acids - bases in waterfaculty.sdmiramar.edu/fgarces/zcourse/all_year/ch201/amy... · 2019-01-23 ·...

June 14 1 Acids - Bases in Water

Acids - Bases in Water …more equilibrium

Dr. Fred Omega Garces Chemistry, Miramar College


Acids (Properties) •Taste Sour •Dehydrate Substances •Neutralizes bases •Dissolves metals

Examples: •Juices: TJ, OJ, AJ •Wine •Banana •Coffee •Vitamin C •Soda

Acids-Bases Characteristics Base (Properties) •Taste Bitter •Denatures Proteins •Neutralizes acids •Turns metal g hydroxides

Examples: •Milk of Magnesia •Lime water •Lye, Drain •Ammonia •blood •Soap


Practice Naming Acids HNO2 HI HF H3PO4 H3PO3 HClO2 HClO H2CO3 HCN HC2H3O2 H2Cr2O7 HClO4

Nitrous Acid

Hydroiodic Acid

Hydrofluoric Acid

Phosphoric Acid

Phosphorous Acid

Chlorous Acid

Hypochlorous Acid

Carbonic Acid

Hydrocyanic Acid

Acetic Acid

Dichromic Acid

Perchloric Acid

Oxy-anions Oxy-acids

Add H+

Add H+

Add H+

Add H+


Svante Arrhenius (1859 - 1927) Acid - Increases H+ (H3O+) concentration Base - Increases OH- concentration Arrhenius acids and bases are limited to aqueous solutions.

Arrhenius Definition

Examples: Acids are substances that are able to ionize to form hydrogen ion and thereby increase the concentration of H+

(aq) ions in aqueous solutions. HNO3 (aq) g H+ (aq) + NO3

- (aq)

Bases are substances that accept (react with) H+ ions. Hydroxide ions, OH-, are basic because they readily react with H+

ions to form water: H+ (aq) + OH - (aq) g H2O (l)


New Definition: Bronsted-Lowry Acids-Bases

Bronsted - Lowry definition

Acid - Proton H+ (H3O+) donor Base - Proton H+ (H3O+) acceptor.

example: acids: HCl (aq) → H+

(aq) + Cl- (aq) Bases: NH3 (aq) + H2O(l) → NH4

+ (aq) + OH- (aq) HCl (aq) + NH3 (aq) → NH4

+ (aq) + Cl- (aq

In an acid - base reaction, H+ & OH- always combine

together to form water and an ionic compound (a salt): HCl(aq) + NaOH(aq) g H2O(l) + NaCl(aq)


Factors Affecting Acid/Base Strength What determines the strength of acids and Base? Dissociation property- Electrolyte - Substances which dissociate in water. Strong electrolyte completely dissociates to ions Weak electrolyte undergoes partial dissociation.

Acid Examples : HClO4 + H2O g H3O+ + ClO4

- 100% Dissociation (Strong Acid) HNO3 + H2O g H3O+ + NO3

- 100% Dissociation (Strong Acid) H2S + H2O D H3O+ + HS- Less 100% Dissociation (Weak Acid)

Base Examples :

NaOH + H2O g OH- + Na+ 100% Dissociation (Strong Base) Ca(OH)2 + H2O g 2OH- + Ca+2 100% Dissociation (Strong Base) NH3 + H2O D OH- + NH4

+ Less 100% Dissociation (Weak Base)


Strong Acids Strong Acids

Earlier we described acids and bases as electrolyte; that is, these substances ionizes (break up to ions) in solution.

Strong electrolyte are substances which completely dissociates (100%).

SA-Strong acids completely dissociates to H + and anion.

HA g H+ + A-

Strong Acids; HX, H2SO4, HNO3, HClO4 HClO3 :

All others are considered weak.

