Transcript chapter

Eldra Solomon
Linda Berg
Diana W. Martin
www.cengage.com/biology/solomon
Chapter 2
Atoms and Molecules:
The Chemical Basis of Life
Albia Dugger • Miami Dade College
Chemistry and Life
• All organisms share fundamental similarities in their chemical
composition and basic metabolic processes
• The structure of atoms determines the way they form
chemical bonds to produce complex compounds
• molecular biology
• Chemistry and physics of the molecules that constitute
living things
Inorganic and Organic Compounds
• inorganic compounds
• Small, simple substances
• Biologically important groups include water, simple acids
and bases, and simple salts
• organic compounds
• Generally large, complex carbon-containing compounds
• Typically, two or more carbon atoms are bonded to each
other to form the backbone, or skeleton, of the molecule
2.1 ELEMENTS AND ATOMS
LEARNING OBJECTIVES:
•
•
Name the principal chemical elements in living things
and provide an important function of each
Compare the physical properties (mass and charge) and
locations of electrons, protons, and neutrons.
Distinguish between the atomic number and the mass
number of an atom
•
Define the terms orbital and electron shell; relate
electron shells to principal energy levels
Elements
• elements
• Substances that can’t be broken down into simpler
substances by ordinary chemical reactions
• Each element has a chemical symbol (Example: C for
carbon)
• Four elements (oxygen, carbon, hydrogen, and nitrogen)
make up more than 96% of the mass of most organisms
• Calcium, phosphorus, potassium, and magnesium, are
present in smaller quantities
• Iodine and copper are trace elements
Functions of Elements in Organisms
Table 2-1, p. 27
Atoms and Matter
• atom
• Smallest unit of an element that retains that element’s
chemical properties
• Made up of tiny subatomic particles of matter
• matter
• Anything that has mass and takes up space
Subatomic Particles
• There are three basic types of subatomic particles:
• An electron carries a unit of negative electric charge
• A proton carries a unit of positive charge
• A neutron is an uncharged particle
• Protons and neutrons compose the atomic nucleus
• Electrons move rapidly around the atomic nucleus
• In an electrically neutral atom, the number of electrons equals
the number of protons
Atomic Number and the Periodic Table
• Every element has a fixed number of protons in the atomic
nucleus (atomic number) which determines an atom’s
identity and defines the element
• The periodic table is a chart of the elements arranged in
order by atomic number and chemical behavior
• Bohr models represent the electron configurations of
elements as a series of concentric rings
The Periodic Table
Chemical symbol
Atomic number
Chemical name
Number
of e– in
each
energy
level
Fig. 2-1, p. 28
Atomic Mass
• The mass of a subatomic particle is expressed in terms of the
atomic mass unit (amu) or dalton
• One amu equals the approximate mass of a single proton or a
single neutron; an electron is about 1/1800 amu
• The atomic mass of an atom equals the total number of
protons and neutrons, expressed in amus or daltons
Characteristics of Subatomic Particles
Particle
Proton
Neutron
Electron
Charge
Positive
Neutral
Negative
~Mass
1 amu
1 amu
~1/1800 amu
Location
Nucleus
Nucleus
Outside nucleus
Isotopes
• Most elements consist of a mixture of atoms with different
numbers of neutrons and different masses
• isotopes
• Atoms of the same element (having the same number of
protons and electrons) with varying numbers of neutrons
• The mass of an element is expressed as an average of the
masses of its isotopes
Isotopes of Carbon
Radioisotopes
• Some isotopes are unstable and tend to break down (decay)
to a more stable isotope (usually a different element)
• radioisotope
• Unstable isotope that emits radiation as it decays
• Example: 14C decays to 14N when a neutron decomposes
to form a proton and a fast-moving electron
• Radioactive decay can be detected by autoradiography, on
photographic film
Radioisotopes in Biology
• Radioisotopes such as 3H (tritium), 14C, and 32P can replace
normal molecules and are used as tracers in research
• In medicine, radioisotopes are used for both diagnosis (such
as thyroid function or blood flow) and treatment (such as
cancer)
Autoradiography
• Tritium (3H) incorporated into the DNA of a fruit fly
Atomic Orbitals and Energy
• Electrons move through characteristic regions of 3-D space
(orbitals), each containing a maximum of 2 electrons
• The energy of an electron depends on the orbital it occupies
• Electrons in orbitals with similar energies (the same principal
energy level) make up an electron shell
• Electrons farther from the nucleus generally have greater
energy than those closer to the nucleus
Valence Electrons
• The most energetic electrons (valence electrons) occupy the
valence shell, represented as the outermost concentric ring
in a Bohr model
• Valence electrons participate in chemical reactions
• An electron can move to a higher orbital by receiving more
energy, or give up energy and sink to a lower orbital
• Changes in electron energy levels are important in energy
conversions in organisms
Atomic Orbitals
Nucleus
1s
(a) The first principal energy level contains
a maximum of 2 electrons, occupying a
single spherical orbital (designated 1s). The
electrons depicted in the diagram could be
present anywhere in the blue area.
