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lab 14 Acids, Bases, Salts and Buffers “Acids and bases are pHun!!!”
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Lab 14: Acids, Bases, Salts and Buffers “Acids and bases are pHun!!!” Introduction According to Smeureanu and Geggier (2017), acids and bases are important in our daily lives. Acids can be identified by their sour taste and bases have bitter taste. Some common acids used in laboratory are hydrochloric acid, sulfuric acid and nitric acid. Acid is a substanve that yields H+ ions and base yields OH- ions. A salt is a product when acid reacts with base in a neutralization reaction. The salt can be either acidic or basic. A buffer is a solution composed of weak acid and its conjugate base. It can resist changes in pH when acid or base is added. The common buffer systems in our body are bicarbonate buffer, phosphate buffer and protein buffer.

Materials - well plate

- thymol blue

- 0.1M FeSO4

- methyl orange

- 0.1M NH4Cl

- methyl red

- 0.1M Na2HPO4

- bromothymol blue

- 0.1M NaH2PO4

- phenolphthalein

- 0.1M Na3PO4

- alizarin yellow

- 0.1M NaCl

- pH paper

- 0.1M NaF

- 150ml beaker

- 0.1M Na2CO3 - 0.1M sodium acetate - Pasteur pipet

- 0.1M acetic acid - two 100ml beakers - two burettes

- 0.1M HCl

- 0.1M NaOH

- tap water

- deionized water

Observations and Experimental Part 1 0.1M solution of

Thymol blue

Methyl orange

Maroon Dark orange red Light Dark NH4Cl yellow orange Dark Light Na2HPO4 blue orange Dark Light NaH2PO4 yellow orange Dark Light Na3PO4 orange blue Dark Light NaCl orange blue Dark Light NaF orange blue Dark Light Na2CO3 orange blue Table 1: color change of solutions FeSO4

0.1M solution Ka Kb of FeSO4 1.8E-2 5.6E-13 NH4Cl 5.6E-10 1.8E-5 Na2HPO4 4.8E-13 1.6E-7 NaH2PO4 6.2E-8 1.3E-2 Na3PO4 4.8E-13 2.1E-2 NaCl NaF 7.2E-4 1.4E-11 Na2CO3 5.6E-11 1.8E-4 Table 2: pH of solutions

Methyl red

Bromothymo l blue

Phenolphthalei n

Pink

Light yellow

White cloudy

Light green

Clear white

Blue

Pink

Light yellow

Clear white

Blue

Pink

Clear pink Light yellow Pink Light yellow Light yellow Light yellow Light yellow

Aqua/greenish blue Aqua/greenish blue

Cloudy white Cloudy white

Blue

Pink

Calculate d pH

Acid / Base

Relative Strength

7.4 5.1 9.8 4.7 12.6 7.0 8.1 11.6

Base Acid Base Acid Base Salt Base Base

Weak base Weak acid Weak base Weak acid Strong base Neutral Weak base Strong base

Alizarin yellow Clear brown Light yellow Light yellow Light yellow Reddish orange Light yellow Light yellow Reddish orange

Approx. pH(indicator ) 4 4.5 8 5 14 4 5 12

Part 2 pH of buffer – 5 HCl added(ml)

pH(HCl added)

NaOH added(ml)

pH(NaOH added)

2.5

5

2.5

5

5

5

5

5

7.5

4.5

7.5

5

10

4.5

10

5

12.5

4

12.5

6

15

4

15

8

17.5

4.5

17.5

12

20

4.5

20

12

22.5

4.5

22.5

12

25

4.5

25

12

Table 3: titration of buffer pH of water - 5 NaOH drops added

Volume added(ml)

pH(NaOH)

HCl drops added

Volume added(ml)

pH(HCl)

10

0.5

6.5

10

0.5

4.5

10

0.5

10

10

0.5

4.5

10

0.5

11.5

10

0.5

4

10

0.5

11.5

10

0.5

4

10

0.5

12

10

0.5

4

10

0.5

12

-

-

-

10

0.5

12.5

-

-

-

10

0.5

12.5

-

-

-

Table 4: pH of water when HCl and NaOH added

pH vs volume of HCl/NaOH added 14 12 10

pH

8 6 4 2 -30

-20

0 0

-10

10

20

30

volume(ml) pH(buffer)

pH(water)

