1. Chapter 2 Lecture and Animation Outline
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2. Nature of Atoms Matter has mass and occupies space
All matter is composed of atoms
Understanding the structure of atoms is critical to understanding the nature of biological molecules
Scanning Tunneling Microscopy
Oxygen atoms on rhodium crystal 2

3. Atomic Structure Atoms are composed of Protons Neutrons Electrons
Positively charged particles
Located in the nucleus
Neutrons
Neutral particles
Electrons
Negatively charged particles
Found in orbitals surrounding the nucleus
Bohr Atomic Model 3

5. Atomic number Number of protons equals number of electrons
Number of protons determines atom’s identity 5

6. Atomic mass Mass or weight? Mass – refers to amount of substance
Weight – refers to force gravity exerts on substance
An object has the same mass despite if it’s on Earth or the Moon which would change the weight
Sum of protons and neutrons is the atom’s atomic mass
Each proton and neutron has a mass of approximately 1 Dalton or atomic mass unit (a.m.u.) 6

7. Mass (in Daltons or a.m.u.)
Properties of Subatomic Particles
Subatomic Particle Summary
subatomic particle
Mass (in Daltons or a.m.u.)
charge
proton 1 +
neutron
neutral
electron
almost 0 -

8. Elements
Element
Any substance that cannot be broken down to any other substance by ordinary chemical means
A given element is made up one type of atom
Iron is an element because it is made up of only iron atoms 8

12. Electrons Negatively charged particles located in orbitals
Neutral atoms have same number of electrons and protons
Ions are charged particles – unbalanced
Cation – more protons than electrons = net positive charge
Anion – fewer protons than electrons = net negative charge 12

13. Electron arrangement
Key to the chemical behavior of an atom lies in the number and arrangement of its electrons in their orbitals
Bohr model – electrons in discrete orbits
Quantum model – orbital as area around a nucleus where an electron is most likely to be found 13

14. Electron Shell Diagram Electron Shell Diagram
Electron Orbitals
Electron Shell Diagram
Corresponding Electron Orbital
Energy
Level K
One spherical orbital (1s)
Electron Shell Diagram
Electron Orbitals y z x 14
Neon
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15. Reactivity of Elements
Valence electrons (e-) – electrons in outermost energy level
Inert (nonreactive) elements have all eight electrons
Octet rule – atoms with 8 e- (2 e- for small atoms) in outer energy level are stable
First energy level can hold 2
All others can hold 8
Atoms want to have full energy levels Ne Li 15
1 valence e-
8 valence e-

16. Reactivity of Elements
Octet rule –
Chemical bonds between atoms are based on unfilled outer energy levels
Atom will want to gain, lose, or share electrons 16

17. Chemical Bonds
Molecules are groups of atoms held together in a stable association
Compounds are molecules containing more than one type of element
Atoms are held together in molecules or compounds by chemical bonds
O2, a diatomic molecule
H2O, a compound 17

19. Covalent bonds Form when atoms share 2 or more valence electrons
Results in no net charge, satisfies octet rule, no unpaired electrons
Strength of covalent bond depends on the number of shared electrons 19

20. covalent bond Single covalent bond Hydrogen gas H H H H H2
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covalent bond
Single covalent bond
Hydrogen gas H H H H H2
Double covalent bond
oxygen gas O O O O O2
Triple covalent bond
Nitrogen gas N N N N N2 20

21. Covalent Bonds Polar covalent bonds Nonpolar covalent bonds
Electrons not shared equally
Example is water
Nonpolar covalent bonds
Electrons shared equally

22. Structural vs. Molecular Formulas
Shows number and types of elements
Structural formulas
Shows arrangement of atoms and types of bonds present 22

23. Carbon and Organic Molecules
Carbon can form up to 4 single covalent bonds
Carbon atoms can form linear chains or rings
These carbon “backbones” form basis of organic molecules
Glucose (C6H12O6) 23

24. Electronegativity Atom’s affinity for electrons
Differences in electronegativity dictate how electrons are distributed in covalent bonds
Nonpolar covalent bonds = equal sharing of electrons
Polar covalent bonds = unequal sharing of electrons 24

25. Water – Polar Covalent Bond

26. 26

27. Chemical reactions
Chemical reactions involve the formation or breaking of chemical bonds
Atoms shift from one molecule to another without any change in number or identity of atoms
Reactants = original molecules
Products = molecules resulting from reaction
6H2O + 6CO → C6H12O O2
reactants products 27

28. Chemical reactions Extent of chemical reaction influenced by
Temperature
Concentration of reactants and products
Catalysts
Many reactions are reversible 28

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30. Water Life is inextricably tied to water
Single most outstanding chemical property of water is its ability to form hydrogen bonds
Weak chemical associations that form between the partially negative O atoms and the partially positive H atoms of two water molecules
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a. Solid
b. Liquid
c. Gas
a: © Glen Allison/Getty Images RF; b: © PhotoLink/Getty Images RF; c: © Jeff Vanuga/Corbis

31. Polarity of water
Within a water molecule, the bonds between oxygen and hydrogen are highly polar
Partial electrical charges develop
Oxygen is partially negative δ+
Hydrogen is partially positive δ– δ δ+ δ+ δ+ 31

32. Hydrogen bonds
Cohesion – polarity of water allows water molecules to be attracted to one another
Attraction produces hydrogen bonds
Each individual bond is weak and transitory
Cumulative effects are enormous
Responsible for many of water’s important physical properties 32

33. Hydrogen atom Water molecule + – Hydrogen bond a. Oxygen atom
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Hydrogen atom
Water molecule + –
Hydrogen bond a.
Oxygen atom
Hydrogen atom
Hydrogen bond + –
An organic molecule 33 b .
Oxygen atom

34. Adhesion – water molecules stick to other polar molecules by hydrogen bonding
Cohesion – water molecules stick to other water molecules by hydrogen bonding 34

35. Properties of water Water has a high specific heat
A large amount of energy is required to change the temperature of water
Water has a high heat of vaporization
The evaporation of water from a surface causes cooling of that surface 35

36. Properties of water Solid water is less dense than liquid water
- Bodies of water freeze from the top down 36

37. 37

38. 4. Water is a good solvent Water dissolves polar molecules and ions38
Hydration shells
Water molecules
Salt crystal
Cl–
Na+ +
 – 38
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39. Both solvent and solute make up a solution
4. Water is a good solvent
When H2O dissolves a polar or ionic compounds like NaCl
water is the solvent
NaCl is the solute
Both solvent and solute make up a solution 39

40. 5. Water organizes nonpolar molecules
Hydrophilic “water-loving”
Hydrophobic “water-fearing”
Water causes hydrophobic molecules to aggregate or assume specific shapes
Hydrophobic molecule 40

42. Acids and bases Pure water
[H+] of 10–7 mol/L
Considered to be neutral
Neither acidic nor basic
[H+] = [OH-]
pH is the negative logarithm of hydrogen ion concentration of solution
pH = -log[H+] pH= –log 10-7 = -(-7)=7 42

43. Acids and Bases - pH Acid (pH < 7) