Transcript Document

LECTURE PRESENTATIONS
For CAMPBELL BIOLOGY, NINTH EDITION
Jane B. Reece, Lisa A. Urry, Michael L. Cain, Steven A. Wasserman, Peter V. Minorsky, Robert B. Jackson
Chapter 2
The Chemical Context of Life
Lectures by
Erin Barley
Kathleen Fitzpatrick
© 2011 Pearson Education, Inc.
Overview: A Chemical Connection to Biology
• Biology is a multidisciplinary science
• living organisms are subject to basic laws of physics and
chemistry
• this chapter deals with the relevance of chemistry to
biology
• one example are the “devil’s gardens”
•
•
•
•
•
stands of Duroia trees
dominates its landscape
not planted or maintained by humans
called devil’s gardens by the locals
who tends the garden???
Dead leaf tissue (cm2)
after one day
RESULTS
16
12
8
4
0
Inside,
unprotected
Inside,
protected
Outside,
unprotected
Outside,
protected
Cedrela saplings, inside and outside devil’s gardens
•
•
•
Stanford biologists studied the question
one hypothesis was that ants living in the trees produce a poison that kills other
plants
field experiments – planted saplings of another tree Cedrela inside the “garden” –
amongst the Duria trees
•
•
•
wanted to see if the pre-existing Duria trees killed the Cedrela transplants
to see if the ants had any affect on viability - applied an insect barrier to some of the
Cedrela saplings planted; the others were unprotected and vulnerable to the Duria
trees
to control for the location – the same Cedrela saplings also planted outside the
garden
Dead leaf tissue (cm2)
after one day
RESULTS
16
12
8
4
0
Inside,
unprotected
Inside,
protected
Outside,
unprotected
Outside,
protected
Cedrela saplings, inside and outside devil’s gardens
•
•
•
•
•
RESULTS: the unprotected Cedrela saplings inside the garden had the most damage
CONCLUSIONS:
1. the location has no effect – inside protected and outside protected Cedrela trees had the
same viability
2. BUT protection was important – unprotected Cedrela trees planted next to the Duria trees
died
• so the ants had an effect
ants were found to kill non-host trees by injecting formic acid into their leaves
Life: Levels of Organization
•Atoms
•Molecules
•Macromolecules
•Organelles
•Cells
•Tissues
•Organs
•Organ systems
•Organism
Concept: 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
•
made up of elements
• an element is a substance that cannot be broken down into other
substances by chemical reactions
• a compound is a substance consisting of two or more elements in a
fixed ratio
• a compound has characteristics different from those of its
elements
Sodium
Chlorine Sodium chloride
The Elements of Life
•
about 20–25% of the 92 natural elements are essential to life = essential
elements
•
•
•
•
•
needed to live a healthy life and reproduce
Table 2.1
of the total amount of natural elements on Earth - carbon, hydrogen,
oxygen, and nitrogen make up 96%
most of the remaining 4% consists of calcium, phosphorus, potassium, and
sulfur
trace elements are those required by an organism in minute quantities
Concept: An element’s properties depend on the structure
of its atoms
Atom = smallest unit of an element that still retains the chemical &
physical properties of that element
i.e. really, really, really tiny thing!
-composed of subatomic particles:
1. protons = one positive charge, 1 atomic mass unit (1.673x10-24g)
2. electrons = one negative charge, no mass (9.109x10-28g)
3. neutrons = no charge, 1 atomic mass unit (1.673x10-24g)
•
•
•
neutrons and protons form the
atomic nucleus
electrons form a cloud that “orbits”
around the nucleus
neutron mass and proton mass are
almost identical and are measured
in Daltons (Da) – but can be
measured in grams
PERIODIC TABLE OF ELEMENTS
-elements are grouped on a Periodic Table of Elements
-the elements are grouped according to physical and chemical characteristics
-on the chart each element is associated with a letter, an atomic number & an
atomic mass
IA
IIA
http://oxford-labs.com/xray-fluorescence/theperiodic-table/
IIIA IVA VA VIA VIIA VIII
Atomic Number and Atomic Mass
• atoms of the various elements differ in number of subatomic
particles
• an element’s atomic number is the number of protons in its
nucleus
•
•
Atomic number = # protons
this equals the number of electrons orbiting when the atom is electrically
neutral
• an element’s mass number is the sum of protons plus neutrons in
the nucleus
•
Mass number = Protons + Neutrons
• atomic mass = the atom’s total mass
•
•
can be approximated by the mass number
so an atom’s mass number is close to its atomic mass (weight)
© 2011 Pearson Education, Inc.
