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TWIRTEENTH EDITION
Enger • Ross • Bailey
CHAPTER 2
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Chapter opener 02
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2.1 Matter, energy and life


Matter is anything that has mass and occupies space.
Energy is the ability to do work.
– There are two types of energy:
– Potential energy
 Stored energy, available to do work
– Kinetic energy
 Energy of motion
– Potential energy can be converted to kinetic energy
to do work.
– A basic understanding of chemistry will help you to
understand living things (See Fig. 2.1)
Fig. 2.1
Law of conservation of energy

Energy is never created or destroyed.
–

Energy can be converted from one form to
another, but the total energy remains
constant.
–
–
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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.
Forms of energy

There are five forms of energy:
1.
Mechanical energy
• Energy of movement (See next slides)
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
• Energy in chemical bonds
Figure 2_02a
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Figure 2_02b
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2.2 What is the nature of matter?


Atoms
– The smallest units of matter that can exist
separately.
Elements
– 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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Text art 2_01
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Atomic structure (I)

Atoms are composed
–
The atomic nucleus

Protons - positively
charged
Atomic number-the
number of protons
– All atoms of the same
element have the
same # of protons.
–

–
Electrons

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Neutrons – no charge
Orbit the nucleus in
energy levels
Are constantly in
motion
Atomic structure (II)
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Elements


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Atoms of the same element have equal number of
electrons and protons.
– Thus, they have a neutral charge.
Isotopes
– 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
– The sum of protons and neutrons in the nucleus.
Isotopes of hydrogen
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Electrons

Electrons occupy specific energy levels around the
nucleus.
–

Energy levels hold specific numbers of electrons.
–
–

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.
Electrons
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2.3 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.
–
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The formula of a compound describes the nature
and proportions of the elements that comprise
the compound.
 H20
2.4 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
speed up, and the temperature will rise.
2.5 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
it’s 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
Figure 2_06c
2.6 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:
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
Ionic bonds
Covalent bonds
Ionic bonds

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Atoms can gain or lose electrons to achieve a full
outermost energy level.
– 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.
An ionic bond
– The attraction between oppositely charged ions
Example: NaCl
– 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.
Ion formation
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How Science Works 2.2 (p. 33)
Text art 02_05
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.
–
–
The atoms sharing electrons sit close enough together so that
their outer energy levels overlap.
Single covalent bond-one pair of electrons is shared.

–
Double covalent bond-two pairs of electrons are shared.

–
Ethylene (See p. 33 for bond form)
Triple covalent bond-three pairs of electrons are shared.
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H2 (See p. 32 for bond form)
N2 (See p. 33 for bond form)
Covalent bonds
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Fig. 2.10 Ethylene (C2H4) and the Ripening Process
2.7 Water: The essence of life

Water has special properties that make it an
essential molecule for life.
–
–
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.
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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This attraction is a hydrogen bond.
Since they do not hold atoms together, they are not
considered true chemical bonds.
Hydrogen bonds are very important in
biology.
–
They stabilize the structure of DNA and proteins.
–
Water molecules can “stick” together with hydrogen bonds.
Hydrogen bonds
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Outlooks 2.1 Water and Life (I) ( p. 36)
Water and life (II)

The following properties of water make it
essential for life:
–
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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
 It requires a lot of heat to break the hydrogen
bonds holding water together.
 Large bodies of water absorb a lot of heat.
– Temperate climates
– Evaporative cooling
Water and life (III)
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–
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Unusual density properties
 Ice is less dense than water, so ice floats.
 Allows aquatic life to survive in cold climates.
The universal solvent
 Water can form hydrogen bonds with any polar
or ionic compound.
 Therefore, many things can be dissolved in
water.
Mixtures and solutions

A mixture
–
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A solution is a homogeneous mixture of ions or molecules
of two or more substances.
–
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–
–
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matter that contains two or more substances that are not in set
proportions.
Components are distributed equally throughout.
The process of making a solution is called dissolving.
The solvent is the substance present in the largest amount.
 Frequently a liquid
The solutes are the substances present in smaller amounts.
Aqueous solutions are solids, liquids or gases dissolved in
water.
Mixtures vs. Pure substances
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2.8 Chemical reactions
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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.
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.
Chemical equations

A chemical equation is a method of describing what
happens in a chemical reaction.
–
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.
Five important chemical reactions in
biology
1.
2.
3.
4.
5.
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Oxidation–reduction
Dehydration synthesis
Hydrolysis
Phosphorylation
Acid–base reactions
Oxidation-reduction reactions
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Oxidation-reduction reactions
–
–
–
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reactions in which electrons (and their energy) are transferred
from one atom to another.
Oxidation

An atom loses an electron.

An atom gains an electron.
Reduction
For oxidation to occur, reduction must also occur.
Example:
–
Respiration


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
Dehydration synthesis reaction


When two small molecules are joined to form a larger
molecule,
– a molecule of water is released.
Example:
– Joining amino acids to form proteins.
NH2CH2CO-OH + H-NH CH2CO-OH  NH2CH2CO-NH CH2CO-OH + H-OH
amino acid 1 + amino acid 2
= protein
+ water
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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:
–
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
Phosphorylation reactions (See p.37)

When phosphate groups are added to other
molecules,
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phosphate groups are clusters of oxygen and
phosphate atoms. (see p. 37)
Bonds between phosphate groups and other
molecules contain high potential energy.
–
–
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
Fig. 2.13 Phosphorylation and Muscle Contractions
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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- →
Hydrocloric + Sodium
acid
hydroxide
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Na+Cl- + H+OHSodium + Water
chloride
Acids, bases and salts

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An acid
– ionic compounds that release hydrogen ions (H+) into
a solution.
– Phosphoric acid, hydrochloric acid
A base
– 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.
– Strong acids release almost all of their hydrogen ions
into water.
– Strong bases release almost all of their hydroxide ions
into water.
Salts
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Neither acids or 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- →
Na+Cl- + H+OH-
Hydrocloric + Sodium
acid
hydroxide
Sodium + Water
chloride
Some common acids, bases and salts
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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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have low pH.
are acidic.
A solution with pH 4 has ten times as many hydrogen ions
as a solution with pH 5.
The pH scale
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