Chemical Level of Organization Chapter 2 Lecture Notes

1
Chemical Level of OrganizationChapter 2
Lecture Notes

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

2
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

3
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)

4
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

5
Introduction

• focus on structure of atoms
• how atoms bond to form molecules
• nature of chemical reactions
• five families of biological molecules

6
Concept 1.2Atomic Structure
7
Chemical Elements

• a substance that cannot be split into simpler
  substance by ordinary chemical means
• 112 recognized element
• 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)

8
Table 2.1
9
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)

10
Figure 2.1
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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 electrons

12
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

13
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?

14
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-

15
Figure 2.2
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Electrons

• outermost shell is called valence
• valance wants eight as magic number
• if eight electrons are present in valence, atom
  is inert
• if fewer than eight electron are present in
  valence, atoms 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 valence

17
Electrons

• outermost shell is called valence
• valance wants eight as magic number
• if eight electrons are present in valence, atom
  is inert
• if fewer than eight electron are present in
  valence, atoms 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

18
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

20
Figure 2.3
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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

23
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

24
Table 2.2
25
Figure 2.4
26
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

27
Figure 2.5
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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

29
Figure 2.6
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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

31
Figure 2.7
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Concept 2.3 Chemical Reactions
33
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

34
Figure 2.8
35
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

36
Energy Laws Chemical Reactons

• 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

37
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)

38
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

39
Figure 2.9
40
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

41
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.

42
Figure 2.10
43
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
46
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

47
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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Figure 2.6
49
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

50
Figure 2.11
51
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

52
Figure 2.21
53
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

54
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

55
Figure 18.1a
56
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

57
Figure 2.12
58
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)

59
Figure 2.13
60
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 a weak acids
• convert strong bases to weak bases
• carbonic acid-bicarbonate buffer system

61
Concept 2.5 Organic Molecules
62
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

63
Figure 2.14
64
Figure 2.15
65
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

67
Figure 2.15
68
Figure 2.16
69
Concept 2.7 Lipids
70
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

71
Table 2.5
72
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

73
Figure 2.17
74
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

75
Figure 2.17c
76
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

77
Figure 2.18
78
Steroids

• structurally unique
• four rings of carbon atoms
• synthesized from cholesterol
• commonly found steroids
• cholesterol
• estrogen
• testosterone
• cortisol
• bile salts
• vitamin D

79
Figure 2.19
80
Concept 2.8 Proteins
81
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

82
Figure 2.6
83
Figure 2.20
84
Figure 2.20
85
Figure 2.20
86
Figure 2.21
87
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

88
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

89
Figure 2.22
90
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)

91
Figure 2.23
92
Concept 2.9 Nucleic Acids
93
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

94
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

95
Figure 2.24
96
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

97
Concept 2.10 ATP
98
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

99
Figure 2.25
100
End Chapter 2