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Basic Chemistry
A basic understanding of chemistry
is extremely important to
understanding biology and
biological systems -- organisms,
their tissues and their cells are
ultimately composed of atoms and
molecules and many important
biological processes are chemical
reactions -- e.g. metabolism,
photosynthesis, growth.
Chemistry organization levels
• Molecule - 2 or more atoms bonded
together
• Atom - smallest unit of an element
• Subatomic Particle - proton, electron,
neutron
Go to PBS Try It, you will need Shockwave
Atom builder
http://www.pbs.org/wgbh/aso/tryit/atom/
terms
I am going to assume that you have had no
chemistry -- we will begin at the beginning
and just cover some very basic chemistry.
• The atom is the basic chemical unit.
Composed of 2 parts:
1. nucleus -- contains subatomic protons (+
charge) and neutrons (no charge).
2. electrons contains subatomic protons (+
charge) and neutrons (no charge).
3. Protons (+ charge)- located in the nucleus
2.1 Atoms
• Organisms are chemical machines
– one must know chemistry in order to
understand biology
• Any substance in the universe that has
mass and occupies space is comprised of
matter
– all matter is made up of atoms
2.1 Atoms
• All atoms have the same structure
– at the core is a dense nucleus comprised of
two subatomic particles
• protons (positively charged)
• neutrons (no associated charge)
– orbiting the nucleus is a cloud of another
subatomic particles
• electrons (negatively charged)
Figure 2.2 Basic structure of an atom
2.1 Atoms
• An atom can be characterized by the number of
protons it has or by its overall mass
– atomic number
• the number of protons in the nucleus
• atoms with the same atomic number exhibit the same
chemical properties and are considered to belong to the
same element
– mass number
• the number of protons plus neutrons in the nucleus
• electrons have negligible mass
2.1 Atoms
• Electrons determine the chemical behavior
of atoms
– these subatomic components are the parts of
the atom that come close enough to each
other in nature to interact
2.1 Atoms
• Electrons are associated with energy
– electrons have energy of position, called potential
energy
– Electrons occupy energy levels, or electron shells, of
an atom, which are actually complex, threedimensional volumes of space called orbitals
• orbitals are where electrons are most likely to be found
2.1 Atoms
• Electron shells have specific numbers of
orbitals/valances/ energy levels that may be
filled with electrons. Bohr Model innermost shell
holds 2 electrons, all other outer shells hold 8.
must fill inner shells before moving to next shell.
– atoms that have incomplete electron orbitals tend to
be more reactive
– atoms will lose, gain, or share electrons in order to fill
completely their outermost electron shell
– these actions are the basis of chemical bonding
2.1 Atoms
• as electrons move to
a lower energy level,
closer to the nucleus,
energy is released
• moving electrons to
energy levels farther
out from the nucleus
requires energy
Figure 2.3 The electrons of
atoms possess potential energy
With electron movement=
electron cloud
Bohr Model
• A = electron
• B= neutron
• C= proton
Elements are star stuff
• There are 92 naturally occurring kinds of
atoms or elements -- differ from one
another by having different number of
protons and neutrons in the nucleus and
different numbers of electrons revolving
around the nucleus.
• They are formed inside of stars
• 1-6 of the periodic table are in all stars the
rest are formed when a star goes
supernova
Go to handouts!
95% of living material is made of
only 4 different atoms:
1. 20% carbon (C)
2. 62% oxygen (0)
3. 10% hydrogen (H)
4. 3% nitrogen (N)
• Remaining 5% made up of 30 different
elements. e.g. magnesium, sodium,
calcium, see back of a multiple
mineral/vitamin bottle
2.2 Ions and Isotopes
• Ions – atoms that have gained or lost one
or more electrons
• Isotopes – atoms that have the same
number of protons but different numbers of
neutrons
– most elements in nature exist as mixtures of
different isotopes
Figure 2.6 Isotopes of the element carbon
2.2 Ions and Isotopes
• Some isotopes are unstable and break up
into particles with lower atomic numbers
– this process is known as radioactive decay
• Radioactive isotopes have multiple uses
– dating fossils
– medical procedures
2.2 Ions and Isotopes
• Short-lived isotopes
– decay rapidly and do
not harm the body
– can be used as
tracers in medical
diagnoses and studies
Figure 2.7 Using a tracer to
identify cancer
MOLECULES
• Atoms are not usually found by
themselves; they combine with other
atoms to form molecules -- e.g. O atoms
combine with one another to form O2-- this
is the form that oxygen generally is found - when different atoms combine a
compound is formed -- e.g. H2O, NH3,
CO2, NaCl.
