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Dee Chemistry of LIFE!
Nature of Matter
• Matter is defined as
anything that has mass
and takes up space.
• All matter is made up of
atoms.
• Matter has 4 states:
–
–
–
–
Solid
Liquid
Gas
Plasma
Basic Atomic Structure:
Atomic Structure
• Atom- smallest unit of
matter: can’t be broken
down by chemical means
• Made of 3 kinds of
particles
– Protons (+)
– Electrons (-)
– Neutrons (neutral)
Atoms actually
consist mostly
of empty space
(called “voids”).
If the nucleus of
an atom were
the size of a
typical marble,
the first
electron level
would be a half
mile away!
Atomic Structure
• Electron- negatively
charged particle that
orbits around the
nucleus.
• Proton- Positively
charged particle
found in the nucleus.
• Neutron- Neutral
particle found in the
nucleus.
What element would this be?
Physical changes
between the states
of matter can be
produced by
heating and
cooling.
Solids
• Lowest energy state
of matter: molecules
are least active.
• Molecules tend to
crystallize and form
lattice-like structures.
• Usually the densest
state of matter.
Lattice Structure Example:
Liquids
• Mid-level energy and
density state of matter.
• Liquid is a fluid in which
the particles are loose
and free flowing.
• Liquid’s shape is
confined to, but not
necessarily determined
by, it’s container. (unlike
a gas which MUST fill
the container it is in!)
Gases
• Highest energy state of
matter.
• Lowest density state
• Particles move freely
and basically randomly.
• Has no definite shape
or volume. Takes
shape of its container.
Plasma
• Distinct phase of matter
referring to an ionized
gas cloud in which a
certain proportion of
electrons are free,
(rather than being bound
to an atom or molecule.
• Most stars are made of
plasma, and voids
contain sparse plasma
as well.
• Plasmas are by far the
most common phase of
matter in the universe,
both by mass and by
volume.
Ions
• An ion is an atom that
has an electrical
charge,
(NO LONGER NEUTRAL)
– Losses e- = + charge
– Gains e- = - charge
• Atoms of opposite
charges are attracted
to each other
Plasma (cont.)
A plasma lamp exhibiting a
characteristic called filamentation.
The colorful filaments are a result of
electrons in excited energy states
relaxing to a lower energy state and
giving off photons of light.
Plasma’s are
electrically
conductive, and
respond strongly to
electromagnetism.
Plasma therefore has
properties quite
unlike those of solids,
liquids or gases and
is considered to be a
distinct state of
matter.
Atomic Structure
• The positive charge of
the proton keeps the
electron from flying off
into space.
• Certain classes of
elements demonstrate
increased reactivity
because of few valence
shell electrons and/or
“the electron shielding
effect”
What is “electron shielding”?
• As electron’s get further & further
away from the nucleus, the pull of the
protons on them are “shielded” by the
other electrons in their orbital’s.
Atomic
Structure
• Element- Pure substance
made of only one kind of
atom.
• Elements differ in the # of
protons their atoms
contain within their
nucleus.
• Elements usually have
the same number of
neutrons in their nucleus
as protons; however
sometimes they don’t…
Isotopes
Two extra neutrons
(white spheres)
• Isotopes- Atoms of one
kind of element which
have differing numbers
of neutrons.
• Example- three
common isotopes of
carbon are C-12, C-13,
& C-14. Each has six
protons, but C-13 has 7
neutrons & C-14 has 8!
Compounds
NaCl
• Compound- substance
made of the joined
atoms of two or more
different elements.
• Every compound can
be identified by their
chemical formula.
• Chemical formula
identifies the elements
in the compound and
their proportions.
Molecules
• Molecule- a group of
atoms held together
by covalent bonds.
