Pharmaceutical Analytical Chemistry (PHCM223-SS16) Lecture...

Pharmaceutical Analytical Chemistry (PHCM223-SS16) Lecture 1

ACID- BASE EQUILIBRIUM-I “pH of acids and bases”

Dr. Rasha Hanafi

PHCM223,SS16 Lecture 1, Dr. Rasha Hanafi. 1

LEARNING OUTCOMES

By the end of this session the student should be able to:

1. Get acquainted with the course description.

2. Identify acids and bases.

3. Determine Acid strength.

4. Apply Kw Calculations.

5. Define the pH scale.

6. Determine the pH of weak/strong acids and bases.

7. Determine the pH of Polyprotic acids.


COURSE DESCRIPTION

• COURSE INSTRUCTORS: 1. Lecturers: Dr. Rasha Hanafi (Lectures# 1-6), B5-101. Dr. Nesrine Elgohary (Lectures# 7-12), B7-207. 2. Teaching assistants: Noura Essam (B7-205), Dina Elsaeed (B1-225), Engy Saad (B7-207), Nadine George (B7-207), Magy Maged (B7-207), Maha Farag (B7-207), Dalia Khalil (B1-207), Salma Mokbel. •TEXT BOOKS 1. Chemistry, 10th ed., Raymond Chang, ISBN 978-0-07-017264-7, McGraw Hill.

2. D. A. Skoog, D.A. West, F.J. Holler, S.R. Crouch, Analytical Chemistry, an introduction, 7th Edition, ISBN 0-03-020293-0.

• INTERNET LINK FOR COURSE MATERIAL, GENERAL INFO AND ANNOUNCEMENTS http://pbt.guc.edu.eg/Courses/Undergrad.aspx


Assessment methods Weight

Contribution in tutorials 10%

Quizzes (Best 2 quizzes out of 3, 10% each) 20%

Midterm exam 30%

Final term exam 40%

All Quizzes will take place during tutorial and will cover up to the previous tutorial session. Each Practical exam will have a theoretical quiz covering its content.

http://pbt.guc.edu.eg/Courses/Undergrad.aspx


Name Room number Office hours (slots)

Dr. Rasha Hanafi B5-101 Sunday 3rd and Monday 3rd

Dr. Nesrine Elgohary B7-207 Tuesday 1st and Thursday 3rd

Salma Mokbel B7-207 Sunday 5th and Tuesday 1st

Engy Saad B7-207 Thursday 3rd and 4th

Dalia Elhelw B1-205 Monday 5th and Wednesday 3rd

Magy Maged B7-207 Monday 3rd and 4th

Nadine George B7-207 Monday 2nd and 4th

Noura Essam B7-205 Tuesday 3rd and Wednesday 3rd

Dina Atef B1-225 Monday 3rd and 4th

Maha Magdy B7-207 Monday 4th and Wednesday 3rd

COURSE INSTRUCTORS

THEORETICAL COURSE CALENDAR

Lecture 1, Dr. Rasha Hanafi.

Lecture no. [instructor]

Week Title of lecture Tutorials Practical sessions Compensations

Lecture 1 [RH] 13/2- 18/2 Acid- Base Equilibrium (ABE)

pH of Acids/bases --

Lecture 2 [RH] 20/2- 25/2 ABE

pH of Salts pH of Acids/bases

Introduction and Acid base titration


pH of Buffers pH of Salts Acid base titration


Titration Curves pH of Buffers Acid base titration


Polyprotic acids+ Double indicators Titration Curves Exam 1


Precipitimertry Polyprotic acids+ Double

indicators Precipitimertry

No teaching 26/3- 31/3

Revision week

No teaching 2/4- 7/4 Midterm exams

Lecture7 [NE] 9/4- 14/4 Complexometry 1 Precipitimertry Complexometry

Lecture8 [NE] 16/4- 21/4 Complexometry 2 Complexometry 1 Complexometry

Lecture 9 [NE] 23/4 - 28/4 Redox 1 Complexometry 2 Exam 2 for 24th 25th and 28th of

April

Lecture 10 [NE] 30/4 - 5/5 Redox 2 Redox 1 Exam 2, cont.

