Chapter 2 Chemistry Comes Alive!
J.F. Thompson, Ph.D.
Definitions: Matter and Mass
Matter
• All living and nonliving things consist of matter


matter - anything which occupies space and has mass
liquid, solid or gas (the phases which matter takes)
• Mass vs. Weight


mass is the actual amount of matter an object contains
weight is a measure of the force of gravity on a mass
o
o
mass is same everywhere
weight varies with the position of an object on or above earth
Definitions: Energy
Energy - the capacity to do work
• Kinds of Energy
 potential energy - inactive or stored
 kinetic energy - energy of motion
• Forms of energy
 chemical energy – the making or breaking of chemical
bonds
 electrical energy – the movement of charged particles
 mechanical energy - movement of matter
 radiant energy – light or other electromagnetic radiation
• Energy conversion is inefficient, with “waste heat”
always lost
Composition of Matter
Elements and Atoms
• Element


a substance which cannot be split into simpler
substances by ordinary chemical reactions
composed of specific types of atoms
• 112 elements (at least)



92 occur naturally in nature
24 occur naturally in the body
most common: H, C, O, N
Composition of Atoms
 Atomic structure
• Protons, neutrons, electrons
 protons positively charged
 neutrons neutral
 electrons negatively charged
• The same number of protons
and electrons are present in
an atom – each atom is
electrically neutral
• Protons and neutrons are
found in the nucleus;
electrons orbit the nucleus
Atomic Number
 The number of protons in the nucleus is the element’s
atomic number
 This is the large number on the periodic table
 The number of protons in the nucleus makes atoms of
one element differ from the atoms of other elements
 Hydrogen: 1 proton, helium: 2 protons, carbon: 6
protons, etc.
How Atoms Are Combined
Molecules and Compounds
• Molecule - 2 or more atoms held together by
chemical bonds (they may be the same element)
• Compound
 Molecules which can be broken down into 2 or more
different elements by chemical means, e.g., H2O
 Demonstrates new characteristics of the combined atoms
which may be quite different that the characteristics of
the individual elements, e.g., H and O are gases at
temperatures at which H2O is a liquid or solid
Mixtures
 Substances composed of two or more components physically
intermixed; they may be in a gas, a liquid or a solid phase.
• Solutions – e.g., salt water and blood plasma
 Homogeneous with each molecule a separate entity in the mix
 Mixture of two or more components (types of molecules)
Solvent – the medium in which the other materials are mixed
o Solute – the various other components in the mix
Colloids – e.g., Jell-O and cell cytoplasm
 heterogeneous mixtures that appear translucent or milky
 smaller particles, clumps of similar molecules, which scatter light
 the particles will not settle out because of gravity
Suspensions - blood (red blood cells suspended in plasma)
 heterogeneous mixture with larger, often visible particles
 particles will settle out due to gravity if the suspension is not in
motion
o
•
•
Mixtures
Mixtures
Concentration – A measure of the
amount of solute dissolved in a
solvent
o Milligrams/deciliter
(mg/dl)
o Molarity (M )– moles/liter
- A mole is the total atomic weight of a
molecule
- Example:
- The atomic weight of glucose is 180.156.
