Chapter 2
Chemical Foundations for
Cells
Chapter Outline

Review of elements and atomic structure


Chemical bonding




Radioactive elements and health/medicine
Ionic
Covalent: nonpolar and polar
Hydrogen “bonding”
Properties of water
 Acids, bases, and buffers
 Chemical change
Elements (2.1, 2.3)
 Living

organisms are composed of matter
Matter is composed of elements
• Element - substance that cannot be broken down
into other substances by chemical means


Elements are made up of atoms.
Atoms join together to make compounds.
Atoms
Compounds
Elements (2.1)
 92


naturally occurring elements
Life requires ~25 of these
~96% of human body is made up of:
•
•
•
•
Carbon (C)
Hydrogen (H)
Oxygen (O)
Nitrogen (N)
Compounds (2.1)
Atoms of one element can join with
atoms of other elements to form
compounds.

•
A given compound is always made of the
same elements combined in the same ways.
•
•
•
NaCl – table salt
H2O - water
C6H12O6 - glucose
Compounds of Life

Only living organisms have the ability to
make the compounds of life:




Carbohydrates: C, H, O
Lipids: C, H, O
Proteins: C, H, O, N, S
Nucleic acids: C, H, O, N, P
Atoms (2.3)
 An
atom is the smallest unit of an element
 Atoms are composed of 3 subatomic
particles:



Protons
Neutrons
Electrons
Subatomic Particles
Subatomic Charge Mass, amu
Particle
Location in
atom
Electron
(e-)
-1
0 amu
Outside of
nucleus
Proton (p)
+1
~1 amu
Nucleus
Neutron (n)
0
~1 amu
Nucleus
Subatomic Particles and the
Elements
 Each
element has a unique number of
protons.

Atomic number - number of protons in an
atom
• Elements are arranged by atomic number on the
periodic table.

Atoms are neutral, therefore # p = # e
Isotopes
 Number
of neutrons is NOT on the
periodic table for most elements….

Isotopes - atoms of a given element that differ
in the number of neutrons in the nucleus
• Mass number – sum of the protons and neutrons in
an atoms’ nucleus
• The periodic table shows the average of the mass
numbers for the isotopes of an element.
Describing Isotopes
Mass number
 Isotopes

12C

13C

14C
12
C
of carbon
carbon-12
carbon-13
carbon-14
__ neutrons
__ neutrons
__ neutrons
• All contain ____ protons and electrons.
Carbon on the
Periodic table
Isotopes and Radioactivity (RA)
 RA


isotope has an unstable nucleus
Nucleus emits energy and particles in an
effort to become more stable
May change the number of protons in the
nucleus and become a different element.
Radioactive Isotopes
 Possible
to target the energy and detect
the radioactivity.
 RA isotopes are used:


in research to track/follow molecules
in medicine to treat cancer and diagnose
disease
• Radiation therapy – treatment of localized cancer
• PET - diagnosis
Radioactive Isotopes
 Overexposure
to RA isotopes is
HARMFUL.

Energy emitted damages cells.
• radiation therapy takes advantage of this, goal is to
damage and kill cancer cells

Exposure to RA can also cause mutations that
lead to cancers
• Eg – exposure to RA element radon is the 2nd
leading cause of lung cancer
Diagnosis - PET Scans
A

radioactive tracer is put into the body.
Often RA glucose
 The
RA glucose goes to the parts of the
body that use glucose for energy.

Cancers use glucose differently from normal
tissue
 As
the radiotracer is broken down
positrons are made. This energy appears
as a 3-dimensional image on a computer
monitor.
Electron Arrangement (2.5)
 When
compounds form, the electrons of
the bonding atoms interact in attempt to
obtain a more stable state.
 Some
electron arrangements are more
stable than others…….see board
Chemical Bonding (2.6-2.7)
bonding – atoms gain,
lose, or share electron(s) to obtain a
stable number of electrons
 Chemical

Can be ionic bond or covalent bond
Chemical Bonding - Ionic
Bond – strong attractive force
between oppositely charged ions
 Ionic

Atoms form ions by losing or gaining enough
electron(s) to obtain a stable # of electrons in
their outer shell
electron transfer
SODIUM
ATOM
11 p+
11 e-
SODIUM
ION
11 p+
10 e-
CHLORINE
ATOM
17 p+
17 e-
CHLORINE
ION
17 p+
18 e-
Ionic
Bonding
Chemical Bonding - Covalent
Bond – bonded atoms share
pair(s) of electrons and form molecules.
 Occurs between nonmetals such as:
C, O, H, N, P, S
 Covalent bonding occurs in
 Covalent
• H2
• O2
• H2O
Two Classes of Covalent Bonds
Covalent Bond – bonded atoms
share electrons equally
 Nonpolar

