1. The Nature of Molecules
Chapter 2

2. Recall from Chapter 1: Levels of Organization
Cellular Organization
cells
organelles
molecules
atoms
The cell is the
basic unit of life.

3. Fig 3

4. Atomic Structure All matter is composed of atoms
Understanding the structure of atoms is critical to understanding the nature of biological molecules:
Bonding
Reactivity
Oxidation/ Reduction pH
Energy

5. Atomic Structure Atoms are composed of:
Protons – positively charged particles
Neutrons – neutral particles
Electrons – negatively charged particles
Protons and neutrons are located in the nucleus of the atom
Electrons are found in orbitals surrounding the nucleus

6. Atomic Structure

7. Atomic Structure
Every different atom has a characteristic number of protons in the nucleus.
Atomic Number = number of protons
Atoms with the same atomic number have the same chemical properties and belong to the same element.

8. Atomic Structure
Each proton and neutron has a mass of approximately 1 dalton.
The sum of protons and neutrons is the atom’s atomic mass.
Isotopes – atoms of the same element that have different atomic mass numbers due to different numbers of neutrons.

9. Atomic Structure: Carbon Isotopes

10. Atomic Structure
Neutral atoms have the same number of protons and electrons
Ions are charged atoms
Cations – have more protons than electrons and are positively charged (+)
Anions – have more electrons than protons and are negatively charged (-)

11. Atomic Structure Atomic Energy Levels
Electrons have discrete energy levels around the nucleus - Energy of Position
Energy Levels are labelled K-M
The distance an electron is from the nucleus is relative to the amount of potential energy of the electron

12. Electrons release energy as they fall from a high
Fig. 2.5
Electrons release energy
as they fall from a high
energy level to a lower
energy level
Electrons climb from a
lower energy level to a
higher energy level as
they absorb energy

13. Atomic Structure Electron Orbitals
At each Energy Level, electrons are located in orbitals described as s or p orbitals
Orbitals are areas where the electron is most likely to be found
Each orbital can contain only 2 electrons

14. Atomic Structure Electron Orbitals
At the K Energy Level, there is one s orbital (referred to as the 1s orbital)
There can be two electrons in the K Energy Level
Fig. 2.4a

15. Atomic Structure Electron Orbitals
At the L Energy Level, there is one s orbital (referred to as the 2s orbital) and 3p orbitals (each referred to as 2p orbitals)
The L Energy Level holds 8 electrons
Fig. 2.4b

16. Atomic StructureElectron Configuration for Neon
Neon has 10 electrons
The K and L Energy Levels are
completely filled
The 1s orbital of K and the 2s and three
2p orbitals of L are completely filled
Fig. 2.4c

17. Fig. 2.4

18. Atomic Structure
Valence electrons are the electrons in the outermost energy level of an atom
Any level (K-M) can be the outermost level, it depends on how many electrons the atom has
An element’s chemical properties depend on interactions between valence electrons of different atoms
These chemical properties include:
Reactivity - atom’s ability to make or break bonds
Electronegativity - atom’s affinity for electrons

19. Fig. 2.2

20. Atomic Structure
Electrons can be transferred from one atom to another, while still retaining the energy of their position in the atom
Oxidation = loss of an electron
Reduction = gain of an electron

21. Atomic Structure LEO the lion says GER
If it Looses Electrons, it’s Oxidized
If it Gains Electrons, it’s Reduced

22. Atomic Structure
Oxidation and Reduction reactions always occur in pairs
If a an atom or molecule is reduced, another atom or molecule must have been oxidized
If a an atom or molecule is oxidizied, another atom or molecule must have been reduced
For this reason Oxidation and Reduction Reactions are known as Redox Reactions

23. Atomic Structure Redox Reactions Some Terminology:
The entity that lost (donates) the electrons in a redox rxn. is oxidized
Therefore, it is known as the Reducing Agent in the Redox Rxn.
It caused the other entity to be reduced by providing the electron