HX: HCl, HBr, HI (all Halogens, except F )

H2SO4 Sulfuric Acid HNO3 Nitric Acid

HClO4 Perchloric Acid HClO3 Chloric Acid

Weak acids, incompletely dissociation in water.

i.e., CH3COOH (acetic acid) only 1 in 100 dissociate


Strong Acids; indicative by acid dissociation constant, Ka

The Ka values shown are the acid dissociation constant for monoprotic acids. The Ka is the equilibrium constant for the ionization of the acid, HA D H+ + A. Strong acids do not have meaningful Ka values and are left out from this table.


Strong Bases Strong Bases

Substances which completely dissociate into cation and hydroxide.

MOH g M+ + OH -

Strong Bases, M(OH)n ; NaOH, KOH, LiOH , Ca(OH)2, Ba(OH)2, Sr(OH)2

Mn O; Li2O, Na2O, K2O

All others are considered weak.

Bases: 3 ways of OH- formation

Extraction of H+ from H2O:

i.e.., reaction of NH3 in water: NH3 (aq) + H2O(l) D NH4+

(aq) + OH- (aq)

Extraction of H+ from H2O by charge

i.e.., reaction of CO32- in water: CO3

2- (aq) + H2O(l) D HCO3- (aq) + OH- (aq)

Dissociation of ionic substance to OH -

i.e.., dissociation of NaOH in water: NaOH(aq) D Na+ (aq) + OH- (aq)


Chemistry of Acids and Bases Acids Base

Strong Acid Weak Acid


Water: Acid-Base Properties Auto-ionization (Self-Ionization) of water Why does water have a pH of 7 ?

Water is Amphoteric (it reacts with itself) 2 in 1 billion water molecules self-ionizes.

H2O (l) + H2O (l) + Energy D H3O(aq) + OH-

(aq)

Keq = Kw

Endothermic reaction

Kw(ion-product constant) = 1•10-14

see that,

[H3O+] = 1•10-7 M

[OH-] = 1•10-7 M

The equation Kw = [H3O+][OH-] is valid in pure water and in any aqueous solution. Kw is temperature-dependent, the auto-ionization rxn is endothermic, so Kw increases with temperature.

°C Kw 10 0.29•10-14 15 0.45•10-14 20 0.68•10-14 25 1.01•10-14 30 1.47•10-14 50 5.48•10-14

Kw


Kw & [H3O+] How does Kw dictate concentration of H3O+ and OH- ?

€

Kw = 1• 10−14 = H3O+# $ %

& ' ( OH−# $ %

& ' ( at 25°C

For pure water,

If [H3O]+ or [OH]- is concentration known, the other can be determine through the mass action expression and a iΔe table.

€

2H2O(l) ! OH- + H3O+

i Excess 0 0Δ - 2x + x + xe excess x x

€

Kw = 1• 10-14 = x2

1• 10-7 = x = [H3O]+ = [OH]-pH g 7.0 pH calculation to be discussed.


Consequences of Auto-ionization • A change in [H3O+] causes an inverse change in [OH-] & vice versa. Higher [H3O+] g lower [OH-] Higher [OH-] g lower [H3O+] • Both ions are present in all aqueous systems.

Acidic solution g [H3O+] > [OH-] Neutral solutions g [H3O+] = [OH-] Basic Solutions g [OH-] > [H3O+]

Kw = [H3O+] • [OH-] or pKw = pH + pOH (as shown later)


pH Calculation

For pure water [H3O+] = 1•10-7 M Since 1•10-7 M = 0.0000001M Concentration is so small, it is much more convenient to use a scale called “p-H” (power of hydrogen or potential of hydrogen)

Example: 1) What is the pH of a hydrobromic acid solution with a molar

concentration of 1.67•10-5 M. (3-Sig figures) 2) Double check your answer by taking the pH answer from #1, and calculate the

molar concentration of the H3O+ for the hydrobromic acid solution. Calculate: pH for [H3O+] = 1.67•10-5 M [H3O+] of pH = 4.777 Calculator Sequence:

Answer : 4.777 (3-Sig figs)

1•10-7 M g pH = 7

pH = - log [H3O+] or [H3O+] = 1•10-[pH]

