I Biology I
Lecture Outline
Basic Chemistry of Life
References (Textbook - pages 20 - 36: Lab Manual - pages 8 - 11)
Matter
Elements
The Atom
Atomic Theory and Definition
Atomic Symbols
Anatomy of an Atom - Subatomic Particles
Atomic Number and Mass Number
Electrons and Energy Levels
The Periodic Table of Elements
Isotopes
Compounds and Molecules
Definitions
Ionic Bonding
Covalent Bonding
Water and Hydrogen Bonding
Acids and Bases
I Biology I
Lecture Notes 3
Basic Chemistry of Life
References (Textbook - pages 20 - 36: Lab Manual- pages 8 - 11)
Organisms (living things) are made of chemicals. Organisms are made of the same
chemicals as inanimate (non-living) things. In other words, many of the same chemicals
that make up your body can be found in rocks, water, asphalt, steel, etc.
Everything in and around you is "chemistry". You might actually call your physical
body a "bag ofchemicab'"
An understanding of the basic principles of Chemistry is necessary to understand
biology. So, we will spend the next lecture or two studying afew important topics of
chemistry.
Webster's Dictionary definition of chemistry - science dealing with composition and
properties of substances, and with the reactions by which substances are produced or
changed.
Matter
1. Matter is defined in our textbook as - anything that occupies space and possesses
mass.
2. Space is defined in Webster's dictionary as - the continuous 3- dimensional expanse in whicb aU things are contained. 3. Mass is defined in Webster's dictionary as - a quantity of matter of indefinite shape and size, a lump. 4. There are 2 kinds of matter.
A. Living - like you and other animals and plants
B. Non-living - like rocks. aspbalt, water, steel, etc.
5. Matter exists in only 3 distinct states. One state can be converted to another by
addition or subtraction of heat.
A. Solid - HOH exists as a solid (ice) below 31 degrees F
B. Liquid - HOH exists as a liquid between 31 degrees F and 111 degrees F
C. Gas - HOH exists as a gas (water vapor) above 111 degrees F
Elements
I. All matter (both living and non-living) is made up of elements.
2. Our textbook defines an element as - a substance that cannot be broken down
into simpler substances with different properties by ordinary chemical means.
3. A property is this definition is a physical or chemical trait such as melJing point,
solubility, density, etc).
4. Chemists have identified 92 naturally occurring elements. Some examples include
hydrogen, helium, nurogen, carbon, etc.
5. Of these 92 naturally occurring elements, only 6 elements are basic to life and make
up about 95% of the body weight of organisms. These six are: (CHNOPS)
A. Carbon - primary element of all organic molecules including carbohydrates,
lipids, and proteins.
B. Hydrogen - a component of most organic molecules and in ionicform
influences the acidity of body fluids.
C. Nitrogen - a component of proteins and nucleic acids (genetic material)
D. Oxygen - major component of both organic and inorganic molecules and as a
gas is essential to the oxidation of glucose and other food fuels during which
cellular energy ATP is captured.
E. Phosphorus - present as a salt and in combination with calcium in bones and
teeth. Also present in nucleic acids.
F. Sulfur - a component of proteins, especially proteins of muscles.
The Atom
l. Atomic Theory and Definition
A. The atomic theory states that ekments consist of tiny particles called atoms.
B. The word atom comes from the Greek word "atomos" which means uncut or
indivisible
c.
Our textbook defines an atom as - the slIUlllest part of an element that displays
the properties of the element.
D.Atoms are SIfUlU. A line of about 1 million wouldjit in the period at the end of a
sentence in our textbook. 100 million arranged in a row would measure about 1
inch.
2. Atomic Symbols
A. Scientists have developed a 1 or 2 letter symbol as an abbreviation for each kind
of atom.
B. These abbreviations are called the Atomic Symbol. Note that the Atomic Symbol
is also used as the symbol for the co"esponding element.
C. Examples of atomic symbols include:
Hydrogen- H
Radon - Ra
Sodium - Na
Carbon - C
3. Anatomy of an Atom - Subatomic Particles
Provide Handout or Model or Hydrogen and Helium Atoms (tbr slmplr!lt two atoms)
A.Physicist's have split atoms apart and identified over 100 kinds of subatomic
particles. We will only study 3.
