HSC101 Biochemistry slides.pptx

THREE IMPORTANT TIPS
1) Make it a practice to
-refer to library/internet;
-make summaries;
-draw diagrams/tables.
2) Write neat/coloured notes and highlight
main points in lecture notes.
3) Encouraged to
-work in groups;
-make consultations;
-understand than memorise. 1

UNITS OF MEASUREMENTS
What is Chemistry?
• Chemistry is the study of the properties and
behaviour of matter.
• Matter is anything that can be touched, tasted,
smelled, seen or felt is made of chemicals.
• Chemists hence study the composition and
properties of chemicals and the way they interact
with each other with the main purpose to benefit
humankind by building new materials e.g. drugs,
clothing, etc.
2

Main Areas of Chemistry
• Analytical Chemistry  examines substances.
Identifies materials, measures quantities and
evaluates properties of elements and compounds.
Thus, uses analytical methods/scientific analysis in
order to find out about something
• Physical Chemistry  concerns energy research.
Looks at chemical and physical changes and
examines the relationships between matter and
energy.
• Organic Chemistry  deals with carbon and carbon
compounds, many of which come from plants and
animals (compounds that make up living things).
Useful in development of drugs, petrochemicals,
fertilisers and plastics. 3

E.g. Vinegar-CH3COOH and methane, CH4.
What is so special with “C”?
• It can form straight, branched chains through covalent
bonding:
n-pentane (straight) 2-methylbutane or
isopentane (branched)
• It can form strong, stable bonds with a lot of elements,
e.g. H, O, N, S, P, halogens. This results in a large
number of compounds.
4

Why is organic chemistry important?
• Organic chemistry is important because it is
the chemistry associated with living matter
(plants and animals) e.g. food (carbohydrates,
fats, proteins, vitamins); medicines (drugs
such as aspirin and paracetamol); clothing
(wool, silk, cotton, nylon); cleaning (perfumes,
detergents, flavours, soaps) and fuel
(gasolines, oils). There are over three million
total known organic compounds.
5

• Inorganic Chemistry  deals primarily with non-
carbon chemistry involving metals, minerals and
electronics.
• Thus, it is the study of chemicals of elements
other than carbon e.g. sodium chloride (salt)-
NaCl and water- H2O.
• Inorganic compounds mainly bond using ionic
bonding.
• Few inorganic covalently bonded cpds are
possible – the longest chain is 12 atoms long.
• Most stable inorganic molecules contain 3 or 4
atoms. 6

• In total there are about 500,000 (½ a million)
inorganic substances.
Properties of organic and inorganic cpds
7
Organic Inorganic
Most flammable Most non-flammable
Most low melting points Most high melting points
Low boiling points High boiling points
Most nonpolar (soluble in
nonpolar liquids)
Most polar (insoluble in
nonpolar liquids)
Insoluble in water Soluble in water
Covalently bonded Ionic bonding
Reactions involve molecule Reactions involve ions
Contain many atoms Contain fewer atoms
Complex e.g. diamond, proteins Simpler e.g. salt

Biochemistry
• The study of biochemistry is an examination of
the chemical reactions that occur in living
organisms.
• The Greek word `bios` from which the word
biochemistry is derived can be translated as “life”
or living organism.
• Thus, the word biochemistry means the
chemistry of living organisms. Thus,
biochemistry is the study of living organisms and
their chemical composition.
• Since biochemistry is concerned with the
chemical rxns of living organisms one might ask:
8

• What types of cpds participate in these rxns?
• What types of cpds are the constituents of
living organisms?
• Four of the major types of cpds found in living
organisms are the carbohydrates, the lipids,
the proteins and the nucleic acids.
• To a large extent, biochemical rxns are the
rxns of these types of compounds and their
supporting systems.
9

Why Study Chemistry?
• Chemistry is the central science, central to the
fundamental understanding of other sciences
& technologies.
• Chemistry offers a variety of job opportunities
such as nursing & midwifery, dietetics,
laboratory technology (medical laboratory
science), inhalation therapy (asthmatics),
dental hygiene (dentists), medicine (medical
doctors), pharmacy and anaesthetics.
10

Chemistry Measurements and Units.
• These are important because science involves
making observations
• These observations could be qualitative such
as:
-observing colours
-evolution of heat
-smell
-sound
11

• They could also be quantitative e.g. making
measurements:
-buying food
-giving doses to patients
Systems of Units
• Systems of units evolved with time.
i) English system
ii) Metric system → evolved to the SI
unit system (international system of
units) [Le’systeme internationale
d’unites]
12

Examples of Measurements and Units
13
Quantity English Symbol Definition SI Unit Symbol Further
Definition
Length inch in 0.0254 m meter m
Volume quart qt 1.14 L m3/L m3/L
Mass pound lb 0.4544 kg kilogram kg
Force dyne dyn 10-5 N Newton N kgms-2 (Jm-1)
Pressure atmosphere atm
(torr/mm Hg)
101325 Pa Pascal Pa kgm-1s-2 (Nm-2)
Temperature oC Kelvin K
Current Ampere A
Energy calorie cal 4.184 J Joule J Kgm2s-2
Frequency Hertz Hz s-1
Power watt W kgm2s-3 (Js-1)

Kinds of Units
Base Units
Length: m; Mass : kg;
Time: s; Current: A;
Temperature: K; Amount: mol (6.02 x 1023)
Derived Units
Area = Length x Width: m2
Volume = Length x Width x Height : m3
Force (N) = Mass x Acceleration (ma): kgms-2
Pressure (Pa): = = = = kgm-1s-2
14

Conversions and Decimal Multipliers
15
Conversions:
1 ft = 12 in
1 yd = 3 ft English
1 mile = 1760 yd
Length
1 in = 2.54 cm
1 yd = 0.9144 m English to Metric
1 mile = 1.609 km
1 lb = 453.6 g
Mass
1 oz = 28.358 g
Volume 1 gal = 3.786 L

Multipliers:
• The SI uses combinations of base units and
various prefixes which denote multiples of
powers of 10 of the SI units themselves.
• E.g. the centimetre (cm) is 1/100 of a meter.
The prefix c can be used with any unit and
always means 1/100 or a multiple of 10-2 of
the unit:
1 cm = 10-2 x (1 m) or more simply 1 cm = 10-2 m;
1 cs = 10-2 x (1 s) or more simply 1 cs = 10-2 s
16

Common SI prefixes
Prefix Name Meaning
G
M
k
d
c
m
μ
n
p
f
giga
mega
kilo
deci
centi
milli
micro
nano
pico
femto
109
106
103
10-1
10-2
10-3
10-6
10-9
10-12
10-15
17

Scientific Notation (Standard Form)
• Scientific notation is useful when dealing with
extremely small/large numbers.
• It is expressed in the form: A X 10n where A
satisfies the condition 1  A  10.
• E.g. 600 = 6.00 x 102; 3251 = 3.251 x 103;
0.0045 = 4.5 x 10-3
18

Density
Density = (kg/m3; g/dm3; g/cm3 = g/ml)
Example 1.1
Alcohol has a density of 0.80 g/mL. How much will 100
mL of it weigh?
Solution 1.1
Density =
Therefore, Mass = Density x Volume
= 0.80 g/mL x 100 mL = 80 g
19

Example 1.2
What is the density of mercury in a
thermometer if 31.2 g of it occupies 2.29 mL?
Solution 1.2
Density = = = 13.6 g/mL.
Specific Gravity (S.G.)
• Specific Gravity is a dimensionless unit as it is
a ratio defined as:
20

S.G. =
=
Density of Substance = S.G. x Density of H2O
Example 1.3
The density of mercury is 13.6 g/mL. What is its Specific
Gravity?
Solution 1.3
• S.G. = = = 13.6
21

Temperature Scales
• Temperature is a measure of the availability of
heat or cold.
• Thus it gives a measure of how hot or cold an
object is.
Scales:
• USA English SI
Fahrenheit (oF) Celsius (oC) Kelvin (K)
22

• The Celsius and Fahrenheit temperature
scales compare with one another as follows:
oF =
oC =
Example 1.4
Convert 80.0 oC to oF:
Solution 1.4
oF = = = 144o + 32o = 176 oF
23

Example 1.5
Change 50 oF to oC:
Solution 1.5
oC = = = = 10 oC
The relationship between oC & K temperatures is given by:
K = oC + 273
Example 1.6
Change 37 oC to K.
Solution 1.6
K = oC + 273; K = 37 + 273; K = 310. Thus, 37 oC = 310 K
24

Use of Numbers: Scientific Notation
(Standard Form)
• Recall: A x 10n where A satisfies the condition 1 
A  10.
Addition and Subtraction
• In adding or subtracting each quantities in
scientific notation we first express each quantity
to the same exponent n, then add or subtract the
A parts of the numbers.
• E.g. (7.41 x 102) + (3.26 x 103)
= (7.41 x 102) + (32.6 x 102) = 40.01 x 102
= 4.00 x 103 (in scientific notation)
25

Multiplication and Division
• In multiplying or dividing, mathematical laws
of indices apply; i.e. add powers of 10 when
multiplying and subtract when dividing.
• E.g. (4.0 x 104) x (1.5 x 102) = 4.0 x 1.5 x 104+2 =
6.0 x 106
Significant Figures
• There are two kinds of numbers we mostly
meet in life:
26

1) Exact numbers:
• These may be counted or defined.
• They are absolutely accurate, e.g. the exact
number of students attending an HSC101
lecture can be counted and there is no doubt
about its size.
2) Numbers from Measurements:
• These numbers are not exact, the
measurements are estimates which depend
on the person taking the reading.
• The readings taken as the result of a
measurement are called significant figures.
27

• When significant figures are counted, the last digit is
understood to be uncertain.
• Significant figures are digits believed to be correct by
the person who makes the measurements.
• Suppose one measures the height of a patient and
reports the measurement as 133.5 cm, what does this
mean?
• In this person`s judgement, the height is greater than
133.4 cm but less than 133.6 cm & the best estimate is
133.5 cm.
• The number 133.5 cm contains four significant figures.
• The last digit is a best estimate & is therefore doubtful,
but it is considered significant.
28

• Often in reporting numbers obtained from
measurements, we report one estimated digit
and no more.
• Because the person making the measurement is
not certain that the digit 5 is correct it would be
meaningless to report the height as 133.53 cm.
• The exactness or precision of the measurement
depends upon the limitation of:
a) The measuring device.
b) The skill with which the device is used.
29

• The precision of a measurement is indicated by the
number of figures used to record it.
• These figures include all those that are known with
certainty plus one more, which is the estimate.
• Quite often when a set of measurements are made we
ask how precise and how accurate the readings are.
What does this mean?
Precision
 Precision refers to how closely individual measurements
of the same quantity agree with each other, i.e.
reproducible a given measurement is?
 Precise measurements often show small range or spread.
30

Accuracy
• Accuracy refers to how closely a measurement
agrees with the correct value.
• The simplest way to determine accuracy of the
measurements is to obtain the average or mean
of the measurement.
• Better results are obtained if we consider the
standard deviation.
• Suppose three HSC101 students X, Y, and Z are
asked to determine the mass of a tablet of
panadol of mass 1.270 g.
• The results of two successive weighings by each
student are as follows:
31

Student: X Y Z
Reading 1: 1.235g 1.243g 1.271g
Reading 2: 1.249g 1.248g 1.273g
• From the readings we can conclude that, Y is
more accurate than X, but Z is not only
accurate but is also the most precise of the
three.
32

