Transcript Chapter 2

Chapter 2 – Chemical Context of Life
• Biology is a multi-disciplined science
– In order to understand biology, an
understanding of basic chemistry is needed
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Some Basic Definitions
• Matter – anything that takes up space and has
mass
• Element – a substance that can’t be broken down
to other substances by chemical reaction
– 92 found in nature
• Compound – substance containing two or more
different elements combined in a fixed ratio
– Has characteristics different from those of its
elements
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Compound
• An example of a compound is table salt (NaCl)
– Sodium is a metal
– Chlorine (chloride) is a poisonous gas (liquid)
+
Sodium
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Chloride
Sodium Chloride
Elements in Biology
•Only 25 of the 92 natural
elements are needed for life
•Of those 25, 4 make up 96%
of all living matter
– O, C, H, N
•Other 4% are Trace Elements
– elements required by an
organism in small
quantities
Table 2.1
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Atoms
•Atom – smallest unit of matter
that still retains the properties
of the element
– Three sub-atomic
particles are of interest to
us
• Protons
• Neutrons
• Electrons
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Sub-Atomic Particles
•Protons
– Positively charged
– Located in the nucleus
– Mass of about 1 dalton
•Each atom of an element has a
unique number of protons in its
nucleus
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Sub-Atomic Particles
•Neutrons
– No charge
– Located in nucleus
– Mass of about 1 dalton
– Number may vary
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Sub-Atomic Particles
•Electrons
– Negative charge
– Located in cloud around the nucleus
– Mass is negligible
– Important to bonding
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• Simplified models of an atom
Cloud of negative
charge (2 electrons)
Electrons
Nucleus
(a) This model represents the
Figure 2.4
electrons as a cloud of
negative charge, as if we had
taken many snapshots of the 2
electrons over time, with each
dot representing an electron‘s
position at one point in time.
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(b) In this even more simplified
model, the electrons are
shown as two small blue
spheres on a circle around the
nucleus.
Atomic Number and Atomic Mass
•Atomic Number – number of protons in
the nucleus
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Atomic Number and Atomic Mass
•Atomic mass – sum of protons and neutrons in
the nucleus of an atom
– Atomic mass – atomic number = # of
neutrons
For our purposes, Atomic mass is the same as Mass number
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• Assume atoms are neutral in charge
– The Atomic number (# of protons or positive
charges) = the number of electrons (or
negative charges).
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Isotopes
• Isotopes – atoms of the same element that
differ in the number of neutrons
• Radioactive isotopes – one in which the
nucleus decays spontaneously, giving off
particles and energy
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Chemical Reactions
• Only electrons are involved in chemical
reactions
• Nuclei never get close enough to one another
to react
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Energy
• Energy – the capacity to cause change (by
doing work)
• Potential energy – the energy that matter
possesses because of location or structure
In order to restore potential energy (slinky to the
top of the stairs), work must be done (slinky
must be brought to the top of the stairs)
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Energy
•The more distant an electron is
from the nucleus, the more potential
energy it has
•An electron’s energy level is related
to its average distance from the
nucleus
– Energy level – the different
states of potential energy that
an electron has in an atom
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Represented
symbolically
by electron
shells
Chemical Behavior of Atoms
• Chemical behavior of atoms is determined by
the distribution of electrons in the atom’s
electron shells
• Electron shells are filled in a specific order:
– First shell holds maximum of 2 electrons
– Second shell holds maximum of 8
– Third shell holds maximum of 8
• Each lower level must be filled before an
electron can be in a higher shell
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• Energy levels
– Are represented by electron shells
Third energy level (shell)
Second energy level (shell)
Energy
absorbed
First energy level (shell)
Energy
lost
Atomic
nucleus
Figure 2.7B
(b) An electron can move from one level to another only if the energy
it gains or loses is exactly equal to the difference in energy between
the two levels. Arrows indicate some of the step-wise changes in
potential energy that are possible.
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Chemical Behavior of Atoms
• Chemical behavior depends mostly on the
number of electrons in the outermost shell,
known as the valence shell and the electrons
in the valence shell are know as valence
electrons
• An atom with a completed valence shell is nonreactive
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Chemical Behavior of Atoms
• Reactivity of atoms arise from unpaired
electrons in the valence shell
• It is these unpaired electrons that interact to
form bonds
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Bonds
• Chemical bonds - attraction between atoms
resulting in either sharing or transferring of
valance shell electrons
• There are several types of bonds
– Covalent
– Ionic
– Hydrogen
– van der Waals interactions
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Covalent Bonds
• Covalent bond – sharing of a pair
of valence electrons by two atoms
• By sharing, each atom has a
completed valence shell for part
of the time
Hydrogen atoms (2 H)
+
+
+
+
• This is the strongest type of
bond
+
+
Hydrogen
molecule (H2)
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Covalent Bond
• A molecule
– Consists of two or more atoms held together
by covalent bonds
• A single bond
– Is the sharing of one pair of valence electrons
• A double bond
– Is the sharing of two pairs of valence electrons
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• Single and double covalent bonds
Name
(molecular
formula)
Electronshell
diagram
(a) Hydrogen (H2).