1.67 EE 5 +/- log +/- 4.78 +/- 10x

1.67•10-5


pH Scale

pH and the Concentration of Acids

Conc [H3O+] Exp [H3O+] pH pOH

1M 1•100 0 14 0.1M 1•10-1 1 13 0.01M 1•10-2 2 12 0.001M 1•10-3 3 11 0.0001M 1•10-4 4 10 0.000001M 1•10-5 5 9 0.0000001M 1•10-6 6 8 0.00000001M 1•10-7 7 7 0.000000001M 1•10-8 8 6 0.0000000001M 1•10-9 9 5 0.00000000001M 1•10-10 10 4 0.000000000001M 1•10-11 11 3 0.0000000000001M 1•10-12 12 2 0.00000000000001M 1•10-13 13 1 0.000000000000001M 1•10-14 14 0


Kw, pH and pOH

pH measures the concentration of [H3O+]

Kw = 1•10-14 = [H3O+] [OH-]

pKw = pH + pOH

14 = pH + pOH

Example: What is the pH of a potassium hydroxide solution with a molar concentration of 3.0•10-4 M (2 sig figs)

Answer: Potassium hydroxide, KOH is a strong base, therefore, the molar concentration of KOH (3.0•10-4 M) is equal to the [OH-] concentration or [OH-] = 3.0•10-4 M

pOH = -log (3.0•10-4 ) = 3.52

pH + pOH = 14.00 g pH = 10.48


Determining pH, pOH, [OH-], [H3O+]

[H3O+] • [OH-] = Kw & pH + pOH = pKw

[H3O+] [OH-] Kw/

Kw/

-log [H3O+]

10-pH 10-pOH

-log

14.0 - pOH

14.0 - pH

[H3O+] • [OH -] = Kw

pH + pOH = pKw= 14

Use this chart to determine acid and base concentration at (25°C) Kw = 1.00 •10-14 & pKw = 14


Kw, pH and pOH Example 2

pH measures the concentration of [H3O+] Kw = 1•10-14 = [H3O+] [OH-] pKw = pH + pOH 14 = pH + pOH

Example: What is the pH of a solution having a [OH-] = 3.0•10-13 M

Way of Multiplying Way of adding / subtracting Using the formula: Using the formula: 1•10-14 = [H3O+] [OH-] 14 = pH + pOH

1•10-14 = [H3O+] [OH-] 14 = pH + pOH 1•10-14 =[H3O+] [3.0•10-13] 14 = pH + 12.52 [H3O+] =3.33•10-2M pH = 1.48 pH = 1.48


Acid Base in Water

Summary Pure water has a low conductivity because it autoionizes to a small extent.

This process is described by an equilibrium reaction whose equilibrium

constant is the ion-product constant for water, Kw (1•10-14 at 25°C). Thus,

[H3O+] and [OH-] are inversely related. In acidic solution, [H3O+] is

greater than [OH-], the reverse is true in basic solution, and the two are

equal in neutral solution. To express small values of [H3O+] more simply, we

use the pH scale (pH = -log [H3O+]). A high pH represents a low [H3O+].

Similarly, pOH = -log [OH-], and pK = - log K. At 25°C, in acidic solutions,

pH < 7.00, in basic solutions, pH > 7.00; and in neutral solutions, pH = 7.0.

The sum of the pH and pOH equals pKw (14.00 at 25°C)

53

Activity 7: Basics of Acid/Base Chemistry

Objective: The purpose of this exercise is to gain proficiency in working with a variety of acid-base problems.

Discussion: Acids are commonly described as sour, tasting like vinegar or lemon juice. Water solutions of bases are said to taste bitter

and feel slippery, but most strong bases like lye (sodium hydroxide) are too bitter and harsh to be tasted. Examples of milder bases are

slaked lime (calcium hydroxide) and ammonium hydroxide. Solutions of ammonia, NH3, in water are often called ammonium hydroxide,

NH4OH, because ammonium ions, NH4+ , and hydroxide ions, OH-, are formed in low concentration in the reaction with water.

