Download Chemistry for BIOS 302

Chemistry for BIOS 302
Life is Chemistry
• We believe that life is just a large set of
chemical reactions. Living organisms obey
the same physical and chemical principles
that non-living things do.
– The scientific world has long since rejected
the theory of vitalism: the idea that a nonmaterial, spiritual force is needed to
convert non-living into living.
• Chemistry of carbon compounds = organic chemistry
• In aqueous (watery) environment, in a very narrow
temperature range.
• The chemical reactions of life are very complex and
intertwined.
• Mostly involve large molecules (macromolecules) composed of
many subunits.
Elements and Atoms
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All matter is composed of atoms.
Elements such as carbon and oxygen are a group of atoms of the same type.
There are 92 naturally occurring elements, plus about 25 artificially-created elements.
Living things are mainly composed of the elements carbon, hydrogen, oxygen, and
nitrogen. Almost as important are sulfur (found in 2 of the 20 amino acids) and
phosphorus (an essential component of all nucleic acids). Another dozen or so elements
are also used: iron, magnesium, calcium, sodium, potassium, chlorine, to name a few.
Atoms
• Atoms have 3 components:
protons, neutrons, and
electrons
• protons and neutrons are in
the nucleus
• Electrons circle around the
nucleus
• NOTE! This nucleus is NOT
the same as the cell’s
nucleus
Properties of Protons, Neutrons, and Electrons
• Protons have a mass of 1 dalton (or 1 atomic mass unit).
Protons also have an electrical charge of +1: protons are
positively charged.
– Daltons are defined as 1/12 the mass of a carbon-12 atom. This
averages out the small weight difference between protons and
neutrons.
– Daltons are related to grams by Avogadro’s number: 6 x 1023 protons
have a mass of 1 gram.
• Neutrons also have a mass of 1 dalton. Most of the mass of
an atom is in its protons and neutrons. Neutrons are neutral
(not electrically charged).
• Electrons weigh about 1/1800 of a dalton, very light. They
have a -1 charge (negatively charged).
Atomic Number and Weight
• The atomic number of an atom is
the number of protons it has. The
atomic number defines the basic
identity of the atom: what element
it is.
– For example. All hydrogen atoms have 1
proton (atomic number = 1), and all
carbon atoms have 6 protons (atomic
number = 6). All gold atoms have 79
protons; uranium atoms have 92 protons.
• The atomic weight of an atom is
the number of protons plus the
number of neutrons.
– For example, most carbon atoms have
6 protons and 6 neutrons, so the
atomic weight is 12.
– Electrons are very light, so they don’t add
significantly to the atomic weight.
Isotopes
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Isotopes are atoms with the same atomic number (same element) but
different atomic weights: that is, isotopes have the same number of protons
but different numbers of neutrons. Most elements have several naturally
occurring isotopes.
Different isotopes of the same element all have the same chemical properties.
Chemical properties are determined by the number of protons only.
Heavier isotopes (too many neutrons) are usually radioactive: the atomic
nucleus is unstable and spontaneously changes into a different configuration
of protons and neutrons. In the process, particles or electromagnetic
radiation are released.
Different isotopes are written like 12C or 14C.
Molecular Weight, Moles and Molarity
• The molecular weight of a chemical compound is the sum of the atomic
weights of the atoms in the compound.
– For example, carbon dioxide (CO2) has 1 carbon (12 daltons) plus 2 oxygens (2 * 16
= 32 daltons), so it has a molecular weight of 44 daltons.
– Note that atomic weights are usually not whole numbers: the fraction comes from
averaging the atomic weights of all the isotopes present in the Earth’s crust.
• Molecular weight can also be thought of as grams per mole (g/mol). A
mole of carbon dioxide has a mass of 44 grams.
• How fast chemical reactions occur depends on the concentration of the
reactants. Most biological chemistry occurs in aqueous solution, so we
care about concentration in terms of moles per liter (molar, symbolized as
M).
– A 1 M solution contains 1 mole of the solute dissolved in a total volume of 1 liter.
• For example, what is the concentration of 16 g of CO2 in 0.5 L?
– 16 g / 44 g/mol = 0.36 moles.
– 0.36 mol / 0.5 L = 0.72 mol/L = 0.72 M
Electron Shells
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Most chemistry is caused by the
electrons of the outermost shell:
atoms have a “desire” to have a
full outer (valence) shell with 8
electrons in it, so they share
electrons to accomplish this.