Protons
Neutrons
Electrons
B. Protons 1) arefound in the nucleus of the atom 2) possess a positive + charge. C.Neutrons 1) are found in the nucleus of the atom 2) possess a neutral 0 charge. D. Electrons
1) Are found moving around the nucleus in an area often called the electron
shell.
2) Modem theory is that electrons fonn a haze or cloud about the atomic
nucleus.
3) Electrons have a negative - charge.
4) Electrons travel around a positively + charged nucleus at varying distances
from the nucleus at high velocities.
E. Note that atoms are mostly empty space.
F. An analogy from our textbook on page 22 says:
I) If an atom were the size of afootbaU freld.
2) the nucleus would be the size of a gumbal/ at the 50 yard line in the center
of the field,
3) electrons would be tiny specks whirling around the upper stands.
4. Atomic Number and Mass Number
A. All atoms of an element have the same number ofprotons in the nucleus.
B. The Atomic Number is the number of protons in the atom's nucleus. Using our
handout, what is the Atomic Number for Helium? (2)
c.
The Mass Number is the number of protons and neutrons in the atom's nucleus.
Using our handout, what is the Mass Number for Helium? (4)
D. For a given atom, the Atomic Symbol, Atomic Number, and Mass Number are
written as follows:
E. Why are these two values important in biology? They help predict how
substances might behave in individual cells, in multi-cellular organisms, and in
the environment.
5. Electrons and Energy Levels (See Handout of Bohr Models of Atoms., page ~ ofTntbook, Mader
]0lIl ed)
A. An atom is electrically neutral when the (+) positive charges of protons in the
nucleus equals the (-) negative charges of electrons orbiting about the nucleus.
B. The Bohr Model uses a 2- dimensional approach to show electron positions and
levels. (Note - thai tire alom ;11 nality is 3-~nsio,.lll)
c. Electrons orbit around the atom nucleus at different levels or shells.
For example
see chart below with electron values taken from Bohr Model Handout.
ELEMENT
Hydrogen
Carbon
Nitrogen
Oxygen
Phosphorus
Sulfur
SheD 1
I electron
2 electrons
2 electrons
2 electrons
2 electrons
2 electrons
SheD 2
SheD 3
N/A
N/A
N/A
N/A
N/A
" electrons
5 electrons
6 electrons
8 electrons
8 electrons
5 electrons
6 electrons
D. Each shell (or level) is restricted to the number of electrons it can contain.
1) Thefirst shell surrounding the nucleus never contains more than 2 electrons.
(Note - hydrogell is tire ollly alom to POSSDS ollly I electron)
2) The second shell may contain up to 8 electrons, but never more than 8. (Note ­
oxygen has only 6 electroll$ ill second shd} while phosplrorus Ir(l$ 8)
3) The Octet Rule states - the outer electron shell is most stable when it has 8
electrons.
(Note - thue art! SOtru! alorm thai corual" 7 dedron slrdls or ,~Is)
4) The outer shell is caned the valence shell.
• The number of electrons in the valence shell determines whether an atom
gives up, accepts, or shares electrons to get to the needed 8 electrons to
achieve stability.
(ReltfDrfbu the Octet JlJJIe)
• Atoms with more than -I electrons in the outer shell have a tendency to
gain or accept eLectrons.
• Atoms withfewer than -I electrons in the outer shell have a tendency to
lose or give up electrons
• The bonding capacity of atoms during chemical reactions is determined
by the number of electrons of its outer orbiJ - the valence orbiJ.
E. Electron Energy Levels
1) Electron shells or orbits can also be thought of as energy levels.
2) For example the sulfur atom has 3 electron shells and therefore 3 energy
levels.
3) Negatively (-) charged electrons are attracted to the (+) positively charged
protons in the nucleus.
4) It takes energy to push electrons away from the (+) positively charged
nucleus and keep them in their own shell or orbiJ.
5) Thefarther away from the nucleus the more energy it takes to keep electrons
in their own shell.
6) Therefore - the farther away from the nucleus the greater the amount of
energy contained in the electron shell.
7) Question - Which orbiJ contains more energy in the sulfur atom? The
innermost electron orbiJ or the outer (3 rd) orbiJ.
