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The Basics of Life: Chemistry
Chapter 2
2-
Levels of Organization of Life
• Cellular Organization
cells
organelles
molecules
atoms The cell is the basic unit of life.
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Life Requires Matter and Energy

Matter is anything that has mass and occupies
space.
–

Energy is the ability to do work.
–
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Matter is composed of atoms
Energy organizes atoms to form higher levels of
organization
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Life Requires Energy

Energy is the ability to do work.

There are two types of energy:
1. Potential energy

Stored energy, available to do work
2. Kinetic energy


2-
Energy of motion
Potential energy can be converted to kinetic
energy to do work.
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Law of Conservation of Energy

Energy is never created or destroyed.
–

Energy can only be converted from one form
to another:
–
–
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
The first law of thermodynamics
An object at the top of a hill has potential energy,
as the object rolls down the hill, the potential
energy is converted to kinetic energy.
Cells convert chemical energy into energy used
for life processes
The sum total energy remains constant.
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Forms of Energy

There are five forms of energy:
1. Mechanical energy

Energy of movement
2. Nuclear energy

Energy from reactions involving atomic nuclei
3. Electrical energy

Flow of charged particles
4. Radiant energy

Energy in heat, light, x-rays and microwaves
5. Chemical energy
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
Energy in chemical bonds
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Life Requires Matter

Matter is anything that has mass and
occupies space.

Matter is composed of atoms.

Atoms
–
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The smallest units of matter that can exist
separately
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Atoms

Atoms are composed of:
–
–
–
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Protons – positively charged particles
Neutrons – neutral particles
Electrons – negatively charged particles

Protons and neutrons are located in the
center of the atom (atomic nucleus)

Electrons are found in orbits surrounding the
nucleus (electron orbitals)
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Elements

Elements are chemical substances
composed of the same kind of atoms
–

The elements are listed and organized on
the periodic table
–

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i.e. atoms with the same numbers of protons and
electrons
each element is represented by a symbol of one
or two letters.
The principal elements that comprise living
things are: C, H, O, P, K, I, N, S, Ca, Fe, and Mg.
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The Periodic Table of the Elements
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Atomic Structure

Atoms are composed of:
A. The atomic nucleus

Protons - positively charged
particles

Neutrons – particles with no
charge (neutral charge)
B. Electrons
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
Negatively charged particles

Orbit the nucleus in energy
levels

The more energy an electron
has, the larger its orbit

AreCopyright
constantly
in motion
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Atomic Structure
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Atomic Structure: Atomic Number
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
Atoms of every element have a characteristic
number of protons in the nucleus.

Atoms with the same atomic number have
the same chemical properties and belong to
the same element.

Atomic Number = number of protons in the
nucleus of an atom
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Atomic Number
• Elements on the
periodic table are
organized by atomic
number
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Atomic Structure: Atomic Mass Number

Mass Number = The sum of protons and
neutrons is the atom’s atomic mass.

Each proton and neutron has a mass of
approximately 1 AMU (atomic mass unit).

Electrons are very small in comparison to
protons and neutrons
–
2-
–
mass of an electron is approx. 1/1,800 of an AMU
electrons are negligible to the atomic mass
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Atomic Number
• Elements on the
periodic table are
organized by atomic
mass number
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Atomic Structure: Isotopes
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
Must all atoms of the same element have the
same atomic number?

Must all atoms of the same element have the
same atomic mass?
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Atomic Structure: Isotopes

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Isotopes
–
Atoms of the same element that have different
numbers of neutrons.
–
The atomic weight on the periodic table is the
weighted average mass of all of the isotopes of an
element.
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Isotopes of Hydrogen
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Isotopes of Carbon
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The Periodic Table of the Elements
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Atomic Structure: Atomic Charge

Atoms of the same element have equal
numbers of electrons and protons.

Stable atoms have an overall neutral charge.
–
–

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Atoms that gain or loose protons will become
charged atom
Atoms that gain or loose electrons will become a
charged atom
Charged atoms are known as ions
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Ions


Neutral atoms have the same number of
protons and electrons
Ions are charged atoms
–
–
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Cations – have more protons than electrons and
are positively charged (+)
Anions – have more electrons than protons and
are negatively charged (-)
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Ions

Oxidation - when an atom looses an electron (e)
–

Reduction - when an atom gains an electron (e-)
–
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forms cations
forms anions
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Atomic Structure: Electrons
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
Electrons occupy specific areas around the
nucleus called orbitals.

