CHEM 30 acid-base note package

CHEM 30

In simple terms Arrhenius acid-base theory can be stated as follows: An acid is a chemical entity that dissociates in water to release hydrogen ions, H+_{(aq) - a strong acid dissociates 100} % at all concentrations while a weak acid tends to dissociate partially at most concentrations.

A base is a chemical entity that dissociates in water to release hydroxide ions, OH-_{(aq) - a} strong base dissociates 100 % at all concentrations (even if it is low solubility). (Weak bases, which do not produce hydroxide ions by dissociation, are not considered by this theory.)

Strong acids: X-(aq) = ClO4-(aq), X= l-(aq) , X=Br-(aq) , X= Cl-(aq) , X= HSO4-(aq) , X= NO3-(aq)

HX(aq) → H+_{(aq) + X}-_{(aq) we assume 100% ionization and dissociation}

Weak acids: A-_{(aq) = CH3COO}-_{(aq), X= CN}-_{(aq) , X=F}-_{(aq) , X= see your data table!}

HA(aq) ⇋ H+_{(aq) + A}-_{(aq) Weak acids do not dissociate 100% So finding the [H}+_{] is harder.}

Strong bases, M(OH)n(aq):Aqueous Group I metal hydroxides, Ba(OH)2(aq) and Sr(OH)2(aq).

M(OH)n(aq) → Mn+_{(aq) + nOH}-_{(aq) we assume hydroxide bases react 100% at this level.}

The Acid Base Table contains the names, formulae, conjugate bases and Ka values of several common acids ranked in order from strongest to weakest. (Acids get stronger as you go up on the right side while the bases get stronger going down the left side).

This is theory was created to help figure out why some solutions that do not contain hydrogen show acidic properties and the same thing for no hydronium containing bases.

The theory was that water was a reactant and it produced the H30+(aq) and OH-(aq) in

certain solutions.

1. Write an equation using the Arrhenius modified acid base definition that helps explain the following experimental observations:

a.* Aqueous solutions of sodium bicarbonate are basic.

b. Aqueous solutions of sulfur dioxide are acidic.

c. Adding lithium oxide to water creates an extremely basic solution.

d.* Aqueous solutions of potassium hydrogen sulfite are acidic.

e. Aqueous solutions of carbon dioxide are acidic.

f. Aqueous solutions of cesium cyanide are basic.

g. Aqueous solutions of sodium carbonate are basic.

h. Aqueous solutions of sulfur trioxide are acidic.

i. Adding potassium phosphate to water creates a basic solution.

j.* Aqueous solutions of sodium hydrogen sulfate are acidic.

In simple terms Brønsted-Lowry acid-base theory (a reaction theory) can be stated as follows:

An acid is a chemical entity that donates/loses a hydrogen ion, H+ _{(aq) ,in a neutralization reaction (acids} are proton donators H+ _{(aq) is a proton).}

A base is a chemical entity that removes/accepts a hydrogen ion, H+ _{(aq) ,in a neutralization reaction} (bases are proton acceptors H+ _{(aq) is a proton).}

A neutralization reaction is the complete or partial transfer of a hydrogen ion, H+, between the strongest acid (SA) and the strongest base (SB) in a reaction mixture.

In general strong acids/bases tend to neutralize other bases/acids completely/100 %/quantitatively. Weak acids/bases are only capable of neutralizing other weak bases/acids partially/<100 %. If the acid lies above the base in an acid base table the products are favoured (>50%). If the acid lies below the base, reactants are favoured (<50%).