Rules
1) All non-zero digits are significant.
2) Zeros between non-zero digits are
significant: 202 and 2.02 both have three
significant figures.
3) Zeros to the left of the first non-zero digit in
a number are non-significant; they merely
indicate the position of the decimal point e.g.
0.00730 is three significant figures.
4) When a number ends in zeros that are to the
right of the decimal point, they are significant.
33

5) When a number ends in zeros that are not to the
right of a decimal point, the zeros are not
necessarily significant e.g. 130 cm (two/three
significant figures) – the way to remove this
ambiguity is to use the standard exponential
notation:
E.g. 10300g = 1.03 x 104g (three SF)
= 1.030 x 104g (four SF)
= 1.0300 x 104g (five SF)
6) The results of counting are exact. There is no
uncertainty in the report “15 students”. It means
exactly 15 not 15 ± 1.
34

Calculations involving Significant Figures
Addition and Subtraction
Rule:
• The number of significant figures to the right of
the decimal point in the final sum or difference is
determined by the lowest number of significant
figures to the right of the decimal point in any of
the original numbers.
• E.g.
29.3 (lowest number of SF)
+ 213.87
243.17 (calculated value)
35

• The answer is reported as 243.2 correct to one
significant figure after the decimal point.
• The value 243.17 has been rounded off to 243.2.
• The rounding off procedure is that when the last
digit in the decimal fraction is 4 and below, we
drop the digit, while when it is equal to or greater
than 5, we increase the preceding digit by one.
• 27.234 rounds of to 27.23 (4 is dropped); 27.236
rounds off to 27.24 (3 is increased by 1).
36

Multiplication and Division
Rule:
• The final answer has similar number of significant
figures as the one with the least number of
significant figures.
• E.g. 102.44 x 0.87 – (product must contain two SF)
= 89.1228 – calculated value
= 89 – reported value.
• In calculations that involve multiple steps, the
original numbers must be rounded off before the
mathematical operations are performed.
37

• Each value with an excess number of
significant figures is rounded off so as to have
one more significant figure than the number
of significant figures required to express the
answer.
• E.g. = =
= 2.167 (calculated value)
= 2.2 (reported value)
38

Errors in measurement
• All people make mistakes; therefore each
measurement is of limited accuracy.
• Mistakes can arise from procedure or
instrument reading.
•  it is necessary to attach an index of reliability
to each measurement e.g. 25.0 ± 0.1 oC.
39

Kinds of errors
a) Systematic (Determinate) Errors
• Some common systematic errors are:
-Instrumental errors (faulty equipment,
uncalibrated weights and glassware)
-Operative errors (personal errors e.g.
misreading scale)
-Errors of the method (most serious &
include side reactions and incomplete rxns)
40

Characteristics
• Systematic errors are:
-consistent
-determinable & can be either avoided or
corrected.
Cure
• Systematic errors can be addressed by:
-patience
-being careful
-being observant with peculiarities of
method and instrument
41

b) Random (Accidental or Indeterminate)
Errors
• Random errors are revealed by small differences
in successive measurements made by the same
analyst under identical conditions and they
cannot be predicted or estimated, e.g. reading
balances or thermometer readings.
• Consider the following balance readings:
2.5124 g 2.5122 g 2.5123 g
• Average = 2.5123 g. Hence, there is a random
error due to instrument fluctuations or limitation
inherent in instrument (manufacturing).
42

Characteristics
• Random errors are:
-unpredictable
-limited by instrument capability
Cure
• Random errors may be addressed by:
-running calibrations
- indicating degree of reliability of each
value in your reported data
43

Illustration
• Suppose we weigh a tablet of panadol five times and
get the following result:
2.5124 g; 2.5122 g; 2.5122 g; 2.5125 g; 2.5123 g.
• These measurements are all very close to the average
value (2.5123 g), so we say there is small random
error- the variation from measurement to
measurement, which sometimes gives a high value
and sometimes a low one.
• When the random error is small, we say that the
measurements are precise.
44

• Now suppose that there is a speck of dust of mass
0.0100 g on the balance pan, our measurements
now might be:
2.5224 g; 2.5222 g; 2.5222 g; 2.5225 g; 2.5223 g
• Average value 2.5223 g.
• The measurements are precise, but there is a
systematic error-an error that appears in every
measurement and does not average out.
• Measurements without a systematic error are said
to be accurate i.e. accurate measurements are
close to the accepted value.
45

NB: Every result (number) reported must have
an index of reliability attached: Form, x ± x
Comments
• x is the average value
• x is the index of reliability
• x tells us the degree of precision or certainty
• x is known as the standard deviation.
• x has the same number of decimal points as
mean value
46

MATTER
• Matter is anything that
-occupies space;
-can be seen (soil, trees) or felt (air);
-has mass or weight;
-possesses energy (has the ability to release or
absorb energy).
Examples of Energy
Kinetic Energy Potential Energy
(energy in motion) (stored energy)
-Mechanical energy from a motor -Food (stored energy)
-Light from a lamp -Dry cell (stored
electric energy)
47

Mass and Weight
• Note that in science we distinguish between mass and weight.
Mass Weight
1) Refers to the amount of 1) Force with an object is an attracted
matter in an object by gravity.
2) Same irrespective of 2) Weight depends on location
location
E.g.
(65 kg same on earth and Earth (acceleration = 9.8m/s2);
on the moon) F = ma = 637 N
Moon (acceleration = 1.6m/s2);
F = ma = 104 N
3) Unit of mass is kilogram (kg) 3) Unit of weight is newton (N) or kgm/s2
48

Reasons for studying matter
• To understand its
-composition;
-properties or characteristics;
-structure;
-chemical and physical changes it undergoes
• Such understanding makes it possible to
-create new materials such as drugs, food and
clothing;
-study living organisms including humans (health
science);
-control environment pollution, degradation
(environmental science) significant
49

Properties of Matter
• These are characteristics that help us distinguish
different kinds of matter e.g. gold and iron.
• Properties of matter are classified as physical and
chemical.
Physical Properties Chemical Properties
Examples
Mass Reactions with acids,
Volume water, air (oxygen)
Length Extensive Properties
Area (depend on size of the sample)
Weight
50

Physical Properties Chemical Properties
Examples
Colour
Melting points
Taste Intensive Properties
Boiling points (independent of size of the sample)
Density
Magnetism
Consider gold and iron
• Among the physical properties for gold are
bright yellow colour and not being magnetic
while iron has a dark silver colour and is
magnetic.
51

• In terms of the chemical properties, there is no
reaction for gold in the presence of air and moisture
whereas iron turns to rust due to occurrence of a
reaction that produces iron oxide, i.e.
Gold + O2 + H2O  No reaction
Iron + O2 + H2O  Iron oxide (rust)
(starting material) (new substance)
• The difference in characteristics between a physical
change and a chemical change is that when a
physical change takes place there is no alteration in
the chemical composition of the substance whereas
a substance changes to a new one different from the
starting material when a chemical change occurs.
52

Conservation of Matter
• Matter is conserved during a chemical or a
physical change.
• This means that the total weights of substances
before they react and after they react remains the
same.
• This is known as the law of conservation of
matter.
Example 2.1

Mercury + Oxygen ⇋ Mercuric oxide
 (red-orange residue)
2.53 g ? 2.73 g
How much oxygen is used up in the reaction? 53

Solution 2.1
Mass of Oxygen = 2.73 g – 2.53 g = 0.20 g
heat ()
 Mercury + Oxygen Mercuric oxide
2.53 g 0.20 g
2.73 g 2.73 g
(before rxn) (after rxn)
“Matter is conserved”
54

States of Matter
• Matter exists in three states.
Solid
Melting Condensation
Freezing (Fusion) Sublimation
Evaporation
Liquid Gas
Condensation
55

Composition of Matter
• Matter has three classes namely elements,
compounds and mixtures.
56

Pure substance
• A pure substance is a kind of matter with
constant composition or mixed in the same
proportion by mass.
• E.g. water: H O H₂O
2 1 ratio by mass
• Pure substances cannot be separated into
other forms of matter by physical means such
as distillation, filtration, magnetism,
decantation, sedimentation or
chromatography.
57

• Components of pure substances are separated by
chemical means such as heat.
2HgO
ℎ𝑒𝑎𝑡
2Hg + O2
• Mercury oxide
∆
Mercury + Oxygen.
Mixture
• A mixture is matter with variable composition of
components e.g. salt water can contain 25% salt
or 2.5% salt.
• Mixturesare separable by physical means e.g. a
mixture of iron and sulphur can be separated by a
magnet while a mixture of salt and water (salt
solution) can be separated by distillation.
58

• Mixtures can be classified into two classes i.e.
homogeneous “true solutions’’ e.g. salt solution
or heterogeneous “pseudo (colloid) solutions”
e.g. oil in water.
Elements
• Elements consist of indestructible particles called
atoms which are the smallest building blocks of
matter.
• Atoms cannot be further subdivided into simpler
units by ordinary chemical means.
• Elements consist of only one kind of atom of each
kind (characterized by properties such as mass).
59

Examples of Elements and Functions
• Calcium (Ca) – Bones
• Phosphorus (P) – Bones/teeth
• Magnesium (Mg) – Enzymes
• Iodine (I) – Thyroid
• Fluorine (F) – Teeth
• Sodium (Na) – Ion transport in the body
Compounds
• Compounds are substances formed from two or
more elements through chemical means in a fixed
ratio (or in definite proportions) e.g. sodium
carbonate, Na2CO3 (Na:C:O ratio is 2:1:3); salt (NaCl)
which contains 39.3% Na and 60.7% Cl.
60

Differences between compounds and mixtures
61
Compounds Mixtures
-Separated by chemical means only.
e.g. 2NaCl 
 
 C
o
800
2Na + Cl2
- Composition is fixed (homogeneous) e.g. water,
ratio of hydrogen to oxygen is fixed, NaCl: 39.3%
Na and 60.7% Cl
-Different properties from those substances from
which they are made e.g.Na and Cl are very
poisonous, NaCl is non-toxic.
-Can be readily separated using physical
techniques.
e.g. Salt water 

 
 n
Evaporatio
Salt + Water
-Composition is variable (inhomogeneous)
e.g. salt and sand mixture can have variable
composition; also recall salt water (25% salt
or 2.5% salt)
-Retain properties of individual components
form which they are made e.g. salt water is no
more toxic than NaCl and water.

STRUCTURE OF MATTER
• An explanation of the structure of matter is
founded on laws.
• The best theory that introduces the atom is
the Dalton’s Atomic Theory summarised
below:
1) Elements are composed of extremely small
particles called atoms (indestructible). Now
we know that an atom has three sub-atomic
particles: electrons, protons and neutrons.
62

2) Atoms of a given element have identical size,
mass and chemical properties. Atoms of one
element are different from atoms of other
elements.
3) Compounds are composed of atoms of more than
one element combined in small whole-number
ratios.
4) A chemical reaction involves only the separation;
combination or rearrangement of atoms; it does
not result in their creation or destruction. This
postulate is true for chemical reactions and not
for nuclear reactions.
63

Errors from Dalton’s Theory
1) Atoms are destructible nowadays e.g. high
voltage can emit sub atomic particles called
electrons.
2) Atoms of the same element do not necessarily
have the same mass. Most elements are a
mixture of two or more substance called
Isotopes.
64

Examples of Isotopes
65
Atomic weights are calculated as follows for isotopes:
Chlorine Atomic Weight = 0.757705 x 35 + 0.242295 x 37
= 35.48 g/mole.

