General Biology. LECTURE 2. THE CHEMICAL FOUNDATION OF LIFE.pptx

Albia Dugger • Miami Dade College
LECTURE 2
THE CHEMICAL
FOUNDATIONOF LIFE
IFE SYVIA C. ANAETOR
B.PHARM, MPH, MSC.Bio
Ianaetor.edu.@gmail.com
+220 7133704, 3662100

THE CHEMISTRY OF LIFE
In other words;
• The Chemistry of The Cell
• The Biochemistry of Life
What is Biochemistry?
• Biochemistry is the study of the structures and roles of the chemical
substances and vital processes occurring in living organisms.
• Biochemists focus heavily on the role and structure of biochemical.
• He studies the chemistry behind biological processes and the
synthesis of biologically active molecules.

Why should we study Biochemistry?
• Life depends on Chemistry
• Living things are composed of chemical compounds,
atoms, ions, and many molecules.
• Life processes depend on the organization and
interaction of these chemical units.
EXAMPLE?
- INGESTION > DIGESTION > ABSORPTION via chemical
reactions to give energy and keep the body alive
- inhalation and exhalation undergoes series of chemical
reactions.
Therefore, the first task of a biologist is to understand the
building blocks of life through study of chemistry of life.

• From the evolution of life which suggested living organisms
to be made up of chemical substances.
• Confirmed after a lot of structural biochemistry research
that all living organisms, from microbes to animals,
• are composed of chemicals.
• Life/organisms are composed of matter.
• What is matter?
Any substance that has mass and takes up space is defined
as matter.
- what is the basic unit of matter?

CONSTITUENTS OF LIVING ORGANISMS
• All matter is composed of extremely small particles
called atoms.
• An atom is the smallest unit of chemical element that still
has the characteristics property of that element.
• The study of chemistry begins with the basic unit of
matter; atom.
• Atoms are incredibly small unit of any substance.
• Despite, its extremely small size, an atom contains
subatomic particles
• Proton > neutron > electron

• At the center of each atom is a small, very dense nucleus consisting of
two other kinds of subatomic particles, protons and neutrons,
except hydrogen atom, the smallest atom, which has only one proton and
no neutrons in its nucleus.
• Oxygen, for example, has eight protons and eight neutrons in its
nucleus.
• The nucleus, the cluster of protons and neutrons is held together by a
force that works only over short subatomic distances.
• Each proton carries a positive (+) charge.
• Neutrons, as their name implies, possess no charge.

• Typically, an atom has one electron for each proton. The number of
electrons outside the nucleus equal the number of protons within the
nucleus.
• The number of protons (the atom’s atomic number) determines the
chemical character of the atom, WHY?
• because it dictates the number of electrons orbiting the nucleus which
are available for chemical activity.
• Every atom possesses an orbiting cloud of tiny subatomic particles
called electrons which spin around the nucleus a far distance away
from the nucleus.
• Electrons are found in orbits or shells at different distances around
the nucleus. Each electron has a negative ( - ) electric charge.

Atomic Structure
•The weight or mass of the atom is the sum of the number of
protons and neutrons in the nucleus.

ATOMIC NUMBER
• All atoms of the same elements have the same
number of protons in their nuclei., THUS THE SAME
ATOMIC NUMBER
• The general symbol is Z, placed at the bottom left of
any symbol.
• A number of elements have an equal number of
protons and neutrons in their nuclei.
• E.g., Oxygen: Z=8, neutrons=8
• Nitrogen: Z=7, neutrons=7
• Neon: Z 10, nuetrons=10

ATOMIC MASS/ MASS NUMBER; (A)
Atomic mass is the approximation of the total mass of an atom.
In other words, the sum of proton and neutron numbers.
The difference btw the mass number and atomic number is the neutron
number.
11
23Na : 11 protons 11 electrons, and 12 neutrons
The contribution of electron(s) to mass number is quite negligible.
WHY?
Because neutrons and protons have a mass very close to 1 dalton while
electrons have a very much smaller mass of approximately 0.0005 amu.
EX.:
• Write the mass number and the symbol of an atom containing
• 58 protons, 58 electrons, and 78 neutrons?
• 17 protons, 20 neutrons, and 17 electrons?
• Cerium= 58
136Ce
• Chlorine= 17
37Cl

The Chemistry of Life
•A substance that cannot be split into simpler substances by
any chemical means and consists entirely of the same type
of atom is called an chemical element.
•Of all the naturally existing elements (over 100 of them),
•only about 25(26%) occur in living organisms.
• Among those present, the most common elements in living
things are FOUR; C, O, H, and N
that constitute about 96.3% of the total weight of living cells.
• in addition with phosphorus and sulfur make up more than
99%.
• The remaining 1% constitutes minor elements which are inorganic in nature
with various functions.
•Life requires about 25 chemical elements. These elements are called
essential elements of life.

Elements Occurring Naturally in the
Human Body
# Symbol Element Atomic Number Percentage of Human
Body Weight
1. O Oxygen 8 65.0
2. C Carbon 6 18.5
3. H Hydrogen 1 9.5
4. N Nitrogen 7 3.3
5. Ca Calcium 20 1.5
6. P Phosphorus 15 1.0
7. K Potassium 19 0.4
8. S Sulfur 16 0.3
9. Na Sodium 11 0.2
10. Cl Chlorine 17 0.2
11. Mg Magnesium 12 0.1

Elements Relative abundance (%)
Hydrogen 60
Oxygen 22.5
Carbon 10.5
Nitrogen 2.4
Calcium 0.22
Phosphorous 0.13
Sulphur 0.13
Potassium 0.03
Sodium 0.03
Magnesium 0.03
Hydrogen, oxygen, nitrogen, carbon, sulfur and phosphorous when in different
Combinations form various building blocks of macromolecules;
carbohydrates, proteins, lipids and nucleic acids which are organic in nature
Table : Constituents of the Human Body.

• The 14 elements exist as trace elements of less than 0.01%
and biological system use them in trace amounts:
• Boron (B) Chromium (Cr)
• Cobalt (Co) Copper (Cu)
• Fluorine (F) Iodine (I)
• Iron (Fe) Manganese (Mn)
• Molybdenum (Mo) Selenium (Se)
• Silicon (Si) Tin (Sn)
• Vanadium (V) Zinc (Zn)
They are also quite important as well as ‘BIG FOUR’,
that a deficiency in any of them will result in either
death, severe developmental abnormalities, or chronic
ailments.

• Iron in the body found in hemoglobin.
• Iron deficiency is marked by fatigue, infections.
• Zinc deficiency results in loss of appetite, lack of
alertness, slower than expected growth (retarded
growth).
• Iodine, essential part of thyroid hormones. It is one of
the very few trace elements that is required by humans and
other vertebrates, but not by other organisms such as
bacteria or plants.
• Fluorine reduces the incidence of dental cavities.
• Sulfur, an essential component of protein.
• Metal such as sodium, potassium ions essential for
the transmission of nerve impulses.
Importance of essential elements

• and
• Calcium is required for muscle contraction.
• Calcium aids in blood clotting.
• Calcium and phosphorous are metals essential for
• teeth and bone formation.
• Potassium regulate electrical balance of body fluids
• and ………………………..
• No other element can take place of any essential element.
• These metals when combined with other elements, molecules,
and compounds can become acids, base or salts.

