Unit
7:
ACIDS
AND
BASES
 Lesson 1: What are Acids and Bases? For thousands of years people have known that vinegar, lemon juice and many other foods taste sour. However, it was not until a few hundred years ago that it was discovered why these things taste sour - because they are all acids. The term acid, in fact, comes from the Latin term acere, which means “sour". In the seventeenth century, the Irish writer and amateur chemist Robert Boyle first labeled substances as either acids or bases (he called bases alkalies) according to the following characteristics: • •

Acids taste sour, are corrosive to metals, change litmus (a dye extracted from lichens) red, and become less acidic when mixed with bases. Bases feel slippery, change litmus blue, and become less basic when mixed with acids.

While Boyle and others tried to explain why acids and bases behave the way they do, the first reasonable definition of acids and bases would not be proposed until 200 years later.

Activity 1: Plant Juice and Litmus Paper as an Indicator In this experiment, you will be testing various acids and alkalis with red cabbage juice and litmus paper indicators. This experiment was first performed in 1664 by Robert Boyle, who from his work with plant juices developed litmus paper. Litmus paper is a form of plant juice indicator, and it still has applications today in the laboratory and in industry. Boyle’s work showed that acids change plant juice and litmus paper from purple to red. Alkalis, on the other hand, change plant juice and litmus paper bluish-green. Using this knowledge, you will determine whether some common everyday items are alkaline or acidic.

To make the red cabbage juice indicator: 1. Boil a few red cabbage leaves in about 100ml of water for 5/10 minutes. 2. Filter the cabbage chunks from the now bluish-purple solution. 3. Allow for the cabbage solution to cool.

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Unit
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 1. From the list of samples that will be tested (see the table below), make a hypothesis as to whether it is acidic, basic, or neutral. 2. Obtain a small amount of each sample listed on Table 1 and place each in an individual beaker. 3. Use your plastic pipette to gather some cabbage juice and squirt a few drops on each sample. 4. Observe the effects of the cabbage juice on the sample. Repeat using phenolphthalein as your indicator. 5. Now using the litmus paper, touch the sample with the litmus paper. 6. Observe the effects of the sample on the litmus paper. Determine if the sample is an acid, a base, or neutral. Table 1

SAMPLE baking soda water vinegar lemon juice lime juice soda pop hand soap dish soap tap water distilled water salt water vegetable oil

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Hypothesis (acid, base, neutral)

Plant Juice

Litmus Paper

Phenolphthalein

Acid, Base, Neutral?

Unit
7:
ACIDS
AND
BASES
 Johann Rudolf Glauber (1604-1670) noted that when a base is added to an acid, effervescence is observed. Glauber concluded that acids and bases perform in a “battle” until they have “slain” one another to produce neither an acid nor a base, but a salt.

Activity 2: Effervescence as an Indicator In this experiment, you will be testing various acids and alkalis using baking soda (sodium bicarbonate). Procedure 1. Obtain a small amount of each sample listed on Table 1 and place each in a small beaker 2. Use your spatula, scoop a small amount of baking soda into the beaker. 3. Continue to add small amounts of baking soda until the exact moment the baking soda solution stops bubbling. 4. Observe the effects of the baking soda on the sample. Record your observations in Table 1. 5. Determine if the sample is an acid. 6. Now using the litmus paper, touch the sample with the litmus paper. 7. Determine if the sample (after adding the baking soda) is an acidic, a basic, or neutral. Table 1 SAMPLE vinegar lemon juice soda pop hand soap dish soap tap water distilled water salt water vegetable oil

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Effervescence?

Acid?

Litmus Paper

Acid, Base, Neutral?

Unit
7:
ACIDS
AND
BASES
 Summary of Properties of Acids and Bases •

ACIDS 1. 2. 3. 4. 5.

•

Taste sour Reach with certain metals (Zn, Fe, etc.) to produce hydrogen gas cause certain organic dyes to change color react with carbonates to produce carbon dioxide React with bases to form salts and water

BASES 1. 2. 3. 4. 5.

Taste bitter feel slippery or soapy react with oils and grease cause certain organic dyes to change color react with acids to form salts and water

Problem: a. Write a complete balanced equation for the reaction of zinc with hydrochloric acid b. Write a complete balanced equation for the reaction of calcium carbonate with hydrochloric acid c. Write a complete balanced equation for the reaction of sodium hydroxide with sulfuric acid.