Fig. 2-4a, p. 30
2s
2py
2px
2pz
(b) The second principal energy level includes four orbitals, each with a
maximum of 2 electrons: one spherical (2s) and three dumbbell-shaped (2p)
orbitals at right angles to one another.
Fig. 2-4b, p. 30
z
1s
2s
y
2py
2px
x
2pz
(c) Orbitals of the first and second principal energy levels of a neon atom
are shown superimposed. Note that the single 2s orbital plus three 2p
orbitals make up neon's full valence shell of 8 electrons. Compare this more
realistic view of the atomic orbitals with the Bohr model of a neon atom at
right.
Fig. 2-4c, p. 30
(d) Neon atom (Bohr model)
Fig. 2-4d, p. 30
ANIMATION: The shell model of electron
distribution
To play movie you must be in Slide Show Mode
PC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE
Key Concepts 2.1
• Carbon, hydrogen, oxygen, and nitrogen are the most
abundant elements in living things
ANIMATION: Atomic number, mass number
To play movie you must be in Slide Show Mode
PC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE
ANIMATION: Electron arrangement in atoms
To play movie you must be in Slide Show Mode
PC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE
Animation: Electron distribution
ANIMATION: Subatomic particles
To play movie you must be in Slide Show Mode
PC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE
2.2 CHEMICAL REACTIONS
LEARNING OBJECTIVES:
• Explain how the number of valence electrons of an atom
is related to its chemical properties
• Distinguish among simplest, molecular, and structural
chemical formulas
• Explain why the mole concept is so useful to chemists
Valence Electrons
• Chemical behavior of an atom is determined by the number
and arrangement of its valence electrons
• Atoms with full valence shells are unreactive
• When the valence shell is not full, an atom tends to lose, gain,
or share electrons to achieve a full outer shell
• Elements in the same vertical column (group) of the periodic
table have similar chemical properties
Compounds and Molecules
• Two or more atoms may combine chemically
• A chemical compound consists of atoms of two or more
different elements combined in a fixed ratio
• Two or more atoms joined very strongly form a stable
molecule
• Example: H20 (water) is a molecular compound
Chemical Formulas
• A chemical formula is a shorthand expression that describes
the chemical composition of a substance
• In a simplest formula (empirical formula), subscripts give the
smallest ratios for atoms in a compound (e.g. NH2)
• A molecular formula gives the actual numbers of each type
of atom per molecule (e.g. N2H4)
• A structural formula shows the arrangement of atoms in a
molecule (e.g. water, H—O—H)
The Mole
• The molecular mass of a compound equals the sum of the
atomic masses of the component atoms of a single molecule
• The amount of a compound whose mass in grams is
equivalent to its molecular mass is 1 mole (mol)
• Example:
• Molecular mass of water (H2O) is (hydrogen: 2 × 1 amu)
+ (oxygen: 1 × 16 amu) = 18 amu
• 1 mol of water is 18 grams (g)
The Mole (cont.)