A. an acid acts as a proton donator :

HCl(aq) + H20(l)  H30+(aq) + Cl-(aq) Acid Base

A base acts as a proton acceptor :

NH3(aq) + H20(l) ⇋ OH-(aq) + NH4+(aq) Base Acid

Water therefore can act as an acid and a base [ AMPHOTERIC SUBSTANCE ]

B All acid and base reactions are reversible ( equilibrium reaction )

From this idea, came the concept of CONJUGATE ACID (CA) & CONJUGATE BASE (CB) and the idea or CONJUGATE PAIRS in an acid – base reaction

1. The reversible ( equilibrium ) reaction.

A. An equilibrium expression could be written about any reversible reaction

w A + x B ⇋ y C + z D

B. The Keq value will describe the extent of the reaction ( the strength of the acid or base ) i. if the Keq value is approximately = 1 , the [ products ] = [ reactants ]

the reaction could be describe as a 50% reaction , A & B is about as strong as CA & CB

ii. if the Keq value is greater than 1 , the [ products ] > [ reactants ]

this reaction reacts to more than 50% , the acid & base is stronger than CA & CB

iii. if the Keq value is less than 1 , the [ products ] < [ reactants ]

this reaction reacts to less than 50% , the acid & base is weaker than CA & CB

2. The amphoteric substance , water ( H2O ), can be both an acid and a base.

a. Since water is an amphoteric substance, water molecules can react with itself

H20(l) + H20(l) ⇋ OH-(aq) + H30+(aq)

b. The Keq will measure the extent of the water reaction

This equilibrium expression has concentration of water as the denominator. Since waters concentration does not change after there is enough water we can change the expression to

The new Keqis designated as Kw and kw= [OH-(aq)][H30+(aq) ]

c. The [ H3O+ ] & [ OH- ] in water was discovered to be 1.0 x 10-7 mol/L , therefore the water constant Kw = 1.0 x 10-14 mol2 / L2

If the [ H3O+ ] is low, then [OH- ] will be high in a basic solution

1.0 x 10

Acidic solutions water basic solution

3. pH & pOH calculation

pH = - log [H30+(aq) ] and pOH = - log [OH-(aq)]

4. The acid equilibrium reaction:

When acids are placed into water, the reaction with water will result in the production of H3O+

An equilibrium expression can be written about this reaction

Because water does not have a concentration, this expression can be changed into

_a

-3

'

eq K

] HF [ ] F [ ] O H [

K  



The equilibrium constant for any acid reaction is defined as Ka

The larger the value of Ka , the stronger the strength of the acid

The range of Ka value of HClO4 to HNO3 is very large to 24 and they are strong acids The range of Ka values of the weak acids is from 0.54 to 1.0 x 10-14

HF + H2O ⇄ H3O+ + F

5. The base equilibrium reaction

When bases are place into water, the reaction with water will result in the production of OH-

NH3 + H2O ⇄ OH- + NH4+

An equilibrium expression can be written about this reaction

Because water cannot have a concentration, this expression can be changed into

The equilibrium constant for any basic reaction is called Kb it can be found by knowing that Ka x Kb = Kw

The larger the value of Kb, the stronger the strength of the base The range of value of Kb will be the same as Ka

] O H [ ] NH [ ] NH [ ] OH [ K 2 3 4 -eq   b 3 -4 ' eq K ] NH [ ] OH [ ] NH [

K  

ø

1. Given the following equation below, label each species as Bronsted- Lowry acid or base

a. HSO3-(aq) + H2O(l)  H3O+(aq) + SO32-(aq)

b. NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

c. HF(aq) + HSO3-(aq) F-(aq) + H2SO3(aq)

d. H2SO3(aq) + HS-(aq) HSO3-(aq) + H2S(aq)

2. Predict the product produced in each of the following equation. a. do not invent unfamiliar species of chemicals

b. label each species in equation as Brønsted – Lowry acid or base c. assume the transfer of only one proton

a. HNO2(aq) + Cl-(aq) ________ + _________

b. CH3COOH(aq) + SO42-(aq) _________ + _________

c. H2S(aq) + NO2-(aq) _________ + ________

d. HCO3-(aq) + S2-(aq) ________ + ________

What name is given to species which act both as acid and as base ?

