14.7 Acid-Base Titrations

14.7 Acid-Base Titrations

  • Acid-Base Equilibria regulates the effects of the carbonate buffering system on the blood's pH.
    • An increase in breathing removes CO2 from the blood through the lungs and lowers the H3O+ in the body.
    • If the blood is too alkaline, a lower breath rate increases CO2 concentration in the blood, driving the equilibrium reaction other way, increasing H+ and restoring an appropriate pH.
  • View the fer system in the water.
  • titrations can be used to analyze solutions for their acid or base concentrations.
    • The changes in the concentrations of the acidic and basic species present in a solution during the process of a titration will be explored in this section.
  • When studying acid-base reactions in solution, we focused on the point at which the acid and base were equivalent.
    • Before, during, or after neutralization, no consideration was given to the solution's pH.
  • Since HCl is a strong acid, we can assume that all of it splits.
    • When the base solution is added, it also provides OH- ion.
  • Only those of the two that were in excess remain, and their concentration determines the pH.
    • The solution is initially acidic, but eventually all the hydronium ion present from the original acid are neutralized, and the solution becomes neutral.
    • The solution becomes basic as more base is added.
  • If n(H+)0 - n(OH-)0 > 0 and so n(H+) > 0.
    • There are no OH- particles left to counteract the hydronium ion that is left from the autoionization of water.
  • The acid and base solutions have the same volumes and concentrations.
  • During a titration, we calculated the pH at four points.
    • There is a detailed sequence of changes in the pH of a strong acid and a weak acid.

  • A strong acid with a strong base is the simplest acid-base reaction.
    • The middle portion of the curve increases rapidly while the first portion increases slowly.
    • The equivalence point for the titration is located at the halfway point of the curve.
    • It shows when equivalent quantities of acid and base are present.
    • The points on the titration curve can be calculated using solution stoichiometry if the acid has a strong base.
  • The NaOH has an equivalent point of 7.00 pH.
    • The NaOH has an equivalent point of 8.72 pH.
  • The titration of a weak acid with a strong base is more complicated than just discussing, but it follows the same general principles.

The curve of the titration is shown in Figure 14.21

  • The acids' initial volume and molarity are the same, but there are differences between the two curves.
    • The titration curve for the weak acid begins at a higher value and continues up to the equivalence point.
    • acetic acid is a weak acid which is partially ionized.
  • The acid-base equilibria can be found after 25.00 mL of the NaOH solution has been added.

  • The equivalence point of this titration is more than 7.
  • CH3 CO2 is converted into 3CO2H.
  • Their concentrations are the same.
  • The result is the same as for the strong acid-strong base titration example, since the amount of strong base added moves the solution past the equivalence point.
  • When the hydronium ion concentration reaches a particular value, certain organic substances change color.
    • Weak organic acids or weak organic bases are acid-base indicators.

  • The nonionized form of HIn is red while the anion of In- is yellow.
    • The equilibrium is shifted toward the nonionized red form when we add acid to a solution of methyl orange.
    • We shift the equilibrium towards the yellow form if we add base.
    • This behavior is similar to the action of buffers.
  • The visible result of the ratio of the concentrations of the two species is an indicator's color.
    • If most of the indicator is In-, we can see the color of the In-ion, which is yellow.
    • The color of the HIn molecule is red if it is present.
  • The ratio of [In-] [HIn] varies with the concentration of hydronium ion.
  • The Henderson-Hasselbalch equation can be used to describe the equilibrium of indicators.
  • There is no change in color when the hydronium ion concentration increases.
  • There are many different acid-base indicators that can be used to determine the approximate pH of an unknown solution by a process of elimination.
  • The ranges of color change are shown in the chart.
  • Titration curves can be used to pick an indicator that will show a sharp color change.
    • An indicator that has a color change interval that brackets the pH at the equivalence point of the titration is the best choice.
  • The color change intervals of three indicators are shown in The equivalence points of the titration of the strong acid and of the weak acid are located in the color- change interval of phenolphthalein.
    • It can be used for titrations of either strong acid with strong base or weak acid with strong base.
  • The curve for the titration of weak acid with strong base is shown in the graph.
    • The shaded areas show the ranges for the color change of phenolphthalein, litmus, and methyl orange.
  • The equivalence point is an indicator for the HCl titration.
    • We should not use litmus for the CH3CO2H titration because it does not leave the range until 25 mL has been added.
    • litmus is useless as an indicator of the equivalence point because the color change is gradual and takes place during the addition of 13 mL of NaOH.
  • We could use methyl orange for the HCl titration, but it wouldn't give very accurate results because it doesn't complete its color change before the equivalence point is reached.
    • The color change begins after about 1 mL of NaOH is added.
    • There is no indication of the equivalence point when the color change is completed long before it occurs.
  • The indicator we use is based on the pH at the equivalence point.
    • At the equivalence point, equimolar amounts of acid and base have been mixed, and the calculation becomes that of the pH of a solution of the salt.