Atoms form Molecules & Compounds
Two or more atoms combine chemically to form molecules;
a new particle. Atoms of the same elements (O2, H2)
A molecule Atoms of different elements (CO2, H2O)
e.g. a.) When two oxygen atoms combine chemically they
form a molecule of oxygen.
b.)when the atoms of different elements bond to form
molecule, they make completely new substance: compound
e.g: CO H O
• Atoms combine by chemical bonding to form molecules
and compounds.
• WHAT IS MOLECULE?
• WHAT IS CHEMICAL COMPOUND?

•A chemical compound consists of two or more different elements combined in
a fixed ratio.
•E.g. Water (H2O) is a chemical compound consisting of hydrogen and
oxygen atoms in a ratio of 2:1
• E.g:
•NaCl = Na, a metal that fizzes violently when it comes in contact with
water
Cl, a green poisonous gas.
• Vitamin C ……….. C6H7O6
• Aspirin …………… C9H8O4
• Rust ………………. FeO3
• Sand (silicon dioxide) ……… SiO2
• Baking soda( sodium hydrogen carbonate) …………… NaHCO3

•Hydroxide ……… (OH). Carbonate ………… (CO3),
nitrate ………… (NO3), Sulfate ……. (SO4)
•The formula of a compound tells you elements chemically
bonded to form a cpd.
•These elements in the cpd cannot be separated again by
any mechanical mean.
•Because compounds are new substances, they are very
different to the elements that make them.

• All atoms of a given element have the same
atomic number, but some of these atoms have
more number of neutrons than other atoms of
the same elements.
• What do you think this would result to?
• Different (greater or lower) mass numbers.
• In only a few instances (e.g; Al, F) do all atoms in a
given element have the same mass.
• What do we call this phenomenon?

Isotopes
• Isotopes are atoms of the same element, that vary only in the
number of neutrons in their nuclei but have the same atomic
number.
• E.g: (a.) hydrogen.
• The tiniest of all atoms
• Only has one electron in its lowest energy level
• Its energy level can be either empty of electrons, or filled with
electrons.
• Present as a gas in the atmosphere only in tiny amounts.
• Found in the water, sun, most of the stars.
• There is a planet mostly composed of hydrogen?
Jupiter

Isotopes of Hydrogen
• There are three isotopes of hydrogen
• H or 1H (protium) :
• The most prevalent hydrogen isotope (99.9%).
• It consists of one proton and one electron but no neutron.
• Does not have any neutrons in its tiny nucleus.
• Not typically found in its monoatomic form, but bonded with itself (H2)
or other elements.
• D or 2H (detrium):
• Consisting of one proton, one neutron, and one electron.
• Used extensively in organic chemistry to study chemical reactions,
• also used in vitamin research
• T or 3H (tritium):
• Consisting of one proton, two neutrons, and one electron.
• Tritium beta radiation does not penetrate the dead outer layer of the
skin.
• Therefore, only poses a health risk if inhaled, ingested, or absorbed

• (b.) Carbon
• There are three isotopes of carbon
• 12C 13C 14C
• ALL ISOTOPES OF THE SAME ELEMENTS, ARE THEY
CHEMICALLY IDENTICAL?
Chemically, all isotopes of the same elements are identical because they have
the same number of electrons
12C 13C 14C
are isotopes of each other and are all chemically identical
o When 14C decays,
one of its neutrons is
converted to a proton and an
electron. In this process, 14C
is converted to
14N, a different element.

Types of Isotopes
• Stable Isotopes
• Unstable isotopes
• Stable isotopes: Those isotopes that have a stable combination of
protons and neutrons, so their nuclei do not have tendency to decay and
lose particles. Non-radioactive form of atoms.
• E.g: Percent Abundance of Stable Isotopes Commonly Employed in Ecosystem
Studies
The most stable nucleus is Fe-56, having lowest mass per nucleon
Fe-58, Fe-56 are the most tightly bound nuclei
N-15, the most common stable isotope used in agriculture
12C 98.89 16O 99.763 H 99.9844 14N 99.64 32S 95.02 84Sr 0.56 39K 93.08 24Mg 78.8
13C 1.11 17O 0.0375 D 0.0156 15N 0.36 33S 0.75 86Sr 9.86 40K 0.0119 25Mg 10.15
18O 0.1995 34S 4.21 87Sr 7.00 41K 6.91 26Mg 11.06
• Why are some nuclei more stable than others?
• What element has the most stable isotopes?,
• What atom has the most stable nucleus?
• Which element is the most stable and why?

Radioactive Isotopes
• Some nuclei are more stable than others depends on the
balance of protons and neutrons in a nucleus.
• Too many neutrons or protons upset this balance disrupting
the binding energy from the strong nuclear forces making the
nucleus unstable.
• Noble gases are the most stable elements due to having
the maximum number of valence electrons their outer shell
can hold.
Unstable Isotopes, otherwise called radioactive Isotopes or
radioisotopes are the ones in which their nuclei decay
spontaneously, giving off particles and energy.
• E.g: 3H, 14C
• All non-natural or synthetic elements are radioactive
isotopes.

The radiation these isotopes give off can be dangerous,
sometimes, useful.
Assignment:
1. State the important biological and scientific uses of
radioisotopes.
Most biological fluids have pH values in the range 6–8.
2. Find the pH of most biological systems such as saliva,
CSF(cerebrospinal fluid), pancreatic juice, stomach juice,
urine, feces, blood, tears,

Atoms of two different elements can have same atomic mass.
True or false???
What is neucleon?
True. Isobars
Ex:
• sulfur (40S) >>>> 16p, 24n =40
• potassium (40K) >>>> 19p, 21n =40
• Calcium (40Ca) >>>> 20p, 20n = 40
• Chlorine (40Cl) >>>>> 17p, 23n = 40
• Argon (40Ar) >>>> 18p, 22n = 40
• Carbon (14C) and Nitrogen (14N)
• Iso=the same; bars= weight >>>> having the same mass
• Isobars are atoms of different elements with the same atomic mass, but
different atomic numbers.
• Neucleon is either proton or neutron
• Total number of nucleon=total number of protons and neutrons=atomic
mass

Define Acid and Base
What element is common to all acids?
• An acid is a substance that increases the hydrogen ion
concentration in a solution.
• Ex: When hydrochloric acid is added to water, hydrogen ions dissociate
from chloride ions: HCl  H+ + Cl−.
• This produces an acidic solution.
• Any substance that reduces the hydrogen ion concentration in a
solution is a base.
 Hydrochloric Acid (HCl)
 All acids contain [H+] in solution.
 Protons are not only found within atomic nucleic, they are also released into
medium whenever a hydrogen atom loses a shared electron.
 Bases don’t contain H+ but what …….. ?
 Sodium Hydroxide (NaOH)
The more OH- is in the base, the less space there is for H+ in solution
 A weak base contains few OH-…
ACIDS , BASES

• Some bases reduce the H+ concentration directly by accepting H+ conc.
• Ammonia (NH3) acts as a base when the nitrogen’s unshared electron
pair attracts a hydrogen ion from the solution, creating an ammonium ion
(NH4
+): NH3 + H+  NH4
+.
• ammonia is a relatively weak base.
• The double arrows in the reaction for ammonia indicate that the binding
and release of hydrogen ions are reversible reactions.
• At equilibrium, there is a fixed ratio of NH4+ to NH3.
• Other bases reduce the H+ concentration indirectly by dissociating to OH−,
which then combines with H+ to form water.
NaOH  Na+ + OH−;
OH− + H+  H2O
HCl and NaOH dissociate completely when mixed with water.
• Hydrochloric acid is a strong acid, and sodium hydroxide is a strong base.

