Biochemistry
Background information
relevant to your study of
biochemistry (you learned it in
your previous science classes)
Properties of Matter

Matter – any substance that has mass and
volume.
– Mass - the quantity of matter in an object.
– Volume - the amount of space that the matter takes
up.

The more properties we can identify for a
substance, the better we can understand its
nature!
Properties of Matter
Physical properties of matter can be
observed and measured without changing
the identity of the matter.
 Physical change - can affect size, shape,
or color of a substance but does NOT
affect the composition of the matter.

Physical Properties of Matter
Mass
 Color
 Volume
 Odor
 Texture
 Taste

Luster
 Hardness
 Melting Point
 Boiling Point
 Phase
 Density

Chemical Properties of Matter
Chemical properties of matter describe a
substance’s ability to change into a NEW
substance as a result of a chemical
change.
 Chemical change - bonds are broken and
new bonds form between atoms.

– Substances display different physical and
chemical properties after the change.
– A chemical change is irreversible!
Signs a chemical change
occurred include…
Production of light
 Production of heat
 Color change
 Gas production (bubbles)
 Odor
 Sound
 A substance was created that wasn’t there before!



Examples: burning coal, ripening banana, baking a cake.
Examples: food is metabolized in body, photosynthesis.
Physical vs. Chemical Changes
Physical Change:
 Reversible
Chemical Change:
 Not reversible
– You can “un-freeze
water”

No new substance is
formed
– Water and ice are both
H2O molecules
– You can’t “un-burn”
wood

A new substance is
formed
– Burning wood results
in CO2, ash, etc.
Phases/States of Matter

Phase of matter - physical property of
matter that describes one of a number of
different states of the same substance.
Phases/States of Matter
Solid - definite shape, definite volume
 Liquid - no definite shape, definite volume
 Gas - no definite shape, no definite
volume
 Plasma - no definite shape, no definite
volume

– highly ionized gas that occurs at high temps
Chapter 2 - Biochemistry
Atoms

Atom - basic unit of matter.
– Greek word “atomos” – unable to cut.
– Atoms are the smallest component of a cell.
Atoms

Atoms compose all living and non living
things.
– Atoms contain subatomic particles: protons
(+), neutrons (neutral), and electrons (-).
– Protons and neutrons are found in the center
of the atom in the atomic nucleus.
– Electrons float around the nucleus in energy
levels and are attracted to the nucleus by the
protons (+’s attract –’s).
Atoms

Atoms are electrically neutral because
their proton number and electron number
balance out their charges.
Atomic Structure
Protons
Neutrons
Electrons
Atomic Structure
The Nature of Atoms

Protons determine the identity of an atom!
– Atomic number – number of protons in the
nucleus of an atom. Each atom has a
different proton number (identity).

Electrons determine how an atom
behaves!
– Electrons float around the nucleus in energy
levels; most of an atom is empty space.
The Nature of Atoms

Mass number is the total number of
protons and neutrons in the nucleus.
– Most of the mass of an atom is in the nucleus!
Atomic Number
Symbol
Name
Atomic Mass (Mass #)
Atomic number equals the number of protons in nucleus.
Atomic mass or mass number equals the number of protons + neutrons.
Atoms have Energy
Electrons in an atom have energy.
 Energy is needed to keep electrons in the
clouds so that they are not pulled into the
nucleus.

Atoms have Energy
•Each energy level can
hold a certain number
of electrons.
•First level: 2 electrons
•Second level: 8 electrons
•Third level: 8 electrons
•Fourth level: 10 electrons
Elements

Elements - substances that are composed
of only one type of atom.
– Cannot be chemically broken down to any
other substances.
– Are represented by chemical symbols on
periodic table.
– More than 100 elements are known, about 25
are found in living organisms.
 6 most abundant include: O, C, H, N, P, S
Isotopes

Isotopes - atoms of the same element that
differ in the number of neutrons.
– Still have the same number of protons (proton number identifies the substance).

Isotopes of an element have the same
chemical properties. They differ by the
number of neutrons (a physical property).
Radioactive Isotopes

Radioactive isotopes - are unstable and from
time to time breakdown releasing radiation
from their nucleus.
– Used to study organisms, diagnose disease (as
tracers), treat disease (kill cancer cells), sterilize
food, measure the ages of rocks.
– Radiation is dangerous! It can kill or damage living
things (i.e. Chernobyl’s radioactive fallout).
Chemical Compounds
A chemical compound is a group of atoms
held together by chemical bonds.
 Compounds are represented by chemical
formulas.- show the proportion of atoms
in a compound
 Examples of chemical formulas:

– NaCl – table salt
– H2O – water
– NH3 – ammonia
– C6H12O6 - glucose
Interactions of Matter
Atoms want to achieve stability – full
outermost energy level (valence shell).
 In order to achieve stability, atoms will
either gain, lose, or share electrons with
other atoms in a process called chemical
bonding.

