Ch 2 The Chemical Context
of Life
I. Overview
All living organisms are subject to the laws of
chemistry & physics.
B. A basic knowledge of both helps us to better
understand how living things work.
1. What class of levers do you find most
often in the human body?
2. Why use those?
3. How do moth species recognize mates of
the same species?
4. Read the test case on the “Devil’s
gardens” in the rain forest in Ch 2.
A.
Section 2.1 The nature of Matter
A. Matter- anything that has mass and takes
B.
up space
1. Matter consists of chemical elements in
pure form or in combinations called
compounds
2. All organisms have mass & take up
space, therefore all organisms are made of
matter
Matter is made of elements.
1. Elements are substances that can’t be
broken down into other substances by
chemical reactions
C.
C.
D.
A compound is a substance consisting of
two or more different elements
chemically combined in a fixed ratio.
1. Compounds will have chemical and
physical characteristics different from
those of their elements.
2. ex. Salt is made of Sodium (Na) and
Chlorine (Cl). Sodium explodes in water,
chlorine is a poison. Salt is a harmless, edible
compound.
The smallest unit of an element is an atom.
The smallest unit of a compound is a
molecule.
Making Salt
II. Essential Elements of Life
A.
B.
25 of the 92 elements are essential to life.
Carbon, hydrogen, oxygen and nitrogen
make up 96% of all living matter.
(CHON) or (CHONS)
C.
Calcium, potassium, phosphorous, and
sulfur makes up most of the remaining 4%.
(CaPSK)
D.
Trace elements make up about 0.05%.
Insufficient nitrogen in
plants
Iodine deficiency in people
What happens when essential
elements are missing?
III.
A.
B.
An element’s properties depend on
the structure of its atoms.
An atom is the smallest unit of matter that
retains the properties of an element.
Atoms are made of subatomic particles which
include:
1. Neutrons: 0 charge, found in the nucleus
2. Protons: + charge, found in the nucleus
3. Electrons: - charge, found in the cloud around the
nucleus
C.
Atomic number- the number of protons in an
element’s nucleus- the number is unique for
each element
D.
E.
F.
All atoms of an element have the same
number of protons, but may have different
numbers of neutrons in the nucleus.
Isotopes- when two atoms of an element
have different numbers of neutrons
Radioactive isotopes- may decay
spontaneously giving off particles & energy.
These isotopes are used for
1. dating fossils
2. diagnosing medical problems
3. tracing atoms through metabolic processes
Highlighted
area
represents
cancerous
throat tissue.
IV. Energy Levels of Electrons
A.
B.
Energy-an ability to do work or cause change
Potential energy- energy due to the position
or structure of matter
1. Electrons’ potential energy is due to their
energy level or position in an electron shell.
2. Electrons losing energy, fall to a lower shell.
C.
D.
The chemical behavior of an atom is due to
the distribution of its electrons in the
electron shells, especially the valence
electrons
An atom’s bonding capacity is called its
valence.
E.
F.
Valence electrons are those in the outermost
or valence shell of the atom.
Elements with a full valence shell are chemically
inert.
V. Electron Orbitals
A.
An orbital is a
three
dimensional
space where
electrons are
found 90% of
the time.
B.
Each electron
shell has a
specific number
of orbitals
Section 2.3 The formation and function of
molecules depend on chemical bonding
between atoms
A.
B.
C.
D.
Atoms with incomplete outer shells (8
electrons) give, take, or share electrons
Such interactions form chemical bonds.
Chemical bonds- attractive forces holding
atoms close together, making molecules
Molecules consist of two or more atoms held
together by chemical bonds.
II. Covalent Bonds
A.
B.
C.
Covalent bond- the sharing of a pair of
valence electrons by two atoms
1. The shared electrons each count as part
of each other’s valence shell
Types of covalent bonds
1. Single covalent bond- made by the sharing of
one pair of electrons
2. Double covalent bond- made by the sharing
of two pairs of electrons
3. Triple covalent bond- made by the sharing of
three pairs of electrons
Covalent bonds can form between atoms
of the same or atoms of different elements
III. Showing chemical
bonds
A.