The pH scale measures the H+ concentration of a.
In any aqueous solution at 25°C, the product of the H+ and OH- concentrations is
constant at 10-14.
This is written as:
[H+] [OH−] = 10-14.
 In such an equation, brackets ([H+] and [OH−]) indicate the molar concentration of
the enclosed substance.
o [H+] [OH−] = 10−14
o In a neutral solution at room temp (25oC),
o [H+] = 10−7 M and [OH−] = 10−7 M.
o Thus, 10-14 is the product of 10-7 x 10-7.
• Adding enough acid to a soln shifts the balance between H+ and OH− toward H+
and leads to an increase H+ (10-5M) and decrease in the OH− concentration by
an equivalent amount to 10-9M.
o If [H+] = 10−5 M, then [OH−] = 10−9 M.
o Note: 10-5 x 10-9 =10-14

 Solutions with
 more OH− than H+ are basic solutions;
 more H+ than OH- are acidic solutions;
 in which the concentrations of OH- and H+ are equal are neutral solutions.
• This constant r/ship expresses the behavior of acids and bases in aqueous soln.
• An acid not only adds H+ to a soln, but also removes OH- because of the tendency for
to combine with OH- to form H2O.
• A base has the opposite effect, increasing OH- conc but also reducing H+ by the
formation of H2O.
o Whenever concentration of either H+ or OH- in aqueous soln is known, the conc of
the other ion would be known.
o The H+ and OH− concentrations of solutions can vary by a factor of 100 trillion or
more.
o To express this variation more conveniently, the H+ and OH− concentrations are
typically expressed with the pH scale.

THE PH SCALE/PH VALUE
The pH scale was invented by the Danish chemist Soren
Sorensen to measure the acidity of beer in a brewery.
 It is measured on a scale of 0 to 14.
Measures the concentration of hydrogen ions in solution;
the value corresponds to the concentration of
hydrogen ions.
 The pH value of a substance is directly related to the ratio
of the hydrogen ion and hydroxyl ion concentrations.
 Low pH values correspond to high concentrations of
H+
 high pH values correspond to low concentrations of
H+
 If the H+ concentration is higher than OH- the material
is acidic.

pH is a unit of measure which describes the
degree of acidity or alkalinity (basic) of a solution
or substance.
pH of any solution is the measure of its [H+]
conc
The value of pH is the negative log of the
hydrogen ion concentration in moles per liter.
pH = -log [H+] ; where H+ = molar conc of
protons
pH = −log [H+] or [H+] = 10−pH
 In a neutral solution, [H+] = 10−7 M and pH = 7.
 The pH decreases as the H+ concentration increases.

 For example pure water H+ ion
concentration is 1 x 10 −7 = 7 M, therefore the
pH would then be 7.
 Each pH unit represents a tenfold difference in H+ and OH−
concentrations
 A small change in pH indicates a substantial change in H+ and OH−
concentrations
 A solution of pH 3 is not twice as acidic as a solution of pH 6 but a
thousand times more acidic.

•0 - 4 = strong acid
• 4.01 – 6.99 = weak acid
• 7 = neutral
• 7.01 – 10 = weak base
• 10.01 – 14 = strong base
 The more hydrogen ions, the stronger the acid. More
H+ → lower the pH = very acidic
 The lower the pH, the ……………… ??
 Opposite is the alkaline
• Few H+ → higher the pH = very basic

pH [H+] in moles per L [OH-] pOH
14 1 x 10-14 0.00000000000001 molar 1 x 10-0 0
13 1 x 10-13 0.0000000000001 molar 1 x 10-1 1
12 1 x 10-12 0.000000000001 molar 1 x 10-2 2
11 1 x 10-11 0.00000000001 molar 1 x 10-3 3
10 1 x 10-10 0.0000000001 molar 1 x 10-4 4
9 1 x 10-9 0.000000001 molar 1 x 10-5 5
8 1 x 10-8 0.00000001 molar 1 x 10-6 6
7 1 x 10-7 0.0000001 molar 1 x 10-7 7
6 1 x 10-6 0.000001 molar 1 x 10-8 8
5 1 x 10-5 0.00001 molar 1 x 10-9 9
4 1 x 10-4 0.0001 molar 1 x 10-10 10
3 1 x 10-3 0.001 molar 1 x 10-11 11
2 1 x 10-2 0.01 molar 1 x 10-12 12
1 1 x 10-1 0.1 molar 1 x 10-13 13
0 1 x 100 1 molar 1 x 10-14 14

•Each numeric change in pOH represents a factor of 10
• A pH of 6 is _______ times more acidic than a pH of 9
• A pOH of 9 is _______ times more basic than a pOH of 6
• A pH of 6 is __1000_____ times more acidic than a pH of 9
• A pH of 3 is _______ times more acidic than a pH of 5
•Tap water does not have neutral pH

pH is Temperature
dependent
ToC pH
0 7.12
10 7.06
20 7.02
25 7.0
30 6.99
40 6.97
pH of pure water decreases as
the temperature increases.
A word of warning!
If the pH falls as temperature
increases, does this mean that
water becomes more acidic at
higher temperatures? NO!
Remember a solution is acidic if there is an
excess of hydrogen ions over hydroxide
ions.
In the case of pure water, there are always
the same number of hydrogen ions and
hydroxide ions.
This means that the water is always neutral

ACTIVITY:
• Acid of a mixture is increased or reduced
to 2.0. what is the H+ conc of the mixture?

 Occasionally, a hydrogen atom participating in a hydrogen bond between two
water molecules shifts from one molecule to the other.
o The hydrogen atom leaves its electron behind and is transferred as a single
proton—a hydrogen ion (H+).
o The water molecule that lost the proton is now a hydroxide ion (OH−).
o The water molecule with the extra proton is now a hydronium ion (H3O+).
 A simplified way to view this process is to say that a water molecule dissociates
into a hydrogen ion and a hydroxide ion: H2O  H+ + OH.
o This reaction is reversible.
 At equilibrium, the concentration of water molecules greatly exceeds the
concentration of H+ and OH−.
o At equilibrium, the concentration of H+ or OH− is 10−7 M (at 25°C).
o In pure water, only one water molecule in every 554 million is dissociated.
o There is only one ten-millionth of a mole of hydrogen ions per liter of pure
water and an equal number of hydroxide ions.

IONIZATION
What is an ION?
• An ion is an electrically charged particle (atom or gp of
atoms), formed when an atom loses or gains one or more
electrons to form a stable outer shell.
• Conducts electricity.
• What happens to the number of electrons and protons of such
atoms?
• Are they equal?
• More or less electrons than protons.
Any ion can either be ……………….. Or ……………………….
• CATION
• ANION Define them!

• Cation is an ion with a positive charge, formed when an atom loses
electron(s) in a reaction
• Lesser electrons than protons
• Hydrogen and metals tend to form cations
• Their atoms have one, two, or three electrons on their outer shell,
• Thus easier to lose, leaving a stable shell underneath than gaining
at least five more.
• Anion is an ion with a negative charge,
• Formed when an atom gains electrons in a reaction.
• More electrons than protons.
• Non metals tend to form anions.
• Their atoms have 5, 6, or 7 electrons on their outer shell.
• Thus, easier to gain electrons to acquire a stable shell than losing
5.

• Hence, because of their electric charges, ions conduct
electricity
• Are very important in many life processes including;
• energy transformation,
• transmission of nerve impulses,
• muscle contractions.
• In medical usage, ions are electrolytes.
• Serum electrolytes level:
• the conc of separate ions (Na+, K+, Cl-, HCO3-) in
circulating blood.

