Download Note 1.1 Chemistry of Life

UNIT 1: Biochemistry
Chapter 1: The Biochemical Basis of life
pg. 6 – 69
1.1: The Fundamental Chemistry of Life
pg. 8 – 18
The properties of life are based on the hierarchical arrangement of chemical
parts.
Matter makes up everything, it has mass and volume.
Elements are composed of matter. Elements are pure substances and can not
be broken down into smaller substances. Elements are composed of atoms.
There are different types of elements they differ in their atomic structure.
Molecules are created when atoms are arranged in fixed ratios, such as; O2,
and H2 gas.
Compounds are created when different elements are combined together in
fixed ratios, held together by chemical bonds (covalent or ionic), such as;
H2O or CO2.
Living organisms are primarily composed of organic compounds (carbon
containing), such as; carbon (C), Oxygen (O), and Hydrogen (H), and can
also contain Nitrogen (N), and Phosphate (P).
Organizational Hierarchy:
Subatomic Particles
Atom
Molecule
Compounds
Organelles
Cell
Tissues
Organ Systems
Organisms
Ecosystem
Biosphere
Atomic Structure
Elements are made up of atoms, which are the smallest unit that maintains
the chemical and physical properties of the element.
An atom is made up of three different sub atomic particles; neutrons (no
charge), protons (positive charge), and electrons (negative charge).
Atomic number is the number of protons found in the nucleus of the atom. It
determines the particular atom identity. (Periodic Table)
Atomic mass is the sum of the number of protons and neutrons found in the
nucleus of an atom.
Electrons are not found within the nucleus and do not contribute to the mass
of the atom, their mass is negligible in comparison to the proton and neutron.
An atom does not have an over all net charge, this is because there is an
equal number of protons (positive) and electrons (negative).
Table 2: Atomic Number and Mass Number
Isotopes
Isotope – is a form of element that differs in its number of neutrons.
All atoms of the same element will have the same number of protons, but
may vary in the number of neutrons found within their nucleus. This will
create atoms of the same atomic number but they may have different atomic
mass.
Radioisotopes and Radioactive Tracers
Radioisotope - is a radioactive isotope of an element.
Some isotopes have a nucleus that is unstable therefore it may breakdown
over time, giving off particles of matter that can be detected as radioactive.
As the atoms nucleus breaks down, the atom is transformed into a different
element.
Examples: Hydrogen and Carbon
Figure 3: Comparison of nuclei of different Isotopes
Since radioisotopes react in the same way as non - radioactive isotopes of
the same elements, they can be used as tracers, because they are easy to
detect by the energy they are releasing.
Radioactive tracers are used to follow a specific chemical through a
chemical reaction. Doctors can trace the path of the chemical as it passes
through cells and different locations of the body.
Melvin Calvin used the radioactive tracer C14 to determine the sequence of
reactions of photosynthesis.
C14 will decomposes into nitrogen. The extra neutrons decay into a
proton. The increases the proton content to seven, which is
nitrogen. Beta admissions.
Half- Life
Radioactive decay occurs at a steady rate. A constant proportion of
radioisotopes atoms break down during a given interval.
P = (0.5) t/h x mass
1. 20 g of Iodine 131 has a half life of 8 days. What is its mass after 32
days?
P = (0.5)
32/8
x 20 g
P = (0.5) 4 x 20 g
P = 0.0625 x 20 g
P = 1.25 g
Electron Arrangement
Orbital - is a region of space that is occupied by electrons located around the
nucleus of an atom.
The arrangement of electrons determines the chemical properties of an atom.
Electrons are directly involved in the forming of breaking of bonds during
chemical reactions.
Electrons are found moving in specific regions (orbital) around the nucleus
of an atom. Only one or two electrons can be found in any one orbital.
Electrons can be found in energy levels, with in the orbital. These are also
known as energy shells and are numbered 1, 2, 3 … as the shells move
further away from the nucleus.
Energy shell 1, (1s orbital), can hold up to 2 electrons only.
Energy shell 2, (2s, 2p orbitals), can hold up to 8 electrons. (2s, 2px, 2py,
2pz). There are also d and f orbitals.
The further away an electron is from the nucleus, the greater the energy.
Table 3: Type of Electron Orbitals.
