Chemical Level of Organization Chapter 2 Lecture Notes

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Chemical Level of OrganizationChapter 2
Lecture Notes

• to accompany
• Anatomy and Physiology From Science to Life
• by
• Gail Jenkins, Christopher Kemnitz, Gerard Tortora

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Chapter Overview

• 2.1 Atomic Structure
• 2.2 Chemical Bonds
• 2.3 Chemical Reactions
• 2.4 Inorganic Compounds
• 2.5 Organic Molecules
• 2.6 Carbohydrates
• 2.7 Lipids
• 2.8 Proteins
• 2.9 Nucleic Acids
• 2.10 ATP

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Essential Terms

• Chemistry
• Study of structure and interactions of matter
• Matter
• Anything that has mass and occupies space
• Mass
• Amount of matter in any object unchanging
• Weight
• Force of gravity acting on matter (changes)

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Introduction

• Body is composed of chemicals
• Body activities are chemical in nature
• Necessary to understand chemistry in order to
  understand human anatomy and physiology
• Chemistry of water
• Nearly 2/3 of body weight
• Important in chemical reactions
• Helps maintain homeostasis

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Introduction

• Focus on structure of atoms
• How atoms bond to form molecules
• Nature of chemical reactions
• Five families of biological molecules

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Concept 1.2Atomic Structure
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Chemical Elements

• A substance that cannot be split into simpler
  substance by ordinary chemical means
• 112 recognized elements
• 92 naturally occurring elements
• Each has a specific chemical symbol
• One or two letters of name in English, Latin, or
  another language
• 26 elements normally present in humans
• O, C, H, N 96 of bodys mass
• Ca, P, K, S, Na, Cl, mg, Fe 3.8
• Remaining 0.2 are trace elements (14 elements)

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Structure of Atoms

• Atom is the smallest unit of matter that retains
  the properties characteristic of an element
• Subatomic particles to be studied
• Protons (positive charge, in nucleus)
• Neutrons (neutral, no charge, in nucleus)
• Electrons (negatively charged, buzzing around
  nucleus)

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Protons

• Positively charged particles in the nucleus
• Element is defined based on the number of protons
• Atomic number is number of protons
• If you remove a proton you change the element
• Neutrons may or may not be the same number as the
  number of protons
• In a neutral atom, the number of electrons equals
  the number of protons

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Atomic Number Mass Number

• Atomic number
• Number of protons
• Mass number
• Number of protons AND number of neutrons
• Number of neutrons varies
• Isotopes are atoms that have different numbers of
  neutrons than the most common number
• Isotopes therefore have different mass number
  than common elements

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Atomic Mass

• The average mass of all naturally occurring
  isotopes
• Measured in daltons
• Proton 1.007 daltons
• Neutron 1.008 daltons
• Electron 0.0005 daltons
• Also called atomic weight
• What is the difference between mass and weight?

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Electrons

• Found in regions called electron shells
• First shell (nearest the nucleus) holds 2
• Second shell holds up to eight
• Third shell can hold up to eighteen but if less
  than 18 are present they will fill up with eight
  then move to a fourth shell
• Fourth and subsequent shells are the same as the
  third
• Short-hand for electron is e-

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Electrons

• Outermost shell is called valence
• Valence wants eight as magic number
• If eight electrons are present in valence, atom
  is inert
• If fewer than eight electrons are present in
  valence, atom is said to be chemically reactive
  and can
• Gain electrons if it has four or more already
• Lose electrons if it has three or less
• Share electrons with another atoms valance

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Ions, Molecules, Compounds

• Ions are atoms that have gained or lost electrons
• Ions can be positively or negatively charged
• Lost electrons positively charged cation
• Gained electrons negatively charged anion
• Molecules form when atoms share electrons
• Compound is a special type of molecule formed
  from two or more different types of atoms

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Ions, Molecules, Compounds

• Free radical
• Electrically charged atom or group of atoms with
  unpaired electron in valence
• Unstable
• Highly reactive
• Can damage other molecules by stealing or
  donating electrons
• Can break apart important biomolecules