• Example- one
molecule of Sulfuric
Acid is made up of:
Sulfuric acid molecule: (H2SO4)
Yellow = Sulfur
Red = Oxygen
White = Hydrogen
–
–
–
–
2 Hydrogen atoms
1 Sulfur atom
4 Oxygen atoms
Chemical formula =
H2SO4
Bonding
•
Not male or
female bonding;
CHEMICAL
BONDING!
1. Hydrogen
Bonding
2. Ionic Bonding
3. Covalent Bonding
Covalent Bonds
• Covalent bonds
are formed when
atoms share
electrons in their
outer electron
shells.
• Covalent bonds are
very strong
chemical bonds
EX: Covalent bonding in methane
X’s = hydrogen electrons
0’s = Carbon electrons
Ionic Bonding
• Positively charged ions are called “cations”
• Negatively charged ions are called “anions”
Ionic Bonding
• Ionic bonding
takes place
when an ion
(charged
particle)
becomes
attracted to an
oppositely
charged
particle and
“sticks” to it.
• The arrangement of their electrons determines
how atoms bond together.
• An atom is said to be “stable” (meaning nonreactive) when its outer electron shell is full.
(Remember the Octet rule)
• All atoms seek to become stable.
Electron Shells
• The outermost
electron shell of an
atom is termed the
“valence shell.”
• Only the valence shell
is used for bonding.
• The first electron shell
can hold only 2
electrons.
• Subsequent electron
shells can hold 8
electrons.
Valence Electrons
• Atoms with relatively
few valence shell
electrons (1-3)
attempt to give them
away.
• Atoms with a lot of
valence shell
electrons (5-7)
attempt to “grab up”
other electrons to fill
the outer shell.
Potassium Bromide is a salt used as an anticonvulsant for animals.
Hydrogen Bonding
• Hydrogen bonds are
weak bonds between
the partially negative
oxygen atom in one
water molecule, & the
partially positive
hydrogen atom in
another water
molecule.
Hydrogen
Bonding
• Two of the notable
effects of the hydrogen
bonding in water are
the phenomenons of
capillary action and
surface tension of
water.
Hydrogen bonding
• Capillary action is
one of the things
that makes life
possible on
planet earth.
Weird water molecule
• The electrons
do not spend
their time
equally around
each atom
(more protons
in the oxygen
atom exerts a
greater affinity
on the e-)
Polar ions dissolving in water
Water and the fitness of
earth for life
Water
• A water molecule (H2O), is
made up of three atoms: one
oxygen and two hydrogen.
Weird water molecule
• The electrons do not
spend their time equally
around each atom (more
time around oxygen)
• This “unequal” sharing of
electrons results in partial
charges on the water
molecule which makes
hydrogen bonding
possible.
Polarity
• Molecules with an uneven
distribution of electrical
charge are referred to as
“polar” molecules
– “polar” basically means
charged
– Because of their inherent
charge, polar molecules
can dissolve in water; nonpolar molecules can’t
(The symbol used here is a
lowercase Greek delta, and
it means partial)
How water dissolves polar
molecules
Hydrogen Bonds:
• Very weak bonds
• Responsible for giving water its unique properties
• Formed as a result of electromagnetic attraction in water
molecules
Victoria Falls, Africa: the
largest waterfall on earth
Weird ways of water
•At sea level, pure water boils at
100 °C and freezes at 0 °C.
•The boiling temperature of water
decreases at higher elevations (lower
atmospheric pressure).
•For this reason, an egg will take
longer to boil at higher altitudes.
5 weird properties of water
1. Cohesion
2. Adhesion
3. High Specific Heat
4. High Heat of Vaporization
5. Less Dense as a Solid!
(Densest at 40 degrees F.)
Why ice floats
1. Cohesion
•Attraction between particles of the
same substance
- why water is attracted to itself
Results in:
Surface tension (a measure of the
strength of water’s surface)
•surface film on water
-allows insects to walk on the
surface of water
Surface Tension
*2. Adhesion*
•Attraction between two different
substances.
- water will make hydrogen bonds
with other surfaces such as glass, soil,
plant tissues, and cotton.
