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Chapter 2:
The Chemical Context of Life
AP Biology
Overview: A Chemical Connection to Biology
•
Biology is a multidisciplinary science
•
Living organisms are subject to basic laws of physics and chemistry
•
•
Ex) The use of formic acid by ants to maintain “devil’s gardens,” stands of
Duroia trees
•
Some patches of forest are almost completely dominated by this
single plant species
•
Researchers showed that ants prevent other plant species from
growing in these” gardens” by injecting intruders with the poisonous
chemical formic acid
This chapter focuses on the chemical components and their interactions that
make up all matter
•
Elements and compounds
•
Atomic structure
•
Chemical bonding and the formation of
molecules
•
Chemical reactions that make and
break chemical bonds
Concept 2.1:
Matter consists of chemical elements
in pure form and in combinations
called compounds
•
Organisms are composed of matter
•
Matter is anything that takes up space and has mass
• Ex) Rocks, metals, oils, gases, humans
•
Matter is made up of elements
–
An element is a substance that cannot be broken down to other
substances by chemical reactions
• Scientists recognize 92 naturally-occurring elements
– Ex) Gold, copper, carbon, oxygen
• Each element has a symbol, usually the first letter or two of its
name
– Some symbols are derived from Latin or German
• Ex) Sodium is “Na” from the Latin word natrium
•
A compound is a substance consisting of two or more different elements in
a fixed ratio
–
A compound has characteristics different from those of its elements
Fig. 2-3
• Ex) Table salt (NaCl) – pure sodium is a metal; pure chlorine is a
poisonous gas
Sodium
Chlorine
Sodium
chloride
Essential Elements of Life
•
About 25 of the 92 elements are essential to life
–
•
Only 4 of these elements together make up 96% of living matter:
•
Carbon (18.5%)
•
Hydrogen (9.5%)
•
Oxygen (65%)
•
Nitrogen (3.3%)
Most of the remaining 4% consists of:
–
Calcium (1.5%)
–
Phosphorus (1.0%)
–
Potassium (0.4%)
–
Sulfur (0.3%)
–
Sodium (0.2%)
–
Chlorine (0.2%)
–
Magnesium (0.1%)
Essential Elements of Life
•
Trace elements are those required by an organism in minute quantities
–
Some trace elements are needed by all forms of life
•
–
Ex) Iron (Fe)
Other trace elements are required by only
certain species
•
Ex) Iodine (I) is an essential ingredient of
a hormone produced by the thyroid gland
in all vertebrates
–
Iodine deficiency causes a condition
called goiter, in which the thyroid
gland grows to an abnormal size
Concept Check 2.1
• 1) Explain how table salt has emergent properties.
• 2) Is a trace element an essential element?
Explain.
• 3) Iron (Fe) is a trace element required for the
proper functioning of hemoglobin, the molecule
that carries oxygen in RBCs. What might be the
effects of an iron deficiency?
Concept 2.2:
An element’s properties
depend on the structure of its atoms
•
Each element consists of unique atoms
–
An atom is the smallest unit of matter that still retains the properties of
an element
• Atoms are symbolized with the same abbreviation used for the
element made up of those atoms
– Ex) “C” represents both the element carbon and a single
carbon atom
•
Atoms are composed of subatomic particles
–
Relevant subatomic particles include:
• Neutrons (no electrical charge)
• Protons (positive charge)
• Electrons (negative charge)
•
Neutrons and protons are packed tightly together at the center of an atom to form the
atomic nucleus
•
•
Protons give the nucleus a positive charge
Electrons form a cloud around the nucleus
•
Attraction between opposite charges keeps these electrons in the vicinity of the
nucleus
Fig. 2-5
Cloud of negative
charge (2 electrons)
Electrons
Nucleus
(a)
(b)
• Neutron mass and proton mass are almost identical (~1.7 x 10 –24
grams)
•
Because the mass of these particles is so miniscule, they are
measured in daltons
• The dalton is the same unit of measure as the atomic mass
unit (amu)
•
Neutrons and protons both have masses close to 1 dalton
• Electron mass can be ignored when computing atomic mass
because it is only ~1/2000 that of a proton or neutron
Atomic Number
•
Atoms of the various elements differ in number of subatomic particles
–
All atoms of a particular element have a unique number of protons,
known as the atomic number, in their nucleus
• The atomic number is written as a subscript to the left of the symbol
for that element
