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Chapter 2
Chemistry
Matter


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Occupies space and has mass
Mass does not depend on gravity,
weight does
Three states of matter
– Solid – definite shape and volume
– Liquid – indefinite shape, definite volume
– Gas – indefinite shape and volume
Energy
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Capacity to do work (put matter into
motion)
Energy cannot be created or destroyed
Does not take up space, has no mass
Measured by its effects on matter
Two types of energy
– Kinetic energy – the energy of motion
– Potential energy – stored energy
Forms of Energy

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Chemical energy – stored in chemical
bonds, especially adenosine triphosphate
(ATP)
Electrical energy – results from movement
of charged particles
Mechanical energy – directly moves matter
Radiant energy – waveform energy, includes
light waves, radio waves, UV waves, and X
rays
Energy Form Conversions
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Energy is constantly converted to
different forms
During conversion, some energy is
“lost” as heat
Atoms and Elements
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Elements CANNOT be broken down
112 different elements, 92 are natural
Main elements in the body = carbon,
oxygen, hydrogen, nitrogen
Atoms are building blocks of elements
Atoms and Elements
(cont.)

Atoms give each element its special
properties
– Physical properties – detectable by our senses
(color, texture, boiling point)
– Chemical properties – interaction of atoms with
other atoms (some elements create rust, some
can be digested, etc.)

Atomic symbol = chemical abbreviation of
element (C = carbon, Ca = calcium, Na =
sodium)
Atomic Structure

The atom has a central nucleus,
containing protons and neutrons
– Protons (p+) – positively charged, have a
mass of 1 atomic mass unit (1 amu)
– Neutrons (n0) – uncharged, have a mass
of 1 amu
– Most of the atom’s mass is in the nucleus
Atomic Structure (cont.)

Electrons orbit the nucleus
– Electrons (e-) – negatively charged, only
1/2000 the mass of a proton (0 amu)

Atoms are electrically neutral, because
number of protons = number of
electrons
Identification of Elements:
Atomic Structure
Figure 2.2
Identifying Elements
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Atomic Number = number of protons
Mass Number = sum of masses of protons
and neutrons in an element
Isotopes = structural variation of an
element, contains different amount of
neutrons
Atomic Weight = average of mass numbers
of all isotopes of an element; dependent on
abundance of the isotopes
Identification of Elements:
Isotopes of Hydrogen
Figure 2.3
Molecules

Molecule = combination of two or more
atoms joined by chemical bonds
– Compounds = combination of two or more
different atoms joined by chemical bonds

Molecules are the smallest parts of
compounds that display characteristics of
that compound
Mixtures

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Mixture = made of two or more mixed
components which are not chemically
bonded
Three types of mixtures are solutions,
colloids, and suspensions
Solutions
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Homogenous mixture; may be gas, liquid,
or solid
Contain solvent (greater amount) and solute
(lesser amount)
Transparent and do not scatter light
Components do not settle out
Examples include salt water and sugar
water
Concentration of
Solutions
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Can be indicated by the percent of
solute in total solution
Can be indicated by molarity (moles/L)
Mole = molecular weight of a
compound weighed out in grams
Avogadro’s number – one mole of any
substance always contains 6.02 X 1023
molecules
Concentration of
Solutions (cont.)
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Glucose = C6H12O6
C = 12.011, H = 1.008, O = 15.999
Total atomic weight = 180.156
To make a 1 molar solution of glucose,
weight out 180.156 grams of glucose, and
add water up to 1 L
Colloids (Emulsions)
– Heterogeneous mixture
– Do not settle out
– Milky in appearance
– Will scatter light
– Can convert from a fluid to solid state
– Examples include gelatin and cytosol
Suspensions
– Heterogeneous mixture
– Visible solutes that settle out
– Will scatter light
– Examples include sand mixed with water
and blood
Comparison of Mixtures
and Compounds
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No chemical bonding in mixtures
Mixtures can be separated by physical
means
Compounds require chemical means to
be separated
All compounds are homogenous,
mixtures can be homogenous or
heterogenous
Chemical Bonds
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Created by interaction of electrons
Electrons travel around nucleus of
atom in electron shells
Atoms can have 7 shells
Electron shells are also called energy
levels
Shell 1 = 2 e-, shell 2 = 8 e-, shell 3 =
18 e-
Chemical Bonds (cont.)
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Shells fill up in order
Only the outermost energy level is
reactive!!
Valence shell = reactive outer energy
level of an atom
If outer shell is full, that atom is not
reactive (inert)
If outer shell is not full, the atom will
react to become stable
Chemical Bonds (cont.)
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Only 8 electrons bond, no matter how
many are in a shell
Octet rule = except in shell 1, atoms
interact with each other to have eight
electrons in their valence shell
Chemically Inert and Reactive
Elements
Figure 2.4a
Figure 2.4b
Types of Chemical Bonds
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Ionic – transfer of electrons
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Covalent – sharing of electrons