Graph 1: pH vs volume of HCl/NaOH added to buffer and water

Discussion and Conclusion According to the results in part 1, FeSO4 pH is around 6 since bromothymol blue changed to light yellow. NH4Cl pH is around 4 because methyl orange changed to light orange. Na2HPO4 pH is around 10 because phenolphthalein changed to pink. NaH2PO4 pH is around 4 because methyl orange changed to light orange. Na3PO4 pH is around 10 because phenolphthalein changed to pink. NaCl pH is around 7 because bromothymol blue changed to aqua/greenish blue. NaF pH is around 7 because bromothymol blue changed to aqua/greenish blue. Na2CO3 pH is around 11 because alizarin yellow changed to reddish orange. There seem to be some limitations as well because the calculated pH and those approximated from the pH paper were not close for 3 substances (FeSO4, NaCl and NaF). There could have been misinterpretation of colors when using pH papers. In part 2, two different solutions, acetate and water, were compared to determine which solution was a better buffer. According to the result, acetate was a better buffer than water because when HCl was added to acetate, the pH level was within the buffer range from 4 to 6. This also happened when NaOH was added till 10ml. Focus Question

1. Without conducting an experiment, how could you predict if a species in a solution is acidic or basic? Give some examples. We can determine by looking at the color of each solution after adding the indicators. We can reference the color of relative pH through acid-base indicator chart. If it is less than 7, it is acidic and if it is higher than 7, it is basic. For example, methyl orange’s pH ranges from 1 to 3, so pH 1 shows red color and pH 3 shows yellow color.

2. How might you explain the different strengths of acids and bases using periodic trends and molecular resonance structures? Use your data from part 1 to explain any relationships. Each bond has different properties and these properties affect the bond strength in a molecule. The bond strength between H and a halogen can determine the strength of an acid. Larger the halogen, the weaker the bond gets in which H will be easily lost making it a stronger acid. An example is NaCl is more acidic than NaF because Cl is larger than F. The charge of a molecule effects how basic it is. The more negatively charged a base is, the more easily it will accept cation. An example is NaH2PO4 and NaH2PO4. NaH2PO4 has more negatively charged PO4- ion. Therefore, it should be more basic than NaH2PO4.

3. What happens when a strong acid or strong base is added to a buffer system? Use chemical equations to support your answer. When strong acid HCl is added to a buffer, weak base A- will react with H+ ion produced from strong acid dissociation and more weak acid HA will be form. A- + H+

HA

When strong base NaOH is added to a buffer, weak acid HA will react with OH- to form more conjugate base A- and water. HA + OHReference

A- + H2O

- Smeureanu, G., & Geggier, S. (2017). General Chemistry Laboratory. New York: LAD Custom Publishing.

Post-Lab Assessment Questions 1. What is the hydronium ion concentration and pH of a 0.1M solution of hypochlorous acid, Ka = 3.5E-8? Ka = 3.5E-8 = x2/0.1 x = 5.916E-5 pH = -log5.916E-5 = 4.2

2. Suppose you have an alkaline buffer consisting of 0.2M aqueous ammonia(NH3) and 0.1M ammonium chloride(NH4Cl). What is the pH of the solution? Kb = 1.8E-2 / Ka = 5.6E-10 pH = -log5.6E-10 + log0.2/0.1 = 9.6

3. Calculate the pH of a 0.5M solution of KCN. Ka for HCN is 6.2E-10. Ka = 6.2E-10 / Kb = 1.61E-5 x = 0.5 * 1.61E-5 = 2.84E-3 pOH = -log2.84E-3 = 2.55 pH = 14 – 2.55 = 11.45

4. How do the concentration/volumes of the buffer affect the buffer capacity? E.g. 50ml of 0.1M acetic acid solution with 50ml of 0.1M sodium acetate solution vs. the buffer you made in the lab (25ml of 0.1M acetic acid solution with 25ml of 0.1M sodium acetate solution) Mole of a substance is equivalent to concentration multiplied by volume by raising the concentration or volume of a buffer. Then the buffer capacity will rise by a proportional amount. Therefore, double the concentration of a buffer will double the buffer capacity....