M2 BIOLAB

• pH of solutions

  Affected by changes in temperatures
• Chemical equilibrium

  Exists in all buffer systems that are usually affected by temperature, concentration, and pressure as learned from Le Chatelier's Principle
• pK changes with temperature

  pH must also change
• Buffered solution

  Mixture of a weak acid and its conjugate base that provides the buffering capacity, its pH changes very little as a small amount of strong acid or base is added into it
• Unbuffered solution

  No buffering capacity so that the pH changes significantly as a small amount of strong acid or base is added into it
• Dissociation constant changes with temperature

  Concentration of the ions also changes
• Buffering effect

  1. If [OH-] increases, then [HA] and [A-] must change

  2. Some of the HA molecules react with the added OH-, making the solution less basic than without the conjugate base(A-)
• Buffering effect

  • If K is small, the buffering effect is small because there's not much HA in the solution
  • If the value of K increases with temperature, then the buffering effect is stronger at a higher temperature
  • If K decreases, then the buffering effect is weaker when it's warmer
• Buffer Effectiveness.

  • An effective buffer should be made of an acid and its conjugate base or a base and its conjugate acid where the Ka value is very similar to the desired pH.
  • The exact ratio of the conjugate base to the acid is determined from the Ka value and the Henderson-Hasselbalch equation for the desired pH. The buffer is most effective when the amounts of acid and its conjugate base are approximately equal.
• pH of solutions

  Directly measured using pH paper and pH meter
• Determining pH by computation

  Using appropriate mathematical formulas and equations such as the Henderson-Hasselbalch Equation
• Henderson-Hasselbalch equation

  pH = pKa + log ([A-] / [HA])
• Effective buffer

  • Made of an acid and its conjugate base or a base and its conjugate acid where the Ka value is very similar to the desired pH
  • Most effective when the amounts of acid and its conjugate base are approximately equal
• Acid dissociation constant (Ka) = [H+][A-]/[HA]
• Buffers can be classified into two types based on their composition:
• Buffer Effectiveness.

  • An effective buffer should be made of an acid and its conjugate base or a base and its conjugate acid where the Ka value is very similar to the desired pH.
  • The exact ratio of the conjugate base to the acid is determined from the Ka value and the Henderson-Hasselbalch equation for the desired pH. The buffer is most effective when the amounts of acid and its conjugate base are approximately equal.
• Strong bases have high Kb values, while weak bases have low Kb values.
• Weak acids have low Ka values, while strong acids have high Ka values.
• In a solution with a pH of 7, there will be more A- than HA because at this pH, the concentration of H+ is relatively small compared to that of OH-. Therefore, the equilibrium lies far to the right, resulting in a higher concentration of A- than HA.
• pH is defined as -log10[H+].
• pKa is defined as -log10(Ka).
• To determine the amount of NaOH needed to raise the pH of a solution from 4.5 to 6.3 using a buffer system, use the Henderson-Hasselbalch equation.
• Buffers work best within their range of effectiveness, which depends on the strength of the acid/base pair.
• A buffer can be used to maintain a constant pH by adding a titrant (acid or base) until the pH reaches the desired level.
• When adding NaOH to a solution containing HCl, the reaction shifts towards the left until it reaches a new equilibrium state.
• To calculate the amount of added NaOH needed to adjust the pH of a given volume of solution, use the formula: moles of NaOH = -(moles of HCl)/1 + (pH - pKa).
• Adding NaOH increases the pH of a solution containing citric acid.
• Increasing the concentration of HCl decreases the pH of a solution containing sodium bicarbonate.
• In a solution with no added acid or base, the concentration of H+ is equal to the concentration of OH-, resulting in a neutral pH of 7.
• Buffers are solutions that resist changes in pH when an acid or base is added.
• The Henderson-Hasselbalch equation can also be used to calculate the pH of a buffer solution given its pKa value and the ratio of conjugate acid (HA) to base (A-) concentrations.
• Acidic buffers contain a weak acid and its salt, while basic buffers contain a weak base and its salt.
• The stronger the base, the weaker the conjugate acid (HA) and vice versa.
• A weak acid has a lower K value than its conjugate base.
• The pH of a buffer solution can be calculated using the Henderson-Hasselbalch equation.
• Acids are proton donors that release hydrogen ions into water, while bases accept hydrogen ions from water.
• The Henderson-Hasselbalch equation relates the pH of a buffer solution to its pKa value and the ratio of conjugate acid to base concentrations.
• The concentration of [HA] decreases while the concentration of [A-] increases when NaOH is added to an acidic solution.
• The concentration of hydrogen ion in a solution determines its pH value.
• Stronger bases are less effective buffers due to their greater tendency to dissociate into hydrogen ions.