Procedure :

1. list major species present in solution

- 6 SA listed as H3O+ & negative ion - major species for WA will be acid solute - ionic bases listed as OH- & positive ion - ionic compounds dissociated

- molecular compounds will not dissociate

2. identify SA & SB ( using data table )

3. predict products – assume one proton transfer ( unless told otherwise ) 4. add some reaction arrows. Rules for arrows

a. quantitative rules : H3O+(aq) + OH-(aq)

: told with words ( quantitative reaction, end point, equivalent point ) with indicators ( the arrow ALWAYS points away from the H30+(aq))

b. favoring rules : if A is above B – favor products if B is above A – favor reactants

Hey---kind of like spont and nonspont REDOX reactions

5. For polyprotic acids/bases repeat steps 1 through 4 for the species list of the products of each reaction. The net acid base reaction is the sum of all the quantitative reactions.

Write net-ionic equation for each of the following given

a. Barium hydroxide solution is mixed with sulfuric acid solution

c. Sodium hydrogen carbonate solution is reacted with hydrofluoric acid

d. Hydrogen sulfide solution is mixed with sodium nitrite solution

e. Methanoic acid is reacted with sodium sulfate solution

f. Two quantitative reactions are observed when oxalic acid is titrated with sodium hydroxide solution

g. One end point was reached when citric acid is titrated with barium hydroxide solution

h. Hydrofluoric acid ( Ka = 6.3 x 10-4 ) is titrated with potassium hydroxide solution using

methyl orange indicator ( Ka = 3.5 x 10-4 )

i. Sodium phosphate solution ( Kb = 2.1 x 10-2 ) is titrated with hydrochloric acid using

j. One equivalent point is reached when sodium carbonate is titrated with hydrochloric acid

k. Given the following information, list the acid from strongest to weakest HA + D- _{< = > A}- _{+ HD ( favors product )}

HD + B- _{< = > D}- _{+ HB ( favors product )}

HA + C- _{< = > A}- _{+ HC ( favors reactant )}

l. Given the following information, rank the four bases from strongest to weakest HNO2 + HCOO- < = > NO2- + HCOOH ( products favored )

HN3 + OBr- < = > N3- + HOBr ( products favored )

HN3 + HCOO- < = > N3- + HCOOH ( reactants favored )

*Since the strong acids ionize and dissociate 100% the [H30+_{] = [ACiD]}

1. Equations :

2. Use :

A. If given concentration of acid ( or [ H+ _{] ) and asked to find pH}

e.g. [ H+ _{] = 0.00043 mol/L , find pH of solution ( pH = - log [ H}+ _{] )}

i. for programable calculators a. punch “ – “ sign

b. punch “ log “ button

c. input number into calculator d. punch “ = “ to get answer

ii. for non- programmable calculators : a. input number into calculator b. punch “ log “ button

c. punch in “ – “ sign by punching “ +/- “

B. if given pH and asked to find [ H+ _]

e.g. pH = 3.366, find [ H+ _{] ? ( [ H}+ _{] = 10}-pH ₎

i. for programmable calculators

input equation exactly as stated and punch in = to find answer

ii. for non- programmable calculators a. punch in number

b. punch in “ – “ sign by punching “ +/- “ c. punch “ 2nd _{function “ “ log “}

depend on calculator , this could also be “ SHIFT “ “ log “

C. if given mass of solute and volume of solvent, and ask to calculate [ H+ _{] , pH , [ OH}- _{] or}

pOH

1. calculate mole of solute ( mol = m / M )

2. calculate concentration of solution ( conc. = mol/ volume ) 3. write dissociation equation

4. calculate [ H+ _{] or [ OH}- _]

5. calculate pH or pOH using pH and pOH equation

pH = - log [ H+ _{] pOH = - log [ OH}- _]

[ H+ _{] = 10}-pH _{[ OH}- _{] = 10}-pOH

pH=pOH= 14.00 [H+_{] [OH}-_{] = 1.0 x 10}-14

- log

- log 0.00043

3.366531544 -

0.00043

- 3.366531544

1. Do the following calculation :

[ H+_{(aq) ] ( mol/L ) is solution acidic or basic ? pH}

1 0.125 2 4.5 x 10-5

3. 3.2 x 10-8

4. 6.7 x 10-13

5. 0.000000789

6 4.25

7 6.59

8 7.38

9 10.59

2.

