Download Biochemistry unit notes part I

11/9/2015
Physical Properties of Matter
Background information
relevant to your study of
biochemistry (you learned it in
your previous science classes)
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Properties of Matter
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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.
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Luster
Hardness
Melting Point
Boiling Point
Phase
Density
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.
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– Substances display different physical and
chemical properties after the change.
– A chemical change is irreversible!
Signs a chemical change
occurred include…
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.
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Chemical Properties of Matter
The more properties we can identify for a
substance, the better we can understand its
nature!
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Mass
Color
Volume
Odor
Texture
Taste
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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.
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Physical vs. Chemical Changes
Physical Change:
 Reversible
Chemical Change:
 Not reversible
– You can “un-freeze
water”
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No new substance is
formed
– Water and ice are both
H2O molecules
Chapter 2 - Biochemistry
– You can’t “un-burn”
wood
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A new substance is
formed
– Burning wood results
in CO2, ash, etc.
Phases/States of Matter
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Phase of matter - physical property of
matter that describes one of a number of
different states of the same substance.
Atoms
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– Greek word “atomos” – unable to cut.
– Atoms are the smallest component of a cell.
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
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– highly ionized gas that occurs at high temps
Atom - basic unit of matter.
Atoms
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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).
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Atoms
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The Nature of Atoms
Atoms are electrically neutral because
their proton number and electron number
balance out their charges.
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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).
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Electrons determine how an atom
behaves!
– Electrons float around the nucleus in energy
levels; most of an atom is empty space.
Atomic Structure
The Nature of Atoms
Protons
Neutrons
Electrons
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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 Structure
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.
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Atoms have Energy
Isotopes
Electrons in an atom have energy.
 Energy is needed to keep electrons in the
clouds so that they are not pulled into the
nucleus.
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– Still have the same number of protons (proton number identifies the substance).
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Atoms have Energy
•First level: 2 electrons
•Second level: 8 electrons
•Third level: 8 electrons
•Fourth level: 18 electrons
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).
Elements
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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 of an element have the same
chemical properties. They differ by the
number of neutrons (a physical property).
Radioactive Isotopes
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•Each energy level can
hold a certain number
of electrons.
Isotopes - atoms of the same element that
differ in the number of neutrons.
Chemical Compounds
A chemical compound is a group of atoms
held together by chemical bonds.
 Compounds are represented by chemical
formulas.
 Examples of chemical formulas:
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– NaCl – table salt
– H2O – water
– NH3 – ammonia
– C6H12O6 - glucose
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Interactions of Matter
Atoms want to achieve stability – full
outermost energy level.
 In order to achieve stability, atoms will
either gain, lose, or share electrons with
other atoms in a process called chemical
bonding.
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– Atoms will bond with other atoms if the
bonding will give both atoms complete
outermost energy levels.
Covalent bonds
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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.
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Molecule – smallest particle of a covalently
bonded compound.
Ionic bonds
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Ionic bonds – chemical bonds that transfer
electrons from one atom to another
forming charged particles called ions.
Intermolecular Forces
Intermolecular forces - also called
molecular attraction.
 Are forces of attraction between stable
molecules (NOT atoms!)
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– Example: hydrogen bonds between water
molecules.
Example: NaCl is a compound
formed by ionic bonds.
Two types of IM forces:
– Na has 1 electron in its outermost energy level.
 When Na looses an electron, it becomes positively
charged (Na+, or a sodium ion).
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Cohesion - intermolecular force of
attraction between LIKE/THE SAME
molecules.
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Adhesion - intermolecular force of
attraction between DIFFERENT
molecules.
– 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.
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Properties of Water
Intermolecular Forces
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Why do they occur?
– Due to differences in charge densities or
uneven distribution of electrons!
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– 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 – this is
why ice floats!
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Polar vs. Non-Polar Molecules
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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!
All cells contain water.
Water is Polar
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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.
– Polarity is the most important property of
water!
 Allows a strong attraction between water
molecules or between water and other polar
molecules!
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.
•Example: Oil and grease.
•Polar solutions will NEVER mix with non-polar
solutions!
•Example: Italian salad dressing.
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Water clings to itself & other
molecules
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.
– Adhesion – Intermolecular force of attraction
between different molecules.
 Water molecules cling to other molecules –
Meniscus in a graduated cylinder.
Water has a large heat
capacity/specific heat
Water is good at forming
mixtures
•Due to slight charge of water molecules.
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•Mixture - substance composed of two or more
elements or compounds that are mixed together
but not chemically combined (are not linked by
chemical bonds).
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•Examples: salt and pepper stirred together;
atmosphere.
– Allows large bodies of water to absorb large amounts
of heat with only a small change in temperature.
– Allows for regulation of cell temperature.
•Two types of mixtures: Solutions & Suspensions
Water has properties of
capillarity
Water’s role in suspensions
– Suspension – a mixture where the solute does not
fully dissolve.
Heat capacity/specific heat – 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).
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 Solute will settle out.
 Example blood (plasma and blood cells).
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.
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Remember, water’s unique properties (the ones that
allow it to sustain life) are ultimately based upon the fact
that is very polar!
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Water’s role in solutions
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Water dissociates - breaks down forming
charged particles called ions (H+ and OH-) when
its bonds are broken.
H2O
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-----
H+
+ OH-
Other compounds also dissociate (break down
into their individual ions) when dissolved in
water.
Ex. NaCl
------>
Na+
+
Cl-
Acids, Bases, and pH
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Acid – a substance that releases hydrogen ions
(H+ ) when dissolved in water.
– Example: HCl ---> H+
pH Scale
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+ Cl
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Base – a substance that releases hydroxide ions
(OH-) when dissolved in water.
pH – measures the amount of hydrogen in a
solution, each measurement of pH represents ten
times.
pH Scale - ranges from 0 to 14.
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–
–
–
– Example: NaOH ---> Na+ + OH-
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.
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.
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– Occurs when H+ ions from strong acids are
mixed in perfect ratios with OH- ions from
strong bases.
H+ + OH- -----> H2O
pH Scale
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Acids - substances that forms hydrogen
ions when dissolved in water.
– The more hydrogen ions (less hydroxide) the
more acidic.
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Bases - substances that forms hydroxide
ions when dissolved in water.
– The more hydroxide ions (less hydrogen) the
more basic or alkaline.
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pH Scale
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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.
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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
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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.
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