Electrons constantly move around the
nucleus within the orbital

Orbitals are areas around the nucleus where
an electron is most likely to be found
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Atomic Structure: Electron Orbitals
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Atomic Structure: Electron Levels

Electrons occupy specific energy levels around the
nucleus.
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–

Energy levels hold specific numbers of electrons.
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–
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Electrons closest to the nucleus have the lowest energy.
Electrons further from the nucleus have higher energy.
The first energy level can have up to 2 electrons.
All other energy levels can have up to 8 electrons.
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Atomic Structure: Electron Levels
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Fig. 2.5
• Electrons release energy
as they fall from a high
energy level to a lower
energy level
• Electrons climb from a
lower energy level to a
higher energy level as
they absorb energy
Valence Electrons


The electrons in the outermost energy level
of an atom are known as valence electrons
Atoms seek to have a full outer energy level.
–

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Atoms that have full outer energy levels are
stable.
Octet Rule: Atoms prefer to have 8 valence
electrons:
 Atoms will gain, loose, steal, or share
electrons to fulfill the octet rule
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Valence Electrons
• Elements on the
periodic table are
organized by
number of valence
electrons
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Valence Electrons
H has 1 e-
O has 6
Valence e-s
C has 4
Valence e-s
Na has 1
Valence e-
Ca has 2
Valence e-s
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Valence Electrons

Become Familiar with the Valence Electron
Configuration for the Elements:
–
–
–
–
–
–
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Hydrogen (H): 1
Carbon (C): 4
Oxygen (O): 6
Nitrogen (N): 5
Chlorine (Cl): 7
Sodium (Na): 1
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Valence Electrons
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The Formation of Molecules
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
Atoms will gain, loose, steal, or share
electrons to fulfill the octet rule

Atoms can fulfill the octet rule by loosing,
gaining or sharing electrons with other atoms

This creates chemical bonds between atoms
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The Formation of Molecules

Molecules consist of two or more atoms
joined by a chemical bond.


A compound is a chemical substance made
of two or more different elements combined
in chemical bonds.
–
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O2 , H 2 O
The formula of a compound describes the nature
and proportions of the elements that comprise the
compound.

H2Copyright
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Molecules and Kinetic Energy


Molecules are constantly in motion.
Temperature is a measure of the average speed of
the molecules in a substance.
–
–

Heat is a measure of the total kinetic energy of
molecules.
–

Measured in calories (amount of heat that will raise 1g of
water 1 degree Celsius).
Heat and Temperature are related.
–
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The greater the speed, the higher the temperature.
Measured in Fahrenheit or Celsius
Add heat energy to a substance and the molecules will
speedCopyright
up, ©and
the temperature will rise.
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Physical Changes and Phases of Matter

Three phases of matter
–
–
–

The phase in which a substance exists depends on
its kinetic energy and the strength of its attractive
forces.
–
–
–
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Solid
Liquid
Gas
Solids strong attractive forces, low kinetic energy, little to
no molecular movement.
Liquid enough kinetic energy to overcome the attractive
forces; more molecular movement.
Gas high kinetic energy, little to no attractive forces;
maximum movement.
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Chemical Changes—Forming New Kinds of Matter

Chemical reactions
–

Creating different chemical substances by
forming and breaking chemical bonds.
There are several types of chemical bonds.



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Ionic bonds
Covalent bonds
Hydrogen Bonds
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Ionic Bonds

An ionic bond
–

Example: NaCl
–
–
–
–
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The attraction between oppositely charged ions
Sodium (Na) has one electron in its outer energy level.
Chloride has seven electrons in its outer energy level.
Sodium donates an electron to chloride, each achieving stability.
The positively charged sodium is attracted to the negatively charged
chloride.
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Ionic Bonds
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Ionic Bonds


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Na easily gives up its single valence electron
to fulfill the octet rule
Because Na lost an electron it now has a
positive (+) charge
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Ionic Bonds


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Cl, with 7 valence electrons, easily picks up
the extra electron to fulfill the octet rule
Because Cl gained an electron it now has a
negative (-) charge
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Ionization: Ion Formation
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Ionic Bonds



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Cation Na+ and Anion Cl- attract
Ionic Bonds form crystals
Ionic Bonds dissociate in water
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Covalent Bonds


Covalent bonds form when atoms share 2
or more valence electrons
Covalent bond strength and length depends
on the number of electron pairs shared by
the atoms




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Single Covalent Bond - one pair of e-’s is shared
Double Covalent Bond - two pairs of e-’s are shared
Triple Covalent Bond - three pairs of e-’s are shared
Covalent Bonds are the strongest bonds and
do not dissociate in water
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Covalent Bonds
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Covalent Bonds
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Chemical Bonds: Electronegativity