[ OH-_{(aq) ] ( mol/L ) is solution acidic or basic ? pOH}

1 0.34 2 0.000045

3 2.89

4 11.65

2.5

pH pOH [H+(aq)] [OH-(aq)] Acid/Base/ Neutral

-1.25

8.23

1.3 x 10-5

1.3 x 10-5

-1.25 8.23

1.0 x 10-7

5.7 x 10-2

3. 15.0 g of hydrogen perchlorate is dissolved into water to make a 1.50 L solution. What is the pH of the solution ?

5. 15.0 g of barium hydroxide is dissolved in water to make a 500 mL solution. What is the pOH of the solution ?

6. A 300 mL solution was found to have a pH of 2.67 . What mass of hydrogen bromide must have been dissolved to make this solution ?

Answers to questions

[ H+_{(aq) ] ( mol/L ) is solution acidic or basic ? pH}

1 0.125

A 0.903

2 4.5 x 10-5 _{A 4.35}

3. 3.2 x 10-8 _{B 7.49}

4. 6.7 x 10-13 _{B 12.17}

5. 0.000000789 A 6.103

6 5.6 x 10-5 _{A 4.25}

7 2.6 x 10-7 _{A 6.59}

8 4.2 x 10-8 _{B 7.38}

9 2.6 x 10-11 _{B 10.59}

2.

[ OH-_{(aq) ] ( mol/L ) is solution acidic or basic ? pOH}

1 0.34 B 0.47

2 0.000045 B 4.35

3 1.2 x 10-3 _{B 2.89}

2.5

pH pOH [H+(aq)] [OH-(aq)] Acid/Base/ Neutral

-1.25 15.25 18M 5.6 x 10-16_{M A}

5.77 8.23 1.7 x 10-5 _{M 5.9 x 10}-9 _{M A}

4.87 9.13 1.3 x 10-5 _{7.4 x 10}-10_{M A}

9.13 4.87 7.4 x 10-10_{M 1.3 x 10}-5 _{M B}

15.25 -1.25 5.6 x 10-16_{M 18M B}

8.23 5.77 5.9 x 10-9 _{M 1.7 x 10}-5 _{M B}

7.00 7.00 1.0 x 10-7_{M 1.0 x 10}-7_{M N}

12.76 1.24 1.7 x 10-13_{M 5.7 x 10}-2_{M B}

3. 15.0 / 100.46 g/mol = 0.149 mol 4. 20.0 g / 98.09 = 0.204 0.149 mol / 1.50 L = 0.0995 M 0.204 / .800 L = 0.255 M [ H+ _{] = 0.0995 pH = 1.002 0.255 M pH = 0.594}

5. 15.0 g / 171.35 = 0.0875 mol 0.0875 / .5 = 0.175 M

0.175 M x2 = 0.350 M pOH = 0.456

6. pH = 2.67 [ H+ _{] = 2.1 x 10}-3 _{M = [ HBr ]} 2.1 x 10-3 _{M x 0.300 L = 6.4 x 10}-4 _mol 6.4 x 10-4 _{mol x 80.91 g/mol = 0.0519 g}

Hydronium Ion and Hydroxide Ion Concntration:

1. What is the molar H3O+ ion concentration in solutions with the following OH- ion

concentrations?

a) 2.0 X 10-4 _mol/L

b) 7.3 X 10-7 _mol/L

c) 3.0 X 10-10 _mol/L

d) 2.5 X 10-8 _mol/L

2. In the problem above (1), indicate whether each solution is acidic, basic, or neutral.

3. What is the molar OH- _{ion concentration in solutions with the following H3O}+ _ion

concentrations?