– Ex) 2He = Helium (2 protons)
–
Unless otherwise indicated, an atom contains the same number of
protons and electrons, making it electrically neutral
Atomic Mass
•
An element’s mass number is the sum of protons plus neutrons in the nucleus
–
The mass number is written as a superscript to the left of the element’s symbol
–
Because the atomic number indicates how many protons are present, the
number of neutrons can be determined by subtracting the atomic number from
the mass number
•
•
Ex)
= Helium (2 protons; 4 - 2 = 2 neutrons)
Atomic mass, the atom’s total mass, can be approximated by the mass number
–
This is because the contribution of electrons to mass is negligible
•
Ex) Sodium (
) is 22.9898, or ~23 daltons
Isotopes
•
All atoms of an element have the same number of protons but may differ in number
of neutrons
–
Isotopes are two atoms of an element that differ in number of neutrons
•
Because neutrons contribute to atomic weight, isotopes have a different
mass than that of the element
– Despite their different masses, isotopes behave identically in chemical
reactions
–
In nature, an element occurs as a mixture of its isotopes
–
•
Ex) There are 3 isotopes of carbon: carbon-12 (99%) and carbon-13
and carbon-14 (1%)
The number usually given as the atomic mass of an element is actually an
average of the atomic masses of all the element’s naturally occurring
isotopes
Isotopes
•
Some isotopes, known as radioactive isotopes, are unstable and have a tendency
to lose particles
–
When this decay eventually leads to a change in the number of protons, the
atom is transformed to an atom of a different element
•
–
Ex) Radioactive carbon decays to form nitrogen
Radioactive isotopes have many useful applications
•
Measurements of radioactivity in fossils can be used to date them
Fig. 2-7
•
Radioactive isotopes are useful as tracers to
follow atoms through metabolism
–
Cells use radioactive atoms in the same
manner as they would nonradioactive
isotopes of the same element, but they can
be easily detected
•
Cancerous
throat
tissue
Ex) Biologists can detect cancer by
using radioactive tracers to monitor the synthesis of DNA and
looking for “hot spots” of activity
The Energy Levels of Electrons
•
When two atoms approach each other during a chemical reaction, only their
electrons are involved directly
•
An atom’s electrons vary in the amount of energy they posses
•
•
Energy is the capacity to cause change
The electrons of an atom have a type of energy called potential energy
because of how they are arranged in relation to the nucleus
•
•
Potential energy is the energy that matter has because of its location
or structure
The negatively-charged electrons are attracted to the positively-charged
nucleus
•
As a result, it takes work to move an electron farther away from the
nucleus
•
This means that more distant electrons have greater potential
energy
The Energy Levels of Electrons
•
Changes in the potential energy of electrons can occur only in steps of fixed
amounts, similar to a ball on a staircase
•
The ball can have different amounts of potential energy, depending on
which step it is on, but it cannot spend much time between steps
•
Similarly, an electron cannot exist in between energy levels
(a) A ball bouncing down a flight
of stairs provides an analogy
for energy levels of electrons
The Energy Levels of Electrons
•
An energy level, or electron shell, is an electron’s state of potential energy
•
An electron’s energy level or shell is correlated with its average distance from
the nucleus
•
In diagrams, shells can be represented by concentric circles
•
Fig. 2-8
The first shell is closest to the
anddown
contains
(a) Anucleus
ball bouncing
a flight electrons with the
of stairs provides an analogy
for energy levels of electrons
lowest potential energy
•
Electrons in the higher shells have increasingly higher amounts of potential
energy
Third shell (highest energy
level)
Second shell (higher
energy level)
First shell (lowest energy
level)
(b)
Atomic
nucleus
Energy
absorbed
Energy
lost
The Energy Levels of Electrons
Fig. 2-8
•
(a) A ball bouncing down a flight
of stairs provides an analogy
for energy levels of electrons
An electron can change the shell it occupies by absorbing or losing an amount of
energy equal to the difference in potential energy between its position in the old shell
and that in the new shell
Third shell (highest energy
level)
•
When an electron absorbs energy, it
