Acid and Base packet Page 1 of 13

Honors Chemistry Acids and Bases

Quiz: Test: Project:

Vocabulary:

1Alkaline 6 diprotic acid 11 neutralization 16 salt 21 triprotic acid

2 amphoteric 7 end point 12 oxyacid 17 standard solution

3 (Arrhenius) acid 8 Hydronium 13 pH 18 Strong acid 22Weak acid

4 (Arrhenius) base 9 Hydroxide 14 Poh 19 strong base

5 Binary acid 10 monoprotic acid 15 polyprotic acid 20 Titration 23 weak base

COMMON ACIDS Memorize

Hydrochloric Acid HCl Nitric Acid HNO3 Acetic Acid HC2H3O2 Sulfuric Acid H2SO4

ACIDS, BASES & SALTS

NAMING ACIDS

Binary vs oxyacids

COMMON STRONG ACIDS (MEMORIZE):

HCl ______________ HNO3 ______________ HClO4 ______________ ________ sulfuric acid

COMMON WEAK ACIDS (MEMORIZE)

_______Hydrofluoric _________ acetic acid H2CO3 _____________________ H3PO4 _______________

HCl and HNO3 are monoprotic meaning ______________________________

Sulfuric acid is ______________________ , 2 Protons (also called polyprotic)

Phosphoric acid is ____________________, 3 protons. (also called polyprotic)

A binary acid contains only 2 elements (H + other).

Examples:

They are named using the prefix hydro- the root of the element and –ic.

Name: HF, HCl, HBr and HI

Oxyacids are composed of H, O and a 3rd element an example is____________

Common acids:

Sulfuric: # 1 industrial chemical, used in making ___________________

Effective dehydrator and is found in acid rain and automobile batteries

Phosphoric also in ______________ Nitric – used to make_______________

Hydrochloric: common lab acid, found in your ________________

Acetic: found in __________________

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PROPERTIES OF ACIDS: 1. Sour taste ex’s: citric acid in lemons, oranges and grapefruit, maleic acid in apples, acetic acid

in vinegar and lactic acid in sour milk.

2. Contain hydrogen, some react with metals to release HYDROGEN gas.

Write the balanced equation with state symbols for the reaction of zinc metal with hydrochloric

acid (HCl):

3. Acids change the color of dyes know as indicators. Example of an indicator: pool pH test strips

* litmus= turns pink for acid, blue for base *phenolphthalein= turns pink with bases

4. Acids react with bases to produce : a salt and water = NEUTRALIZATION reaction

5. Acids are electrolytes, - they will conduct electricity as a solution.

Definitions for acid:

Arrhenius (traditional definition): cmpd containing hydrogens and ionizes in aqueous soln to form

hydronium ions:

Example: nitric acid HNO3 + H2O H3O+ + NO3-

hydronium ion (H+ or H3O+)

Strong acids are strong electrolytes – they ionize completely in aqueous soln (completely

breakdown). Remember electrolytes – conduct a current in solution.

Weak acids are weak electrolytes – they ionize partially in aqueous soln (in equilibrium – partially

break down). Remember electrolytes – conduct a current in solution. We can write a Ka for their

dissociation

PROPERTIES OF BASES: 1. Have a bitter taste, ex: soap

2. Feel slippery

3. Change color of indicators

4. React with acids to make salts and water

5. Are electrolytes, conduct a current as a solution

Traditional definition (Arrhenius def) of a base: Contains hydroxide: ______________

In a neutralization reaction an acid and a base react to form a salt and water. Neutralization

reactions are often double displacement reactions.

Practice neutralization reactions: name the acid and circle the salt in the final answer!

1. HCl + NaOH __________ + ___________ acid name: ___________________________

2. H2SO4 + KOH __________ + ___________ acid name: ___________________________

3. HBr + Mg(OH)2 __________ + ___________ acid name: ___________________________

4. HC2H3O2 + NaOH __________ + ___________ acid name: ___________________________

5. Ca(OH)2 and H3PO4 ____________ + _____________ acid name: ___________________________

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Types of bases:

Strong: hydroxides of metals in I and II Ex: __________, ___________ and __________. Strong electrolytes

Weak: hydroxides of many transition metals ex Fe(OH)3 and organic compounds containing nitrogen

Example: ammonia= NH3 .