24. Atomic Structure Redox Reactions Conversely:
The entity that gained (accepted) the electrons in a redox rxn. is reduced
Therefore, it is known as the Oxidizing Agent in the Redox Rxn.
It caused the other entity to be Oxidized by accepting the electron

25. Atomic Structure Redox Reactions
Differentiate Oxidation/ Reduction and Oxidizing Agent/ Reducing Agent
Remember, LEO says GER describes Oxidation/ Reduction
The Oxidizing Agent is Reduced, The Reducing Agent is Oxidized
Redox Reactions are very important in biological processes (and easy test questions)
Cellular Respiration, Photosynthesis

26. Elements
The Periodic Table arranges all elements according to their atomic number and their number of Valence Electrons
Therefore, the Periodic Table identifies and groups elements with similar chemical properties
The number of Valence Electrons of an element can be determined by counting across the columns of the Periodic Table

27. Periodic Table of the Elements
O has 6
Valence e-s
H has 1 e-
C has 4
Valence e-s
Na has 1
Valence e-
Ca has 2
Valence e-s

28. Elements There are 90 naturally occurring elements
Only 12 elements are found in living organisms in substantial amounts
Twelve or so others are found in trace amounts: I, for example
Four elements make up 96.3% of human body weight: carbon, hydrogen, oxygen, nitrogen

29. Elements Become Familiar with the Valence Electron
Configuration for the Elements:
Hydrogen (H): 1
Carbon (C): 4
Oxygen (O): 6
Nitrogen (N): 5
Chlorine (Cl): 7
Sodium (Na): 1

30. (H and He are exceptions)
Elements
Octet rule: States that Atoms typically
attempt to achieve 8 valence electrons
(H and He are exceptions)
Eight valence electrons will fill the outer energy
level (H needs 1e-, He needs 2 e-s)
Atoms will steal, give up, or share valence
electrons in order to fulfill the Octet Rule
Atoms with full energy levels are more
stable and less reactive than atoms with
unfilled energy levels

31. Elements

32. Chemical Bonds
Atoms can fulfill the octet rule by loosing, gaining or sharing electrons with other atoms, creating chemical bonds
Atoms are held together in molecules or compounds by chemical bonds
Molecules are groups of atoms held together in a stable association
Compounds are molecules containing more than one type of element

33. Chemical Bonds We will study 3 types of chemical bonds Ionic Bonds
Covalent Bonds
Hydrogen Bonds

34. Chemical Bonds 1. Ionic bonds
Ionic Bonds are formed by the attraction of oppositely charged ions
Cation Na+ and Anion Cl- attract

35. Chemical Bonds 1. Ionic bonds
Na easily gives up its single valence electron to fulfill the octet rule
Because Na lost an electron it now has a positive (+) charge

36. Chemical Bonds 1. Ionic bonds
Cl, with 7 valence electrons, easily picks up the extra electron to fulfill the octet rule
Because Cl gained an electron it now has a positive (-) charge

37. Chemical Bonds 1. Ionic bonds
Cation Na+ and Anion Cl- attract
Ionic Bonds form crystals
Ionic Bonds dissociate in water

38. Chemical Bonds 2. Covalent Bonds
Covalent bonds form when atoms share 2 or more valence electrons
Covalent bond strength and length depends on the number of electron pairs shared by the atoms
Single Covalent Bond - one pair of e-’s is shared
Double Covalent Bond - two pairs of e-’s are shared
Triple Covalent Bond - three pairs of e-’s are shared
Covalent Bonds are the strongest bonds and do not dissociate in water

39. Chemical Bonds 2. Covalent Bonds
Covalent bond strength depends on the number of electron pairs shared by the atoms
Single Covalent Bond - one pair of e-’s is shared
C C
Double Covalent Bond - two pairs of e-’s are shared
C C
Triple Covalent Bond - three pairs of e-’s are shared
C C

40. Chemical Bonds 2. Covalent Bonds
Covalent bond strength depends on the number of electron pairs shared by the atoms
Single Covalent Bond - one pair of e-’s is shared
Double Covalent Bond - two pairs of e-’s are shared
Triple Covalent Bond - three pairs of e-’s are shared
Bond Strength:
Bond Length:
single
bond
double
bond
triple
bond
<
<
single
bond
double
bond
triple
bond
>
>

41. Chemical Bonds 2. Covalent Bonds

42. Chemical Bonds Electronegativity is an atom’s affinity for electrons
Atoms of each of the elements differ in electronegativity
Atomic size, mass, nuclear charge, electron configuration all attribute to the differences in electronegativity
Based on what you know about the Na atom, do you think Na has a high or low electronegativity?
What about Cl?

