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Chapter 2
Lecture
Outline
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The Chemistry of Life
• Atoms, Ions and Molecules
• Water and Mixtures
• Energy and Chemical
Reactions
• Organic compounds
Atoms, Ions and Molecules
•
•
•
•
•
The chemical elements
Atomic structure
Isotopes and radioactivity
Ions, electrolytes and free radicals
Molecules and chemical bonds
The Chemical Elements
• Element = simplest form of matter with
unique chemical properties
• Atomic number = # of protons in nucleus
– periodic table
• elements arranged by atomic number
– 24 elements have biological role
• 6 elements = 98.5% of body weight
• trace elements in minute amounts
Minerals
• Inorganic elements absorbed from soil
by plants
• Equals 40% of body weight
– structure (teeth, bones, etc)
– enzymes
Structure of an Atom
• Nucleus = center of atom
– protons: single (+) charge, mass = 1 amu
– neutrons: no charge, mass = 1 amu
• Electron shells surround the nucleus
– electrons: single negative charge, little mass
– valence electrons in the outermost shell
• interact with other atoms
• determine chemical behavior
Planetary Models of Elements
p+ represents protons, no represents neutrons
Isotopes and Radioactivity
• Isotopes
– differ in # of neutrons
– extra neutrons increase atomic weight
– isotopes of an element are chemically
similar
• have same valence electrons
• Atomic weight
– average atomic mass of the isotopes
• Carbon
Isotopes
• Contain the same
number of protons
and different
number of
neutrons
Radioisotopes and
Radioactivity
• Isotopes
– same chemical behavior, differ in physical behavior
– breakdown gives off radiation
• Radioisotopes
– unstable isotopes
– every element has at least one radioisotope
• Radioactivity
– radioisotopes decay to stable isotopes releasing
radiation
– we are all mildly radioactive
Marie Curie
• First woman in world to
receive a Ph.D.
• First woman to receive
Nobel Prize (1903)
– discovered radioactivity of
radium
– trained physicians in use of X
rays and radiation therapy as
cancer treatment
• Died of radiation poisoning
at 67
Ions and Ionization
• Ions - carry a charge due to an unequal
number of protons and electrons
• An ion is a charged atom
• Ionization =
transfer of
electrons from one
atom to another
• Forms an ionic
bond
• ( stability of
valence shell)
Anions and Cations
• Anion
– atom that gained electrons (net negative
charge)
• Cation
– atom that lost an electron (net positive
charge)
• Ions with opposite charges are attracted to
each other
Anions and Cations
Electrolytes
• Salts that ionize in water to form body
fluids
– capable of conducting electricity
• Electrolyte importance
– chemical reactivity
– osmotic effects (influence water movement)
– electrical effects on nerve and muscle tissue
• Imbalances cause muscle cramps, brittle
bones, coma and death
Free Radicals
• Particle with an odd number of electrons
• Produced by
– normal metabolic reactions, radiation,
chemicals
• Causes tissue damage
– reactions that destroy molecules
– causes cancer, death of heart tissue and
aging
• Antioxidants
– neutralize free radicals
– in diet (vitamin E, carotenoids, vitamin C)
Molecules and Chemical Bonds
• Molecules
– two or more atoms covalently bonded
• Compounds
– two or more atoms of different elements
covalently bonded
• Molecular formula
– elements and how many atoms of each
• Structural formula
– location of each atom
– structural isomers revealed
Water Molecule
Structural Formula
Molecular Compound
H20
Chemical Bonds
• Ionic bonds
• Covalent bonds
• Hydrogen bonds
• Van der Waals
force
Ionic Bonds
• Attraction of oppositely charged ions
• No sharing of electrons
• Weak bond (easily dissociates in water)
Covalent Bonds
• Formed by sharing of valence electrons
• Types of covalent bonds
– single = sharing of single pair electrons
– double = sharing of 2 pairs
– nonpolar
• shared electrons (equal time around each nucleus)
• strongest of all bonds
– polar
• negative charge where electrons spend most time
Single Covalent
Nonpolar Covalent
Double Covalent
Polar Covalent
Single Covalent Bond
• One pair of electrons are shared
Double covalent bonds:
Two pairs of electrons are shared each C=O bond
Nonpolar /Polar Covalent
Bonds
electrons
shared
equally
electrons
shared
unequally
Hydrogen Bonds
• Weakest bond = no sharing of electrons
• Attraction between polar molecules
– positive hydrogen atoms to negative
oxygen atoms in a 2nd molecule
• Physiological importance
– properties of water created by shapes of
large complex molecules
– determined by folding due to hydrogen
bonds
Hydrogen Bonding in Water
Van der Waals Forces
• Weak attractions between neutral atoms
• Fluctuations in electron density create
polarity
• Only 1% as strong as a covalent bond
– folding of large molecules
– significant when 2 large surfaces meet
Mixtures and Water
• Substances
physically but not
chemically combined
• Mixtures in our bodies
contain water
• Water 50-75% of body
weight
– depends on age, sex,
percentage body fat,
etc.