A solution that contains OH- from a soluble base is referred to as:______________

___anhydrous__ means without water, so __anhydrides___ are acids or bases that have had water

removed.

Metallic oxides generally form bases and nonmetallic oxides generally form acids.

See ref packet pg 6

Water is amphoteric – it can act as an acid or a base

Buffer: solution that can resist changes in pH, usually made up of a weak acid or base and a salt of

the weak acid or base. Examples: blood, sea water

The pH Scale

The pH scale ranges from 0 to 14. It measures the acidity or basicity of a solution. A pH of 7 means it is

a neutral solution. Pure water has a pH of 7. A pH of less than 7 means the solution is acidic. A pH of

more than 7 means the solution is basic. The less pH, the more acidic the solution is. The more pH, the

more basic the solution is.

pH stands for the power of H, or the amount of H+ ions acids or bases take or contribute in solution. pH

equals the negative log of the concentration of H+. pH = -log[H+]

Acid and Base packet Page 4 of 13

WHAT REALLY HAPPENS

Acids are compounds that break into hydrogen (H+) ions and another compound when placed in

an aqueous solution. Bases are compounds that break up into hydroxide (OH-) ions and another

compound when placed in an aqueous solution.

Let's change the wording a bit. If you have an ionic compound and you put it in water, it will break

apart into two ions. If one of those ions is H+, the solution is acidic. If one of the ions is OH-, the solution

is basic. There are other ions that make acidic and basic solutions, but we won't be talking about

them here.

That pH scale we talked about is actually a measure of the number of H+ ions in a solution. If there

are a lot of H+ ions, the pH is very low. If there are a lot of OH- ions, that means the number of H+ ions

is very low, so the pH is high.

Notes on pH Aqueous solution:

Water is self-ionized (slightly) breaks down partially on its own.

Example: 2 H2O H3O+ + OH-

Reversible reaction – reaches equilibrium – some of the water breaks down into hydronium H3O+ (or

H+ ) and hydroxide OH –

NOTE: Hydronium can be written as either H3O+ or its abbreviated form H+

Pure water has: 1 x 10 -7 M H+ and 1 x 10 -7 M OH-

Because [H+] = [OH-] in pure water it is neutral (remember [ ] indicates molar concentration)

if [H+] > [OH-] = acidic if [OH-] > [H+] = basic

In all aqueous solutions: [H+] [OH-] = 1 x 10-14 (at 25 0C)

Example: Calculate [H3O+] and [OH-] for 1 x10-4 M nitric acid.

Example 2: Calculate [H3O+] and [OH-] for 0.02 M Ba(OH)2

You try: Calculate [H+] and [OH-] for 1 x 10-3 M H2SO4

Acid and Base packet Page 5 of 13

Mathematically, it is possible to have a pH below 0 or above 14. We will not go outside of that range

All six equations: in reference packet!

pH = -log [H+] pOH = -log [OH-]

[H+] = 10 –pH [OH-] = 10 –pOH

[H+] [OH-] = 1 x10 -14 pH + pOH = 14

logarithm = the power to which 10 must be raised to get that number.

Example: log 107 = 7 log 10-3 = -3 etc.

You can calculate the log by inspection only when you have a perfect base 10 – otherwise use the

log button on your calculator

Example: What is the pH of 0.001 M NaOH? Copy work

Example 2: what is the pH of 3.4 x 10 -3 M H2SO4 ?

You try: What is the pH of 1.2 x 10-3 M Mg(OH)2?

To find concentration from pH use [H+] = 10 –pH (in packet)

Use the 10x key – second function above the log button DO NOT use the exp key!!!

Example: what is the [H+] of a solution with a pH of 7.52?

You try : what is the hydroxide concentration of a solution with a pH of 3.2?

pOH scale opposite of the pH scale – on the pOH scale bases have low numbers and acids have

high numbers!

