CHM 123Chapter 15 - Acids, Bases, Buffers and Titration

CHM 123Chapter 15 - Acids, Bases, Buffers and Titration

CHM 123Chapter 15 - Acids, Bases, Buffers and Titration

  1. Strong acids and strong bases
  2. acids that ionize 100% are called strong acids

HCl(aq)  H+(aq) + Cl-(aq)

  1. bases that ionize 100% are called strong bases

NaOH(aq)  Na+(aq) + -OH(aq)

  1. Weak acids

•Only a few molecules dissociate.

•Most of the weak acid remains as the undissociated (molecular) form of the acid.

•The concentrations of the H3O+and the anion (A-) are small.

HF(aq) + H2O(l) H3O+(aq) + F-(aq)

E.g Weak bases

•are most other bases.

•dissociate only slightly in water.

•form only a few ions in water.

NH3(g) + H2O(l) NH4+(aq)+ -OH(aq)

  1. Buffer Solution and Le Chatelier’s principle

A solution which contains a weak acid and its conjugate base and resists drastic changes in pH.

•typically has an equal concentration of a weak acid and its salt.

•may also contain a weak base and a salt of the conjugate acid.

Weak acid + conjugate base OR Weak base + Conjugate acid

The Henderson-Hasselbalch Equation

HA(aq) + H2O(l) H3O+(aq) + A-(aq)

Buffer capacity:

-refers to the maximum amount of either strong acid or strong base that can be added before a significant change in the pH will occur. This is simply a matter of stoichiometry. The maximum amount of strong acid that can be added is equal to the amount of conjugate base present in the buffer. The maximum amount of base that can be added is equal to the amount of weak acid present in the buffer

How does the buffer resist the drastic change in pH?

HA(aq) + H2O(l) H3O+(aq) + A-(aq)

When a strong acid (H3O+) is added to a buffer solution the conjugate base present in the buffer consumes the hydronium ion converting it into water and the weak acid of the conjugate base.

The net-ionic neutralization reaction isA-(aq) + H3O+(aq) H2O(l) + HA(aq)

This results in a decrease in the amount of conjugate base present and an increase in the amount of the weak acid. The pH of the buffer solution decreases by a very small amount.

When a strong base (OH-) is added to a buffer solution, the hydroxide ions are consumed by the weak acid forming water and the weaker conjugate base of the acid. The amount of the weak acid decreases while the amount of the conjugate base increases. This prevents the pH of the solution from significantly rising, which it would if the buffer system was not present.

The net-ionic neutralization reaction is OH-(aq) + HA(aq)  H2O(l) + A-(aq)

Equation writing review

Write an hydrolysis equation for the following acids and bases

HNO2(acid)

CF2SO2H(acid)

CH5N(base)

HS-(base)

Four types of acid-base neutralization reactions:

  1. Strong acid-strong base
  2. Weak acid-strong base
  3. Strong acid-weak base
  4. Weak acid-weak base

E.gWrite balanced molecular and net-ionic equations for the neutralization of equal molar amounts of the following acids and bases.

1.HNO2 and KOH

2.NaOH and HClO4

15.5Titration

A procedure for determining the concentration of a solution by allowing a carefully measured volume to react with a solution of another substance (the standard solution) whose concentration is known.

Titrant: solution that is added from the buret

Equivalence Point: The point at which stoichiometrically equivalent quantities of acid and base have been mixed together.

Endpoint: is the point at which the titration is complete (usually by a sudden color change), as determined by an indicator

Acid-Base Titration Curves

Titrations are often plotted as an “acid-base titration curves” that displays the pH of the reaction (titration) mixture versus the volume of titrant (added base or acid). We will consider titrations curves for three types of titrations: strong acid-strong base, weak acid-weak base and strong acid-weak base

titrations

The indicator should be chosen so that pH range over which it changes color is in the vertical region of the titration in question. This ensures that when the color change occurs, the volume of titrant added will be as close as possible to the amount corresponding to the equivalence point.

Acid-base indicators are used to signal the end of acid-base titrations

Selecting Solutions for Acid-Base Titrations

If you are titrating an acid, make sure you use a base so that your titration reaction is a neutralization. It should have at least one STRONG reactant so it will go to completion.

For example, if you are titrating the acid CH3COOH (WA), use a STRONG BASE like NaOH, KOH etc. You could not use another acid (like HCl etc.) . Also, since CH3COOH is a WEAK acid, you cannot use a weak base (like NH3)

Also, the concentration ...... of your standard should be relatively close to the concentration of the solution you are titrating so that the volumes used are comparable.

Example:In titrating 25.00 mL samples of NH3 which is approximately 0.1 M, which of the following solutions should be used to determine the [NH3]?

a) 0.00100M HCl b) 0.125 M HCl c) 6.00 M HCl d) 12.0 M HCl e) 0.100M NaOH

TitrationCurve StrongAcid gif

TitrationCurve WeakAcid gif

Calculating the pH of the Titration Solution

You must be able to calculate the pH at any point in the titration of all three titration types

  • Before any titrant has been added
  • Before the equivalence point
  • At the equivalence point
  • After the equivalence point

The calculations for strong acid-strong base titration are significantly different than the calculations for weak acid-strong base or strong acid-weak base. However, the calculations for these latter two types are very similar.

To do this well, you must first identify the type of titration in a given problem. You must have memorized your list of common strong acids and strong bases.

To understand calculations involving weak acids or base, you will need to have a good grasp of these review concepts: Ka, Kb, buffer, neutralization reactions, and hydrolysis.

E.gWrite a balanced ionization equation of HClO. What the equilibrium expression Ka for this acid ionization?

What equilibrium expression should you write for the following reaction

NH3(aq) + H2O(l) -OH(aq) + NH4+(aq)

15.6Strong acid-Strong base titrations

Example: A 60.00 mL of 0.150 M HNO3 is titrated with 0.405 M NaOH at 25oC.

  1. What is the pH before any titrant has been added?
  2. This is not really a titration problem. Remember, the “concentration” of the strong acid HCl is actually the concentration of H3O+ (WHY?) which can be directly converted to pH.
  1. Where is the “equivalence point” of the titration?
  • That is, what volume of titrant (NaOH) must be added to reach the equivalence point?
  • Always calculated with MaVa = MbVv
  1. What is the pH after 10.00 mL of titrant have been added?
  2. Consider neutralization and dilution before calculating [H3O+]
  1. What is the pH at the equivalence point?
  2. At 25oC, the answer is always pH = 7.00. At this point, the acid (HCl) has been completely “neutralized” by adding an “equivalent” amount of base (NaOH). At the equivalence point, a strong acid and a strong base are entirely converted to H2O and “salt” (NaCl in this case). The salt produced will always be neutral (i.e., does not hydrolyze in water), the pH is just the result of the self-ionization of water at 25oC
  1. What is the pH after a total of 50.00 mL of titrant have been added?
  2. Consider neutralization and dilution before calculating [H3O+]

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