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Chapter 2
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2-1
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Matter, Energy, and Life


Matter is anything that has mass and occupies
space.
Energy is the ability to do work.
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There are two types of energy:

Potential energy
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
Kinetic energy
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Stored energy, available to do work
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.
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
Energy can be converted from one form to
another, but the total energy remains
constant.
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–
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The first law of thermodynamics
An object at the top of a hill has potential energy
based on its location.
When the object rolls down the hill, the potential
energy is converted to kinetic energy.
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Forms of Energy

There are five forms of energy:
1.
Mechanical energy
•
2.
Nuclear energy
•
3.
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Flow of charged particles
Radiant energy
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5.
Energy from reactions involving atomic nuclei
Electrical energy
•
4.
Energy of movement
Energy in heat, light, x-rays and microwaves
Chemical energy
•
Energy in chemical bonds
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What is the nature of matter?

Atoms
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
Elements
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The smallest units of matter that can exist
separately
Chemical substances composed of the same kind
of atoms
Listed on the periodic table
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
2-6
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Atomic Structure

Atoms are composed of:
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The atomic nucleus

Protons - positively
charged
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–

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Neutrons – no charge
Electrons


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Atomic number-the
number of protons
All atoms of the same
element have the same
number of protons.
Orbit the nucleus in energy
levels
Are constantly in motion
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Atomic Structure
2-8
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Elements

Atoms of the same element have equal numbers
of electrons and protons.
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
Isotopes
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Atoms of the same element that have different
numbers of neutrons.
Atomic weight-the average of all of the isotopes in a
mixture.
Mass number
–
2-9
Thus, they have a neutral charge.
The sum of protons and neutrons in the nucleus.
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Isotopes of Hydrogen
2-10
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Electrons

Electrons occupy specific energy levels around the
nucleus.
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
Energy levels hold specific numbers of electrons.
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The first energy level can have up to 2 electrons.
All other energy levels can have up to 8 electrons.
Atoms seek to have a full outer energy level.
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Electrons closest to the nucleus have the lowest energy.
Atoms that have full outer energy levels are inert.
Other atoms seek to fill their outer energy levels through
chemical bonds.
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Electrons
2-12
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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 elements combined in
chemical bonds.
–
The formula of a compound describes the nature
and proportions of the elements that comprise the
compound.

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H 2O
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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.
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Heat is a measure of the total kinetic energy of
molecules.
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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
speed up, and the temperature will rise.
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Kinetic Energy, Physical Changes
and Phases of Matter

Three phases of matter
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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.
Remember: Atoms form chemical bonds to fill
their outermost electron energy levels, achieving
stability.
There are several types of chemical bonds.
–
We will discuss:

2-16

Ionic bonds
Covalent bonds
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Ionic Bonds

Atoms can gain or lose electrons to achieve a full outermost energy
level.
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An ionic bond
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The attraction between oppositely charged ions
Example: NaCl
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Atoms with charge are called ions.
When an atom gives away an electron, it ends up with more protons than
electrons and gains a positive charge; cation.
When an atom accepts an electron, it ends up with more electrons than
protons and gains a negative charge; anion.
This process is called ionization.
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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Ion Formation
2-18
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Covalent Bonds


Atoms can achieve full outermost energy levels by
sharing electrons instead of exchanging them.
A covalent bond is formed by the sharing of electrons.
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The atoms sharing electrons sit close enough together so that
their outer energy levels overlap.
Single covalent bond- one pair of electrons is shared.

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Double covalent bond- two pairs of electrons are shared.
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ethylene
Triple covalent bond- three pairs of electrons are shared.