Electronegativity is an atom’s affinity for
electrons - how badly an atom wants e-’s
Atoms of each of the elements differ in
electronegativity
–


2-
Atomic size, mass, nuclear charge, electron
configuration attribute to differences in electronegativity
Based on what you know about the Na atom, do
you think Na has a high or low electronegativity?
What about Cl?
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Chemical Bonds: Electronegativity
50
Electronegativity(EN) Trend in the Periodic Table
Increasing EN
Greatest
EN
Lowest
EN
51
Electronegativity(EN) Trend in the Periodic Table
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Covalent Bonds: Polar Covalent Bonds
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
A covalent bond is formed by a sharing of electrons,
but electrons are not always shared equally

Differences in electronegativity dictate how electrons
are actually distributed in covalent bonds

In a covalent bond between a highly electronegative
atom and a weakly electronegative atom, electrons
will be pulled more towards the highly
electronegative side of the bond
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Covalent Bonds: Polar Covalent Bonds
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
The water molecule is an important example of a
covalent bond between a highly electronegative
atom and a weakly electronegative atom

As a result of this unequal sharing of electrons, one
side of the molecule (the highly electronegative
oxygen side) has a partial negative charge and the
other side (the hydrogen side) has a partial positive
charge
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Covalent Bonds: Polar Covalent Bonds




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A covalent bond with an unequal sharing of electrons
is known as a polar covalent bond
Polar because it has two ‘poles’, a partial positive
pole and a partial negative pole
A covalent bond with an equal sharing of electrons is
known as a nonpolar covalent bond
A Carbon to Carbon covalent bond is an example of
a nonpolar covalent bond
 C
C
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Hydrogen Bonds

The positive hydrogen end of one polar molecule is
attracted to the negative end of another polar
molecule.
–

Hydrogen bonds hold molecules together.
–

Since they do not hold atoms together, they are not
considered true chemical bonds.
Hydrogen bonds are very important in biology.
–
–
2-
This attraction is a hydrogen bond.
They stabilize the structure of DNA and proteins.
Water molecules can “stick” together with hydrogen bonds.
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Hydrogen Bonds
2-
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Hydrogen Bonds
2-
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Water and Life

The following properties of water make it essential
for life:
–
High surface tension


–
Water molecules stick to each other via hydrogen bonds.
Capillary action moves water through streams, soil, animals
and plants.
High heat of vaporization


A lot of heat is required to break the hydrogen bonds holding
water together.
Large bodies of water absorb a lot of heat.
–
Temperate climates
– Evaporative cooling
2-
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Water and Life

Unusual density properties
–
–

The universal solvent
–
–
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Ice is less dense than water, so ice floats.
Allows aquatic life to survive in cold climates.
Water can form hydrogen bonds with any polar or
ionic compound.
Therefore, many things can be dissolved in water.
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Introduction to Life Science
Chapter 1
Begin 1/20/2011
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Mixtures and Solutions

Mixture
–

Homogeneous mixture
–

Components are distributed equally throughout.
Heterogeneous mixture
–
2-
Matter that contains two or more substances
Components are not distributed equally throughout.
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Mixtures vs. Pure Substances
2-
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Mixtures and Solutions
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
Solutions
– Homogeneous mixture of ions or molecules of two
or more substances.
– The solvent is the substance present in the largest
amount.
– The solutes are the substances present in smaller
amounts.

Aqueous solutions
– water is the solvent
– solid, liquid or gas solutes dissolved in water.
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Concentration of Solutions

Molarity (M) is the unit for concentration



Amount of solute measured in moles


2-
M= Amount of solute
Liters of solution
1 mole = 6.022x1023 particles (atoms, molecules)
of substance
Atomic mass on the periodic table is
measured in grams/mole (g/mol)
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Chemical Reactions

A chemical change:
–
–

2-
Bonds between atoms are made or broken to
make new materials with new properties
Happens via chemical reactions.
In a chemical reaction the elements remain
the same, but the compounds they form and
their properties are different.
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Five Important Chemical Reactions in
Biology
1.
2.
3.
4.
5.
2-
Oxidation–Reduction (Redox Reactions)
Dehydration Synthesis Reactions
Hydrolysis Reactions
Phosphorylation Reactions
Acid–base Reactions
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Chemical Reactions and Energy

Chemical reactions produce new compounds
with less or more potential energy.
–
Catabolic Reactions


–
Anabolic Reactions


2-
Breakdown large molecules into smaller ones
Release energy
Build larger molecules from smaller ones
Require energy
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Chemical Equations

