1. Chemistry
Part 1

2. Has mass & takes up space States of
Matter
Has mass & takes up space
States of
Solid – definite shape & volume
Liquid –definite volume, changeable shape
Gas – changeable shape & volume

3. Conversion between forms Conservation of energy
Capacity to do work
Types of energy
Kinetic – motion/action
Potential – position; stored energy
Forms
Chemical – atomic bonds
Electrical – movement of charged particles
Mechanical – moving matter
Radiant – energy traveling in waves
Conversion between forms
Conservation of energy

4. Subatomic particles Atoms Elements Atomic symbols
Composition of Matter
Subatomic particles
Electrons (e-)
Neutrons
Protons
Atoms
Unique arrangements of subatomic particles
Can’t break down chemically
Elements
Matter of one type of atom
Atomic symbols

5. Properties of Elements
Periodic table
Elements grouped by properties
Unique physical & chemical properties
Physical properties –detected by senses
Chemical properties –way atoms interact

6. Elemental Composition of the Human Body
Major Constituents
O ~ 65%
C ~ 18.5%
H ~ 9.5%
N ~ 3.2 %
Remaining ~ 3.8%
Ca, P, K, S, Na, Cl, Mg, I, Fe
Trace elements
Zn, Mn, Cu

7. Nucleus Electron orbitals Atomic Structure Neutrons no charge
mass = 1 atomic mass unit (amu)
Protons
+1 charge
mass = 1 amu
Electron orbitals
Electrons
surround nucleus in energy levels (orbitals)
-1 charge
mass = amu

8. Electrons move around the nucleus in fixed, circular orbits
Atomic Models
Planetary Model
Electrons move around the nucleus in fixed, circular orbits
Orbital Model
Regions around the nucleus in which electrons are most likely to be found

9. Characteristics of Atoms & Elements
Atomic number –
# protons
Atomic mass –
# protons + # neutrons
Isotope –
neutron # can vary
different # of neutrons
Atomic weight – average mass of all isotopes
11C, 12C, 13C, 14C or 235U, 238U
Radioisotopes – atoms that undergo spontaneous decay called radioactivity
131I, 99mTc, 60Co, 14C, or 235U

10. Subatomic Configuration of Elements
Figure 2.2

11. Isotopes of Elements
Figure 2.3

12. Chemically Inert Elements
Inert elements have full outer e- orbitals
Full = 8 e- (except for He)
Figure 2.4a

13. Chemically Reactive Elements
Reactive elements have unfilled outer orbitals
Figure 2.4b

14. Molecule – ≥ 2 atoms bonded
Molecules & Compounds
Molecule – ≥ 2 atoms bonded
Compound – ≥ 2 different kinds of atoms bonded
ALL Compounds Are Molecules
BUT
All Molecules NOT Compounds

15. Chemical bonds formed by e- in outer orbitals (valence shell)
The Octet rule
Atoms interact to have 8 e-s in valence shells
(except H only 2 e-s)
lose, gain or share
Types of Bonds
Ionic – loss or gain e-s
Covalent – sharing e-s

16. Example: NaCl (sodium chloride)
Ionic Bonds
Ions
charged atoms
Anions - charge = gained e-
Cations + charge = lost e-
Example: NaCl (sodium chloride)

17. Formation of an Ionic Bond
Ionic compounds form crystal structures

18. Single bond each atom donates 1 e-
Covalent Bonds
Sharing e-s fill outer orbitals
Single bond each atom donates 1 e-

19. Double & Triple Covalent Bonds
Atoms share 2 or 3 e-s

20. Electronegative & electropositive atoms form ionic bonds
Polar & Nonpolar Bonds
Polar bonds
Unequal e- sharing
Electronegative & electropositive atoms form ionic bonds
Atoms w/ 6 -7 valence shell e-s = electronegative
Atoms w/ 1-2 valence shell e-s = electropositive
Nonpolar bonds
Equal sharing of e-s
Atoms w/ 3-5 valence e-s form covalent bonds