B.
Structural
formula- a
notation used to
represent atoms
& bonding in a
molecule
ex. H H
Molecular
formula- an
abbreviated
formula for a
compound
ex. H2
IV. Polarity
A.
B.
C.
D.
Electronegativity- how strongly an atom
attracts electrons; it depends on the element’s
position in the Periodic Table; the greater the
electronegativity, the stronger the pull on the
electrons
Nonpolar covalent bond- electrons are shared
equally between the two atoms
Polar covalent bond- one atom is more
electronegative than the other; the electrons
are unequally shared
Polar compounds- compounds that have polar
covalent bonds. Due to unequal electron
sharing, one atom has a slight negative charge,
the other has a slight positive charge
E. Water- a polar molecule
1.
-Indicates a partially
positive charge on the
atom
2.
-Indicates a partially
negative charge on the
atom
3.So, the oxygen atom
is slightly negative (it
has the electrons most
of the time) and the
hydrogen atoms are
partially positive.
V. Ionic bonds
A.
B.
C.
D.
E.
F.
When there is a large difference in the
electronegativity of the two atoms, one may
take the valence electron(s) from the other.
After the transfer of electrons, the two
atoms have charges
Ion- a charged atom or molecule
Cation- a positively charged ion (J)
Anion- a negatively charged ion (L)
Ionic bond- a chemical bond formed by the
attraction between an anion and a cation.
(Breaks easily to form ions)
IONIC BONDS ILLUSTRATION
The red electron on sodium is transferred to Chlorine.
This leaves 8 electrons in the remaining outer electron
shell for sodium and completes the valence shell for
chlorine with the 8th electron, making an ionic bond.
VI. Weak bonds
A.
B.
C.
Chemical bonds vary in strength. Going from
strongest to weakest, the order is
covalentionichydrogenVanderwaals
Covalent bonds form most of a cells bonds.
Weak bonds are used to
1. reinforce the shapes of large molecules
2. help molecules adhere to each other
B.
C.
Hydrogen bonds- form when the hydrogens of
one polar molecule are attracted to strongly
electronegative atoms of a nearby polar
molecule.
In cells, nitrogen and oxygen are electronegative
Hydrogen bonds between polar
molecules
D.
E.
F.
Unequally shared electrons in a molecule
or atom cause “hot spots” of positive or
negative charge
Vanderwaals interactions- are attractions
between molecules that are close together
as a result of these “hot spots”.
This is what gives us the strong attracti ve
forces between the ridges on a geko’s toe
pads & a wall.
VII. Molecular shape & function
The function of a molecule depends on its shape.
Examples include DNA & proteins.
B. The shape of a molecule depends on the positions
of the valence electron orbitals for each of the
atoms in the molecule.
C. In covalent molecules, the s & p orbitals may affect
each other (hybridize) to make specific shapes.
A.
Molecular Shapes
Biological
molecules
recognize & react
with each other
based on the
specific shape of
the molecules.
E. Molecules with
similar shapes can
have similar
biological effects
(see picture to the
left)
D.
VIII. Section 2.4 Chemical Reactions
A.
B.
C.
D.
E.
F.
Chemical reactions- making & breaking
chemical bonds to make new molecules
Reactants- starting molecules of a reaction
Products- final molecules of a reaction
Some chemical reactions go to completion
& all reactants are converted to products
Some reactions are reversible; the
products of the forward reaction become
the reactants of the reverse reaction.
Chemical equilibrium- is when the forward
& reverse reactions occur at equal rates.
G.
H.
G.
Some reactions are not reversible but may
be paired with an opposite reaction. An
example is photosynthesis & respiration.
Photosynthesis- sunlight powers the
conversion of carbon dioxide and water
into glucose and oxygen.
CO2 + H2O  C6H12O6 + O2
Respiration- the breakdown of glucose
using oxygen into carbon dioxide, water,
and usable energy
C6H12O6 + O2  CO2 + H2O
Acknowledgements: This presentation is drawn almost entirely from the materials provided by Reese Campbell 8 th ed. DVD materials & notes