Valence Electron - is an electron in the outermost energy level or shell of an
atom.
Atoms that do not have a complete octet in their outer energy shell, or
valence shell, are chemically reactive. Inert atoms, such as; helium and neon
have complete octets and are non reactive.
Atoms in groups 1, 2 and 17 are more reactive. These have a tendency to
lose or gain electrons to complete their valence shells. When metals and non
metals combine, they form ionic compounds.
When non metals combine, they share electrons to complete their valence
shells, and form covalent bonds. They form hybridized electron orbitals.
Biological molecules are created when covalent bonds form between atoms
of C, H, O, and N.
Table 4: Valence Electrons Shells
Chemical Bonds
Ionic Bonds
Ionic Bond - is a bond that results from the attraction between two
oppositely charged atoms or molecules.
Cation - is an ion that has a positive charge
Anion - is an ion that has a negative charge.
An ionic bond is form between two elements that have lost or gain electrons,
creating charged particles (ions), and oppositely charged particles attract.
Alkali and Alkaline Earth metals take on positive charges and Halogens take
on a negative charge. (Metals are attracted to non metals)
Covalent Bonds
Covalent bond - is a bond between two elements that share two or more
pairs of electrons.
The strength of these bonds is determined by the electronegativity of each
atom.
Electronegativity - is the measure of an atom’s attraction to share electrons.
Molecular diagrams use dashes or dots to represent covalent bonds between
atoms. The number of covalent bonds formed between atoms depends on the
number of valence electrons available to be shared.
Carbon has four valence electrons, therefore carbon can form four single
covalent bonds, with hydrogen, or two double covalent bonds with oxygen.
Figure 6: Lewis dot diagram showing the structure of methane CH4
Table 5: Orbital and VSEPR models
VSEPR – valence shell electron pair repulsion is a theory used to predict the
arrangement of the bond angles, developed by Ronald L Gillespie.
Polar Molecules
Polar covalent bond - is a bond between two atoms, make up of unequally
shared electrons.
Polarity is partial positive or negative charge at the ends of a molecule.
Covalent bond deals with the sharing of electrons between atoms, there can
be an unequal sharing between these atoms. The greater the electronegativity
that atom has, the greater the attraction to an electron from another atom.
The unequal sharing of electrons between two atoms, with two different
electronegativities, will result in a polar covalent bond.
An atom which attracts the valence electron more strongly then the other
atom will carry a partial negative charge, while the other will have a partial
positive charge. This will create a molecule with a non-uniform charge
distribution.
A water molecule is a polar molecule, created by the unequal sharing of
electrons. Oxygen takes on a negative charge and the hydrogen is slightly
positive.
Polar molecules are attracted to other polar molecules. These molecules are
also very soluble in water.
Intermolecular Forces
Intermolecular forces are forces of attraction between two molecules and
these forces are known as van der Waals forces.
Intermolecular force - is the force of attraction between two molecules.
Van der Waals forces - are very weak attractions between two molecules, or
parts of molecules, when they are close together. (London force, dipoledipole force, and hydrogen bonds)
Hydrogen Bonds
Hydrogen bond - is the attractive force between a partially positively charge
hydrogen atom and a partially negatively charged atom in another molecule.
When a hydrogen atom is covalently bonded to a strong electronegative
atom in one molecule, a polar molecule is created. When two polar
molecules are attracted to each other because of the positively charged
hydrogen of one molecule is attracted to the negatively charge oxygen,
nitrogen, or fluorine found in another molecule, a hydrogen bond is created.
Other Van Der Waals Forces
London forces – is the weakest of all van der Waals forces. It is the
attraction between non – polar molecules
Dipole-dipole – occur between polar molecules.
Summary of van der Waals Forces
Chemical Reactions
Dehydration reaction - is a chemical reaction in which subunits of a larger
molecule are joined by the removal of water, also called a condensation
reaction.
Hydrolysis reaction - is a chemical reaction in which water is used as a
reactant to split a larger molecule into smaller subunits.
Neutralization reaction - is a reaction which an acid and a base combine to
create a salt and water.
Redox reaction - is an electron transfer reaction.
Oxidation - is a reaction in which a molecule loses electrons.
Reduction - is a reaction in which a molecule gains electrons.