Two hydrogen
atoms can form a
single bond.
(b) Oxygen (O2).
Two oxygen atoms
share two pairs of
electrons to form
a double bond.
Figure 2.11 A, B
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Structural
formula
H
H
O
O
Spacefilling
model
Covalent Bond
• Electronegativity
– Is the attraction of a particular kind of atom for
the electrons in a covalent bond
• The more electronegative an atom
– The more strongly it pulls shared electrons
toward itself
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Covalent Bond Types
• Non-polar covalent bond
– Equal sharing of electrons is
– Atoms have similar electronegativities
– Ex: H-H
O=O
• Polar covalent bond
– Unequal sharing
– Resulting molecule has areas of relative
charge
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Polar Covalent Bond
• In a polar covalent bond
– The atoms have differing electronegativities
– Share the electrons unequally
Because oxygen (O) is more electronegative than hydrogen (H),
shared electrons are pulled more toward oxygen.
d–
This results in a
partial negative
charge on the
oxygen and a
partial positive
charge on
the hydrogens.
O
Figure 2.12
d+
H
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H
H2O
d+
Polar Covalent Bond
• Oxygen is one of the most electronegative
elements, attracting shared electrons more
strongly than hydrogen does
• Result is an unequal sharing
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Ionic Bonds
• In some cases, two atoms are so unequal in
their attraction for valence shell electrons, the
more electronegative atom strips the electron
away from its bonding partner
• This results in two ions or charged particles
– Cation – positively charged particle (has less
electrons)
– Anion – negatively charged particle (has more
electrons)
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Ionic Bonds
• An ionic bond
– Is an attraction between anions and cations
The lone valence electron of a sodium
atom is transferred to join the 7 valence
electrons of a chlorine atom.
1
2 Each resulting ion has a completed
valence shell. An ionic bond can form
between the oppositely charged ions.
+
Na
Na
Figure 2.13
Sodium atom
(an uncharged
atom)
Cl
Cl
Chlorine atom
(an uncharged
atom)
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Na
Na+
Sodium on
(a cation)
–
Cl
Cl–
Chloride ion
(an anion)
Sodium chloride (NaCl)
Ionic Bonds
• The transfer of an electron is not the
formation of a bond, but it allows an ionic
bond to form because the resulting ions are
attracted to each other
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Ionic Bonds
• Ionic compounds
– Are often called salts, which may form crystals
Na+
Cl–
Figure 2.14
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Ionic Bonds
• Ionic bonds are not as strong as covalent
bonds
• One reason is that the bond can be affected by
the envionment
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Weak Chemical Bonds
• Several types of weak chemical bonds are
important in living systems
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Hydrogen Bonds
• A hydrogen bond
– Forms when a hydrogen atom covalently
bonded to one electronegative atom is also
attracted to another electronegative atom
d–
d+
H
Water
(H2O)
O
H
d+
d–
Ammonia
(NH3)
N
H
d+
Figure 2.15
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H
H
d+
d+
A hydrogen
bond results
from the
attraction
between the
partial positive
charge on the
hydrogen atom
of water and
the partial
negative charge
on the nitrogen
atom of
ammonia.
Van der Waals Interactions
• Van der Waals interactions
– Occur when transiently positive and negative
regions of molecules attract each other
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Weak Bonds
• Weak chemical bonds
– Reinforce the shapes of large molecules
– Help molecules adhere to each other
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Molecular Shape and Function
• The precise shape of a molecule
– Is usually very important to its function in the
living cell
– Is determined by the positions of electrons in
the molecule’s atoms
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• Molecular shape
– Determines how biological molecules
recognize and respond to one another with
specificity
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Carbon
Nitrogen
Hydrogen
Sulfur
Oxygen
Natural
endorphin
Morphine
(a) Structures of endorphin and morphine. The boxed portion of the endorphin molecule (left) binds to
receptor molecules on target cells in the brain. The boxed portion of the morphine molecule is a close match.
Natural
endorphin
Brain cell
Figure 2.17
Morphine
Endorphin
receptors
(b) Binding to endorphin receptors. Endorphin receptors on the surface of a brain cell
recognize and can bind to both endorphin and morphine.
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Chemical Reactions
• A Chemical reaction
– Is the making and breaking of chemical bonds
– Leads to changes in the composition of matter
– Can not create or destroy matter
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•Chemical reactions
– Convert reactants to products
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Chemical Reactions
• Some reactions continue to completion
– That is, all the reactants are used up to make
products
• Most reactions are reversible
– That is, the products of the forward reaction
become the reactants of the backward reaction
– 3H2 + N2
2NH3
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Chemical Reactions
• Chemical equilibrium
– Is reached when the forward and reverse
reaction rates are equal
– The reactions are still occurring, but there is no
net effect on the concentrations
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