Hydrogen Atomic Wt = 0.99986 x 1 + 0.00045 x 2
= 1.0076 g/mole
Atomic weights are obtained from Relative
Atomic Masses:
• Hydrogen (H) 1.008 g/mole
• Mercury (Hg) 200.59 g/mole
Example 3.1
1. Find the weight of each hydrogen atom.
2. How many atoms of hydrogen weigh 1g?
66

Solutions 3.1
1.
Compare with 1 dozen = 12
• 1 mole of H –atoms: weighs 1.008 g
• 1 H-atom weighs =
• 1 atom = 1.67 x 10⁻24 g
• Also 1 atomic mass unit (amu)
• 1 amu = = 1.67 x 10-27 kg = 1.67 x 10-24 g.
• 1 H – atom weight = 1 amu.
67

2. 1 atom weighs 1.67 x 10-24 g (= 1 amu)
 number of atoms in 1g hydrogen.
= = 5.99 x 1023 atoms
Convenience:
• Express 5.99 x 10²³ atoms in moles
• 5.99 x 1023 atoms = ? moles.
• 1 mole = 6.022 x 1023 atoms.
•  5.99 x 10²³ atoms = 0.994 moles
• Hence, 1 g hydrogen has 0.994 moles of
atoms or 5.99 x 1023 atoms.
68

Atomic Structure
Names of elements
• There is no rule or systematic way of naming the
elements.
• Some are ancient e.g. the name of copper is derived
from Cyprus where it was once mined.
• The word gold is derived from an old English word
meaning yellow.
• Some names are based on a characteristic property
of the element e.g. Chlorine is a yellow green gas
and its name is derived from the Greek word
meaning “yellow-green”.
69

• Recently elements have been named after their
discoverers or honour names of places or people
e.g. americium, berkelium, californium,
einsteinium and curium.
• The International Union of Pure and Applied
Chemistry (IUPAC) is the international body that
currently among other duties, approves names
for elements.
• Chemists have a useful system that saves writing
out the full names of the elements. Each element
is represented by a chemical symbol made up of
one or two letters. Many of the symbols are the
first one or two letters of the element’s name:
70

hydrogen H carbon C nitrogen N oxygen O
helium He aluminium Al nickel Ni silicon Si
• The first letter of a symbol is always uppercase
and the second letter always lowercase.
• Some elements have symbols derived from the
first letter of the name and a later letter:
magnesium Mg chlorine Cl zinc Zn plutonium Pu
• Other symbols are taken from the element’s
name in Latin, German or Greek:
71

Elements Symbol Latin Appearance
Copper Cu Cuprum Reddish metal
Gold Au Aurium Yellow metal
Iron Fe Ferrum Silver-white
Lead Pb Plumbum Bluish white
Mercury Hg Hydrargyum Silver white (liquid metal)
Potassium K Kalium Soft silver white
Silver Ag Argentum Silver white
Sodium Na Natrium Soft silver white
Tin Sn Stantum White silver
--------------------------------------------------------------------------------------------------
English
Helium He Colourless
Chlorine Cl Greenish-yellow
Oxygen O Colourless
72

Shape of atom and its constituents
• An atom is made up of 3 sub-atomic
particles: electrons, protons & neutrons:
Shell or orbit or energy level
e  Nucleus (p + n)
Electrons orbit around nucleus on each shell
73

Properties of subatomic properties
• The lightest particle is the electron.
• The proton is  2000 times heavier than the
electron.
74
Particles Symbol M
ass (kg) Charge
electron
prot on
neutron
e-
p
n
9. 1095 x 10- 31
1. 6726 x 10- 27
1. 6750 x 10- 27
negative (-1)
positive (+1)
neutral (0)

Sub-atomic particle numbers
• Atomic Number (Z) represents the number of
protons.
• Mass Number (A), is the number of protons plus
the number of neutrons.
•  A = Z + No. of neutrons
• Mass Number is the Atomic Mass. Z and A for an
atom X are denoted as follows:
• A = Mass number, Z = Atomic number and X =
Atomic symbol
75

Example
Given the atom , state
(i) the number of protons,
(ii) the number of neutrons,
(iii) the number of electrons,
(iv) the number of electrons in the ion .
Solution
(i) p = Z = 11 (ii) n = A  Z = 12
(iii) e = p = Z = 11 in a neutral atom
(iv) e in the ion = 11  1 = 10
76

Note:
• A positively charged ion implies that electrons
have been lost.
• A negatively charged ion indicates gain of
electrons.
• Nucleons are the species found in the nucleus.
Thus, for , the number of nucleons = A =
23.
77

Electronic Arrangement in Atoms
(Electronic Configurations)
n=2
 n=1
n=3
Nucleus
• n represents shell number or orbit number or
energy level.
• Electrons are located in energy levels.
• The maximum number of electrons per shell =
2n2
78

Illustration
Shell Number of Electrons
1 2
2 8
3 18
4 32
5 50
• The maximum number of electrons in any
outer shell is “8” for main group elements.
This is known as the Octet Rule.
79

• The first energy level must be filled before
filling the second shell.
• The same applies with the second and third
shells.
• n is at times known as the Principal Quantum
Number and determines the energy as well as
position of an electron in an atom. n is an integer
and can have the values 1, 2, 3, 4, 5,…etc.
80
El e m
ent At o m
i c Num
ber Shells
n =1 n = 2 n = 3
Hydrogen ( H
) 1 1
Heli um( H
e) 2 2
Lit hi um( L
i) 3 2 1
Fl uorine (F) 9 2 7
Sodi um( N
a) 11 2 8 1
Ar gon ( A
r) 18 2 8 8

Quantum Numbers
• An electron in an atom has some location relative to
the nucleus and is associated with some energy.
• The region of space in which the probability if finding
the electron is maximum is called orbital or atomic
orbital.
• Hence, an orbital is a graphical representation of the
electron probability around the nucleus.
• Both the energy & the probability distribution of an
electron in an atom are described by a set of numbers
called Quantum Numbers.
• The allowed values and general meaning of each of the
four different quantum numbers of an electron in an
atom are as follows:
81

Principal Quantum Number (n)
• Determines the energy level of an electron in
an atom.
• The smaller the value of n, the lower the
energy.
• Also represents the average distance of an
electron from the nucleus in a particular orbit.
• The larger the value of n, the greater the
average distance of an electron in the orbital
from the nucleus and  orbital is larger and
less stable.
82

• The Principal Quantum Number can have any
positive integral values 1, 2, 3, 4…….etc.
• The inner most orbit is given the n of 1;
second has n=2, etc.
• These shells are designated by the letters K, L,
M, N, O,…..etc.
• The maximum number of electrons in an atom
with the given value of n is 2n2. Thus,
83

84
Azimuthal or Angular Momentum or
Subsidiary Quantum Number (l)
• Represents the Shape of the orbital.
• Its value depends upon the value of n.
• For a given value of n, l can have values from 0
to (n-1). Thus, if

• A shell consists of one or more subshells or
suborbital.
• Each subshell is represented by the letters s, p,
d, f…etc.
• The No. of subshells in a given shell is equal to
the value of n.
85
n = 1; l = 0
n = 2; l = 0to (n-1)
l = 0to (2-1)
l = 0, 1
n = 3; l =0, 1, 2

Value of l: 0 1 2 3 4 5……..
Subshells Designations: s p d f g h……..
• The first four letters originate from the sharp,
principal, diffuse and fundamental series of
lines in spectra of alkali metals.
• Starting from the letter f, the subshell
designation is in the alphabetical order.
• There is a slight difference in the energies of s,
p, d and f subshells.
• The order of the energy is s < p < d < f.
86

Magnetic Quantum Number (ml)
• Determines the orientation of the orbital in
space when placed in a magnetic field.
• Its value depends upon the value of l.
• For a given value of l the maximum number of ml
values is (2l + 1).
• The ml can take any integral value from +l
through 0 to –l i.e. +l…..0…..l.
• , if l = 0; ml = 0 (one value); Thus, there is only
one s-orbital present.
87

• For l = 1; ml = +1, 0, 1 (three values).
• Thus, there are three p-orbitals having
different orientations along three Cartesian
axes x, y, and z; Hence px, py and pz.
• For l=2; ml = +2, +1, 0, 1, 2, (five values)
represented by dxy, dyz, dxz, dx
2
y
2 and dz
2
.
• For l=3; ml = +3, +2, +1, 0, 1, 2, 3 (seven
values).
88

Spin Quantum Number (ms)
• Determines the spin of an electron on its own
axis and orientation of the magnetic field
produced by the spin of that electron.
• The two possible values of ms are +½ (for the
spin in the clockwise direction, ) and ½ (for
the spin in the anticlockwise direction, )
89

Information on Quantum Numbers and
Atomic Orbitals
n l Subshell
(nl)
ml Total
No. of
Orbital
s
Orbital
Designation
1 0 1s 0 1 1s
2 0 2s 0 4 2s
1 2p +1, 0, 1 2px, 2py, 2pz
3
0 3s 0
9
3s
1 3p +1, 0, 1 3px, 3py, 3pz
2 3d +2, +1, 0,
1, 2
3dxy, 3dyz,
3dxz, 3dx
2-y
2,
3dz
2
90

Introduction to Organic Chemistry
• The study of compounds whose molecules
contain carbon is called Organic Chemistry.
• Because C has 4 valence electrons
([He]2s22p2), it forms 4 bonds in all its
compounds.
• When all 4 bonds are single bonds, the
electron pairs are arranged in a tetrahedral
arrangement.
91

• When there is one double bond, the arrangement
is trigonal planar.
• With a triple bond, the C compound is linear.
• Almost all organic compounds contain C-H bonds.
• The valence shell of H can only hold 2 electrons, H
forms only 1 covalent bond.
• Hence H atoms are always located on the surface of
organic molecules whereas C-C bonds form the
backbone, or skeleton, of the molecule
C C C
H
H
H
H
H
H
H
H 92

The stabilities of organic compounds
• C forms strong bonds with a variety of
elements, esp. H, O, N & the halogens.
• It can also bond with itself and has the ability
to form long chains.
• Without this property, large biomolecules
such as proteins, lipids, carbohydrates, and
nucleic acids could not form.
93

Introduction to hydrocarbons
• The simplest class of organic compounds is the
hydrocarbons.
• They are characterised by stable C-C & C-H
bonds.
• C is the only element that can form stable,
extended chains of atoms bonded thru single,
double or triple bonds.
94

Classification of HCs
• There are four basic types of hydrocarbons:
– Alkanes (CnH2n+2): E.g. Ethane, CH3CH3
– Alkenes or Olefins (CnH2n):
E.g. Ethene (ethylene), CH2=CH2
– Alkynes (CnH2n2): E.g. Ethyne (acetylene), CHCH
– Aromatic hydrocarbons (where the Cs are connected in
a planar ring structure):
E.g. Benzene, C6H6
95

• Each type has different chemical properties.
• The physical properties are similar in many
ways.
• HCs are non-polar, therefore they are almost
completely insoluble in water.
• Many HCs are familiar because they are used
widely:
– CH4 is a major component of natural gas.
– Propane is a major component of bottled gas
where natural gas is not available.
– Alkanes with 5-12 Cs are used to make gasoline.
96

Structures of alkanes
• There are different ways of writing chemical
formulas (structural & condensed).
• Condensed structures reveal the way in which
atoms are bonded to one another but does
not require drawing all the bonds
STRUCTURAL FORMULAR
98

Nomenclature of alkanes
• Methane CH4
• Ethane CH3CH3
• Propane CH3CH2CH3
• Butane CH3CH2CH2CH3
• Pentane CH3CH2CH2CH2CH3
• Hexane CH3CH2CH2CH2CH2CH3
• Heptane CH3CH2CH2CH2CH2CH2CH3
• Octane CH3CH2CH2CH2CH2CH2CH2CH3
• Nonane CH3CH2CH2CH2CH2CH2CH2CH2CH3
• Decane CH3CH2CH2CH2CH2CH2CH2CH2CH2CH3
99