–
A chemical equation describes what
happens in a chemical reaction.
Photosynthesis is described by the following
equation:
–
6CO2 + 6H2O → C6H12O6 + 6H2O




Reactants substances that are changed, usually on the
left side of the equation.
Products new chemical substances formed, usually on
the right side of the equation.
Chemical equations must be balanced
–
2-
They must have the same number of atoms of each
element on both sides of the equation
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Five Important Chemical Reactions in
Biology
1.
2.
3.
4.
5.
2-
Oxidation–reduction
Dehydration synthesis
Hydrolysis
Phosphorylation
Acid–base reactions
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Oxidation-Reduction Reactions

Reactions in which electrons (and their energy)
are transferred from one atom to another.
–
–

Redox Reactions are very important in
biological processes

2-
Oxidation - an atom loses an electron.
Reduction - an atom gains an electron.
Cellular Respiration, Photosynthesis
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Oxidation-Reduction Reactions

LEO the lion says GER
–
–
2-
If it Looses Electrons, it’s Oxidized
If it Gains Electrons, it’s Reduced
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Oxidation-Reduction Reactions

Oxidation and Reduction reactions always
occur in pairs:
–
–

2-
If a an atom or molecule is reduced, then another
atom or molecule must have been oxidized
If a an atom or molecule is oxidizied, then
another atom or molecule must have been
reduced
For this reason Oxidation and Reduction
Reactions are known as Redox Reactions
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Oxidation-Reduction Reactions

In Biology, hydrogen atoms are commonly
transferred during redox reactions.

Cellular Respiration
-
-
The sugar glucose is oxidized to form carbon
dioxide and oxygen is reduced to form water.
Energy is released in the process.
–C6H12O6

2-
–
+ 6O2 → 6H2O + 6CO2+ Energy
Sugar + oxygen → water+ carbon dioxide + energy
(glucose)
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Oxidation-Reduction Reactions
2-
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Dehydration Synthesis Reaction

Biological molecules are typically
Marcomolecules
–
–

Macromolecules are Polymers assembled
from smaller Monomers
–
–
2-
very large molecules with high molecular weights
DNA over a meter long
Monomers - small, identical or similar subunits
Polymers - covalently bonded monomers
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Dehydration Synthesis Reaction

Proteins
–

Nucleotides
–

Nucleotide monomers polymerize to form DNA
and RNA Macromolecules
Carbohydrates
–
2-
Amino Acid monomers polymerize to form
proteins
Simple sugar monomers polymerize to form
complex sugars
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Dehydration Synthesis Reaction

Two small molecules are joined to form a larger
molecule

Example:
–
Joining amino acids to form proteins.
–
NH2CH2CO-OH + H-NH CH2CO-OH  NH2CH2CO-NH CH2CO-OH + HOH
amino acid 1

+ amino acid 2
= protein
water

2-
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+
Dehydration Synthesis Reaction
2-

Dehydration Synthesis - the chemical
reaction for how living cells synthesize
macromolecules

A covalent bond is formed between
monomers and water is produced as a
product of the reaction

As the name implies, water is lost during the
reaction - dehydration
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Dehydration Synthesis
Reaction
• Monomers covalently bond together to
form a polymer with the removal of a
water molecule
– A hydroxyl group is removed from one monomer
and a hydrogen from the next
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Dimer
Monomer 1
Monomer 2
O
OH HO
H+ + OH–
H2O
(a) Dehydration synthesis
Figure 2.15a
2-67
Dehydration Synthesis Reaction

A hydroxyl (-OH) group is removed from one
monomer, and a hydrogen (H+) from another

A new covalent bond is formed between the
monomers
 -OH
–
2-
and H+ come together to form water (H2O)
Water is released as a by-product
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Hydrolysis Reactions



When a larger molecule is broken down into is
smaller smaller subunits
Opposite of a dehydration synthesis
Example:
–
Digesting proteins into amino acids.
–
NH2CH2CO-NH CH2CO-OH + H-OH  NH2CH2CO-OH + H-NH CH2CO-OH
Protein
–
+ water = amino acid 1 + amino
acid 2
2-
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Hydrolysis Reactions



The reaction for the separation of joined
monomers
“Splitting with water”
Opposite of dehydration synthesis
–
–
–
–
2-
a water molecule ionizes into –OH and H+
the covalent bond linking one monomer to the
other is broken
the –OH is added to one monomer
the H+ is added to the other
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Hydrolysis Reactions
• Splitting a polymer (lysis) by the addition of a water
molecule (hydro)
– a covalent bond is broken
• All digestion reactions consists of hydrolysis reactions
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Dimer
Monomer 1
Monomer 2
OH
O
H2O
HO
H+ + OH–
(b) Hydrolysis
Figure 2.15b
2-68
Biological Molecules
85
Phosphorylation Reactions