Check for Understanding:

1. Bases cause phenolphthalein to turn: d. Violet e. Orange f.

Colorless

g. Green

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2. Which of the following is a property of acids? a. Turn red litmus blue b. Feel slippery c. React with carbonates to from CO2 gas d. Taste bitter 3. Bases taste a. Bitter b. Sweet c. Sour d. Salty 4. Indicator

Results

phenolphthalein

colorless

Red litmus

Remains red

Blue litmus

Turns red

An unknown substance produced the experimental results noted above. Based on this data, the unknown is: a. A sugar b. Impossible to identify c. An acid d. A base

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5. When acids react with metals, the gas produced is a. Hydrogen b. Carbon dioxide c. Water vapor d. Oxygen

6. A positive test for a base occurs when: a. Red litmus remains red b. Blue litmus turns red c. Blue litmus remains blue d. Red litmus turns blue

7. Acids taste a. Sour b. Sweet c. Salty d. Bitter

8. A positive test for an acid occurs when: a. Red litmus remains red b. Red litmus turns blue c. Blue litmus remains blue d. Blue litmus turns red

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Unit
7:
ACIDS
AND
BASES
 Lesson 2:

In the late 1800s, the Swedish scientist Svante Arrhenius suggested that acids are compounds that contain hydrogen and can dissolve in water to release hydrogen ions into solution. For example, hydrochloric acid (HCl) dissolves in water as follows: HCl

H2 O

H+(aq)

+

Cl-(aq)

Arrhenius defined bases as substances that dissolve in water to release hydroxide ions (OH-) into solution. For example, a typical base according to the Arrhenius definition is sodium hydroxide (NaOH): H2O

NaOH

Na+(aq)

+

OH-(aq)

Check for Understanding: For each of the following compounds, show the dissociation of the substance into its ions and underline the ion which makes that substance acidic or basic according to the Arrhenius definition. • • • •

Potassium hydroxide Hydrobromic acid Magnesium hydroxide Nitric acid

Watch BrainPOP “Acids and Bases” and then take the quiz.

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Another way to determine if a substance is an acid or a base is by use of a special scale xdcalled the pH scale. The pH scale • • • • • •

•

measures how acidic or basic a substance is. ranges from 0 to 14. A pH of 7 is neutral. A pH less than 7 is acidic. A pH greater than 7 is basic. The pH scale is logarithmic and as a result, each whole pH value below 7 is ten times more acidic than the next higher value. For example, pH 4 is ten times more acidic than pH 5 and 100 times more acidic than pH 6. The same holds true for pH values above 7, each of which is ten times more basic than the next lower whole value. For example, pH 10 is ten times more basic than pH 9 and 100 times more basic than pH 8. The formula you can use to determine pH is

pH = -log [H+] where [H+] is the concentration of hydrogen ions

In the next activity, you will investigate how changing the concentration of an acid (and therefore the concentration of the H+ ion) affects the measured pH. Page
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Unit
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Activity 3: Investigating pH and [H+] Make sure you wear eye protection during this lab. a Number seven test-tubes 1–7. b Half-fill test-tube 1 with the hydrochloric acid solution. c Transfer 1 cm3 of the 1M hydrochloric acid into the measuring cylinder. Add distilled water to the measuring cylinder, up to the 10 cm3 mark. d Pour some of the resulting diluted solution from the measuring cylinder into test-tube 2, enough to come to a similar height as the solution in test-tube 1. e Carefully, pour away all but 1 cm3 of the solution remaining in the measuring cylinder. Now add distilled water to the measuring cylinder up to the 10 cm3 mark. Pour the resulting solution into test-tube 3. Continue in this way until you have solutions in testtubes 1 to 6. Put only distilled water into test-tube 7. f – j Repeat instructions a – e using the 1M sodium hydroxide solution instead of hydrochloric acid. Number the test-tubes 8–13. k Put the two racks of test-tubes together so that the solutions are in order 1 to 13. The test-tubes now have solutions in them with pH 1 (test-tube 1) to pH 13 (test-tube 13). l Add a drop of Universal indicator to each test-tube. Rock each test-tube from side to side to mix the contents. Add more Universal indicator solution to each test-tube if needed to allow the colors to be seen more clearly. Be sure to add the same number of drops of indicator to each test-tube. m Compare the colors of the solutions with the pH indicator chart.