moves to a shell farther out from the
nucleus
•
•
Second shell (higher
energy level)
First shell (lowest energy
level)
Ex) Light energy can excite an
(b)
electron to a higher energy level,
as seen in the first step of photosynthesis
Energy
absorbed
Energy
lost
Atomic
nucleus
When an electron loses energy, it “falls back” to a shell closer to the nucleus,
with the lost energy usually released to the environment as heat
•
Ex) The hot temperature of a car’s surface on a sunny day is the result of
sunlight exciting electrons at the car surface to higher energy levels, and
then those electrons falling back to their original levels
Electron Distribution and Chemical Properties
•
The chemical behavior of an atom is determined by the distribution of electrons in the
atom’s electron shells
•
The periodic table of the elements shows the electron distribution for each
element
•
The elements are arranged in rows, known as periods, that correspond to
the number of electron shells in their atoms
•
•
Ex) Nitrogen = 2nd row = 2 electron shells
The left-to-right sequence of elements in each row corresponds to the
sequential addition of electrons
and protons
•
Ex) Oxygen has +1
electrons and
protons than
nitrogen
•
Electrons (like all matter) tend to exist in the lowest available state of potential
energy
–
In an atom, the lowest such state is in the 1st shell
•
The 1st shell, however, can only hold 2 electrons
–
•
Atoms with more than 2 electrons must therefore use higher shells
because the 1st shell is full
The 2nd shell holds a max of 8 electrons
–
Atoms with more than 10 electrons (2 in the 1st, 8 in the 2nd) must
therefore begin
filling the next
energy level
•
The chemical behavior of an atom depends mostly on the number of electrons in its
outermost shell
–
Electrons in the outermost shell are called valence electrons
•
–
The outermost electron shell in called the valence shell
•
•
Ex) Lithium’s valence shell is its 2nd shell
Atoms with the same number of electrons in their valence shells exhibit similar
chemical behavior
–
•
Ex) Lithium has 1 valence electron
Ex) Fluorine and chlorine both have 7 valence electrons; both form compounds
when combined with the element
sodium
An atom with a completed valence shell
is chemically unreactive, or inert
–
Such atoms are found at the far
right of the periodic table
(He, Ne, Ar)
Electron Orbitals
•
•
The electron shells of atoms are often symbolized as 2-dimensional concentric-circle diagrams
–
Each concentric circle, however, represents only the average distance between an
electron in that shell and the nucleus
–
Thus, the exact location of an electron in an atom can never be known
We can instead describe the space in which an electron spends most of its time
–
An orbital is the threedimensional space where an
electron is found 90% of
the time
Electron Orbitals
•
Orbitals can be thought of as components of an electron shell
–
There are a specific number of orbitals of distinctive shapes and orientations in
each electron shell
•
The 1st electron shell has only 1 spherical s orbital, called 1s
•
The 2nd shell has 4 orbitals:
– One large spherical s
orbital, called 2s
–
•
Three dumbbellshaped p orbitals,
called 2p orbitals
The 3rd shell and other
higher electron shells also
have s and p orbitals,
along with orbitals of more
complex shapes
Electron Orbitals
•
•
No more than 2 electrons can occupy a single orbital
–
The 1st electron shell (1 orbital) can therefore only hold 2 electrons
–
The 2nd electron shell (4 orbitals) can hold up to 8 electrons
The reactivity of atoms arises from the presence of unpaired electrons in one or
more orbitals of their valence shells
–
Atoms interact in such a way
that completes their valence
shells, which involves finding
partners for their unpaired
electrons
Concept Check 2.2
•
1) A lithium atom has 3 protons and 4 neutrons. What is its atomic mass in
daltons?
•
2) A nitrogen atom has 7 protons, and the most common isotope of nitrogen
has 7 neutrons. A radioactive isotope of nitrogen has 8 neutrons. Write the
atomic number and mass number of this radioactive nitrogen.
•
3) How many electrons does fluorine have? How many electron shells?
Name the orbitals that are occupied. How many electrons are needed to fill
the valence shell?
•
4) If 2 or more elements are in the same row, what do they have in
common? If 2 or more elements are in the same column, what do they have
in common (see Figure 2.9, pp. 36)?
