Download Ionic bonds

Chapter 2:
The Chemical Context of Life
Reminder from Chapter 1:
Organisms are natural systems to which basic
concepts of chemistry and physics apply. One of
the main themes of biology is the organization
of life on a hierarchy of structural levels, with
additional properties emerging at each
successive level. In this chapter, we will see how
the theme of emergent properties applies to the
lowest level of biological organization.
Matter consists of chemical elements in pure form and in
combinations called compounds
•All life is composed of matter.
Matter = anything that takes up space and has mass.
Matter exists in diverse forms, each with its own
characteristics.
•All matter is made up of chemical elements.
•Element = a substance that cannot be broken down to
other substances by chemical reactions.
•Compound = combination of two or more elements in a
fixed ratio. Compounds have emergent properties - i.e.
characteristics different from those of its elements.
Life requires about 25 elements
•Six elements make up 97.6% of living matter
(carbon, oxygen, hydrogen, nitrogen,
phosphorous, sulfur)( Table 2.1). These elements
form stable covalent bonds.
•Trace elements = those required by an organism
in minute quantities.
•Deficiencies in trace elements can cause illness.
e.g. lack of iron and iodine cause anemia and
goiter, respectively.
Atomic structure determines the
behavior of an element
•Atoms = fundamental unit of matter. Smallest possible
amount of an element that retains that element’s
properties.
•Each element consists of a certain kind of atom, which
is different from the atoms of other elements.
•Atoms are composed of even smaller parts. Only three
kinds of subatomic particles are relevant from a
biological perspective: protons, neutrons, and electrons.
•Neutrons and protons are packed together tightly at the
center of an atom to form a nucleus. The electrons move about
this nucleus at almost the speed of light.
•Electrons (-) and protons (+) are electrically charged,
whereas the neutron is neutral. Protons give the nucleus a
positive charge, and it is the attraction between opposite
charges that keeps the rapidly moving electrons orbiting
around the nucleus.
•The unit of measurement for atomic particles is the dalton.
Neutrons and protons have a mass of 1 dalton each. Electron
mass is negligible.
•Atoms of the various elements vary in their number of subatomic
particles. All atoms of a particular element have the same number of
protons in their nuclei. This is their atomic number.
•Mass number = sum of protons and neutrons in the nucleus of an
atom.
•Atomic mass (weight) = since neutrons and protons have a mass
close to 1 dalton, the mass number tells us the approximate mass of
the whole atom.
•Isotopes = variant forms of elements. Have same number of protons
and electrons, but different number of neutrons. •The nucleus of 14C
is unstable and therefore radioactive.
•Radioactive isotopes = nucleus decays spontaneously giving off
particles and energy. Dangerous to life because it causes mutations in
DNA. However, radioactive isotopes can also be useful in biological
research and medicine as tracers. Living cells cannot distinguish
radioactive isotopes from nonradioactive atoms of the same
elements.
Energy levels
•Atoms are mostly empty space. When two atoms approach each
other during a chemical reaction, their nuclei do not come close
enough together to interact. Only electrons are directly involved in
the chemical reactions between atoms.
•An atom’s electrons vary in the amount of energy they possess (Fig
2.7).
•Energy = ability to do work.
•Potential energy = energy that matter stores because of its position
or location.
•Matter has a natural tendency to move to the lowest possible state
of potential energy. Electrons of an atom also have potential energy
because of their position in relation to the nucleus. The negatively
charged electrons are attracted to the positively charged nucleus.
The more distant the electrons are from the nucleus, the greater
their potential energy.
Energy levels (continued)
•Changes in the potential energy of electrons can only occur in
steps of fixed amounts. The different states of potential energy for
electrons in an atom are called energy levels or electron shells.
Electrons in first shell closest to nucleus have the lowest energy.
Electrons in the second shell have more energy, electrons in third
shell have more energy still, and so on.
•An electron can change its shell, but only by absorbing or losing an
amount of energy equal to the difference in potential energy
between the old shell and the new shell.
Electron orbitals
•We can never know the exact trajectory of an electron.
Instead, we describe the volume of space in which an
electron spends most of its time (Fig 2.9).
•Orbital = three-dimensional space where an electron is
found 90% of the time.
•No more than two electrons can occupy the same
orbital.
•First shell has a single spherical orbital and can hold
only 2 electrons. An atom with more electrons must use
higher shells.
•The second electron shell can hold 8 electrons, two in
each of four orbitals (1 spherical and 3 dumbbellshaped).
Electron configuration and chemical properties
•The chemical properties of an atom depend mostly on
the number of electrons in its outermost shell (Fig 2.8).