  • A compound that can give a hydrogen ion to another compound is called a Bronsted-Lowry acid.
    • The compound that accepts the protons is called a Bronsted-Lowry base.
    • The conjugate base of the acid is the species remaining after a loss of a protons.
    • The conjugate acid of the base is what formed the species.
    • The formation of the conjugate base of the reactant acid and the formation of the conjugate acid of the reactant base is what occurs in an acid-base reaction.
    • Amphiprotic species can act as both donors and acceptors.
    • The most important amphiprotic species is water.
  • The concentration of H3O+ can be expressed as the pH of the solution.
    • The concentration of OH- can be expressed as the pOH of the solution.
    • The acid or base ionization constants can be used to determine the strengths of the acids and bases in the solution.
  • Strong acids are completely ionized because their conjugate bases are weaker than water.
    • Weak acids have conjugate bases that are strong enough to compete with water for possession of protons.
    • Strong bases react with water.
    • Weak bases don't give much hydroxide ion.
    • The strengths of the two acids increase from left to right across a period of the periodic table.
    • As the oxidation number of the element increases, the strengths of oxyacids that contain the same central element increase.
    • As the electronegativity of the central element increases, the strengths of oxyacids increase as well.
  • The properties of aqueous solutions of Bronsted-Lowry acids are due to the presence of hydronium ion.
    • The neutralization occurs when the solution of acids and bases results from the reaction of hydronium and hydroxide to form water.
    • The product solutions may be slightly acidic or basic due to the salts formed in neutralization reactions.
  • The amount of the ion in the solution determines the pH.
  • A polyprotic acid has more than one ionizable protons.
    • The acids ionize in steps.
  • A buffer solution is a mixture of an acid and its conjugate base.
    • When a small amount of acid or base is added to the buffer solution, the hydronium ion concentration doesn't change much.
    • The acid in the buffer reacts with the base.
  • A titration curve is a graph that shows the change in the basic solution's pH.
  • The titration curve has characteristics that are dependent on the solution being titrated.
    • The solution's pH may be greater than or less than 7.00 at the equivalence point.
    • The range of the color change of the indicator is one of the factors that can affect the choice of an indicator.
  • NH3 is both a conjugate acid and a conjugate base.
  • 2 PO4 is both an acid and a base.
  • Write chemical equations to show the amphiprotic character of the following species.

  • The OpenStax book is free and can be found at http://cnx.org/content/col11760/1.9.
  • acetic acid, CH3CO2H, is a weak acid that causes the smell of vinegar.
  • The ammonia solution in the water is weak.
  • hydrochloric acid is found in gastric juice.
    • Milk of Magnesia can be used to counteract excess stomach acid.
    • Write a balanced equation for the neutralization reaction.
  • In order to prevent the growth of algae in swimming pools, nitric acid reacts with insoluble copper(II) oxide.
  • Explain your reasoning for each of the following pairs if you know which acid is stronger.
  • Predict which compound in each of the following pairs is more acidic.
  • In order to increase acidity or basicity, rank the compounds in each of the following groups.
  • In order to increase acidity or basicity, rank the compounds in each of the following groups.
  • There is a limited amount of ionize in water.
    • The active ingredient in aspirin is C6H4OH(CO2H).
    • The weak acid is the carboxyl group.
    • The OH group bonds to an aromatic ring and acts as a weaker acid.

  • NH3 is a stronger acid than C6H5NH2.

  • acetic acid is 5.0% by mass solution.
  • Lactic acid, CH3CH(OH)CO2H, an acid found in the blood after strenuous exercise, is 1.36 x 10-4.
  • The solution of household ammonia has a pH of 11.
  • Novocaine is the salt of the base procaine and hydrochloric acid.
    • 7 x 10-6 is the constant for procaine.
    • Assuming the density of the solution is 1.0 g/mL, what are the H3O+, OH-, and pH of a 2.0% solution?
  • There is no need for calculations to answer this question.

  • CH3CO2C6H4CO2H is acetylsalicylic acid.
    • There is a saturated solution of salicylic acid and a solution of aspirin.
  • The ion is an amphiprotic species that can act as either an acid or a base.

  • The solution was reduced to 0.100 L.
  • A curve is drawn for a series of solutions.
    • The plot shows the total on the vertical axis and the total concentration on the horizontal axis.
  • On the horizontal axis, there is a total concentration of NH3 that is ionized and nonionized.