•Capillary action-water molecules will
“tow” each other along - i.e. transpiration
process by which plants and trees remove
water from the soil; and paper towels
soak up water.
•(Demonstration)
.
Capillary Action
3. High Specific Heat
Amount of heat needed to raise or lower
1g of a substance 1° C.
•Water resists temperature change, both
for heating and cooling, because of the
many hydrogen bonds.
•Water can absorb or release large
amounts of heat energy with little change
in actual temperature.
Water resists temperature change
Avg. temps of San Francisco Bay
area vs. U.S. average
• Why water takes a long
time to boil (has to
absorb enough energy to
break all the hydrogen
bonds)
• Large bodies of water
(like oceans) tend to
moderate air temps.
(resist wide temperature
fluctuations.) {next slide}
4. High Heat of Vaporization
Amount of energy to convert 1g of a
substance from a liquid to a gas
• In order for water to evaporate,
hydrogen bonds must be broken. As
water evaporates, it removes a lot of
heat with it! (sweating)
Sweating off the heat
• Sweat is an efficient
cooling mechanism that
takes advantage of the
High Heat of Vaporization
of water.
• Water must absorb a lot of
heat to gain enough
energy to evaporate; this
lowers the body
temperature back down.
5. Water is Less Dense as a Solid
1.Ice is less dense as a solid than as a liquid (ice floats)
[40 degree rule]
2.Liquid water has hydrogen bonds that are constantly
being broken and reformed.
3.Frozen water forms a crystal-like lattice whereby
molecules are set at basically fixed distances.
Ice-Water
• Property of water
being less dense as a
solid than a liquid is
CRITICAL to life on
earth!
• What would happen
if ice was more dense
than liquid water?
• Hint-think about our
lakes
Ice-Water
• Most marine animals would die!
• As lake ice freezes, the top layer of
ice protects the rest of the body of
water from the cold air temps
above.
• If ice was more dense than liquid
water, it would sink as it froze.
This would continue until the
entire body of water was frozen
solid, and the marine life dead!
Water is Less Dense as a
Solid
Water
Ice
Hydrogen donors
The pH Scale
• The pH scale is based on the
concentration of Hydrogen ions
within a solution. (called
hydronium)
• The scale is inversely
proportional to the number of
hydronium ions present in the
solution.
– The more hydrogen ions, the
lower the pH
– Scale is BASE 10!!!!
• pH of 2 is 100 times more acidic
than a pH of 4… NOT twice!
Hydrogen acceptors
Acids are hydrogen donors,
Bases are Hydrogen acceptors.
• Thus, acids often
have a chemical
formula that starts
with “H”, while bases
often have a chemical
formula that ends with
“OH”.
• OH- is called a
“hydroxide ion”
• H+ is a hydrogen ion
(or just a proton)
Some strong Acids
• HCl
– (Hydrochloric Acid)
• HBr
– (Hydrobromic Acid)
• H2SO4
– (Sulfuric Acid)
• HNO3
– (Nitric Acid)
Some Strong Bases
• NaOH
– (Sodium Hydroxide)
• KOH
– (Potassium
Hydroxide)
• Ca(OH)2
– (Calcium Hydroxide)
• Ph scale is used to measure the acidity or alkalinity of
a solution.
• pH scale is defined as the negative logarithm of the
hydrogen ion concentration in a solution.
• pH scale therefore increases and decreases by
powers of 10.
Buffers
• Buffers keep the pH of
a system from
fluctuating very much.
• Without buffers, pH of a
system can fluctuate
greatly as well as VERY
quickly!
Buffers
• The internal pH of most
living cells is very close to
7, and blood pH is 7.40
– You cannot survive more than
a few minutes if blood pH
fluctuates very far.(+ or - .05)
• The most important
buffering system for
humans is the carbonic
acid-bicarbonate system
present in the blood.
• Most buffering systems are
weak acid-base pairs.