•valence electrons = electrons in outer most shell
•valence shell = outermost energy shell.
•Valence = an atom’s bonding capacity (# of electrons
needed to fill outer shell).
•Atoms with a complete valence shell are not reactive. All
other atoms are chemically reactive because they have
incomplete valence shells with unpaired electrons.
Atoms combine by chemical bonding to form molecules
•When atoms with incomplete outer shells react, each
atom gives up or acquires electrons so that partners end
up with completed outer shells.
•Atoms do this by either sharing (covalent bonds) or
transferring outer electrons (ionic bonds) resulting in
chemical bonds.
•The strongest chemical bonds are covalent bonds and
ionic bonds.
Covalent bonds
•Covalent bond = two atoms sharing one or more pairs of outer shell
electrons ( Fig 2.10 and 11).
•Molecule = two or more atoms held together by covalent bonds.
•The number of single covalent bonds an atom can form is equal
to the number of additional electrons needed to fill its outer shell
(i.e. it's valence).
•Double bond = sharing of 2 pairs of electrons. Stronger than
single bonds. •Atoms in a covalently bonded molecule are
constantly in a tug-of-war for the electrons of their covalent
bonds.
•Electronegativity = an atom’s attraction for the shared electrons of the bond. The more
electronegative an atom, the more strongly it pulls electrons towards itself.
•Nonpolar covalent bonds = electrons shared equally between the atoms of equal
electronegativity (H2, O2, CH4 ). •Water is made up of 2 kinds of atoms with differing
electronegativity (O>H). Oxygen attracts electrons more strongly than hydrogen.
•Polar covalent bond = chemical bond in which shared electrons are pulled closer to the
more electronegative atom, making it partially negative and the other atom partially
positive (Refer to Fig 2.12). •H2O, even though is neutral overall, has a slightly
negative pole and two slightly positive poles, making it a polar molecule.
Ionic bonds
•Refer to Fig 2.13
Ionic bonds = attractions between ions of opposite charge
(e.g. table salt, NaCl). Much weaker than covalent bonds.
•When atoms of chlorine and sodium collide, chlorine atom
strips sodium’s outer electron away. This results in sodium
having a positive charge and chlorine having a negative
charge. Two ions of opposite charge attract each other;
when the attraction holds them together, it is called an ionic
bond.
Ion = atom or molecule with an electrical charge resulting
from a gain or loss of one or more electrons.
anion = ion with negative charge; cation = ion with a
positive charge
•NaCl is a type of salt. Salts are ionic compounds that often
exist as crystals in nature.
Weak chemical bonds play important roles in the
chemistry of life
•Weak bonds, unlike covalent bonds, allow interactions
between molecules to be brief; molecules may come
together, change in some way and then separate.
•The most important weak bond in living matter is the
hydrogen bond.
Hydrogen bond = occurs when a hydrogen atom
covalently bonded to one electronegative atom is also
attracted by another electronegative atom. In living cells,
the electronegative partner involved are usually oxygen
and nitrogen atoms. (refer to Fig 2.15)
•Hydrogen bonds, ionic bonds, and other weak bonds,
form between and within molecules. Although these
bonds are individually weak, their cumulative effect can
re-enforce the 3-D shape of a large molecule.
A molecules biological function is related to
its shape
•Molecular shape is important in biology because it is the
basis for how most molecules of life recognize and
respond to one another.
•Recognition and binding of neurotransmitters to cell
surface receptors in synapses of brain cells is basis on
intercellular communication in the nervous system (Fig
2.16 and 17).
Chemical reactions make and break chemical bonds
•Living matter is not static. There is constant flux, as new molecules
are being built and others are being broken down. The goal of
biochemistry is not simply to catalogue the molecules that make up
the living world, but to understand how these molecules are
transformed into others in biochemical pathways. These
transformations always involve chemical reactions.
•In a chemical reaction, reactants interact, atoms rearrange, and
products result.
•Matter is conserved in a chemical reaction. Reactions cannot create
nor destroy matter but can only rearrange it.
•Living cells carry out thousands of chemical reactions that rearrange
matter in significant ways.
•Some chemical reactions go to completion, others are reversible.
•Chemical equilibrium = point at which rate of forward reaction equals
that of reverse reaction
Chapter 3:
Water and the Fitness of the Environment
The polarity of water molecules results in hydrogen bonding
•Oxygen is more electronegative than hydrogen. Consequently,
the electrons of the polar bonds spend more time near the
oxygen atom. This makes water a polar molecule.