Concept 2.3: The formation and
function of molecules depend on
chemical bonding between atoms
• Atoms with incomplete valence shells can interact with other atoms in such a
way that completes their valence shell, either by:
• Sharing valence electrons
• Transferring valence electrons
• These interactions usually result in atoms staying close together, held by
attractions called chemical bonds
• The strongest kinds of chemical bonds are:
• Covalent bonds
• Ionic bonds
Covalent Bonds
•
A covalent bond is the sharing of a pair of valence electrons by two atoms
–
When 2 atoms approach one another, their valence orbitals overlap and allow
them to share electrons
–
In a covalent bond, the shared electrons count as part
of each atom’s valence shell
–
Two or more atoms held together by covalent bonds
form a molecule
•
Ex) H2: the 1s orbitals (both containing 1 electron)
overlap to complete the valence shell with 2 electrons
•
Molecules can be represented in different ways:
–
Lewis dot structure: uses element
symbols, with dots representing the
outermost valence electrons
•
Ex) H:H
–
Electron distribution diagram: shows the
locations of all electrons within their
appropriate orbitals, along with
overlapping valence orbitals containing
shared electrons
–
Space-filling model: 3-D model that
comes closest to representing the
actual shape of a molecule
•
Molecules can be represented in different ways:
–
Structural formula: uses element
symbols, with lines to represent the
sharing of pairs of electrons
–
–
Ex) H-H
•
One line represents a single bond,
which is the sharing of one pair
of valence electrons
•
Two lines represent a double bond,
which is the sharing of two pairs of
valence electrons
Molecular formula: indicates the type
and number of atoms present in a
molecule but not electron sharing
•
Ex) H2
• Atoms that can share valence electrons have a bonding
capacity that corresponds to the number of covalent bonds they
can form
– This bonding capacity is called an atom’s valence
– It usually equals the number of unpaired electrons required
to complete the atom’s outermost (valence) shell
• Ex) Oxygen, with 6 electrons in its outermost shell, has
a valence of 2
• Covalent bonds can form between atoms of the same element
or atoms of different elements
– Molecules consisting of the same types of atoms are
known as pure elements
• Ex) H2
– A compound is a combination of two or more different
elements
• Ex) H2O
•
The attraction of a particular kind of atom for the electrons of a covalent
bond is called its electronegativity
–
The more electronegative an atom, the more strongly it pulls shared
electrons toward itself
• In a covalent bond between 2 atoms of the same element, both
atoms have equal electronegativity and thus share electrons
equally
– This type of bond in known as a nonpolar covalent bond
• Ex) O2
• In compounds where one atom is bonded to a more electronegative
atom, the electrons are not shared equally
– This type of bond in known as a polar covalent bond
• Ex) H2O
•
Unequal sharing of electrons causes a partial positive or negative charge
for each atom or molecule
–
Electrons spend more time near the more electronegative atom,
leading to a partial negative charge (δ-)
–
As a result, the less electronegative atom has a partial positive charge
(δ+)
Ionic Bonds
•
In some cases, 2 atoms are so unequal in electronegativity that the more
electronegative atom strips a valence electron completely away from its
partner
–
•
Ex) the transfer of an electron from sodium (11Na) to chlorine (17Cl)
After the transfer of an electron, both atoms have charges
–
A charged atom (or molecule) is called an ion
•
•
Depending on its charge, an ion can be classified as:
–
A cation: a positively charged ion
–
An anion: a negatively charged ion
An ionic bond is an attraction between an anion and a cation
–
The transfer of an electron is not the formation of a bond itself; rather it
simply allows an ionic bond to form
–
Ions participating in an ionic bond with each other need not have
acquired their charge by electron transfer with each other
Animation: Ionic Bonds
•
Compounds formed by ionic bonds are called ionic compounds, or salts
–
Salts (ex: sodium chloride (table salt)) are often found in nature as crystals of
various sizes and shapes
–
Each salt crystal is an aggregate of many cations and anions bonded by their
electrical attraction and arranged in a 3-D lattice
–
Ionic compounds do not consist of molecules having a definite size and number
of atoms
•
The formula for an ionic compound (ex: NaCl) indicates only the ratio of
elements in a crystal of the salt
•
Not all salts have equal numbers of cations and anions