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Concept 2.2 Chemical bonds
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Chemical Bonds

• Forces that hold atoms together
• Depends on number of electrons in valence
• Atoms of most biologically important molecules
  have less than eight electrons in valence (see
  again figure 2.2)
• Octet rule helps explain why and how atoms react
  to form ionic, covalent, or hydrogen bonds

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Ionic Bonds

• Occur between ions of opposite charges
• Opposites attract
• Ionic compounds generally found as solid crystals
• Ionic compound that dissociates in body water are
  called electrolytes
• Called electrolytes because solution can conduct
  electricity

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Covalent Bonds

• Stronger than ionic bonds
• Occurs when atoms share valence electrons
• Co-valence
• Can form between atoms of same or different
  elements
• Most common chemical bonds in body
• Compounds that result from them form most body
  structures
• Can be simple, double, or triple bonds
• Can be equally or unequally shared

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Polar Nonpolar Covalent Bonds

• Polar covalent
• Electrons shared UNEQUALLY
• Results in
• Partial negative end where electron spends most
  of its time
• A partial positive end where electron is rarely
  found
• Nonpolar covalent
• Electrons shared EQUALLY

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Hydrogen Bonds

• Form as a result of partial positive and negative
  charges of polar covalent molecules
• Weak compared to covalent bonds
• Several together can be strong
• Hydrogen bonds form attraction between water
  molecules making it cohesive
• Cohesion of water molecules give water surface
  tension
• Give biomolecules 3D shape

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Concept 2.3 Chemical Reactions
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Chemical Reactions

• Occur when new bonds form or existing bonds break
• Foundation of all life processes
• Starting substances are reactants
• Ending substances are products
• Metabolism is sum of all chemical reactions in
  the body

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Energy

• Capacity to do work
• Potential energy
• Energy stored by matter due to its position
• Kinetic energy
• Energy associated with movement of matter
• Chemical energy
• Form of potential energy stored in bonds of
  compounds and molecules

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Energy Laws Chemical Reactions

• Energy
• Can neither be created nor destroyed
• Left over energy released as heat
• When chemical bonds are broken
• Some energy is wasted
• Heat is thus given off
• Some of which is used to maintain normal body
  temperature

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Energy Transfer

• Exergonic reactions
• Release more energy than they absorb
• Endergonic reactions
• Absorb more energy than they release
• Key feature of bodys metabolism is the coupling
  of exergonic and endergonic reactions (energy
  released from exergonic reactions used to drive
  endergonic reactions)

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Activation Energy

• Energy of activation
• Energy required to activate a chemical reaction
• Molecules, ions, atoms have kinetic energy when
  moving
• Continually moving, colliding
• Forceful enough collision can disrupt valence
  electrons and break chemical bonds
• Kinetic energy needed to break chemical bonds is
  the activation energy

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Activation Energy

• Both concentration and temperature of matter
  influence chance of collision of particles
• Increased concentration increases likelihood of
  collision
• Decreased concentration decreases likelihood
• Increased temperature (increased movement of
  particles) increases likelihood of collision
• Decreased temperature decreases likelihood of
  collision

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Catalysts

• Catalysts increase the likelihood of chemical
  reactions without depending on concentration or
  temperature changes
• Remember
• Metabolism is sum of all chemical reactions
• Body activities are chemical in nature
• We disrupt homeostasis if we increase
  concentrations too high
• We can die if our body temperature goes too high,
  because proteins will denature, etc
• Enzymes are biological catalysts.

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Types of Chemical Reactions

• Synthesis Reactions - Anabolism
• To synthesize is to put together
• A B AB
• Decomposition Reactions - Catabolism
• To compose is to build
• To decompose is to break apart
• AB A B
• Exchange Reaction
• Old bonds broken AND new bonds formed
• AB CD AD BC

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Types of Chemical Reactions

• Reversible Reactions
• Some chemical reactions are reversible
• This is shown by either two arrows each pointing
  the opposite direction in between products and
  reactants
• Or by using one arrow with two heads
• A B AB
• If special conditions are required they are
  written above or below the arrow.