•The unique (emergent) properties of water arises from
attractions among these polar molecules.
•Each water molecule can hydrogen bond (H-bond) to a max of
four neighbors. •H-bond = electrostatic attraction between a
hydrogen in a polar bond to an electronegative atom of
another molecule.
•The charged regions of a polar molecule are attracted to
opposite charges of neighboring polar or ionic molecules.
Organisms depend on the cohesion of water molecules
•Water molecules stick together as a result of H-bonding.
H-bonds form, break, and reform very frequently. At any
given time, a substantial portion of all molecules are
bonded to their neighbors, giving water more structure
than most liquids.
•Cohesion = tendency of molecules to stick together.
Much stronger for water than for other liquids. Important
in water transport in plants.
•Adhesion = the clinging of one substance to another.
Also important in water transport in plants.
•Surface tension = a measure of how difficult it is to
stretch or break the surface of a liquid. Higher for water
than for most liquids.
Water contributes to earth's habitability by
moderating temperatures
•Water stabilizes air temperature by absorbing heat from
air that is warmer and releasing the stored heat to the air
that is cooler.
•Water can store a lot of energy (heat) with only a slight
increase in its own temperature.
•Heat = measure of the total quantity of kinetic energy
(energy of motion) due to molecular motion in a body of
matter.
•Temperature = measures intensity of heat due to the
average kinetic energy of the molecules.
•Calorie = amount of heat energy it takes to raise the
temp of 1 gram of water by 1°C (Food calorie = 1000
calories)
Water's high specific heat
•Specific heat = amount of heat that must be absorbed or
lost for 1 gram of that substance to change temperature
by 1 °C.
•Compared to most substances, water has an unusually
high specific heat (10x that of Fe) This is due to Hbonding. •A calorie of heat causes a relatively small change in
temperature of water because much of that heat energy is used to
disrupt H-bonds before water molecules can begin to move faster.
•Conversely, when the temperature of water drops
slightly, many additional H-bonds form, releasing a lot of
heat energy.
•Water buffers against extreme changes in temperature.
Evaporative cooling
•Molecules in a liquid stay close together because they
are attracted to one another. Molecules moving fast
enough to overcome these attractions can depart from
the liquid and enter into gas state.
•Heat of Vaporization = quantity of heat a liquid must
absorb for 1 gram of it to be converted from liquid to
gaseous state.
•Water's high heat of vaporization helps moderate earth's
climate. A considerable amount of solar heat absorbed by
tropical seas is consumed during evaporation of surface
water. Thus, as moist tropical air circulates poleward, it
releases heat as it condenses to form rain.
•Evaporative cooling also helps moderate temperature in
lakes and ponds, and prevents terrestrial organisms from
overheating.
Ice floats
•Water is one of the few substances that is less dense as
a solid than as a liquid. (Fig 3.5)
•If ice sank, eventually all bodies of water would freeze
solid since floating ice insulates liquid water below.
•Ice floats because as temperature decreases, there is
less energy to break H-bonds, so eventually all water
molecules are H-bonded to one another resulting in a
crystal lattice structure in which water molecules are less
densely packed.
Water as the solvent of life
•Water dissolves more solutes than any other liquid called “Universal Solvent.”
• Solution = liquid that is a homogenous mix of 2 or more
substances
•Solvent = dissolving agent
•Solute = substance that is dissolved
•Aqueous solution = one in which water is solvent•The
versatility of water as a solvent is based on its polarity.
•Ions and polar water molecules have a mutual affinity
through electrical attractions. E.g. Cl- attracted and
surrounded by positive part of water and Na+ attracted
and surrounded by negative part of water molecules - the
resulting sphere of water molecules around each ion is
called the hydration shell.
•A compound does not have to be an ion to be dissolved
by water. Polar compounds are also water-soluble ( E.g.
proteins, carbohydrates, nucleic acids).(refer to Fig 3.8:
hydration of soluble protein) and become surrounded by
a hydration shell as well.
•Hydrophilic = any substance with an affinity for water
(ions, polar molecules), even if that substance does not
dissolve (E.g. cellulose)
•Hydrophobic = any substance that neither dissolves nor
has an affinity for water (nonpolar). E.g fats, waxes,
etc...
Solute concentrations in Aqueous solutions
Knowing concentrations is important in biology. This
allows the combinations of substances in fixed ratios to
make chemical solutions.
•Molarity = Number of moles/Liter
•A mole (mol) is a quantity = 6.02 x 1023 (Avogadro’s
Number) molecules
•Molecular weight(mass) = sum of the weight of all the
atoms in a molecule.