Solvency
• Solvency - ability to dissolve other
chemicals
– Hydrophilic (charged substances) dissolve
easily in water
– Hydrophobic (neutral substances) do not
easily dissolve in water
• Water = universal solvent
– metabolic reactions and transport of
substances
Hydrophilic & Hydrophobic
Water as a Solvent
• Polar water molecules overpower the ionic
bond in Na+Cl– forming hydration spheres around each ion
– water molecules: negative pole faces Na+,
positive pole faces Cl-
Adhesion and Cohesion
• Adhesion – tendency of one
substance to cling to
another
• Cohesion – tendency of like
molecules to cling to each
other
– water is very cohesive due to
its hydrogen bonds
– surface film on water formed
by surface tension
Thermal Stability of Water
• Water stabilizes internal temperature
– has high heat capacity
• hydrogen bonds inhibit temperature increases
by inhibiting molecular motion
– water absorbs heat without changing temperature
– effective coolant
• 1 ml of perspiration removes 500 calories
– calorie: amount of heat required to raise temperature
of 1g of water by 1°C
Solutions
• Mixture of a solute
into a solvent
• Small solute
particles
– pass through cell
membranes
• Solution transparent
• Remains mixed
Colloids
• Mixture of protein
and water
– change from liquid
to gel state within
and between cells
• Particles too large
to pass through
cell membranes
• Cloudy
• Remains mixed
Suspensions and Emulsions
• Suspension
– particles suspended in
a solvent
– particles exceed 100nm
• too large to pass through
a cell membrane
– cloudy or opaque
appearance
– separates on standing
• Emulsion
– suspension of one
liquid in another
– fat in breast milk
Acids, Bases and pH
• An acid is proton donor (releases H+ ions)
• A base is proton acceptor (accepts H+ ions)
• pH = the concentration of H+ ions in
solution
– a pH of less than 7 is acidic solution
– a pH of greater than 7 is basic solution
– a pH of 7.0 is neutral pH
pH
• pH = measurement of molarity of H+ [H+]
on a logarithmic scale
– pH = -log [H+] thus pH = - log [10-3] = 3
– a change of one number on the pH scale
represents a 10 fold change in H+
concentration
• a solution with pH of 4.0 is 10 times as acidic as
one with pH of 5.0
• Our body uses buffers to prevent change
– pH of blood ranges from 7.35 to 7.45
– tremors, paralysis or even death
pH Scale
Work and Energy
• Energy - capacity to do work
• Kinetic energy - energy of motion
– heat is kinetic energy of molecular
motion
• Potential energy- energy due to
object’s position (ions on one side
only of cell membrane)
– chemical energy - potential energy
stored in the molecular bonds
Chemical Reaction
• Process that forms or breaks an
ionic or covalent bond
• Symbolized by chemical equation
– reactants  products
Classes of reactions
• Decomposition reactions
• Synthesis reactions
• Exchange reactions
Decomposition
Reactions
• Large molecules
broken down into
smaller ones
• AB  A + B
Synthesis
Reactions
• Two or more small
molecules combine to
form a larger one
• A + B  AB
Exchange Reactions
• Two molecules collide and exchange
atoms or group of atoms
• AB+CD  ABCD 
AC + BD
Stomach acid (HCl)
and sodium
bicarbonate (NaHCO3)
from the pancreas
combine to form NaCl
and H2CO3.
