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Chapter 2
Atoms and Molecules: The Chemical
Basis of Life
Inorganic vs. Organic

Inorganic compounds - simple
substances that do not contain
carbon


Ex. water
Organic compounds – carboncontaining compounds that are
large and complex

Ex. Glucose
Organic compounds all contain:
Ca
r
c
0%
bo
n
0%
Zin
6%
en
4.
Ni
tro
g
3.
94%
n
2.
Oxygen
Nitrogen
Carbon
Zinc
Ox
yg
e
1.
Elements


Elements – substances that cannot be
broken down into simpler substances
4 elements responsible for more than
96% of the mass of most organisms:





Oxygen
Carbon
Hydrogen
Nitrogen
See table 2-1 in book p.24
Other elements that make up living
organisms







Calcium
Phosphorus
Potassium
Magnesium
Sodium
Iron
Sulfur
What is the most abundant
element in the human body?
12%
en
Ni
tro
g
Hy
dr
og
en
6%
bo
n
4.
29%
Ca
r
3.
53%
n
2.
Oxygen
Carbon
Hydrogen
Nitrogen
Ox
yg
e
1.
Chemistry Quick Review


Atom – the smallest portion of an
element that retains its chemical
properties
Subatomic particles:






Electron – negative charge
Proton – positive charge
Neutron – uncharged
# of electrons = # of protons
Nucleus – protons and neutrons
Electrons – move rapidly through space
around nucleus
Atomic Number


Each kind of element has a fixed
number of protons in the atomic
nucleus
Written as a subscript to the left of the
chemical symbol



Example: 8O
Oxygen nucleus contains 8 protons
Determines the atom’s identity and
defines the element
Which element has an atomic
number of 7?
He
liu
m
0%
en
0%
Ni
tro
g
0%
bo
n
4.
Ca
r
3.
og
en
2.
Hydrogen
Carbon
Nitrogen
Helium
Hy
dr
1.
100%
The Periodic Table


Chart in which elements are
arranged in order by atomic number
Can be used to determine electron
configurations

Bohr model – shows the electrons
arranged in a series of concentric
circles around the nucleus
Bohr Model
What element
is this?
Atomic Mass



Mass of protons + neutrons
Mass of electron = 1/1800 the mass of
a proton or neutron
Atomic mass number is a superscript
to the left of the symbol

Example:
16O
Isotopes


Atom with different number of
neutrons (different masses)
Most elements mixture of isotopes


Ex. Carbon-12, Carbon-14
Mass of element is average of the
masses of its isotopes

Atomic mass of Carbon = 12.011
Isotopes

Radioisotopes – unstable isotopes



Tend to break down (decay) to a more
stable isotope
Emit radiation when they decay
Ex. Carbon-14 decays to Nitrogen
Electrons move in orbitals



Orbitals are more like “electron
clouds”
The farther away from the nucleus,
the more energy the electrons have
Valence electrons – the most
energetic electrons

Occupy valence (outer) shell
Chemical Reactions



Valence electrons participate in
chemical reactions
When valence shell is full, it is
stable
When valence shell is not full,
atoms tend to lose, gain, or share
electrons
To be full, the first electron shell
has how many electrons?
0%
0%
6%
0%
18
5.
8
4.
4
3.
94%
2
2.
1
2
4
8
18
1
1.
Compounds and Molecules


Atoms combine to form compounds
and molecules
Compounds - 2 or more different
elements combined in a fixed ratio


Ex. NaCl (table salt)
Molecules - 2 or more atoms
combine chemically

Ex. O2, DNA
Are all molecules compounds?
71%
29%
No
2.
Yes
No
Ye
s
1.
Molecule or Compound? O2
24%
Bo
th
Co
m
po
un
d
le
12%
cu
3.
ol
e
2.
Molecule
Compound
Both
M
1.
65%
Chemical Formulas


Represents chemical composition
Simplest formula – most simple ratio


Molecular formula – actual numbers
of each type of atom per molecule


Ex. NH2
Ex. N2H4
Structural formula – shows
arrangement of atoms

Ex. Water
H–O–H
Chemical Equations



Reactants – participate in reaction
Products – formed by the reaction
Example – cellular respiration
C6H12O6 + 6O2 -> 6CO2 + 6H2O + Energy
Chemical Bonds