- A 1M solution of glucose would contain 180.156
g of glucose in 1 L of solvent
Types of Chemical Bonds
From Strongest to Weakest:
• Covalent
• Ionic
• Hydrogen
• van der Walls = hydrophobic
Covalent Bonds
 Electrons are shared
among two or more
atoms
 Carbon + 4 Hydrogen
= Methane
• carbon - 4 valence
•
electrons to share
hydrogen - 1 valence
electron to share
 Single, double or
triple bonds refer to
the number of
electrons shared
between 2 atoms
Types of Covalent Bonds
 Polar and nonpolar
molecules
• If electrons are shared
equally = nonpolar
covalent bond
• Carbon dioxide has two
nonpolar covalent bonds
• If electrons are shared
unequally = polar
covalent bond
• Water is a dipole because
it has two polar covalent
bonds
Ionic Bonds
Electrons can be completely
separated and are transferred
from one atom to another –
forming positive and negative
ions
Cations: Electron donors that
have a net positive charge
Anions: Electron acceptors that
have a net negative charge
Oppositely charged ions attract
each other: ionic bonds
Ionic Bonds
 Example: table salt (sodium chloride)
• Sodium atom has one electron in its outer valence shell it can lose 1
electron for greater stability
• Chlorine has 7 electrons in its outer shell – it needs 1 more for greater
stability
• Sodium transfers one electron to chlorine

Cl- becomes an anion and Na+ becomes an cation
• Ions are created and they attract each other forming an ionic bond
Compare Covalent and Ionic Bonds
molecular stability
Hydrogen Bonds
 Hydrogen Bonds
• Weak attractions between
different molecules which
contain polar covalent
bonds
• Covalently bound
hydrogen attracts other
charged atoms
 Water
• Forms many H-bonds
• Surface tension
• Very important to water’s
properties which support
life
Chemical Reactions
The formation, rearrangement, or breaking of
chemical bonds
Energy is transferred in the process; waste
heat is lost
Chemical equations - written to describe the
chemical reactions involved:
A+B→C
reactants
product
Synthesis = Anabolic Reactions
 Atoms or molecules combine to form larger, more
complex molecules
• Always involves some type of chemical bond formation
• Always involves the input of chemical energy
• A+B  C
• Amino acids  protein
Decomposition = Catabolic Reactions
 Larger molecule is broken down into smaller molecules
or constituent atoms
• Chemical bonds are broken
• Chemical energy is always released, but that energy may or
may not be useful for another purpose
•D  E + E + E …+ E
• Glycogen (animal starch)  glucose
Chemical Exchange Reactions
 Both synthesis and decomposition occur in the same
reaction
• Chemical bonds are made and broken; molecular parts are
exchanged
• AB + CD ⇌ AD + CB
• ATP + Glucose ⇌ Glucose-6-phosphate + ADP
Oxidation-Reduction (redox) Reactions
• Electrons are exchanged between reactants
• Hydrogen ions often follow the electrons in the reaction
• Reactant losing electron(s) [& H+ ions]
 Electron Donor
 Oxidized
 Energy is released
 Decreased in potential (useful chemical) energy of the
oxidized molecule
• Reactant gaining electron(s) [& H+ ions]
 Electron Acceptor
 Reduced (its positive charge is reduced)
 Energy is absorbed
 Increased in potential (useful chemical) energy of the
reduced molecule
Energy Flow in Chemical Reactions
• Exergonic Reactions - release energy
 the products have less energy in their chemical bonds
than the reactants
 the reaction releases chemical energy, which may be
usable
 the reaction releases waste heat = exothermic
• Endergonic Reactions - require energy
 products have more potential energy in bonds than
reactants
 the reaction requires energy input = endothermic
 often coupled to exergonic reactions which provide the
energy to drive the endergonic reaction
Reversibility of Chemical Reactions
All chemical reactions are theoretically
reversible
 more difficult when energy is released during reaction
 energy must be added back to reverse the reaction
All chemical reactions tend toward an
equilibrium state
 the rates of the forward and reverse reactions are equal
 the products and reactants are in balanced proportions
 a change in energy relationships or a change in the
concentrations of reactants or products will cause a
proportionate shift in the other components of the
reaction
Factors Influencing Rate of Reactions
Molecular Collisions are required for
chemical reactions to occur
The Probability of collisions is affected by:
 Temperature - changes in kinetic energy; hotter
particles move faster
 Particle size - smaller particles move faster
 Concentration - collisions are more likely to occur
when more molecules are present
 Catalysts - bring reactants together more rapidly
or in more favorable orientations for reactions to
take place, or both
Inorganic Compounds
usually lack carbon atoms (a few exceptions,
e.g., carbon monoxide and carbon dioxide and
carbides)
usually smaller molecules
water (H2O)
gases, e.g., O2, CO2, N2
various ions = electrolytes (derived from
dietary minerals and certain salts, acids, and
bases)
Water Has Special Properties Due to its