Example: What is the [H+] of a solution with a pOH of 8.2?

You try: What is the pOH of a solution with an [H+] =6.5 x 10 -5?

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Acid Rain Reading Questions

1. At what pH does acidic rainfall

become a concern?

2. How are plants harmed by acid

rain (according to the author)?

3. Why would non-polluted rural

areas be concerned about acid rain?

4. What acid anhydride normally exists

in non-polluted air making rain acidic?

NAMES TO KNOW

Here are a couple of definitions you should know:

Acid: A solution that has an excess of H+ ions. It comes from the Latin word acidus that means

"sharp".

Base: A solution that has an excess of OH- ions. Another word for base is alkali.

Aqueous: A solution that is mainly water. Think about the word aquarium. AQUA means water.

Strong Acid: An acid that has a very low pH (0-4).

Strong Base: A base that has a very high pH (10-14).

Weak Acid: An acid that only partially ionizes in an aqueous solution. That means not every molecule

breaks apart. They usually have a pH close to 7 (3-6).

Weak Base: A base that only partially ionizes in an aqueous solution. That means not every molecule

breaks apart. They usually have a pH close to 7 (8-10).

Neutral: A solution that has a pH of 7. It is neither acidic nor basic.

Acid and Base packet Page 7 of 13

Titration notes

In a titration you have two solutions (typically one is an acid and one is a base). You know the

concentration of one solution (the standard) and you are trying to calculate the concentration of the

unknown. This is done by using an indicator and slowly mixing the two solutions together until you get

a color change (which indicates the end point!). When the color changes you use your two volumes

(the volume of the standard and the volume of the unknown) along with the concentration of the

known to calculate the unknown’s concentration. We will do math calculations in class and then in

our formal lab we will complete an actual titration.

End point (equivalence point): Rapid change in pH, there are equivalent amounts of H+ and OH, the

equivalence point does not always occur at a pH of 7 – see graphs below. (depends on whether you

have a strong or weak acid/base)/

Indicators & the pH scale: the best indicator in a titration is based on the pH at the equivalence pt:

Examples of titration curves! Graphs of data from titration experiments.

Which indicator should we use for curve A _____________ curve B ? _____________

Which indicator would be pink when the [H+] = 1 x 10 -4 ?

You try which indicator would be blue when [H+] = 3.2 x10 -12 ?

We will use a buret in our titrations – they measure the volume released from the buret – they are

read from the top down (zero is at the top and 50.00 ml is at the bottom):

Ex: ___________ you try: ______________

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Working titration problems

you can use the formula: MaVa na = MbVb nb

M = molarity (of acid or base) V = volume of acid or base and

na = moles of H+ in the acid and nb = moles of OH- in the base.

Example: In a titration, 27.4 ml of a standard solution of Ba(OH)2 is added to 20.0 ml of HCl. The

concentration of the Ba(OH)2 is 0.0154 M. What is the Molarity of HCl?

You try: it takes 25 ml of 0.20 M H2SO4 to neutralize 37 ml of NaOH – what is the molarity of the NaOH?

EXTRA PRACTICE: Calculate the pH of the solutions below

1. 0.01 M HCl 4. 0.0030 M H2SO4

2. 0.0010 M NaOH

3. 0.050 M Ca(OH)2 5. 0.150 M KOH

6. What is the [H3O+] of a solution with a pH of 6.7? (remember [H3O+] = [H+])

7. What is the [OH-] of a solution with a pH of 9.5?

8. What is the pH of a solution with a pOH of 12.5?

9. What is the [H+] of a solution with a pOH of 3.5?

10. What is the pH of a solution with a [OH-] of 2.3 x 10-6? Is this an acid or a base?

11. What is the pOH of a solution with a [H+] of 1.7 x 10-6 ?

12. What is the [OH-] of a solution with a pH of 6.3?

13. What is the pH of a 0.012 M solution of Be(OH)2?

14. What is the [H+], [OH-], pH and pOH of 0.22 M H2SO4 ?

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Titration practice sheet To determine the concentration of an acid (or base), we can react it with a base (or acid of known

concentration until it is completely neutralized. The point of complete neutralization is the endpoint –

noted by the color change of the indicator.