Question: How closely do your solution colors match the colors on the indicator chart? If they are slightly different, can you suggest a reason?

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Unit
7:
ACIDS
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BASES
 Another useful formula in acid base chemistry is

pH + pOH = 14

pOH is determined the same way as pH except [OH-] is used instead of [H+] ___________________________________________________________________________ Worked problem 1: What is the pH of a 3.75 x 10-5 M solution of nitric acid? Step 1: Write the equation for the dissociation of HNO3. HNO3 (aq) → H+ (aq) + NO3- (aq) Step 2: Determine the concentration of H+ Since the [HNO3] is 3.75 x 10-5 M and each HNO3 produces one H+ ion, the concentration of H+ is 3.75 x 10-5 M Step 3: Calculate the pH pH = - log [H+] = - log (3.75 x 10-5) = 4.43 ______________________________________________________________________________ Worked problem 2: What is the pH of a 1.0 x 10-4 M solution of potassium hydroxide? Step 1: Write the equation for the dissociation of potassium hydroxide KOH(aq) →

K+ (aq) + OH- (aq)

Step 2: Determine the concentration of OHSince the [KOH] is 1.0 x 10-4 M and each KOH produces one OH- ion, the concentration of OH- is 1.0 x 10-4M Step 3: Calculate pOH then the pH pOH = -log [OH-] = - log (1.0 x 10-4) Page
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Unit
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BASES
 =4 To calculate pH, use pH + pOH = 14 so pH = 10 Checking for Understanding 1. For each of the following pairs of solutions, determine which is the more acidic. a. Blood of pH = 7.4 and brass polish of pH=9.5 b. Black coffee of pH = 5.0 and vinegar of pH = 2.8

2. For each of the following pairs of solutions, determine which is more alkaline a. Toothpaste of pH = 8.0 and milk of magnesia of pH = 10.5 b. Orange juice of pH = 3.5 and lemon juice of pH = 2.3

3. If the [H+] of an acidic solution is 1.0 x 10-5 mol/L, state the pH of the solution.

4. Solution A is 1000 times more acidic than solution B. If solution A has a pH=2, state the pH of solution B.

5. Determine the pH of solutions of the following concentrations: a.[NaOH] = 3.34 x 10-5 M b. [Mg(OH)2] = 2.50 x 10-4 M

6. 25.0 ml of 8.00M nitric acid is diluted to 200ml with distilled water. Calculate the pH of the resultant solution.

7. 3.73 g of potassium hydroxide are dissolved in 250 ml of water. Calculate the pH of the resultant solution.

8. Calculate the pH of a solution formed by bubbling 0.225 mol of HCl gas through 1.5 litres of water.

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Unit
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Lesson 3: Neutralization: Acids release H+ into solution and bases release OH-. If we were to mix an acid and base together, the H+ ion would combine with the OH- ion to make the molecule H2O, or plain water: H+(aq) +

OH-(aq)

H2 O

The neutralisation reaction of an acid with a base will always produce water and a salt, as shown below: Acid

Base

Water

Salt

HCl

+

NaOH

H2O

+

NaCl

HBr

+

KOH

H2O

+

KBr

Check for Understanding: Write balanced equations for the following neutralization reactions: 1. 2. 3. 4.

Hydrochloric acid with potassium hydroxide Nitric acid with sodium hydroxide Sulfuric acid with magnesium hydroxide Hydrochloric acid with aluminum hydroxide

Though Arrhenius helped explain the fundamentals of acid/base chemistry, his theory was limited because it could not explain why some substances, such as common baking soda (NaHCO3), can act like a base even though they do not contain hydroxide ions. In 1923, the Danish scientist Johannes Brønsted and the Englishman Thomas Lowry published independent yet similar papers that refined Arrhenius' theory.

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Unit
7:
ACIDS
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BASES
 In Brønsted's words, "... acids and bases are substances that are capable of splitting off or taking up hydrogen ions, respectively." The Brønsted-Lowry definition broadened the Arrhenius concept of acids and bases. The Brønsted-Lowry definition of acids is very similar to the Arrhenius definition- any substance that can donate a hydrogen ion is an acid The Brønsted definition of bases is, however, quite different from the Arrhenius definition. The Brønsted base is defined as any substance that can accept a hydrogen ion. In essence, a base is the opposite of an acid. NaOH and KOH would be considered bases because they can accept an H+ from an acid to form water. However, the Brønsted-Lowry definition also explains why substances that do not contain OH- can act like bases. Baking soda (NaHCO3), for example, acts like a base by accepting a hydrogen ion from an acid as illustrated below: Acid HCl

Base +

NaHCO3

Salt H2CO3

+

NaCl

Watch BrainPOP “pH scale ” and then take the quiz.

