Term
| What happens when NaC2H3O2 is added to a solution of HC2H3O2? |
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Definition
| Because C2H3O2- is a weak base, the pH of the solution increases; that is, [H+] decreases. |
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Term
| In which direction will the equilibrium of this reaction shift to? |
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Definition
| This equilibrium will shift to the left, thereby decreasing the equilibrium concentration of [H+]. |
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Term
| What is the common-ion effect? |
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Definition
| The extent of ionization of a weak electrolyte is decreased by adding to the solution a strong electrolyte that has an ion in common with the weak electrolyte. |
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Term
| Is the ionization of a weak base also decreased by the addition of a common ion? |
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Definition
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Term
| What is a buffered solution (or buffer)? |
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Definition
| A solution which contains a weak conjugate acid-base pair that can resist drastic changes in pH upon the addition of small amounts of strong acid or strong base. |
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Term
| Why does a buffer resist changes in pH? |
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Definition
| A buffer resists changes in pH because it contains both an acidic species to neutralize OH- ions and a basic one to neutralize H+ ions. |
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Term
| Why does a buffer resist changes in pH? |
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Definition
| A buffer resists changes in pH because it contains both an acidic species to neutralize the [OH-] ions and a basic one to neutralize [H+] ions. |
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Term
| How are buffers prepared? |
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Definition
| Buffers are prepared by mixing a weak acid or a weak base with a salt of that acid or base. |
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Term
| Under what conditions will buffers most effectively resist a change in pH in either direction? |
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Definition
| When the concentrations of weak acid and conjugate base are about the same. |
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Term
| What are the two important characteristics of a buffer? |
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Definition
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Term
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Definition
| Buffer capacity is the amount of acid or base the buffer can neutralize before the pH begins to change to an appreciable degree. |
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Term
| What determines a buffer's capacity? |
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Definition
| The amount of acid and base from which the buffer is made. |
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Term
| What determines the pH of the buffer? |
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Definition
| The pH of the buffer depends on the Ka for the acid and on the relative concentrations of the acid and base that comprise the buffer. |
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Term
| What is the Henderson-Hasselbach equation? |
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Definition
| pH = pKa + log ([base] / [acid]) |
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Term
| Explain why a mixture of HC2H3O2 and NaC2H3O2 can act as a buffer while a mixture of HCl and NaCl cannot. |
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Definition
In a mixture of HC2H3O2 and NaC2H3O2, HC2H3O2 reacts with added base and C2H3O2- combines with added acid, leaving [H+] relatively unchanged. Although HCl and Cl- are a conjugate acid-base pair, Cl- has no tendency to combine with added acid to form undissociated HCl. Any added acid simply increases
[H+] in an HCl-NaCl mixture. |
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Term
| Do reactions between strong acids and weak bases or strong bases and weak acids essentially proceed to completion? Under what conditions? |
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Definition
| Yes, as long as the buffering capacity of the buffer is not exceeded, we can assume the strong acid or strong base is completely consumed by reaction with the buffer. |
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Term
| What are the steps to calculate how the pH of a buffer will respond to the addition of a strong acid or a strong base? |
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Definition
1. Consider the acid-base neutralization reaction, and determine its effect on [HX] and [X-]. (Stoichiometry calculation)
2. Use Ka and the new concentrations of [HX] and [X-] from step 1 to calculate [H+]. (Equilibrium calculation) |
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Term
| What is a pH titration curve? |
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Definition
| It is the graph of the pH as a function of the volume of the added titrant. |
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Term
For a strong acid-strong base titration, what determines the initial pH of the solution before the addition of a strong base (or strong acid)?
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Definition
| The pH of the solution is determined by the inital concentration of the strong acid, if the titrant is a strong base, or the initial concentration of the strong base, if the titrant is a strong acid. |
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Term
| For a strong acid-strong base titration, what determines the pH of the solution after a strong base or strong acid is being added but before the equivalence point? |
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Definition
| The pH of the solution is determined by the concentration of acid or base that has not yet been neutralized. |
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Term
| For a strong acid-strong base titration, what is the pH of the solution at the equivalence point? |
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Definition
| The pH of the solution is 7.00 (neutral). |
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Term
| For a strong acid-strong base titration, what determines the pH of the solution after the equivalence point? |
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Definition
| The pH of the solution is determined by the concentration of the excess strong base or excess strong acid in the solution (concentration of titrant). |
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Term
| What is the end point of a titration? |
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Definition
| It is the point in a titration where the indicator changes color to distinguish it from the actual equivalence point that it closely approximates. |
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Term
| For a weak acid-strong base titration, what determines the pH of the solution prior to the addition of the strong base titrant? |
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Definition
| The inital pH is determined by the pH of just the weak acid. |
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Term
| For a weak acid-strong base titration, what determines the pH of the solution just prior to the equivalence point? |
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Definition
| The pH of the solution just prior to reaching the equivalence point is determined by the concentrations of the weak acid and its conjugate-base. |
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Term
| For a weak acid-strong base titration, what is the pH of the solution? What determines this pH? |
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Definition
| The pH of the solution at the equivalence point is greater than 7.00, which is determined by the conjugate-base of the weak acid, a weak base. |
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Term
| For a weak acid-strong base titration, what determines the pH of the solution after the equivalence point? |
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Definition
| The pH of the solution after the equivalence point is determined by the concentration of [OH-] from the excess strong base. |
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Term
| How does a pH titration curve for a weak acid-strong base titration differ from a strong acid-strong base titration? (There are 3 differences.) |