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Concept 2.4 Inorganic Compounds
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Inorganic Compounds

• Lack carbon
• Exception carbon dioxide and bicarbonate ion
• Are structurally simple
• Held together by ionic or covalent bonds
• Include
• Water (55 60 lean adults body mass)
• Salts
• Acids
• Bases

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Water

• Most important and abundant inorganic compound in
  all living things
• Nearly all bodys chemical reactions occur in a
  watery medium
• Polarity of water makes it
• Excellent solvent for both ionic and polar
  substances
• Makes water molecules cohesive
• Allows water to resist temperature changes

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Water as Solvent

• Solvent
• Substance that will dissolve another
• Usually in highest concentration
• Solute
• Substance dissolved in solvent
• Usually in lowest concentration
• Solution
• Solute dissolved in solvent
• Water said to be universal solvent
• Hydrophilic substances dissolve easily
• Polar and ionic compounds and molecules
• Hydrophobic substances do not dissolve easily if
  at all.
• nonpolar compounds

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Water in Chemical Reactions

• Hydrolysis reactions
• Hydro water
• Lysis to break apart
• Hydrolysis is to break apart with addition of
  water
• Allow decomposion reactions to occur
• Dehydration synthesis
• Hydrate add a water molecule
• Dehydrate remove a water molecule
• Synthesis to build a new something
• Dehydration synthesis
• To bring together by the removal of a water
  molecule

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Thermal Properties of Water

• Hydrogen bonds give water molecules tremendous
  cohesiveness
• Compared to most substances, water can absorb
  tremendous amounts of heat energy without
  changing temperature
• Due to hydrogen bonding and polarity of water
  molecules
• Helps modulate body temperature

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Mixtures

• Combination of elements/compounds physically
  blended together but not bound by chemical bonds
• Solutions
• Solutes remain equally distributed
• Appears transparent
• Colloids
• Solutes bigger, large enough to scatter light
• Usually appear opaque or translucent
• Suspensions
• Unlike solutions and colloids particles will
  settle out

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Acids, Bases, and Salts

• Acids
• Have excess hydrogen ions (H)
• Will dissociate into H and one or more anions
• Have low ph
• Bases
• Have excess hydroxide ions (OH-)
• Will dissociate into OH- and one or more cations
• Have high ph
• Salts
• Have neither OH- nor H
• Will dissociate into other cations and anions
• Will form when acids mixed with bases

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pH

• pH expresses solutions acidity or alkalinity
• Acids have low ph (0 - 6.99)
• Bases high ph (7.01-14)
• Pure water is neutral (7.00 exactly)

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Buffer Systems

• pH of fluids inside and outside cells needs to
  remain almost constant
• Strong acids and bases continually taken in and
  formed by body
• Buffer systems function stabilize ph of a
  solution
• Remove or add hydrogen ions
• Convert strong acids to weak acids
• Convert strong bases to weak bases
• Carbonic acid-bicarbonate buffer system

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Concept 2.5 Organic Molecules
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Organic Molecules

• Four main types
• 1. Carbohydrates 3. Proteins
• 2. Lipids 4. Nucleic acids - ATP
• Can be very large (macromolecules or polymers) or
  very small (monomers)
• Monomers covalently bonded via dehydration
  synthesis reactions
• Contain carbon and hydrogen bonded together often
  with oxygen and nitrogen

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Concept 2.6 Carbohydrates
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Carbohydrates

• 2-3 of total body mass
• Sugars, starches, glycogen, cellulose
• In animals function mainly as source of energy to
  produce ATP
• Contain carbon, hydrogen, and oxygen
• Generally CH2O
• Three main groups based on size
• Simple sugars
• Monosaccharides
• Disaccharides
• Complex carbohydrate
• Polysaccharides
• Glycogen and starch

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Concept 2.7 Lipids
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Lipids

• 18-25 of body mass in lean adults
• Contain carbon, hydrogen, oxygen
• Less oxygen than carbohydrates
• Fewer polar covalent bonds
• Insoluble in water hydrophobic
• Includes
• Triglycerides
• Phospholipids
• Steroids
• Fatty acids
• Fat soluble vitamins

