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Define 'analyte' in titration.
The solution with the unknown concentration in a titration.
Define 'titrant' in titration.
The solution with the known concentration that is added to the analyte in a titration.
Define 'equivalence point'.
The point in a titration where the moles of titrant added equals the moles of analyte in the solution.
Define 'titration curve'.
A plot of pH change versus the volume of titrant added during a titration.
Define 'half-equivalence point'.
The point in a weak acid/base titration where half of the acid/base has been neutralized, resulting in pH = pKa or pOH = pKb.
Define 'buffer'.
A solution that resists changes in pH due to the presence of a weak acid and its conjugate base, or a weak base and its conjugate acid.
What is the effect of adding excess OH- after the equivalence point?
The solution becomes basic, and the pH increases.
What is the effect of having a buffer solution during a titration?
The solution becomes less responsive to pH changes.
What happens if you use an indicator with a pH range that does not include the equivalence point?
The endpoint of the titration will not accurately reflect the equivalence point, leading to inaccurate results.
What is the effect of using a burette that is not properly calibrated?
The volume measurements will be inaccurate, leading to errors in the calculated concentration of the analyte.
What is the effect of not balancing the chemical equation before performing titration calculations?
The stoichiometric coefficients will be incorrect, leading to an incorrect calculation of the unknown concentration.
What happens when a weak acid is titrated with a strong base?
A conjugate base is formed, affecting the pH at the equivalence point, making it basic.
Describe the steps to find the concentration of HF when titrating with NaOH, given the equivalence point occurs when 20mL of 0.1M NaOH is added to 10mL of HF.
1. Use the equation $M_aV_a = M_bV_b$. 2. Plug in the known values: $M_a(10 \text{mL}) = (0.1 \text{M})(20 \text{mL})$. 3. Solve for $M_a$: $M_a = (0.1)(20)/10 = 0.2 \text{M}$.
What are the steps to determine the pH of a solution formed from titrating a weak acid with a strong base?
1. Write the balanced reaction equation. 2. Determine the initial moles of weak acid and strong base. 3. Perform stoichiometry to find the remaining moles of weak acid and conjugate base after the reaction. 4. Use the Henderson-Hasselbalch equation to calculate the pH.
Outline the general steps of acid-base titration.
1. Prepare the solutions of titrant and analyte. 2. Carefully add titrant to the analyte using a burette. 3. Monitor the pH change using an indicator or pH meter. 4. Stop adding titrant when the equivalence point is reached (indicated by a color change or a sharp pH change). 5. Calculate the unknown concentration using stoichiometry.
Describe the steps to determine the pH of a solution formed from titrating a weak base with a strong acid.
1. Write the balanced reaction equation. 2. Determine the initial moles of weak base and strong acid. 3. Perform stoichiometry to find the remaining moles of weak base and conjugate acid after the reaction. 4. Use the Henderson-Hasselbalch equation (for bases) to calculate the pOH. 5. Convert pOH to pH using pH = 14 - pOH.
What are the steps to calculate the pH at the equivalence point of a weak acid-strong base titration?
1. Determine the moles of conjugate base formed at the equivalence point. 2. Calculate the concentration of the conjugate base. 3. Set up an ICE table for the hydrolysis of the conjugate base. 4. Calculate the hydroxide ion concentration using the Kb expression. 5. Calculate the pOH and then the pH.