The Chemistry of Life
Objectives
• Relate the particle structure of an
atom to identify elements.
• Relate the formation of covalent
and ionic bonds to the stability of
atoms
• Distinguish mixtures and
solutions
• Define acids and bases and relate
their importance to the biological
systems
Chemistry
• All matter is made of
mass
• Atoms are the smallest
particle of an element
that cannot be broken
down to anything
smaller without losing
its properties as an
element
Elements…
•
Are pure
substances that
cannot be
broken down
into smaller
particles by
ordinary
chemical
means
Atom1. Contains 3 subatomic particles:
Charge
Location
Mass
a) Proton
+
Nucleus
1 amu
b) Neutron
0
Nucleus
1 amu
c) Electron
-
Orbital
0 amu
• 1 amu- Atomic Mass Unit
• 1 amu= 1x10 -24 g
• 0 amu= 1x10 -28 g
Atom
2. Locations in an Atom:
• A) Nucleus: located in the center of an atom.
Contains the protons and neutrons
• B) Orbital (shell): Path around the nucleus that
contains electron(s)
Periodic Table
• -The table of all known elements in their
non-reacting state.
• Position in the table indicates certain
information or properties about the
element
• Ex) non-metals are located more on the right
of the table
• Number in each box
that increases by 1
when the table is read
left to right is the
atomic number
• The other number
found in the box
(usually larger than
the atomic number
and is a decimal) is the
atomic mass
3 Rules
1. Atomic number =
# of protons
2. # of protons= # of
electrons in a nonreacting atom
3. Atomic Mass =
#of protons + # of
neutrons
• Each atom has a
set number of
orbitals in a nonreacting state and
each orbital can
hold a certain
number of
electrons
Orbitals
1. 1st orbital can have at most
2 electrons
2. 2nd orbital can have at most
8 electrons
3. 3rd orbital can have at most
8 electrons
• Last element in each
row of the periodic
table has a full outer
orbital
Ex) HE- atomic #2 has 2
electrons
• Atoms in a nonreacting state are
neutral (# protons=# of
electrons)
Isotopes
• Naturally occurring
atoms with the
same # of protons
and electrons but a
different number of
neutrons (different
mass)
• Usually radioactive (unstable)- used in
medicine and dating objects
Ex)
• 1H-mass 1 amu (Protium) normal
hydrogen
• 2H- mass 2 amu (Deuterium)  isotope of
hydrogen
• 3H- mass 3 amu (Tritium)  isotope of
hydrogen
Orbital Diagrams
• Diagrams used to show placement
and number of subatomic particles
found in an atom of an element
(Helium and Magnesium examples on board)
Ions
• Ions are atoms with a net electrical charge
• Atoms with a full outer orbital are called
stable and will not react with other atoms
(Fluorine and Aluminum examples on board)
• Naturally
occurring stable
atoms are all in
the last column
of the periodic
table and are called noble gases
• Ex) He, Ne, Ar, etc.
• Naturally all other atoms strive to
become stable and to do this they
need to gain or lose electrons which
will make the atom charged
Bonds
• When atoms become stable by losing
or gaining electons the tend to form
bonds or attractions with other atoms
• There are 3 major types of bonds:
ionic, covalent, and hydrogen
(draw picture)
1) Ionic Bonds
• Bond formed between 2 ions of opposite charges
• Ex) Na- 2, 8, 1 electrons in its orbital
• Ex) Cl- 2, 8, 7 electrons in its orbital
• Na needs to lose 1 electron to become stable
• Cl needs to gain 1 electron to become stable
Na +1 -----x---- Cl-1 NaCl join in an ionic bond
forming table salt
• Ionic bonds tend to be strong but can be easily
broken when put in water
2) Covalent Bonds
• Bond formed by atoms sharing electrons
• Ex) H- 1proton, 1electron, 0neutrons
(draw picture)
• Covalent bonds formed by atoms sharing electrons
equally is called a non-polar covalent bond
Ex. Cl 17 protons, 17 electrons, 18 neutrons
H- 1 proton, 1 electron, 0 neutrons
(draw picture)
•
Covalent bond formed by atoms sharing
electrons unequally is called a polar
covalent bond
Covalent bonds are shown by lines
between the atoms
•
•
Ex. H-H, H-Cl
• Each line represents one pair of
shared electrons
• Oxygen needs 2 electrons to fill outer orbital
O=O called double covalent bond
Non-polar due to atoms sharing electrons
equally
• Nitrogen needs 3 electrons to fill an outer
orbital
• N N is called a triple covalent bond
• Non-polar due to atoms sharing electrons equally
• Carbons need 4 electrons to fill the outer orbital
(draw picture)
3) Hydrogen Bonds
• Bond formed between 2 water molecules
• Not formed directly due to atoms trying to
become stable
• Hydrogen bond is very weak and is constantly
made and broken
(draw picture)
Effects of the Hydrogen bond
a) Surface tension: Bond holds together
separate H2O molecules. (water insects
and paper clip)
b) Capillary Action:
allows H2O
Molecules to
“defy” gravity and
move up very
small tubes using a
combination of
cohesion and
adhesion. (plants
use this motion to
get water up from
roots)
c) Resistance to Temperature Change: very
high resistance. Keep bonds of H2O
from freezing too quickly and heating
up too quickly
d) High Vaporization Temperature: a high
temperature is needed to evaporate
water. Will take a lot of energy (heat)
from the sun or other source.
e) Lower Density in Solid Form than in Liquid Form:
ice floats as H2O starts freezing, molecules move
closer together becoming more dense (normal), but
at about 4oC molecules start spreading out to
accommodate the hydrogen bonds making it less
dense
Chemical Reaction
• Process of
breaking bonds,
creating bonds,
or a combination
of both and at the
same time
rearranging
atoms
(draw example)
Chemical Formula
• Method used to
indicate # of
atoms of each
element in a
molecule.
• Ex) H20, CO2,
CH4, C6H12O6
Structural Formula
•
Method used to
indicate # of
atoms of each
element the
shape of the
molecule, and
type of bonds
(draw example)
Balancing Chemical Reaction Equations
•
•
•
An equation is balanced when it has
the same number of atoms of each
element on each side of the arrow
Ex. H2 + Cl2 HCl
Not balanced- 2 atoms of H and 2
atoms of Cl on left and only one of
each on the right.
Rules to Balancing Equations
1. Cannot change subscripts
Ex. I need 4 atoms H to balance and
equation
H2  cannot change this number to 4
(H4 NO WAY)
Rules to Balancing Equations
2. Can only change coefficient and it
effects the whole molecule
Ex. I need 6 atoms of hydrogen to
balance an equation
1 H2 3H2 3 H20  will give you 6
atoms of H+3 of oxygen
Rules to Balancing Equations
3. Cannot put a coefficient in the middle of a
molecule
Ex. I need 2 atoms of oxygen to balance an
equation
H20 I could not do the following H220
I could do 2 H20