Ability to Form Hydrogen Bonds
1. High heat capacity
•
•
Considerable kinetic energy must be added or subtracted
to change the temperature of water by a degree
Contributes to temperature stability for the body
2. High heat of vaporization
•
•
•
The change from liquid to gas phase requires breaking
many hydrogen bonds
Considerable energy is used to boil (vaporize) water
Sweating is a similar transition to gas phase that takes
heat (energy) away from the body
Water Has Special Properties
3. Polarity/solvent properties – polar covalent bonds
•
•
Water is a good solvent
Water brings solute molecules into close proximity, making
reactions more likely
4. Reactivity
•
•
Water participates in many chemical reactions
Added/removed from molecules in breakdown or synthesis
reactions
5. Lubrication and cushioning
•
•
•
Serous and Mucous fluids in many body locations lubricate
the movements of various parts
Cerebrospinal fluid cushions the brain and spinal cord
Amniotic fluid cushions the fetus
Salts
 Contain cations and
anions -- neither of
which is H+ or OH-
• Examples - NaCl, Ca2CO3
 Dissociate (ionize) into
component ions when
dissolved in water
 Ions are called
electrolytes
Acids & Bases
Compounds which break their ionic bonds when
dissolved in H2O, altering the number of
hydrogen ions in the solution
• Acids
 proton donors (H+ ion donors)
 dissociate to release hydrogen ion(s) (H+) and anion(s)
 HCl, H2CO3 H2SO4, H3PO4
• Bases
 proton acceptors (H+ ion acceptors)
 generally dissociate into hydroxyl ion (OH-) and one or
more cations
 NaOH, Ca(OH)2, NH3OH (ammonia)
pH Scale
 Concentration of hydrogen ions
is measured in pH units
 pH ranges from 0 to 14
• Based on the concentration of H+
ions = [H+] in solution
• Neutral pH is 7, [H+] = [OH-]
• A 1 pH unit change = a 10x
change in [H+]
• Below pH 7, solutions are acidic;
have more H+ ions
• Above pH 7, solutions are basic;
have more OH- ions
Acid-Base Balance
 It is very important for internal homeostasis
to keep acids/bases in balance in body fluids
 Biochemical reactions are very sensitive to pH
 pH maintained at 7.35-7.45
 Humans use 3 homeostatic mechanisms to
regulate the acid-base balance
1. Buffers
2. Respiratory System
3. Renal System
Acids & Bases
Neutralization
• Acids and bases react to form salts
• HCl + NaOH  H2O + NaCl
Buffers – resist changes in pH
• Common body buffers: HCO3-, HPO4-, proteins
• The renal and respiratory systems control pH by
regulating buffer concentrations
• H2O + CO2  H2CO3  H+ + HCO3-
 the bicarbonate buffer system is especially important
General Properties of Organic
Compounds
 ALWAYS contain C, H, and generally O
 S, N, P and a few other elements may be present
 Carbon can form 4 covalent bonds
 Carbon atoms can form chains with other carbons
 Polar organic compounds dissolve well in water
 Non-polar organic compounds do not dissolve well in
water
 Covalent bonds in organic compounds may be broken
to release useful chemical energy
• adenosine triphosphate (ATP) transfers this energy
General Properties of Organic Compounds
Organic molecules may be quite large
Organic molecules make good structural
components
Simple molecules = monomers are joined to
form macromolecules = polymers
Four classes of biologically important polymers:
• Carbohydrates
• Lipids
• Proteins
• Nucleic Acids
Carbohydrates
Sugars, starch, glycogen, cellulose, chitin
Classified by size
Chemical formula = [CH2O]N (multiples of
CH2O)
 e.g., C6H12O6 = glucose
Carbohydrates have many functions
 Structural components of molecules (e.g., DNA, RNA),
cells and tissues
o
o
cellulose is the most abundant organic substance on earth
we cannot digest it, but it is an important part of our diet
 Broken down for chemical energy production; often ATP
formation
 Chemical energy is stored as glycogen for future use
Monosaccharides
 Monomers = simple sugars = monosaccharides
 Single chain or ring structures
• glucose – the preferred source of energy in the body
• fructose
• galactose
• deoxyribose
• ribose
Disaccharides
 2 monosaccharides joined together
• dehydration synthesis – a water molecule is removed
during bond formation
• hydrolysis - add a water molecule to liberate the
monosaccharides
• examples: sucrose, maltose, lactose
Polysaccharides
 Large, relatively insoluble carbohydrate polymers
 10’s or 100’s of monosaccharides bonded together
• glycogen - storage of glucose for energy in humans
 in liver cells, skeletal muscle cells and endometrial lining cells
• starch – storage of glucose for energy in plants
• cellulose – structural fiber in plants; “roughage” in our diet
• chitin – structural fiber in some fungi, insects, crustaceans
Lipids = Fats
Contain carbons, hydrogens, oxygens (fewer
oxygens per carbon)
Roughly 18-25% of body weight
Fewer polar covalent bonds
Insoluble in water = hydrophobic
Dissolve in lipid solvents, e.g., alcohols,
detergents, oils
Major energy storage compounds
Structural components of cell membranes
Some serve as molecular signal compounds
• steroid hormones, prostaglandins, etc.