• 1 mol of any substance has exactly the same number of
atoms or molecules: 6.02 × 1023 (Avogadro’s number)
• Avogadro’s number allows scientists to calculate the number
of atoms or molecules in sample simply by weighing it
• A 1 molar solution (1 M) contains 1 mol of a substance
dissolved in a total volume of 1 liter (L)
Chemical Reactions
• Chemical reactions, such as the reaction between glucose
and oxygen, are described by chemical equations:
C6H12O6 + 6 O2 → 6 CO2 + 6 H2O + energy
• Substances that participate in the reaction (reactants) are
written on the left side of the arrow
• Substances formed by the reaction (products) are written on
the right side
Chemical Reactions (cont.)
• Many reactions proceed forward and reverse simultaneously
• At dynamic equilibrium, the rates of forward and reverse
reactions are equal
CO2 + H2O ↔ H2CO3
• When this reaction reaches equilibrium, there will be more
reactants (CO2 and H2O) than product (H2CO3)
Key Concepts 2.2
• The chemical properties of an atom are determined by its
highest-energy electrons, known as valence electrons
ANIMATION: Chemical bookkeeping
To play movie you must be in Slide Show Mode
PC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE
Animation: Covalent bonds
2.3 CHEMICAL BONDS
LEARNING OBJECTIVE:
• Distinguish among covalent bonds, ionic bonds,
hydrogen bonds, and van der Waals interactions
• Compare them in terms of the mechanisms by which they
form and their relative strengths
Chemical Bonds
• Atoms can be held together by chemical bonds
• Valence electrons dictate how many bonds an atom can form
• bond energy
• Energy necessary to break a chemical bond
• Two types of strong chemical bonds: covalent and ionic
Covalent Bonds
• Covalent bonds involve sharing electrons between atoms in
a way that fills each atom’s valence shell
•
A molecule consists of atoms joined by covalent bonds
• Example: hydrogen gas (H2)
• Unlike atoms linked by covalent bonds form a covalent
compound
Lewis Structure
• A simple way of representing valence electrons is to use dots
placed around the chemical symbol of the element:
• Oxygen (6 valence electrons) shares electrons with two
hydrogen atoms to complete its valence shell of 8 – each
hydrogen atom completes a valence shell of 2
Carbon Bonds
• Carbon has 4 electrons in its valence shell, all of which are
available for covalent bonding (e.g. methane, CH4)
• Each orbital holds a maximum of 2 electrons
Single, Double, and
Triple Covalent Bonds
• When one pair of electrons is shared between two atoms, the
covalent bond is called a single covalent bond
• A double covalent bond is formed when two pairs of
electrons are shared (represented by two parallel solid lines)
• A triple covalent bond is formed when three pairs of
electrons are shared (represented by three parallel solid lines)
Electron Sharing
in Covalent Compounds
Hydrogen (H)
Hydrogen (H)
Molecular hydrogen (H2)
or
H H
(a) Single covalent bond formation. Two hydrogen atoms achieve stability by
sharing a pair of electrons, thereby forming a molecule of hydrogen. In the
structural formula on the right, the straight line between the hydrogen atoms
represents a single covalent bond.
Fig. 2-5a, p. 33
Oxygen (O)
Oxygen (O)
Molecular oxygen (O2)
(double bond is formed)
or
O
O
(b) Double covalent bond formation. In molecular oxygen, two oxygen atoms
share two pairs of electrons, forming a double covalent bond. The parallel
straight lines in the structural formula represent a double covalent bond.