• The ability of particular atoms to
combine or react with other atoms is
determined by the atom’s electron
configuration -- electrons move around
nucleus in levels called
shells, orbits, or valence levels
which representing different
energy levels. + or –
Example oxygen has 6 outer elections so
it is a +6
Shells/Orbits/Levels and electrons
• The inner most
shell/orbit/valence
level= energy level
can hold 2 electrons.
The second shell/orbit
etc and all other outer
shells/orbits etc can
hold a maximum of 8
electrons
• Atoms that have outer
electron shells that are
filled are inert =
nonreactive
• Other atoms, including
H, C, O, and N, do not
have filled outer shells
and they are reactive - they readily combine
with other atoms.
• All atoms, while having a matching
number of protons and electrons, want to
have a full outer shell/orbit. This leads to
atoms combining, sharing outer shell
electrons, to be happy. But since their
outer shell is not full all the time they stay
in balance and stable. Creating a stable
molecule
2.3 Molecules Form Bonds
•
A molecule is a group of atoms held together
by energy in the form of a chemical bond
There are 3 principal types of chemical bonds
•
1. Ionic
2. Covalent
3. Hydrogen
•
van der Waals forces are a kind of weak chemical
attraction (not a bond) that come into play when
atoms are very close to each other
2.3 Molecules
• Ionic bonds involve
the attraction of
opposite electrical
charges
• Molecules comprised
of these bonds are
often most stable as
crystals
Fig. 2.8(a) The formation of
ionic bonds in table salt
Ionic Bonds
• Ionic bonding: One way that atoms may achieve electron
stability in their outer shell is by losing or gaining
electrons and becoming ions.
• Example: sodium (Na) and chlorine (Cl) often form ions
(Na+, Cl-) -- sodium losses an electron and chlorine
picks up an electron -- since they are ions with opposite
charges, they are attracted to one another and an ionic
bond is formed -- the resulting molecule is a compound
called sodium chloride (NaCl) (table salt).
• Ionic bonds are strong when the compound exists as a
dry solid -- in water, these compounds dissolve readily
back into ions.
Consider sodium and chloride
• Consider sodium and chlorine.
Sodium is in column I so it has
one electron in its outer shell..
If it could get rid of that one,
lonely electron from level 3,
then its outer level would be
level 2 which is nice and full.
Chlorine is in column VII, so it
has seven electrons in its outer
shell, one short of a full outer
shell.
• Thus, it would be more stable if
it could grab an electron from
somewhere to fill up that one,
last spot.
2.3 Molecules
• Covalent bonds form between two atoms
when they share electrons
– the number of electrons shared varies
depending on how many the atom needs to fill
its outermost electron shell
– covalent bonds are stronger than ionic bonds
are directional
Figure 2.9 Covalent bonds
Covalent Bonds
• Covalent bonding: rather than gaining or
losing electrons, atoms may share
electrons with other atoms to fill both
their outer shells and give them
stability.
• Two types of covalent bonds: nonpolar
and polar covalent bonds.
• Nonpolar covalent bonds: equal sharing
of electrons -- e.g. H2, O2, N2
2.3 Molecules
• Some atoms may be better at attracting
the shared electrons of a covalent bond
– this creates tiny partial negative and positive
charges within the molecule, now called a
polar molecule
– polar covalent bonds form when the shared
electrons of a covalent bond spend more time
in the vicinity of a particular atoms
• Polar covalent bonds: unequal sharing of
electrons -- often occurs when atoms are of
different sizes -- e.g. NH3, CO2, H2O
• So charged sides
Polar covalent
2.3 Molecules
• Water molecules
are polar and can
form hydrogen
bonds with each
other
Figure 2.11 Hydrogen
bonding in water molecules
• Hydrogen bonds -- differ from ionic
and covalent bonds.
• Ionic and covalent bonds are bond
between atoms to form molecules - but, hydrogen bonds form
between molecules.
2.3 Molecules
• Hydrogen bonds are weak electrical attractions
between the positive end of one polar molecule
and the negative end of another
– each atom with a partial charge acts like a magnet to
bond weakly to another polar atom with an opposite
charge
– the additive effects of many hydrogen bonding
interactions can add collective strength to the bonds
• Hydrogen bonding gives water many of its
biologically important characteristics -- e.g.
high specific heat, surface tension,
cohesiveness.