atomic
symbol
atomic
mass (weight)
7
3
Li
e.g. # protons (e-) = 3
# pr+3 + #No 4 = 7
atomic
number
39
19
K
e.g. # protons (e-) = 19
# pr+19 + #No 20 = 39
Isotopes
• all atoms in an element have the same number of
protons but may differ in number of neutrons
• the number of protons defines what the element is
• e.g. 6 protons = carbon and only carbon
• isotopes are two atoms of an element that differ in
number of neutrons
• in nature an element is a mixture of its isotopes
© 2011 Pearson Education, Inc.
Radioactivity
• the protons and neutrons are held together in the nucleus by
a kind of nuclear “glue”
• when the number of neutrons increase – the nucleus
becomes unstable
• the breakup of the nucleus releases particles with energy in
the form of radioactivity
•
•
also known as radioactive decay
three different kind of particles released – each with different energy
levels
•
•
alpha (helium atom), beta and gamma
the decay can eventually change the # protons – transform one atom
into another
pr+:
e-:
No:
12C
13C
6
6
6
6
6
7
14C
6
6
8** radioactive
Radioactive isotope uses:
1. carbon dating - 14C decay
2. radioactive imaging - e.g. PET scanning
-use of FDG – radioactive glucose tracer
-18F radioactive isotope (2-fluoro-deoxy-glucose)
3. cancer treatment - 60Co, 131I
Cancerous
throat
tissue
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
• differ in their location around the nucleus
• an electron’s state of potential energy (i.e. its
location around a nucleus) is called its energy level
or electron shell
© 2011 Pearson Education, Inc.
Electron Configurations
• “bed check” for electrons
• description on how are electrons organized
around the nucleus of protons and neutrons
• Bohr model: Nils Bohr proposed electrons
“orbit” around the atom’s nucleus in specific
energy levels or orbits (electron shells)
– these shells have a specific energy level
– closer the electron is to the nucleus the less
energy it needs to “orbit”
– to move up to a new electron shell requires
an input of energy
– as the electrons returns to its correct shell
(its “ground” state) – it releases this energy
– his model only works for smaller atoms
– larger atoms are described by quantum
mechanics
(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
• the Bohr model proposed that not only
are there electron shells that surround
the nucleus
• but each shell is comprised of subshells
called orbitals
• an orbital is the three-dimensional
space where an electron pair is found
90% of the time
• each electron shell consists of a specific
number of orbitals
• s, p and d orbitals
First shell
Neon, with two filled
Shells (10 electrons)
Second shell
(a) Electron distribution diagram
First shell
Second shell
y
x
1s orbital
2s orbital
z
Three 2p orbitals
(b) Separate electron orbitals
1s, 2s, and
2p orbitals
(c) Superimposed electron orbitals
–
–
–
–
–
–
–
–
each orbital holds a pair of electrons
s orbital = 2 electrons maximum
p orbitals = 6 electrons maximum
d orbitals = 10 electrons maximum
1st shell – closest to the nucleus - holds 2 electrons (1 s orbital only)
2nd shell can hold 8 (1 s and 3 p orbitals – 2 + 6 electrons)
3rd holds 18 (1 s, 3 p and 5 d orbitals – 2 + 6 + 10 electrons)
4th holds 18 (1 s, 3 p and 5 d orbitals – 2 + 6 + 10 electrons)
Electron Distribution and Chemical Properties
• the chemical behavior of an atom is determined by the distribution
of electrons in electron shells
•
the reactivity of certain atoms results from the presence of unpaired
electrons in one or more orbitals
• an atom will always try to complete its outermost shell = valence
shell
• basis for bonding reactions
• the number of electrons in the outer most electron shell
involved in bonding reactions = valence electrons
• chemists really only consider the electrons in the s and p orbitals
as valence electrons
• once an atom completes fills up its valence orbitals – it is chemically
inert
•
unable to participate in bonding reactions
© 2011 Pearson Education, Inc.