Hydrogen – weak attractions
Ionic Bonds
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Ion = charged particle formed from gain or
loss of electrons
Ionic bond = electrons are transferred from
one atom to another
If an atom accepts electrons, it has a
negative charge = anion
If an atom loses electrons, it has a positive
charge = cation
Formation of an Ionic Bond
Figure 2.5a
Ionic Bonds (cont.)
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Ionic compounds are usually salts
Salts form crystals
Figure
2.5b
Covalent Bonds
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Covalent bond = sharing of pairs of
electrons between two or more atoms
Shared electrons will orbit the whole
molecule
Single Covalent Bonds
Figure 2.7a
Double Covalent Bonds
Figure 2.7b
Triple Covalent Bonds
Figure 2.7c
Polar and Nonpolar
Covalent Bonds

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Nonpolar molecules are formed from
equal electron sharing (symmetrical)
Polar molecules are formed from
unequal electron sharing
(nonsymmetrical)
– Electronegative atoms attract electrons
– Electropositive atoms donate electrons
Comparison of Ionic, Polar Covalent, and
Nonpolar Covalent Bonds
Figure 2.9
Hydrogen Bonds
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Formed when hydrogen is covalently
linked to one electronegative atom
and is attracted to another
electronegative atom
Hydrogen bonding is common in
water, causing surface tension
Hydrogen bonding is also found in
proteins and DNA
Hydrogen Bonds
Figure 2.10a
Chemical Reactions
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Occur when chemical bonds are
formed, rearranged, or broken
Reactions are written as chemical
equations
– Equations show reactants and products
Types of Chemical
Reactions
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Synthesis, or combination – atoms or
molecules join to form a more complex
molecule
Also called anabolic
A + B  AB
Types of Chemical
Reactions (cont.)
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Decomposition reaction = breakdown
of a molecule into smaller particles
Also called catabolic
AB  A + B
Types of Chemical
Reactions (cont.)

Exchange reactions = both synthesis
and breakdown occur
AB + C  AC + B
AB + CD  AD + CB
Oxidation-Reduction
Reactions
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Redox reactions are how the body
breaks down food for energy
Electrons are transferred between
reactants
Reactant that loses electrons (electron
donor) is oxidized
Reactant that gains electrons (electron
acceptor) is reduced
Energy of Reactions

Exergonic reactions = reactions that release
energy
– Products have less energy than reactants
– Energy produced can be used

Endergonic reactions = reactions that
absorb energy
– Products have more energy than reactants
– Energy is put in to the reaction
Reversibility of Reactions

All reactions are reversible
A + B  AB
AB  A + B

Chemical equilibrium = when a
reaction proceeds equally in both
directions
Factors Influencing Rate
of Reactions

Atoms have to collide to react
– Temperature – increasing temperature speeds
up reactions
– Concentration – increasing reactant amounts will
speed up reactions
– Particle size – smaller particles move faster,
speeds up reactions
– Catalysts – increase rate of reactions without
being changed themselves
Biochemistry

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Study of chemistry as related to living
matter
Organic compounds
– Contain carbon
– Have covalent bonds

Inorganic compounds
– Waters, salts, acids, and bases
– Do not contain carbon
Water
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1.
2.
60-80% of all cells
High heat capacity – can absorb and
release large amounts of heat
without large temperature change
High heat of vaporization – when
water evaporates, a lot of heat is
released (perspiration)
Water (cont.)
3.
Polar solvent properties
– Universal solvent
– Ions will dissociate in water
– Water surrounds larger molecules,
keeps them from settling out of solution
– Major transport medium of the body
Water (cont.)
4.
Reactivity
– Water plays a part in many chemical
reactions
– Decomposition using water = hydrolysis
– Synthesis using water = dehydration
synthesis
5.
Cushioning – protects internal organs
Salts
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Inorganic compounds
Contains cations other than H+ and
anions other than OHSalts dissolve in water to create
electrolytes
Common salts in body = NaCl, calcium
phosphates
Acids and Bases
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Dissociate into ions in water
Acid – releases hydrogen ions in solution, proton
donor (HCl, H2CO3)
HCl  H+ + Cl –

Base – releases hydroxide ions in solution, proton
acceptor (NaOH, HCO3-, NH3)
NaOH  Na+ + OH–
pH
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More H+ free in solution, more acidic
More OH- free in solution, more basic
(alkaline)
Neutral solutions, H+ = OHConcentration of H+ in solution = pH
Acid-Base
Concentration (pH)
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Acidic: pH 0–6.99
Basic: pH 7.01–14
Neutral: pH 7.00
pH change of 1
represents 10-fold
change in [H+]
Figure 2.13
Neutralization

Mixture of acids and bases in solution
causes formation of water and salt
HCl + NaOH
NaCl + H2O
Buffers
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Prevent large changes in pH of body
fluids
Release H+ when pH rises
Bind H+ when pH falls

pH of a solution due only to free H+
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Strong acids and bases = completely
dissociate in water
Buffers (cont.)
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Weak acids and bases = do not
dissociate completely in water, so less
effect on pH of solution
Weak acids will bind H+ if acid is
added, and will release H+ if base is
added
Weak bases will bind OH- if base is
added, and will bind more H+ if acid is
added
Organic Compounds
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Contain carbon
Carbon is electroneutral (always shares
electrons), and can form 4 bonds
Four types of organic compounds
– Carbohydrates
– Lipids
– Proteins
– Nucleic acids
Carbohydrates
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Sugars and starches
Contain carbon, hydrogen, and oxygen
2:1 ratio of hydrogen and oxygen
Larger carbohydrates are less soluble
in water
Carbohydrates (cont.)
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Easily available source of fuel
Classes of carbohydrates
– Monosaccharides
– Disaccharides
– Polysaccharides
Monosaccharides
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Building blocks of carbohydrates
Single rings or chains with 3-7 carbons
General formula = (CH2O)n
Most important in body are pentoses
and hexoses
Isomers = same chemical formula,
different arrangement of atoms
Monosaccharides
Figure 2.14a
Disaccharides
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Formed from joining of two
monosaccharides by dehydration
synthesis
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Water is released in this reaction
Must be broken down by hydrolysis to
be absorbed from intestines
Disaccharides
Figure 2.14b
Polysaccharides
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Long chains of monosaccharides
Important storage molecules
Two main ones important for humans
= starch and glycogen
Figure 2.14c
Lipids
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Do not dissolve in water
Contain C, H, and O, but the proportion of
oxygen in lipids is less than in carbohydrates
Also can contain phosphorus
Examples:
–
–
–
–
Neutral fats or triglycerides
Phospholipids
Steroids
Eicosanoids
Triglycerides
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Composed of three fatty acids bonded to a
glycerol molecule
Fats when solid, oils when liquid
Nonpolar molecules
Figure 2.15a
Triglycerides (cont.)
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Saturated = fatty acid with single
bonds only
– Straight hydrocarbon chains
– Solid form

Unsaturated = fatty acid with double
bonds between carbons
– Kinks in hydrocarbon chains
– Oil form at room temperature
Phospholipids
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Phospholipids – modified triglycerides with
two fatty acid groups and a phosphorus
group
Phosphate group is polar, rest is nonpolar
= amphipathic or amphiphilic
Figure 2.15b
Other Lipids
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Steroids – flat molecules with four interlocking
hydrocarbon rings
Eicosanoids – 20-carbon fatty acids found in cell
membranes
Figure 2.15c
Functions of Lipids
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Triglycerides – insulation
Phospholipids – form cell membranes
Steroids – part of cell membranes,
also used in hormone production
Eicosanoids – prostaglandins most
important; used for blood clotting and
inflammation
Proteins
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Amino acids are building blocks of
protein
Contain C, H, O, and N
Amino group NH2
Carboxyl group COOH
Amino acids can be acidic or basic
Amino Acids
Figure 2.16a–c
Protein

Macromolecules composed of
combinations of 20 types of amino acids
bound together with peptide bonds
Peptide bond
H
H
R
O
N
C
C
OH
H
Amino acid
+
H
H
R
O
N
C
C
OH
H
Amino acid
Dehydration H O
2
synthesis
Hydrolysis
H
H2O
H
R
O
H
R
O
N
C
C
N
C
C
H
H
OH
Dipeptide
Figure 2.17
Structural Levels of
Proteins
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Primary = amino acid sequence
Secondary = twisting and bending
creates alpha helices and beta chains
Tertiary – folding of secondary
structures
Quaternary – separate polypeptide
chains linked together
Structural Levels of Proteins
Figure 2.18a–c
Structural Levels of Proteins
Figure 2.18b,d,e
Fibrous and Globular
Proteins

Fibrous proteins (structural)
– Strandlike, some have only secondary structure
– Insoluble in water, very stable
– Examples: keratin, elastin, and collagen

Globular proteins (functional)
– Spherical proteins with tertiary and quaternary
structures
– Water soluble, chemically active
– Examples: antibodies, hormones, and enzymes
Protein Denaturation
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Proteins must fold into a specific
pattern to function
Destruction of hydrogen bonds by
temperature or pH changes leads to
denaturation
Denaturation = unfolding of proteins
which causes loss of function
Protein Denaturation
(cont.)

Active sites
must be intact
for protein to
function
Figure 2.19a
Protein Denaturation

Irreversibly denatured proteins cannot
refold, and cannot function
Figure 2.19b
Enzymes
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Globular proteins; act as catalysts
Chemically specific
Speed up reactions
Usually end in –ase
Often have to be activated
Lower activation energy of reactions
Enzymes Decrease
Activation Energy
Figure 2.20
Enzyme Action
1.
2.
3.
Enzyme binds with substrates
Enzyme-substrate complex
rearranges to form products
Enzyme releases the products;
enzyme is unchanged
Active site
Amino acids
+
Enzyme (E)
Substrates (S)
Enzyme-substrate
complex (E-S)
H2O
Free enzyme (E)
Peptide bond
Internal rearrangements
leading to catalysis
Dipeptide product (P)
Figure 2.21
Nucleic Acids

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Composed of C, N, H, O, and P
Built from nucleotides
– Nucleotides have N-containing base,
pentose sugar, and a phosphate group

5 bases = adenine (A), guanine (G),
cytosine (C), thymine (T), and uracil (U)
Nucleic Acids (cont.)

Adenine and guanine are purines
– Large, have 2 rings

Cytosine, thymidine, and uracil are
pyrimidines
– Smaller, have 1 ring

Two major classes of nucleic acids =
DNA and RNA
Deoxyribonucleic Acid (DNA)

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
Found in nucleus
Contains genetic material which is the
blueprint for protein production
Sugar is deoxyribose
Replicates before cell division
Forms a double helix
Structure of DNA
Figure 2.22a
Structure of DNA
Figure 2.22b
Ribonucleic Acid (RNA)

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Single-stranded
Usually located outside of nucleus
Uses uracil instead of thymine
Important for creation of proteins
Sugar is ribose instead of deoxyribose
3 kinds: mRNA, tRNA, and rRNA
Adenosine Triphosphate (ATP)





Stores energy in phosphate bonds
Readily usable by body cells
Adenine-containing nucleotide with three
phosphate groups
Unstable
Terminal phosphate group is transferred to
other molecules – called phosphorylation
Structure of ATP
Figure 2.23
Membrane
protein
Pi
P
Solute
Solute transported
(a) Transport work
ADP
+
Pi
ATP
Relaxed smooth
muscle cell
Contracted smooth
muscle cell
(b) Mechanical work
Pi
X
P
X
Y
+ Y
Reactants
Product made
(c) Chemical work
Figure 2.24