Fig. 2-5b, p. 33
Bonds Found in
Biologically Important Molecules
•
•
•
•
•
•
Atom
Hydrogen
Oxygen
Carbon
Nitrogen
Phosphorus
Sulfur
Symbol
H
O
C
N
P
S
Covalent Bonds
1
2
4
3
5
2
Molecular Shape and Function
• The functions of molecules in living cells are determined
largely by their geometric shapes
• When atoms form covalent bonds, orbitals in valence shells
may become rearranged (orbital hybridization), affecting the
shape of the resulting molecule
• Example: In methane (CH4), the hybridized valence shell
orbitals of the carbon form a tetrahedron
Tetrahedron: Methane (CH4)
• Geometric shape of a
molecule provides the
optimal distance
between atoms to
counteract repulsion of
electron pairs
Methane (CH4)
Fig. 2-6, p. 34
Polar and Nonpolar Covalent Bonds
• electronegativity
• A measure of an atom’s attraction for shared electrons in
chemical bonds (e.g. oxygen has high electronegativity)
• nonpolar covalent bond
• When covalently bonded atoms have similar
electronegativities, electrons are shared equally
• polar covalent bond
• Covalent bond between atoms that differ in
electronegativity; electrons are pulled closer to the nucleus
of the atom with greater electron affinity
Polar Molecules
• A polar covalent bond has two dissimilar ends (poles), one
with a partial positive charge and the other partially negative
• A polar molecule has one end with a partial positive charge
and another end with a partial negative charge
• Example: Water has a partial positive charge at the hydrogen
end and a partial negative charge at the oxygen end, where
“shared” electrons are more likely to be
Water: A Polar Molecule
Hydrogen (H)
Oxygen (O)
Hydrogen (H)
Fig. 2-7, p. 34
Oxygen part
Hydrogen parts
Partial
negative
charge at
oxygen end
of molecule
Partial
positive
charge at
hydrogen
end of
molecule
Water molecule (H2O)
Fig. 2-7b, p. 34
Ions
• An atoms or group of atoms with 1 or more units of electric
charge is called an ion
• Atoms with 1, 2, or 3 valence electrons tend to lose electrons
to other atoms and become positively charged cations
• Atoms with 5, 6, or 7 valence electrons tend to gain electrons
from other atoms and become negatively charged anions
Functions of Ions
• The properties of ions are different from those of the
electrically neutral atoms from which they were derived
• Electric charges of cations and anions provide a basis for
energy transformations within the cell, transmission of nerve
impulses, muscle contraction, and other biological processes
Ions and Biological Processes
• Sodium, potassium, and
chloride ions are
essential for a nerve cell
to stimulate muscle
fibers, initiating a
muscle contraction
• Calcium ions in the
muscle cell are required
for muscle contraction
Muscle fiber
Nerve
Fig. 2-8, p. 35
Ionic Bonds
• An ionic bond is formed by attraction between the positive
charge of a cation and the negative charge of an anion
• An ionic compound is a substance consisting of anions and
cations bonded by their opposite charges
• Example: Sodium chloride (NaCl), an ionic compound
• When sodium reacts with chlorine, sodium’s single
valence electron is transferred completely to chlorine
• Sodium becomes a cation (Na+); chlorine becomes an
anion (Cl−)
Ionic
Bonding
11 protons
17 protons
and
11 electrons
Sodium (Na)
10 electrons
Sodium ion (Na+)
17 electrons
Chlorine (Cl)
18 electrons
Chloride ion (Cl–)
Sodium chloride (NaCl)
Arrangement of atoms
in a crystal of salt
Fig. 2-9, p. 35
ANIMATION: Ionic bonding
To play movie you must be in Slide Show Mode
PC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE
Ionic Compounds in Solution
• In the absence of water, ionic bonds are very strong
• Example: Electrical attraction in ionic bonds holds Na+ and Cl−
together to form NaCl (sodium chloride, table salt)
• When placed in water, ionic compounds (such as sodium
chloride) tend to dissociate into individual ions :