– Atoms will bond with other atoms if the
bonding will give both atoms complete
outermost energy levels.
– Valence electrons- electrons in outermost
energy level
Chemical Reactions

Chemical reaction – a process that
changes one set of chemicals into another
set of chemicals; involves the breaking
and reforming of chemical bonds.
– Reactants - chemicals that undergo a change
(left side of equation).
– Products - chemicals that are the result of a
change (right side of equation).
A
+
B
--------->
C
+
D
Energy in Chemical Reactions

Energy is stored within chemical bonds.
– When bonds are broken, energy is released.
– All living organisms must have a source of
energy to carry out chemical reactions!

Two types or reactions deal with the
energy stored in chemical bonds:
 Endergonic reactions
 Exergonic reactions
Endergonic Reactions

Endergonic reactions – reactions that
absorb energy.
– Need a source of energy to trigger the
reaction (don’t occur spontaneously).
– Reactions tend to feel cold.
Exergonic Reactions

Exergonic reactions – reactions that
release energy.
– Energy is released as heat, light, or gas.
– Can occur spontaneously.
– Often feel warm.
Ionic bonds

Ionic bonds – chemical bonds that transfer
electrons from one atom to another
forming charged particles called ions.
Example: NaCl is a compound
formed by ionic bonds.
– Na has 1 electron in its outermost energy level.
 When Na looses an electron, it becomes positively
charged (Na+, or a sodium ion).
– Cl needs 1 electron to fill its outermost energy
level.
 When Cl gains an electron from Na, it produces a
negatively charged ion, Cl-.
– The two oppositely charged ions are attracted
to one another and form NaCl through
transferring electrons in ionic bonding.
Covalent bonds

Covalent bonds – chemical bond formed
by the sharing of electrons so that each
atom fills its outermost energy level.
– Most bonds in living organisms are covalent.
– Examples: H2O, CO2, NH3, C6H12O6.

Molecule – smallest particle of a covalently
bonded compound.
Dogs Teaching Bonding:
http://www.youtube.com/watch?v=_M9khs87xQ8&sns=em
Intermolecular Forces
Intermolecular forces - also called
molecular attraction.
 Are forces of attraction between stable
molecules.

– Example: hydrogen bonds (see section 33).
Two types of IM forces:

Cohesion - intermolecular force of
attraction between LIKE molecules.

Adhesion - intermolecular force of
attraction between DIFFERENT
molecules.
Intermolecular Forces

Why do they occur?
– Due to differences in charge densities or
uneven distribution of electrons!
Acids, Bases, and pH

Acid – a substance that releases hydrogen ions
(H+ ) when dissolved in water.
– Example: HCl ---> H+

+ Cl-
Base – a substance that releases hydroxide ions
(OH-) when dissolved in water.
– Example: NaOH ---> Na+ + OH-
Acids, Bases, and pH
Water is a neutral solution - water
separates forming an equal number of
hydrogen and hydroxide ions.
 Neutralization reaction - Hydrogen ions
and hydroxide ions react to form water.

– Occurs when H+ ions from strong acids are
mixed in perfect ratios with OH- ions from
strong bases.
H+ + OH- -----> H2O
pH Scale
pH – measures the amount of hydrogen in a
solution, each measurement of pH represents ten
times.
 pH Scale - ranges from 0 to 14.

–
–
–
–
Less than 7 is for acids (more H+ than OH-).
Greater than 7 is for bases (more OH- than H+).
7 is neutral (equal amounts of H+ and OH- in solution).
Most cells have a pH of 6.5-7.5.
 Controlling pH is an example of homeostasis.
pH Scale

Acids - substances that forms hydrogen
ions when dissolved in water.
– The more hydrogen ions (less hydroxide) the
more acidic.

Bases - substances that forms hydroxide
ions when dissolved in water.
– The more hydroxide ions (less hydrogen) the
more basic or alkaline.
pH Scale

What happens when acid is added to a solution?
– As more acid is added the pH will go down, but the
H+ concentration goes up.

