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Matter  anything that takes up space and has mass
Element  a substance that cannot be broken down
(periodic table)
Compound  2 or more different elements (Ex. H20,
NaCl); O2 is NOT a compound
4 elements that make up 96% of
living things: C, H, N, O
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Essential Elements  necessary for living
things (C, H, N, O); make up 96% of all living
things
Trace Elements  required by an organism but
only in minute quantities
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Atom – smallest unit of matter that still has all
the properties of that element
- made up of a nucleus (protons (+)
and neutrons (no charge)) and an
electron cloud (electrons (-))
Atomic Number – number of protons in an atom;
unique to each element
Mass Number – sum of protons plus neutrons
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Carbon-14 Dating
Isotopes are atoms of the same element that
have different numbers of neutrons.
Radioactive isotopes are useful in carbon
dating. As they decay, it converts a neutron
to a proton…thus changing the identity of
that atom. Ex. C14 dating (carbon decays
into nitrogen at a fixed rate)
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Energy  ability to cause change by doing work
Potential Energy  energy that matter stores due to
structure or location (stored in BONDS!)
Kinetic Energy  energy of motion
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Electrons FURTHER away from nucleus have MORE energy
Potential Energy –
Energy that matter
stores due to
POSITION or
LOCATION
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When the outer shell is filled, the atom is
stable and unreactive. When it is not full,
the atom is reactive and can bond with
other atoms.
Electrons are arranged in orbitals that give
an atom its 3-D shape
Electrons can move energy levels by
absorbing or releasing energy
Valence # = # of
UNPAIRED electrons in
outer shell; this is typically
the number of bonds the
atom can form
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Covalent Bonding =
SHARING ELECTRONS
Covalent bonds form
molecules. Know how
many bonds each
molecule forms:
Carbon – 4
Nitrogen – 3
Oxygen – 2
Hydrogen – 1
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Electronegativity = how strongly an
atom “pulls” the electrons in a
covalent bond
STRONG electronegative molecules:
Oxygen, Nitrogen, Chlorine
WEAK electronegative molecules:
Hydrogen, Carbon
WRITE THIS
DOWN!!!
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Polar Molecules = Uneven
distribution of charge; leads to one
side slightly more positive and one
side slightly more negative…overall it
is uncharged though
Polar Covalent Bonds  e- NOT
shared equally, found between strong
and weak electronegative atoms
NonPolar Covalent Bonds  eshared equally; found between either
the same atom (O=O) OR between
atoms of similar electronegativities
(CH3)
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Ionic Bonds = Giving away/
Taking electrons; Results in ions
(charged atoms) which are opposite
charges and thus attracted to each
other
GENERALLY….ionic
bonds are between a
metal and non-metal.
Called SALTS!
Cations = positively
charged ions
Anions = negatively
charged ions
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ONE water
molecule can
form H
bonds with
up to FOUR
other water
molecules.
Hydrogen Bonds = Attraction of the
hydrogen atom of one molecule (positively
charged) to the oxygen/ nitrogen of another
molecule (negatively charged)
Hydrogen bonds and polarity are what give
water a lot of its properties.
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Because e- are in constant motion,
they can form “hot spots” in
molecules; interactions between + and
– hot spots are van der waals
interactions
These can occur between different
parts of the same molecule, or
between 2 different ones.
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Chemical Reactions = making and
breaking chemical bonds
Reactants = what you start with
Products = what you end up with
Chemical Reactions are affected by the
concentration of the reactants….the more
reactants, the more collisions, and therefore
more product is produced faster!
Chemical Equilibrium = the rate of
formation of products is the same as the rate
of formation of reactants; no net change on
either side of the equation
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¾ of the Earth’s surface is covered in water
Cells are 70-95% water
Water is unusual because it exists in 3 states of
matter: solid (ice), liquid, gas (water vapor)
Water is a reactant in many chemical reactions
Water is the reason we have life on Earth!
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Remember, water is a POLAR molecule, so
the oxygen side has a slight NEGATIVE
charge and the hydrogen sides have a slight
POSITIVE charge.