NaCl (in H2O) → Na+ + Cl–
Water as a Solvent
• Water is an excellent solvent, capable of dissolving many
substances (solutes)
• Because of their polarity, water molecules easily dissolve
polar or ionic substances
• In solution, each cation or anion is surrounded by oppositely
charged ends of the water molecules (hydration)
Hydration of an Ionic Compound
Salt
Fig. 2-10, p. 36
ANIMATION: Spheres of hydration
To play movie you must be in Slide Show Mode
PC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE
Hydrogen Bonds
• Hydrogen bonds are relatively weak bonds (easily formed
and broken) that are very important in living organisms
• When hydrogen combines with a relatively electronegative
atom, it acquires a partial positive charge
• Hydrogen bonds form between an atom with a partial
negative charge and a hydrogen atom that is covalently
bonded to oxygen or nitrogen
• Water molecules interact with one another extensively
through hydrogen bond formation
Hydrogen Bonding
ANIMATION: Examples of hydrogen bonds
To play movie you must be in Slide Show Mode
PC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE
van der Waals Interactions
• Adjacent molecules may interact in transient regions of weak
positive and negative charge
• The resulting attractive forces (van der Waals interactions)
operate over very short distances and are weaker and less
specific than other types of interactions
• They are important when they occur in large numbers and
when molecular shapes permit close contact between atoms
Key Concepts 2.3
• A molecule consists of atoms joined by covalent bonds
• Other important chemical bonds include ionic bonds
• Hydrogen bonds and van der Waals interactions are weak
attractions
ANIMATION: How atoms bond
To play movie you must be in Slide Show Mode
PC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE
2.4 REDOX REACTIONS
LEARNING OBJECTIVE:
• Distinguish between the terms oxidation and reduction,
and relate these processes to the transfer of energy
Redox Reactions
• Many energy conversions involve oxidation–reduction (redox)
reactions in which an electron (and its energy) is transferred
from one substance to another
• oxidation
• Chemical process in which an atom, ion, or molecule loses
one or more electrons
• reduction
• Chemical process in which an atom, ion, or molecule
gains one or more electrons
Redox Reactions (cont.)
• Redox reactions occur simultaneously; one substance
accepts electrons that are removed from the other
• The oxidizing agent accepts electrons and is reduced
• The reducing agent gives up electrons and is oxidized
• In cells, oxidation often involves removal of a hydrogen atom
(an electron plus a proton) from a covalent compound
• Reduction often involves the addition of a hydrogen atom
Key Concepts 2.4
• The energy of an electron is transferred in a redox reaction
2.5 WATER
LEARNING OBJECTIVE:
• Explain how hydrogen bonds between adjacent water
molecules govern many of the properties of water
Importance of Water
• About 70% of our total body weight is water
• Water (via photosynthesis) is the source of oxygen in air,
and hydrogen atoms used in many organic compounds
• Water is a solvent for biological reactions, and a reactant
or product in many chemical reactions
• Water is a principal environmental factor for organisms
• Water’s unique properties are essential to life
Effects of Water on an Organism
(a) Commonly known
as "water bears,"
tardigrades, such as
these members of the
genus Echiniscus, are
small animals (less
than 1.2 mm long) that
normally live in moist
habitats, such as thin
films of water on
mosses.
Fig. 2-12a, p. 37
(b) When subjected to desiccation
(dried out), tardigrades assume a
barrel-shaped form known as a
tun, remaining in this state,
motionless but alive, for as long
as 100 years. When rehydrated,
they assume their normal
appearance and activities.
