The document discusses pH and acid-base chemistry, including the definitions of acids and bases, acid-base reactions, pH and pKa scales, autoionization of water, acid and base strength, and methods for measuring pH. Common acid-base concepts like conjugate acid-base pairs, strong vs weak acids, and calculations involving hydronium and hydroxide ion concentrations are explained. The properties and types of solutions such as concentration, solubility, dilution, and factors that affect solubility are also covered.

1. 1
pH and solutions
1

2. 2
 In aqueous chemistry, an acid is a substance that increases the concentration of
H3O+ (hydronium ion) when added to water.
 Conversely, a base decreases the concentration of H3O+.
 Brønsted-Lowry acid: proton donor
 Brønsted-Lowry base: proton acceptor
e.g.
 A Lewis acid is a substance that can accept a pair of electrons to form a
covalent bond.
 A Lewis base is a substance that can donate a pair of electrons
e.g.

3. 3
Conjugate Acids and Bases
 Conjugate acids and bases are related to each other by the gain or loss of one H+
The Nature of H+ and OH-
 The proton does not exist by itself in water.
 The simplest formula found in some crystalline salts is H3O+
 We will ordinarily write H+ in most chemical equations, although we really mean
H3O+

4. 4
Strong and Weak Acids
 A strong acid dissociates completely into ions in water:
 HA(g or l) + H2O (l) → H3O+(aq) + A-(aq)
 A dilute solution of a strong acid contains no HA molecules
 A weak acid dissociates slightly to form ions in water:
HA(aq) + H2O(l) H3O+(aq) + A-(aq)
 In a dilute solution of a weak acid, most HA molecules are undissociated.
 The Acid Dissociation Constant, Ka of HA(aq) + H2O(l) H3O+(aq) + A–(aq)
 The value of Ka is an indication of acid strength
[H3O+][A-]
[HA][H2O]
Kc = has a very small value.
Kc[H2O] = Ka =
[H3O+][A–]
[HA]
[H3O+][A–]
[HA][H2O]
Kc =
Stronger acid larger Ka
higher [H3O+]
Weaker acid smaller Ka
lower % dissociation of HA

5. 5
The Relationship between Ka and pKa
 pKa = –logKa
Example
Acid Name (Formula) Ka at 25°C pKa
Hydrogen sulfate ion (HSO4
-) 1.0x10–2 1.99
Nitrous acid (HNO2) 7.1x10–4 3.15
Acetic acid (CH3COOH) 1.8x10–5 4.75
Hypobromous acid (HBrO) 2.3x10–9 8.64
Phenol (C6H5OH) 1.0x10–10 10.00
 A low pKa corresponds to a high Ka.

6. 6
Autoprotolysis
 Water undergoes self-ionization, called autoprotolysis, in which it
acts as both an acid and a base
 the collisions between water molecules are energetic enough for a
reaction to occur.
 The autoprotolysis constant for H2O has the special symbol Kw,
where “w” stands for water:
 Autopprolysis

7. 7
 The value Kw = 1.0 x 10-14 at 250C
 Any aqueous solution in which [H+] and [OH−] are equal is a neutral solution.
Example. Calculate the concentrations of [H+] and [OH−] in pure water at 250C
Solution: Molar ratio 1:1 , then their concentrations must be equal
 The concentrations of [H+] and [OH−] are both 1.0 10-7 M in pure water
 What is the concentration of [OH−] if [H+] 1.0 10-3 M? (From now on, assume that
the temperature is 250C unless otherwise stated.)

8. 8
 pH is the negative logarithm of the H+ concentration
 where pOH = - log[OH-], just as pH = –log[H+].
 At 250C, an acidic solution has a pH below 7 and
 a basic solution has a pH above 7
In an acidic solution, [H3O+] > [OH–]
In a neutral solution, [H3O+] = [OH–]
In a basic solution, [H3O+] < [OH–]

9. 9
 What is the relationship b/n pH, [H+] and [OH-]?
When [H+] is
given in the
format 1 × 10–n,
it’s easy to find
the pH. It’s just
the absolute
value of the
exponent n. Also,
note that [H+] ×
[OH–] always
equals 1 × 10–
14.

10. 10
Example. Calculate the hydronium and hydroxide ion concentrations
and the pH and pOH of 0.200 M aqueous NaOH at 25°C.

11. 11
 Questions
 The pH of an unknown solution is 6.35. What is the hydrogen-ion concentration of
the solution?
 What is the pH of a solution if [OH−] = 4.0 × 10−11M?
Measuring pH
 Either acid-base indicators or pH meters can be used to measure pH.
 Acid - Base indicators (also known as pH indicators) are substances which
change color with pH.
 They are usually weak acids or bases, which when dissolved in water dissociate
slightly and form ions.
 An indicator (HIn) is an acid or a base that dissociates in a known pH range.
 Indicators work because their acid form and base form have different colors in
solution

12. 12
 The acid form of the indicator (HIn) is dominant at low pH and
high [H+].
 The base form (In−) is dominant at high pH and high [OH−].
 The change from dominating acid form to dominating base form
occurs within a narrow range of about two pH units.
 At all pH values below the range, you would see only the color of
the acid form.
 At all pH values above this range, you would see only the color of
the base form.