Reversible Reactions
• Go in either direction (symbolized with
double-headed arrow)
• CO2 + H2O
H2CO3
HCO3- + H+
– most common equation discussed in this
book
• Law of mass action determines
direction
– side of equation with greater quantity of
reactants dominates
Reaction Rates
• Basis for reactions is molecular motion and
collisions
– reactions occur when molecules collide with enough
force and the correct orientation
• Reaction Rates affected by:
– concentration
• more concentrated, more collisions, faster rate
– temperature
• higher temperature, greater collision force, faster rate
– Catalysts (enzymes)
• speed up reactions without permanent change to itself
• holds reactant molecules in correct orientation
Metabolism
• The sum of ll the chemical reactions of the
body
• Catabolism
– energy releasing (exergonic) decomposition
reactions
• breaks covalent bonds, produces smaller molecules,
releases useful energy
• Anabolism
– energy storing (endergonic) synthesis reactions
• requires energy input
Oxidation-Reduction
Reactions
• Oxidation
– molecule gives up electrons and releases
energy
– accepting molecule is the oxidizing agent
• oxygen is often the electron acceptor
• Reduction
– molecule gains electrons and energy
– donating molecule is the reducing agent
• Oxidation-reduction (redox) reactions
– Electrons are often transferred as hydrogen
atoms
Organic Chemistry
• Study of compounds containing carbon
• 4 categories of carbon compounds
– carbohydrates
– lipids
– proteins
– nucleotides and nucleic acids
Organic Molecules and
Carbon
• Only 4 valence electrons
– bonds readily to gain more valence
electrons
• Forms long chains, branched molecules
and rings
– serve as the backbone for organic
molecules
• Carries a variety of functional groups
Functional
Groups
• Atoms attached
to carbon
backbone
• Determines
chemical
properties
Monomers and Polymers
• Macromolecules = very large molecules
• Polymers = macromolecules formed
from monomers bonded together
• Monomers = an identical or similar
subunit
Polymerization
• Bonding of monomers together to form
a polymer
• Formed by dehydration synthesis
– starch molecules are a polymer of 3000
glucose monomers
– protein molecules are a polymer of amino
acids
Dehydration Synthesis
• 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
Hydrolysis
• Splitting a polymer (lysis) by the addition of a
water molecule (hydro)
– a covalent bond is broken
• All digestion reactions consists of hydrolysis
reactions
Organic Molecules:
Carbohydrates
• Hydrophilic organic molecule
• General formula
– (CH2O)n
n = number of carbon atoms
– for glucose, n = 6, so formula is C6H12O6
– 2:1 ratio of hydrogen to oxygen
• Names of carbohydrates
– word root sacchar- or the suffix -ose often
used
• monosaccharide or glucose
Monosaccharides
• Simple sugars
• General formula is C6H12O6
– structural isomers
• Major monosaccharides
– glucose, galactose and
fructose
– produced by digestion of
complex carbohydrates
• glucose is blood sugar
Disaccharides
• Sugar molecule
composed of 2
monosaccharides
• Major disaccharides
– sucrose = table sugar
• glucose + fructose
– Lactose = sugar in milk
• glucose + galactose
– Maltose = grain products
• glucose + glucose
Polysaccharides
• Chains of glucose subunits
• Starch: energy storage in plants
– digestible by humans for energy
• Cellulose: structural molecule of plant cell
walls
– fiber in our diet
• Glycogen: energy storage in animals
– liver synthesizes after a meal and breaks down
between meals
Carbohydrate Functions
• All digested carbohydrates converted to glucose
and oxidized to make ATP
• Conjugated carbohydrate = bound to lipid or
protein
– glycolipids
• external surface of cell membrane
– glycoproteins
• external surface of cell membrane
• mucus of respiratory and digestive tracts
– proteoglycans
• gels that hold cells and tissues together
• joint lubrication
• rubbery texture of cartilage
Organic Molecules: Lipids
• Hydrophobic organic molecule
– composed of carbon, hydrogen and oxygen
• Less oxidized and thus has more
calories/grams
• Five primary types in humans
– fatty acids
– triglycerides
– phospholipids
– eicosanoids
– steroids
Fatty Acids
• Chain of 4 to 24 carbon atoms
– carboxyl (acid) group on one end, methyl group on the other
and hydrogen bonded along the sides
• Classified
– saturated - carbon atoms saturated with hydrogen
– unsaturated - contains C=C bonds without hydrogen
Triglycerides (Neutral Fats)
• 3 fatty acids bonded to glycerol molecule
(dehydration synthesis)
• Function - energy storage, insulation and
shock absorption
Phospholipids
• Triglyceride with
one fatty acid
replaced by a
phosphate group
• Amphiphilic
character
– fatty acid “tails”
are hydrophobic