What happens when base is added to a
solution?
– As more base is added the pH will go up, but the H+
concentration goes down.
Buffers

Buffers – weak acids or bases that can
react with strong acids or bases to prevent
sharp, sudden changes in pH.
– Are important for maintaining homeostasis in
living organisms.
 Ex. Carbonic acid and sodium bicarbonate buffer
your blood’s pH.
Properties of Water

All cells contain water.
– About two thirds of the molecules in our body
are water.
Water provides a medium in which other
molecules can interact.
 Water exists as all three states/phases of
matter.
 Water expands when it freezes!!!!

Water is Polar

Water is a polar molecule - molecule has
slight charge (+ or -) on each end due to
uneven distribution of electrons.
– Oxygen pulls hydrogen’s electrons closer to it
therefore the oxygen atom is slightly negative
and the hydrogen becomes slightly positive.
– This is the most important property of water!
 Allows a strong attraction between water
molecules or between water and other polar
molecules!
Polar vs. Non-Polar Molecules

Polar - unequal distribution of charge means a great
amount of attraction between molecules.

Non-Polar - equal distribution of charge means a
weak attraction between molecules.
Do Polar and Non-Polar
Solutions Mix?
•Polar solutions mix with other polar solutions!
•Example: Milk and water.
•Non-polar solutions mix with other non-polar
solutions!
•Example: Oil and grease.
•Polar solutions will NEVER mix with non-polar
solutions!
•Example: Italian salad dressing.
Water clings to itself & other
molecules
-Cohesion – Intermolecular force of attraction
between like molecules.
 Water molecules cling to other WATER molecules
(hydrogen bonding) – Beading of water on a
smooth surface.
– Adhesion – Intermolecular force of attraction
between different molecules.
 Water molecules cling to other molecules –
Meniscus in a graduated cylinder.
Water is good at forming
mixtures
•Due to slight charge of water molecules.
•Mixture - substance composed of two or more
elements or compounds that are mixed together
but not chemically combined (are not linked by
chemical bonds).
•Examples: salt and pepper stirred together;
atmosphere.
•Two types of mixtures: Solutions & Suspensions
Water’s role in suspensions
Suspension – a mixture where the solute does not
fully dissolve.
 Solute will settle out.
 Example blood (plasma and blood cells).
Water’s role in solutions
Solution – small particles are dispersed in mixture,
all components are evenly distributed.
 Solute the substance that is dissolved.
 Solvent the substance that does the dissolving.
 Water acts as a solvent to dissolve solutes (ex.
sugar) forming solutions.
Water’s role in solutions

Water dissociates - breaks down forming
charged particles called ions (H+ and OH-) when
its bonds are broken.
H2O

-----
H+
+ OH-
Other compounds also dissociate (break down
into their individual ions) when dissolved in
water.
Ex. NaCl
------>
Na+
+
Cl-
Water has a large heat capacity
Heat capacity – amount of heat required to
change a substance’s temperature by a given
amount.
 Is a result of the multiple hydrogen bonds
between water molecules.
 A large amount of heat energy is required to
cause the molecules to move faster (which is
how the temperature of the water is raised).

– Allows large bodies of water to absorb large amounts
of heat with only a small change in temperature.
– Alllows for regulation of cell temperature.
Water has properties of
capillarity

Capillary action– the interplay of cohesion and
adhesion to hold a solution in a thin tube against
the force of gravity.
– Draws water out of the roots of plants and up into the stems
and leaves.
– Helps move blood through the body.
Biological Macromolecules
Inorganic vs. Organic
Compounds
C, H, N, and O make up almost all
chemical compounds in living organisms.
 Organic compounds - contain carbon.

– Carbon can form long carbon chains by
bonding to other carbon atoms.
 Unique because they are very strong/stable!

Inorganic compounds - do not contain
carbon.
– Exception: CO2
Carbon Compounds Polymerize

Polymerization - process by which large
compounds are constructed by joining
together smaller compounds (monomers).
– Monomers are joined by chemical bonds to form
polymers.
– Very large polymers are called macromolecules.
Building macromolecules
Dehydration synthesis or condensation
reactions - reactions that joins two
monomers into a polymer and involves the
loss of water.
 Hydrolysis - reaction that breaks a
polymer into monomers by using a water
molecule.

Carbon
Carbon is a key component of biological
macromolecules for two reasons:
 1. Carbon atoms have 4 valence electrons.

– Allows them to form strong covalent bonds
with many other elements.