C. 1.0 X 10-7 _{mol/L D. 5.0 X 10}-8 _mol/L

4. In the problem above (3), indicate whether each solution is acidic, basic, or neutral.

pH:

5. Calculate the pH of solutions with the following hydronium ion concentrations. a) 1.0 X 10-4 _mol/L

b) 1.0 X 10-9 _mol/L

c) 0.000010 mol/L

d) 0.0000000010 mol/L

e) 4.0 X 10-2 _mol/L

f) 5.0 X 10-7 _mol/L

g) 3.0 X 10-3 _mol/L

h) 3.0 X 10-3 _mol/L

i) 3.000 X 10-3 _mol/L

6. What is the hydronium ion concentration associated with each of the following pH values?

a) 3.0 b)5.7

c) 2.43 d)3.43

7. Solution A has [OH-_{] = 4.3 X 10}-4 _{mol/L. Solution B has [H3O}+_{] = 7.3 X 10}-10

mol/L.

a) Which solution is more basic?

b) Which solution has a lower pH?

8. A solution has a pH of 4.500. What will be the pH of this solution if the hydronium ion concentration is

a) doubled

b) quadrupled

c) increased by a factor of 10

d) increased by a factor of 1000

9. In which of the following pairs of solutions does the first listed solution have a lower pH than the second listed solution?

a) 0.1 mol/L HCl and 0.2 mol/L HCl

b) 0.1 mol/L HCl and 0.1 mol/L H2SO4

c) 0.2 mol/L H2SO4 and 0.2 mol/L H2CO3

10.What would be the pH of a solution that contains 0.1 mole of each of the solutes NaCl, HNO3, HCl, and NaOH in enough water to give 3.00 L of solution?

pH Calculations Answers

Hydronium Ion and Hydroxide Ion Concentration: 1.

a) 5.0 X 10-11 mol/L b) 1.4 X 10-8 _mol/L c) 3.3 X 10-5 mol/L d) 4.0 X 10-7 mol/L 2. a) basic b) basic c) acidic d) acidic 3.

a) 3.7 X 10-11 mol/L b) 1.3 X 10-6 mol/L c) 1.0 X 10-7 mol/L d) 2.0 X 10-7 mol/L 4. a) acidic b) basic c) neutral d) basic pH: 5. a) 4.0 b) 9.0 c) 5.0 d) 9.0 e) 1.4 f) 6.2 g) 2.5 h) 2.52 i) 2.5229 6.

a) 1 X 10-3 _mol/L b) 2 X 10-6 mol/L c) 3.7 X 10-3 mol/L d) 3.7 X 10-4 mol/L e) 3.5 X 10-8 mol/L

7.

a) Solution A b) Solution B 8. a) 4.199 b) 3.900 c) 3.500 d) 1.500 Additional Problems: 9.

a) 0.2 mol/L HCl(aq) b) 0.1 mol/L H2SO4(aq) c) 0.2 mol/L H2SO4(aq)

For the reaction HA(aq) + H20(l) ⇄ A-(aq) + H30+(aq) the percent ionization can be described as the concentration of H30+(aq) at EQ divided by the original acids concentration.

says that if the Ka x 1000 < [HA] or [HA(aq)]÷ Ka >1000 the

assumption can be made. This will always be the case if given a Ka and asked to find

the pH of an acid.