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Atoms are usually neutral in electrical charge,
which means that they have the same number
of electrons (- charge) as protons (+ charge).
Electrons circle the nucleus at defined positions
called shells.
The innermost shell of every atom holds 2
electrons.
The next two shells hold up to 8 electrons. The
first three shells thus hold 2 + 8 + 8 = 18
electrons, which corresponds to argon.
– Things get more complex above this point,
but nearly all the biologically important
elements use only the first 3 electron shells.
– Big exception: metals such as iron, zinc,
manganese, copper.
Electrons fill in from the nucleus out, so the
inner shells are always full (except hydrogen,
which only has 1 electron).
In most atoms, the outermost shell (the valence
shell) is not full.
Ions
• The number of electrons in an atom is usually
the same as the number of protons. Since
electrons have a -1 charge and protons have
a +1 charge, atoms are electrically neutral.
• Ions are atoms where the number of
electrons is different from the number of
protons.
• Ions have an electrical charge: positive charge
if more protons than electrons, and negative
charge if more electrons than protons.
• Cations are positively charged: Na+ and Mg2+
are examples.
• Anions are negatively charged: Cl- for
example.
– The prefix “an-” means “not” or “negative”:
anonymous = no name; anarchy = no government;
anhydrous = no water.
For example, sodium atoms
have 11 protons. Neutral
sodium atoms also have 11
electrons, but sodium ions
have only 10 electrons: 11
protons plus 10 electrons,
giving a total charge of +1.
Chemical Bonds
• Atoms can combine with each other to form molecules. Very few elements
exist naturally in an uncombined state: mostly they are joined into molecules.
• A molecule is a defined number of atoms grouped into a defined spatial
relationship. For example, water, H2O, is 2 hydrogen atoms connected to an
oxygen atom. The oxygen is in the middle, and the hydrogens are attached at
an angle to it. The bonds have specific lengths and angles relative to each
other.
• A large group of the same molecule is called a compound (just as a large group
of the same atom is called an element).
• Molecules are held together by chemical bonds.
• Chemical bonds are the result of 2 forces:
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1. The octet rule, which means that atoms want to have 8 electrons in their
outer shell (2 in the case of hydrogen).
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2. The attraction between atoms of opposite electrical charge.
• The three main types of chemical bond are: ionic bond, covalent bond, and
hydrogen bond.
Outer Shell Electrons
• Hydrogen has 1
electron in its outer
shell, and needs a total
of 2 electrons to be
stable.
• All other biologically
relevant atoms need 8
electrons in their outer
shell
– Thus, carbon has 4
electrons and wants 4
more, nitrogen has 5
electrons and wants 3
more, and oxygen has 6
and wants 2 more.
Ionic Bonds
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In an ionic bond, one atom gives an electron
to another atom. This makes both atoms
ions, and they are held together because
their opposite charges attract each other.
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In sodium chloride (table salt), sodium starts
out with 1 electron in its outer shell. The
next shell down has 8 electrons, so by giving
1 electron away, the sodium atom gets a full
outer shell. It then has a +1 charge.
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Chlorine starts out with 7 electrons in its
outer shell. By gaining one more electron, it
gets 8 in the outer shell, and a -1 charge.
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The + charged sodium and the – charged
chlorine attract each other, and they pack
together in salt crystals.
Ionic Bonds in Biology
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Ionic bonds are very strong in many
environments. But in an aqueous
environment, they are greatly weakened by
interactions with water molecules, and are
much weaker than covalent bonds.
– However, unlike covalent bonds, ionic bonds
don’t have a fixed geometry. They are equally
strong regardless of how the 2 atoms are
oriented.
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In biology, ionic bonds are mainly used to
stabilize the three-dimensional shapes of
molecules and to help molecules interact
with each other.
We usually speak of electrostatic
interactions between charged molecules
(both attraction and repulsion), rather than
use the term “ionic bond”.
– Hydrogen bonds are another example of
electrostatic interactions that we will discuss
soon.
Covalent Bonds
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Covalent bonds occur when 2 atoms
share a pair of electrons. The
electrons spend part of their time
with both atoms, so the octet rule is
satisfied sufficiently.
