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Chapter 15: Problem 41

Why do we normally not quote \(K_{\mathrm{a}}\) values for strong acids such as \(\mathrm{HCl}\) and \(\mathrm{HNO}_{3}\) ? Why is it necessary to specify temperature when giving \(K_{\mathrm{a}}\) values?

Short Answer

Expert verified
We don't quote \(K_{\mathrm{a}}\) values for strong acids like \(\mathrm{HCl}\) and \(\mathrm{HNO}_{3}\) because they completely dissociate in water, hence their \(K_{\mathrm{a}}\) values would be extremely large. It's necessary to specify the temperature when giving \(K_{\mathrm{a}}\) values because these values are temperature-dependent, and an increase in temperature results in a larger \(K_{\mathrm{a}}\) value.

Step by step solution

01

Understanding Acid Strength and \(K_{\mathrm{a}}\) values

The strength of an acid is determined by its ability to donate protons (or \(H^{+}\) ions) to a solution. In water, strong acids like \(\mathrm{HCl}\) and \(\mathrm{HNO}_{3}\) completely dissociate into their ions, meaning nearly every molecule of these acids donate a proton to the solution. Therefore, their \(K_{\mathrm{a}}\) values would be extremely large and it's not necessary to quote these values. It's more relevant to quote \(K_{\mathrm{a}}\) values for weak acids, which only partially dissociate in water.
02

Temperature and \(K_{\mathrm{a}}\) Values

The value of \(K_{\mathrm{a}}\) (like all equilibrium constants) is temperature-dependent. This is why it's necessary to specify the temperature when giving \(K_{\mathrm{a}}\) values. An increase in temperature would increase the rate of the forward reaction (the dissociation of the acid), thus leading to a larger \(K_{\mathrm{a}}\) value as more products (ions) are formed. Therefore, \(K_{\mathrm{a}}\) values are only true at the specified temperature.

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Key Concepts

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

strong acids
Strong acids, such as hydrochloric acid (HCl) and nitric acid (HNO鈧�), are renowned for their complete dissociation in water. This means that, when dissolved, these acids fully break apart into hydrogen ions (H鈦�) and their respective anions (Cl鈦� for HCl, NO鈧冣伝 for HNO鈧�).
This complete ionization leads to them being termed as 'strong' because they completely donate their protons, leaving no molecular form of the acid in the solution.
  • Complete Dissociation: Almost all molecules release their protons.
  • Extremely High Ka Values: Because they fully dissociate, their acid dissociation constant ( K_a) is very large, often not practically defined.
  • Functional in Solutions: Strong acids significantly affect pH levels of solutions they are added to.
Interestingly, the strength of a strong acid isn't typically depicted by its Ka value. This is because the Ka value is so large that it isn't practical to measure or quote it for such fully dissociated acids.
temperature dependence
Temperature plays a crucial role in the behavior of acid dissociation constants, Ka. These constants are not fixed and can vary with changes in temperature, much like other chemical equilibria.
To fully understand the behavior of acids under different conditions, it's vital to note that the value of Ka will increase with rising temperature. This is because increased temperature typically speeds up the forward reaction, thus leading to more dissociation.
  • Temperature Impacts the Rate of Dissociation: Higher temperatures can lead to more significant ion formation in the solution.
  • Equilibrium State Changes: An adjustment in temperature can shift the equilibrium between the acid and its dissociated ions, necessitating specification of temperature when giving Ka values.
For students, it's important to remember: always consider the temperature at which a Ka value is given. It influences the understanding of acid strength and overall reaction equilibrium.
weak acids
Unlike their stronger counterparts, weak acids do not completely dissociate in water. Instead, a significant portion of the weak acid still remains in its molecular form in solution.
This partial dissociation means that weak acids have a much smaller Ka value compared to strong acids. The Ka of a weak acid provides a good measure of its strength because it quantifies the degree of ionization.
  • Partial Dissociation: Only some molecules donate protons to the solution.
  • Smaller Ka Values: Reflect the incomplete ionization which is typical for weak acids.
  • Adjusts pH Gently: The presence of molecular as well as ionized form moderates the pH level of the solution less drastically than strong acids.
Thus, for weak acids, quoting the Ka value becomes essential for chemists and students alike, as it gives insight into how acidic the solution will become with a certain concentration of the acid.

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Most popular questions from this chapter

List four factors that affect the strength of an acid.

How does the strength of an oxoacid depend on the electronegativity and oxidation number of the central atom?

About half of the hydrochloric acid produced annually in the United States (3.0 billion pounds) is used in metal pickling. This process involves the removal of metal oxide layers from metal surfaces to prepare them for coating. (a) Write the overall and net ionic equations for the reaction between iron(III) oxide, which represents the rust layer over iron, and \(\mathrm{HCl}\). Identify the Br酶nsted acid and base. (b) Hydrochloric acid is also used to remove scale (which is mostly \(\mathrm{CaCO}_{3}\) ) from water pipes (see Chemistry in Action essay "An Undesirable Precipitation Reaction" in Section 4.2 ). Hydrochloric acid reacts with calcium carbonate in two stages; the first stage forms the bicarbonate ion, which then reacts further to form carbon dioxide. Write equations for these two stages and for the overall reaction. (c) Hydrochloric acid is used to recover oil from the ground. It dissolves rocks (often \(\mathrm{CaCO}_{3}\) ) so that the oil can flow more easily. In one process, a 15 percent (by mass) HCl solution is injected into an oil well to dissolve the rocks. If the density of the acid solution is \(1.073 \mathrm{~g} / \mathrm{mL},\) what is the \(\mathrm{pH}\) of the solution?

From the relationship \(K_{\mathrm{a}} K_{\mathrm{b}}=K_{\mathrm{w}},\) what can you deduce about the relative strengths of a weak acid and its conjugate base?

Calculate the concentrations of \(\mathrm{H}^{+}, \mathrm{HCO}_{3}^{-},\) and \(\mathrm{CO}_{3}^{2-}\) in a \(0.025 \mathrm{M} \mathrm{H}_{2} \mathrm{CO}_{3}\) solution.
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