Structural isomers
• The alkanes we have seen are straight chain
HCs.
• Alkanes consisting of 4 or more carbons can
also form branched chains called branched
chain HCs.
• The branches are often called side chains.
100

• Compounds that have the same molecular
formula but different structures are called
structural isomers.
• Thus C4H10 has 2 structural isomers & C5H12
has 3.
C C C
H
H
H
H
H
H
C
H
H
H
H
C C C
H
H
H
H
C
H
H
H
H
H
H
101

C C C
H
H
H
H
C
H
C
H
H
H
H
H
H
H
C C C
H
H
H
H
H
H
C
H
H
C
H
H
H
H
n-pentane 2-methylbutane
(isopentane)
2,2-dimethylpropane
(neopentane)
C C C
H
H
H
C
C
H
H
H
H
H
H
H
H
H 102

• The number of possible structural isomers
increases rapidly with the number of C atoms
E.g. C8H18 has 18 isomers while C10H22 has 75.
• The physical properties of structural isomers
differ slightly from one another E.g. mpts &
bpts
103

Common names
• There are common names & IUPAC names.
• The common name of the unbranched chain
begins with n (indicating normal).
• When one CH3 group branches off the major
chain, the common name starts with iso-
• When 2 CH3 groups branch off, the name begins
with neo-
• With an increase in the # of isomers, it becomes
impossible to find suitable prefixes.
• Hence the need of a systematic naming system.
104

IUPAC – International Union of Pure &
Applied Chemistry
• The systematic names have 3 parts:
– Base: This tells how many carbons are in the longest
continuous chain.
– Suffix: This tells what type of compound it is.
– Prefix: This tells what groups are attached to the chain.
prefix base suffix
What substituents? How many carbons? What family?
105

IUPAC Naming branched alkanes
• The key steps in the naming of more
complicated branched alkanes are as follows:
1) Identify the longest continuous chain of carbon
atoms.
2) Name this longest root chain using standard
naming rules.
3) Name each side chain by changing the suffix of
the name of the alkane from "-ane" to "-yl”.
106

Condensed Structural Formulas and
Common Names for Several Alkyl Groups
Group Name
CH3 Methyl
CH3CH2 Ethyl
CH3CH2CH2 Propyl
CH3CH2CH2CH2 Butyl
107

4) Number the root chain so that sum of the
numbers assigned to each side group will be as
low as possible (i.e. Number the chain from the end
nearest the first substituent encountered).
5) Number and name the side chains before the
name of the root chain.
6) If there are multiple side chains of the same
type, use prefixes such as "di-" and "tri-" to
indicate it as such, and number each one.
7) Add side chain names in alphabetical
(disregarding "di-" etc. prefixes) order in front of
the name of the root chain.
108

• 8) Use commas between numbers and a dash
between a number and a letter.
Example: Name the following molecule
Problems: Name the following compounds:
(a) CH3CH2CH(CH3)CH(CH3)CH3
(b) CH3CH(CH3)CH(CH3)CH2CH(C2H5)CH(CH3)CH2CH3
Solutions:
(a) 2,3-dimethylpentane
(b) 5-ethyl-2,3,6-trimethyloctane 109

Cycloalkanes
• Simple cycloalkanes (General Formula = CnH2n)
have a prefix "cyclo-" to distinguish them from
alkanes.
• Cycloalkanes are named with respect to the # of
C atoms, e.g., cyclopentane (C5H10) is a
cycloalkane with 5 Cs just like pentane (C5H12),
but they are joined up in a five-membered ring.
pentane
cyclopentane 110

Other examples:
cyclopropane
propane
butane cyclobutane
111

Alkanes can also be represented using the
condensed form
• CH3CH2CH2CH3 CH3(CH2)2CH3
• CH3CH2CH2CH2CH3 CH3(CH2)3CH3
• CH3CH2CH2CH2CH2CH3 CH3(CH2)4CH3
112

Or the bond-line formula
• CH3CH2CH2CH3
• CH3CH2CH2CH2CH3
• CH3CH2CH2CH2CH2CH3
butane
pentane
hexane
113

Alkenes
• Because alkanes have the highest # of possible H
atoms per C, they are said to be saturated HCs.
• Alkenes, alkynes & aromatic HCs contain
multiple bonds.
• Hence they contain less Hs than alkanes with the
same # of Cs.
• They are called unsaturated HCs.
• Unsaturated molecules are more reactive than
saturated ones. 114

Alkenes
• Have at least 1 double bond
C C
H
H
H
H
The simplest member
115

• All the alkenes with 4 or more carbon atoms in
them show structural isomerism.
• This means that there are two or more
different structural formulae that you can
draw for each molecular formula.
C C
H
H3C
CH3
H
C C
H3C
H3C
H
H
C C
H
H
CH2CH3
H
116

Geometric isomers (2-Butene)
• The carbon-carbon double bond doesn't allow
any rotation about it.
• Hence it is possible to have the CH3 groups on
either end of the molecule locked either on
one side of the molecule or opposite each
other.
117

Geometric isomers cont.
C C
H
H3C
CH3
H
C C
H
H3C
H
CH3
trans-But-2-ene
or trans-2-butene
cis-But-2-ene
or cis-2-butene
118

Nomenclature of alkenes
• The names are based on the longest
continuous chain that contains the double
bond.
• Change the corresponding alkane from ane- to
–ene.
• Indicate the position of the double bond by a
prefix # (smallest #)
119

Examples:
1) CH3CHCH2; Propene
2) CH3CH2CHCH2; 1-Butene
3) CH3CHCHCH3; 2-Butene
4)
4-Methyl-2-pentene
120

Alkynes
• HCs that contain at least a Triple CC bond.
• Names of compounds containing CC bonds
end with yne and the rest follows the same
procedure as for alkenes by indicating position
of the triple bond.
Examples:
HC CCH3; Propyne
CH3C CCH3; 2-Butyne
121

Organic functional groups
• A group of atoms like C-O-H group, which
determines how an organic compound reacts
or functions, is called a functional group.
• A functional group is the centre of reactivity in
an organic molecule.
• There are many functional groups in organic
molecules.
• Each undergoes characteristic reactions,
regardless of the molecule size & complexity
of the molecule. 122

• Thus the chemistry of organic molecules is
largely determined by the functional groups.
• The functional groups are usually bonded to
one or more alkyl groups designated as R.
• If more than 1 alkyl group are present, they
are designated as R, R’, R’’ etc
123

Functional Groups
Class of
Compound
Functional
Group
Typical Aliphatic Example
Formula Name
Alkane CH3CH3 CH3CH2CH3 Propane
Alkene CH2CH2 CH3CHCH2 Propene
Alkyne CHCH CH3CCH Propyne
Alcohol ROH CH3OH Methanol
(methyl alc)
Ether ROR’ CH3CH2OCH2CH3 Ethyl Ether
Aldehyde RCHO CH3CHO Ethanal
(Acetaladehyde)
124

Functional Groups cont.
Class of
Compound
Functional
Group
Typical Aliphatic Example
Formula Name
Ketone RCOR’ CH3COCH3 Propanone
(Acetone)
Acid RCOOH CH3COOH Ethanoic acid
(Acetic acid)
Amine RNH2 CH3NH2 Methylamine
Amide RCONH2
CH3CONH2 Acetamide
Ester RCOOR’ CH3COOCH2CH3 Ethyl acetate
where R/R’ are alkyl substituents 125

Alcohols
• An alcohol is a compound in which a H of an
alkane has been replaced by an OH group.
• They are named by changing the last letter in the
name to –ol (e.g. ethane to ethanol).
• Where necessary, the location is designated by a
numerical prefix (e.g. CH3CH2CH2OH;
1-Propanol).
• A study of the properties of the –OH group is
important because of the vast industrial
importance of this functional group and because
of its wide occurrence in biological molecules.
126

TYPES OF ALCOHOLS
PRIMARY (1O) ALCOHOLS
Ones that contain an –OH attached to a C that
has one or no C atoms attached to it. E.g.
H H H
  
H  C  OH H  C  C  OH
  
H H H
Methanol Ethanol
127

SECONDARY (2O) ALCOHOLS
Ones in which the –OH is attached to a C atom
having two other C atoms attached to it. E.g.
H H H
  
H  C  C  C  H
  
H OH H
2-Propanol
128

TERTIARY (3O) ALCOHOLS
Ones in which the –OH is attached to a C atom
that has three C atoms attached to it. E.g.
CH3

CH3  C  CH3

OH
2-Methyl-2-Propanol
129

CARBOHYDRATES
• Contain the elements C, H & O.
• Form a class of organic compounds that include
sugars, starches and cellulose.
• Defined as Polyhydroxyaldehydes or
Polyhydroxyketones or substances that yield these
compounds on hydrolysis.
• Polyhydroxy means “containing several alcohol
groups”.
• Thus simple carbohydrates are alcohols and are
also either aldehydes or ketones (they contain a
carbonyl group).
130

CLASSIFICATION
• Carbohydrates are divided into three
categories based upon hydrolytic possibilities:
Monosaccharides:
• (Mono- means one).
• These are simple sugars.
• They cannot be changed into simpler sugars
upon hydrolysis (reaction with water).
131

Disaccharides:
• (Di-means two).
• These are double sugars on hydrolysis, they
yield two simple sugars.
Disaccharide 2Monosaccharides
• Trisaccharide refers to three linked sugars and
an oligosaccharide to N-linked sugars.


 
Hydrolysis
132

Polysaccharides:
• (Poly-means many).
• These are complex sugars.
• On hydrolysis they yield many simple sugars.
Polysaccharide Many simple sugars
• Monosaccharides (simple sugars) are called
either Aldoses or Ketoses, depending upon
whether they contain an aldehyde (-CHO) or a
ketone (RCOR’) Group.