Valence = 1
Valence = 4
Valence = 3
Valence = 2
VIII
I
Hydrogen
1H
Electron
distribution
diagram
First
shell
II
Lithium
3Li
Beryllium
4Be
III
Boron
5B
IV
Helium
2He
V
VI
VII
Carbon
6C
Nitrogen
7N
Oxygen
8O
Fluorine
9F
Neon
10Ne
Silicon
14Si
Phosphorus
15P
Sulfur
16S
Chlorine
17Cl
Argon
18Ar
Second
shell
Sodium Magnesium Aluminum
11Na
12Mg
13Al
Third
shell
•
the periodic table of the elements tells you the electron distribution for each element
•
•
•
•
•
by its row and column position
1st row – valence electrons in the first electron shell
2nd row – valence electrons in the second electron shell
1st column – 1 valence electron
3rd column – 3 valence electrons
Concept: The formation and function of molecules
depend on chemical bonding between atoms
• atoms with incomplete valence shells can share or transfer
valence electrons with certain other atoms to form molecules
• molecule - particle formed by the union of more than one atom
•
e.g. same kind of atom - O2
•
e.g. different types of atoms - H20
• formed by held by attractions called chemical bonds
• two types of chemical bonds
© 2011 Pearson Education, Inc.
Ionic Bonds
• atoms sometimes strip electrons from their bonding partners
•
•
•
•
•
•
an example is the transfer of an electron from sodium to chlorine
after the transfer of an electron, both atoms have charges
a charged atom (or molecule) is called an ion
a cation is a positively charged ion
an anion is a negatively charged ion
an ionic bond is an attraction between an anion and a
cation
Na
Sodium atom
Cl
Chlorine atom
+
–
Na+
Sodium ion
(a cation)
Cl–
Chloride ion
(an anion)
Sodium chloride (NaCl)
•
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
HINT: elements in columns I and II form ionic bonds with the
elements in column VII
Na+
Cl–
© 2011 Pearson Education, Inc.
Covalent Bonds
• if it isn’t favorable for an atom to gain or lose an electron
- it will have to share it with another
• covalent bond = bond in which atoms share electrons
• atoms that like to form covalent bonds
• oxygen
• nitrogen
• carbon
-usually forms when one atom has
to lose or gain three or more
electrons
Hydrogen atoms (2 H)
e.g. carbon would have to gain 4 valence
electrons to complete its outer shell, nitrogen
would have to gain 3 valence electrons
-can also form between two identical atoms
e.g. nitrogen (N3), oxygen gas (O2),
hydrogen gas (H2)
Hydrogen molecule (H2)
Covalent Bonds for the Biologist
•
•
•
•
•
•
•
covalent bonding by some common biological elements
Hydrogen = valence 1; 1 electron needed; 1 covalent bond
Oxygen = valence 2; 2 electrons needed; 2 covalent bonds
Sulfur = valence 2; 2 electrons needed; 2, 4 or 6 covalent bonds
Nitrogen = valence 3; 3 electrons needed; 3 or 4 covalent bonds
Carbon = valence 4; 4 electrons needed; 4 covalent bonds
Phosphorus = valence 3; 3 electrons needed; 5 covalent bonds
© 2011 Pearson Education, Inc.
• so we can now modify the definition of a molecule
further
• a 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
•
is notated in the structural formula as H-H
• a double covalent bond, or double bond, is the sharing
of two pairs of valence electrons
•
is notated in the structural formula as C=C
© 2011 Pearson Education, Inc.
Figure 2.12
Name and
Molecular
Formula
(a) Hydrogen (H2)
(b) Oxygen (O2)
(c) Water (H2O)
(d) Methane (CH4)
Lewis Dot SpaceElectron
Distribution Structure and Filling
Structural
Model
Diagram
Formula
Polar and Non-Polar Covalent Bonds
• 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
• modifies the covalent bond
• what are the two categories of covalent bond?
