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 7: ACIDS AND BASES
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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Unit 7: ACIDS AND BASES
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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Unit 7: ACIDS AND BASES
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:
H2O
HCl
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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Unit 7: ACIDS AND BASES
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.
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Unit 7: ACIDS AND BASES
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?
Another useful formula in acid base chemistry is
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pH + pOH = 14
Unit 7: ACIDS AND BASES
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)
=4
To calculate pH, use pH + pOH = 14 so pH = 10
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Unit 7: ACIDS AND BASES
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 7: ACIDS AND BASES
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)
H2O
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.
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,
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Unit 7: ACIDS AND BASES
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
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.
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Unit 7: ACIDS AND BASES
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
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Unit 7: ACIDS AND BASES
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
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
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Unit 7: ACIDS AND BASES
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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Unit 7: ACIDS AND BASES
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.
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.
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Unit 7: ACIDS AND BASES
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
1st Trial
2nd Trial
3rd Trial
4th Trial
10.0 ml
10.0 ml
10.0 ml
10.0 ml
10.0 ml
Initial
Volume of
HCl
Final volume
of HCl
Volume of
HCl added
Volume of
NaOH
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Unit 7: ACIDS AND BASES
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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