• Metal such as sodium and potassium ions essential
for the transmission of nerve impulses.
• Calcium and phosphorous ions are essential for teeth
and bone formation.
• Conc of various electrolyte levels can be altered by
many diseases,
• in which loss of electrolytes from the body:
• Ex:. In excess vomiting (hyperemesis),
• Excess diarrhea
• Or accumulation of electrolytes
• As in renal failure.

• When electrolytes concs are severely diminished;
• they can be corrected by administering the appropriate
substance
• by mouth (orally),
• by intravenous (IV) drips.
• When excess of electrolytes exists, it may be removed by
hemodialysis using dialyser or artificial kidney,
• or by use of enema (qty of fluid infused into the rectum through a
tube passed into the anus.

REDOX REACTION.
OXIDATION AND REDUCTION REACTION
• What is Oxidation?
• What is Reduction?
• Oxidation and Reduction reactions occur when electrons are
transferred.
• Oxidation is gain of oxygen to a cpd, molecule.
• Loss of electron from atom or compound.
• Loss of H+
• Increased in the oxidation state.
• Reduction is loss of oxygen
• Gain of electron
• Reducing/ decrease in oxidation state of an atom, ions, or

• The molecule that is oxidized, loses an electron.
• The molecule that is reduced gains the electron lost by the oxidized molecule.
• Both oxidation and reduction occur simultaneously.
• Some elements (such as metals-Na, Mg, Fe) lose electrons more easily than
others
• (easily oxidized).
• Non metals elements (including N, O2, Cl) are more difficult to lose electrons
• (not easily oxidized).
• Oxidized molecule or cpd carries a charge.
• Example:
• When an Iron object undergoes oxidation, O2 steals e- from Fe
• Fe get oxidized and O2 becomes reduced
• The Fe is transformed because it has lost electrons (Fe3+).
• The result compd is iron III oxide/ ferric oxide (Fe2O3) or rust.

• It is important to minimize the exposure of iron to oxygen and moisture.
• Iron will continue to lose electrons to O2 as long O2 as is present.
• Oxidation also takes place in fruits and other substances such as
sealed/canned food or drink.
• Oxidation causes browning of fruit.
• When the flesh of fruit is exposed to O2 in the air,
• they become oxidized, causing them to break down and turn to
brown.
• Oxidized form of food or substance is different from un-oxidized form
which is unfortunately unappealing to eat, similar to rusting.
• Many ‘ superfoods’ are advertised containing antioxidants.
• What is antioxidant???

• An antioxidant is a compound that reduces the oxidation of
other compounds.
• Consuming antioxidants helps the body fight off the harmful
effects of oxidation.
• Antioxidants
• scavenge free radicals from the body cells,
• prevent or reduce the damage caused by oxidation.
• Thereby keeping cells and enzymes happy and healthy.
• Antioxidant preservatives also play a crucial role in food
production by increasing shelf life.

• Examples of antioxidants include:
• vitamins C and E
• Vitamin C, commonly known as ascorbic acid, is the most powerful water-
soluble antioxidant found in blood plasma
• Glutathione is the most powerful and important among the antioxidants the
body produces, helps combat free radicals.
• Your cells contain glutathione,
• a substance made from three amino acids: cysteine, glutamate, and
glycine.
• In other words, eating food like blackberries, blueberries, walnuts,
cranberries, raspberries, strawberries, red grapes, peaches, red currants,
figs, cherries, pears, guava, oranges, apricots, and carotenoids help keep
the body from looking like brown fruits.

No. INDICATOR pH range Colour in acidic pH Colour in basic pH
1 Phenophthalein 9.3-10.5 colourless pink
2 Methyl orange 3.1-4.6 red yellow
3 Bromophenol blue 3.0-4.6 yellow blue
4 Methyl red 4.4-6.2 Red yellow
5 Phenol red 6.8 – 8.4 yellow red
6 Litmus 4.5-8.3 red Blue
SOME IMPORTANT INDICATORS USED IN A CLINICAL LABORATORY

indicator pH range
a. Litmus 5 - 8
b. methyl orange 3.1 - 4.4
c. Phenolphthalein 8.3 - 10.0
What is the molarity of hydronium ions in a solution that has a pOH of 9.6?
pH + pOH = 14
pH + 9.6 = 14
pH = 4.4
Acidic
pH = -log[H3O+]
4.4 = -log[H3O+]
-4.4 = log[H3O+]
[H3O+] = 4.0 × 10-5 M
pH = -log[H3O+]
pOH = -log[OH-]
pH + pOH = 14
What is the pH of 0.050 M HNO3?
pH = -log[H3O+]
pH = -log[0.050]
pH = 1.3
Acidic or basic? Acidic

 Although the dissociation of water is reversible and statistically rare, it is very
important in the chemistry of life.
o Because hydrogen and hydroxide ions are very reactive, changes in their
concentrations can drastically affect the chemistry of a cell.
 Adding certain solutes, called acids and bases, disrupts the equilibrium and
modifies the concentrations of hydrogen and hydroxide ions.
 Controlling pH is important for maintaining homeostasis.
 What is homeostasis?
Homeostasis is a physiological process by which the internal systems of the body
(e.g; B.P, body temp, acid-base balance) are maintained at equilibrium, despite
external variations.
• Biological system (organs, tissues, cells) are protected from pH fluctuations.
• Acid-base balance/ homeostasis : maintenance of normal pH (proper balancing of
acid and bases) within the body systems
• Which system performs this task?

BUFFER SYSTEM
• A compound that reacts with free hydrogen ions or hydroxyl ions, thereby
resisting or minimizing changes in the conc. of H+ and OH- in a solution.
• Buffer solutions usually contain a weak acid together with its conjugate base,
which combine reversibly with H+.
• Buffers allow a relatively constant pH in biological fluids despite the addition or
subtraction of acids or bases.
• How do they do so?
• They do so by absorbing excess H+ in the solution or donating OH- to it when
depleted.
• The body produces more acids than bases
• Acids taken in with foods
• Acids produced by metabolism of lipids and proteins
• Cellular metabolism produces CO2.

There are several buffers or
mechanisms that contribute to pH
stability in the human body.
• Ways to control acids in the
body.

ControlofAcids.
Changes Against in H+ Concentration
3. Kidney Excretion
1.Chemicalacid-base
buffersystem
2. Respiratory mechanisms

• 1. Buffer systems
(first line of defense)
In the ECF, the main chemical buffers are bicarbonate and
plasma proteins.
i.) Bicarbonate: (NaHCO3) and carbonic acid (H2CO3)
• The bicarbonate buffering system is especially key, as CO2 can be shifted
through carbonic acid (H2CO3) to H+ and bicarbonate (HCO3
-)
• : H2CO3 ↔ HCO3
- + H+
• If more H+ is added to this solution, it simply shifts the equilibrium to the left,
absorbing H+,
• so the [H+] remains unchanged.
• If H+ is removed (e.g. by adding OH-) then the equilibrium shifts to the right,
• releasing H+ to keep the pH constant.

• The chemical equilibrium btw
carbonic acid and bicarbonate acts as
a pH regulator.

ii.) Phosphate buffer.
• Major intracellular buffer:
• H+ + HPO42- ↔ H2PO4-
• OH- + H2PO4- ↔ H2O + H2PO42-
iii.) Protein buffer
• Includes hemoglobin.
• Carboxyl group gives up H+
• Amino Group accepts H+

2. Respiratory mechanisms
(secondlineofdefenseinacid-base disturbances)
62
• Body pH canbeadjustedbychangingrate anddepth
of breathing
• Exhalation ofcarbon dioxide
• Powerful, butonlyworkswithvolatile acids
• Doesn’taffectfixedacidslike lactic acid
• CO2 + H20 ↔ H2CO3 ↔ H+ +HCO3
-
• Increase in ventilation eliminates CO2 from extracellular fluid,
by mass action, reduces the H+ concentration.
• Conversely, decreased ventilation increases CO2, thus also
increasing H+ concentration in the extracellular fluid.