Occurs between like atoms or between atoms
with a similar ability to attract shared electrons
Covalent Bond – unequal sharing of
electrons by the bonded atoms
 Polar

Occurs between atoms with very different
ability to attract shared electrons
Two hydrogen atoms,
each with one proton,
share two electrons
in a single nonpolar
covalent bond.
molecular hydrogen (H2)
H—H
Fig. 2-8b(1), p.25
water (H2O)
H—O—H
Two oxygen
atoms share
four electrons in
a nonpolar
double
covalent bond.
molecular oxygen (O2)
O=O
Types of Covalent Bonds
covalent – bonded atoms
share the electrons equally
 Nonpolar

Examples of nonpolar bonds:
• H2
 Atoms
with different electronegativity
values form polar covalent bonds.
• Electronegativity (EN) – measure of an
atom’s ability to attract shared electrons in
a covalent bond
• Oxygen and nitrogen have fairly large EN
values – often d • Carbon and hydrogen have low EN values
– often d +
Covalent – unequal pull on shared
electrons by the bonded atoms
 Results in partial charges on the bonded
atoms
d-
 Polar
O
H
d+
H
d+
Common Polar Covalent Bonds
O-H
N-H
C-O
C=O
Label the polarity in each bond.
Forces between Molecules
 Molecules
are weakly attracted to each
other by intermolecular (IM) forces,
 The most important IM force in biology is
the hydrogen “bond” (2.8)

Attractive force between d + H and d – O, N
or F
 Hydrogen
“bond” is a weak attractive
force between a d + hydrogen and a
d- O, N, or F in a second polar bond
Water is a polar molecule
capable of hydrogen bonding.
Properties of Water
1.
Water is cohesive and has high surface
tension.
• Cohesion – ability of molecules to
stick together
• Surface tension - ability to resist
rupturing when under tension
Properties of Water
2)
Water resists changes in
temperature.
• When heat is applied to an aqueous
solution much of the heat (energy) is used
to break hydrogen bonds, not to increase
the movement of the molecules.
Properties of Water
Solid water (ice) is less dense than
liquid water
3)
•
Ice floats
• Therefore, ice forms on the top of lakes
and insulates the liquid water below.
4)
Water is a good solvent for ionic
compounds and small polar
molecules.
Water H bonds to polar
molecules like ethanol
Water as solvent
 Water
pulls ions apart and hydrates them
Related Terms

Hydrophilic
• Water loving
• Capable of hydrogen bonding to water
(polar)

Hydrophobic
• Water “fearing”
• Cannot hydrogen bond to water
(nonpolar)
Acids, Base, and Buffers (2.14)
 Many
ions are dissolved in the fluids
in/outside of cells – called electrolytes


Na+, Ca+2, K+
H+
 Level

of each ion is critical
Our focus is on H+ (hydrogen ions)
Acids, Base, and Buffers
Substance that produces H+ when
dissolved in water……….
 Acid:

Examples:
• Hydrochloric acid – stomach acid
• Lactic acid – made when cells run out of oxygen
• Amino acids – building blocks of proteins
Acids, Base, and Buffers
substance that accepts H+1
(hydrogen ions) in water
 Base:

Examples:
• Sodium hydroxide - NaOH
• Most nitrogen containing compounds



Ammonia – NH3
Urea – in urine
Amino acids – building blocks for proteins
Acids, Base, and Buffers
 Classify
substances as acid, base or
neutral by their pH



Acids: pH < 7
Base: pH > 7
Neutral: pH = 7
• Pure water has a pH of 7
• See page 28
Acids, Base, and Buffers
 How



the pH scale works
The lower the pH the more acidic
The higher the pH the more basic (alkaline)
A difference of 1 pH unit is a 10-fold
difference in acidity or alkalinity
Why is pH important?
 Most
cells require a pH near 7.
 Above or below this pH for too long and
they die.

Proteins function only at specific pHs.
• In lab you will determine the optimal pH for a
protein that is needed to breakdown hydrogen
peroxide in cells
Acids, Base, and Buffers
 Buffers:
solution that resists changes in
pH even when acid or base is added


Buffers can both produce H+ and neutralize
H+
Buffers are key to maintaining pH
homeostasis
• Most body solutions are buffered
Why is pH important?
 Blood

has a pH of 7.3 – 7.4
If the pH is above or below this range for
more than a couple of days death occurs.
• The blood buffer system helps keep blood pH in a
range that supports life.