• Now a message about water yes, is to
science dihydrogenoxide or H2O
2.4 Unique Properties of Water
• Water is essential for life
– the chemistry of life is water chemistry
• Water is a polar molecule
– water can form hydrogen bonds
– hydrogen bonding confers on water many
different special properties
2.4 Unique Properties of Water
• Heat Storage
– water temperature changes slowly and holds temperature well
• Ice Formation
– few hydrogen bonds break at low temperatures
• water becomes less dense as it freezes because hydrogen bonds
stabilize and hold water molecules farther apart
• High Heat of Vaporization
– water requires tremendous energy to vaporize because of all the
hydrogen bonds that must be broken
– when water vaporizes, it takes this heat energy with it, allowing
for evaporative cooling
Figure 2.12 Ice formation
2.4 Unique Properties of Water
• Water molecules are
attracted to other polar
molecules
– cohesion – when one water
molecule is attracted to
another water molecule
– adhesion – when polar
molecules other than water
stick to a water molecule
Figure 2.13 Cohesion
2.4 Unique Properties of Water
• High polarity
– in solution, water molecules tend to form the
maximum number of hydrogen bonds
• hydrophilic molecules are attracted to water and
dissolve easily in it
– these molecules are also polar and can form hydrogen
bonds
• hydrophobic molecules are repelled by water and
do not dissolve
– these molecules are nonpolar and do not form hydrogen
bonds
Figure 2.14 How salt dissolves in water
2.4 Unique Properties of Water
• Water is essential for life
– the chemistry of life is water chemistry
• Water is a polar molecule
– water can form hydrogen bonds
– hydrogen bonding confers on water many
different special properties
2.4 Unique Properties of Water
• Heat Storage
– water temperature changes slowly and holds temperature well
• Ice Formation
– few hydrogen bonds break at low temperatures
• water becomes less dense as it freezes because hydrogen bonds
stabilize and hold water molecules farther apart
• High Heat of Vaporization
– water requires tremendous energy to vaporize because of all the
hydrogen bonds that must be broken
– when water vaporizes, it takes this heat energy with it, allowing
for evaporative cooling
Figure 2.12 Ice formation
2.4 Unique Properties of Water
• Water molecules are
attracted to other polar
molecules
– cohesion – when one water
molecule is attracted to
another water molecule
– adhesion – when polar
molecules other than water
stick to a water molecule
Figure 2.13 Cohesion
2.4 Unique Properties of Water
• High polarity
– in solution, water molecules tend to form the
maximum number of hydrogen bonds
• hydrophilic molecules are attracted to water and
dissolve easily in it
– these molecules are also polar and can form hydrogen
bonds
• hydrophobic molecules are repelled by water and
do not dissolve
– these molecules are nonpolar and do not form hydrogen
bonds
Figure 2.14 How salt dissolves in water
2.5 Water Ionizes
• The covalent bond within a water molecule
sometimes breaks spontaneously
H2O 
Water
OH
Hydroxide
+
+
H
Hydrogen
• This produces a positively hydrogen ion (H+) and
a negatively charged hydroxide ion (OH-)
2.5 Water Ionizes
• The amount of ionized hydrogen from
water in a solution can be measured as pH
• The pH scale is logarithmic, which means
that a pH scale difference of 1 unit actually
represents a 10-fold change in hydrogen
ion concentration
Figure 2.15 The pH scale
2.5 Water Ionizes
• Pure water has a pH of 7
– there are equal amounts of [H+] relative to [OH-]
• Acid – any substance that dissociates in water
and increases the [H+]
– acidic solutions have pH values below 7
• Base – any substance that combines with [H+]
when dissolved in water
– basic solutions have pH values above 7
2.5 Water Ionizes
• The pH in most living cells and their
environments is fairly close to 7
– proteins involved in metabolism are sensitive to any
pH changes
• Organisms use buffers to minimize pH
disturbances
– a buffer is a chemical substance that takes up
or releases hydrogen ions
Review of water and pH
Water
constitutes two-thirds the mass of most
organisms & 75% earth's surface.
•
•
•
•
Importance of Water to Life
1. Coolant: Has high heat of vaporization,
aids body cooling.