43. Periodic Table of Elements

44. Chemical Bonds Electronegativity is an atom’s affinity for electrons
The Periodic Table of Elements discloses an Electronegativity Trend

45. Electronegativity(EN) Trend in the Periodic Table of Elements
Greatest EN
Increasing EN
Lowest EN

46. Table 2.2

47. Chemical Bonds
A covalent bond is formed by a sharing of electrons, but electrons are not always shared equally
Differences in electronegativity dictate how electrons are distributed in covalent bonds
In a covalent bond between a highly electronegative atom and a weakly electronegative atom, electrons will be drawn more towards the highly electronegative side of the bond

48. Chemical Bonds
The water molecule is an important example of a covalent bond between a highly electronegative atom and a weakly electronegative atom
As a result of this unequal sharing of electrons, one side of the molecule (the highly electronegative side) has a partial negative charge and the other side has a partial positive charge

49. Fig. 2.11a

50. Chemical Bonds
A covalent bond with an unequal sharing of electrons is known as a polar covalent bond
Polar because it has two ‘poles’, a partial positive pole and a partial negative pole
A covalent bond with an equal sharing of electrons is known as a nonpolar covalent bond
A Carbon to Carbon covalent bond is an example of a nonpolar covalent bond
C C

51. Chemical Bonds
A molecule with an unequal sharing of electrons is known as a polar molecule
Water is an important polar molecule
A molecule with an equal sharing of electrons is known as a nonpolar molecule
A Carbon to Carbon molecule (like an oil molecule) is an example of a nonpolar molecule
Polar molecules mix only with other polar molecules, nonpolar molecules mix only with other nonpolar molecules
‘Like dissolves like’

52. Chemical Bonds
Chemical reactions involve the formation or breaking of chemical bonds
The making and breaking of bonds
Whether a chemical reaction occurs is influenced by many factors:
temperature
concentration of reactants and products
availability of a catalyst

53. Chemical Bonds Chemical reactions are written with the
reactants first, followed by the products
6H2O + 6CO C6H12O O2
reactants products
Chemical reactions are often reversible
C6H12O O H2O CO2

54. Water Chemistry All living organisms are dependent on water
The structure of water is the basis for its unique properties
The most important property of water is the ability to form hydrogen bonds

55. Water Chemistry

56. Water Chemistry 3. Hydrogen Bonds
Hydrogen bonds are weak attractions between the partially negative oxygen of one water molecule and the partially positive hydrogen of a different water molecule.
Hydrogen bonds can form between water molecules or between water and another charged molecule.

57. Water Chemistry 3. Hydrogen Bonds
Within a water molecule, the bonds between oxygen and hydrogen are highly polar.
Partial electrical charges develop:
oxygen side is partially negative
hydrogen side is partially positive
The partially negative oxygen side of one
water molecule is weakly attracted to the
partially negative hydrogen side of another
water molecule to form a hydrogen bond

58. Water Chemistry

59. Water Chemistry
The polarity of water causes it to be cohesive and adhesive.
Cohesion: water molecules stick to other water molecules by hydrogen bonding
Adhesion: water molecules stick to other polar molecules by hydrogen bonding

60. Water Chemistry Surface Tension

61. Water Chemistry Capillary Action
Look for the meniscus

62. Properties of Water 1. Water has a high specific heat
A large amount of energy is required to change the temperature of water.
2. Water has a high heat of vaporization
The evaporation of water from a surface causes cooling of that surface.