2. Carbon atoms can bond to other carbon
atoms.
– Gives the ability to form chains, rings,
multiple bonds, and millions of different large,
complex structures.
Carbohydrates

Carbohydrates - macromolecules that are
composed of the atoms carbon, hydrogen,
and oxygen in the proportion of 1:2:1.
– 1 carbon : 2 hydrogen : 1 oxygen.
– Examples: sugars and starches.
Carbohydrates - Monosaccharides

Monosaccharide – simple, single sugar
molecule.
– Examples: glucose (produced by green plants),
fructose (fruits), and galactose (milk).
– Sugars are important for living things because
they contain a great deal of energy.
Carbohydrates - Disaccharides

Disaccharide – 2 sugar molecules bonded
together by a covalent bond.
– Examples: lactose or sucrose.
Carbohydrates - reactions

Dehydration synthesis - reaction that joins
two monosaccharides into a disaccharide
and involves the loss of water.
– Hydrolysis - reaction that breaks a
disachharide into monosaccharides by using a
water molecules.
Carbohydrates - Polysaccharides

Polysaccharides – macromolecules formed from
linking many monosaccharides together.
– Ex. Starch – a polysaccharide plants use to store
energy; many glucose molecules bonded together.
– Ex. Glycogen - stored form of glucose from starch;
stored for energy in liver of animals.
– Ex. Cellulose – chains of glucose, structurally different
from starch, tough flexible molecule found in plants.
Nucleic Acids

Nucleic acids - Polymers made of building
blocks (monomers) called nucleotides.
– Contain H, O, P, C, and N.
– Made up of nitrogenous base, 5-C sugar, 1-3
phosphate groups
 Nitrogenous bases: Adenine, Thymine, Guanine,
Cytosine, Uracil (RNA only)
– Nucleic acids store and transmit hereditary
information.
– Example – DNA and RNA.
 DNA has a deoxyribose sugar, RNA has a ribose
sugar.
Lipids

Lipids - organic compounds that are oily or
waxy.
–
–
–
–
Common examples: fats, oils, and waxes.
Lipids are made of C, H, and O (no ratio H to O).
Multiple rings of C
Lipids function in energy storage, form biological
membranes, and act as chemical messengers.
 Lipids have more energy than carbohydrates because lipids
have more hydrogens bonded to the carbon chain.

Lipids have a water loving portion, and a water
hating portion.
Types of Lipids

Lipids are polymers made of monomers of
fatty acids and glycerol.
– Ex fatty acids: oleic acid, palmitic acid (produced by liver), linolenic acid
(essential)
– Saturated Lipids:
 contain the maximum number of carbon to
hydrogen bonds.
 Example – animal fats.
 Also called “bad fats”.
Types of Lipids
– Unsaturated Lipids:
 Contain carbon to carbon double bonds; less
hydrogen.
 Example – plant oils (corn oil, vegetable oil)
 Also called “good fats”.
Types of Lipids

Sterols - ringed structures that play roles
in building cells and carrying messages.
– Example – cholesterol; hormones.

Phospholipids - contain parts that dissolve
well in water and parts that do not.
– Spontaneously form bilayers to keep water
hating portions protected and water loving
portions in contact with water.
Proteins

Proteins - polymers made of building blocks
(monomers) called amino acids.
– Amino acids have an amino group, carboxyl group
and an R group.
 Differences in R groups make each of the 20 amino acids
different.
Peptides
Peptides - short polymers of amino acids linked
by peptide bonds.
 Peptide bonds are covalent bonds that join
together amino acids.
 Once a polypeptide (long chain of amino acids)
is formed, it must be folded into a 3-D shape
before it is called a protein.

– The shape is important for recognition of the protein
by the cell and for the actions of the protein.
Proteins
Used to form skin, muscle, hair.
 Proteins play a role in metabolism, help
fight disease, used to assist chemical
reactions (enzymes), and signaling other
cellular functions.
 Enzymes are special proteins.


Enzymes
Enzymes – proteins that act as biological
catalyst and speed up the rate of a
chemical reaction.
– Enzymes are not changed by the reaction (so
they can be re-used).
– Enzymes are very specific – they will only
speed up one chemical reaction.
– Enzymes speed up chemical reactions by
lowering the “start-up” energy of a reaction.
– The names of most enzymes will end in “ase”
such as ligase, amylase, polymerase.
Enzymes

Enzymes will bind to the reactants of the
chemical reaction that it will catalyze.
– The reactants that enzymes bind to are called
substrates.
– The site in which the substrates are brought
to is called the active site.
– Substrates will fit into the active site like a
lock and key.
 If the substrates do not fit in the active site, it is
the wrong enzyme and it will not catalyze a
reaction!
Enzymes
– Enzymes function in
regulating chemical
pathways, making
materials that cells
need, releasing
energy, and
transferring
information.