Lecture 11[NE] 7/5 - 12/5 Gravimetry Redox 2 Redox for 1st and 2nd of May

Lecture 12 [NE] 14/5 - 19/5 Applications on different types of equilibria

Gravimetry Exam 3

No teaching Revision week

PHCM223,SS16 5


pH in living systems Compartment pH

Gastric Acid 1

Human Skin 5.5

Water 7

Blood 7.35-7.45

Mouth pH 5-7

Small intestine

pH 6-7.5

Large intestine

pH 5-7

Stomach

pH 1-3


Acid-Base Chemistry

Reversible reaction

Equilibrium

Proton Exchange

Donates protons

Acids

Accepts protons

Bases

strong

weak

Neutralization reaction

Titration

]reactants

products[k

AMPHOTERIC NATURE OF WATER


Aqueous acid-base reactions occur in water. Water can behave as an acid or as a base, by donating a proton (H+) or accepting a proton, respectively.

CH3COOH + H2O H3O+ (aq) + CH3COO-

(aq) acid

acid base

base

conjugate conjugate

EQUILIBRIUM CONSTANT OF ACIDS, Ka


HA + H2O H3O+ (aq) + A-

(aq)

acid Conjugate acid base Conjugate base

][

]][[

]][[

]][[constant mequilibriu The

2

3

HA

AH

OHHA

AOHka


ACIDS

Classified by the equilibrium position of its dissociation reaction

STRONG

equilibrium lies far to the right, i.e. almost all HA is

dissociated (ionized).

yields a weak conjugate base (that has low affinity

to H+).

Ex: HCl, H2SO4, HNO3, HClO4

WEAK

equilibrium lies far to the left. HA dissociates only to

a small extent (<1%).

Yields a strong conjugate base (that has high affinity to

H+).

H3PO4, CH3COOH, C6H5COOH

AUTOIONIZATION OF WATER

The equilibrium expression of water is

Kw = [H+][OH-]= 1.0x10-14 at 25°C

Kw is “the ion-product constant” or “dissociation constant for water”.


-log Kw= -log [H+] – log[OH-]

Kw depends on T: At 0C, Kw = 0.114x10-14 At 50C, Kw= 5.47x10-14 At 100C, Kw = 49x10-14

In aqueous solutions, if [H+] = [OH-] Neutral soln. [H+] > [OH-] Acidic soln. [H+] < [OH-] Basic soln.

The term –log is commonly referred to “p”, so

pKw = pH + pOH Since Kw = 1.0x10-14

pH + pOH = 14.00

H2O + H2O OH- + H3O+

Test yourself

Since [OH-] and [H+] are formed only from the dissociation of water (there is no acid nor base in the solution), hence, the concentrations of [OH-] and [H+] must be equal: [H+]=[OH-]=

wK


Calculate the hydrogen and hydroxide ion concentrations of pure water at 25C and 100C.

At 25C, Kw= [OH-][H+] = 1.00x10-14

So, [OH-]=[H+] = (1.00x10-14)1/2 = 1.00x10-7M

THE pH SCALE

pH = - log[H+] Thus for solution of [H+] = 1.0x10-7, pH = -log(10-7) = 7.00 (2nd decimal is a significant figure). N.B.: A solution of pH 3 has a [H+]concentration 10 times more than that of a solution of pH 4 and 100 times more than that of a solution of pH 5.


The pH scale is a compact way to represent acidity of solutions.


Test yourself

Calculate [H+] or [OH-] as required for each of the following solutions at 25°C and state whether the solution is neutral, acidic or basic.

1.0x10-5 M OH-

[H+] =1.0x10-14/ 1.0x10-5 = 1.0x10-9M

[OH-] > [H+]

The soln. is basic.

1.0x10-7 M OH-

[H+] = 1.0x10-14/ 1.0x10-7 = 1.0x10-7M

[H+] = [OH-]

The soln. is neutral

10.0 M H+

[OH-] = 1.0x10-14/ 10.0 = 1.0x10-15M

[H+] > [OH-]

The solution is acidic.