8) Electron energy levels play an important role in nature in a critical biological
process called photolynthesis. (Why is pltolosymhesis ;mporttull to life on earth '/)
• During photosynthesis, energy from the sun is absorbed by certain
atoms in the plant cell and electrons are boosted to a higher electron
orbit.
(a higltu energy IneJ) (We wiU study photosynthesis in mo,e dl!1l1illaJe, in tile
semester)
• Later these electrons return to a lower electron orbit (a lower energy
level) and energy is released that is used by the plant cell to complete
photosynthesis and produce oxygen.
• Our textbook states on page 25 - "Our (humans) very existence is
dependent on the energy ofelectrons".
The Periodic Table of Elements
(See Handout ofPeriodic Table from page 9 ofLab Manum)
1. The Periodic Table ofElements is a table that organizes all the known elements
(atoms) according to certain characteristics like their atomic number, the number of
electron shells, and other characteristics.
Isotopes
I
Textbook definition -Isotopes are atoms of the same element that differ in the
"umber of neutrons. (RI!nII!mbu that aU atoms of'he same element have 'he same 1IIImber ofprotons)
2. Another definition could be - Isotopes are atoms of the same element which have the
same Atomic Number, but different Mass Numbers. WHY? (Remembuthat Ihe MaY
Number ofan atom ;s Ihe 1IIImber of protons plus the numbe, ofneutrons)
3. Isotopes differ physically, but behave in a similar fashion both chemically and
physically.
4. Many elements found in nature have several isotopes; however, a great many
isotopes have been made by man.
5. The carbon atom has three common isotopes that can be written as:
Carbon 12 -
- (has 6 neurtons)
Carbon 13 -
- (has 7 neutrons)
Carbon 14 -
- (has 8 electrons, radioactive, and unstable)
6. There are many medical uses of radioactive isotopes.
A. To diagnose cancer and other diseases
B. To treat cancer.
C. To sterilize medical and dental products
Compounds and Molecules
1. Compounds and molecuks are the result of chemical reactions where atoms of different elements bond together. 2. Atoms do not normally exist singly or alone.
3. Atoms are most often joined to other atoms to form molecules.
4. Atoms forming molecules are held together by chemical bonds.
5. Recall these two facts we have already discussed.
A. Chemical reactions take place primarily between electrons in the outer shell or
outer energy level of separate atoms.
B. The bonding capacity of atoms is determined by the number of electrons in its
outer shell - the valence shell.
6. Defmitions
A. Textbook definition of compound - a compound exists when 2 or more elements
have bonded together. (EumpiH are NaCI - "ble ,aIt. H2O- Water, CH4 - methane)
B. Textbook states - a molecule is the smaDest part of a compound that still has the
properties of the particular compound.
C. A molecule is composed of 2 or more atoms.
D. A glass of water can be said to contain the compound water, including lOs of
thousands of water molecules
E. In practice these two terms (compound and molecule) can be used interchangeably. F. A chemicalformula tells the number of each kind of atom in a molecule.
G. One molecule of glucose can be written as
Indicating 6 atoms of carbon
12 atoms of hydrogen
6 atoms of oxygen
H. The electron bonds contain energy.
I. During chemical reactions where glucose is broken down, electrons shift their
energy levels and energy is released.
7. Ionic Bonding
(see HaDdout of Figure 2.7, page 26 of TeItbook, Mader 10'" Ed.)
A. Textbook defines an ionic bond as - a chemical bond in which ions are attracted
to one another by opposite charges.
B. Textbook definition of ion - a charged particle that carries a negative or positive
charge.
C. Example of Ionic Bonding - Na (sodium) and CI (chloride) 1) Na (sodium) atom • Has only 1 electron in its outer (3 rd) sheD and tends to be an electron
donor
• Once it gives up this electron its 2"" shell with 8 eleclrons becomes the
outer shell (remember the octet rule).
2) CI (chloride) atom
• Has 7 electrons in its outer (3 Td) shell and tends to be an electron acceptor
• By accepting 1 more electron it its outer shell it will have 8.
3) When a Na atom and CI atom come together - an electron is trans/erred
from the Na atom to the Cl atom. 4) Now both have 8 electrons in their outer shell 5) But now they are ions. WHY??? • Before bonding each atom (Na and CI) were electrically neutral
because each bad the same number of protons and neutrons.