 
Hydrolysis
133

• Aldoses and Ketoses are further classified
according to the number of carbon atoms they
contain.
• An Aldopentose is a 5-C simple sugar containing
an aldehyde group.
• A Ketohexose is a 6-C simple sugar containing a
ketone group.
• Although there are simple sugars with 3Cs
(Trioses), 4Cs (Tetroses), and 5Cs (Pentoses), the
Hexoses (6-C simple sugars) are the most
common in terms of the human body because
they are the body’s main energy-producing
compounds. 134

ORIGIN
• Plants pick up CO2 from the air and H2O from the
soil and combine them to form carbohydrates in
a process called Photosynthesis.
• Enzymes, chlorophyll and sunlight are necessary.
• The overall reaction:
6CO2 + 6H2O C6H12O6 + 6O2 (1)
Glucose &
other carbohydrates


 

Enzymes
l
Chlorophyl
Sunlight
135

• During Photosynthesis O2 is given off into the
air, thus renewing our vital supply of this
element.
• The carbohydrate produced in reaction (1)
above, C6H12O6 is a monosaccharide.
• Plant cells also have the ability to combine
two molecules of a monosaccharide into one
of a disaccharide:
2C6H12O6 C12H22O11 + H2O (2)
monosaccharide disaccharide
136

• Reaction (2) above is the reverse of hydrolysis.
• Water is removed when two molecules of a
monosaccharide combine.
• Plant (& animal) cells can also combine many
molecules of monosaccharide into large
polysaccharide molecules:
nC6H12O6 (C6H10O5)n + nH2O (3)
monosaccharide polysaccharide
• The n in equation (3) represents a number larger
than 2.
• This is an example of a Polymerization reaction.
137

• Polysaccharides occur in plants as cellulose in
the stalks and stems and as starches in the
roots and seeds.
• Monosaccharides and disaccharides are
generally found in plants in their fruits.
• Plants as well as animals are able to convert
carbohydrates into fats and proteins.
138

O2-CO2 Cycle in Nature
• Although plants have the ability to pick up CO2
from the air and H2O from the ground to form
carbohydrates, animals are unable to do this
and must rely on plants for their carbohydrates.
• Animals oxidise carbohydrates in their bodies
to yield CO2, H2O & energy:
C6H12O6 + 6O2 6CO2 + 6H2O + energy
139

• Note that this overall reaction during metabolism is the
reverse of the one taking place during photosynthesis.
• Both reactions can be summarised by Eqn (4):
Animal metabolism
energy + 6CO2 + 6H2O ⇋ C6H12O6 + 6O2 (4)
Plant photosynthesis
• Thus, there is a cycle in nature.
• During photosynthesis, plants pick-up CO2 from the air
and give off O2.
• Both plants and animals pick-up O2 from the air and give
off CO2.
140

• During photosynthesis, the energy from the
sun is needed for the reaction (endothermic
reaction).
• During metabolism of these carbohydrates in
animals this same amount of energy is
liberated (exothermic reaction).
141

Stereoisomerism
• Stereoisomers are compounds with the same
molecular formula but different structures
that are mirror images of one another.
• The mirror image of an object is the reflection
of that object in a plane mirror e.g. left hand
whose mirror image is the right hand and vice
versa.
• Stereoisomers are not superimposable e.g.
left foot & right foot.
142

• Some molecules are superimposable with
their mirror images and others are not.
CHO CHO
 
H  C  OH HO  C  H
 
CH2OH CH2OH
mirror
glyceraldehyde mirror image of
glyceraldehyde
143

• Glyceraldehyde and its mirror image are not
superimposable; they are said to be chiral.
• An object that is superimposable on its mirror
image is said to be achiral e.g. methane, CH4.
H H
 
H  C  H H  C  H
 
H H
mirror
CH4 molecule mirror image of CH4
144

• In general, any object or molecule with a
plane of symmetry is achiral.
• CH4 has a plane of symmetry; it is achiral.
H

H  C  H Plane of symmetry

H
CH4 molecule
145

• Conversely, if an object or molecule does not
have a plane of symmetry, it is chiral.
• Another method for determining whether a
molecule is chiral is to see if there are four
different groups attached to a central C atom.
• The central C atom in glyceraldehyde is chiral-
it has four different groups attached to it.
• They are -CHO (aldehyde); -OH (2o alcohol);
-CH2OH (1o alcohol) & -H.
146

• Stereoisomers which are chiral are called
enantiomers and are optically active.
• They rotate in the plane of polarised light
equally but in opposite directions.
• Hence, glyceraldehyde can exist in two
optically active forms.
• Enantiomers have 3-D structure & can be
represented by bonds that extend toward the
front as solid wedges and those projecting
toward back as dotted wedges.
147

CHO CHO
H OH HO H
CH2OH CH2OH
structures of glyceraldehyde
• A 2-D method of indicating the structure of an
enantiomer is called a Fischer projection.
• The horizontal lines indicate bonds extending
forward & the vertical lines indicate bonds
extending backward.
• The Fischer projection formulas are always
written with the aldehyde (or ketone) group at
the top. 148

CHO CHO
 
H  C  OH HO  C  H
 
CH2OH CH2OH
Fischer structures of glyceraldehyde
• In these Fischer formulas, the –H and –OH
groups project forward and the –CHO and
–CH2OH groups project backward. 149

D and L Enantiomers
• Enantiomers with the OH group on the left side
of the chiral C are called L (for levo, Latin for left)
compound.
• One with the OH group on the right side of the
chiral C atom are called the D (for dextro, Latin
for right) compound.
D-glyceraldehyde L-glyceraldehyde
150

• Most carbohydrates have longer C chains than
glycelaldehyde & contain more than one chiral
atom.
• In such cases, the carbonyl group is again written
at the top of the structure and the CH2OH group
at the bottom.
• The position of the OH group on the chiral C
atom farthest from the carbonyl group
determines whether the compound will be of
the L or D type.
151

• The number of optical isomers depends on
the number of chiral C atoms present in a
compound and can be calculated by using the
formula 2n, where n = No. of chiral carbons.
• Thus, glyceraldehyde with 1 chiral C has 21 = 2
optical isomers.
• Glucose has 4 chiral C & so 24 = 16 optical
isomers.
• Of these 16 isomers, 8 belong to the D series
and 8 to the L series (one set of eight is the
mirror image of the other set).
153

Aldohexoses: C6, four chiral carbons, sixteen stereoisomers
Aldopentoses and Aldohexoses.
Aldopentoses: C5, three chiral carbons, eight stereoisomers
154

• Stereoisomerism is of great importance in the
body because many enzymes will interact with
only one particular enantiomer.
• In the human body, the D series is the primary
configuration for carbohydrates, whereas the
L series is the primary one for proteins.
• Stereoisomers that are mirror images are
called enantiomers (they are not
superimposable) e.g. D- & L-glyceraldehyde.
• Stereoisomers that are not mirror images are
called diastereomers (they are not
superimposable) e.g. D-ribose & D-lyxose. 155

Single Enantiomers of Chiral Drugs & Their Uses
Enantiomer Use
Ibuprofen Pain
Dexfenfluramine Obesity
Indinavir AIDS
Levofloxacin Antibiotic
Levomoprolol Hypertension
Lisinopril Hypertension
Paclixatel Ovarian cancer
Paroxetine Psychiatric depression
156

Monosaccharides
• Are simple sugars & cannot be broken down
into other sugars.
• Categorised according to the No. of Cs they
contain e.g. trioses (3-Cs), tetroses (4-Cs).
Pentoses
• Are 5-C sugar molecules.
• The most important of these are Ribose and
Deoxyribose which are found in Nucleic acids.
157

• Ribose forms part of ribonucleic acid (RNA) &
deoxyribose forms part of deoxyribonucleic
acid (DNA).
• Both DNA and RNA are components of every
cell nucleus and cytoplasm.
CHO CHO
H OH H H
H OH H OH
H OH H OH
CH2OH CH2OH
D-ribose D-deoxyribose
158

• The prefix de- means without, so deoxy- means
without oxygen.
• Note that deoxyribose has one less oxygen
atom than does ribose.
Hexoses
• These are 6-C sugars & are the most common
of all the carbohydrates.
• The most important as far as the human body is
concerned are glucose, galactose & fructose.
159

• All three of these hexoses have the same
molecular formula, C6H12O6, but different
structural formulas; they are isomers.
Glucose
• Glucose (C6H12O6) is an Aldohexose & can be
represented structurally as:
1CHO
H 2 OH
HO 3 H
H 4 OH
H 5 OH
6CH2OH
D-glucose 160

• The Fischer projection representations for D-ribose,
D-deoxyribose & D-glucose shown above are called
open-chain structures.
• Medically, glucose means the D-isomer because
that is the biologically active isomer.
• Likewise, other hexoses, the D-isomer is commonly
called by name only without the prefix D.
RING STRUCTURES
• Monosaccharides exist in solution mainly as ring
structures in which the carbonyl (aldehyde or
ketone) group has reacted with a hydroxyl group in
the same molecule to form a 5- or 6-membered
ring. 161

• Note that the rxn of an aldehyde or a ketone
with an alcohol yields cpds known as
hemiacetals or hemiketals, respectively.
H H
 
R  C  O + R’OH ⇋ R  C  OH

OR’
aldehyde alcohol hemiacetal
162

OR”

R  C  R’ + R”OH ⇋ R  C  R’
ǁ 
O OH
ketone alcohol hemiketal
 Hence, In solution carbohydrates form cyclic
structures (Haworth projection)
163

-D-glucose -D-glucose
(Chemical names: -D-glucopyranose & -D-glucopyranose
respectively for organic chemists) 164

• Hence, the oxygen that was on the hydroxyl
group is now part of the ring, & the original
carbonyl C, which now contains an –OH group,
has become the anomeric C atom.
• A hydroxyl group on the anomeric C drawn down
below the ring is in the  position; drawn up
above the ring, it is in the  position.
• Cyclization of carbohydrates to the hemiacetal
creates a new chiral center. The hemiacetal or
hemiketal carbon of the cyclic form of
carbohydrates is the anomeric carbon.
• Carbohydrate isomers that differ only in the
stereochemistry of the anomeric carbon are
called anomers e.g. - & -glucose.
165

• In the actual 3-D structure, the ring is not
planar but usually takes a “chair”
conformation in which the hydroxyl groups are
located at a maximum distance from each
other.
166

Mutarotation
• In solution, the hydroxyl group on the anomeric C
spontaneously (nonenzymatically) changes from
the - to -position through a process called
mutarotation.
• When the ring opens, the straight-chain aldehyde
or ketone is formed.
• When the ring closes, the hydroxyl group may be
in either the  or the  position (Figure below).
• This process occurs more rapidly in the presence
of cellular enzymes called mutarotases. 167

• Enzymes are specific for  or  bonds between
sugars and other molecules & react with only one
type. 168

• Glucose is known commonly as dextrose or
grape sugar.
• It is a white crystalline solid, soluble in water
& insoluble in most organic liquids.
• It is found , along with fructose, in many fruit
juices.
• It can be prepared by the hydrolysis of
sucrose, a disaccharide, or by the hydrolysis of
starch, a polysaccharide.
169

• Glucose is the most important of all the
monosaccharides.
• It is normally found in the bloodstream & in the
tissue fluids.
• Glucose requires no digestion & can be given
intravenously to patients who are unable to take
food by mouth.
• Glucose is found in the urine of patients suffering
from diabetes mellitus & is an indication of this
disease.
• The presence of glucose in the urine is called
glycosuria.
170

Galactose
• Galactose is an isomer of glucose & is also an aldohexose.
• The structures are:
1
-galactose (For -galactose, OH
on C1 above plane)

• Glucose and galactose differ from each other
only in the configuration of the H & OH about a
single C atom.
• Two sugars that differ only in the configuration
about a single C atom are called epimers.
• Galactose is converted to glucose in the liver by
a specific enzyme called an epimerase.
• Galactose is present in some glycoproteins &
glycolipids.
• Obtained from lactose, a disaccharide. 172

• Galactosemia, a severe inherited disease,
results in the inability of infants to metabolise
galactose because of a deficiency of either the
enzyme galactose 1-phosphate uridyl
transferase or the enzyme galactokinase.
• The galactose concentration increases in the
blood & urine (galactosuria).
173

Fructose
• Fructose is a ketohexose. Its molecular
formula is also C6H12O6 & can be represented
by the following structures:
174

• Fructose (levulose or fruit sugar) occurs
naturally in fruit juices & honey.
• It can be prepared by the hydrolysis of sucrose,
a disaccharide.
• Fructose is the most soluble sugar & also the
sweetest of all sugars, being 75% sweeter than
glucose.
• Fructosemia, fructose intolerance, is an
inherited disease due to a deficiency of the
enzyme fructose 1-phosphate aldolase.
175

• An infant suffering from this disease
experiences hypoglycemia, vomiting & severe
malnutrition.
• Such a condition is treated by placing the
infant on a low fructose diet.
176