So in summary

1.What is the pH and % Rxn of a 1.0 mol/L acetic acid solution? ( 2.37, 0.42%)

2. What is the pH and % Rxn of a 0.10 mol/L HOCl (aq)? ( 4.20, 6.3 x 10-2 _%)

4. What is the pH and % Rxn of a 1.2 mol/L hydrofluoric acid solution? (1.55, 2.3%)

5. What is the pH and % Rxn of a 0.85 mol/L methanoic acid solution? (1.91, 1.5%)

6. Estimate the pH and % Rxn of a 3.0 mol/L aqueous sodium hydrogen sulphate solution? (0.76, 5.8%)

7. What is the pH and % Rxn of a 0.20 mol/L ammonium chloride solution? (4.97, 0.0054%)

9. What is the the pH of 1.00 mol/L aqueous sodium hydrogen oxalate?

10. What is the pH of 0.25 mol/L abietic acid (Ka = 2.4 x 10-8_)?

11. What is the pH of 0.85 mol/L acetoacetic acid (Ka = 2.6 x 10-4_)?

12. What is the pH of 0.10 mol/L aqueous allantoin (Ka = 1.1 x 10-9_)?

13. What is the pH of 0.10 mol/L vitamin B-12 (Ka = 2.3 x 10-8_)?

15. What is the Ka of a weak acid, if its 1.20 mol/L aqueous solution has a pH of 1.69? Suggest the weak acid’s identity.

16. What is the Ka of a monoprotic weak acid if its 0.010 mol/L aqueous solution has a pH of 3.22? Suggest the weak acid’s identity.

Never us the rule of 1000 or the assumption when working backwards or solving for the Ka. If the pH is given it was measured and we can use this to figure out the [H3O+(aq)].

How to Solve for the pH and pOH of weak bases…. First we must remember

that:

K

x K

= K

K

= 1.0 x 10

so the Kb = Kw/Ka but you must use the Ka of the bases conjugate acid

So to find the Kb of HC03-(aq) you must divide the Kw by the Ka of H2CO3(aq)

So what are the Kb of:

S042-(aq)

HP042-(aq)

2. Estimate the pH of a 0.25 mol/L solution of aqueous sodium cyanide. (11.30)

3. Estimate the pH of a 0.50 mol/L solution of aqueous sodium nitrite. (8.48)

4. Estimate the pH of a 0.10 mol/L solution of aqueous sodium fluoride. (8.10)

6. Estimate the pH of a 1.0 mol/L solution of aqueous sodium bicarbonate. (10.17)

7. Estimate the pH of a 2.0 mol/L solution of aqueous sodium methanoate. (9.02)

A polyprotic acid will become an amphoteric substance after donating 1 proton A polybasic species will become an amphoteric substance after receiving a proton

H2CO3  HCO3-  CO32-

polyprotic amphoteric polybasic acid substance species

another example is

H3PO4  H2PO4-  HPO42-  PO43-

polyprotic amphoteric polybasic acid substances species

Will the amphoteric substance exhibit acidic or basic properties when dissolved in water ?

Is NaHCO3(aq) an acid or base ?

Is NaH2PO4(aq) or Na2HPO4(aq) an acid or base ?

We can tell by comparing the value of the Ka against the Kb of the amphoteric substance.

If the Ka is larger than the Kb , then the amphoteric substance will exhibit acidic properties. The

So now you can figure out the pH of amphoteric substances.

What is the pH of 1.2M sodium carbonate ( baking soda) ?

Indicators are organic species which are usually extracted from plants. This class of chemicals are used in titration to indicate the end points of titrations

4 Important points in using indicator for titration:

1. names & formulae : 3. Weaker acid or base : 2. color properties 4. Mid range color

1. Names & formulae :

Names given to indicators usually contain its natural color e.g. Indian Blue, Congo Red etc

Since their chemical compositions are fairly complex, we will use their names as formula therefore bromothymol blue = Bb , methyl orange = Mo , cresol red = Cr

