I just did a lab on Acids and Bases and I am having trouble with the lab questions I was given to complete.
Concentration of NaOH: 0.09999 M
Volume of NaOH: 3.25 mL
# moles of NaOH: 0.325 moles
Initial Concentration of Weak Acid: ? (How do we calculate this)
Part 2: Hydrolysis of Salts
a) Potassium Chloride has a pH of 7 and a yellow-green universal colour indicator. Explain with an equation to show hydrolysis
b) Sodium Acetate has a pH of 8 and a dark green universal colour indicator. Explain using equation
c) Ammonium Chloride has a pH of 5.5 and orange universal colour indicator. Explain using equations.
Part III: Buffers
Concentration of acetic acid (0.1 M volume used) : 20 mL
Concentration of sodium acetate (0.1 M volume used) : 25 mL
Concentration of hydrochloric acid (0.1 M volume used) : 5 mL
pH Measured of buffer solution: 4.55
pH Measured of buffer +5 mL of 01 M HCl: 4.20
We have to now determine the pH calculated for both solutions above.
I would greatly appreciate the help or any guidance to these questions. Thanks!
Concentration of NaOH: 0.09999 M
Volume of NaOH: 3.25 mL
# moles of NaOH: 0.325 moles
Initial Concentration of Weak Acid: ? (How do we calculate this)
There isn't enough information to work the problem here. (weak acid) = moles NaOH/volume weak acid used in the titration. All this is ok if the acid is a monoprotic acid.
Part 2: Hydrolysis of Salts
a) Potassium Chloride has a pH of 7 and a yellow-green universal colour indicator. Explain with an equation to show hydrolysis
b) Sodium Acetate has a pH of 8 and a dark green universal colour indicator. Explain using equation
c) Ammonium Chloride has a pH of 5.5 and orange universal colour indicator. Explain using equations.
KCl doesn't hydrolyze; therefore, the pH is 7.0.
Part b. Sodium acetate. Sodium is not hydrolyzed but acetate is producing acetic acid and hydroxide ion. You can write the equation and see that the solution is basic.
part c. NH4Cl. Cl is not hydrolyzed. NH4^+ is. I will let you write the equation to obtain H3O + ?? showing that the solution is acidic.
Part III: Buffers
Concentration of acetic acid (0.1 M volume used) : 20 mL
Concentration of sodium acetate (0.1 M volume used) : 25 mL
Concentration of hydrochloric acid (0.1 M volume used) : 5 mL
pH Measured of buffer solution: 4.55
pH Measured of buffer +5 mL of 01 M HCl: 4.20
We have to now determine the pH calculated for both solutions above.
Use the Henderson-Hasselbalch equation.
Thanks for the help
Sure! I can help you with your lab questions. Let's start with the first question:
To calculate the initial concentration of the weak acid, you need to utilize the balanced equation for the reaction that occurred during your lab. The balanced equation will give you the stoichiometric relationship between the acid and the base.
For example, if the reaction is:
Weak Acid (HA) + NaOH --> H2O + NaA
You should know the concentration and volume of the NaOH, and the number of moles of NaOH that reacted. From these values, you can determine the stoichiometric ratio between NaOH and the weak acid (HA).
For instance, if the number of moles of NaOH is 0.325 moles, you can use the stoichiometry to calculate the number of moles of the weak acid (HA) that reacted. Once you have the moles of HA, divide it by the initial volume (in liters) to obtain the initial concentration of the weak acid.
Moving on to Part 2: Hydrolysis of Salts,
a) Potassium Chloride (KCl) has a pH of 7, which indicates that it is neutral. The hydrolysis equation for KCl is:
KCl + H2O ⇌ K+ + Cl-
In this case, since the pH is 7, the concentration of both K+ and Cl- ions are equal, resulting in a neutral pH.
b) Sodium Acetate (CH3COONa) has a pH of 8, indicating that it is slightly basic. The hydrolysis equation for NaAc is:
CH3COONa + H2O ⇌ CH3COO- + Na+
In this case, the CH3COO- ion acts as a weak base and accepts a proton from water, resulting in the release of OH- ions, which makes the solution slightly basic (pH > 7).
c) Ammonium Chloride (NH4Cl) has a pH of 5.5, indicating that it is slightly acidic. The hydrolysis equation for NH4Cl is:
NH4Cl + H2O ⇌ NH4+ + Cl-
In this case, the NH4+ ion acts as a weak acid and donates a proton to water, resulting in the release of H+ ions, which makes the solution slightly acidic (pH < 7).
Lastly, for Part III: Buffers,
To determine the calculated pH of the buffer solutions, you can use the Henderson-Hasselbalch equation:
pH = pKa + log ([A-]/[HA])
Here, pKa is the acid dissociation constant of the weak acid (acetic acid, CH3COOH) used in the buffer system, [A-] is the concentration of the conjugate base (sodium acetate, CH3COONa), and [HA] is the concentration of the weak acid.
By plugging in the values for the pH Measured of the buffer solution and the buffer solution with added HCl, as well as the known concentrations of acetic acid and sodium acetate, you can solve for the pKa and then calculate the pH of both solutions.
I hope this helps you understand how to approach these questions! Let me know if there's anything else I can assist you with.