Checking for Understanding: 1.

For the following reactions, identify which reactant is acting as a Bronsted-Lowry acid and which is acting as a Bronsted-Lowry base. a. HCO3- + Clb. HSO4- + NH3 c. HCl + H20

CO32- + HCl SO42- + NH4+ H30+ + Cl-

Real world application: During autumn, the colors of leaves change. Leaves may change from green to bright red, orange and yellow. The short cool days of autumn bring an end to the production of chlorophyll – the green pigment in leaves. As chlorophyll breaks down, the colors of the more stable carotenoid ( yellow/orange) and anthocyanin ( red/blue/purple) pigments become visible. The color of anthocyanins depends on acidity. Anthocyanins are water Page
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Unit
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BASES
 soluble and are dissolved in the cell sap rather than bound to the membranes as chlorophyll is. If the cell sap is quite acidic, the anthocyanins impart a vivid red color but if it is less acidic the color may appear purple.

Lesson 4: Acid Base titrations

An acid base titration is a method used in chemistry that relies on the neutralization reactions of acids and bases. We know that indicators can have one color in an acid but a different color in a basic solution. Phenolphthalein, you will recall, is colorless in an acid but pink in a base. So if you have a basic solution and add some phenolphthalein the solution will turn pink. If you now add acid slowly to that solution, the color will remain pink for some time until you reach the point where the acid exactly neutralizes the entire base present. What do you think will happen if you now add one more drop of acid? The solution will turn colorless. This gives us an indication of neutralization. We can make use of this process to track changes in pH as an acid is added to a base. We can also use the process of titration to determine the concentration of an acid ( or a base).

The following applet allows you to practice (on line) the process of titrating. http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/s toichiometry/acid_base.html Page
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Unit
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Worked titration problem: Titration reveals that 11.6 ml of 3.0 M sulfuric acid are required to neutralize the sodium hydroxide in 25.00 ml of NaOH solution. What is the molarity of the NaOH solution? Two different solution methods are shown below: 1. Using dimensional analysis H2SO4 + 2NaOH → 2H20 + Na2SO4

2. An alternate method H2SO4 + 2NaOH → 2H20 + Na2SO4 Calculate the amount ( in mol) of the known substance. n(H2SO4 ) = cV = 3.0 M x 11.6 ml/1000ml = 3.48 x 10-2 mol of H2SO4 Determine the mole ratio of unknown to known from the equation n(NaOH) = 2 x n(H2SO4 ) = 2 x 3.48 x 10-2 mol = 6.96 x 10-2 mol of NaOH

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Unit
7:
ACIDS
AND
BASES
 Determine the concentration of the unknown C = n/V = 6.96 x 10-2 mol of NaOH/ 25.00ml/1000ml = 2.784 ≈ 2.8 M NaOH

Checking for Understanding: 1. If 14.7 ml of 0.102 M NaOH is required to titrate 25.00 ml of HCl solution, what is the molarity of the hydrochloric acid solution? ( Answer: 0.0579 M HCl) 2. If 19.1 ml of 0.118 M HCl is required to neutralize 25.00 ml of a sodium hydroxide solution, what is the molarity of the sodium hydroxide? ( Answer: 9.02 x 10-2 M NaOH) 3. If 7.3 ml of 1.25 M HNO3 is required to neutralize 25.00 ml of a potassium hydroxide solution, what is the molarity of the potassium hydroxide? ( Answer: 0.044 M KOH) 4. If 12.0 ml of 1.34 M NaOH is required to neutralize 25.00 ml of a sulfuric acid solution, what is the molarity of the sulfuric acid? ( Answer: 0.322 M H2SO4) 5. A solution containing ammonia requires 18.0 ml of 0.100 M HCl to reach equivalence. Calculate the amount ( moles) of ammonia that reacted with the HCl. If the volume of the ammonia solution was 2.0 L, what is the concentration of the ammonia in the solution? 6. How many ml of 0.100 M HCl are required to neutralize 25.00 ml of Ba(OH)2? 7. What is the molarity of a hydrochloric acid solution 30.0 ml of which is just neutralized by 48.0 ml of 0.100 M NaOH?