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Definition
1. The solution of the weak acid has a higher inital pH than a solution of a strong acid of the same concentration.
2. The pH change at the rapid-rise portion of the curve near the equivalence point is smaller for the weak acid than it is for the strong acid.
3. The pH at the equivalence point is above 7.00 for the weak acid-strong base titration. |
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Term
| Why is the choice of indicator for a weak acid-strong base titration more critical than it is for a strong acid-strong base titration? |
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Definition
| The pH change near the equivalence point becomes smaller as Ka decreases. |
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Term
| What is the solubility-product constant (or simply the solubility product)? |
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Definition
| It is the equilibrium constant (Ksp) for the equilibrium that exists between a solid ionic solute and its ions in a saturated aqueous solution. |
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Term
| Do solids, liquids, and solvents appear in the equilibrium-constant expressions for heterogeneous equilibria? |
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Definition
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Term
| Why is the concentration of undissolved solid not explicitly included in the expression for the solubility-product constant? |
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Definition
| The concentration of undissolved solid does not appear in the solubility product expression because it is constant as long as there is solid present. |
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Term
| What is the difference between solubility and the solubility-product constant? |
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Definition
| The solubility of a substance is the quantity that dissolves to form a saturated solution. The solubility-product constant (Ksp) is the equilibrium constant for the equilibrium between an ionic solid and its saturated solution. |
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Term
| Besides temperature, name three other factors that affect the solubility of ionic compounds? |
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Definition
| The presence of common ions, the pH of the solution, and the presence of complexing agents. |
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Term
| How is the common-ion effect related to the solubility of a slighly soluble salt? |
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Definition
| The solubility of a slightly soluble salt is decreased by the presence of a second solute that furnishes a common ion. |
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Term
| The concentrations of ions calculated from Ksp sometimes deviate appreciably from those found experimentally. What three factors contribute to this deviation? |
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Definition
| The deviations are due to electrostatic interactions between ions in solution (which can lead to ion pairs), ignoring other equilibria that occur simultaneously in the solution, and the assumption that ionic compounds dissociate completely into their component ions when they dissolve. |
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Term
| Will the solubility of any substance whose anion is basic be affected to some extent by the pH of the solution? |
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Definition
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Term
| The solubility of almost any ionic compound is affected if the solution is made sufficiently acidic or basic. Under what condition are the effects noticeable? |
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Definition
| The effects are very noticeable when one or both ions involved are at least moderately acidic or basic. |
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Term
| What effect does an increase in [H+] (as pH is lowered) have on the solubility of slightly soluble salts containing basic anions? |
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Definition
| The solubility of slightly soluble salts containing basic anions increases as [H+] increases (as pH is lowered). The more basic the anion, the more the solubility is influenced by pH. Salts with anions of negligible basicity (the anions of strong acids) are unaffected by pH changes. |
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Term
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Definition
| A complex ion is an assembly of a metal ion and the Lewis bases bonded to it. |
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Term
| How is the stability of a complex ion in aqueous solution judged? |
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Definition
| Stability of a complex ion in aqueous solution can be judged by the size of the equilibrium constant for its formation from the hydrated metal ion. |
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Term
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Definition
| Kf is the formation constant. |
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Term
| What effect does the presence of suitable Lewis bases (NH3, CN-, or OH-, for example) have on the solubility of metal salts? |
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Definition
| The solubility of metal salts increases in the presence of suitable Lewis bases, if the metal forms a complex with the base. |
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Term
| Some metal hydroxides and oxides that are relatively insoluble in neutral water dissolve in what kind of solutions? |
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Definition
| Strongly acidic and strongly basic solutions. |
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Term
| Substances that are soluble in strong acids and bases because they capable of behaving as either an acid or base are called what? |
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Definition
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Term
| Name four amphoteric substances that are metal hydroxides and oxides. |
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Definition
| The hydroxides and oxides of Al3+, Cr3+, Zn2+, and Sn2+. |
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