2. Color properties :

A. Each indicator has 2 two forms. An acid form and a base form. When it is in its acid form, the chemical formula used is HIn When it is in its basic form, the chemical formula used will be In

e.g HBb and Bb- _{will be the acid and base form of bromothymol blue}

HMo and Mo- _{will be the acid and base form of methyl orange indicator}

B. The two forms of the indicator must have distinctively different colors than each other to be a good indicator.

e.g. HBb is yellow in color and Bb- _{is blue}

HMo is red in color and Mo- _{is yellow}

C. The names, formula of the two forms, colors and Ka are given on page 10 of data table

Note : acid form color is given first and then base form color

3. must be a weaker acid or base:

The indicator used in a titration must be a weaker acid or base than the acid or base titrated

If an acid is to be titrated, and the acid has a Ka value of 3.6 x 10-2

Then the indicator used must have a Ka value lower than that

A. titration of monoprotic acid :

if HF is the acid titrated and its Ka is 6.3 x 10-4 , then indicator used to monitor this

reaction should have their Ka less than that of HF ( any value below that )

B. titration of diprotic acid :

if a diprotic acid such as oxalic acid is titrated and two quantitative reactions are cited HOOCCOOH(aq) + OH-(aq)  HOOCCOO-(aq) + H2O(l)

HOOCCOO-_{(aq) + OH}-_{(aq)  OOCCOO}2-_{(aq) + H2O(l)}

The indicator used for the first titration should have Ka between 5.6 x 10-2 to 1.5 x 10-4

H2Tb ( Ka = 2.2 x 10-2 ), HMo ( Ka = 3.5 x 10-4 ) could be used for this titration

The indicators used to indicate the second titration should have Ka lower than 1.5 x 10-4 HBg or below should be OK.

C. Only 1 indicator can be used for each titration. If two indicators are used in a titration, the natural colors of the indicators may interfere with each other and the color change may not be observed distinctly.

D. Same concepts are applied to base titrations. The only difference is the use of Kb instead of Ka

HOOCCOOH Ka = 5.6 x 10-2

H2Tb Ka = 2.2 x 10-2

HMo Ka = 3.5 x 10-4

HOOCCOO- Ka = 1.5 x 10-4

HBg Ka = 1.3 x 10-5

These two indicators can be used for HOOCCOOH titration

4. Mid range color of indicator

using bromothymo blue as an example :

Other mid range color : red – yellow = orange ; blue – red = purple Yellow blue

6.0 14

pH = 0.0

Y

Y

Y Y

Y

Y

B

B

B Y

B

B Y

Y

B

7.6

1. Find the pH range and average hydronium ion concentration of the following solutions – please!

Solution A : After addition to separate samples of the solution, methyl violet was blue, methyl orange was yellow, methyl red was red and phenolphthalein was colorless.

Solution B : After addition to separate samples of the solution, indigo carmine was blue, phenol red was yellow, bromocresol green was blue and methyl red was yellow.

Solution C: After addition to separate samples of the solution, phenolphthalein was colorless, thymol blue was yellow, bromocresol green was yellow and methyl orange was orange.

Use this page to analyze the titration of a strong acid with a strong base titrant.

Plot the data collected from on the titration of 100 mL of HCl(aq) with 0.1 M NaOH(aq) pH vs. volume graph below.

pH vs volume of NaOH(aq)added

1. How does your graph show:

a. the sample was a strong acid? b. The equivalence point?

c. This is a reaction of a strong acid with a strong base?

d. The volume of titrant required to consume all of the sample?

2a. Shade in the area atop your graph line where each of these indicators would change:

i. methyl orange ii. bromothymol blue iii. phenolphthalein

b. Which of the three indicators will show an endpoint that corresponds to the equivalence point?

3. Use the initial pH from your graph to predict the concentration of the original acid.

4. Explain why the pH changes so slowly in the first part of the graph and so quickly in pH

Titration curves and buffering regions

pH pH = 7.0

13

10

7

4

1

End point volume

no buffer pH pH = 7.0

End point volume

no buffer 13 10 7 4 1

Titration curves for ( left ) a strong acid titrated with a strong base ( right ) a strong base titrated with a strong acid.

