The Chemical Context of Life
Earth supports an enormous variety of
organism. The structure & function of
all living things are governed by the
laws of CHEMISTRY
Overview: A Chemical Connection to Biology
• Biology is a multidisciplinary science
• Living organisms are subject to basic laws of
physics and chemistry
Bombardier Beetle
Elements and Atoms
• Everything in the universe is made of
matter
• Matter is anything that occupies space
& has mass
Mass -vs.- Weight
• The term mass and weight are often used but
have slightly different meanings.
• Mass refers to the amount of a substance
• Weight refers to the force gravity exerts on a
substance.
Ex: an object has the same mass on earth & moon,
but its weight will be different b/c of the
gravitational force on earth is greater than on the
moon
Calculate your mass on earth
• Find your mass by taking your weight and
divide by 2.2 kg
•
•
•
•
•
Ex: 300 lbs / 2.2 kg = 136.4 kg
W=mxa
W = 136.4 kg x 9.8 m/s2
W=
NOW- FIND YOUR MASS ON EARTH!
ON THE MOON
• Use the same procedures and calculate
your mass on the moon if kg = 1.6 m/s2
Elements and Atoms
• Elements are substances that cannot be
broken down chemically into simpler
kinds of matter
– More than 100 different elements have
been identified
– Fewer than 30 are important to living
things
– 90% of the mass of all living things comes
from 4 elements (CHON)
– Carbon, Hydrogen, Oxygen & Nitrogen
Kinds of Atoms
• 92 naturally occurring elements-each with a
different number of protons and a different
arrangement of electrons are found naturally
on Earth
• 11 are found in organisms in more than trace
amounts (.01% or higher)
• 4 of them make up the largest part of
our body weight (96.5%)
Families or Groups
• Inert elements-full outer shell (noble gases)
• Halogens-7 valence electrons-tend to react by
gaining one electron
• Alkali metals tend to give up their one
valence electron; very reactive
• Alkaline Earth metals give up two valence
electrons
Ions
• Atoms that have lost or gained
electrons, thus no longer neutral
• Cations-atom that has more protons
than electrons-will have a net positive
charge (Na +1)
• Anion-atom that has fewer protons than
electrons-will have a net negative
charge (Cl –1)
Structure of the Atom
• In 1913 a Danish physicist Niels Bohr
proposed every atom possesses an
orbiting cloud of tiny subatomic
particles called electrons which rotates
around the nucleus of the atom
Elements and Atoms
• Each element consists of a certain type
of atom, different from the atoms of
any other element
• Atom is the simplest particle of an
element that retains all the properties
of that element
• Atoms are composed of protons (+),
Neutrons (0), and Electrons (-)
Structure of the Atom
• Nucleus: makes up the bulk of the mass of
the atom & consist of 2 subatomic particles
1. Protons: positive charge; also called
the atomic number
2. Neutrons: neutral charge
• Mass Number: is equal to the total number of
P(+), & N(0)
– An atoms mass & subatomic particles are
measured in the units called Daltons.
– (6.02 x 1023) Daltons = 1 gram
Structure of the Atom
• Electrons: (-) are negatively charged;
are equal to the number of protons.
– Are high energy particles that have very
little mass
– They move around the nucleus at very
high speeds and are located in orbital
Isotopes
• Atoms of the same element have the
same number of protons but have
different numbers of neutrons
• Additional neutrons change the mass of
the element
• Most elements are made of a mixture of
isotopes
Radioactive Isotopes
• Radioactive isotopes decay spontaneously,
giving off particles and energy
• Some applications of radioactive isotopes
in biological research are
– Dating fossils
– Tracing atoms through metabolic processes
– Diagnosing medical disorders
Radioactive Isotopes
• Radioactive isotopes decay-nucleus
breaks up over time
• Exposure to radioactive elements is
very controlled. They can cause
mutations or even cell death.
• Half-life: time it takes for one half of
the atoms in a sample to decay-C-14
has a half life of about 5600 years
Cancerous
throat
tissue
A PET Scan
uses radioactive
Isotopes to detect
location of
intense chemical
activity in the
body
The yellow area
indicates cancerous
tissue
Carbon-12, Carbon-13, & Carbon 14
Carbon 12: 6 protons, 6 neutrons
Carbon 13: 6 protons, 7 neutrons
Carbon 14: rarest carbon isotope, unstable. Its nucleus breaks
up into elements with lower atomic numbers called nuclear
break up which emits radioactive decay.