NH3(aq) + H2O(aq) D NH4+(aq) + OH-(aq)

Acids are characterized by the presence of H+ ions (actually H3O+ ions) in water solutions, and bases are typically those compounds

having OH- ions. The H+ ions of acids and the OH- ions of bases react with each other to form water, so acids and bases are said to

neutralize each other. For this reason they are considered to be opposites. In this exercise, you will learn to measure and to express in

concise terms the degree of acidity or basicity of a solution. The famed pH scale was devised precisely for this purpose.

An understanding of acidity and basicity in water solution is grounded on the concept of the auto-ionization of water. Pure water auto-

ionizes only about 0.000001%, or 2 in 1 billion; the auto-ionization reaction is shown in the following equation.

H2O + H2O D H3O+ + OH-

The hydronium ion, H3O+ is often written as H+. The ionization of water can be simplified to

H2O D H+ + OH-

The equation shows that in pure water, there are as many H+ ions as there are OH- ions; that is, the H+ ion concentration, [H+], and the

OH- concentration, [OH-], are equal. The actual concentration of each is very small. In pure water at 20 °C, these are both known

experimentally to be only 1 x 10-7 (0.0000001) mole/L, or 1 x 10-7 molar (M).

[H+] = [OH-] = 1 x 10-7 mole/L

In acid solution, the H+ ion (or hydronium ion) concentration is greater than 10-7 mole/L, and the OH- ion (or hydroxide ion)

concentration is less than 10-7: [H+ ] > [OH-]. In basic solution, the opposite is true, and OH- ions predominate: [OH-] > [H+].

The dissociation of water can be described by an equilibrium expression, just like any other reaction at equilibrium. The equilibrium

constant for water, Kw, is found by multiplying the concentration of H+ and OH- ions in pure water, 1 x 10-7 mole/L for each.

Kw = [H+] [OH-] = (1 x 10-7)2 = 1 x 10-14

In any aqueous solution at room temperature, the product of the H+ and OH- ion concentrations is always equal to 10-14. As the

concentration of H+ ions becomes greater (in acidic solutions), the OH- ion concentration necessarily becomes smaller; the product of

the two always equals 10-14. Both kinds of ions are present in all water solutions. The acidity or basicity of the solution depends on

which ion predominates. If their concentrations are equal (at 10-7mole/L), the solution is neutral. For example, if the H+ ion

54

concentration is 10-3 mole/L, then the OH- ion concentration can be found from the equation for the equilibrium constant for water.

Since [H+] [OH-] = 10-14,

10-3 x [OH-] = 10-14

[OH-] = 10-11 mol/L.

The pH scale constitutes a highly convenient method for specifying the acidity (or basicity) of a solution. The pH is actually the

exponent of the H+ ion concentration (the power to which 10 is raised) with its sign changed. (pH is more elegantly defined

mathematically as the negative logarithm of the H+ ion concentration, pH = - log [H+] ). A solution with an H+ ion concentration of 10-3

mole/L has a pH of 3. A pH of 6 means an H+ ion concentration of 10-6 mol/L, which also means an OH- ion concentration of 10-8

mole/L.

The equation below shows the relationship of the quantities pH, [H+], [OH-] and pOH. Given any one of these parameters, the other

three can be calculated, given the Kw of water. This is also summarized in Figure A8.1

Table A7.1 shows the relationships of H+ ion concentrations, OH- ion concentrations, pH and pOH values.

Table A7.1 pH range for solutions in aqueous medium

[H+] pH pOH [OH- ]

101 -1

Acid

15 10-15 100 0 14 10-14 10-1 1 13 10-13 10-2 2 12 10-12 10-3 3 11 10-11 10-4 4 10 10-10 10-5 5 9 10-9 10-6 6 8 10-8 10-7 7 Neutral 7 10-7 10-8 8

Base

6 10-6 10-9 9 5 10-5 10-10 10 4 10-4 10-11 11 3 10-3 10-12 12 2 10-2 10-13 13 1 10-1 10-14 14 0 100 10-15 15 -1 101