21. Comparison of Ionic, Polar Covalent, & Nonpolar Covalent Bonds

22. Crucial bond for life functions
Hydrogen Bonds
Crucial bond for life functions
Allows reversible interactions between molecules
Unequal sharing of H’s e- with N or O
Responsible for properties of H2O
Bonds between large macromolecules (ie proteins & nucleic acids)

23. Hydrogen Bonds in H2O
Figure 2.9

24. Properties of Water
High heat capacity – absorbs & releases large amounts of heat before changing temperature
High heat of vaporization – boiling requires large amounts of heat
Polar solvent properties – dissolves ionic substances, forms hydration layers around large charged molecules, & serves as the body’s major transport medium
Reactivity – is an important part of hydrolysis & dehydration synthesis reactions

25. Mixtures & Solutions
Mixtures – two or more components physically intermixed but not chemically bonded
Solutions – homogeneous mixtures of components
Solvent – substance present in greatest amount
Solute – substance(s) present in smaller amounts
Colloids - heterogeneous mixtures whose solutes do not settle out
Suspensions - heterogeneous mixtures with visible solutes that settle out

26. Concept of Concentration
Concentration refers to the amount of a substance in a given volume
A critical concept to master

27. Mass (weight) Review of Metric Units Kg g mg
1Kg = 1000g OR 1g = 0.001Kg
1g = 1000mg OR 1mg = 0.001g
1mg = 1000 mg OR 1 mg = 0.001mg

28. Volume Review of Metric Units L ml 1 L = 1000ml OR 1 ml = 0.001 L
1 ml = 1000 ml OR 1 ml = 0.001ml
Less common units:
dL = deciliter
1dL = 100ml OR 1dL = 0.1L

29. Units of Concentration
Percent – parts per hundred
PPM – parts per million
PPB – parts per billion
Molarity – moles per liter
Molality – moles particles per kg solvent
Mass/volume
g/ml (gram per milliliter)
mg/ml (milligram per milliliter )
g/ml (microgram per milliliter)
Mass/Mass
mg/kg body mass
g/kg body mass

30. Molecular Weight, Moles & Molarity
The mass of a mole of atoms or molecules
Has units of g/mole
= atomic weight of an atom in grams
1 mole of C = 12 g OR C is 12g/mole
Molecule’s molecular weight = sum of the atomic weights of its atoms
1 mole of H2O = 1g + 1g + 16g = 18 g
MW of H2O = 18g/mole

31. Molecular Weight, Moles & Molarity
What’s a mole??
The number of molecules in the gram molecular weight of that molecule
MW of C = 12
there are 12g C/mole C
there are 6.02 x 1023 C atoms in 12g of C
MW of H2O = 18
there are 18g H2O/mole H2O
there are 6.02 x 1023 H2O molecules in 18g of H2O
Always 6.02 x 1023
This magical number is Avogadro’s Number

32. Molarity – The Standard of Chemical Concentration Terms
Molarity = moles of solute per liter (L) of solvent
Abbreviated with M
A 5M NaCl solution contains 5 moles of NaCl molecules per liter of solution
So 1 L of a 5 M NaCl solution contains 3.01 x 1024 molecules of NaCl
6.02 x 1023 molecules/mole x 5 moles = 3.01 x 1024
What mass of NaCl is in 1 L of a 5 M solution?
MW NaCl = = 58.5g/mole
5 moles x 58.5g/mole = 292.5g NaCl

33. Chemical formula for glucose is C6H12O6
Calculating Molarity
Chemical formula for glucose is C6H12O6
MW = 180 g/mole
(6 x 12)+ (12 x 1) + (6 x 18) =180
What is concentration of glucose in a can of coke?
42g of glucose/355ml H2O
How many moles of glucose?
42g 180g/mole = moles
How many liters of coke?
355 ml  1000ml/L = L
How many moles/liter?
0.233 moles/0.355 L = M