Example of Half-Life
• A sample of carbon containing 1 gram
of carbon-14 today would contain about
0.5 grams of carbon 14 after 5,600
years
• 0.25g 11,200 years from now
• 0.125g 16,800 years from now
Scientists who work with radioactivity
• X-Ray technologists wear radiationsensitive badges to monitor amounts of
radioactivity exposures
• Badges are collected each month for
study
The Energy Levels of Electrons
• Energy is the capacity to cause change
• Potential energy is the energy that matter
has because of its location or structure
– The electrons of an atom differ in their amounts of
potential energy
– Changes in potential energy occur in steps of fixed
amounts
• An electron’s state of potential energy is
called its energy level, or electron shell
(a) A ball bouncing down a flight
of stairs provides an analogy
for energy levels of electrons.
Third shell (highest
energy level in this
model)
Second shell (higher
energy level)
First shell (lowest
energy level)
(b)
Atomic
nucleus
Energy
absorbed
Energy
lost
Distribution of Electrons within Energy
Levels
•
•
•
•
•
1st energy level can hold up to 2 electrons
2nd energy level can hold up to 8 electrons
3rd energy level can hold up to 18 electrons
4th energy level can hold up to 32 electrons
***No more than 8 electrons can be in the
last energy level of an atom: Octet Rule
• Electrons are found in different electron shells,
each with a characteristic average distance from
the nucleus
– The energy level of each shell increases with
distance from the nucleus
– Electrons can move to higher or lower shells by
absorbing or releasing energy, respectively
• The chemical behavior of an atom is determined
by the distribution of electrons in electron shells
• The periodic table of the elements shows the
electron distribution for each element
Figure 2.6
2
Hydrogen
1H
Atomic number
He
Atomic mass
First
shell
4.00
Helium
2He
Element symbol
Electron
distribution
diagram
Lithium
3Li
Beryllium
4Be
Boron
5B
Carbon
6C
Nitrogen
7N
Oxygen
8O
Fluorine
9F
Neon
10Ne
Sodium
11Na
Magnesium
12Mg
Aluminum
13Al
Silicon
14Si
Phosphorus
15P
Sulfur
16S
Chlorine
17Cl
Argon
18Ar
Second
shell
Third
shell
• Chemical behavior of an atom depends
mostly on the number of electrons in its
outermost shell, or valence shell
• Valence electrons are those that occupy
the valence shell
• The reactivity of an atom arises from the
presence of one or more unpaired electrons
in the valence shell
• Atoms with completed valence shells are
unreactive, or inert
Electrons Determine the Chemical
Behavior of Atoms
• Orbital is a 3-dimension region around the
nucleus that indicates the location of the
electron
• Electrons in orbitals farther away from the
nucleus have greater energy than electrons
closer to the nucleus
• When all orbitals are combined, there is a
cloud of electrons surrounding the nucleus
– Energy levels Sublevels  Orbitals
Orbitals correspond to specific
energy levels
• Each orbital can hold only a certain
number of e• Ex: the 1st orbital can hold only 2 e• There are 4 orbitals in the second elevel & that can hold up to 8 total e- w/
a max of 2 e- per orbital
• Sublevels: S, P, D, & F
• Spherical orbitals-near the nucleus
• Dumbbell shaped orbitals-near the nucleus (p
orbitals)
f2-4_electron_orbitals_cb.jpg
Compounds
Compounds are made up of atoms of 2
or more elements <H2O
Covalent Bonds
• form when 2 atoms share one or more
pairs of electrons H2
• In a covalent bond, the shared
electrons count as part of each atom’s
valence shell
• Two or more atoms held together by
valence bonds constitute a molecule
Covalent Bonds Build Stable Molecules
• When two or more atoms share one or
more pairs of valence electrons
1. It has no net charge
2. The octet rule is satisfied for each atom
3. It has no free electrons
• Hydrogen gas is a diatomic molecule meaning
that it is composed of 2 hydrogen atoms
Strength of a covalent bond
• The strength depends on the number of
shared electrons
• Can be single, double, or triple bond (share
one pair of electrons, 2 pair or 3 pair)
– The notation used to represent atoms and
bonding is called a structural formula
• For example, H—H
• For example, H2
FORMULAS
• Structural:
• H-H (line represents on pair of shared
electrons)
• O=O (two pair of electrons being shared)
• N---N (each line between two bonded atoms
represent the sharing of one electron) the
strongest
– Molecular formula: H2 and O2
Name and
Molecular
Formula
(a) Hydrogen (H2)
Electron
Distribution
Diagram
Structural
Formula
SpaceFilling
Model
Name and
Molecular
Formula
(b) Oxygen (O2)
Electron
Distribution
Diagram
Structural
Formula
SpaceFilling
Model
Diatomic Molecules
• Molecules of 2 atoms that are chemically
bonded together
• Can be the same element or 2 different
elements
• Common diatomic elements include:
• H2, N2, F2, O2, I2, Cl2, Br2
Electronegativity
• Atoms in a molecule attract electrons to
varying degrees
• Electronegativity is an atom’s
attraction 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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Nonpolar vs Polar covalent bonds
• Nonpolar covalent bond, the atoms share
the electron equally
• Polar covalent bond, one atom is more
electronegative, and the atoms do not share
the electron equally
• Unequal sharing of electrons causes a partial
positive or negative charge for each atom or
molecule
© 2014 Pearson Education, Inc.