3. Kidney excretion
63
• Can eliminatelargeamountsof acid.
• Can also excretebase.
• Can conserve andproducebicarbonate ions.
• Most effectiveregulatorof pH.
• If kidneysfail, pH balance fails.

Rates of correction
65
• Buffers function almost instantaneously
• Respiratory mechanisms takeseveral
minutesto hours
• Renal mechanisms maytakeseveral hours
to days

• The internal pH of most living cells is close 7
• The pH of the fluid within most cells in the human
body must generally be kept between 6.5 – 7.45.
• Most enzymes function only with narrow pH
ranges.
• Just a slight changes in pH (H+ conc) can produce major
disturbances.
• This results in harmful changes in blood or body tissue pH.
• Disturbing acid-base balance in the plasma or tissue
• Affecting the ionic state of biological molecules or affect biological

• Acid-base balance can also affect electrolytes (Na+, K+, Cl-)
• Can also affect hormones
• A pH imbalance in the mouth can lead to tooth decay.
• Accumulation of acid in the stomach beyond the protective
sphincter in the esophagus, can cause heart burn, a sour taste.
• Some can trigger cancer
• A person cannot survive the blood pH 7
• or < 6.8 or > 8.0 death occurs

• The plasma containing too much H+ (pH as low as < 7.35) is
called acidemia or acidosis.
• Containing too few H+, pH is too High > 7.5, called
alkalemia/alkalosis.
• Acidosis refers to disorders that lower cell/
tissue pH to < 7.35
• Acidemia refers to an arterial pH < 7.3
• Alkalosis refers to disorders that elevate
pH to > 7.45.
• Alkalemia refers to an arterial pH > 7.45.

• The Two Basic Causes Of Disturbance
Of The Body’s Acid-base Balance:
Respiratory Causes
Metabolic causes

• RespiratoryAcidosisandRespiratoryAlkalosis
• CO2 mostlycausesacidemia.
• ThearterialpartialpressureofCO2(PaCO2)isinverselyproportional
toalveolarventilation.
• Therelativelysmallchangesinventilationhaveaprofoundeffecton
H+conc.andpH.
• When thelungscouldnotremoveenoughofCO2producedbythe
body.
• ItresultstoaccumulationofCO2,whichcauses thepHofbloodand
otherbodilyfluidstodecrease,
• Makingthemtooacidic,termedRespiratoryacidosis.
• WhatisRESPIRATORYACIDOSIS????
RESPIRATORY CAUSES

• Respiratoryacidosis isamedicalemergencyin
which decreasedventilationcausesincreasedblood
carbon dioxideconcentrationandultimatelyleadsto
decreaseinthepHlevel.
• DuringAlveolarhypoventilationthereisan
increasein CO2thusleadstoanincreased
PaCO2.
• ietosay;PaCO2isabovethelimitof reference
range >6kpa(45mmHg).
• Hypercapnia– highlevelsofpCO2 in blood
Respiratory Acidosis

TYPES OF RESPIRATORY ACIDOSIS
14-03-
7
ACUTE RESPIRATORYACIDOSIS
CHRONIC RESPIRATORY ACIDOSIS

ACUTE RESPIRATORY ACIDOSIS
• occurs whenanabrupt failure of ventilationoccurs
• failure inventilation maybe causedby depression of the
central respiratory center (respiratorycenter inthebrainthatcontrols
breathingrate)– drugsorhead trauma.
• inability to ventilate adequatelydueto
• neuromusculardisease, or
• airway obstruction related to asthma,bronchitis,sleep
apneaor
• COPD (chronic obstructive pulmonarydisease) exacerbation.
• Adult Respiratory Distress Syndrome
• Pulmonary edema
• Pneumothorax

CHRONIC RESPIRATORY ACIDOSIS
• Chronic respiratory acidosis maybesecondary to
• manydisorders, includingCOPD
• Chronic respiratory acidosis alsomaybesecondaryto
• Pickwickiansyndrome,
• neuromusculardisorders, and
• severe restrictive ventilatory defects asobservedininterstitial
fibrosisandthoracicdeformities.
• Paralysis ofrespiratoryorchest muscles
• Asthma,
• Pneumonia,
• Emphysema

Acid-Base Imbalances
• Disturbanceof thebody’sacid-basebalancewillresult
incompensatorymechanism.
The body response to acid-base imbalance
is called compensation.
• Maybecompleteifbroughtbackwithin normallimits.
• Partialcompensationifrangeis still outsidenorms.
• Ifunderlyingproblemismetabolic, hyperventilationor
hypoventilationcan help:respiratorycompensation.

CompensationMechanismforRespiratory Acidosis
• Acute respiratory acidosis, compensationoccurs in 2 steps:
• The initial responseis cellular buffering that occurs
over minutesto hours.
• Cellular buffering elevates plasmabicarbonate
(HCO3
−)onlyslightly,
• approximately 1 mEq/L for each 10-mm Hg increase
in PaCO2.

• The second step is renal compensation
• A mechanism by which the kidneys can regulate the plasma
pH.
• It is slower than respiratory compensation, occurs over 3–5
days.
• But has a greater ability to restore normal values.
• With renal compensation,
• Kidney produces and excretes ammonium (NH4+) and
monophosphate
• Generating and retaining bicarbonate ion in the process
• Eliminate hydrogen ion while clearing acid..
CompensationMechanismforRespiratory Acidosis

Signs andSymptomsofRespiratory Acidosis
80
• Breathlessness
• Restlessness
• Lethargyanddisorientation
• Tremors,convulsions,coma
• Rapidrespiratoryrate,whichgraduallydepressed
• Warmandflushedskinduetovasodilation causedbyexcess
CO2
TreatmentofRespiratory Acidosis
• Restoreventilation
• IV lactatesolution
• Treatunderlyingdysfunctionor disease

Respiratory Alkalosis
• ↑respiration/↑O2→↓pCO2 (<35mmHg)and↑pH
• A medical condition in which increased respiration
elevates the blood pH beyond the normal range (7.35-
(7.35-7.45) with a concurrent reduction in arterial
levels of CO2 (hypocarpia).
• Most common acid-base imbalance
• This alters the dynamic chemical equilibrium of
carbon dioxide in the circulatory system,
• The system reacts according to Le Chatelier's
principle.
• Chatelier’s principle is also known as “The

• Chatelier’s principle states that when a system such as
temp, conc, pressure, experiences a disturbance, it will
respond to restore a new equilibrium state.
• i.e; a behavior of a system due to changes in temp, press,
conc

Types of respiratory alkalosis
 Acute Respiratory Alkalosis
 Chronic Respiratory Alkalosis
Acute Respiratory Alkalosis occurs rapidly
• For every 10 mmHg drop in PCO2 in arterial blood,
there is a corresponding 2 mEq/L drop in bicarbonate
ion due to acute compensation.
Chronic respiratory alkalosis
• More long-standing condition.
• For every 10 mmHg drop in PCO2 in arterial blood,
there is a corresponding 5 mEq/L drop in

• Primary cause is hyperventilation
• Pulmonary disease and Congestive heart failure – caused
by hypoxia.
• Drug use
Causesof Respiratory Alkalosis.
Conditions that stimulate respiratory center:

CompensationofRespiratory Alkalosis
85
• This is termed metabolic compensation.
• Byrenal system.
• Kidneysconservehydrogenion.
• Excretebicarbonateion.
• Less HCO − is reabsorbed, thus lowering the
pH.
TreatmentofRespiratory Alkalosis
• Treatunderlying cause
• IV Chloridecontainingsolution– Cl-ions replace
lostbicarbonate ions

Metabolic Acidosis
• A condition that occurs when the body produces
excessive quantities of acid or,
• When the kidneys are not removing enough acid
out from the body.
• Bicarbonate deficit- blood concentrations of bicarb drop
below 22mEq/L.
• On long standing (chronic situation), the body fails to
form sufficient bicarbonate in the kidney
• If untreated it leads to Acidemia/Metabolicacidosis.
• The potential consequences are so serious leading to
coma and death.