• Molar Mass = molecular weight(mass) (in grams) of a
particular substance per 1 mole of that substance
•e.g. 1 mole of sucrose has 6.02 x 1023 molecules and
weighs 342 g. 1 mole of ethanol also has 6.02 x 10-23
molecules, but weighs only 46 g.
Acids and Bases
•In pure water, [H+] = [OH-] = 10-7
•Acid = substance that increases the [H+] of a solution. A
strong acid, such hydrochloric acid, dissociates completely
when mixed with water.
Base = A substance that reduces the [H+] in a solution. A
strong base, such as sodium hydroxide, dissociates completely
when mixed with water.
Some bases, such as ammonia, reduce [H+] directly by
accepting hydrogen ions
Weak acids and bases do not dissociate completely.
Ex: Carbonic Acid
The pH scale
•pH = -log [H+]
•For a neutral solution [H+] is 10-7 M, therefore -log 10-7
= -(-7) = 7
•Each pH unit represents a 10 fold change in
concentration.
Buffers
•Because biological systems are very sensitive to pH,
they need to minimize changes in pH. They do this with
buffers.
•Buffers = compounds that resist changes to their own pH
when acids or bases are introduced.
•A buffer works by accepting hydrogen ions from the
solution when they are in excess and donating hydrogen
ions to the solution when they have been depleted. Most
buffers are weak acids or bases. (acid-base pairs)
Examples of physiological buffers: carbonate and
phosphate buffers
Chapter 4: Carbon and the molecular diversity of life
Biological diversity reflects molecular diversity. Of
all chemical elements, carbon is unparalleled in
its ability to form molecules that are large,
complex, and diverse. This chapter focuses on
the concepts of molecular architecture that
highlight carbon's importance to life.
Carbon atoms are the most versatile building blocks of
molecules
•Valence of the major elements (Fig 4.4).
•The chemical characteristics of an atom depend on its
valence electrons, by determining the number of bonds
an atom will form with other atoms.
The tetravalence of carbon explains its versatility at
making large, diverse, complex molecules possible
Variation in carbon skeletons contributes to diversity of
organic molecules
•Variation in carbon skeletons is important source of
molecular complexity and diversity.
•Hydrocarbons = organic molecules consisting of only
carbon and hydrogen(Fig 4.5).
•Major component of petroleum, a fossil fuel.
•Nonpolar, therefore hydrophobic (Fig 4.6)
•can vary in length•Position of double bonds can vary
•may be branched or unbranched
•can be arranged in rings
•Isomers = compounds with same molecular formula but
different structures and hence properties.
•Types of isomers (Fig 4.7):
•1. Structural isomers - differ in covalent arrangement of
their atoms•contributes significantly to their variation e.g.
18 types for C8H18, and 366,319 types for C20H42
•2. Geometric isomers - have same covalent relationships,
but differ in spatial arrangements (cis- and trans-) •arise
from inflexibility of double bonds •have distinct biological
activities (ex: retinal)
•3. Enantiomers - molecules that are mirror images of
each other •asymmetric carbon attached to 4 different
atoms or groups of atoms. •thalidomide was a mix of 2
enantiomers: one with sedative properties, the other
caused birth defects •L-Dopa vs D-Dopa is another
example (fig 4.8).
Functional groups
•The distinctive properties of organic molecules depend
not only on arrangement of carbon skeleton, but also on
other molecular components attached to that skeleton.
•Functional groups = groups of atoms attached to carbon
skeleton which are commonly involved in chemical
reactions.
•All functional groups studied here are hydrophilic.
•Hydroxyl group (-OH) •compounds containing -OH are alcohols.
Eg. ethanol, sugars, phenol •-OH group is polar. Therefore such
compounds dissolve in water (sugars)
•Carbonyl group (C=O) •aldehyde = carbonyl group on end carbon
of chain •ketone =- carbonyl group attached to internal carbon
•Variation in locations of functional groups along carbon chain is a
source of variation.
•Carboxyl group (-COOH) •compounds containing carboxyl group
are called carboxylic acids. Eg. formic acid, acetic acid, amino acids
•they are weak acids because they are a source of H+ ions
•Amino group (-NH2) •compounds containing amino group are
called amines. Eg. amino acids Note: amino acids are both amines
and carboxylic acids. •Amino group can act as a base.
•Sulfhydryl group (-SH) •organic compounds containing sulfhydryls
are called thiols •important in stabilizing protein structure.
•Phosphate group •Phosphate is an anion formed by dissociation of
an inorganic acid called phosphoric acid (H3PO4) •Functions in
energy transfer between organic molecules