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Lesson 5: Titration procedure In today’s class you will get a chance to do a titration yourself. The technique is simple but it must be performed precisely. The purpose of this titration is to determine the unknown concentration of a sodium hydroxide solution.

Titration: quantitative analysis of an acid/base reaction 1. Obtain the following equipment: a 10 mL pipette, pipette filler, 50 mL burette, 125 mL Erlenmeyer flask, and two 50 mL beakers. Be careful when handling the burettes – they are very expensive. 2. Rinse all equipment with tap water. For the burette: a) close the stopcock, b) pour water into the end from a beaker, c) open the stopcock for a few seconds to rinse out the tip, d) rotate the burette as you dump out the remaining water (this rinses the sides of the burette), e) open and then close the stopcock to allow water in the tip to escape. For the pipette: a) use a pipette filler to fill the pipette part way with water, b) remove the filler, c) rotate the pipette as you dump the water, d) dry the end. 3. Dry the 50 mL beakers. Label one “base”, and the other “acid”. Fill the acid beaker with 0.175 M HCl. Half fill the base beaker with “unknown” NaOH. To conserve chemicals, take only what you need. 4. Rinse the burette (see above) with a small amount of acid. Remember to rinse, & then drain, the tip. 5. Get a burette clamp from the front of the room. Using the clamp, attach the burette to your retort stand. 6. Fill the burette with acid to the 0 mL mark. Ensure that the tip of the burette is filled with acid – not air (to get rid of air bubbles, run some acid through the burette) 7. Rinse the pipette with base by taking up NaOH to the 0 mL mark and then emptying it into the sink.

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Unit
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 8. Using a pipette, add 10 mL of base to the flask (i.e. fill pipette to 0 mL mark, then drain to 10 mL). 9. Get a dropper bottle of phenolphthalein. Add three drops of phenolphthalein to the NaOH in the flask.

10. Add about 10 mL of distilled water to the flask (no need to measure it - just estimate). The extra water will not influence the reaction; it will simply increase the total volume, making the reaction easier to see. 11. With the tip of the burette in the neck of the flask, add acid – about 1 mL every 3-5 seconds. Swirl the contents of the flask constantly. To see the reaction better you can place a sheet of paper under the flask. Stop the acid flow immediately when you see the colour change (called the “endpoint”). This will happen very quickly so be careful not to “over-titrate”. Note the volume of acid used. 12. You will now repeat step 11; to get a more accurate reading, you will be adding acid slower near the endpoint. Rinse out the flask. Add 10 mL of NaOH (using the pipette) & 3 drops phenolphthalein to the flask. Fill the burette back to the 0 mL mark with acid. Drain acid into the flask (from the burette) until the volume reads approximately 3 mL less than what was used in step 11. Now, add acid slowly (1 drop every 2-3 seconds) – stop when the endpoint is reached. Record the volume below. 13.Repeat the titration until you have three values for the volume of acid added that are very close . 14. Clean up. Rinse all equipment well with tap water (including burette & pipette). For burettes, remember to rinse and drain the tip. Return the burettes with the stopcock in the “open” position. 15. Use your data to determine the concentration of the NaOH solution. Data Table: Table 1: Volumes of sodium hydroxide and hydrochloric acids used in an acid base titration Rough trial Initial Volume of HCl Final volume of HCl

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1st Trial

2nd Trial

3rd Trial

4th Trial

Unit
7:
ACIDS
AND
BASES
 Volume of HCl added Volume of NaOH

10.0 ml

10.0 ml

10.0 ml

10.0 ml

10.0 ml

w Lesson 6 and 7 : Designing a lab Your task is to design a lab procedure that will enable you to determine the concentration of acetic acid in a sample of vinegar. The equation for the reaction of acetic acid with sodium hydroxide is shown here:

CH3COOH(aq) + NaOH(aq) → H2O(l) + NaCH3COO(aq) For the lab, you will be provided with : •

A vinegar solution

•

The sodium hydroxide you used in your last experiment ( you now know its concentration).

•

Phenolphthalein indicator

•

Distilled water

•

Burettes

•

10 ml pipettes and pipette fillers

Include a complete data table as part of your design.

Once your design is approved by your teacher, you may proceed with your titration.

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Unit
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