2-19
H2
N2
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Covalent Bonds
2-20
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Hydrogen Bonds

The positive hydrogen end of one polar molecule is
attracted to the negative end of another polar
molecule.
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Hydrogen bonds hold molecules together.
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Since they do not hold atoms together, they are not
considered true chemical bonds.
Hydrogen bonds are very important in biology.
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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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Water: The Essence of Life

Water has special properties that make it an
essential molecule for life.
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H2O
Electrons are shared unequally by hydrogen and
oxygen.
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This is a polar covalent bond.
Oxygen has more protons than hydrogen.
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The electrons spend more time around oxygen than around
hydrogen.
The oxygen end of water is more negative.
The hydrogen end of water is more positive.
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Hydrogen Bonds
2-23
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Mixtures and Solutions

A mixture
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Matter that contains two or more substances that are not in set
proportions
A solution is a homogeneous mixture of ions or
molecules of two or more substances.
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Components are distributed equally throughout.
The process of making a solution is called dissolving.
The solvent is the substance present in the largest amount.
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Frequently the solvent is a liquid.
The solutes are the substances present in smaller amounts.
Aqueous solutions are solids, liquids or gases dissolved
in water.
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Mixtures vs. Pure Substances
2-25
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Water and Life

The following properties of water make it essential
for life:
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High surface tension
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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
2-26
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Water and Life
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Unusual density properties
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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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Chemical Reactions

A chemical change:
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When the bonds of compounds are made or
broken, new materials with new properties are
produced.
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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Chemical Reactions and Energy

Chemical reactions produce new compounds
with less or more potential energy.
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Energy is released when compounds are made with
less potential energy.
Energy is used to make compounds with more
potential energy.
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Chemical Equations

A chemical equation is a method of
describing what happens in a chemical
reaction.
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For example, photosynthesis is described by the
following equation:
Energy + 6CO2 + 6H2O → C6H12O6 + 6H2O
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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.
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Five Important Chemical Reactions
in Biology
1.
2.
3.
4.
5.
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Oxidation–reduction
Dehydration synthesis
Hydrolysis
Phosphorylation
Acid–base reactions
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Oxidation-Reduction Reactions

Oxidation-reduction reactions
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Reactions in which electrons (and their energy) are transferred from
one atom to another.
Oxidation
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Reduction
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An atom loses an electron.
An atom gains an electron.
For oxidation to occur, reduction must also occur.
Example:
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Respiration
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
Sugar is oxidized to form carbon dioxide and oxygen is reduced to form
water.
Energy is released in the process.
C6H12O6 + 6O2 → 6H2O + 6CO2+ Energy
Sugar + oxygen → water+ carbon dioxide + energy
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Dehydration Synthesis Reaction

When two small molecules are joined to form a
larger molecule,
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A molecule of water is released.
Example:
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Joining amino acids to form proteins.
NH2CH2CO-OH + H-NH CH2CO-OH  NH2CH2CO-NH CH2CO-OH + H-OH
amino acid 1
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+ amino acid 2
= protein
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+ water
Hydrolysis Reactions

When a larger molecule is broken down into
smaller parts,
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A water molecule is split
Opposite of a dehydration synthesis
Example:
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Digesting proteins into amino acids.
NH2CH2CO-NH CH2CO-OH + H-OH  NH2CH2CO-OH + H-NH CH2CO-OH
Protein
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+ water = amino acid 1 + amino acid 2
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Phosphorylation Reactions

When phosphate groups are added to other molecules,
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Bonds between phosphate groups and other molecules
contain high potential energy.
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Phosphate groups are clusters of oxygen and phosphate
atoms.
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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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.
H+Cl-
+ Na+OH- → Na+Cl- + H+OH-
Hydrocloric + Sodium
acid
hydroxide
2-36
Sodium
chloride
+ Water
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Acids, Bases and Salts

An acid
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A base
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Compounds that release hydroxide ions (OH-) into a solution
Sodium hydroxide, ammonia
Because bases are negatively charged, they will react
with a positively charged hydrogen in solution.
The strength of an acid or base is determined by how
completely it will dissociate in water.
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Ionic compounds that release hydrogen ions (H+) into a solution
Phosphoric acid, hydrochloric acid
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
2-38
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Salts


Neither acids nor bases
Usually formed when acids and bases react
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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
2-39
Sodium
hydroxide
Na+Cl- + H+OHSodium
chloride
+ Water
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pH
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
A measure of hydrogen ion concentration
Solutions with high hydrogen ion concentrations
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Solutions with low hydrogen ion concentrations
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Have a high pH
Are basic
There is a 10-fold difference in hydrogen ion
concentration between solutions that differ by one
pH unit.
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2-40
Have low pH
Are acidic
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
2-41
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