Download Ch. 2: Biochemistry

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Chapter 2
The Chemical Context of Life
+ Overview: A Chemical
Connection to Biology
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Living organisms follow basic laws of physics and
chemistry
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use of formic acid by ants to
maintain “devil’s gardens,”
stands of Duroia trees
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
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Concept 2.1: Matter consists of chemical elements
in pure form and in combinations called
compounds
Organisms composed
of matter
Matter: anything that
takes up space and has
mass
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Made up of elements
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Element: substance that cannot be
broken down to other substances by
chemical reactions
Compound: substance consisting of
two or more elements in a fixed ratio
Compound may not have same
characteristics as elements it’s made of
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
Essential
Elements
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of Life
 Only
25 of 92 elements
essential to life
 96% of living matter:
 Carbon, hydrogen,
oxygen, and nitrogen
 Remaining
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4%:
Calcium, phosphorus,
potassium, and sulfur
 Trace
elements: those
required by an organism
in minute quantities
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
Fig. 2-4
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(a) Nitrogen deficiency
(b) Iodine deficiency
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Concept 2.2: An element’s properties
depend on the structure of its atoms
Each element consists of unique
atoms
Atom: smallest unit of matter that
still retains the properties of an
element
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Neutron mass & proton mass are
almost identical
 measured in daltons
Subatomic Particles
 Neutrons (no electrical charge)
 Location: Atomic Nucleus
 Protons (positive charge)
 Location: Atomic Nucleus
 Electrons (negative charge)
 Location: In cloud around nucleus
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
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Atomic Number and Atomic Mass
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Atomic number: number
of protons in nucleus
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Mass number: sum of
protons plus neutrons in
nucleus
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Atomic mass: atom’s total
mass AKA mass number
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Same element = same number of
protons but may differ in number of
neutrons
Isotopes: two atoms of an element that differ
in number of neutrons
Radioactive isotopes: decay spontaneously,
giving off particles and energy
 GREAT for labeling proteins and cells
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Applications:
 Dating fossils
 Tracing atoms
through metabolic
processes
 Diagnosing
medical disorders
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
+ The Energy Levels of Electrons
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Energy: capacity to cause change
Potential energy: that matter has because of its location or
structure
Electrons of atom
differ in amounts
of potential
energy
Electron’s state
of potential
energy is called
its energy level,
or electron shell
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
+Electron Distribution and Chemical Properties
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Valence electrons: in the outermost shell, or valence shell
Elements with full valence shell are chemically inert
Chemical behavior of atom determined by distribution of electrons in
electron shells, MOSTLY by valence electrons
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Periodic table of the elements shows electron distribution for each element
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
+ Electron Orbitals
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An orbital is the three-dimensional space where an electron
is found 90% of
the time
Each electron
shell consists of
a specific
number of
orbitals
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
+ Concept 2.3: The formation and function of
molecules depend on chemical bonding
between atoms
• Incomplete valence
shells can share/transfer
valence electrons with
other atoms
• usually result in
chemical bonds
• atoms staying
close together,
held by attraction
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
 Covalent
bond:
sharing of a pair of
valence electrons by
two atoms
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Molecule: consists of two or more
atoms held together by covalent
bonds
A single covalent bond, or single
bond, is the sharing of one pair of
valence electrons
A double covalent bond, or double
bond, is the sharing of two pairs of
valence electrons
Structural formula: notation used
to represent atoms and bonding
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Ex: H–H
Molecular formula: indicates the
amount and type of atoms in a
molecule
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Ex: H2
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
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Covalent bonds can form between same OR different elements
Compound: combination of two or more different elements
Valence: Bonding capacity of atom
Electronegativity: atom’s attraction for electrons in covalent bond
More electronegative an atom = the more strongly it pulls electrons
Nonpolar covalent bond: electrons equally shared
Polar covalent bond: one atom is more electronegative  electrons
unequally shared
Unequal sharing  partial positive or negative charge
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
Fig. 2-14-2
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Na
Cl
Na
Cl
Na
Sodium atom
Cl
Chlorine atom
Na+
Sodium ion
(a cation)
Cl–
Chloride ion
(an anion)
Sodium chloride (NaCl)
+ Ionic Bonds
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Ion: charged atom or molecule
After the transfer of an electron  both atoms have charges
Cation: positively charged ion
Anion: negatively charged ion
Ionic bond: attraction between anion and cation
Ionic compounds, or salts: compounds formed by ionic
bonds
Salts like NaCl
often found in
nature as crystals
Animation: Ionic Bonds
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
+ Weak Chemical Bonds
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Most of the strongest bonds in organisms are covalent bonds
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form a cell’s molecules
Weak chemical bonds:
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reinforce shapes of large molecules
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help molecules adhere to each other
Hydrogen bond: forms when a
hydrogen atom covalently
bonded to one electronegative
atom is also attracted to another
electronegative atom
In living cells electronegative
partners usually oxygen or
nitrogen atom
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
+ Van der Waals Interactions
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Electrons are distributed
asymmetrically“hot spots” of positive or
negative charge
Van der Waals interactions: attractions
between molecules that are close together
as a result of these charges
Together can be strong, as
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Ex: between molecules of a gecko’s toe hairs
and a wall surface
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
+ Molecular Shape and Function
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Shape = Function
Shape determined by positions of atoms’
valence orbitals
In covalent bond s and p orbitals may
hybridize  specific molecular shapes
Biological molecules recognize and
interact based on molecular shape
Similar shapes = similar biological
effects
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
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Concept 2.4: Chemical reactions make and
break chemical bonds
Chemical reactions: making and breaking of chemical bonds
Reactants: starting molecules of chemical reaction
Products: final molecules of a chemical reaction
Photosynthesis IMPORTANT chemical reaction
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Sunlight powers the conversion of carbon dioxide + water  glucose + oxygen
6 CO2 + 6 H20 → C6H12O6 + 6 O2
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All chemical reactions reversible
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products of the forward reaction
become reactants for the reverse
reaction
Chemical equilibrium reached
when the forward and reverse
reaction rates are equal
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
+ You should now be able
1.
Identify
to: the four major elements
Distinguish between the following pairs of terms:
2.
1.
2.
3.
neutron and proton,
atomic number and mass number,
atomic weight and mass number
Distinguish between and discuss the biological importance of
the following:
3.
4.
nonpolar covalent bonds,
polar covalent bonds,
ionic bonds,
hydrogen bonds, and
5.
van der Waals interactions
1.
2.
3.
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings
Fig. 2-UN3
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Nucleus
Protons (+ charge)
determine element
Neutrons (no charge)
determine isotope
Electrons (– charge)
form negative cloud
and determine
chemical behavior
Atom
Fig. 2-UN5
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Single
covalent bond
Double
covalent bond
Fig. 2-UN9
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Fig. 2-UN10
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Fig. 2-UN11
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