• After bonding the Na atom now has 1 more proton (11) than electrons
(10) and therefore has a +1 charge. It is positively charged, can be
called the sodium ion and is written as Na+.
• After bonding, the CI atom now has 1 more electron (18) than protons
(17) and has a -1 charge. It is a negatively charged, can be called the
chloride ion and is written as CI-.
6) NaCI or table salt is an ionic compound held together by ionic bonding - the
attraction between positive and negative ions.
7) Table salJ occurs in soldforln as a 3-dimensionallottice of crystals.
8) When mixed with water, table salt dissolves and the ions separate in the
water.
8. Covalent Bonding (see Handout of Figure 2.8, plge 27 of Tutbook, Mader IO&ao E)
A. Definition - covalent bond is a chemical bond where atoms share pairs of electrons. B. Example of Hydrogen gas - H2
J) The H atom has only J electron shell which needs 2 electrons to be
complete.
2) When 2 hydrogen atoms come together they share their electrons so that
each atom now has a complete outer shell (R~member tJr~ I" eJ~ctron shell is complete
with only 2 electrons)
3) 1 pair of electrons is shared and this is called a single covalent bond.
C. Example of Oxygen gas - 02.
1) Oxygen atoms come together and share 2 pair (4 electrons) to achieve an
outer shell of 8 electrons (called an octet) tor each oxygen atom.
2) 2 pairs of electrons are shared and this is called a double covalent bond.
D. Example of Methane - CH4.
1) 4 pairs of electrons are shared between 4 atoms of hydrogen and 1 atom
of carbon.
2) Bonds result in all 4 hydrogen atoms having a 2 electron outer shell and
carbon atom having 8 electrons in its outer shell.
Water
Water is truly a remarkable molecule and compound.
It has several properties that make it essential for the survival of life on earth.
1. Hydrogen Bonding
A. The water molecule is V-shaped (DTtlW 011 blackboaTrl) and this shape results in a
polarity (uneven charges) with the 2 hydrogen atoms being slightly (+) positive
and the oxygen atom being slightly (-) negative.
(­
B. This polarity allows a water molecule to form a weak bond with other water
mokcules and other atoms and molecules.
C. The slightly (+) positive hydrogen (Draw on bladboord) atom in one water molecule
is attracted to the sligh,tlyj -J oxygen atom in an adjacent water molecule.
I li'1e"
t.~r
/
'"
~t-
(:'y;)
~~
0
~~
(~
rJi;~ ~+
~
+
D. ..
~,d, Dte1J ~'~r,)I1
\ ~ ' f (J
~~
~
T
Due to hydrogen bonding, water molecules cling together and give water m any
of its important properties. (your latbook has tJ very good discussion of5 o/Iheu properties)
E. All living things are made of 70% to 90% waleI'.
F. Recall - we said that nonnally water melJs at 32 degrees F (0 degrees C) and
boils at 212 degrees F (100 degrees C). You can say that water remains a liquid
in a temperature range of 180 degrees F.
G. Without hydrogen bonding - water molecules would mell at -148 degrees F (­
100 degrees C) and boil at -132 degrees F (-91 degrees C). You can say that
water would only remains a liquid in a much smaller temperature range of 16
degrees F.
H. So - without hydrogen bonding most of the water on earth would be steam!!!
I. Note that hydrogen bonding is not unique to water - but occurs in other
molecules (like DNA) that possess hydrogen atoms.
1. Textbook definition of hydrogen hond - the attraction of a slightly positive
hydrogen atom to a slightly negative atom in the vicinity.
Acids and Bases
1. Chemical suhstances can be classified as being either acids or bases.
2. Acids - suhstances that dissociate in water, releasing hydrogen ions (H+). (RtMeMI defi"ition ofion - a charged (HUtick t/tal carries Q " egative or positive charge) 3. Examples of acid solutions are lemon juice, vinegar, coffee, hydrochloric acid.
4. Acids can be either weak or strong depending on the number of (H+) ions
released. Example - hydrochloric acid is a stronger acid than vinegar.
5. Bases are substances that either take up (H+) ions or release hydroxide (OH-)
ions.
