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Chapter 2—Chemical Context
of Life
Atoms, Elements, Compounds,
and Molecules
Hierarchy of Biological Order
Emergent Properties
Figure 2.2
I. Chemical Elements &
Compounds
• Element
– Cannot be broken down to other substances
by chemical reactions
– Examples: carbon (C), sodium (Na), oxygen (O)
• Compound
– Substance containing 2 or more elements
combined in a fixed ratio
– Examples: H2O, NaCl, C6H12O6 (emergent properties)
Which 4 are the most common elements in the human body?
Life requires ~25 chemical elements
• Four elements make up 96% of living
matter:
• carbon (C) • hydrogen (H)
• oxygen (O) • nitrogen (N)
• Four more elements make up most of
remaining 4%:
• phosphorus (P) • calcium (Ca)
• sulfur (S) • potassium (K)
• Trace elements (<0.01%)
II. Atoms & Molecules
• Atomic structure determines the behavior
of an element
– Atom
• Smallest unit of matter that retains the properties
of an element
– C (atom) vs. C (element)
Subatomic Particles
Particle
Charge
Location
Mass
(amu/dalton)
proton
+
nucleus
1
neutron
0
nucleus
1
electron
-
Cloud outside
nucleus
0
Atomic nucleus vs. cell nucleus?
(negligible)
Simplified Model of a Helium (He)
Atom
Figure 2.5
Atoms are mostly empty space—
(nucleus = golf ball, electron cloud = 1 km)
Atomic Number and Mass
• Atomic number
– # of protons in nucleus
of an atom
– Also = # of electrons
– Unique for a particular
atom
• Mass number
– The sum of protons +
neutrons in nucleus
Isotopes
• How are isotopes different than ‘regular’
atoms?
– Isotope
• An atom with more neutrons than usual
(larger mass)
• Behaves the same in chemical reactions
– Examples: carbon-13, carbon-14 (99% = carbon-12)
– Why is the atomic mass of carbon 12.011, not 12?
Use of Radioactive (unstable) Isotopes
Substances are ‘labeled’
with isotopes in order to:
- follow metabolic
processes
- find their locations within
cells
- to use as diagnostic
tools in medicine
Figure 2.6
Electron Energy Levels
Figure 2.9
• Electrons have potential energy due to position in relation to nucleus
• Electrons exist only at fixed levels of potential energy (electron shells)
Electron Energy Levels
Ball on a
staircase…
Electron energy levels (shells) have different states of potential energy
(Higher levels have more energy)
Electron Configurations & Chemical
Properties
•Atoms with completed
valence shell are
unreactive (noble
gases) (Ne & Ar)
Figure 2.10
•Chemical behavior/bonding of an atom depends on # of electrons in
its outermost shell (valence shell)
•Atoms with same # of valence electrons behave similar chemically
(F & Cl) (O & S)
Chemical Reactivity
• Atoms tend to complete a partially filled
valence shell
or
• empty a partially filled valence shell
– This tendency drives chemical
reactions…and creates bonds
Electron Orbitals
 Orbital = 3-dimensional space where an electron is found 90% of the time
 Rule—no more than 2 electrons per orbital
 One electron shell/level may contain multiple orbitals
Strangers getting
on a bus….
Figure 2.11
Atoms combine by chemical
bonding
• Chemical Bonds
– Attraction between 2 atoms due to:
• sharing of outer shell electrons (covalent bonds)
– or
• The presence of opposite charges on the atoms
(ionic bonds)
– Bonded atoms gain complete outer electron
shells
Covalent Bonding forms MOLECULES
•Covalent Bond = 2
atoms sharing a pair
of valence electrons
•Valence = bonding
capacity of an atom
(# of unpaired e-)
(Single, double, &
triple bonds
possible)
Figure 2.12
Nonpolar Covalent Bonds
• Electronegativity
– Attraction of an atom for the electrons in a
covalent bond
• The more electronegative, the more strongly it
pulls
• Nonpolar covalent bond
– electrons are shared equally between atoms
(equal tug of war)
– Examples: O2, H2, CH4
Type of Bonding?
Water Molecule
(polar covalent bonds)
Polar covalent
bond = electrons
are not shared
equally between
atoms
(e- spend more
time closer to the
more
electronegative
atom)
i.e. H2O
Figure 2.13
Type of Bonding?
Ionic Bonding
Figure 2.14
 Transfer of electron from one atom to another causes ions to form
cation—ion with positive charge
anion—ion with negative charge
 Opposite charges attract = ionic bond
Ex. NaCl
Ionic Compounds (salts)
Why is an ionic
compound not
called a
molecule?
(no definite size
or number of
atoms, only a
ratio of
elements)
Example: MgCl2
Figure 2.15
Type of Bonding?
Hydrogen Bonding
(weak chemical bond)
 A hydrogen atom (+)
from one molecule is
attracted to an
electronegative atom (-)
in another molecule
 Attraction = hydrogen
bond
Figure 2.16
Van der Waals interactions
 weak attractions between molecules or parts
of molecules due to localized charge
fluctuations
 Due to random chance
 Molecules/atoms must be very close together
The function of a molecule is related to
its shape
Figure 2.17
Specific molecular
shapes allow for
molecule to molecule &
cell to cell
communication
(lock & key)
Molecular Shape & Brain Chemistry
Molecular Shape & Brain Chemistry
Figure 2.18
Molecular Mimics
Figure 2.19
Chemical Reactions—making and
breaking chemical bonds
Law of Conservation of Mass—same # of each atom on both sides
• Most reactions are reversible
• Example: 3H2 + N2 ↔ 2NH3
• Chemical Equilibrium
– the point at which the rate of the forward
reaction equals the rate of the reverse
reaction
– Concentrations of products/reactants stop
changing