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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
(negligible)
Atomic nucleus vs. cell nucleus?
Simplified Model of a Helium (He)
Atom
Atoms are mostly empty space—
(nucleus = golf ball, electron cloud = 1 km)
Figure 2.5
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
• Electrons have potential energy due to position in relation to nucleus
• Electrons exist only at fixed levels of potential energy (electron shells)
Figure 2.9
Electron Energy Levels
Electron energy levels (shells) have different states of potential energy
(Higher levels have more energy)
Ball on a
staircase…
Electron Configurations & Chemical
Properties
•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)
•Atoms with completed
valence shell are
unreactive (noble
gases) (Ne & Ar)
Figure 2.10
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
Figure 2.11
Strangers getting
on a bus….
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
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
Figure 2.14
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
Specific molecular
shapes allow for
molecule to molecule &
cell to cell
communication
(lock & key)
Figure 2.17
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