1.What reagent is used as the indicator for the neutralization?
a)citric acid
b)sodium hydroxide
c)sodium citrate
d)phenolphthalein
2.
What is the mole ratio of NaOH to citric acid for this neutralization reaction?
a)2 mole NaOH/1 mole citric acid
b)1 mole NaOH/1 mole citric acid
c)1 mole NaOH/2 mole citric acid
d)1 mole NaOH/3 mole citric acid
e)3 mole NaOH/1 mole citric acid
3.If 35.00 mL of 0.1104 M NaOH were used to reach the neutralization point, how many moles of NaOH were dispensed?
a)3.154 x 10-3 mol NaOH
b)1.104 x 10-4 mol NaOH
c)258.8 mol NaOH
d)3.864 x 10-3 mol NaOH
4. Using the above mole ratio, convert the moles of NaOH in the previous question to moles of citric acid.
a)1.288 x 10-3 mol citric acid
b)1.159 x 10-2 mol citric acid
c)3.864 x 10-3 mol citric acid
d)1.932 x 10-3 mol citric acid
5.
How many grams of citric acid were neutralized in the question above? (MM = 192.12 g/mol)
a)0.7425 g citric acid
b)0.2475 g citric acid
c)0.0825 g citric acid
d)6.704 x 10-6 g citric acid
These are my answers
1. I think it's d (phenolphthalein)
2. I think it's e ( 3:1)
3. I think it's e
4. I don't know
5. I don't know
If you could double check my answers for #1-3 that would be great! But most importanly if you could help with 4-5 because I am very confused on those. I have also included a picture of my lab beucase these are pre-lab qustions, but I don't think you'll need it. Thank you so much
![**Educational Content: Citric Acid Titration**
**Experiment Overview:**
Many foods contain weak acids. For instance, sour milk contains lactic acid, vinegar contains acetic acid, soda drinks contain phosphoric acid, and fruit juices/drinks contain citric acid. In this experiment, you will prepare a solution of citric acid and then verify its concentration by titrating it with a NaOH solution of known molarity. The equivalence point will be detected by an indicator called phenolphthalein that turns from colorless to pink at its endpoint. Some acids are polyprotic, meaning that each molecule of acid produces more than one H⁺ ion in the solution, in which case the equations need to be balanced accordingly. The molecular equation for the neutralization reaction that occurs during the titration today is:
\[ \text{(citric acid) } \text{H}_3\text{C}_6\text{O}_7\text{H}_{5(aq)} + 3 \text{NaOH}_{(aq)} \rightarrow \text{Na}_3\text{C}_6\text{O}_7\text{H}_{5(aq)} + 3 \text{H}_2\text{O}_{(l)} \]
**Procedure:**
The goal of these titrations is to determine the molar concentration (in mass percent) of citric acid in the solution that you prepare.
- **Prepare 200 mL of 0.030 M citric acid solution** using a beaker, solid citric acid, a graduated cylinder, a balance, and water. Record your procedure (including the mass of citric acid that you measured) in the report sheet.
- In a test trial, titrate your aqueous citric acid solution (known concentration) with a standardized sodium hydroxide solution. Follow the procedure below, steps 1-5. Compare your result with the expected concentration of citric acid (about 0.030 M). If your result agrees reasonably well with the expected value, continue to the next set of titrations.
1. **Use a 20 mL volumetric pipette** to measure 20 mL of your citric acid solution into a 125 mL Erlenmeyer flask.
2. **Do not forget this step!** Add 2-3 drops of phenolphthalein indicator solution to the Erlenmey](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F1bf8ae83-2003-4837-b9fe-b672e6e8d8b8%2F0f120101-e4c0-4a1d-8ca6-a298418d8a2b%2Fo1itlsl.jpeg&w=3840&q=75)
![Sure, here is the transcription of the document suitable for an educational website:
---
### Titrating a Weak Acid (Citric Acid) with a Strong Base
**Reminders:**
- Read this document completely before coming to lab.
- Watch the videos associated with this activity before coming to lab.
- Complete the prelab questions before coming to lab.
- Wear the correct protective equipment in the lab.
- Record all data in ink.
- Make sure that you clean up correctly after the lab.
**Learning Outcomes:**
- Demonstrate the use of balances, burets, pipettes, graduated cylinders.
- Demonstrate the preparation of solutions of known concentrations.
- Measure solution concentrations by titration.
**Introduction:**
A *weak acid* is an acid that only partially dissociates when dissolved in water. There is an equilibrium established between the acid in its molecular form (HA) and the ions it produces (H⁺, A⁻):
\[ HA_{(aq)} \rightleftharpoons H^+_{(aq)} + A^-_{(aq)} \]
A *strong base* is a base that completely dissociates in water. The base is present only in the form of ions:
\[ NaOH_{(aq)} \rightarrow Na^+_{(aq)} + OH^-_{(aq)} \]
When a strong base and a weak acid are mixed together, what can be thought of as a two-step process occurs. First, the acid dissociates as shown above. Then the H⁺ from the acid is neutralized by the OH⁻ from the base:
\[ H^+_{(aq)} + OH^-_{(aq)} \rightarrow H_2O_{(l)} \]
The complete and net ionic equation for the overall acid-base reaction are:
\[ HA_{(aq)} + Na^+_{(aq)} + OH^-_{(aq)} \rightarrow H_2O_{(l)} + A^-_{(aq)} + Na^+_{(aq)} \]
\[ HA_{(aq)} + OH^-_{(aq)} \rightarrow H_2O_{(l)} + A^-_{(aq)} \]
A quantitative analysis technique that takes advantage of solution reactions is called a *titration*. In a titration, a solution of known molarity, called a standard solution, is dispensed from a buret and allowed to react with an unknown amount](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F1bf8ae83-2003-4837-b9fe-b672e6e8d8b8%2F0f120101-e4c0-4a1d-8ca6-a298418d8a2b%2Fd1s0oe3.jpeg&w=3840&q=75)
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