63. Properties of Water 3. Solid water is less dense than liquid water
Bodies of water freeze from the top down
4. Water is a good solvent
Water dissolves polar molecules - “like dissolves like”
Water dissolves ionic molecules
Water Soluble

64. Properties of Water Water dissolves ionic molecules

65. Properties of Water 5. Hyrophobic Interactions
Water organizes nonpolar molecules
Polar molecules are hydrophilic: “water-loving”
Nonpolar molecules are hydrophobic: “water- fearing”
Water causes hydrophobic molecules to aggregate or assume specific shapes
“Like dissolves like”, oil and water don’t mix
A Micelle

66. Properties of Water 6. Water can form 3 ions: H2O  OH-1 + H+1
water hydroxide ion hydrogen ion
H2O H  H3O+1
water hydronium ion

67. Properties of Water 6. Water can form 3 ions:
1. Hydrogen Ion
H2O  OH H+1
water hydroxide ion hydrogen ion
The hydrogen ion is a lone proton with a +1 charge

68. Fig. 2.2

69. Properties of Water 6. Water can form 3 ions:
2. Hydroxide Ion
H2O  OH H+1
water hydroxide ion hydrogen ion
The hydoxide ion retains the electron yielding a -1 charge

70. Fig. 2.2

71. Properties of Water 6. Water can form 3 ions:
3. Hydronium Ion
H2O H  H3O+1
water hydronium ion
A water molecule can pick up a hydrogen ion (a proton, H+1) to become a hydronium ion with +1 charge

72. Table 2.3

73. Is water an Acid or a Base?
Acids and Bases
Acid: a chemical that releases H+1 ions
Base: a chemical that accepts H+1 ions
The acidity/basicity of a solution is measured in terms of the pH Scale
Is water an Acid or a Base?

74. Is water an acid or a base?
Acids and Bases
Is water an acid or a base?
H2O  OH H+1
H2O H  H3O+1
Water acts as an acid and a base
Pure water exists as a balance of H2O, OH-1, and H3O+1
Water has a neutral pH

75. Acids and Bases

76. Acids and Bases The pH Scale
Hydrogen ion concentration [H+1] is the basis of the pH scale
pH = -log [H+1] or pH =
log [H+1]
Greater H+1 concentration = lower pH (acidic)
Lower H+1 concentration = higher pH (basic)

77. Acids and Bases The pH Scale
Hydrogen ion concentration [H+1] of pure water is 10-7 moles [H+1] / L
pH = pHwater =
log [H+] log [10-7]
pHwater = 7
The pH scale is based on the pH of pure water

78. Acids and Bases

79. Acid: a chemical that releases H+1 ions
Acids and Bases
Acid: a chemical that releases H+1 ions
Hydrochloric acid, HCl
HCl  H Cl-
Sulfuric acid, H2SO4
H2SO4  2H SO4-2
Carbonic acid, H2CO3 (Example from text)
H2CO3  H HCO3-
Bicarbonate ion

80. Base: a chemical that accepts H+1 ions
Acids and Bases
Base: a chemical that accepts H+1 ions
Sodium Hydroxide, NaOH
NaOH  Na+1 + OH-1
OH-1 + H+1  H2O
Bicarbonate, HCO3-
HCO H+  H2CO3
carbonic acid

81. Acids and Bases Acid-Base Pairs
Acids and Bases come in pairs
H2CO3  H HCO3-
Carbonic acid Bicarbonate
HCO H+  H2CO3
Bicarbonate Carbonic acid
H2CO H HCO3-

82. Acids and Bases Buffer Systems
Buffer: a chemical that resists a change in pH
Buffers accept/release H+1 as necessary to keep pH constant
Buffers accept hydrogen ions in acidic solutions and release hydrogen ions in basic solutions, minimize a change in hydrogen ion concentration
The key buffer in human blood is the carbonic acid/ bicarbonate acid-base pair

83. Acids and Bases
Most biological buffers consist of a pair of molecules, one an acid and one a base.

84. Acids and Bases

85. Table 2.1

86. The Nature of Molecules
End Chapter 2