Each water
molecule
can H-bond
with up to 4
neighbors
Four main properties of water:
- Cohesion
- Ability to stabilize temperatures
- Expansion upon freezing
- Versatility as a solvent
-Polarity
-Hydrogen Bonds
- H2O
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Cohesion  water sticking to water; it is due to hydrogen bonds; this helps to
allow water to move up through the roots in plants
Adhesion  water sticking to another substance; ex. Water sticking to the inside of
xylem in plants in order to be transported up
Surface Tension  the force needed to break the surface of a liquid; this allows
some organisms to “walk on water”; water has a very HIGH surface tension
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Kinetic Energy = Energy of motion (all
molecules are in constant motion, so they
have kinetic energy)
Heat = TOTAL quantity of kinetic energy of a
body of matter
Temperature = AVERAGE kinetic energy in a
body of matter
Specific Heat = Amount of heat needed to
change 1 gram of a substance by 1 ºC
Measuring Temperature
Calorie = amount of heat it takes to raise 1g of water 1º C
Kilocalorie (kcal) = amount of heat it takes to raise 1000g (1 kg) of water 1º C
Joule = 1 J = 0.239 cal
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Vaporization or Evaporation  liquid to gas  must move fast enough to
overcome attraction to other molecules
Adding heat increases the rate
Heat of Vaporization  amount of heat a liquid absorbs for 1g of it to be
converted to a gas; Water has a HIGH heat of vaporization
Evaporative Cooling (surface cools/ sweating)
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Water – the hydrogen bonds are constantly breaking and reforming so the
molecules can be packed more tightly together….therefore more dense than ice
Ice – the hydrogen bonds are stable, so the molecules are more spread
out….therefore less dense – so they FLOAT on liquid water; due to the structured
crystalline lattice; ice can also act as an insulator (think of lakes in the winter)
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Solution = Homogeneous mixture of 2
ore more substances
Solvent = Dissolving agent (usually
water)
Solute = substance getting dissolved
Aqueous Solution = solution where
water is the solvent
Water is sometimes considered the
“Universal Solvent” ….however , this
is inaccurate because it does not
dissolve everything. However, it is a
great solvent due to the polarity of the
water molecules.
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Hydrophilic = water-loving; ionic and polar
bonds are hydrophilic; will dissolve in water
Hydrophobic = water-fearing; non-ionic and
non-polar substances; does not mix with water
(ex. oil)
Waxy cuticle of
a plant is
hydrophobic
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Mole (mol) = grams of a substance equal to the molecular weight
Molecular weight = sum of all the weights of all the atoms in a molecule
The advantage of measuring in moles is that one mole of any substance has
exactly the same number of molecules as 1 mole of any other substance
- The # of molecules in a mole = Avogadro’s Number = 6.02 x 1023
SO….one mole of any substance will have the same # of molecules, but will
weigh different amounts!
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How to make a 1M solution of Sucrose 
- First find the molecular weight of the molecule
- Then you would dissolve 342g of sucrose in water until the whole
solution reached 1 liter. That would give you a 1M solution.
- IF you wanted to make a 2M solution, you would multiply the
molecular weight by 2 (684g), dissolve it in water, and bring the
solution to 1 liter.
- IF you wanted to make a .5M solution, you would multiply the
molecular weight by .5 and then add that much solvent (it would be
171g), etc
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Make a 1M solution of C6H12O6
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C = 12 daltons x 6 = 72
H = 1 dalton x 12 = 12
O = 16 daltons x 6 = 96
180 g = molecular weight
So, dissolve 180g of C6H12O6 in water and fill it up
to 1 liter.
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Make a .5M solution of C12H22O11
C = 12 daltons x 12 =
144
H = 1 dalton x 22 =
22
O = 16 daltons x 11 = 176
342 g = molecular weight
So, multiply 342g by .5 (171g) and then dissolve
171g of C12H22O11 in water and fill it up to 1 liter.
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Make a 1.5 M solution of C4H6NO2
C = 12 daltons x 4 =
 H = 1 dalton x 6 =
 N = 14 daltons x 1 =
 O = 16 daltons x 2 =

48
6
14
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100 g = molecular weight
So, multiply 100g by 1.5 (150g) and then dissolve
150g of C4H6NO2 in water and fill it up to 1 liter.
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-This is a reversible reaction, but does not occur frequently in nature.
-H+, OH-, and H3O+ are reactive and can affect cells
This occurs when the hydrogen atom that is attached to one water molecule
leaves and goes with another water molecule. The electron does NOT come
with it….so the molecule it left is now negative (lost the positive of the H)
and the molecule it joined is now positive!
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pH scale – measures the H+ concentration in a
solution; 1-14; a change in one number
represents a 10 fold difference…6 is 10 times
more acidic than 7; 6 is 100 times more acidic
than 8
Acids – donate H+ to solutions (increase H+
concentration); the more acidic a substance is,
the lower the pH (1-6)
Bases – accept H+ to solutions (decreases H+
concentration); bases have a higher pH (8-14);
have higher OH- concentrations
What is the pH??
[H+] = 10 - 8
[OH-] = 10 - 3
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Buffers are solutions that can act as
acids or bases (donate or accept H+)
to minimize changes in pH. They
are very important in living systems.
One common buffer in human blood is
carbonic acid. It acts as an acid and base
pair either accepting OR donating H+
depending on the need of the blood.
Can act as an acid and DONATE H+
Can act as a base and ACCEPT H+
Response to a RISE in pH →
H2CO3
↔
HCO3-
←Response to a DROP in pH
H+ donor
(acid)
Carbonic Acid
H+ acceptor
(base)
Bicarbonate ion
+
H+
Hydrogen
ion
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