The region from the start to the end point is not buffered. The end-point pH is always 7.0

Titration curves for ( left ) a weak acid titrated with a strong base ( right ) a weak base titrated with a strong acid.

Between the start and the end point is a buffer region ( mark with an X ) due to the presence of a conjugate acid-base mixture. Because of the hydrolysis of the products, the end point pH is NOT 7.0

Note : hydrolysis means reaction with water.

e.g. the hydrolysis of HCl means HCl + H2O  H3O+ + Cl

In the case of the the graph on the left, ( assume the weak acid to be HF(aq)) After the HF is neutralized by OH‾, ( HF(aq) + OH-(aq)  F‾(aq) + H2O(l) ) The product F‾(aq) can hydrolyze with water to produce a basic solution,

therefore the equivalent point is above pH of 7.

In the case of the graph on the right, ( assume the weak base to be NH3(aq) ) After the NH3 is neutralized by OH-, ( NH3(aq) + H3O+(aq) NH4+(aq) + H2O(l) )

pH > 7.0 pH 13 10 7 4 1

End point volume

buffer

X

X

pH < 7.0 pH 13 10 7 4 1 buffer region

Titration curve for the titration of phosphoric acid with a strong base

BUFFERS

A buffer is a system that is resistant to change. To have a buffer we must have a weak acid or a weak base and its appropriate conjugate in the system at the same time. Since there will reach equilibrium with each other the system will react to any stress that you apply. IE : removal or addition of H30+(aq) or OH-(aq)

Type of titration

endpoint pH

Strong acid + Strong base pH = 7 Weak acid + Strong base pH > 7 Strong acid + weak base pH < 7

pH

End point 2 End point 1

13

10

7

4

1

H3PO4

and

H2PO4 buffer

zone

H2PO4

and

HPO42

buffer

zone

HPO42

and

PO43

buffer

zone

Titration graph of a triprotic acid ( such as H3PO4(aq) ) with NaOH(aq) Titration graph give 2 equivalent points

1st _{: H3PO4(aq) + OH‾(aq)  H2PO4}2-_{(aq) + H2O(l)} 2nd _{: H2PO4‾(aq) + OH‾(aq)  HPO4}2_{‾(aq) + H2O(l)}

the 3rd _{equivalent point becomes impossible because the acid is so weak it} cannot be quantitatively neutralize by OH‾(aq)

HPO42‾(aq) + OH-(aq) ⇋ PO43‾(aq) + H2O(l)

The three regions where the three buffering pairs are exerting their influences are indicated on the graph

We can use ALL the endpoints to determine the concentration of an acid or a base during the titration

Here is a classic buffer set up

What happens to the [H30+(aq)] when you add a base to the system? Which way does the EQ shift

to fix this stress?

What happens to the [H30+(aq)] when you add an acid to the system? Which way does the EQ

shift to fix this stress?

Provide Brønsted-Lowry predictions of all the reactions that are likely to occur in the following scenarios:

a. Sulfuric acid is titrated with sodium hydroxide in the presence of methyl orange indicator.

b. Hydrochloric acid is continuously added (i.e. it is always present) to an aqueous solution ofsodium citrate in the presence of phenolphthalien.

c. An aqueous solution of sodium carbonate is titrated with nitric acid (which means it is added continuously). Which of phenol red, bromothymol blue or methyl orange would be the best choice for indicating: (i) the 2nd quantitative reaction? (ii) the 1st quantitative reaction?

e. A solution of sodium hydroxide is continuously added to citric acid containing some bromocresol green.

f. A sample of aqueous sodium sulfite containing orange (IV) is titrated to the first equivalence point with perchloric acid.

g. Nitric acid is continuously added to a solution of sodium phosphate. Suggest indicators for the 1st and 2nd neutralizations.

h. Sodium hydroxide is continuously added to a solution of phosphoric acid.