Polar and Non-Polar Covalent Bonds
• in a nonpolar covalent bond, the atoms share the
electron equally
• carbon frequently forms non-polar covalent bonds
Polar and Non-Polar Covalent Bonds
• in a 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
•
•
annotated with the greek letter delta
positive and negative ends are known as dipoles
• oxygen is very electronegative – creates polar bonds
• nitrogen is also very electronegative
–
• \
O
+
H
H
H2O
+
Weak Chemical Bonds
• the strongest bonds in organisms are covalent bonds
that form a cell’s molecules
• BUT weak chemical bonds, such as ionic bonds and
hydrogen bonds, are also important
• weak chemical bonds reinforce shapes of large molecules
• help molecules adhere to each other
• weak non-covalent bonds are transient
• constantly breaking and reforming at room temperature
• but together multiple, weak non-covalent bonds can
produce highly stable structures
• most common weak chemical bonds?
A Weak Bond: Hydrogen Bond
• 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
•
•
e.g. seen between the bases of the
DNA double helix
e.g. seen between two water
molecules
• not really a bond but an
interaction between two dipoles
• weaker than covalent and ionic
© 2011 Pearson Education, Inc.
+
–
Water (H2O)
+
Hydrogen bond
–
Ammonia (NH3)
+
+
+
Hydrogen bonds are
intermolecular bonds
Concept: Polar covalent bonds in water molecules result in
hydrogen bonding
•
•
the water molecule is a polar molecule: the opposite ends have opposite
charges
polarity allows water molecules to form hydrogen bonds with each other

Hydrogen
bond
+
+
Polar covalent
bonds


+
© 2011 Pearson Education, Inc.
+

Another Weak Bond: Van der Waals
• if electrons are distributed asymmetrically in
molecules or atoms, they can result in “hot
spots” of positive or negative charge
• Van der Waals interactions are attractions
between molecules that are close together
as a result of these charges
•
•
named after the physicist Johannes Diedrick Van
der Waals
often used to describe the totality of
intermolecular forces
• collectively, these bonds can be strong
•
e.g. as between molecules of a gecko’s toe hairs
and a wall surface
van der Waals bonds are intermolecular bonds
Molecular Shape and Function
• a molecule’s shape is usually very important to its
function
• molecules such as H2 or O2 are always linear
• but others like H2O have a 3D conformation
s orbital
•
•
a molecule’s shape is determined
by the positions of its atoms’
valence orbitals
when a covalent bond forms – the
valence electrons undergo
rearrangement
•
•
•
this creates specific molecular
shapes
for methane - creates a tetrahedron
of four tear-drop like orbitals
•
•
specifically the electrons in the s and p
orbitals may hybridize with each other as
they shift and rearrange within the
molecule
Four hybrid orbitals
z
Three p orbitals
x
y
Tetrahedron
(a) Hybridization of orbitals
Space-Filling
Model
Ball-and-Stick
Model
Unbonded
Electron
pair
Water (H2O)
symmetrical shape
for water – creates a V
•
assymetrical shape (angle of 104.5)
Hybrid-Orbital Model
(with ball-and-stick
model superimposed)
Methane (CH4)
(b) Molecular-shape models
Carbon
Hydrogen
Natural endorphin
•
•
biological molecules
recognize and interact with
each other based on
molecular shape
creates specificity
•
•
Nitrogen
Sulfur
Oxygen
Morphine
e.g. interaction of a hormone
with its receptor
e.g. interaction of an enzyme
with its substrate
(a) Structures of endorphin and morphine
•
•
molecules with similar
shapes can have similar
biological effects
IN OTHER WORDS – you
can predict function by
looking at 3D structure
Natural
endorphin
Brain cell
Morphine
Endorphin
receptors
(b) Binding to endorphin receptors
Concept: Chemical reactions make and break
chemical bonds
• Chemical reactions are the making and breaking of chemical
bonds
• The starting molecules of a chemical reaction are called
reactants
• The final molecules of a chemical reaction are called products
2 H2
+
Reactants
O2
2 H2O
Reaction
Products
Chemical reactions:
3 types:
1. Synthesis - A + B
2. Decomposition - AB
reactions)
3. Exchange - AB + CD
AB (Anabolism reactions)
A + B (Catabolism
AD + BC
-these equations must be balanced
-Law of conservation of Mass or “chemical book keeping”