At 25 °C, all pH values less than 7 indicate acidic solutions, and all pH values greater than 7 indicate basic solutions. A neutral solution

has a pH of 7 since the H+ and OH- ion concentrations are equal. It should also be emphasized that each pH unit on the scale means a

tenfold increase or decrease in H+ ion concentration from the previous number. Thus, a pH range of O to 14, represents a range of

concentration of H+ (or OH-) ions from 1 mol/L to 10-14 mol/L (from 1 to 0.00000000000001), a tremendous range. Finally, the pH is

not restricted to whole numbers. Calculations of fractional pH values require the use of logarithms. Some fractional pH values and the

corresponding concentrations of H+ and OH- ions are shown in Table A7.2

€

1∗10−14 = [H3O+] ∗[OH−]14 = pH + pOH

€

pH = -log[H3O+]pOH = -log[OH-]

55

Table A7.2 Some fractional pH values

[H+] [OH- ] pH pOH Type solution

2.0 x 10-3 5.0 x 10-12 2.0 12.00 Acidic

7.7 x 10-5 1.3 x 10-10 4.11 9.89 Acidic

5.0 x 10-9 2.0 x 10-6 8.30 5.70 Basic

3.0 x 10-12 3.3 x 10-3 11.52 2.48 Basic

Example: What is the pH of a solution having a [OH-] = 3.0•10-13 M?

Long Way Short way

1•10-14 = [H3O+] [OH-]

1•10-14 = [H3O+] [3.0•10-13]

[H3O+] = 3.33•10-2 M pH = 1.48

pOH = -log [3.0•10-13] = 12.52 14 = pH + pOH 14 = pH + 12.52 pH = 1.48

Figure A7.1 Useful equation for acid – base calculations.

Acid-base indicators are compounds that exhibit one color in an acid solution and a different color in a basic solution. Not all indicators

change color at the same pH; some indicators change colors at pH 7, others at pH 4, pH 5, or pH 6, and some at pH 8, pH 9, or pH 10.

The color shift of an indicator from its "acidic color" to its "basic color," requires about 2 pH units, or 1 full pH unit on either side of

the midpoint. Indicators are useful for determining when solutions change from acidic to basic, and vice versa. In a titration where an

acid and a base are gradually mixed together, it is convenient to have an indicator present to signal when enough of one has been added

to neutralize the other. See Table A8.3, for a list of common indicators, their color ranges, and their transition points.

Generally, indicators are available in laboratories as aqueous solutions. Usually a few drops of indicator in the analyte will show a definite

color change at the endpoint. Sometimes indicators are impregnated in strips of paper that can be dipped into a solution. Litmus papers,

which turn blue in base and red in acid, are common examples. Increasing in use are universal indicator papers, called pH papers or

hydro-ion papers. They show a continuous change of colors over a wide pH range, usually from very red for pH 1, to very blue for pH 14.

There are also available narrow-ranged pH papers, which indicate by color hue the pH of a solution to within a few tenths of a pH unit.

56

Table A7.3 The useful pH range for several common indicators. Note that most indicators have a useful range or effective pH range

of about two pH units, as predicted by the expression pka + 1.

Laboratories that requires rapid accurate pH measurements usually use an electronic device called a pH meter. The theory of how a pH

meter works is complicated, but it is easy to use. After calibrating the pH electrode with buffer solutions, it can be used to give an

instant reading of the pH for any aqueous solution. The measurement is based upon the experimental fact that a change in hydrogen ion

concentration can cause a change in the voltage of an electrochemical cell. The pH meter is fast and consequently is especially valuable

when testing many samples. It is also preferred for dark or colored solutions where chemical indicator colors may be

obscured. Remember however that the accuracy of the pH meter is only as good as the calibration solution and calibration process.