Triglycerides
 Storage form of fats for energy
 Non-polar, hydrophobic
 Glycerol head and 3 fatty acid side-chains
 Saturated fatty acids vs. unsaturated fatty acids
• double bonds in unsaturated fatty acids
• the more unsaturated bonds, the more fluid at room temp.
Note dehydration synthesis
Phospholipids
 Modified triglyceride
• 2 fatty acids and a phosphate group attached to glycerol
 Phosphate group is a polar group
• dipole
• amphipathic - has polar and non-polar regions
• can hydrogen bond with water due to phosphate’s polarity
 Found in cell membranes
hydrophilic head
hydrophobic tail
Other Lipid Compounds
 Steroids
• 4 rings of carbon atoms
• Non-polar
• Fat soluble/hydrophobic
• Cholesterol is a membrane
structural component
• Cholesterol-derived steroid
hormones
 Eicosanoids (Local
Hormones)
• Prostaglandins, leukotrienes
• Regulate blood clotting,
inflammation, immune
reactions, labor, etc.
3
1
2
4
Proteins




contain carbon, hydrogen, oxygen, nitrogen, sulfur
12 - 18% of body mass
a wide variety of structural and functional roles
amino acids are the building blocks of proteins
• 20 different AA's
• each AA has a hydrogen (H), an amino group (NH2), carboxyl group
•
(CO2H), and side group (R)
some R groups (amino, carboxyl groups) ionize; these buffer body fluids
Peptide Bond Formation
 amino acids (monomers) combine in a linear sequence
to form a polypeptide (polymer)
 AA's covalently bonded together by “peptide bonds”
between carboxyl and amino groups
 a water molecule is created as well
 form dipeptides  tripeptides  polypeptides
Hierarchy of Protein Structure
 Primary – amino acid sequence
 Secondary - due to hydrogen
bonds between adjacent amino
acids
• random coil,
sheet
α-helix, or ß-pleated
 Tertiary – due to various types of
bonds between more distant amino
acids
• produces a unique 3D shape, for each
•
protein, and causes a unique function
fibrous vs. globular shapes
 Quaternary - 2 or more
polypeptide chains bound together
to form a functional whole protein
Protein Denaturation
 Protein structure is generally
stable
 Strong structure/function
relationship is dependent on
the 3D shape of the protein
 High temperatures and
strong pH alter proteins
 Structural changes reduce or
completely stop activity
 Structural changes are
irreversible at temperature
and pH extremes
Enzymes
Most have names ending with –ase
Globular proteins that are organic catalysts
• apoenzyme (protein portion)
• cofactors and coenzymes may assist in the enzymecatalyzed reaction
Enzyme-catazyzed reactions are thousands to
millions times more likely than reactions caused
by random molecular collisions
Enzymes are highly specific for their substrates
and the reactions they catalyze
Enzyme Action – Catalytic Mechanism
 Activation energy – the
energy needed to trigger
a reaction to occur
 Enzymes (catalysts)
decrease the activation
energy required
• increase interactions
between substrates
• without increasing
temperature or pressure
• properly orient molecules
Enzyme Action and Regulation
 Three basic steps
1. bind at the active site to
form an enzymesubstrate complex
2. cause internal bond
rearrangements
3. release the product(s) &
repeat
 Many regulatory
controls
•
•
Substrate specificity
Many cellular regulation
systems (feedback
controls)
Nucleic Acids
 Deoxyribonucleic Acid (DNA) & Ribonucleic acid (RNA)
 DNA, RNA made of nucleotides monomers- 3 parts to
a nucleotide
• nitrogenous bases: adenine, thymine (DNA only), cytosine,
guanine, uracil (RNA only)
• pentose sugar (5 carbon): deoxyribose or ribose
• phosphate group
Nucleic Acids
 DNA is the Genetic Material
 4 nucleotide monomers
• adenine (A), guanine (G), cytosine
(C), thymine (T)
 complimentay bases
connect the two strands:
• A-T, C-G
 forms a double helix, i.e., a
double stranded coil
 RNA carries hereditary
information from nuclear
DNA to the cytoplasm
(inside cells)
• uracil (U) replaces T
• single stranded
Adenosine Triphosphate
 three phosphate groups,
attached to a ribose
sugar and adenine
high energy bonds
(adenosine nucleotide)
 ATP
 ADP + Pi
 high-energy phosphate
groups are attached by
aerobic or anaerobic
catabolic reactions
 ATP is the major cellular
energy transfer
compound
End Chapter 2
Note: You will find additional slides after this “end”
slide with additional details of chemistry processes,
which you may review on your own.