Reactions of the Hexoses
• Hexoses, which are either aldoses or ketoses
show reducing properties.
• This reducing property is the basis of the test
for the sugar in the urine & in the blood.
• When a reducing agent is treated with an
oxidising agent such as Cu2+ complex ion
(Fehlings soln: alkaline Cu2+/tartrate; Benedict’s
soln: alkaline Cu2+/citrate; deep blue colour) a
red-orange ppt of copper(I) oxide (Cu2O) is
formed.
177

• The unbalanced eqn for the rxn of an aldehyde
with Cu2+ can be written as follows:
178

• In this rxn the aldehyde is oxidised to the
corresponding acid.
• Laboratory tests for the presence of glucose in
urine use Benedict’s soln or Fehling’s soln,
both of which contain Cu2+ complex ion.
• Clini-test tablets, which also contain a Cu2+
cpx, give a rapid quantitative measurement of
the concn of glucose present.
• If the blue liquid turns green, a trace of sugar
(glucose) is present recorded as +.
179

• A yellow colour, indicated by ++, indicates up
to 0.5% sugar; an orange colour, +++, 0.5-
1.5%; & a red colour, ++++, over 1.5% sugar.
• Glucose does not normally appear in the urine
for any extended period of time.
• Its persistent presence usually indicates that
something is wrong with the metabolism of
carbohydrates such as diabetes mellitus.
• Another laboratory test for the presence of a
reducing sugar uses Tollens’ reagent which
contains Ag+ cpx ion. 180

• In this rxn glucose is oxidised to gluconic acid as
before and the Ag+ cpx ion is reduced to free
silver which appears as a bright shiny mirror on
the inside of the test tube.
glucose + Ag+ gluconic acid + Ag (s) + H2O
Tollens’ silver
reagent mirror
181

Oxidised & reduced sugars
• The C-C & C-O groups are described as
“oxidised” or “reduced” according to the No.
of e’s around the C atom.
• Oxidation is the loss of e’s & results in the loss
of H atoms together with one or two e’s, or
the gain of an O atom or hydroxyl group.
• Reduction is the gain of e’s & results in the
gain of H atoms or loss of an O atom.
182

Oxidation
• Sugars can be oxidised at the aldehyde C to form
an acid.
• The cpd is no longer a sugar & the ending on its
name is changed from “-ose” to “-onic acid” (e.g.
gluconic acid above).
• If the alcohol at the end opposite the aldehyde
(the other end of the molecule) is oxidised, the
product is called a “-uronic acid”.
• The oxidation of the alcohol end of glucose yields
glucuronic acid.
• Glucuronic acid is a minor product of glucose
metabolism. 183

• If both ends of the glucose are oxidised at the
same time, the product is called saccharic acid
(check structure).
184

Reduction
• The aldohexoses can be reduced to alcohols.
• D-Glucose is reduced to D-glucitol (also called
D-sorbitol) using hydrogenation (H2 and a
metal catalyst).
Sorbitol
accumulation in the
eye is a major factor in
the formation of
Cataracts due to
Diabetes.
185

Fermentation
• Glucose ferments in the presence of yeast,
forming ethyl alcohol & CO2.
• This rxn will not readily occur in the absence
of yeast.
• Yeast contains certain enzymes that catalyse
this particular rxn.
The net reaction is:
C6H12O6  2C2H5OH + 2CO2
glucose ethyl alcohol
186

• Fructose will also ferment; galactose will not
readily ferment.
• Pentoses do not ferment in the presence of
yeast.
Disaccharides
• There are three common disaccharides:
sucrose, maltose & lactose, all of which are
isomers with the molecular formular
C12H22O11.
• On hydrolysis these disaccharides yield two
monosaccharides as follows:
187

C12H22O11 + H2O C6H12O6 + C6H12O6
disaccharide monosaccharides
sucrose glucose + fructose
maltose glucose + glucose
lactose glucose + galactose
The disaccharides are white, crystalline, sweet
solids.
Sucrose is very soluble in H2O; maltose is fairly
soluble; & lactose is only slightly soluble.
188

Reducing Properties
• Of the 3-disaccharides only maltose & lactose
show reducing properties with alkaline Cu2+ cpx
ions.
• Sucrose is not a reducing sugar.
Fermentation
• Sucrose & maltose will ferment with yeast owing
to the presence of the enzymes sucrase &
maltase.
• Lactose will not ferment with yeast because of
the absence of the enzyme lactase.
189

• The identity of a disaccharide can be deduced on the
basis of its fermentation rxn & its reducing properties.
Question
Suppose that a test tube contains a disaccharide, C11H22O11.
Is it sucrose, lactose, or maltose?
Solution
• The identity can be determined by the following method:
(1) Mix the unknown disaccharide with alkaline
Cu2+ cpx & warm gently. If there is no rxn, the
disaccharide must be sucrose. In this case, no
further test is necessary to prove the identity of the
disaccharide.
190

(2) If the unknown disaccharide gives a positive
test with alkaline Cu2+ cpx, it must be either
maltose or lactose. In this case, another
sample of the disaccharide is mixed with
yeast & allowed to stand to observe
whether or not fermentation takes place. If
it ferments, then it must be maltose & if
not, then it must be lactose.
 The same two laboratory tests can be
performed in reverse order with the same
results.
191

Sucrose (Cane Sugar)
• Sucrose is the sugar used ordinarily in the
home & is produced commercially from sugar
cane & sugar beets.
• When sucrose is hydrolysed, it forms a
mixture of glucose & fructose.
• This 50:50 mixture of glucose & fructose is
called invert sugar.
• Honey contains a high % of invert sugar.
192

Maltose (Malt Sugar)
• Maltose is present in germinating grain & produced
commercially by the hydrolysis of starch.
Lactose (Milk Sugar)
• Lactose is present in milk & differs from the
preceeding sugars in that it has an animal origin.
• Certain bacteria cause lactose to ferment, forming
lactic acid.
• When this rxn occurs, the milk is said to be sour.
193

Polysaccharides
• Are polymers of monosaccharides & yield
monosaccharides upon hydrolysis.
• Polysaccharides have a high mol. wt., are
insoluble in H2O, are tasteless, & give negative
tests for reducing sugars.
• These properties are the opposite of those for
monosaccharides & disaccharides.
• Three common polysaccharides are starch,
cellulose & glycogen.
194

Starch
• Plants store their food as starch which is
insoluble in H2O.
• Starch gives a characteristic deep blue colour
with iodine.
• This test is used to detect the presence of starch
because it is conclusive even when only a small
amount of starch is present i.e. if I2 is added to
an unknown & a blue colour is produced, starch
is present.
195

Cellulose
• Plants use cellulose as supporting & structural
parts.
• Wood, cotton & paper are composed primarily
of cellulose.
Glycogen
• Is present in the body & is stored in the liver &
muscles, where it serves as a reserve supply of
carbohydrates.
196

• Glycogen forms colloidal dispersion in water &
gives a red colour with I2.
• It gives no test with alkaline Cu2+ cpx.
• It is formed in the body cells from molecules
of glucose (glycogenesis).
• When glycogen is hydrolysed into glucose, the
process is called glycogenolysis.
Glycogenesis
Glucose Glycogen
Glycogenolysis 197

Lipids
• A 2nd group of organic cpds that serves as food for the
body is the lipids.
• Lipids are organic cpds of biologic origin & in general
1. are insoluble in water.
2. are soluble in nonpolar organic solvents such
as ether, acetone, & CCl4.
3. contain C, H & O; sometimes contain N & P.
4. yield fatty acids on hydrolysis or combine
with fatty acids to form esters.
5. Take part in plant & animal metabolism. 198

Fatty acids
• Fatty acids are straight-chain organic acids.
• The fatty acids found in natural fats usually
contain an even number of C atoms.
• Fatty acids can be either saturated (contain
only single bonds btn C atoms) or unsaturated
(contain a few double bonds btn C atoms).
• Polyunsaturated fatty acids contain many
double bonds.
199

Common Fatty Acids
Saturated fatty acids
200
Name Formula Source
Butyric C3H7COOH Butter fat
Caprylic C7H15COOH Coconut oil
Capric C9H19COOH Palm oil
Palmitic C15H31COOH Palm oil, lard (pig
fat), cottonseed oil
Stearic
(m.p. 70 oC)
C17H35COOH Plant & animal fats
such as lard, peanut
oil
Arachidic C19H39COOH Peanut oil

Unsaturated fatty acids
201
where DB means Double Bonds

• Unsaturated fatty acids have lower m.p. than the
corresponding saturated fatty acids.
• The greater the degree of unsaturation the
lower the m.p.
• The need to lower the amount of saturated fat
has been well publicised & many individuals as
well as commercial establishments have switched
to vegetable oils for food prepn.
• Saturated fats are found in meat & dairy
products & oils such as palm oil.
• Dietary saturated fats increase the blood levels of
low-density lipoproteins (LDL) which aid in the
deposition of cholesterol on artery walls.
202

Classification of Lipids
• Lipids are divided into 3 main categories: simple,
complex and precursor & derived.
Simple Lipids
• Are esters of fatty acids. The hydrolysis of a
simple lipid may be expressed as
simple lipid + H2O fatty acid(s) + alcohol
• If the hydrolysis of a simple lipid yields 3 fatty
acids & glycerol, the simple lipid is called a fat or
an oil. 203

• If it yields a fatty acid & a high mol. wt. monohydric
alcohol, the simple lipid is called a wax.
Complex Lipids
• On hydrolysis yield one or more fatty acids, an
alcohol & some other type of compound.
• In this category are phospholipids & glycolipids.
Phospholipids + H2O Fatty acid + alcohol
+ phosphoric acid
+ a nitrogen compound
204

Precursor & Derived Lipids
• Are cpds produced when simple & complex
lipids undergo hydrolysis.
• They include such substances as fatty acids,
glycerol & other alcohols.
• Derived lipids are formed by metabolic
transformation of fatty acids.
• They include sterols (solid alcohols having a
high mol. wt.) & fatty aldehydes.
205

Fats & Oils
Structure
• Fats are esters formed by the combination of a
fatty acid with one particular alcohol, glycerol.
• If 1 molecule of glycerol reacts with 1 molecule
of stearic acid (a fatty acid), glyceryl
monostearate is formed.
stearic acid + glycerol glyceryl monostearate + H2O
206

H H
C17H35COOH + HO C H C17H35COO C H + H2O
HO C H HO C H
HO C H HO C H
H H
stearic acid glycerol glyceryl monostearate
The product of this rxn can react with a 2nd
molecule & then with a 3rd molecule of stearic
acid.
207

H H
C17H35COOH + C17H35COO C H C17H35COO C H + H2O
HO C H C17H35COO C H
HO C H HO C H
H H
stearic acid glyceryl distearate
H H
C17H35COOH + C17H35COO C H C17H35COO C H + H2O
C17H35COO C H C17H35COO C H
HO C H C17H35COO C H
H H
stearic acid glyceryl tristearate (a fat)
208

• Glyceryl tristearate (tristearin) is formed by
the rxn of one molecule of glycerol with 3
molecules of stearic acid.
• Since stearic acid is a saturated fatty acid, the
product is a fat.
• As the degree of unsaturation of the fatty
acids increases, the m.p. decreases.
• Fats with a m.p. below room temp. are called
oils.
• The glycerol molecule contains 3-OH groups &
so combines with 3 fatty acids.
209

• However, these fatty acids do not have to be
the same.
• Fats & oils can contain 3 different fatty acid
molecules, which can be saturated,
unsaturated, or some combination of these.
Example:
Write out the structural formula for the
triglyceride containing a linoleic acid, a palmitic
acid, and a capric acid.
210