Ionic Bonds
• Ionic Bonds: when an anion and cation
attract each other
NaCl (gain/lose electrons)
• A cation is a positively charged ion
• An anion is a negatively charged ion
Ionic Bonds
• Transfer of electron(s) from one atom
to another
• Results in ions (charged atoms)
• Valence electron is unpaired (free) &
can bond with another electron
– Example Table Salt NaCl; where Na +1,
and Cl –1
– Opposite charges form an ionic compound
Ionic Compounds
• Compounds formed by ionic bonds are
called ionic compounds, or salts
– Salts, such as sodium chloride (table salt),
are often found in nature as crystals
Na
Cl−
Weak Chemical Bonds
• Most of the strongest bonds in
organisms are covalent bonds that form
a cell’s molecules
• Weak chemical bonds, such as ionic
bonds and hydrogen bonds, are also
important
• Many large biological molecules are held
in their functional form by weak bonds
© 2014 Pearson Education, Inc.
Hydrogen Bonds
• is the force of attraction between a
hydrogen molecule with a partial
positive charge and another atom or
molecule with a partial or full negative
charge.
Van der Waals Forces
• When molecules come close together, the
attractive forces between slightly positive and
negative regions pull on the molecules and
hold them together
• The strength of the attraction depends on the
size of the molecule, its shape, and its ability
to attract electrons
Van der Waals Interaction
• Each gecko toe has hundreds
of thousands of tiny hairs
with projections at the hairs
tip that increase surface area
• Hair tips and wall surface are
numerous and can support a
geckos weight
Energy and Chemical Reactions
•4 Types of Energy and Chemical Reactions
•1. Chemical Reaction
•2. Energy Transfer
•3. Activation Energy
•4. Electron Transfer
1. Chemical Reactions
• Chemical reactions are the making and
breaking of chemical bonds
• Reactants: the
starting molecule;
are shown on the
left side of an
equation
• Products: the final
molecule; are shown
on the right side of
an equation
Living things undergo thousands of chemical reactions.
2H2 + O2 2 H2O
2. Energy Transfer
• Exergonic Reaction: energy is provided
to your body by sugars from foods &
your body breaks down these sugars &
other substances releasing energy.
• The Chemical reaction that involve a
net release of free energy is called
Exergonic reaction
2. Energy Transfer
• Endergonic Reaction: the reactions that
involve a net absorption of free energy
3. Activation Energy
• Before endergonic & exergonic reactions can
start, energy must be added to the reactants
• Activation Energy is the amount of energy
needed to start the reaction
• Catalysts: reduce the amount of activation
energy that is needed for a reaction
• Enzymes are important classes of catalysts
-a single organism may have thousands of
different enzymes, each made for a different
chemical reaction
4. Oxidation-Reduction Reaction
• Also called redox reactions
• During some chemical reactions,
electrons are transferred from one atom
to another
Oxidation-Reduction Reaction
• Oxidation-loss of electron becoming more
positive in charge
• Reduction-gain of an electron becoming more
negative in charge
– When electrons are transferred this way it keeps it
energy of position
• Energy levels represent different distances
from the nucleus in which electrons are
located
Example Redox Reaction
• Na atom loses an e- (Na+1) to form a
Na ion
• Thus Na +1 undergoes an oxidation
• Cl atom gains an e- (Cl-1) to form a Cl
ion
• Thus Cl –1 undergoes a reduction
Redox Reaction
-
+
Oxidation
+
Reduction
The factors that influence the rate of
chemical reactions
1. Temperature (heat increases rate)
2. Concentration of Reactants and
Products (proceed more quickly when
more reactants are available)
3. Catalysts (substance that increases the
rate of reaction-Enzymes)
Chemistry of Water
The body of this Moon jellyfish
is almost 99% water
Water is considered to be polar
1. They form covalent bonds but do not share
e- equally
2. The region where the O-2 atom is located it
has a partial negative charge
3. The regions of the molecule where each of
the 2 H+1 has a partial positive charge
4. The charge on a water molecule is neutral
& evenly distributed across the water
molecule
5. This makes water Polar
Polar Molecules of water
• Hydrophilic: “water loving” molecules
which readily form H bonds with water
– (polar molecules)
• Polar molecules which form hydrogen bonds with
water cause the molecules to take normal shape
Nonpolar Molecules of water
• Hydrophobic: “fear water” molecules
that shrink from contact with water
– (nonpolar molecules)
• Water molecules always form w/ the maximum
number of H-bonds
• When non-polar molecules (oils) are placed in water,
they exclude
• The non-polar molecules are “forced” into association
with one another which minimizes their disruption of
the bond causing them to shrink
4 Properties of Water
• Are important to Earth’s suitability for
life:
– 1.
