Matter, Elements,
and Compounds
Matter: Anything that takes up space and has
mass.
Element: A substance that cannot be broken
down by ordinary means. The material making
up matter.
• There are 92 naturally occurring
elements, of these 25 are
essential to life. 4 of these make
up 96% of living matter (Carbon,
Hydrogen, Oxygen, and
Nitrogen).
• The rest are called
trace elements. These
are required in minute
amounts (Zinc, Cobalt,
Iron, and Magnesium).
Small units of matter are called
atoms. Protons (+), neutrons
(0), and electrons (-) are the
subunits of atoms. Their mass
is measured
in units called Daltons.
Atomic Number: The total number of protons
in an atom.
Atomic Mass: The total number of neutrons
and protons in an atom.
Isotopes: Different atomic forms caused by
varying the number of neutrons.
Example: Normal carbon is 12, carbon isotope
is 14.
Some isotopes are
radioactive, they undergo
a transformation to gain a
stable condition. This
transformation is called
the half-life of the isotope.
Energy Levels: All
electrons have the
same mass and
charge. Electrons
closer to the
nucleus contain less
potential energy.
The farther away
from the nucleus,
the more potential
energy they contain.
Energy Levels
• These different states of potential energy are called
energy levels.
• Each level is divided into subunits called orbitals. No
more then 2 electrons can occupy the same orbital.
• The orbital closest to the nucleus contains 2 e-.
• The next level is divided into 4 orbitals for a total of 8
electrons.
Energy Levels
• Valence Electrons:
These electrons occupy
the last energy level of
an atom.
• It stands to reason that
chemical bonds will
occur here in a chemical
reaction, since it is here
where atoms come in
contact with each other.
Energy Levels
• The maximum number of valence electrons any atom
can contain is 8.
• Any number less than 8 will allow that atom to take or
give up electrons to become stable.
• Atoms that give electrons will become positive ions,
while ions that receive electrons will become negative
ions.
Chemical Bonding: Atoms will form bonds depending
on their incomplete valence shell. There are several
types of chemical bonds:
– Covalent Bond: These bonds are the strongest bonds.
They are formed by the sharing of valence electrons.
– Ionic Bond: These bonds are formed by the taking of
electrons. Anion: negative ion, Cl- and (OH). Cation:
positive ion, Na+ and (NH4)+.
– Hydrogen Bond: This bond is formed when hydrogen that
is covalently bonded to an electronegative atom is attracted
to another electronegative atom on another molecule.
Covalent Bonds
Ionic Bonds
Hydrogen Bonds
Chemical Reactions: The combination of 2 or more
elements forming a different product or products.
• Each reaction contains reactants and products. The
reactants are written on the left side of the equation,
while the products are written on the right side.
• The reactants and products must contain the same
number of atoms making the reaction balanced.
Reactants  Products
Water
and the Fitness of
the Environment
Hydrogen Bonding of Water Molecules:
Due to polar covalent bonds that hold a water
molecule together, hydrogen bonds form where the
negative oxygens and the positive hydrogens are
located.
The results of these
bonds are as follows:
1. Cohesion
• Cohesion is the
sticking together of
similar molecules.
• Water is very
cohesive.
2: Surface Tension
Cohesion allows water to
pull together and form
droplets or form an
interface between it and
other surfaces. The
measure of how hard it
is to break this interface
is surface tension.
3: Adhesion
Adhesion is the sticking of one
substance to another. Water
very adhesive. It will cling to
many objects and act as a glue.
Capillary action is an example
of cohesion and adhesion
working together to move water
up a thin tube.
4: Imbibition
Imbibition is the process
of soaking into a
hydrophilic substance.
Some examples are
water being taken into a
sponge, a seed, or
paper towels.
5: High Specific Heat
•
Specific heat of a substance is the heat needed (gained
or lost) to change the temperature of 1g of a substance 1
degree Celsius. A Kilocalorie equals 1,000 small calories.
It takes 1,000 calories to raise 1,000g of water 1 degree
Celsius.
• This high specific heat allows water to act as a heat sink.
Water will retain its temperature after absorbing large
amounts of heat, and retain its temperature after losing
equally large amounts of heat.
• The reason for this is that hydrogen bonds must absorb
heat to break.
• They must release heat when they form. The Ocean acts
as a tremendous heat sink to moderate the earth’s
temperature.
High Specific Heat
QuickTime™ and a
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• Because of high specific heat capacity, water during
the day is cooler than land. Rising air above warm
land is replaced by cooler air pushed in from the
lake. The reverse happens at night, when the
land's temperature has fallen below that of the lake;
the lake's temperature drops, too, at night, but not as
much as the land's.
6. High Heat of Vaporization
Water must absorb a certain amount of
additional heat to change from a liquid to
a gas. This extra heat is called heat of
vaporization. In humans, this value is
576 cal/g. This results in evaporative
cooling of the surface.
Alcohol has a value of 237 cal/g and
chloroform 59 cal/g.
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7: Freezing and Expansion of Water
Water is most dense at at 4
degrees C. At 0 degrees
C it is less dense. Ice
floats because maximum
hydrogen bonding occurs
at 0 degrees C.
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8: Versatile Solvent
•
Water is a major solvent in nature.
When water and another substance are
mixed, the resulting solution is call an
aqueous solution.
• Any solution that contains the following
parts: Solute (what’s being dissolved) +
Solvent (what’s doing the dissolving) =
Solution.
• Solute Concentration: The
concentration of the dissolved
materials in relation to the solvent.
This is always measured in moles.
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8: Versatile Solvent
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TIFF (Uncompressed) decompressor
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QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
• A mole is the amount of a substance whose mass in grams
is numerically equivalent to its molecular weight in daltons.
• One must first find the atomic weights of the substances
involved and add them together for the representative
molecule and change the value to grams.
• Molarity occurs when the mole (gram atomic weight of the
substance) is placed in a container nd dissolved in one liter
of water.
pH
pH: refers to the dissociation of water molecules.
The pH constant is Kw = 1.0 x 10 -14 (mol/L)
• When we have an even split of H+ and OH- ions, the pH
value is 7.
• Given the pH constant, in a solution with pH 7, the
concentration of the H+ ion is 1 x 10-7 and the
concentration of the OH- ion is also 1 x 10 -7 .
• The true definition of pH is
pH = - log [H+]
QuickTime™ and a
TIFF (Uncompressed) decompressor
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QuickTime™ and a
TIFF (Uncompressed) decompressor
are needed to see this picture.
Practice problems with pH
Problems:
1.) H+ conc = 1 x 10 -10 mol/L. Determine the pH.
2.) OH- conc = 1 x 10 -2 mol/L. Determine the pH.
Buffers
• Abrupt changes in pH are
harmful to the cell and any
living organism. In order to
minimize this harm, cells
contain buffering systems.
• In order to change the pH of
a solution, H ions must be
added or taken from a
solution
• Buffers maintain equilibrium
by adding or removing H+ as
needed.
Quick Time™a nd a
TIFF ( Unco mpre ssed ) dec ompr esso r
ar e nee ded to see this pictur e.