–
Ex) MgCl2 has 2 chloride ions for each magnesium ion; magnesium must lose 2
valence electrons to obtain a complete valence shell, making it a cation with a
net charge of 2+
•
–
One magnesium cation can therefore form ionic bonds with 2 chloride ions
The term ion also applies to entire molecules that are electrically charged
•
Ex) In the salt ammonium chloride (NH4Cl), an ion with an electrical charge
of 1+, the anion is a single chloride ion (Cl-) and the cation is ammonium
(NH4+)
Weak Chemical Bonds
•
In organisms, most of the strongest chemical bonds are covalent bonds
–
•
The bonds link atoms to form cell’s molecules
Weaker bonding within and between molecules is also essential to the cell
•
–
Ex) Ionic and hydrogen bonds
Most important large biological molecules are held in their functional form by
weak bonds
–
In addition, when 2 molecules in a cell make contact, they may adhere
temporarily by weak bonds
•
This reversibility of weak bonding is advantageous because 2 molecules
can come together, respond to one another in some way, and then
separate
Hydrogen Bonds
•
A hydrogen bond forms when a hydrogen atom covalently bonded to one
electronegative atom is also attracted to another electronegative atom
–
In living cells, the electronegative partners are usually oxygen or
nitrogen atoms
Van der Waals Interactions
•
Even molecules with nonpolar covalent bonds may have positively and negatively
charged regions
–
Electrons are not always evenly distributed in such molecules
•
At any instant, they may accumulate by chance in one part of the molecule
or another
–
This results in continually
changing regions of positive
and negative charge that allow
all atoms and molecules to
stick together
•
These are called van der
Waals interactions
Van der Waals Interactions
•
Van der Waals interactions are weak and occur only when atoms or molecules are
very close together
–
Despite their weakness, they have recently been
shown to be responsible for ability of gecko lizard
to walk up a wall
•
Each gecko has 100s of 1000s tiny hairs with
multiple projections at hair’s tip that increase
surface area
•
These hairs are so numerous that,
despite their individual weakness, together they
can support the gecko’s body weight
Molecular Shape and Function
•
Together, van der Waals interactions, hydrogen bonds, ionic bonds in water,
and other weak bonds may form not only between molecules but also
between different regions of a single large molecule, such as a protein
–
Though these bonds are individually weak, their cumulative effect is to
reinforce 3-D shape of a large molecule
–
The characteristic shape and size of a molecule is very important to its
function
•
A molecule’s shape is determined by the positions of its atoms’ valence orbitals
–
When an atom forms covalent bonds, the orbitals in the valence shell rearrange
•
Molecules made up of only 2 atoms (H2, O2) are always linear
•
Molecules with more than 2 atoms have more complicated shapes
because the s and p orbitals may hybridize
–
For atoms with valence electrons in both s and p orbitals, the single s and 3 p
orbitals hybridize to form 4 new hybrid orbitals shaped like identical teardrops
extending from the atomic nucleus
•
If we connect the
larger ends of the
teardrops with lines,
we have the outline of
a geometric shape
called a tetrahedron (similar to a pyramid)
•
3 models representing molecular shape are shown for water and methane
–
Water: 2 of the hybrid orbitals in the oxygen’s valence shell are with hydrogen
atoms
•
–
•
This results in a molecule shaped roughly like a V, with its 2 covalent
bonds spread apart at an angle of 104.5 degrees
Methane: has shape of a completed
tetrahedron because all 4 hybrid orbitals
of carbon are shared with hydrogen
atoms
•
Nucleus of carbon is at the center
•
4 covalent bonds radiate to
hydrogen nuclei at corners of
tetrahedron
-Larger molecules containing multiple carbon
atoms have a more complex overall shape
–
However, the tetrahedral shape of
carbon bonded to 4 other atoms is a
common theme in many other the
molecules that make up living matter
•
Molecular shape is important in biology because it determines how
biological molecules recognize and respond to one another with specificity
–
Only molecules with complementary shapes can form weak bonds with
one another
–
Molecules with similar shapes can
have similar biological effects
• Ex) We can see this specificity in
the effects of opiates (drugs
derived from opium)
•
Opiates relieve pain and alter mood by binding to specific receptor molecules on
surface of brain cells (Why would brain cells carry receptors for opiates, compounds
NOT made by our bodies?)