13. 13
Colorless (Acid) pink (Base)
 Phenolphthalein is a colorless, weak acid which dissociates in water
forming pink anions.
 Under acidic conditions, the equilibrium is to the left and the concentration
of the anions is too low for the pink color to be observed.
 However, under alkaline conditions, the equilibrium is to the right, and the
concentration of the anion becomes sufficient for the pink color to be
observed.

14. 14
A pH meter is used to make rapid, continuous measurements of pH.
 The measurements of pH obtained with a pH meter are typically accurate to
within 0.01 pH unit of the true pH.
 If the pH meter is connected to a computer or chart recorder, the user will
have a record of the pH changes
 A pH meter can be easier to use than liquid
indicators or indicator strips.

15. 15
TYPES OF SOLUTIONS
Why substances dissolve?
 A solution is a homogeneous mixture of two substances: a solute and
a solvent.
 Solute: substance being dissolved; present in lesser amount.
 Solvent: substance doing the dissolving; present in larger amount.
 Solutes and solvents may be of any form of matter: solid, liquid or
gas

16. 16
 Solutions form between solute and solvent molecules because of
similarities between them. (Like dissolves Like)
 Ionic solids dissolve in water because the charged ions (polar) are
attracted to the polar water molecules.
 Non-polar molecules such as oil and grease dissolve in non-polar
solvents such as kerosene.
 Solubility refers to the maximum amount of solute that can be
dissolved in a given amount of solvent.
 Many factors such as type of solute, type of solvent and temperature affect
the solubility of a solute in a solution.

17. 17
 Solubility of most solids in water increases as temperature increases
 Solubility of gases in water decreases as temperature increases
 At higher temperatures more gas molecules have the energy to escape from
solution.
 Saturated: it contains the maximum amount of dissolved solute at a given
temperature.
 Unsaturated solution that contains less
concentration of dissolved solute
 Unsaturated .a solution that contains more
than the equilibrium concentration
of dissolved solute
The
Properties
of
Solutions

18. 18
Concentration
 Concentration is a general measurement unit stating the amount of
solute present in a known amount of solution
 Expressing concentration of solutions:
Molarity (M)
 molarity express concentration as moles of solute per liter of solution.
 The most common unit of concentration used in the laboratory

19. 19
 Common Units for Reporting Concentration

20. 20
 What is the molarity of a solution prepared by dissolving 60.0 g of NaOH in 0.250
L of solution?
Solution:
 What is the mass % (m/m) of a solution prepared by dissolving 30.0 g of NaOH in
120.0 g of water?
Solution:
30g
150g
x100 = 20%

21. 21
 Weight percent (% w/w), volume percent (% v/v) and weight-to-
volume percent (% w/v) express concentration as units of solute per
100 units of sample.
 A solution in which a solute has a concentration of 23% w/v contains
23 g of solute per 100 mL of solution
 Parts per million (ppm) and parts per billion (ppb) are mass ratios of
grams of solute to one million or one billion grams of sample,
respectively
 Question. A concentrated solution of aqueous ammonia is 28.0% w/w
NH3 and has a density of 0.899 g/mL. What is the molar concentration
of NH3 in this solution?

22. 22
Example 1: If there is 0.6 g of Pb present in 277 g of solution, what is the Pb
concentration in parts per thousand?
Solution: Use the definition of parts per thousand to determine the concentration.
Substituting:
2. The concentration of Cl– ion in a sample of H2O is 15.0 ppm. What mass of Cl–
ion is present in 240.0 mL of H2O, which has a density of 1.00 g/mL?
Solution First, use the density of H2O to determine the mass of the sample

23. 23
Dilution
 The process of preparing a less concentrated solution from a more concentrated
solution
 Since the total amount of solute is the same before and after dilution,
Co × Vo = Cd × Vd
Example. A laboratory procedure calls for 250 mL of an approximately 0.10 M
solution of NH3. Describe how you would prepare this solution using a stock solution
of concentrated NH3 (14.8 M).
Solution: substituting in the above equation
14.8 M × Vo = 0.10 M × 0.25 L= 1.69 × 10–3 L, or 1.7 mL
 What is serial dilution?

24. 24
 serial dilution is the process of stepwise dilution of a solution with an
associated dilution factor.
 The objective of the serial dilution method is to estimate the
concentration (number of organisms, bacteria, viruses, or colonies)
of an unknown sample by enumeration of the number of colonies
cultured from serial dilutions of the sample.
 Example: prepare a serial dilution of (0.1,0.25, 0.5,1 and 1.5M)
from 2M NaOH solution.