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A molecule of hydrogen gas, H2, has 2
hydrogen atoms. Each atom provides
1 electron, so in the bond each atom
shares 2, a complete shell for
hydrogen.
– The bond is symbolized as a line
connecting the 2 H’s: H-H
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Covalent bonds are what hold most
biological molecules together.
Covalent Bond Orientation
• The 8 electrons of the outer
electron shell arrange themselves
into 4 groups of two, which stay as
far apart as possible. This produces
a tetrahedron shape (triangular
pyramid). Covalently bonded
molecules mostly fit this pattern.
• In water (H2O), the oxygen has 6
electrons in its outer shell, and it
shares one with each of the 2
hydrogens, giving 8 shared
electrons for oxygen and 2 for each
hydrogen. The hydrogens form 2
corners of the tetrahedron, and the
two pairs of unbonded electrons
form the other 2 corners.
Single, Double, and Triple Bonds
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In a single bond, a pair of electrons (2
electrons) is shared. H2 gas and water
are examples of this. Most covalent
bonds are single bonds.
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In a double bond, 2 pairs of electrons
(total of 4 electrons shared) are
shared. In oxygen gas (O2), each
atom has 6 electrons, a total of 12.
Each atom contributes 2 electrons to
the bond, so 4 are shared. Thus each
atom has 4 unshared and 4 shared
electrons, satisfying the octet rule.
Carbon dioxide is another example.
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In a triple bond, 3 pairs of electrons
(6 electrons) are shared. There are
very few triple bonds in biological
molecules: nitrogen gas and cyanide
(CN) are examples.
Double Bonds Give Planar Configuration
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Double bonds are shorter and stronger than
single bonds.
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Chemical groups attached to single bonds can
rotate freely around the bond axis.
Groups attached to double bonds are forced
into a planar configuration and can’t rotate.
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Triple bonds are even shorter and stronger.
This gives rise to cis and trans isomers for
compounds with double bonds. Cis isomers
have both side groups coming out of the same
side of the double bond; trans isomers have
the side groups coming out opposite sides.
Here, cis-2-butene has different chemical
properties than trans-2-butene.
This is important with fatty acids, where
artificially created trans isomers have been
shown to be unhealthy relative to the naturally
occurring cis isomers.
Carbon dioxide is non-polar because the two
double bonds between C and O create a linear
molecule. The partial negative charges on the
oxygens cancel out.
Polar Covalent Bonds
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Sometimes the electrons in a covalent bond aren’t shared equally, because one atom attracts electrons more
strongly than the other.
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Formally: the atom that attracts electrons more strongly has a greater electronegatvity.
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When this happens, the electrons spend more time with one atom, and that atom becomes slightly negatively
charged. The other atom becomes slightly positively charged. This is a polar covalent bond, because the
atoms form positive and negative poles.
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Rule: Carbon and hydrogen share electrons equally. Oxygen and nitrogen also share equally. But, oxygen and
nitrogen attract electrons more strongly than carbon or hydrogen.
– That is, oxygen and nitrogen are more electronegative than carbon and hydrogen.
– Or: in a covalent bond, an oxygen or nitrogen will have a slightly negative charge, and carbon or hydrogen
will have a slightly positive charge.
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Water is a polar compound, because the oxygen is slightly negative and the hydrogens slightly positive.
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Note that the total charge on the molecule is balanced, same number of electrons as protons, but within the
molecule the charges are slightly separated.
Bonds where the electrons are shared equally are called non-polar. The C-H bond is non-polar.
Hydrogen Bonds
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The slight + and – charges in polar bonds
attract each other. In biological
molecules, it is common for the partial +
charge on a hydrogen to attract the partial
– charge on a nearby oxygen or nitrogen.
This attraction is called a hydrogen bond.
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Hydrogen bonds are very weak compared
to covalent bonds, but large numbers of
them can add up to a strong bond. The
strands of DNA are held together by
hydrogen bonds.
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Hydrogen bonds also form between
different parts of the same molecule, and
between water and other molecules.
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NaCl dissolves as H-bonds form between
water molecules and the Na+ and Cl- ions.
Water molecules surround each ion
completely, separating the ions from the
salt crystal and dispersing them
throughout the water.
More on Bonds
• Covalent bonds are by far the strongest bonds for biological molecules.
• Ionic bonds are really just extreme cases of polar bonds: the sharing is
so unequal that one atom never gets the electron.