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Triglycerides

• Most plentiful in body and diet
• Liquid or solid at room temperature
• Liver converts excess carbohydrates, proteins,
  fats and oils to triglycerides
• Stored in adipose tissue
• Unlimited storage possible
• Most highly concentrated form of chemical energy
• Twice as much chemical energy

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Fats

• Saturated fats
• Fully saturated fatty acid tail
• No double bonds
• Tails are straight (uniform)
• Solid at room temperature
• Unsaturated fats
• Not fully saturated fatty acid tail
• One (monounsaturated) or more (polyunsaturated)
  double bonds
• Tails are kinked at each double bond
• Liquid at room temperature

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Phospholipids

• Similar to triglycerides
• Glycerol and fatty acids
• Different from triglycerides
• Only two fatty acids
• Third fatty acid replaced by phosphate group and
  charged functional group
• Amphipathic
• Fatty acid end hydrophobic
• Phosphate/glycerol end hydrophilic
• Main component of plasma membrane

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Steroids

• Structurally unique
• Four rings of carbon atoms
• Synthesized from cholesterol
• Commonly found steroids
• Cholesterol
• Estrogen
• Testosterone
• Cortisol
• Bile salts
• Vitamin D

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Concept 2.8 Proteins
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Proteins

• 12 18 total body mass in lean adult
• Contain carbon, hydrogen, oxygen, nitrogen
• Amino acids bound together by peptide bonds
• Forms between carboxyl end of one and amino end
  of next amino acid
• 20 different amino acids
• Compare to 26 letters of American alphabet
• Nearly limitless combinations
• Many amino acids bound together called a
  polypeptide
• Functional polypeptide is protein

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Four Levels of Protein Structure

• Primary
• Sequence of amino acids
• Secondary
• Stabilized by hydrogen bonds formed along
  backbone of polypeptide
• Alpha helix
• Repeated clockwise twists
• Beta pleated sheets
• Repeated folding

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Four Levels of Protein Structure

• Tertiary
• Lend unique 3D shape (3rd level 3D!!)
• Determines function
• Several types of bonds
• Strong disulfide bonds
• Weaker hydrogen bonds, ionic bonds, hydrophobic
  interactions
• Quaternary
• Present only when protein has more than one
  polypeptide chain

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Enzymes

• Proteins
• Biological catalysts
• Lower activation energy
• Keeping temperature stable
• Often named for substratesuffix -ase
• Important properties
• Specificity (binds only to substrate)
• Efficiency (very fast reactions)
• Control (can be controlled by cell)

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Concept 2.9 Nucleic Acids
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Nucleic Acids

• NA of DNA and RNA
• Deoxyribonucleic acid DNA
• Ribonucleic acid RNA
• Built from nucleotides
• Three parts of a nucleotide
• One of five nitrogenous bases
• Adenine, thymine, guanine, cytosine, uracil
• Five carbon sugar
• Ribose in RNA
• Deoxyribose in DNA
• Phosphate group

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DNA

• Double helix (discovered by Watson and crick)
• Resembles a spiral ladder
• Alternating sugar (deoxyribose) and phosphate
  group as uprights
• Nitrogenous bases in the middle forming rungs
• Double ring purine (A G) binds to single ring
  pyrmidine (C T)
• A binds to T and T binds to A
• G binds to C and C binds to G
• Serves as template for new DNA synthesis and RNA
  synthesis

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RNA

• Single strand
• Alternating sugar (ribose) and phosphate group as
  backbone
• Nitrogenous bases attached as fringe
• Double ring purine (A G) binds to single ring
  pyrmidine (C U)
• A in DNA binds to U in RNA
• T in DNA binds to A in RNA
• G in DNA binds to C in RNA
• C in DNA binds to G in RNA

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Concept 2.10 ATP
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ATP

• Structurally similar to nucleic acids
• Ribose, adenine, three phosphates
• Energy currency of the cell
• Is spent when third phosphate removed from ATP
  forming ADP free energy
• ATP synthesized by enzyme called ATP synthase
• Requires input of energy

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End Chapter 2