10 μm
Fig. 2-12b, p. 37
Hydrogen Bonds
Between Water Molecules
• Water molecules are polar
• Hydrogen bonds form between the partial positive charge
(hydrogen) of one water molecule and the partial negative
charge (oxygen) of a neighboring water molecule
• Each water molecule can form hydrogen bonds with up to four
neighboring water molecules
Hydrogen Bonding of Water Molecules
Fig. 2-13, p. 38
Cohesion and Adhesion
• cohesion
• Tendency of water molecules to stick to one another, due
to hydrogen bonds among molecules
• Any force exerted on part of a column of water is
transmitted to the column as a whole
• Major mechanism of water movement in plants
• adhesion
• The ability of water to stick to other substances,
particularly those with charges on their surfaces
• Explains how water makes things wet
Capillary Action and Surface Tension
• capillary action
• The tendency of water to move in narrow tubes, even
against the force of gravity
• A combination of adhesive and cohesive forces
• surface tension
• Molecules at water’s surface crowd together, producing a
strong layer as they are pulled downward by the attraction
(cohesion) of water molecules beneath them
Capillary Action
• Adhesion between
water and glass in a
narrow tube pulls other
water up by cohesion
• In the wider tube,
adhesion is not strong
enough to overcome the
cohesion below
Fig. 2-14, p. 38
Surface Tension
• Water striders (Gerris) supported by surface tension of water
Fig. 2-15, p. 38
Interactions with Water
• Because water molecules are polar, water dissolves many
kinds of substances, and excludes others
• Hydrophilic (“water-loving”) substances interact easily with
water (polar and ionic compounds)
• Hydrophobic (“water-fearing”) substances are not soluble in
water (nonpolar molecules)
• Hydrophobic interactions occur between groups of
nonpolar molecules, which cluster together in water
Water and Temperature
• Water exists in three states, which differ in degree of
hydrogen bonding: gas (vapor), liquid, and ice (crystalline)
• Adding heat energy makes molecules move faster (increases
kinetic energy) and breaks hydrogen bonds
• heat
• The total kinetic energy in a sample of a substance
• temperature
• The average kinetic energy of the particles
Water and Temperature (cont.)
• Much of the heat energy added is used to break hydrogen
bonds – less energy is available to speed the movement of
water molecules (increasing temperature)
• heat of vaporization
• Amount of heat energy required to change 1 g of a
substance from liquid phase to vapor phase
• calorie (cal)
• Amount of heat energy – equivalent to 4.184 joules (J) –
required to raise 1 g of water 1 degree Celsius (C)
Water and Temperature (cont.)
• evaporative cooling
• As water is heated, some molecules move faster than
others and are more likely to evaporate, taking heat with
them and lowering the temperature of the water
• Humans dissipate excess heat as sweat evaporates
• specific heat
• A large amount of energy is required to raise the
temperature of water (1 cal/g of water per degree Celsius)
• Maintains constant environmental temperatures
Water and Temperature (cont.)
• Hydrogen bonding causes ice to have unique properties with
important environmental consequences
• Liquid water expands as it freezes, making ice about 10%
less dense than liquid water – ice floats on denser cold water
• Ice insulates liquid water below it, retarding freezing and
permitting organisms to survive without freezing
Three Forms of Water
212°F 100°C
(a) Steam becoming water vapor (gas)
50°C
(b) Water (liquid)
32°F 0°C
(c) Ice (solid)
Fig. 2-16, p. 39
ANIMATION: Structure of water
To play movie you must be in Slide Show Mode
PC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE
Key Concepts 2.5
• Water molecules are polar, having regions of partial positive
and partial negative charge that permit them to form hydrogen
bonds with one another and with other charged substances
ANIMATION: Dehydration synthesis and
hydrolysis
To play movie you must be in Slide Show Mode
PC Users: Please wait for content to load, then click to play
Mac Users: CLICK HERE
3D Animation: Dissolution
2.6 ACIDS, BASES, AND SALTS
LEARNING OBJECTIVES:
• Contrast acids and bases, and discuss their properties
• Convert the hydrogen ion concentration (moles per liter)
of a solution to a pH value and describe how buffers help
minimize changes in pH
• Describe the composition of a salt and explain why salts
are important in organisms
Ionization of Water
• In pure water, a small number of water molecules dissociate
into hydrogen ions (H+) and hydroxide ions (OH−)
HOH ↔ H+ + OH−
• The concentrations of hydrogen ions and hydroxide ions in
pure water are exactly equal
• Such a solution is said to be neutral – neither acidic nor basic
Acids and Bases
• acid
• Substance that dissociates in solution to yield hydrogen
ions (H+) and anion; a proton donor
• Acid → H+ + anion
• base
• Substance that dissociates in solution to yield a hydroxide
ion (OH−) and a cation; a proton acceptor
• NaOH → Na+ + OH• OH- + H+ → H2O
Acids and Bases (cont.)