• Hydrogen bonds are the most common form of polar bonds.
• Van der Waals forces are very weak interactions that cause atoms to
weakly attract each other in the absence of all other forces.
Water
• Water, H2O, is a polar compound.
The 2 hydrogens are held at an angle
to each other, and so the oxygen end
of the molecule is partially negative
and the hydrogen end is partially
positive.
• Water forms many hydrogen bonds
with other water molecules and with
other polar substances. This causes
cohesion: water molecules sticking
together. Also adhesion: water
sticking to other (polar) things.
Hydrophilic and Hydrophobic
• Polar (and ionic) substances dissolve in water,
because water forms hydrogen bonds with the
polar molecules. Thus, polar substances are
called hydrophilic or “water-loving”.
• Non-polar substances don’t dissolve in water
because they can’t form hydrogen bonds, so they
are called hydrophobic or ‘water-fearing”. Oils
and fats are examples of non-polar substances.
– Non-polar substances dissolve in non-polar
solvents. Hexane and toluene are examples.
• “Like dissolves like”.
• Lipids are the main hydrophobic molecules in the
cell. Membranes are composed of lipids. The
inside and outside of the cell are polar, aqueous
environments, but they are separated by a
hydrophobic barrier that prevents them from
mixing freely.
Hydrophobic molecules
tend to clump together, to
minimize their interactions
with water molecules. This
is a major driving force in
the assembly of
membranes and in protein
folding.
Acids and Bases
• Water molecules tend to dissociate into H+
and OH- ions. The H+ ion is just a proton.
– It’s an equilibrium: water molecules are
constantly dissociating and associating.
• In practice, free protons are not found in
water. Instead, a proton combines with
another water molecule, generating H3O+
(hydronium ion) and OH- (hydroxyl ion).
– In pure water, the concentrations of
hydronium and hydroxyl are each 10-7 M. This
is a pH of 7.
• Acids produce H+ ions when dissolved in
water. For example, HCl (hydrochloric acid).
• Bases absorb H+ ions when dissolved in
water. A substance that releases OH- ions is a
base because the OH- combine with the H+ in
the water.
• Water is neutral, neither acid nor base,
because it always has equal numbers of H+
ions and OH- ions.
Strong vs. Weak
• Strong acids and bases completely
dissociate as soon as they are dissolved
in water. Examples: HCl and NaOH.
• Weak acids and bases do not completely
dissociate: some molecules retain the H+
or OH-. The percentage that are
dissociated depends on the dissociation
constant of the acid or base.
• Most biological acids and bases are
weak.
• Biological acids contain the –COOH
group (carboxylic acid)
• Biological bases mostly contain the –NH2
group (amino)
– The amine group removes a H+
(proton) from water, generating OH-
Dissociation Constants
• The dissociation constant is an
equilibrium constant: it describes the
relative amounts of the different
molecules at equilibrium.
• We write the relevant reaction, then put
the concentrations of the products
multiplied together in the numerator,
and the concentration of the original
compound in the denominator.
– We ignore the concentration of water.
– For acids, we use the acid dissociation
constant, Ka. This is easily mistaken for the
association constant.
– For bases, we use the base dissociation
constant, Kb, which is exactly the same thing
as Ka except that [OH-] is used instead of [H+]
is used.
More Dissociation Constants
• Dissociation constants are often written
as pKa, the negative logarithm of the Ka.
This is the same as pH, which is the
negative log of the hydrogen ion
concentration.
• For a single H+ or OH-, the pKa is the pH
at which half the molecules are
dissociated.
pH Scale
• Acidity is measured on the pH scale, which
indicates the amount of H+ ions present. The
scale runs from 0 (very acidic) to 14 (very
basic).
• pH is defined as the negative logarithm of
the H+ concentration.
– Water, which is neutral, has a pH of 7 because its
H+ concentration is 10-7
• Acids have lower pHs: the hydrochloric acid
in your stomach has a pH of about 2. Eating
food stimulates your stomach to secrete
more acid. Antacids (like Tums or Rolaids)
neutralize some of this acid.
• Bases have a higher pH: the lye in oven
cleaner has a pH of about 14.
• Body fluids are slightly basic, pH 7.4
Buffers and Salts
• Too much acid or base is harmful. The body needs to protect
itself against large pH shifts. It uses buffers, pairs of weak
acids and weak bases to absorb excess H+ and OH- and keep
the body’s pH near neutral.