-i.e. the number of atoms of each element is the same before
and after a chemical reaction
• All chemical reactions are reversible:
products of the forward reaction become
reactants for the reverse reaction
• Chemical equilibrium is reached when the
forward and reverse reaction rates are equal
Water: The Solvent of Life
• major component of blood plasma, interstitial fluid, CSF, cytosol
etc…
• hydrogen bonding allows the even distribution of dissolved
substances throughout our system – so water is an excellent
transport medium
 role in : transporting chemicals, waste products
•solvent – substance that dissolves solutes
• can assist in chemical reactions
•water is a polar solvent
+
• dissolves polar solutes
•water is the universal solvent for polar compounds – facilitates
most chemical reactions in the body
O
H
H
+
Concept: Four emergent properties of water contribute to its
importance for life
• four of water’s properties that facilitate an
environment for life are
–
–
–
–
Cohesive behavior
Ability to moderate temperature
Expansion upon freezing
Versatility as a solvent
#1: Cohesion of Water Molecules
• collectively, hydrogen bonds hold
multiple water molecules together
in a well-ordered structure
– this attractive phenomenon is
called cohesion
– at any given time many of
water’s molecules are linked by
hydrogen bonds
• cohesion: attraction between the
same molecules
• explains the capillary action of
water we see in plants
cohesive force helps the transport of water and
its dissolved substances against gravity in plants
- pulls water molecules up the plant following evaporation
from the leaf surface
Property #1: Cohesion of Water Molecules
• DON’T CONFUSE COHESION WITH ADHESION!!!
• adhesion is an attraction between different molecules
–
for example, between water and plant cell walls
• cohesion can also be observed in other liquids
–
e.g. mercury – cohesive attraction between Hg molecules is stronger than the adhesive
attraction to another surface
• collectively – all the water molecules connected to each other creates a cohesive
force
• surface tension is a measure of how hard it is to stretch or break the surface of a
liquid
–
water has a greater surface tension than most liquids
• surface tension is related to cohesion
–
–
at the water-air interface is a thin film of well-ordered water molecules – all hydrogen
bonded to one another
creates a strong “invisible film” at this interface = known as surface tension
Property #2: Moderation of Temperature by Water
• water absorbs heat from warmer air and releases stored heat
to cooler air
• water can absorb or release a large amount of heat with only
a slight change in its own temperature
• so water is great at stabilizing temperature
• BUT – what is heat and temperature?
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Heat and Temperature
• kinetic energy is the energy of motion
• heat is a measure of the total amount of kinetic energy due to molecular
motion
• high motion = higher heat
• temperature measures the intensity of heat due to the average kinetic energy
of molecules
–
–
–
•
the unit of heat expended in a reaction is a calorie (cal)
–
–
•
temperature on Earth is measured with the Celsius scale (ºC)
but the true measure of temperature is in Kelvin
absolute zero (0ºK) – when all molecular motion stops
the amount of heat required to raise the temperature of 1 g of water by 1 ºC
the “calories” on food packages are actually kilocalories (kcal), where 1 kcal =
1,000 cal
the joule (J) is another unit of energy where
1 J = 0.239 cal, or 1 cal = 4.184 J
© 2011 Pearson Education, Inc.
Water’s High Specific Heat
• the ability of water to stabilize temperature comes from its
relatively high specific heat
• specific heat = amount of heat that must be absorbed or lost for
1 g of that substance to change its temperature by 1ºC
–
–
–
–
specific heat of water is 1 cal/g/ºC
specific heat of ethanol is 0.6 cal/g/ºC
specific heat of iron – ten times less than water
so its easier and faster to change the temperature of iron vs. water
• so because of this - water changes its temperature very little
when it absorbs or gives off heat
• think of specific heat as how well a substance resists a
temperature change
© 2011 Pearson Education, Inc.