• Some bases do not dissociate to yield hydroxide ions directly
• Example: Ammonia (NH3) acts as a base by accepting a
proton from water, producing an ammonium ion (NH4+) and
releasing a hydroxide ion:
• NH3 + H2O → NH4+ + OH−
pH
• A solution’s acidity is expressed in terms of pH
• pH
• The negative logarithm (base 10) of the hydrogen ion
concentration [H+] (expressed in moles per liter):
• pH = −log10[H+]
• The negative logarithm corresponds to a positive pH value
• Pure water has a hydrogen ion concentration of
0.0000001 (10—7 mol/L)
• Logarithm = −7; pH is 7
pH of Solutions
• neutral solution (pH 7)
• Equal concentrations of hydrogen ions and hydroxide ions
(concentration of each is 10−7 mol/L)
• acidic solution (pH <7)
• Hydrogen ion concentration is higher than hydroxide ion
concentration
• basic solution (pH >7)
• Hydrogen ion concentration is lower than hydroxide ion
concentration
Calculating pH Values and
Hydroxide Ion Concentrations
Table 2-2, p. 41
pH Values
• pH of most plant and
animal cells (and their
environment) ranges
from around 7.2 to 7.4
pH scale
0
Battery acid 0.0
1
Hydrochloric acid 0.8
Stomach acid 1.0
2
Stomach gastric juice 2.0
Increasing 3
acidity
Vinegar 3.0
4
5
Neutrality
Beer 4.5
Black coffee 5.0
6
Rainwater 6.25
Cow milk 6.5
7
Distilled water 7.0
Blood 7.4
8
Seawater 8.0
9
Bleach 9.0
10
Mono Lake, California 9.9
Increasing 11
alkalinity
Household ammonia 11.5
12
13
Oven cleaner 13.0
14
Lye 14.0
Fig. 2-17, p. 41
Animation: The pH scale
Buffers Minimize pH Change
• Homeostatic mechanisms maintain appropriate pH values
• Example: pH of human blood is about 7.4 and must be
maintained within very narrow limits
• buffer
• Substance that resists changes in pH when an acid or
base is added
• A buffering system includes a weak acid or a weak base
A Buffering System
• In blood, carbon dioxide reacts with water to form carbonic
acid, a weak acid that dissociates to yield H+ and bicarbonate:
CO2 + H2O ↔ H2CO3 ↔ H+ + HCO3–
• Addition of excess hydrogen ions shifts the system to the left,
as H+ combine with bicarbonate ions to form carbonic acid
• Addition of hydroxide ions shifts the system to the right
Formation of Salts
• When an acid and a base are mixed in water, anions from the
acid and cations from the base combine to form a salt
• salt
• Compound in which the hydrogen ion of an acid is
replaced by some other cation
• Example: Sodium chloride (NaCl) is a salt in which the
H+ of HCl has been replaced by the cation Na+
• HCl + NaOH → H2O + NaCl
Salts (cont.)
• When a salt, acid, or base is dissolved in water, its
dissociated ions (electrolytes) can conduct an electric current
• Animals and plants contain a variety of dissolved salts
(important mineral ions) essential for fluid balance and acid–
base balance
• Homeostatic mechanisms maintain concentrations and
relative amounts of various cations and anions
Key Concepts 2.6
• Acids are hydrogen ion donors; bases are hydrogen ion
acceptors
• The pH scale is a convenient measure of the hydrogen ion
concentration of a solution