Buffered systems are solutions in which pH changes are minimized when either acid or base is mixed into the solution. The mechanism

for buffered solutions is that either the conjugate base reacts with the excess H+ so the hydronium ion concentration does not

increase, or the conjugate acid reacts with excess OH-, so that the hydroxide ion does not increase. This is the case because conjugate

pairs exist in the same solution. When either of the conjugate pair in a buffered solution is exhausted because of the reaction of

either hydronium or hydroxide, then the buffer capacity is reached for the solution. The buffer capacity is the amount of acid or base

that can be added to a buffered solution before a significant pH change, i.e., 1 pH unit.

http://chemcollective.org/activities/tutorials/buffers/buffers5

WWW Links to acid-base chemistry concepts (Accessed Jan 2019)

1. http://www.ch.ic.ac.uk/vchemlib/course/indi/indicator.html

2. https://www.mccrone.com/mm/handbook-of-acid-base-indicators/

57

Activity 7a: Basics of Acid/Base Chemistry; Part 1

_____ /___Score Name (last)____________________(first)____________________

Lab Section: Day _________ Time _______

Read and follow these directions: For numerical calculations, show your complete work on a separate piece of paper, box your answers, and then write the answer on this worksheet. Remember to use the correct number of significant figures. i. For each chemical listed, write the reaction of these ions in water. Include the phases of all species. Identify and label the

conjugate pairs in your reactions and write the equilibrium constant for each of the reactions you propose. (Use Appendix 9).

a. HPO32-

b. F-

c. HCO3-

ii. (a) Phenoxide ion (C6H5O-) is a weak base, with Kb = 7.70 • 10-5. What is the pKa of the phenol, C6H5OH (conjugate of phenoxide

ion)? Calculate the pH of a 0.50 M aqueous solution of [C6H5O-]. ii. (b) Do you think the pH of 1,0 M tri-methyl ammonium (CH3)3NH+, pKa = 9.80, will be higher or lower than that of 1.0 M phenol,

C6H5OH? What is the difference in pH values for the two acids? iii. Hydrosulfuric acid, also known as hydrogen sulfide, is a diprotic acid. Its two stages of ionization are shown below:

H2S (aq) D H+ + HS-(aq) Kal 5.70 • 10-8

HS- (aq) D H+ + S2-(aq) Ka2 = 1.00 • 10-9

a) Calculate the concentration of HS- ion in a 0.222 M H2S solution. [HS-] =______________(Answer) b) Determine the pH of the solution. pH = ______________(Answer) c) Determine the S2- concentration. [S2-] = ______________(Answer)

iv. Methyl red is a common acid-base indicator. It has a pKa equal to 5.0•10-6. Its un-dissociated form is red and its anionic form is

yellow. a) What color would a methyl red solution have at pH = 7.00? Show your calculations and explain your answers (in words).

Hint: Determine the concentration ratio of methyl red and its conjugate.

59

Activity 7b: Acid/Base Chemistry, Common Ion Effect and Buffers; Part 2

_____ / ___Score Name (last)____________________(first)____________________

Lab Section: Day _________ Time _______

Show your calculations for full credit. Reference: https://labs.chem.ucsb.edu/zhang/liming/pdf/pKas_of_Organic_Acids_and_Bases.pdf v. a) Define buffer capacity and effective pH range.

b) What is the effective pH range for a sodium acetate/acetic acid buffer? Ka = 1.80 • 10-5. Explain your answer. ______________(Answer)

c) Calculate the pH of a solution that is 0.45 M CH3COOH and 0.75 M NaCH3COO. ______________(Answer)

vi. HCN is a weak acid with a Ka = 4.90 • 10-10. A 50.00 mL sample of this HCN solution (0.250 M) is titrated with 0.500 M NaOH. a) What is the pH of the original HCN solution?

______________(Answer)

b) What is the pH after 15 mL of NaOH is added? ______________(Answer) c) What is the pH at the equivalence point? ______________(Answer) vii. Rank the following in order of increasing pOH and justify: 0.10 M solutions of chloroacetic acid, carbonic acid, and citric acid Write out the order instead of placing 1, 2 ... viii. Rank the following in order of increasing pKa and justify: 0.10 M mandelic acid, 0.20 M maleic acid, 0.40 M malonic acid. Write out the order instead of placing 1, 2 ...

acids - bases in waterfaculty.sdmiramar.edu/fgarces/zcourse/all_year/ch201/amy... · 2019-01-23 ·...

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