Causes:–
• Loss of excessive alkali (bicarbonate) from the body through
diarrhea or renal dysfunction.
• Formation of excessive quantities of metabolic acid in the
body.
• Failure of kidneys to excrete H+.
• DM, Accumulation of acids (lactic acid-(lactic acidosis) or
ketones-(ketoacidosis).
• Ingestion of acid and Poisoning by organic acids eg. Acetyl salicylates
(aspirin), ethanol, methanol, formaldehyde, ethylene glycol, paraldehyde)
• chronic renal failure (accumulation of sulphates, phosphates,
urea)
• intoxication:
• Sulphates, metformin(Glucophage)
• massive rhabdomyolysis
whatisKussmaulbreathing??

SymptomsofMetabolic Acidosis
Symptomsarenotspecific.
• Extreme acidemia leads to neurological and
cardiac complications;
• Neurological: altered mental status such as
severe anxiety due to hypoxia,
• lethargy, stupor, coma, seizures,
• headache
• Cardiac: arrhythmias (ventricular
tachycardia), decreased response to
epinephrine; both lead to hypotension (low
blood pressure),

Diagnosed using
• arterial blood gas sample (ABG).
• Anion gap.
• The anion gap is the difference or gap in the
measured cations and the measured anions
electrolytes in
• serum,
• plasma, or
• urine.
Diagnosis

SymptomsofMetabolic Acidosis
Others may include;
• decreased visual acuity
• nausea and vomiting
• abdominal pain
• altered appetite and weight gain
• muscle weakness
• bone pain and joint pain.
• diarrhea
• Coma
• Death
Treatment of Metabolic Acidosis
• IV lactate solution

CompensationforMetabolic Acidosis
The body regulates the acidity of the blood by four buffering
mechanisms.
• Bicarbonate buffering system.
• Intracellular buffering by absorption of hydrogen atoms by
various molecules, including proteins, phosphates and
carbonate in bone.
• K+exchangeswithexcessH+inECF (H+intocells,K+out ofcells)
• Respiratory compensation (Increasedventilation).
• Renal compensation (excretionofhydrogenionsandreductioninamt
ofHCO3
- if possible).
Mechanism:
• ↓pH→↑respiration→↓pCO2 to
matchthe lowered HCO3
- →pH backstowardsnormal.

Causes:
 Excess vomitingresultsinthelossofHClacid
(
stomachacid)
 Excessive use ofalkalotic drugs/agents-suchas
bicarbonate(administratedincasesofpepticulceror
hyperacidity).
 Heavy/excessingestionof antacids
 Certain diuretics(suchasloopdiureticsand
thiazides)
 Endocrine disorders
 Severe dehydration

TreatmentofMetabolic Alkalosis
93
• Electrolytestoreplacethose lost
• IV chloridecontainingsolution
• Treatunderlyingdisorder
Generally,inanydisorder:
• Decide the cause of the problem which value,
pCO2 or HCO3- , is outside the normal range.
• If the cause is a change in pCO2, the problem is
respiratory.
• If the cause is HCO3- the problem is metabolic.

Diabetic ketoacidosis
• Normally, the cells of the body take in and use sugar (glucose-
blood sugar made by the liver ) as a source of energy (a fuel
source).
• Glucose moves through the body in the blood.
• A peptide hormone, insulin helps the cells take in the glucose
from the blood and use it for fuel. .
• In diabetes or when the signal from insulin in the body is so
low , the cells cannot take in and use this glucose in a normal
way.
• This may be because the body does not make enough insulin.
Or it may be because the cells do not respond to it normally.
• As a result, glucose builds up in the bloodstream and does not
reach the cells resulting in high blood sugar levels.
• Without the cells to use glucose to function properly, the
body's cells precisely, the liver burns/breaks down some of the
body fat instead of glucose into a fuel for energy.

• That fuel is called ketones (When cells burn fat, ketones are
made).
• When ketones are produced too quickly and build up in the
blood can be toxic.
• A high level of ketones in the urine or blood is called ketosis
which can poison the body.
• High levels of glucose can also build up in the blood and the
excess can alter/change the chemical balance of the blood and
throw off the entire system and cause other symptoms. This is
known as acidosis.
• These include electrolytes such as sodium, potassium, and
bicarbonate to be imbalance which can lead to other
problems.
• The concurrent occurrence of both ketosis and acidosis are
together known as Ketoacidosis which happens most often in

• People with type 1 diabetes are at risk for
ketoacidosis, since their bodies do not make any
insulin.
• Ketones can also go up when :
• Missed a meal
• Sick or stressed
• Have an insulin reaction
• Have not injected enough insulin
• In rare cases, ketoacidosis can happen in a person
with type 2 diabetes.
• It can happen when they are under stress, such as
when they are sick.

• Diabetic ketoacidosis, also known as DKA, is an elevated
blood ketones (ketosis) and acidity of the blood (acidosis).
• Ketoacidosis shouldn’t be confused with ketosis, which is
harmless.
• These ketones are normally used by the muscles and the heart
• Ketosis can occur as a result of an extremely low carbohydrate
diet, known as a ketogenic diet, or fasting.
• When the body excretes these in urine, they can make the
urine smell like popcorn.
• The more ketones in the blood, the more ill a person with
diabetic ketoacidosis will become.
• DKA is a serious complication of diabetes

• Left untreated, diabetic ketoacidosis can be life-threatening,
causes potentially fatal complications, such as severe
dehydration, coma and swelling of the brain
• But it usually takes many hours to become that serious.
• Diabetic patient can be in DKA within 24-48 hours.
• Beyond that, mortal outcomes would likely occur within days to
to perhaps a week or two.
• In summary, diabetic ketoacidosis occurs because there is not
not enough insulin to move sugar (glucose) into the cell where it
where it can be used for energy.

• In older type 2 diabetic patients, it’s more likely to have a
condition with some similar symptoms called HHNS
(hyperosmolar hyperglycemic nonketotic syndrome).
• Hyperosmolar hyperglycemic state (HHS) is one of two
serious metabolic derangements that occur in patients with
diabetes mellitus (DM).
• Any illness that makes you dehydrated or reduces your insulin
activity can lead to HHS.
• It's commonly a result of unmanaged or undiagnosed
diabetes.
• DKA is characterized by ketoacidosis and hyperglycemia,
HHS usually has more severe hyperglycemia but
no ketoacidosis.
• Each represents an extreme in the spectrum of hyperglycemia

• Failure to monitor and manage blood glucose levels can also
lead to HHS.
• It is a life-threatening emergency that, although less common
than its counterpart, diabetic ketoacidosis (DKA), has a
much high mortality rate, reaching up to 5-10%.
• Serum ketones are not present because the amounts of
insulin present in most patients with type 2 diabetes are
adequate to suppress ketogenesis.
The organs affected by ketoacidosis
• DKA can cause complications such as:
• Low levels of potassium (hypokalemia)
• Swelling inside the brain (cerebral edema)
• Fluid inside your lungs (pulmonary edema)
• Damage to your kidney or other organs from your fluid loss.