6. Examples of basic solutions are milk ofmagnesia and ammonia.
7. The pH scale is used to indicate the degree of acidity or basicity of a solution.
8. The pH scale ranges from a pH of0
to a pH of14.
9. A pH of 7 indicates neutrality.
10. Solutions with a pH of below 7 are considered acidic
II. Solutions with a pH of above 7 are considered basic
12. In living organisms, the pH of body fluids is maintained at a narrow range. The
pH of human blood is always around 7.4 which is slightly basic.
13. Buffers are chemicals that take up excess (H+) ions or (OH-) ions and keep the pH
of body fluids within normal limits. Examples of buffers ?
~
-r ~:1
'-:J '.
o :: ...:4
~ ~1
""t-t­
.~
l~
".-....
~
'>
~
.:5
=i ~0
<~
\.J
0
+­
"
.:u
(
I
~
'/
U
~
q
)1
.-­
~
~
:S
.J
"­
.­.:::
-
---::t:: ~
I)
~
~
~
~
,.... ~ ~
w
~
::t--r6
cr
-&::"
J
"'---/
-.......) C
C"
~
~
-rr
UJ n:
~
:::>
II
~
~
cv
- 'i
:) i
<: -
~
~
'1­
Qj
7' ,
u .- :>
II
UO ~
/'
--
/
I
~
I
'\)
•
\
-=r: ~
IJ
..;
~
J
V
I
-.
<)
V
o~
.....
~
\I
I
~
t,;
~
A.JL "'-'
c
~
n
!
~1
a
r
11
~~~
-:::t::
-
I (
0
~
It
-
--.<)
:1
~
--,­?l
~
~
-
-::t:
c:
Q)
en Z
0 ..,.r-.
...
~
c:
...
:l
:l
(I)
"J"J
N,Q
~-
(I) ...
::J
o
.J::.:..
Q. -
IF'J
(I) ~-
o
a.
.r:
<l>
;)
-=-=,
;­
..l.
:>
...)
.J::.
CI)
c
0\ ­
c
0\ ­
-u
Q)
Q)
()
CI)
:::l
<l>
u
<l> :::l
<l> C
c:
Q)
m_
o _
... -­
'2
(5l~
CD
i
-0
"0
m
:D
o
c
n
~
6;
r
m
Q
."
~
:l:
m
m
r
!i!m
z
~
:e
=i
:l:
m
r
~
z
~
z
m
»
i:
m
en
H
4. 001
1.0010
1I,IIunl
N~
~>
I-'I
6 .. I-- Atom ic Number
4
'.m
a. rllium
Sodium
M.gnul um
t~
29.100
,o.01
POlllslum
C.l,lum
21
5c
44."
tJ
47.'0
S"ndlum Th.nlum
rib
j~
15.41
I7.U
".92
91.21
Rubidium
Strontiu m
Yttrium
Zirconi um
S7
72
S6
Y
40
Ir
Ba
La
HI
11UI
131.16
IU.2
t7UO
CIII~m
hr lum
l.nlhln um
Hlfnlu m
.7
Fr
II
Ra
89
104
(Auuiln
Propo u l
UnoUie il \)
Ac
1m)
122.)
(221)
Fllnclum
Rldium
ACiinium
'Q
t:
2S
Mn
26
Fe
~~
2.
Ni
l~
n.54
sus
SUI
sue
sus
51."
51.11
Vanadium
Chrom ium
Mangan.1I
/ron
Cob.1I
Nick.1
Copper
Nb
42
R~
46
47
.2.91
43
44
Mo
Tc
Ru
.5.•5
(")
101.1
Niobi um Molvbdenum T.chn.t lum RUlh.niu m
75
74
73
Ta
W
Re
IUJ6
116. 22
6~
IIUS
Tlnilium
WoU rim
Rh.n lum
ItU
Dlmium
58
59
60
61
Ce
Pr
Nd
Pm
Pd
Ag
30
\6.000
IUO
1O.11J
I'I hrog,n
Olyg. n
Fluor l".
1'1.""
1~
15
16
17
AI
5,
21.0'
30.m
32.066
H.m
n.'41
Al uminum
Silicon
Phosp horu l.