The addition of phosphate groups to other
molecules,
–

Phosphate groups (P) are clusters of oxygen and
phosphate atoms.
Functions of Phosphorylation:
–
–
Phosphorylation is used to turn enzymes on or off
Phosphorylation is used to transfer energy
–
2-
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87
Phosphorylation Reactions

Bonds between phosphate groups and other
molecules contain high potential energy.
–
–
–
2-
When these bonds are broken, the energy that is released can
be used by the cell to do work.
Phosphorylation reactions are commonly used to transfer
potential energy.
Q-P + Z  Q + Z-P
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Acids and Bases

An acid
–
–

A base
–
–

Compounds that release hydroxide ions (OH-) into a solution or
accept hydrogen ions (H+) from solution
Sodium hydroxide, ammonia
Acids and Bases react to neutralize each other producing
water and salts
–
2-
Ionic compounds that release hydrogen ions (H+) into a solution
Phosphoric acid, hydrochloric acid
The -OH from bases are negatively charged, they will react with a
positively charged H+ from an acid to produce water.
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Acids and Bases

Acid: a chemical that releases H+ ions
- Hydrochloric acid, HCl
 HCl  H+ + Cl-
- Sulfuric acid, H2SO4

H2SO4  2H+ + SO4-2
- Carbonic acid, H2CO3

2-
H2CO3  H+ + HCO3-
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
Acids and Bases

Base: a chemical that releases -OH ions or
accepts H+1 ions
 Sodium Hydroxide, NaOH

NaOH  Na+1 + OH-1
 OH-1 + H+1  H2O
 Bicarbonate, HCO3
2-

HCO3- + H+  H2CO3
carbonic acid
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Acid-Base Reactions
Occurs when ions from an acid interact with
ions from a base.
 This type of reaction allows harmful acids
and bases to neutralize one another.
 HCl
+ NaOH → NaCl + HOH
 H+ Cl+ Na+ OH- → Na+ Cl- + H+ OH


2-
Hydrocloric +
acid
Sodium
hydroxide
Sodium
chloride
+ Water
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Salts

Usually formed when acids and bases react
–
–
–



2-
The dissociated hydrogen ions and hydroxide ions join to
form water.
The remaining ions form ionic bonds, creating a salt.
This is an example of neutralization:
H+Cl- + Na+OH- →
Hydrocloric +
acid
Sodium
hydroxide
Na+Cl- + H+OHSodium
+ Water
chloride
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Acid-Base Reactions
 Acids and Bases come in pairs:
H2CO3  H+1 + HCO3Carbonic acid
Bicarbonate




HCO3- + H+  H2CO3
Bicarbonate
Carbonic acid


2-
H2CO3
H+1 + HCO3-
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Acids and Bases
2-

The strength of an acid or base is
determined by how completely it will
dissociate in water.

Strong acids release almost all of their
hydrogen ions into water.

Strong bases release almost all of their
hydroxide ions into water.
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Some Common Acids, Bases and Salts
2-
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Acids and Bases
2-

The acidity/basicity of a solution is measured
in terms of the pH Scale

pH is a measure of hydrogen ion
concentration [H+] in a solution

Solutions with high hydrogen ion
concentrations are acidic

Solutions with low hydrogen ion
concentrations are basic (alkaline)
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The pH Scale
 Hydrogen ion concentration [H+1] is
the basis of the pH scale:

pH = -log [H+]

or
pH =
1
log [H+]
 Greater H+1 concentration = lower pH
(acidic)
2-
 Lower H+1 concentration = higher pH (basic)
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pH

The pH Scale makes large numbers easier to use

Log Function:
– The Log of 10000000 = 7
– The Log of 0.0000001 = -7

There is a 10-fold difference in hydrogen ion
concentration between solutions that differ by one
pH unit.
–
2-
A solution with pH 4 has ten times as many hydrogen ions
as a solution with pH 5.
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The pH Scale
2-
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The pH Scale
 Hydrogen ion concentration [H+1] of
pure water is 10-7 moles [H+] / L

pH =

1
log [H+]
pHwater =
1
log
[10-7]
pHwater = 7
 The pH scale is based on the pH of

2-
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The pH Scale

Is water an acid or a base?


H2O  OHH 2O
+
+
H+ 
H+
H 3O +
 Water acts as an acid and a base
 Pure water exists as a balance of H2O,
OH-1, and H3O+1
2-
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The pH Scale
2-
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