–
The discovery of endorphins answered this question in 1975
•
Endorphins are signaling molecules made by pituitary gland that bind to
receptors, relieving pain and
producing euphoria during times
of stress (intense exercise)
–
Opiates have shapes similar to
endorphins (see boxed portions) and
mimic them by binding to endorphin
receptors in brain
•
This is why opiates and endorphins
have similar effects
Concept Check 2.3
• 1) Why does the following structure fail to make sense
chemically?
H-C=C-H
• 2) Explain what holds together the atoms in a crystal of
magnesium chloride (MgCl2).
• 3) If you were a pharmaceutical researcher, why would you
want to learn the three-dimensional shapes of naturally
occurring signaling molecules?
Concept 2.4:
Chemical reactions make and break
chemical bonds
•
The making and breaking of chemical bonds, which leads to changes in the composition of
matter, are called chemical reactions
–
When we write a chemical reaction, we use an arrow to indicate the conversion of the
starting material (reactants) to the final molecules (products)
•
-Coefficients indicate the number of molecules involved
–
Ex) Reaction between hydrogen and oxygen that forms water
•
This reaction breaks covalent bonds of H2 and O2 and forms the new bonds of H2O
–
The coefficient 2 in front of H2 means that the reaction starts with 2 molecules
of hydrogen
•
Notice that all the atoms of the reactants (amounts) must be accounted for in the
products (matter is conserved in a chemical reaction) – Stoichiometry
–
Reactions cannot create or
destroy matter – they can
only rearrange it!
•
Photosynthesis is another important chemical reactions
–
The raw materials of photosynthesis are
carbon dioxide (taken from air) and water
(absorbed from soil)
•
Within plant cells, sunlight powers the
conversion of these ingredients to sugar
(glucose) and oxygen molecules (released
as by-product to surroundings)
–
Although photosynthesis is actually a sequence of many chemical reactions,
we still end up with the same number and kinds of atoms we had at the start
•
Matter has simply been rearranged, with an input of energy provided by
sunlight
6 CO2 + 6 H20 → C6H12O6 + 6 O2
•
All chemical reactions are reversible, with products of forward reaction becoming
reactants for reverse reaction
–
Ex) Hydrogen and nitrogen molecules can combine to form ammonia (NH3),
but ammonia can also decompose to regenerate hydrogen and nitrogen
•
2 opposite-headed arrows indicate that the reaction is reversible:
3H2 + N2
–
2 NH3
One factors affecting the rate of a reaction is the concentration of reactants
•
Greater concentration = more frequent collisions of reactant molecules with
one another = more opportunities to react and form products
–
As products accumulate, however, collisions resulting in reverse
reaction become more frequent
•
Eventually, forward and reverse reactions occur at same rate, so relative
concentrations of products and reactants stop changing , a condition
known as chemical equilibrium
•
Chemical equilibrium is dynamic
–
Reactions are still going on, but with no NET effect on concentrations of
reactants and products
–
Equilibrium does NOT mean that reactants and products are equal in
concentration, but that their concentrations have stabilized at a
particular ratio
• In some chemical reactions, equilibrium point may lie so far to the
right (toward products), that these reactions go essentially to
completion (virtually all reactants are converted to products)
Concept Check 2.4
• 1) Refer to the reaction between hydrogen and oxygen that forms
water, shown with ball-and-stick models on pp. 42. Draw the Lewis
dot structures representing this reaction.
• 2) Which types of chemical reactions occur fastest at equilibrium:
the formation of products from reactants or reactants from products?
• 3) Write an equation that uses the products of photosynthesis as
reactants and uses the reactants as products. Add energy as
another product. This new equation describes a process that occurs
in your cells. Describe this equation in words. How does this
equation relate to breathing?
You should now be able to:
1. Identify the four major elements
2. Distinguish between the following pairs of
terms: neutron and proton, atomic number
and mass number, atomic weight and
mass number
3. Distinguish between and discuss the
biological importance of the following:
nonpolar covalent bonds, polar covalent
bonds, ionic bonds, hydrogen bonds, and
van der Waals interactions
Copyright © 2008 Pearson Education, Inc., publishing as Benjamin Cummings