• water’s high specific heat can be traced to hydrogen bonding
– heat must be absorbed for hydrogen bonds to break
– heat is released when these hydrogen bonds reform
• the high specific heat of water minimizes temperature fluctuations
to within limits that permit life
Burbank
90°
Santa Barbara 73°
Los Angeles
(Airport) 75°
70s (°F)
80s
San Bernardino
100°
Riverside 96°
Santa Ana
84°
Palm Springs
106°
Pacific Ocean 68°
90s
100s
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San Diego 72°
40 miles
Water and Evaporative Cooling
• molecules of a liquid stay close together due to cohesive attraction
• if they move fast enough (due to increased kinetic energy) – they can overcome
this attraction and be released
–
–
i.e. increase in kinetic energy can result from an increase in heat
molecules are released as a gas
• transformation from liquid to gas occurs = evaporation
• heat of vaporization is the heat a liquid must absorb for 1 g to be converted to
gas
–
water has a high heat of vaporization due to the strength of its hydrogen bonding
• as a liquid evaporates – the remaining surface cools - a process called
evaporative cooling
–
the “hottest” molecules leave as a gas and the remaining liquid cools
• evaporative cooling of water helps stabilize temperatures in organisms and
bodies of water
© 2011 Pearson Education, Inc.
Property #3: Expansion of Ice Upon Freezing
• water is one of the few substances that is less dense as a solid than as a liquid
–
water expands when heated and contracts as it cools
• in other words ice floats in liquid water
• below 4°C water begins to freeze – molecules move too slowly to break their
hydrogen bonds – molecules become locked in place (crystalline formation)
• distance between molecules when frozen is greater than in liquid form
–
makes ice about 10% less dense vs. liquid form
• when ice melts – the hydrogen bonds are broken and the molecules ‘slip’ closer to
each other
Hydrogen bond
Liquid water:
Hydrogen bonds
break and re-form
Ice:
Hydrogen bonds
are stable
*** if ice sank, all bodies of water would eventually freeze solid, making life impossible
on Earth****
Property #4: Water - The Solvent of Life
• a solution is a liquid that is a homogeneous mixture of 2
or more solutes
– the solvent is the dissolving agent of a solution
– the solute is the substance that is dissolved
• an aqueous solution is one in which water is the solvent
• water is a universal solvent because it dissolves just about
everything
– except non-polar compounds
– e.g. oil and water
• water is a versatile solvent due to its polarity
– allows it to form hydrogen bonds easily with other charged molecules
• when an ionic compound is dissolved in water - each ion becomes surrounded by
water molecules - known as a solvation or hydration shell
-water + salt: the electronegative O- of water attracts the +ve sodium in the
salt crystal and pulls it away
- the electropositive H+ of water attracts the –ve chlorine and pulls it away
-the crystal lattice of salt is eventually broken up and each Na+ and Clbecomes surrounded by water molecules
• water can also dissolve compounds made of polar molecules
• even large polar molecules such as proteins can dissolve in water if they have
ionic and polar regions
+


+
Hydrophilic and Hydrophobic Substances
• a hydrophilic substance is one that has an affinity for water
• a hydrophobic substance is one that does not have an affinity
for water
• oil molecules are hydrophobic because they are relatively
nonpolar in terms of their bonds
Solutions, Colloids and Suspensions
•
mixture = two or more types of elements or molecules physically blended together in a solvent
without the formation of physical bonds between them
•
1. solution = homogenous (same) mixture of substances dissolved in a solvent
–
–
–
•
e.g. sugar + water
the mixture is the same no matter where you sample it
substances are very small and are not acted upon by gravity – remain suspended in the solution
2. colloid = solution of larger components called dispersed-phase particles
–
–
–
their particles are larger than that of solutions
have the potential to settle out due to gravity
BUT: these particles all carry the same charge (repel each other) – so they remain suspended in the solvent
•
•