• The main buffer in the body is carbonic acid, which dissociates
into H+ and HCO3-. If H+ is added by an acid, it gets converted
into the neutral H2CO3. Similarly, excess OH- combines with
the H+, leaving the much less basic HCO3-. These opposing
reactions keep the pH at the proper level.
Buffers and Salts
Cl-
Na+
• Salts are ionic compounds that don’t release H+ or OH- when they dissolve. Thus,
sodium chloride is a salt because it dissolves to form Na+ and Cl- , while hydrogen
chloride (hydrochloric acid, HCl) is an acid because it dissolves to form H+ and Cl-.
• Acids usually have an associated salt. An example of this is glutamic acid, an
important component of proteins, and monosodium glutamate (MSG), which is
used as a flavor enhancer in food. They are the same thing chemically, except that
MSG has a sodium where glutamic acid has a hydrogen.
Simple Organic Chemistry
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Carbon forms 4 bonds, because it has only 4 electrons in its outer shell and needs 4
more to become stable. This means carbon can form many different compounds.
– Nitrogen forms 3 bonds, oxygen forms 2 bonds, hydrogen and halogens (F, Cl, Br, I) form 1
bond.
• Carbon bonds are quite stable at typical temperatures.
• When drawing organic molecules, the hydrogens are often left out
completely
• Also, the lines connecting carbons are drawn, but not the actual “C”s.
Other atoms are represented by their letter symbols.
More Organic Chemistry
• Since the 4 single bonds at a carbon atom are formed in a tetrahedron,
chains of carbons are often drawn as a squiggly line. Hydrogens are left off:
their existence is implied by the need for each C to have 4 bonds.
• Many ring-shaped molecules have alternating single and double bonds,
which are actually all identical, and thus can be drawn as a circle inside the
carbon ring. The simplest example is benzene. Molecules of this type are
called aromatic, and they have special properties.
– Hydrocarbons that don’t have aromatic rings are called aliphatic.
Functional Groups
• Many organic compounds have functional
groups attached to them. Functional
groups are groups of atoms that have
specific functions and properties.
• The –OH group is called hydroxyl, and
molecules with it are called alcohols, with
names that end in –ol.
• The C=O group is called carbonyl, and
molecules containing it are aldehydes or
ketones.
– A carbonyl group that has another carbon
attached to one side is called an acyl group.
• Also important: -CH3 (methyl group)
• Also –PO4 (phosphate group). At
biological pHs, phosphates carry a -2
charge, or a -1 charge if they are
connected to 2 other groups.
Stereoisomers
• If a compound contains a carbon atom is attached
to 4 different groups, there will be a left handed
version and a right handed version, mirror images
of each other.
– These molecules are called chiral molecules, and the
carbon atom is called an asymmetric carbon atom or a
chiral center.
•
These are stereoisomers, molecules with the same
chemical and structural formula but which are not
superimposable.
– Stereoisomers don’t exist for atoms that join with 3 or
fewer others, or with carbons that have 2 or more identical
side groups.
• It is possible for a molecule to have several chiral
centers, which give rise to a whole series of
stereoisomers.
– Stereoisomers that have all chiral centers in the opposite
configuration as each other are called enantiomers.
– Stereoisomers that are not enantiomers are called
diastereomers.
Erythrose has D- and L- forms
(enantiomers). Erythrose is a
diastereomer with threose: they
are not mirror images, but they
have the same chemical and
structural formulas.
More Stereoisomers
• In biology, the amino acids used in proteins are all left-handed (L-amino
acids: L stands for “levo”). Also, the sugars used in the cell are right
handed (D-sugars: D stands for “dextro”); this includes the ribose and
deoxyribose used in nucleic acids.
– If a beam of light is shined on a solution of a stereoisomer, the plane of polarization is
rotated, the to left or to the right. For this reason, stereoisomers are considered to be
optically active.
– However, the D and L refer to a compound’s chemical relationship to the original
stereoisomers of glyceraldehyde, not how the beam of is actually rotated.
• Living systems usually use just one stereoisomer. On the other hand,
when a chiral compound is synthesized in the laboratory from optically
inactive components, it usually appears as a mixture of D and L forms.
This is called a racemic mixture.