Sullur
Ch lo rin'
I\lIt0 11
33
34
As
Se
35
Br
11."
S.'. niu m
8,omin.
7UO
74.91
Zln'
Ge rma nium
Arlanic
1m)
Pa
.U
IWI.
21&.01
Thorium Pro llClinium Uranium
49
nhodiu m
Indium
.0
Hg
Ir
Pt
It2.2
It5.0'
"7. 0
200.61
Pi llinum
Go ld
Mer cu ry
63
64
Sm
Eu
Gd
65
Tb
Np
Pu
Iml
1m)
Am
rm)
N.pl unl um Plu lon lum Am.rlclum
Cm
(W)
Curium
11
201.39
Th,lllum
66
Dy
8'2
Pb
98
Bk
IWI
12W
36
Kr
13.10
Klvpton
.~4
Xe
12 1.76
127,t1
126.91
Il l .lD
Ant imony
T.llurlu m
Iod in.
X,null
84
15
83
Bi
fe
Po
At
86
Rn
207"1
20'.00
12101
12 10)
\",1
Llld
Bismuth
Polonium
Aml ine
It " Ill II
67
Ho
\6U4
IU.SI
Oysp roliu m Hol mium
Cf
7"'16
Ar
I
111.70
Tin
114.12
CI
53
5b
11 2.4 1
5
52
51
Sn
Ca dmium
62
50
In
Slim
Iridiu m
Ge
Cd
101.110
Au
32
P
Gallium
157.26
(147)
152.0
ISI.'J
ISUS
144.17
140.n
ICO."
Cer ium Priliodymium N.odymlum Promlthium Slm.rium Eur opium G. dolinlvm Tllblu ni
97
96
95
94
93
92
91
90
Th
31
IOU
79
I.
13
Go
n.71
41
-
10
N,
14.001
In
n.lI
9
F
C
P.lladium
18
N
•
0
lUll
r..rbon
10UI
77
7
B
,....
Mg
6
10 .n
Boron
11
2U2
~:
5
C ...I-- Sv mbol
12.011-! - Ato mic We ight
C.rb on ~ Name
Be
2Ufl
i
it
I'D
KEY
Hy dr ov· n
i
L~
i
6.940
A L1thlurn
~
He
99
Es
11.14)
eerk.llt,m C.II/ornlum Eln lllinium
68
Er
69
Tm
'U.f.
"'.21 Thulium
Erbium
100
101
fm
Md
IlU)
Fer mium
1256)
l!I~nd.'.vlum
70
Yb
71
Lu
173.04
114."
Yutrblum
lu"IIUfIl
102
10]
No
Lr
IIU)
251
Nob. ,ium l iWI, ntll,.,, 1
FIGURE 2.7
r@~&
~ J
\
J}
/
~y
chlorine atom (Cil
sodium atom (Nil)
j
sodium Ion (Na·)
I
chloride Ion (CI-) I
I
sodium chloride (NaCO
a.
b.
Formation of sodium chloride (table salt).
<t. During the formation of sodium chloride, an electron is transferred fro m the sodium atom
to the chlorine atom. At the completion of [he reaction, each atom l1las eight electrons in
the outer shell, but each also carries a charge as shown. b. In a sodium chloride Cl"ystal. ionic
bonding between Na+ and CI- causes the atoms to assume a three-dimensionallanice in
which each sodium ion is surrounded by six chloride Ions, and each chloride ion is surrounded
by six sodium ions . The res Lilt
IS
crystals of salt as in table salt .
Electron Model
~r\
(\~
I
Structural
Formula
H-
H
Molecular
Formula
H~
a. Hydrogen gas
b. Oxygen gas
H
I
I
H-C -H
CH ;
H
C.
Melhane
FIGURE 2.8
Covalently bonded molecules.
In a covalent bond, atoms share electrOnS, allowing each atom to have a com pleted
shell. a. A molecule of hyd rogen (H,) conoins two hydrogen atoms sharing a
pair of electrons. This single covalem bond can be represented in any of the th ree
ways shown . b. A molecule of oxygen (01) concains two oxygen aroms sharing [Wo
pairs of electrons. This results in a double covalent booet Co A molecule of methane
(CHJ concains one carbon atom bonded to four hydrogen atO<T1S­
Outel'