34. Examples using Molarity
Ion concentrations in cells & body fluids
[Na] = 0.15M outside cells & 0.015M inside cells
[K] = 0.005M outside cells & 0.15M inside cells
How much more Na is outside than inside?
How much more K is inside than outside?
Molar is often converted to millimolar (mM)
M = moles/L
mM = millimoles/L
0.15M = 150mM
multiply M by 1000 to convert M to mM

35. Mass/Volume & Percent Concentration Terms
Many molecules in the body are measured in mass/volume
Normal glucose = 100mg/dL
or 100mg/100ml
or 1mg/ml or 1g/L
or 0.001g/ml
or 0.1g/100ml
Cholesterol should be below 200mg/dL
or 0.2g/dL
or 0.2g/100ml
Many substances are described in percentages
Percentage is g/100g or g/100ml
Blood [glucose] would = 0.1%
Blood [cholesterol] would = 0.2%

36. Forming or breaking chemical bonds Chemical equations show:
Chemical Reactions
Forming or breaking chemical bonds
Chemical equations show:
Reactants & products & their relative amounts

37. Patterns of Chemical Reactions
Combination reactions: Synthesis reactions which always involve bond formation
A + B  AB
Decomposition reactions: Molecules are broken down into smaller molecules
AB  A + B
Exchange reactions: Bonds are both made & broken
AB + C  AC + B

38. Oxidation-Reduction (Redox) Reactions
C6H12O6 + 6O CO H2O
glucose
carbon dioxide
Reactants losing electrons are electron donors & are oxidized
Reactants taking up electrons are electron acceptors & become reduced

39. Energy Flow in Chemical Reactions
Exergonic reactions – reactions that release energy
Endergonic reactions – reactions whose products contain more potential energy than did its reactants

40. Rates & Reversibility of Chemical Reactions
A + B AB
Chemical reactions proceed with measurable rates
All reactions are theoretically reversible
Equilibrium (dynamic)
Forward & reverse reactions proceed at same rate

41. Factors Influencing Rate of Chemical Reactions
Temperature
higher temperatures increase rates
Concentration
higher [reactant] increases rates
Catalysts
Molecules that increase reaction rates
Enzymes
Biological catalysts with high specificities

42. Chemistry Comes Alive Part 2

43. Categories of Molecules
Organic molecules
Contain C
Covalent bonds
Produced by living or once living organisms
Inorganic molecules
Typically do not contain C
Include:
Water, ammonia, salts, & some acids & bases

44. Examples: NaCl, KCl, CaCl2, Ca3(PO4)2
Salts
Electrolytes
conduct electrical currents
Examples: NaCl, KCl, CaCl2, Ca3(PO4)2

45. Acids release H+ Bases release OH– Acids & Bases HCl  H+ + Cl –
NaOH  Na+ + OH–

46. Acid-Base Concentration (pH)
pH describes the concentration of H+ in a solution
pH = -log[H+]
pH scale based on [H+] of H2O
H2O  H+ + OHˉ
[H+] =1x10-7M & [OHˉ] = 1x10-7M
Therefore H2O has a pH = 7
Acidic solutions have [H+] higher than 1x10-7M & therefore a pH < 7
Alkaline solutions have lower [H+] concentrations & therefore a pH > 7

47. Acid-Base Concentration (pH)
Acidic: pH 0–6.99
Neutral: pH 7.00
Basic: pH 7.01–14
Figure 2.12

48. Solutions of molecules that resist changes in pH
Buffers
Solutions of molecules that resist changes in pH
pH of blood is maintained by carbonic acid-bicarbonate buffering system
Carbonic acid dissociates, reversibly releasing bicarbonate ions & protons
The chemical equilibrium between carbonic acid & bicarbonate resists pH changes in the blood
H2O + CO H2CO H+ + HCO3ˉ

