This document summarizes key concepts about acids, bases, pH and buffers. It defines acids as substances that can donate protons and discusses strong acids like HCl that fully dissociate in water versus weak acids like acetic acid that only partially dissociate. Bases are defined as substances that can accept protons. Salts are formed when the hydrogen of an acid is replaced by a metal. The pH scale measures the concentration of hydrogen ions from 0-14, with lower values being more acidic and higher more basic. Buffers resist changes in pH using mixtures of weak acids and their salts. Diffusion and osmosis are also summarized as processes by which substances move across membranes down concentration gradients.

1. Chemistry of Life
Lecture 3
Prepared by IJn

2. Acids, bases, pH and buffers
H+
Acids
A hydrogen atom consists of one electron and one proton.
Loss of electron Proton – a hydrogen ion, H+
An acid is a substance which can act as a proton donor.
or,
An acid is defined as a substance which ionizes in water to give H+ ions as the cation
(positive ion).
A strong acid Hydrochloric acid (undergoes almost complete
dissociation)
HCl + Cl-
A weak acid Acetic acid (a small proportion of the acid dissociates to
give hydrogen ions)
CH3COOH CH3COO- + H+

3. Acids
• Strong Acid = pH 1-3
• High in H+
ions
• Lower number of OH-
ions

4. Bases
• Substance that can accept protons, so they can raise the
pH of fluids and make them basic, or alkaline.
Examples: NaOH, KOH etc
• Strong Base = pH 11 – 14
• High in OH-
ions
• Lower in number of H+
ions

5. Salts
A salt is a substance in which replaceable hydrogen of an
acid has been partly or completely replaced by a metal.
For example, Sodium Chloride (NaCl) where the hydrogen
atom of hydrochloric acid has been replaced by an atom of
sodium.
HCl + NaOH NaCl + H2O
(Salt)

6. Acids, Bases and Salts
a) In water, Hydrochloric acid (HCl) dissociates into H+ and Cl-.
b) Sodium hydroxide (NaOH) a base, dissociates into OH- and Na+ in water.
c) In water, Table salt (NaCl) dissociates into positive ion (Na+) and negative ion
(Cl-), neither of which are H+ or OH-.

7. pH
The pH is defined as the negative logarithm to the base 10 of the
hydrogen ion concentration; determined in moles per liter [H+].
Pure water contains 1x107 moles of hydrogen ions per liter. The
pH of water is therefore,
pH = -log1010-7 = 7
• pH is a measure of the concentration of hydrogen ions in a
solution.

8.  pH = 7 : A neutral solution (H+ and OH- are equal)
 pH < 7 : an acidic solution (H+ concentration> OH-
concentration)
 pH > 7 : an alkaline solution (OH- concentration > H+
concentration)
• The pH scale ranges from 0 - 14
• A change of one whole number represents a tenfold change
from the previous concentration.
• A solution of pH 1 has 10 times more H+ than a solution of
pH 2 and 100 times more H+ ions than a solution of pH 3.

9. pH scale

10. The pH of few common substances
Approximate pH Common Examples
Strong Acids 0-2
Stomach acid (HCl),
battery acid (H2SO4)
Weak Acids 3-6 Lemon juice, vinegar
Neutral 7 Pure water
Weak Bases 8-11 Bicarbonate solution
Strong Bases 12-14 Solutions of NaOH, KOH
Human blood pH is 7.4 – Mild Basic
Water pH is 7.0 -- Neutral
Gastric juice pH is 2.0 --- Strong acid

11. Buffers
Living organisms Takes up
nutrients
Excretes waste
Chemical reactions
Balance of acids and bases change
Change of pH
Living organisms overcome this adverse effect by means of pH buffers
A buffer solution is a solution containing a mixture of a weak acid and its soluble
salt. It acts to resist changes in pH. Such changes can be brought about by dilution
or addition of acid or alkali.
Increased acidity More H+
Free anion (negative ion)
from salt
Removal of H+ from
solution
Drop in pH

12. At increased alkalinity
Decrease in acidity Tendency to release
hydrogen ions
Thus buffer solution tends to maintain a constant, balanced hydrogen ion concentration
Example,
NaHCO3 Na+ + HCO3
-
Sodium hydrogen
carbonate
Sodium ion Hydrogen carbonate ion
HCO3
-
Hydrogen carbonate ion
+ H+ H2CO3
Carbonic acidHydrogen Ion
(removal of hydrogen
ions from the solution)
Lowering solution’s acidity
HCO3
-
Hydrogen carbonate ion
+ OH-
CO3
2- + H2O
(removal of Hydroxyl ions
from the solution)
Lowering solution’s
alkalinity

13. Diffusion
• The difference in concentration of a substance between two areas is called a
concentration gradient.
• Particles move down a concentration gradient by diffusion, until they are
spread evenly.
• Diffusion is a passive process: it requires no input of energy.

14. Examples of diffusion across concentration gradients in organism
Place Particles move From To
Gut Digested Gut cavity Blood in capillary of villus
Food products
Lungs Oxygen Alveolar air space Blood circulating around the lungs

15. Diffusion in water

16. What affects the rate of diffusion?
• Concentration gradient: The greater the difference in the
concentration of a substance in two areas, the faster the rate
of diffusion
The rate of diffusion is
directly proportional to
the concentration
gradient

17. The larger the surface area the higher the rate of diffusion (e.g.
in gases diffusing into/out of leaves)

18. • Osmosis is the diffusion
of water.
• It is the net movement
of solvent (water)
molecules from a
region of their higher
concentration to a
region of their lower
concentration, through
a partially permeable
membrane.
Osmosis