Kw = [H+][OH-]= 1.0x10-14

pH OF WEAK ACIDIC SOLUTIONS


STRONG ACIDS WEAK ACIDS

Fully ionize in water Less than 1% ionization

HCl H+ + Cl-

nM n M n M

CH3COOH CH3COO- + H+

(n – x) M x M x M

pH = -log[H+]= -log[n] pH = -log[H+]= - log[x]

It is easy to get [H+] because it is the same concentration of the acid

To get [H+], not only the concentration of the acid needs to be known, but

how much exactly dissociates from the acid, i.e. the dissociation constant (ka)

needs also to be known

][

]][[

3

3

COOHCH

COOCHHka

x2 = (7.2x10-4)(1.00)

x= 2.7x10-2 = [H+]

pH = -log(2.7x10-2) = 1.57

pH of weak acid = -log

] [ acidweakKa


pH OF WEAK ACIDIC SOLUTIONS

00.1

.

][

]][[ xx

HF

FHka

Example: to calculate the pH of 1.00M soln. of the weak acid HF, the dissociation constant needs to be known. It is Ka = 7.2x10-4.

HF H+ + F-

(1-x) M x M x M

(1-x) ≈1 because it is < 1% of x, thus


BASES

Classified by the equilibrium position of its dissociation reaction

STRONG

equilibrium lies far to the right, i.e. almost all the base is dissociated

(ionized).

yields a weak conjugate acid (that has low ability to donate H+).

All hydroxides of group 1A (LiOH, NaOH, KOH, …etc) and 2A alkaline earth

hydroxides (Ca(OH)2, Mg(OH)2) are strong bases.

NaOH and KOH are common lab reagents.

WEAK

equilibrium lies far to the left. The base dissociates only to a small

extent (<1%).

Yields a strong conjugate acid

(that has ability to donate H+).

Ammonia (NH3), pyridine

Calculate the pH for a 15.0 M NH3 (Kb = 1.8x10-5)

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


pH OF WEAK BASIC SOLUTIONS

5-

3

x108.100.15

.

][

]][[4

xx

NH

OHNHkb

x = [OH-] = 1.6x10-2 M [H+] = Kw/[OH-] = 1.0x10-14/1.6x10-2 = 6.3x10-13 M pH = -log(6.3x10-13) = 12.20 N.B.: pOH can be calculated directly from [OH-] and then subtract it from 14.

Polyprotic acids (ex: H2SO4, H3PO4…etc) dissociate in a stepwise manner, yielding one proton at a time. The successive Ka values for the dissociation equilibria are designated Ka1 and Ka2.

Carbonic acid (important in maintaining blood pH at 7.35)


pH OF POLYPROTIC ACIDIC SOLUTIONS

The conjugate base HCO3- of the first dissociation equilibrium

becomes the acid in the second step.

STEPWISE DISSOCIATION CONSTANTS FOR SOME COMMON POLYPROTIC ACIDS


For weak polyprotic acids, Ka1>Ka2>Ka3 and the relative acid strengths in case of H3PO4 for example would be H3PO4» H2PO4

-»HPO42-.

This means that in a solution prepared by dissolving H3PO4 in water, only the first dissociation step makes an important contribution to [H+] and this greatly simplifies the pH calculations for polyprotic acids in general.

pH= -log ][acidKa

REFERENCES

1. S.S. Zumdahl, S.A. Zumdahl, Chemistry 6th Ed., Houghton Mifflin Company, ISBN 0-618-22156-5, (Chapters 14, 15)

2. D. A. Skoog, D.A. West, F.J. Holler, S.R. Crouch, Analytical Chemistry, an introduction, 7th Edition, ISBN 0-03-020293-0. (Chapters 4, 8, 10, 11, 12, 13, 14, 15, 18)

3. Lecture 1, PHCM223, by Prof. Rasha Elnashar, GUC, SS 2014.

4. Introduction to acid-base chemistry (http://www.chem1.com/acad/pdf/c1xacid1.pdf)


http://www.chem1.com/acad/pdf/c1xacid1.pdf

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