– Phosphate “head”
is hydrophilic
Eicosanoids
• Derived from arachidonic acid (a fatty acid)
• Hormone-like chemical signals between cells
• Includes prostaglandins – produced in all
tissues
– role in inflammation, blood clotting, hormone
action, labor contractions, blood vessel diameter
Steroids and Cholesterol
• Steroid = lipid with carbon atoms in
four rings
– all steroids are derived from cholesterol
• cortisol, progesterone, estrogens, testosterone
and bile acids
• Cholesterol
– important component of cell membranes
– produced only in animal liver cells
• naturally produced by our body
Organic Molecules: Proteins
• Protein = polymer of amino acids
• Combination determines structure and
function
• Amino acid = carbon with 3 attachments
– Amino (NH2), carboxy (COOH) and radical
group (R group)
• 20 unique amino acids
– -R groups differ
– properties determined by -R group
Protein Structure
and Shape
• Primary structure
– amino acid sequence
• Secondary structure
– coiled or folded shape
– hydrogen bonds between
negative C=O and positive NH groups
• Tertiary structure
– further folding and bending
into globular and fibrous
shapes
• Quaternary structure
– associations of two or more
separate polypeptide chains
Protein Conformation and
Denaturation
• Conformation – unique 3-D shape crucial
to function
– ability to reversibly change their conformation
• opening and closing of cell membrane pores
• Denaturation
– conformational change that destroys function
• extreme heat or pH
Protein Functions
• Structure
– collagen, keratin
• Communication
– some hormones, cell receptors
• Membrane Transport
– channels, carriers
• Catalysis
– enzymes
Protein Functions
• Recognition and protection
– antigens, antibodies and clotting proteins
• Movement
– molecular motor = molecules that can change
shape repeatedly
• Cell adhesion
– proteins bind cells together
Enzymes
• Proteins as biological catalysts
– promote rapid reaction rates
• Substrate - substance an enzyme acts
upon
• Naming Convention
– named for substrate with -ase as the suffix
• amylase enzyme digests starch (amylose)
• Lowers activation energy = energy needed
to get reaction started
– enzymes facilitate molecular interaction
Enzymes
– Enzymes are organic catalysts generally
consisting of proteins
– Catalysts are molecules that speed up
reactions by lowering energies of activation
Enzymes and Activation Energy
Steps of an Enzyme Reaction
• Substrate approaches enzyme molecule
• Substrate binds to active site forming
enzyme-substrate complex
– highly specific
• Enzyme breaks bonds in substrate
• Reaction products released
• Enzyme repeats process over and over
Enzymatic
Reaction
Steps
Enzymatic Action
• Reusability of enzymes
– enzymes are unchanged by the reactions
• Astonishing speed
– millions of molecules per minute
• Temperature and pH
– change shape of enzyme and alter its ability to bind
– enzymes vary in optimum pH
• salivary amylase works best at pH 7.0
• pepsin works best at pH 2.0
– temperature optimum for human enzymes = body
temperature
Cofactors and Coenzymes
• Cofactors
– nonprotein partners (iron, copper, zinc,
magnesium or calcium ions)
– bind to enzyme and change its shape
– essential to function
• Coenzymes
– organic cofactors derived from water-soluble
vitamins (niacin, riboflavin)
– transfer electrons between enzymes
Coenzyme
+
NAD
• NAD+ transports electrons from one
metabolic pathway to another
Organic Molecules:
Nucleotides
• 3 components
– nitrogenous base
– sugar (monosaccharide)
– one or more phosphate groups
• Physiological important nucleotides
– ATP = energy carrying molecule
– cAMP = activates metabolic pathways
– DNA = carries genetic code
– RNA = assists with protein synthesis
ATP (Adenosine Triphosphate)
ATP contains adenine, ribose and 3 phosphate groups
ATP
• Holds energy in covalent bonds
– 2nd and 3rd phosphate groups have high
energy bonds ~
• ATPases hydrolyze the 3rd high energy
phosphate bond
– separates into ADP + Pi + energy
• Phosphorylation
– addition of free phosphate group to another
molecule
Overview of
ATP
Production
• ATP consumed within
60 seconds
• Continually replenished
Other Nucleotides
• Cyclic adenosine monophosphate (cAMP)
– formed by removal of both high energy Pi’s
from ATP
– formation triggered by hormone binding to cell
surface
– cAMP becomes “second messenger” within
cell
– activates effects inside cell
Nucleic Acids
• DNA (deoxyribonucleic acid)
– 100 million to 1 billion nucleotides long
– contains genetic code
• cell division, sexual reproduction, protein synthesis
• RNA (ribonucleic acid) – 3 types
– transfer RNA, messenger RNA, ribosomal RNA
– 70 to 10,000 nucleotides long
– involved in protein synthesis coded for by
DNA