– 2.
– 3.
– 4.
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Cohesion
Adhesion
Temperature Moderation
Density
Properties of Water
1.Cohesion-water molecules stick
together b/c of H-bonds
B/C water is cohesive, it is a liquid &
not a gas at moderate temperature
– Example: helps leaves pull water upward
from the roots; seeds swell & germinate
Type of Cohesion
• Surface tension:
results from h-bonds
pulling downward to
the liquid acting like
a thin “skin on the
surface”
• A type of cohesion
Water Strider
Properties of Water
2. Adhesion
• The attractive force
between 2 particles
of different
substances
• Ex: water molecules
& glass molecules
Properties of Water
3. Temperature Moderation Water has the
ability to absorb a relatively large
amount of energy as heat and the
ability to cool surfaces through
evaporation.
• Example: High specific heat
Properties of Water
4.Lower Density of Ice: solid water is denser
than liquid (b/c of the shape of the molecule
& h-bonding)
• water molecules in an ice crystal are spaced
relatively far apart because of H bonding
– Ex: ice is less dense than water, lakes freeze from
the top – down
– Ice insulates the water below from the cold air so
fish can live
Ionization of Water
• Water molecules move about & bump
into each other sometimes causing a
chemical change: -• --1 water molecule loses a proton, and
the other gains a proton.
2 steps how water ionizes
Water can dissociate to form H+ and OHa.
H2O  H+ + OH-
OH- hydroxide ion
b.
H+ Free hydrogen
H+ + H2O H3O
H3O hydronium ion
Alkalinity or Acidity is measured by the
amount of hydronium ions & hydroxide
ions dissolved in a solution
Ionization of Water
• Sometimes a hydrogen ion (H) is
transferred from one water molecule to
another, leaving behind a hydroxide ion
(OH−)
– The proton (H) binds to the other water
molecule, forming a hydronium ion (H3O)
– By convention, H is used to represent the
hydronium ion
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Acids & Bases
An important aspect of a living system
is the degree of its acidity or alkalinity
Is it an acid or base?
• Acid: when the number of hydronium
ions is greater than the number of
hydroxide ions
• Sour Taste, Corrosive
• HCl  H+ + Cl-
Is it an acid or base?
• Base: contains more hydroxide ions
than hydronium ions (alkaline)
• Have bitter taste, slippery
• NaOH  Na + OH
pH
• A scale for comparing concentrations of
hydronium ions and hydroxide ions in a
solution is the pH scale
• It ranges from 0 to 14
• 0 is acidic, 7 neutral, 14 basic
• pH is measured on a logarithmic scale
• Every increase or decrease in units of pH
represents a 10-fold increase or decrease
pH = -log[H+]
pH Logarithmic scale example
• Urine (pH 6) has 10x the H3O than
water (pH 7)
• Vinegar (pH 3) has 1,000x more H3O
than urine (pH 6) & 10,000 more H3O
than water (pH7)
Buffers
• Are chemical substances that neutralize
small amounts of either an acid or base
added to a solution
• Buffers control pH in living systems
• The internal pH of most living cells must
remain close to pH 7
Buffers in your blood
• kidneys and the lungs work together to
help maintain a blood pH of 7.4
• Three different buffer systems exist in
blood, 1. bicarbonate buffer
2. phosphate buffer made of "simple"
chemicals
3. In addition the carbonyl groups (COOH) and the amide group (-NH2)
Acidification: A Threat to Our Oceans
• Human activities such as burning fossil
fuels threaten water quality
• CO2 is the main product of fossil fuel
combustion
– About 25% of human-generated CO2 is
absorbed by the oceans
• CO2 dissolved in seawater forms
carbonic acid; this causes ocean
acidification
• As seawater acidifies, H ions combine
with CO32− ions to form bicarbonate
ions (HCO3–)
• It is predicted that carbonate ion
concentrations will decline by 40% by
the year 2100
• This is a concern because organisms
that build coral reefs or shells require
carbonate ions
Figure 2.24
CO2
CO2  H2O  H2CO3
H2CO3  H  HCO3−
H  CO32−  HCO3−
CO32−  Ca2  CaCO3