Mass Number
 Mass number = total number of protons + neutrons
• Mass number may vary among the atoms of an element
because of different numbers of neutrons (isotopes)
• All isotopes of an element have:
 the same number of protons
 the same chemical characteristics
• Radioactive isotopes are unstable isotopes which “decay”
into other isotopes, even into other types of elements
Identifying Elements
 Isotopes of an element:
• Have different numbers of neutrons
• Have the same number of protons
• Have the same chemical characteristics
• The mass number indicates the number of protons &
neutrons
• Radioactive isotopes are unstable and “decay” into other
isotopes, even into other types of elements
The Role of Electrons
 Electrons form clouds called
shells
• Each shell contains one or more
orbitals
• The first three shells hold 2, 8, and
18 electrons, respectively
 The outer shell is the valence shell
• Atoms are stable (inert) when the
valence shell is filled with electrons
• Atoms are chemically reactive if the
valence shell is not full of electrons
Composition of Atoms
 Electron structure
• Electrons orbit nucleus in
shells
• The first three shells hold 2, 8, and
18 electrons, respectively
• Shells contain subshells
• The outer shell is the valence shell
• Atoms are the most stable when
the valence shell is filled with
electrons
End Chapter 2
Note: The following slides refer you to specific exercises you
may access on-line at your textbook publisher’s website, if
you wish additional review of the chemistry from CH 2. This
same suggested exercises can be found on a page linked
from the Exam 1 Review page on Dr. T’s webpages.
Anatomy and Physiology Place Assignment
 PART I: BASIC CHEMISTRY
 Definition of Concepts: Matter and Energy (25-26)
•
Animation: Energy Concepts
•
Animation: Atomic Structure How Matter Is Combined: Molecules and Mixtures
(30-31)
 Composition of Matter: Atoms and Elements (27-30)




Chemical Bonds (31-36)
Chemical Reactions (36-40)
Part II: BIOCHEMISTRY
Inorganic Compounds (40-43)
•
InterActive Physiology®: Introduction to Body Fluids Organic Compounds (4359)
 Animations: Disaccharides | Polysaccharides | Fats
Art Labeling: Lipids (fig. 2.15, p. 47)
Animations: Structure of Proteins | Primary and Secondary Structure |
Tertiary and Quaternary Structure
Animation: How Enzymes Work
Art Labeling: Mechanism of Enzyme Action (fig. 2.21, p. 55)
Art Labeling: Structure of DNA (fig. 2.22, p. 56)
Memory: Important Molecules Chapter Summary
Anatomy and Physiology Place
 Chapter Quizzes
Art Labeling Quiz
Matching Quiz
Multiple-Choice Quiz (Level I)
Multiple-Choice Quiz (Level II)
True-False Quiz
Crossword Puzzles
Crossword Puzzle 2.1
Crossword Puzzle 2.2
Crossword Puzzle 2.3
Get Ready for A&P Activities
 Your Starting Point: Pre-Quiz
 Atomic Structure 165
• Atoms and Isotopes
• Build an Atom
 Period Table of Elements 171
• Periodic Table
 Chemical Bonding 177
• Hydrogen Bonding
Nonpolar and Polar Molecules
 What Did You Learn? Post Quiz
End Chapter 2