ACID RAIN/PRECIPITATION
All life on earth depends on water.
Consider when water bodies (rivers, seas, lakes) are
contaminated, ?????
• Causes a dire environmental problem.
• One of the most serious assaults on water quality.
• Rain is considered acidic when it is below 5.5.
• Uncontaminated/normal rain is naturally acidic (5.5 or above) from
carbonic acid (H2CO3).
• Acid precipitation is indicated when rain, snow, sleet, or fog has a pH
lower.
• HOW/ WHY ????
• Reaction of carbondioxide and water forms carbonic acid

• Acid precipitation/ Corrosive downpour is caused
primarily by
• the presence in the atmosphere of sulfur oxides and
nitrogen oxides.
• The major sources of these oxides are,
• Emissions from burning fossil fuels (such as coal, oil,
gas) in factories, electrical power plants, eruption of
volcanoes and even some automobiles into the
atmosphere.
• These gaseous compds react with water in the air to
Causes of Acid Rain

• Winds carry the pollutants away.
• In the presence of sunlight, N and S both react
with O2 to form NO2 and SO2 respectively.
• SO2 will further react with O2 to form SO3.

• NO2 and SO3 will dissolve in H2O found in the cloud
to form nitric acid and sulfuric acid (HNO3 and
H2SO4) respectively which fall back as acid
precipitation
• Acid rain may fall hundreds of kilometers away from
industrial centers.
• This is referred to as negative externalities, external
cost of production, or negative spillover effect.
Causes of Acid Rain

ACID RAIN EFFECTS
•Acid rain does NOT affect humans, only the NOx and SOx do
•Acid rain releases aluminum from soils and eventually the
aluminum can enter streams and lakes.
•Some aquatic life is more affected than others by the lowering
pH.
•Ex- Trout and snails have lower tolerance
•Leaches away certain mineral ions such as calcium and
magnesium ions, that ordinarily help buffer the soil solution and
are essential nutrients for plants.
•Causes slower growth, injury and death to forests
•Can corrode metals and chemically weather rocks

REDUCING ACID RAIN
•Education and awareness:
• reducing energy use lowers SOx
• reducing transportation use lowers NOx
•Use natural gas, oil and nuclear power plants as oppose to coal
•Use cleaner coal. Install wet scrubbers in coal power plants
•Catalytic converters reduce NOx from automobiles
The pH in acidic lakes can be raised by adding limestone (a base)
to the lake
As rain water becomes more acidic what happens to its pH?
1.The pH increases
2.The pH decreases
3.The pH stays the same
4.It turns into lemonade!

 Biological chemistry is “wet” chemistry, with most reactions
involving solutes dissolved in water.
 Chemical reactions depend on collisions of molecules and therefore
on the concentrations of solutes in aqueous solution.
 When carrying out experiments, we use mass to calculate the number
of molecules.
o We know the mass of each atom in a given molecule, so we can
calculate its molecular mass, which is the sum of the masses of all
the atoms in a molecule.
o We measure the number of molecules in units called moles.
• Mole is a unit for measuring the number of molecules in a substance.
Substances are measured in units called ????
• Molarity is the amount of moles of solute per liter of solution, used
by biologist as the unit of conc for aqueous solutions.

 To calculate the number of molecules, mass is used.
 The actual number of objects/molecules in a mole is called Avogadro’s number,
 6.02 × 1023.
 The advantage of using a mole as a unit of measure is that
 a mole of one substance has exactly the same number of molecules as a mole
of any other substance.
 I mole = Avagadro’s number
o If substance A has a molecular weight of 180 daltons and substance B has a
molecular weight of 10 daltons, then 180g of A will have the same number of
molecules as 10g of B.
 A mole (mol) is equal to the molecular weight of a substance but scaled up from
daltons to grams.
 To illustrate, how can we measure 1 mole of table sugar—sucrose (C12H22O11)?
o A carbon atom weighs 12 daltons, hydrogen 1 dalton, and oxygen 16 daltons.
o One molecule of sucrose weighs 342 daltons; the sum of the weights of all the
atoms in sucrose, or the molecular weight of sucrose.
It is important to know how to calculate conc
(amt) of solutes in an aqueous solutions.

 A mole of NaCl also contains 6.02 × 1023 molecules, but has only 58.5g as its
mass.
 This is because the mass of sucrose is greater than that of a molecule of NaCl.
 To make a liter of solution consisting of 1 mol of sucrose dissolved in water.
 342g of sucrose would be measured out, and then gradually add water, while
stirring, until the sugar is completely dissolved.
 Then, enough water would be added to bring it up to the total volume of the
solution 1L.
 At this point, 1M solution of sucrose.
 To get 1 mole of glucose, we would weigh out 342 g.
 TRY THIS:
 NaCl to i.) 1M ii.) 0.5M

WHAT IS ELECTRONEGATIVITY?
• Some atoms are so strongly electronegative that they can capture
electrons from other atoms during a chemical reaction.
• Electronegativity is the ability of an atom in a molecule to attract
electrons to itself than the other.
• E.g:
• When the elements such as Na a silver (colored metal) and Cl (a
toxic gas) are mixed, the single electron in the outer shell of each
sodium atom migrates to the electron- deficient chlorine atom.
• As a result, these two atoms are transformed into charged ions and
presented in crystals; sodium chloride, or table salt.
• This brings us to Bonding.

Chemical Bonding
• What is chemical bond?
• Why bonding or why do atoms bond?
• Bond formation involves the electrons that surround each
atomic nucleus (outer layer of each different atom).
• Atoms bond to acquire a stable configuration, a complete
outer shell.
• A chemical bond is formed when two atomic nuclei attract the
same electron(s) at the outermost shell.
• A force/link that holds two or more atoms of same or different
elements together to form a molecule or compound, called

Valence Electrons
• The electrons that are available to form bond are called …..?
• Valence electrons are electrons that are always at the outer
shell of an atom, they are ready to form bond.
• Valence electrons can exist as paired electrons or
unpaired electrons lying at the outermost shell of an atom.
• The reactivity of an atom arises from the existence of
unpaired electrons in the valence shell.
• They are lost by atoms in ionic bonding and metallic bonding
but shared with other atoms in covalent bonding.
• Valence electrons determine the chemical properties of the
atom because chemical reactions results in the loss, gain, or
sharing (rearrangement) of these electrons.

• Those electrons that are not involved in chemical
behavior (reaction) are called ……………
• Chemical bonding provides the energy necessary
to hold two different atoms together as part of a
chemical compound.
• Energy is released when a bond is formed, and
energy must be supplied to break a bond.
• Strength of the bond depends on the molecules or
atoms involved in the process of bond formation.

Types of Chemical Bonds.
•  Ionic bond
•  Covalent bond
•  Coordinate covalent bond
•  Hydrogen bond
•  Metallic bond:
• This is the electrostatic force of attraction between positively
charged ions and delocalized outer electrons. A bond in
which metallic cations are held together by sea of
delocalized electrons.
• Metal + Metal (bonds hold particles together in a metal).