Solution:
• H2COOC17H31
HCOOC15H31
H2COOC9H19
Iodine Number
• Unsaturated fats & oils will readily combine
with I2, whereas saturated fats & oils will not
do so very readily.
• The more unsaturated the fat or oil, the more
I2 it will react with. 211

• The iodine number of a fat or oil is the number of
grams of I2 that will react with the double bonds
present in 100 g of that fat or oil.
• The higher the I2 No. the greater the degree of
unsaturation of the fat or oil.
• In general, animal fats have a lower I2 No. than
vegetable oils.
• This indicates that vegetable oils are more
unsaturated.
• This increasing unsaturation is also accompanied
by a change of state: 212

• Animal fats are solid.
• Vegetable oils are liquid.
• Fats have I2 Nos. below 70; oils above 70.
Uses of Fats in the Body
• Fats serve as a fuel in the body, producing more
energy per gram than either carbohydrates or
protein.
• Metabolism of fat produces 9 kcal/g, whereas the
metabolism of either carbohydrates or protein
produces 4 kcal/g.
213

• Fats also serve as a reserve supply of food &
energy for the body.
• Fat is stored in the adipose tissue and serves
as a protector for the vital organs; i.e. fats
surround the vital organs to keep them in
place & also act as shock absorbers.
• Fats in the outer layers of the body act as heat
insulators, helping to keep the body warm in
cold weather.
214

Physical Properties
• Pure fats & oils are generally white or yellow
solids & liquids respectively.
• Pure fats & oils are odourless & tasteless.
• However, over a period of time fats become
rancid; they develop an unpleasant odour &
taste.
• Fats & oils are insoluble in water but are
soluble in organic liquids such as benzene,
acetone & ether.
215

• Fats are lighter than water & have a greasy
feeling.
• Fats & oils form a temporary emulsion when
shaken with H2O.
• The emulsion can be made permanent by the
addition of an emulsifying agent such as soap.
Chemical Reactions
Hydrolysis
• When fats are treated with enzymes, acids, or
bases, they hydrolyse to form fatty acids &
glycerol.
216

• E.g. when tripalmitin (glyceryl tripalmitate) is
hydrolysed, it forms palmitic acid & glycerol &
requires 3 molecules of water.
• Recall: In the formation of a fat, water is a
product.
H H
3H2O + C15H31COO C H HO C H + 3C15H31COOH
C15H31COO C H HO C H
H H
tripalmitin glycerol palmitic acid
217

• When fats are hydrolysed to fatty acids &
glycerol, the glycerol separates from the fatty
acids & can be drawn off & purified.
• Glycerol is used both medically & industrially.
Saponification
• Is the heating of a fat with a strong base
such as NaOH to produce glycerol & the
salt of a fatty acid (soap).
218

H H
3NaOH + C17H35COO C H HO C H + 3C17H35COONa
C17H35COO C H heat HO C H
H H
tristearin glycerol sodium stearate
(a soap)
Hydrogenation
• Fats & oils are similar cpds except that oils are
more unsaturated i.e. oils contain many double
bonds.
• These double bonds can change to single bonds
upon addition of H2.
219

• Vegetable oils can be converted to fats by the
addition of H2 in the presence of a catalyst.
• This process is called Hydrogenation.
• Margarine is prepared by the hydrogenation
of certain fats & oils with the addn of
flavouring & colouring agents, plus vit. A & D.
• Cpds that give butter its characteristic flavour
are sometimes added.
220

H H
3H2 + C17H33COO C H C17H35COO C H
C17H33COO C H catalyst C17H35COO C H
C17H33COO C H C17H35COO C H
H H
triolein, an oil tristearin, a fat
(contains double bonds) (contains single bonds)
 In actual practice, vegetable oils are not completely
hydrogenated.
 Enough H2 is added to produce a solid at room temp.
 If the oil were completely hydrogenated, the solid fat
would be hard & brittle & unsuitable for cooking
purposes. 221

Acrolein Test
• Is a test for the presence of glycerol & is
sometimes used as a test for fats & oils since
all fats & oils contain glycerol.
• When glycerol is heated at a high temp.
especially in the presence of a dehydrating
agent such as potassium bisulphate (KHSO4), a
product called acrolein results.
• This substance is easily recognised by its
strong, pungent odour.
222

H
H C OH heat H C  O
H C OH H C + 2H2O
H C OH KHSO4 H C
H H
glycerol acrolein
• When fats or oils are heated to a high temp.
or are burned, the disagreeable odour is that
of acrolein.
223

Rancidity
• Fats develop an unpleasant smell & taste when
allowed to stand at room temp for a short period
of time i.e. they become rancid.
• Rancidity is due to two types of rxns: hydrolysis
& oxidation.
• O2 present in air can oxidise some unsaturated
parts of fats & oils.
• If this oxidn rxn produces short-chain acids or
aldehydes, the fat turns rancid as evidenced by a
disagreeable smell & taste.
224

• Since oxidn as well as hydrolysis occur more
rapidly at higher temp; fats & foods containing a
high % of fats should be stored in a cool place.
• Oxidn of fats can be inhibited by addn of
antioxidants (substances that prevent oxidn).
• Two naturally occuring antioxidants are vit. C &
vit. E.
• When butter is allowed to stand at room temp.
hydrolysis occurs btn the fats & water present in
butter yielding fatty acids & glycerol.
225

• One of the fatty acids produced, butyric acid
has the unpleasant smell that causes one to
say that the butter is rancid.
• The catalysts necessary for the hydrolysis rxn
are produced by the action of micro-
organisms present in the air acting on butter.
• At room temp. this rxn proceeds rapidly so
that the butter soon turns rancid.
• This effect can be overcome by keeping the
butter refrigerated & covered.
226

Soaps
• Soaps are produced by the saponification of
fats.
• Soaps are salts of fatty acids.
• When the saponifying agent used is NaOH, a
sodium soap is produced.
• Sodium soaps are bar soaps.
• When the saponifying agent is KOH, a
potassium soap is produced.
• Potassium soaps are soft or liquid soaps. 227

C17H35COOH + NaOH C17H35COONa + H2O
stearic acid sodium stearate (a soap)
• Various substances can be added to soap to
give them a pleasant colour & smell.
• Germicidal soaps contain a germicide.
• Scouring soaps contain some abrasive.
• Ca & Mg ions present in hard water react with
soap to form insoluble Ca & Mg soaps.
228

2Na soap + Ca2+ Ca soap(s) + 2Na+
2Na soap + Mg2+ Mg soap(s) + 2Na+
• The soap “precipitate” is mostly organic &
floats to the top rather than sinking to the
bottom as most precipitates do.
• This precipitated soap is seen as “the ring
around the bath tub”.
• More soap is required to produce a lather in
hard water than soft water.
229

Cleansing Action
• Soaps are cleansing agents.
• Consider a soap molecule such as sodium stearate,
CH3(CH2)16COONa.
• The long-chain aliphatic part is non-polar, whereas
the carboxylate part is polar.
Note on Polarity of bonds & Partial Charges
• Polar bonds are covalent bonds in which the
electron cloud is more dense around one atom (the
atom with greater electronegativity) than the other.
230

• O is more electronegative than C, & a CO bond
is polar, with the O atom carrying a partial
negative charge () & the C atom a partial
positive charge (+) .
• In nonpolar CC bonds & CH bonds, the two
electrons in the covalent bond are shared almost
equally.
• In general, nonpolar cpds dissolve in nonpolar
liquids & polar cpds dissolve in polar liquids.
• If soap is added to a mixture of water & oil &
then shaken rapidly, the nonpolar end of the
soap molecule will dissolve in the oil, a nonpolar
liquid.
231

• At the same time, the polar end of the soap
molecule will dissolve in the water, a polar
liquid.
• The nonpolar end of the soap molecule is said
to be hydrophobic (water repelling).
• The polar end is hydrophilic (water loving).
• The carboxylate end of the soap molecule,
which is in water, yields Na+ ions, which are free
to move about generating a micelle.
• Note that the oil drop has a negative charge
because of the negative ends of the soap
molecule sticking out into the water. 232

• This negatively charged oil drop will repel all other oil drops, which
will have acquired a like charge.
• The oil will have become emulsified with the soap acting as the
emulsifying agent.
• This is the manner in which soap cleanses, since most dirt is held on
skin & clothing by a thin layer of grease or oil.
• Mechanical washing causes the oil or grease to break up into small
drops, the soap then emulsifies that oil or grease, which can then
be easily washed away.
233
water
Na+
water Na+
Oil
Soap in an oil-water mixture

Detergents
• Detergents (syndets) are synthetic cpds used as
cleansing agents.
• They work like soaps but are free of several of the
disadvantages that soaps have:
1. Detergents work as well in hard water as they
do in soft water i.e. Ca & Mg salts of detergents
are soluble & do not precipitate out of soln.
2. Detergents are generally neutral cpds compared
with soaps which are usually alkaline or basic
substances.
234

3. Detergent containing straight chains are
biodegradable & do not cause water
pollution whereas those containing
branched chains are nonbiodegradable &
cause pollution.
• Detergents are sodium salts of long-chain alcohol
sulphates e.g. sodium lauryl sulphate.
C11H23CH2OH + H2SO4 C11H23CH2OSO3H + H2O
lauryl alcohol lauryl hydrogen sulphate
C11H23CH2OSO3H + NaOH C11H23CH2OSO3Na + H2O
lauryl hydrogen sulphate sodium lauryl sulphate (a detergent)
235

Waxes
• A wax is a cpd produced by the rxn of a fatty
acid with a high mol. wt. monohydric alcohol
such as myricyl alcohol (C30H61OH) or ceryl
alcohol (C26H53OH).
• Beeswax is largely C15H31COOC30H61 (an ester
of myricyl alcohol).
• Note that waxes are primarily esters of long-
chain fatty acids with an even No. of C atoms
& long-chain alcohols also with an even No.
of C atoms.
236

• The No. of C atoms is usually 2634.
• The alcohol may also be a steroid such as
lanosterol producing lanolin widely used in
cosmetics & ointments.
• Waxes are insoluble in water, nonreactive &
flexible; hence waxes make excellent protective
coatings for plant leaves, skin lubrication,
“waterproofing” feathers of birds.
237

Complex Lipids
Phospholipids
• Are phosphate esters and can be divided into
two categorises  phosphoglycerides &
phosphosphingosides depending on whether
the alcohol is glycerol or sphingosine.
• Phospholipids also contain a nitrogen
compound.
• Phospholipids are found in all tissues in the body
& occur in the membranes of all cells.
238

• Are responsible for passage of various
substances into & out of the cells.
Phosphoglyceride
fatty acid
fatty acid
phosphoric acidnitrogen compound
 At carbons 1 & 2 of the glycerol there are esters
of fatty acids. 239
g
l
y
c
e
r
o
l

• At carbon 3 there is a phosphate group
bonded to a nitrogen cpd.
• There are many different phosphoglycerides,
depending on the types of fatty acids bonded
to the glycerol & also on the identity of the N
cpd bonded to the phosphate gp.
• Most phosphoglycerides have a satd fatty acid
connected at C1 & an unsatd fatty acid at C2.
• The phosphate gp & the N cpd are polar
substances whereas the fatty acid molecules
are nonpolar. 240

Glycero-
phospholipids
Phosphatidic acid
4 major
glycerophospholipids
are polar/charged:

• The fatty acid chains are hydrophobic (they
point away from water) while the one
containing N cpd & phosphoric acid is
hydrophilic (dissolves in water).
• Molecules with a hydrophobic (nonpolar) & a
hydrophilic (polar) end are said to be
amphipathic, e.g. soap.
• Phosphoglycerides can be subdivided into
several types, depending on the nitrogen
compound present e.g. lecithins & cephalins.
242

