroxide and hydrogen peroxide.

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Key Concepts

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

Alkali Metals
Alkali metals are a group of elements found in group 1 of the periodic table. This family of metals includes lithium, sodium, potassium, rubidium, cesium, and francium. They are known for being highly reactive, especially with water. This reactivity is due to their single valence electron, which they readily lose to form positive ions (cations).
Some characteristics of alkali metals include:
  • They are soft and can be cut with a knife.
  • They have low melting points compared to most other metals.
  • They form ionic compounds such as metal hydroxides when reacting with water.
These metals also display unique colors when burned, due to their atomic emission spectrum. For instance, when potassium is burned, it exhibits a lilac or purple flame. This flame color helps in identifying the specific metal. Overall, understanding the properties of alkali metals is essential in predicting their reactions with various substances, such as oxygen and water.
Potassium Superoxide
Potassium superoxide ( KO鈧�) is a compound formed from the reaction of potassium with oxygen. It's interesting due to its role in generating oxygen and removing carbon dioxide when used in respiratory devices. In this context, however, it's important to focus on its white solid appearance and its reactivity with water.
When potassium superoxide reacts with water, it produces potassium hydroxide (KOH) and hydrogen peroxide ( H鈧侽鈧�). This reaction is quite fascinating because it involves the transformation of a superoxide into a hydroxide and an essential, yet unstable compound, hydrogen peroxide.
This process can be summarized in the chemical equation:
  • KO鈧� (s) + H鈧侽 (l) 鈫� KOH (aq) + H鈧侽鈧� (aq)
The behavior of potassium superoxide showcases the fascinating characteristics of the superoxide ion ( O鈧傗伝). Understanding this is crucial in applications where potassium superoxide is utilized, such as in the self-contained breathing apparatuses for producing oxygen.
Balancing Chemical Equations
Balancing chemical equations is a fundamental skill in chemistry that ensures the law of conservation of mass is upheld. This principle states that matter cannot be created or destroyed in a closed system. Hence, a balanced chemical equation has the same number of atoms of each element on both sides.
To balance equations, follow these steps:
  • Write the unbalanced equation with correct formulas of reactants and products.
  • Count and compare the number of atoms of each element on both sides of the equation.
  • Add coefficients in front of the chemical formulas to equalize the number of atoms for each element.
  • Continue adjusting the coefficients as needed, ensuring that you never change the subscripts in the formulas.
  • Double-check to make sure the equation is balanced.
For instance, in the reaction given: 2KO鈧� (s) + 2H鈧侽 (l) 鈫� 2KOH (aq) + H鈧侽鈧� (aq),
you see that the number of K, O, and H atoms is the same on both sides. This ensures that conservation of mass is strictly followed. A balanced equation allows for accurate predictions of reactants and products in any chemical reaction.

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Most popular questions from this chapter

Give three examples of +2 ions that have an electron configuration of \(n d^{10}(n=3,4,5 \ldots)\)

Write balanced equations for the following reactions: (a) sulfur dioxide with water, (b) lithium oxide in water, \((\mathbf{c})\) zinc oxide with dilute hydrochloric acid, (d) arsenic trioxide with aqueous potassium hydroxide.

Using only the periodic table, arrange each set of atoms in (a) \(\mathrm{Cs}\), Se, Te; order of increasing radius: (b) \(\mathrm{S}, \mathrm{Si}, \mathrm{Sr} ;\) (c) P, Po, Pb.

Which of the following chemical equations is connected to the definitions of (a) the first ionization energy of oxygen, (b) the second ionization energy of ox ygen, and (c) the electron affinity of oxygen? (i) \(\mathrm{O}(g)+\mathrm{e}^{-} \longrightarrow \mathrm{O}^{-}(g)\) (ii) \(\mathrm{O}(g) \longrightarrow \mathrm{O}^{+}(g)+\mathrm{e}^{-}\) (iii) \(\mathrm{O}(g)+2 \mathrm{e}^{-} \longrightarrow \mathrm{O}^{2-}(g)\) (iv) \(\mathrm{O}(g) \longrightarrow \mathrm{O}^{2+}(g)+2 \mathrm{e}^{-}\) \((\mathbf{v}) \mathrm{O}^{+}(g) \longrightarrow \mathrm{O}^{2+}(g)+\mathrm{e}^{-}\)

Potassium superoxide, \(\mathrm{KO}_{2},\) is often used in oxygen masks (such as those used by firefighters) because \(\mathrm{KO}_{2}\) reacts with \(\mathrm{CO}_{2}\) to release molecular oxygen. Experiments indicate that 2 mol of \(\mathrm{KO}_{2}(s)\) react with each mole of \(\mathrm{CO}_{2}(g) .\) (a) The products of the reaction are \(\mathrm{K}_{2} \mathrm{CO}_{3}(s)\) and \(\mathrm{O}_{2}(g) .\) Write a balanced equation for the reaction between \(\mathrm{KO}_{2}(s)\) and \(\mathrm{CO}_{2}(g) .(\mathbf{b})\) Indicate the oxidation number for each atom involved in the reaction in part (a). What elements are being oxidized and reduced? (c) What mass of \(\mathrm{KO}_{2}(s)\) is needed to consume \(18.0 \mathrm{~g} \mathrm{CO}_{2}(g)\) ? What mass of \(\mathrm{O}_{2}(g)\) is produced during this reaction?
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