49. Unique to living systems - hence organic
Organic Molecules
Unique to living systems - hence organic
Small molecules & macromolecules (biomolecules)
Macromolecules are polymers of smaller organic molecules
Major biomolecule groups
Carbohydrates
Lipids
Proteins
Nucleic Acids

50. Monosaccharides or simple sugars
Carbohydrates
Major functions
energy source
structural support (plants, fungi & bacteria)
parts of other macromolecules
Monosaccharides or simple sugars

51. Carbohydrates
Disaccharides

52. Carbohydrates Polysaccharides - polymers of simple sugars
Each monosaccharide is a residue
Glycogen - energy storage in animals
Starch – energy storage in plants
Cellulose – support structures in plants
Figure 2.13c

53. Carbohydrates in other Biomolecules
Glycoproteins & proteoglycans
Proteins containing sugar residues
Glycolipids
Phospholipids with sugar residues
Ribose & Deoxyribose sugars part of nucleotides & nucleic acids

54. Lipids Representatives Neutral fats – triglycerides - adipose tissue
Phospholipids – chief component of cell membranes
Steroids – cholesterol, bile salts, vitamin D, hormones
Vitamins A, E & K
Eicosanoids – prostaglandins

55. Fatty Acids & Triglycerides

56. Phospholipids
Figure 2.14b

57. Phospholipids in membranes
Phospholipids make up cellular membranes (lipid bilayers)

58. Steroid Lipids

59. Used to make prostaglandins
Eicosanoids –
Used to make prostaglandins

60. Polymers of amino acids
Protein
Polymers of amino acids
Figure 2.16

61. Amino Acids
Building blocks of protein, containing an amine group, a carboxyl group, & a variable side chain

62. Structural Levels of Proteins

63. Structural Levels of Proteins

64. Structural Levels of Proteins
Figure 2.17d, e

65. Fibrous & Globular Proteins
Fibrous proteins
Extended & thread-like proteins
Examples: keratin, elastin, collagen, myosin, actin
Globular proteins
Compact, spherical proteins with tertiary & quaternary structures
Examples: antibodies, peptide-hormones, & enzymes

66. Characteristics of Enzymes
Proteins that are biological catalysts
Chemically specific
Usually named for the reaction they catalyze
Names often end in suffix -ase

67. Protein Function - Enzymes
Protein & substrate fit together in a specific way due to H-bonds, ionic bonds & non-polar interactions
Figure 2.18a

68. Mechanism of Enzyme Action
Active site
Substrates 1
Enzyme (E)
Substrates (s)
Enzyme-substrate complex (E–S)
H20
Free enzyme (E) 2 3
Covalent bond
Internal rearrangements leading to catalysis
Product (P)

69. Polymers of nucleotides Nucleotide is
Nucleic Acids
Polymers of nucleotides
Nucleotide is
N-containing base
pentose sugar
phosphate group
DNA & RNA

70. Nucleotides – The Bases

71. Nucleotides – The Sugars
for RNA
for DNA

72. Nucleosides – Sugar + Base
Deoxyadenosine

73. Nucleotide – Nucleoside + Phosphates

74. Nucleic Acids – Polymers of Nucleotides

75. Complementary base-pairing A-T G-C
Structure of DNA
Complementary base-pairing
A-T
G-C
Figure 2.21a

76. Deoxyribonucleic Acid (DNA)
Double-stranded helical molecule
Constitutes chromosomes in nucleus
Replicates ensuring genetic continuity
Provides instructions for protein synthesis

77. Ribonucleic Acid (RNA)
Contains the base uracil in place of thymine
Made from a DNA template
Three major varieties of RNA:
mRNA – encodes a protein
tRNA – conveys amino acid to ribosome as directed by mRNA
rRNA – joins amino acids together to form protein as directed by mRNA

78. Adenosine Triphosphate (ATP)
Source of immediately usable energy for the cell
Figure 2.22

79. How ATP Drives Cellular Work
Figure 2.23