Ionic Bond
• The electrostatic attraction or a bond formed by the transfer of
valence electron(s) from metal atom to the outer most shell of non
metals atom.
• Key:
Metal + Nonmetal

Example of Ionic Bond
• A classic example of ionic bonding is between Na and Cl.
• Na is a silvery metal.
• It has 1 valence electron.
• Cl is a yellow-green gas,
• and it needs 1 electron to fill its valence shell.
• The ions now have opposite charges and are attracted to each other
by electrostatic forces.
• They form a crystal with the rock salt structure.
• Other ex: CaCl2, CaO, LiBr, BaI2, Al2S3

Ionic Bond contd
Ionic compound

Covalent Bond
• A bond formed by the mutual sharing of valence electron(s)
between the atoms in a molecule, hence, each atom acquires a stable
outer shell of eight electrons.
• Electrons are shared in pairs, and are called electron pair.
• A pair of shared electron is the covalent bond.
• Key:
• A non metal + A non metal,
• A metalloid + a non metal.
• The number of bonds an atom can form depends on the number of
electrons needed to fill its outer shell.

Covalent Bond contd.
• Consequently, it is further classified into:
• single, (e.g:.,H2, Cl2, F2, ….)
• double, (e.g:., O2, CO2) and
• Triple (e.g:., CC, NO, CN, NN,C2H2) covalent bond
• with respect to mutual sharing of one, two, and three bonds
respectively.
CO2

Covalent Bond contd.
• Covalent interaction would be difficult to disrupt when compounds
are put into water. Therefore, the strongest bonding.
• The type of bond between atoms in a molecule has important
consequences in determining the shape of molecules.
• E.g:
• Atoms joined by a single bond are able to rotate relative to one
another.
• Whilst the atoms of double and triple bonds are rigid.

Nonpolar covalent bond
• A bond formed by the mutual sharing of electrons often between atoms of same
element (H2, O2, C4…).
• Same forces of attraction onto the electron pair.
• The electron pair is situated at same distance between the two atoms.
Attractive forces (A = B)

Polar covalent bond
• A bond formed by the mutual sharing of electrons often between the
atoms of different elements (H2O, NH3, CH4 …)
• Different forces of attraction onto the electron pair.
• The electron pair (shared electrons) tend to be situated more closely
to the atom with greater attractive force (more electronegative atom)
than other.
• Attractive force of B > attractive force of A
• Electronegativity of B > electronegativity of A
• Donor atom (δ+) = A
• Acceptor atom (δ-) = B

Polar Molecules
• Among the atoms most commonly present in biological molecules, nitrogen and
oxygen are strongly electronegative.

Hydrogen bond
• A hydrogen bond is an attraction (attractive force) between polar
molecules which occurs when a hydrogen (H) atom is bound to a
highly electronegative atom (acceptor atom(δ-)) of one molecule.
• That hydrogen (donor δ+) is also attracted towards some other
nearby highly electronegative atom of another molecule.
• Each hydrogen atom is covalently bonded to the oxygen via a shared
pair of electrons.
• Usually the electronegative atom has a partial negative charge such
as
• Oxygen (O),
• Nitrogen (N),
• Fluorine (F).

Hydrogen bonds contd
• These Hydrogen bonds can occur
• between different molecules (intermolecular) such as
• water, or
• within different parts of the same molecule (intramolecular)
such as
• DNA.
• The hydrogen bond is weaker than covalent and ionic bonds.
• WATER
DNA

Biological Importance of Chemical bond.
• Ionic Bond
• Compounds with ionic bonds split into ions in water.
• Ions conduct electricity.
• Gives specialized cells (nerve, muscle) excitable
properties.
• Covalent Bond
• This type of bond holds together the long chains of
macromolecules (DNA, RNA and Proteins).

Biological Importance of Chemical bond.
• Hydrogen Bond
• Water:
• H- bonds make water molecules stick together.
• Responsible for many of the strange properties of water.
• Proteins:
• H-bonds cause protein chains to be spiral and bent, giving
unique shapes.
• DNA:
• H-bonds hold together the 2 chains to form the double helix.
• Allow chains to "unzip" for replication and transcription.

Polar and Non polar Molecules
• When two or more atoms form a bond, the entire resulting molecule is either
polar (asymmetrical) or non polar (symmetrical).
• Molecules that have an asymmetric distribution of charge, (or dipole) such as
water, are referred to as Polar Molecules.
• Polarity is the condition of having two poles with opposite qualities.
• E.g:.,
• Water’s single oxygen atom attracts electrons much more forcefully than do
either of its hydrogen atoms.
• As a result, the O-H bonds of a water molecule are said to be polarized,
• Such that the opposite ends of the molecule have opposite charges (one
of the atoms has a partial negative charge and the others a partial positive
charge).

• Polar substances such as HCl and NaCl readily dissolve in a polar solvent like
water.
• Polar molecules of biological importance contain one or more electronegative
atoms, usually O, N, S, or/ Cl.
• Molecules that lack both electronegative atoms and strongly polarized bonds are
said to be non polar.
• Such molecules as oils, fats, and waxes consist entirely of carbon and hydrogen
atoms.
• Non polar molecules are relatively inert, will generally dissolve only in non polar
solvents.
• Non polar substances do not dissolve in water, therefore called
hydrophobic substances.
• Polar substances readily dissolve in water, referred to as hydrophilic
substances.

• Weak attraction exists btw non polar substances.
• Strong attraction exists btw polar molecules.
• Some of the more interestingly biological molecules including
proteins and phospholipids, contain both polar and non polar
regions, which behave very differently.

pH of solutions of weak concentrations
Weak Acid
pH of a 1M solution of ethanoic acid with a Ka value of 1.8 x 10-5
pH = -Log10
pH = -Log10
pH = 2.3723

Ionization of Water
H2O + H2O >> H3o+ + OH-
[H3O+][OH-] = 1.0x 1014
Find the hydroxide ion concentration of 3.0 × 10-2 M HCl.
[H3O+][OH-] = 1.0 × 10-14
[3.0 × 10-2][OH-] = 1.0 × 10-14
[OH-] = 3.3 × 10-13M
Acidic or basic? Acidic
[H3O+] and [OH-] balance each other, more of one means less of
the other

pH of solutions of weak concentrations
Weak Base
pH of a 0.2M solution of ammonia with a Kb value of 1.8
x 10-5
pOH = -log10
pOH = -log10
pOH = 2.7319
pH = 14 – 2.7319
pH = 11.2681
Strong acids dissociate completely in
solution
Strong bases also dissociate completely in
solution

pH Exercises
a) pH of 0.02M HCl
pH = – log10 [H+]
= – log10 [0.020]
= 1.6989
= 1.70
b) pH of 0.0050M NaOH
pOH = – log10 [OH–]
= – log10 [0.0050]
= 2.3
pH = 14 – pOH
= 14 – 2.3
=11.7
c) pH of solution where [H +]
is 7.2x10-8M
pH = – log10 [H+]
= – log10 [7.2x10-8]
= 7.14
(slightly basic

 molecular weight of 100 daltons, then we
know that 10 g of substance A has the same
number of molecules as 100 g of substance
B.
 A mole of sucrose contains 6.02 × 1023
molecules and weighs 342 g, while a
 the same number of molecules as a mole of
any other substance.
 Carbonic acid is a weak acid, which
reversibly releases and accepts hydrogen
ions.
o H2CO3  HCO3
- + H+
Carbonic acid Bicarbonate ion
Hydrogen ion

i.Oxygen
ii.molecular nitrogen
(N2)

General Biology. LECTURE 2. THE CHEMICAL FOUNDATION OF LIFE.pptx

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