1. Why do you need to use an indicator when titrating a solution?

2. If 20.0 ml of 0.0100 M HCl is required to neutralize 30.0 ml of NaOH determine the molarity of the

NaOH solution.

3. Suppose that 20.0 ml of 0.10 M Ca(OH)2 is required to neutralize 12.0 ml of HCl – what is the

molarity of the acid?

4. How much 2.2 M Ba(OH)2 is required to neutralized 15.0 ml of 1.75 M H2SO4?

5. What is the concentration of 50 ml of sulfuric acid that requires 25 ml of 2.0 M potassium hydroxide

to be completely neutralized?

6. What volume of 1.25 M Ca(OH)2 is required to titrate 36 ml of 1.77 M HCl?

7. Given the reaction: Ba(OH)2 + 2HClO4 Ba(ClO4) 2 + 2HOH

95.5 mL sample of Ba(OH)2 is neutralized to completion with 75.0 mL of 0.8 M HClO4. Determine the

concentration of Ba(OH)2.

8. A 10.0 mL sample of H2SO4 was exactly neutralized by 13.5 mL of 1.0 M KOH what is the molartiy of

the sulfuric acid?

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Review on Acids, Bases, pH, and Titrations

Find the pH of the following solutions.

1. a 0.08 M HCl solution

2. a 0.11 M Mg(OH)2

3. a solution with a pOH of 8.3

4. a solution with [H3O+] = 5.6 x 10-5

5.What is the pOH of a solution with a [H+] = 2.3 x 10-8?

6.What is the OH- concentration of a solution with a pH = 3.66?

7.What is the H+ concentration of a solution with a hydroxide ion concentration of 4.7 x 10-3M?

8. What is the [OH-], [H3O+], pH and pOH of a 0.050 M H2SO4?

9.What is the molarity of 1.5L of a hydrochloric acid solution that requires 790 ml of 2.8 M sodium

hydroxide to be titrated?

10. If it requires 90 ml of 1.0 M sulfuric acid to titrate 150 ml of potassium hydroxide, what is the

concentration of the KOH?

11. What is the molarity of 1.75 L of HNO3 titrated with 560 ml of 1.28 M LiOH?

12. How many liters of 0.5 M RbOH are required to completely titrate 750 ml of 0.1 M HCl?

13. List the 5 properties of acids

14. List the 5 properties of bases

15. Write the formula or name for A. HCl, B. H2SO4, C. HNO3, D. HC2H3O2, E. phosphoric acid,

F. perchloric acid, G. Hydrofluoric acid and H. carbonic acid

16. Which of the above are diprotic acids?

17. Which of the above are oxyacids?

18. Which of the above are binary acids?

19. Which of the above is found in your stomach?

20. Which of the above is used to make explosives?

21. Which of the above is found in vinegar?

22. Which of the following is a salt: A. KOH B. KCL C. HOH D. HCL ?

Answers 1) 1.1 2) 13.3 3) 5.7 4) 4.25 5) 6.36 6) 4.6 x 10-11 7) 2.13 x 10-12

8) 1 x 10-13, 0.1, 1, 13 9) 1.5 M HCl 10) 1.2 M KOH 11) .41 M HNO3 12) .15L RbOH

13. Sour taste, react with metals to produce hydrogen gas, changes the colors of indicators, reacts

with bases to form salts and water (neutralization), are electrolytes.

14.Bitter taste, feel slippery, changes the colors of indicators, reacts with acids to form salts and

water (neutralization), are electrolytes.

15 A. Hydrochloric B. Sulfuric C nitric D acetic E. H3PO4 F. HClO4 G. HF

H H2CO3

16 B and H 17. B, C, D, E, F, H 18 A and G

19 A 20. C 21. D 22 B – ionic copmpound

Acid and Base packet Page 11 of 13

Acid-Base Titration Lab Name:___________________ pd:___________

Pre-lab Questions- Use notes on pg 7 to answer

1. What is a titration?

2.What is an equivalence point?

3.Why are titrations done?

