Introduction to Biochemistry

1. INTRODUCTION
TO
BIOCHEMISTRY
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2. ATOMS, MOLECULES and
BONDS
• Introduction
• Atomic structure
• What is the difference between atoms?
• Isotopes
• What determines the chemical characteristics of
elements?
• Ionic bonds
• Covalent bonds
• Molecular properties of water
• Hydrogen bonds

3. Introduction
• What is Biochemistry?

4. Introduction
• An understanding of some basic chemical principles is
necessary to ensure an understanding of the chemistry
of living systems i.e. BIOCHEMISTRY
• All living and non-living compounds of the universe are
composed of chemical elements. There are 118 different
elements with 96 occurring naturally.
• Each element is designated by a chemical shorthand
consisting of the first one or two letters of the English or
Latin name for that element. For example:
– Oxygen : O
– Carbon : C
– Iron : Fe (Latin = ferrum)
– Sodium: Na (Latin = Natrium)
– Potassium: K (Latin = Kalium)

5. Atomic structure
• Bohr atomic theory:-
– The basic unit of an element is an atom.
– An element is a quantity of matter all
composed of the same atoms.

6. Sub atomic particles
• There are three types of sub-atomic
particles
– Protons (+ve charge)
– Neutrons (no charge); These occur in the
nucleus
– Electrons -ve charge
– These circulate around the nucleus.
• The number of protons = the number of
electrons

7. Sub atomic particles
• In this highly simplified diagram of a
carbon atom, 6 protons and 6 neutrons
make up the nucleus and 6 electrons
circulate in two orbits around the nucleus.

8. What is the difference between
atoms?
• The number of protons (or electrons) is fixed for
a particular element.
• Each element has a different number of protons
(and hence electrons) from every other element.
• Mass of proton = mass of neutron.
– Each is given an arbitrary mass value of 1. The mass
of an electron is negligible compared to these. In fact
the mass of an electron is only about 1/2000 that of a
proton, so even in the largest atoms, the total mass of
all the electrons is not equal to even one proton.
• The atomic number of an element = the
number of protons (or electrons)
• Atomic mass = Number of protons + number of
neutrons

9. Isotopes
• The number of neutrons can vary, however,
within a population of atoms of an element.
• Those atoms with a neutron number different
from the majority are called isotopes.
• Note: isotopes will also have a different
atomic mass from the majority of atoms of
that element. The nuclei of isotopes are less
stable than normal atoms.
• The nuclei rapidly change to a more stable form
and release energy.
– This property is referred to as radioactivity.

10. Chemical properties of elements
• The number of protons = the number of
electrons in any atom.
• The number of protons (and electrons) is
constant for all atoms of an element.
• This number determines the chemical
characteristics of that element.
– In fact, the chemical characteristics or properties of
an element are determined by the number and
arrangement of the electrons in the atoms.

11. What determines the chemical
characteristics of elements?
• The term "chemical properties" of elements
refers to how they combine with each other.
• When atoms combine with each other they form
chemical bonds between the atoms.
• As previously mentioned, electrons circulate
around the nucleus in orbits.
• These orbits have progressively higher energy
levels the further they are from the nucleus.

12. Energy levels
• There is a maximum number of electrons
which can be contained in each energy
level.

13. Energy levels and chemical bonds
• When the atoms of two elements combine
they attempt to fill the outermost energy
level with the maximum number of
electrons.
• This stabilizes both combining atoms.
• This requirement restricts the range of
elements that can combine with each
other.

14. Energy levels and chemical bonds
• To achieve the maximum number of electrons in
the outermost energy level, atoms can either :
– donate,
– accept, or
– share
• Electrons to obtain a stable outermost energy
level containing 2 or 8 electrons.
• The mechanism by which the atoms attain a
stable electron configuration will determine the
type of chemical bond formed between the
atoms.

15. Chemical bonds: Ionic Bonds
• Sodium has atomic number 11 i.e. it has 11 protons and
11 electrons.
• Its electrons will be arranged in three energy levels.
• Closest to the nucleus will be the first containing 2
electrons, then the second level containing 8 electrons,
leaving 1 electron in the outermost energy level - as
shown below.

16. • The easiest way for sodium to gain a stable
electron configuration is to donate the single
outermost electron to another atom.
• This will result in a +1 charge for this new entity -
a sodium ion (Na+). The positively charged
sodium ion is called a cation.

17. • Chlorine has atomic number 17. In
contrast to sodium it has 7 electrons in its
outermost energy level.
– It can readily gain a stable electron
configuration by gaining 1 electron from
another atom; achieving a net charge of -1.
• The product resulting from the gain of one
electron is a chloride ion (Cl-).
• The negatively charged chloride ion is
called an anion.

18. • Characteristically, 1 sodium atom will
combine with 1 chlorine atom to form a
chemical compound called sodium
chloride.
•Opposite charges attract so the ions in sodium chloride are held
together by the attraction between Na+ and Cl -. This forms an ionic
bond.
•When compounds containing ionic bonds are added to water they
dissociate into their component ions. This results in them dissolving in
water. When solid sodium chloride is added to water (and briefly
stirred) it dissolves to form a solution of sodium ions and chloride ions.

19. •Ions are produced when atoms can obtain a
stable number of electrons by giving up or
gaining electrons.
•For example Na (sodium) can donate an
electron to Cl (chlorine) generating Na+ and Cl-.
The ion pair is held together by strong
electrostatic attractions.

20. The ability of ions and other molecules to dissolve in water is due to
polarity.
For example, in the illustration below sodium chloride is shown in its
crystalline form and dissolved in water.

21. Covalent bonds
• An alternative type of chemical bond is
called a covalent bond. In this type of
bond 1, 2 or 3 pairs of electrons are
shared between participating atoms.
• The shared electrons now circulate about
both atoms participating in the bond.
• Covalent bonds are relatively strong
• Covalent bonds are much more common
in organic compounds (and therefore in
the biological world)

22. Covalent bonds
• Carbon (which is a fundamental element
in all organic chemistry) always forms
covalent bonds.
• The number of covalent bonds a particular
atom forms is determined by the number
of electrons in the outermost energy level.
• The number of electrons in the outermost
energy level determines the valency of the
element and this value represents the
number of covalent bonds formed.

23. Electronegativity and covalent
bonds
• Electronegativity refers to the tendency for
atoms to bind electrons.
– Oxygen (0) with an electronegativity of 3.5 has a
strong affinity.
– Hydrogen (H)(2.1) and carbon (C)(2.5) each have
lower affinities.
– A bond between C and H will have nearly equal
sharing of electrons.
– Oxygen and hydrogen form a highly polar bond
because of the much stronger affinity for electrons by
O.
– NOTE: Highly electronegative atoms are: Fl,
O,S,P, N

24. Covalent bonds
• In some covalent bonds the electrons are
shared equally between the component
atoms giving an even charge distribution
over the whole molecule - called a non-polar
covalent bond.
• In some molecules one atom attracts the
electrons more than another resulting in
an uneven charge distribution - called a
polar covalent bond.

25. Non polar covalent bonds
• Methane has four covalent bonds between
carbon (C) and hydrogen (H). The figure below
shows the methane molecule in four different
views. Notice how these different views
represent the atoms and their bonds differently.

26. Polar covalent bonds
• Molecules of this type interact with each
other such that positive regions in one
molecule are attracted to negative regions
in adjacent molecules.
• Water contains polar covalent bonds
therefore it will interact with other
compounds with polar covalent bonds.
This can be used to explain many of
the important properties of water.