H
H C OOCC17H35
C15H27COO C H
O
H C O P O CH2 CH2 N+ (CH3)3
H O
• Lecithins are good sources of phosphoric acid
needed for the synthesis of new tissue.
• Lecithin is abundant in egg yolk & soybeans &
commercially used as an emulsifying agent in dairy
products & in the manufacture of mayonnaise.
243
fatty acid
fatty acid
phosphoric
acid
choline
glycerol
A Lecithin

• Cephalins are similar to lecithins except that
another N cpd, ethanolamine (CH2CH2N+H3)
is present instead of choline.
• Cephalins are important in the clotting of the
blood & also are sources of phosphoric acid
for the formation of new tissue.
Phosphosphingosides
• Also called sphingolipids differ from
phosphoglycerides in that they contain the
alcohol sphingosine in place of glycerol e.g.
sphingomyelin.
244

A sphingolipid
fatty acid
phosphoric acidcholine
General formula for a sphingolipid
245
s
p
h
i
n
g
o
s
i
n
e

CH3 (CH2)12
C H
H C
O HO C H
H31C17C NH C H O
H C O P O CH2 CH2 N+ (CH3)3
H O
sphingomyelin
• Note that the fatty acid in sphingomyelin is
bonded to an NH2 gp rather than an OH gp as in
phosphoglycerides.
246
fatty acid
phosphoric
acid
choline
sphingosine

Glycolipids
• Are similar to sphingomyelins except that they
contain a carbohydrate, often galactose in
place of the choline & phosphoric acid.
• Glycolipids produce no phosphoric acid on
hydrolysis because they do not contain this
compound.
• Glycolipids are also called cerebrosides
because they are found in large amounts in
the brain tissue.
247

Derived Lipids
Eicosanoids
• Are a biologically active gp of cpds derived from
arachidonic acid.
• They are extremely potent cpds with a variety
of actions.
• Among the eicosanoids are the prostaglandins,
the thromboxanes, & the leukotrienes.
248

249
• Three classes of the Eicosanoid class of lipids:
1. Prostaglandins - isolated from prostrate
gland, found in nearly all tissues.
2. Thromboxanes - 6-membered rings with
oxygen - may help in blood clotting.
3. Leukotrienes - isolated from leukocytes
(white blood cells), cause contraction of
smooth muscle.
• All are derived from polyunsaturated 20-
carbon fatty acid, arachidonate (Recall:
C19H31COOH; 4DBs)

Three classes of the Eicosanoid lipids are all derived from
polyunsaturated 20-carbon fatty acid, arachidonate
250
•PGE2 induces wakefulness
•PGD2 promotes sleep
tri-refers to 3
alternate sets of DBs

The Prostaglandins
• Consist of 20-carbon unsaturated fatty acids
containing a 5-membered ring & 2-side chains.
• One side chain has 7-C atoms & ends with an acid
group (COOH).
• The other chain contains 8-C atoms with an –OH gp
on the 3rd C atom from the ring.
• The E series of prostaglandins has, in addn to 4 chiral
C atoms, a trans-configuration.
• The abbrev. PGE1 refers to prostaglandin E with 1
double bond (DB) (PGE2 has 2 DBs).
251

Steroids
• Are high mol. mass tetracyclic (four-ring) cpds.
Steroid structures have four fused rings, A, B, C, and D.252

• Those containing one or more OH groups &
no C=O gps are called sterols.
• The most common sterol is cholesterol found
in animal fats but not in plant fats.
• Cholesterol levels in humans should be in the
range of 200220 mg/dL.
• Elevated cholesterol levels should be
controlled, usually by diet.
• In extreme cases, cholesterol-lowering drugs
such as pravastatin or lovastatin may be
prescribed.
253

• Other steroids include the sex hormones &
the hormones of the adrenal cortex.
a) Cholesterol molecular structure - a steroid
254

255
Structures of bioactive products produced from cholesterol.
b) Estradiol - a female sex hormone.
c) Testosterone - a male sex
hormone.
d) Cortisol - a regulator of glucose
metabolism.

Anabolic Steroids
• Anabolic steroids are hormones that control
the synthesis of larger molecules from smaller
ones.
• Athletes have used these substances (illegally)
to increase muscle mass, & hence body
strength e.g. the male hormone testosterone.
• While it does increase muscle mass, it has
several undesirable side effects.
256

• In men these side effects include:
(1)Impotence
(2)Hypercholesterolemia
(3)Breast growth
(4)Liver cancer.
• Women using anabolic steroids will develop
the following side effects:
(1) Increased masculinity
(2) formation of greater amount of hair
(3) deepening of the voice
(4) menstrual irregularities.
257

• Another drawback of the use of such anabolic
steroids is that they cannot be taken orally;
they must be injected.
• It is standard practice nowadays during
athletic competitions to routinely test (WADA-
World Anti-Doping Agency) an athlete’s urine
for the presence of these illegal substances.
258

Proteins
• Other than H2O, proteins are the chief
constituents of all cells of the body.
• They are much more complex than either fats
or carbohydrates.
• All proteins contain C, H, O & N.
• Most proteins also contain S, some contain P
& a few such as haemoglobin contain some
other element e.g. Fe.
259

Sources
• Plants synthesise proteins from inorganic
substances present in the air & in the soil.
• Animals cannot synthesise proteins from such
materials.
• Animals must obtain proteins from plants or
from other animals who in turn have obtained
them from plants.
• Proteins function in the body in the building of
new cells, the maintenance of existing cells &
the replacement of old cells.
260

• Thus, proteins are the most important type of cpd in the
body.
• Proteins are also a valuable source of energy in the body (1
g of protein yields 4 kcal; just as does oxidn of 1 g of
carbohydrate).
• Proteins are involved in
(1) regulation of metabolic processes (hormones),
(2) catalysis of biochemical rxns (enzymes),
(3) transportation of O2 (haemoglobin),
(4) body’s defense against infection (antibodies),
(5) transmission of impulses (nerves),
(6) transmission of hereditary characteristics
(nucleoprotein);
(7) muscular activity (contraction).
261

Molecular Masses
• Proteins have very high molecular masses.
Amino Acids
• Proteins are polymers built up from simple units
called amino acids.
• Hydrolysis of proteins yields amino acids.
• There are 20 known amino acids that can be
produced by the hydrolysis of protein.
• All these amino acids, except glycine, which has
no chiral C, have the L-configuration.
262

Fischer Projections of Amino Acids
Amino acids
• Are chiral except for glycine.
• Have Fischer projections that are stereoisomers.
• That are L are used in proteins.
L-alanine D-alanine L-cysteine D-cysteine
263
CH2SH
H2N H
COOH
CH2SH
H NH2
COOH
CH3
H NH2
COOH
CH3
H2N H
COOH

Composition
• An amino acid is an organic acid that has an
amine (NH2) group attached to a chain
containing an acid gp.
• Although the amine gp can be anywhere on the
chain, amino acids found in nature usually have
the amine gp on the alpha () C, i,.e. the C atom
next to the acid gp. [Note: 2nd C = beta (); 3rd C =
gamma () and 4th C = delta ()].
• -Amino acids are represented by the general
formula below, where R can be many different
gps.
264

COOH CHO
H2N C H HO C H
R CH2OH
L-amino acid L-glyceraldehyde
• Amino acids are grouped according to whether
their side chains (R group attached to the -
carbon) are:- acidic, basic, polar or nonpolar.
• If the R gp is nonpolar, then the amino acid will
be less soluble in water than one containing a
polar gp (-OH, -SH, NH2 or –COOH).
265

• The body can synthesise some but not all of
the amino acids it needs.
• Those that it cannot synthesise must be
supplied from the food consumed.
• These are called the nutritionally essential
amino acids (Arginine, Isoleucine, Leucine,
Lysine, Methionine, Phenylalanine,
Threonine, Tryptophan, Valine, Histidine).
266

Amphoteric Nature
• Amino acids contain the COOH gp, which is
acidic & the NH2 gp which is basic.
• In soln, the carboxyl gp can donate a
hydrogen ion (H+) to the amino gp, forming a
dipolar ion, called a zwitterion.
R CH COOH R C COO
NH2 NH3
+
amino acid zwitterion form of an amino acid
 Amino acids are amphoteric cpds; i.e. they can react
with either acids or bases. 268

• When an amino acid is placed in a basic soln, it forms a
negatively charged ion that will be attracted toward a +vely
charged electrode.
• In an acidic soln, it forms a positively charged ion that will be
attracted toward a vely charged electrode.
H+ OH
R CH COOH ⇋ R CH COO ⇋ R CH COO
NH3
+ NH3
+ NH2
+vely charged ion zwitterion vely charged ion
(in acid soln) (in basic soln)
269

• Since amino acids are amphoteric, proteins which are
made up of them are also amphoteric which accounts
for their ability to act as buffers in the blood; they
can react with either acids or bases to prevent an
excess of either.
• At a certain pH the amino acids will be neutral (equal
no. of +ve & ve ions).
• This point is called the isoelectric point (pI).
• At a pH above the isoelectric point, a protein has
more ve than +ve charges & below the isoelectric
point a protein has more +ve than ve charges.
270

Examples of Amino Acids
271
H
+ │
H3N—C—COO
│
H glycine
CH3
+ │
H3N—C—COO
│
H alanine

272
Types of Amino Acids
Amino acids are classified as
 Nonpolar (hydrophobic)
with hydrocarbon side
chains.
 Polar (hydrophilic) with
polar or ionic side chains.
 Acidic (hydrophilic) with
acidic side chains.
 Basic (hydrophilic) with
–NH2 side chains.
Nonpolar Polar
Acidic
Basic

A nonpolar amino acid has an R group that is
H, an alkyl group, or aromatic.
273
Nonpolar Amino Acids

Polar Amino Acids
A polar amino acid has an R group that is an
alcohol, thiol, or amide.
274

Acidic and Basic Amino Acids
An amino acid is Acidic with a carboxyl R group (COO−);
Basic with an amino R group (NH3
+).
275
Acidic Amino Acids
Basic Amino Acids
* Isoelectric Point

Learning Check
Identify each as (P) polar or (NP) nonpolar.
+
A. H3N–CH2–COO− (Glycine)
CH3
|
CH–OH
+ │
B. H3N–CH–COO− (Threonine)

Solution
Identify each as (P) polar or (NP) nonpolar.
+
A. H3N–CH2–COO− (Glycine) (NP) nonpolar
CH3
|
CH–OH
+ │
B. H3N–CH–COO− (Threonine) (P) polar

A zwitterion
• Has charged —NH3
+ and COO- groups.
• Forms when both the —NH2 and the —COOH groups
in an amino acid ionize in water.
• Has equal + and − charges at the isoelectric point (pI).
O O
║ + ║
NH2—CH2—C—OH H3N—CH2—C—O–
Glycine Zwitterion of glycine
Summary of Zwitterions and
Isoelectric Points

In solutions more basic than the pI,
 The —NH3
+ in the amino acid donates a proton.
+ OH–
H3N—CH2—COO– H2N—CH2—COO–
Zwitterion Negative ion
at pI pH > pI
Charge: 0 Charge: 1−
Summary of Amino Acids as Acids

In solutions more acidic than the pI,
 The COO− in the amino acid accepts a proton.
+ H+
+
H3N—CH2—COO– H3N—CH2—COOH
Zwitterion Positive ion
at pI pH< pI
Charge: 0 Charge: 1+
Summary of Amino Acids as Bases

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