2. High specific heat: ( H20 = 1 cal / g /
degree C ) - bodies of water stay constant
temp.
3. Transport: polarity - dissolves many
substances.
4. Habitat: Major component of internal
(organism) & external environments
Properties of Water
Property
1. Transparency
Meaning
Light passes thru water.
Importance of Property
Light reaches chloroplasts in cells &
aquatic plants.
Dissolved compounds can be brought to
Many compounds dissolve in
2. Universal Solvent
cells (via sap or blood) or move about
water.
cell cytoplasm.
3. Cohesion
4. Adhesion
Water molecules stick together
Small animals may walk on water.
due to H bonds.
Capillary action. Water pulled to top of
Water molecules stick to other
trees.
molecules.
5. Heat Capacity
Water bodies have stable temperatures.
Large amounts of energy are
Body temps can be maintained. Transfer
needed to raise temp of water.
of heat from warm to cool body parts.
Much energy needed to pull
In nature, water rarely boils so life is
6. High boiling point. water molecules apart.
spared.
7. Density of Ice
8. Evaporation
Ice is less dense than water so
Ice insulates organisms living beneath.
it floats.
Evaporation (boiling) requires
Evaporation can cool warm cells.
much energy.
Concepts of acids and bases
• Some compounds release large numbers of H+ ions
when they are in water -- these are called acids -- e.g.
HCl, H2SO4 are acids.
• On the other hand, some compounds release OH- ions
when they are in water -- these compounds are called
bases -- e.g. NaOH, KOH are bases.
• Strengths of acids and bases are measured by the pH
scale --ranges from 0 to 14 -- measures relative
concentrations of H+ and OH- ions -- pure water is
neutral and has pH = 7.0 -- anything with a pH less than
7 is an acid; pH greater than 7 is a base.
pH scale
• pH is the measure
of the
concentration of
H ions.
7 neural
0-just below 7 acidic
Just above 7 to 14
basic or alkaline
pH of Common Items
Acid rain
5.2
Drinking 6.5 – Sea
water
8
water
Ammonia 11.6 eggs
7.6-8 Orange
water
juice
7.36
8.21
3-4
Apple
juice
2.9
–
3.3
Grapes
3.54.5
2.3
Baking
soda
8.0
Human
blood
7.35- Stomach 1- 3
7.45 juice
Borax
9.2
Distilled
water
7.0
Lemon
juice
Bread
Human
saliva
6.36.6
5-6
Milk of
10.5
magnesia
• pH is important biologically because
chemical reactions in cells are pH specific
-- disruptions in pH will adversely affect
organisms -- may effect entire groups of
organisms -- e.g. acid rain and effect on
lakes.
• Organisms must maintain relatively
constant internal pH (homeostasis) to live.
• pH of blood = 7.4 -- slightly basic --serious
illness or death may result if this pH isn't
maintained within narrow limits -- -- yet H+
and OH- may be added to blood through
various metabolic processes -- how is "pH"
balance maintained?
• Blood also contains special chemical
compounds called "buffers" -- buffers pick
up excess H+ or OH- ions or release them
if they are in short supply -- by doing this,
they act to keep pH in blood (and cells)
from changing.
In the bonds we store energy
• We speak of energy as kinetic or potental
• Kinetic is the energy of movemen
While potential is energy in storage for later
use.
Energy is the ability to move things or do
work
Energy can not be created or destroyed it
can however be changed between forms.
• First law of thermodynamics
In its simplest form, the First Law of
Thermodynamics states that neither matter nor
energy can be created or destroyed. The
amount of energy in the universe is constant –
energy can be changed, moved, controlled,
stored, or dissipated. However, this energy
cannot be created from nothing or reduced to
nothing. Every natural process transforms
energy and moves energy, but cannot create or
eliminate it.This principle forms a foundation for
many of the physical sciences
Just in case you were wondering
• They is a second and even a third law
• Second Law of Thermodynamics - Increased Entropy
The Second Law of Thermodynamics is commonly
known as the Law of Increased Entropy. While quantity
remains the same (First Law), the quality of
matter/energy deteriorates gradually over time. How so?
Usable energy is inevitably used for productivity, growth
and repair. In the process, usable energy is converted
into unusable energy. Thus, usable energy is
irretrievably lost in the form of unusable energy.
- See more at: http://www.allaboutscience.org/secondlaw-of-thermodynamics.htm#sthash.n7axZkbS.dpuf