Objectives:

To complete an acid- base titration To practice using a buret

To use lab data to calculate the concentration of an unknown solution.

Safety: Full MSDS sheets are in the MSDS folder in the lab room. MSDS health information:

HC2H3O2, NaOH and phenolphthalein

Health Effects: may cause skin irritation, eye irritation, gastrointestinal irritation

First Aid: eyes: flush with water for 15 minutes.

Skin – flush with soap and water for 15 minutes

Ingestion: give 2-4 cups of milk or water – get medical attention

Equipment:

Beakers, Flask, pipet, buret, buret clamp, ring stand, phenopthalein, 0.250 M NaOH, ? M HC2H3O2

Procedure:

1. Use small graduated cylinder & place 10.0 ml of the acid - HC2H3O2 into a clean Erlenmeyer flask.

2. Add 2-3 drops of phenolphthaleinindicator to the acid solution. Set to the side.

3. Record the initial buret reading.

4. Place the flask under the buret (place a piece of clean white paper under the flask to make it

easier to observe).

5. Begin titrating by adding small amounts of base to the flask (slightly turn the stopcock to allow

small amounts of base out – be sure to add the base slowly!)

Horizontal = CLOSED Vertical = FULLY OPEN 6. Gently swirl the solution after each addition.

7. When localized pink spots appear and then disappear, reduce the size of additions to dropwise.

Ms. Clark will show you how to add a “half” drop.

8. When one drop of base produces a faint pink color that persists with swirling the endpoint has

been reached. Record the final buret reading.

9. Repeat the titration with a second sample of acid – redo steps 3-5; your buret does not need to

be refilled. Just record the initial buret reading before starting the second titration process.

10. Clean up excess acid and base may go down the drain with plenty of water. If you are not the

last class leave the burets set up for the next class!

11. 6th pd drain burets = Ms.Clark come around with vinegar water to rinse burets out.

NOTE: You do NOT want your sample to look HOT pink. You want a

pale pink color to remain. You get two shots to get it right!

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Data Table: include units and use correct significant digits

Trial One Trial Two Trial three

1. M of NaOH (Mb) 0.250 M 0.250 M 0.250 M

2. Volume of HC2H3O2 (Va) 10.0 ml 10.0 ml 10.0 ml

3 Intial buret reading of NaOH _______ _______ _______

4. final buret reading _______ _______ _______

fill in this part of the data table after the lab! Show work below in question section!

5. volume of NaOH used in titration _______ ________ ________

(subtract #4 - # 3) (Vb)

6. Calculated M of HC2H3O2 (Ma) ______________ _____________ _____________

7. Average Ma of HC2H3O2

(avg of three answers in #6)

Calculations: Show all work – credit will not be given without supporting work. Include units!

8. Calculate the volume of base used in each trial. (Show how you calculated answer for #5)

9. Write a balanced equation for the reaction between the acid and base. (water IS NOT a

reactant –use the state symbol aq to indicate water has been added) Use state symbols!

10. Calculate the M of HC2H3O2 for each trial – (when using your na , use na = 1 because HC2H3O2

only loses the first H in a titration.) (Show how you calculated answer for #6)

11. Calculate the average M of HC2H3O2 (Show how you calculated answer for #7)

12. What is the concentration of 75.0 ml of NaOH titrated by 100 ml of 1.2 M hydrochloric acid?

13. How much 0.68 M carbonic acid would be required to titrate 250 ml of 0.2 M lithium hydroxide?

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Questions:

14. Why did Ms. Clark rinse the buret out with a 5 ml portion of the base before beginning the titration.

15. Why would your sample turn bright pink if you added too much base?

16. What would happen if you forgot to add the phenopthalein to your sample?

Conclusion:

17. In a titration you have two solutions (typically one is an acid and one is a base). You know the

concentration of one solution called the _____________________ and you are trying to calculate the

concentration of the unknown. This is done by using an indicator and slowly mixing the two solutions

together until you get a color change (which indicates the end point!). My favorite part of this lab

was_____________________________________________________________________________________________

__________________________________________________________________________________________________

____________________________________________________________________________.