Valence electrons dictate # of
bonds
2 types of chemical bonds:


Covalent – atoms share electrons
Ionic - attraction between positive
cations and negative anions

Transfer electrons
Covalent Bonds

Ex. H2 gas




Each atom has 1 electron
2 electrons fill valence shell
Both atoms attract the electrons
(share)
Valence shell is full w/ 2 electrons
Types of Covalent Bonds



Single covalent bond – 1 pair of
electrons is shared
Double covalent bond – 2 pairs of
electrons shared
Triple covalent bond – 3 pairs of
electrons shared
Covalent Bonds





Electronegativity - measure of atom’s
attraction for shared electrons in chemical bonds
Oxygen, Nitrogen, Fluorine, Chlorine very
electronegative
Can be polar or nonpolar
Similar electronegativities = nonpolar bonds
Different electronegativities = polar bonds

Electrons are pulled closer to the nucleus of the atom
with the higher electronegativity
Polar Molecules



Molecule with one
or more polar
covalent bonds
One end with a
partial positive
charge and other
end a partial
negative charge
Ex. Water (p.31)
A water molecule is polar because
24%
..
ll .
ta
or
bi
el
Th
e
Th
e
el
ec
tro
ns
ec
tro
ns
or
bi
or
bi
tt
tt
he
he
...
...
0%
ct
ro
ns
3.
76%
el
e
2.
The electrons orbit
the H atoms more
closely
The electrons orbit
the O atom more
closely
The electrons orbit
all atoms equally
Th
e
1.
Ionic Bonds


Ionic compound – consists of anions
and cations bonded together
Ex. NaCl (p.31 & 32)



Na – 1 valence electron
Cl – 7 valence electrons
Cl takes electron from Na to complete
valence shell
Hydrogen Bonds


Weak attractions
Important in determining the 3-D
structure of large molecules


DNA
Proteins
Why are hydrogen bonds essential
to the function of DNA?
33%
n.
..
st
ro
ar
Th
ey
lo
al
Th
ey
e
w
th
th
e.
..
e.
..
33%
ep
3.
33%
ke
2.
They keep the 2
strands tightly
bonded together
They allow the 2
strands to
separate for
replication
They are strong
bonds
Th
ey
1.
Redox Reactions




Reaction that involves electron
transfer
Cellular Respiration and
Photosynthesis
Oxidation – atom/ion loses electron
Reduction – atom/ion gains electron
Water



70% of total body weight
Reactant/product in many chemical
reactions
Solvent for most biological reactions


Hydrophilic – react with water
Hydrophobic – not disrupted/dissolved
by water
Which of the following substances
is hydrophobic?
33%
Oi
l
33%
Su
ga
r
3.
t
2.
Salt
Sugar
Oil
Sa
l
1.
33%
Properties of Water



Cohesive – water molecules stick to
each other
Adhesive – water molecules stick to
other substances
Capillary action – cohesion and
adhesion working together


Water will move against gravity in a
narrow tube
In plants, water moves from soil to roots
Properties of Water

Surface tension – water molecules
crowd together at the surface


High specific heat


strong layer
Maintains a stable temperature
High heat of vaporization

Much heat required to change to water
vapor
Acids and Bases

Acids – proton donors






Acid -> H+ + Anion
Acidic solutions have higher hydrogen
ion concentration
Turn blue litmus paper red
Sour taste
HCl – inorganic acid
Acetic Acid – from vinegar, Lactic Acid
– from sour milk (organic acids)
Acids and Bases

Bases – proton acceptors






Base -> OH- + Cation
Basic solutions have lower hydrogen
ion concentration
Turn red litmus paper blue
Feel slippery to the touch
Ex. Sodium Hydroxide, Ammonia –
inorganic
Purine and Pyrimidine – organic
pH scale




Logarithmic expression of the
hydrogen ion concentration of
solution
7 = neutral
Below 7 = Acid
Above 7 = Base