e.g. plasma proteins within the blood
3. suspension = solution of even larger components that are solid and can settle out
–
–
larger particles than that of colloid
if left undisturbed these particles will settle out to form a solid
•
e.g. red blood cells within blood
Solute Concentration in Aqueous Solutions
• most biochemical reactions occur in water
• chemical reactions depend on collisions of molecules and
therefore on the concentration of solutes in an aqueous
solution
• the number of molecules is usually measured in moles, where 1
mole (mol) = 6.02 x 1023 molecules
– known as Avogadro’s number
– the unit Dalton (Da) is defined such that 6.02 x 1023 daltons = 1 g
• Molarity (M) is the number of moles of solute per liter of
solution
• to calculate the number of moles – you need to know molecular
mass = total mass of all atoms in a molecule
Concept: Acidic and basic conditions affect living
organisms
• between two water molecules something unique can happen
• a hydrogen atom in a hydrogen bond can shift from one to the
other
– the hydrogen atom leaves its electron behind and is transferred as a
proton, or hydrogen ion (H+)
– the molecule with the extra proton is now a hydronium ion (H3O+),
though it is often represented as H+
– the molecule that lost the proton is now a hydroxide ion (OH–)
+
2 H 2O
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Hydronium
ion (H3O+)

Hydroxide
ion (OH)
Acids and Bases
• an acid is any substance that increases the H+
concentration of a solution
• a base is any substance that reduces the H+
concentration of a solution
• 1. release H+  Acids
e.g. HCl  H+ + Cl2. release ions to combine with H+  Bases
e.g. NaOH Na+ + OH3. acids + bases  Salts
e.g. HCl + NaOH  H20 + NaCl
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The pH Scale
• biologists use something called the pH scale to describe whether a solution
is acidic or basic (the opposite of acidic)
• in any aqueous solution at 25C the product of H+ and OH– is constant and
can be written as
[H+][OH–] = 10–14
• the pH of a solution is defined by the negative logarithm of H+
concentration, written as
pH = –log [H+]
• for a neutral aqueous solution, [H+] is 10–7, so
pH = –log(–7) = 7
H+
H+
 H+
H+ OH
+
OH H H+
+
H H+
Acidic
solution
1
Battery acid
2
Gastric juice, lemon juice
3
Vinegar, wine,
cola
4
Tomato juice
Beer
Black coffee
5
6
OH
OH
H+ H+ OH

OH OH +
+
H
H
H+
Neutral
solution
OH
OH
OH H+ OH

OH OH

H+ OH
Basic
solution
Neutral
[H+] = [OH]
7
8
Increasingly Basic
[H+] < [OH]
• Acidic solutions have
pH values less than 7
• Basic solutions have pH
values greater than 7
• Most biological fluids
have pH values in the
range of 6 to 8
Increasingly Acidic
[H+] > [OH]
pH Scale
0
Rainwater
Urine
Saliva
Pure water
Human blood, tears
Seawater
Inside of small intestine
9
10
Milk of magnesia
11
Household ammonia
12
13
Household
bleach
Oven cleaner
14
Buffers:
• biological fluids need to remain relatively neutral
• chemical or compound that keeps the pH of a solution within
a normal range
• resists pH change by taking up excess H+ or OH- ions
H20
H2CO3
e.g. blood = pH 7.4
“Bicarbonate buffering system”
H+ +
HCO3-
carbonic acid
excess OH-
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excess H+
Acidification: A Threat to Water Quality
• human activities such as
burning fossil fuels threaten
water quality
• CO2 is the main product of
fossil fuel combustion
• about 25% of humangenerated CO2 is absorbed by
the oceans
• CO2 dissolved in sea water
forms carbonic acid; this
process is called ocean
acidification
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CO2
CO2 + H2O
H2CO3
H2CO3
H+ + HCO3
H+ + CO32
CO32 + Ca2+
HCO3
CaCO3
• as seawater acidifies, H+ ions combine with carbonate ions to
produce bicarbonate
–
carbonate is required for calcification (production of calcium carbonate) by many
marine organisms, including reef-building corals
• BUT the burning of fossil fuels is also a major source of sulfur oxides
and nitrogen oxides
–
these compounds react with water in the air to form strong acids that fall in rain or
snow
• Acid precipitation is rain, fog, or snow with a pH lower than 5.2
–
acid precipitation damages life in lakes and streams and changes soil chemistry on land
(a)
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(b)
(c)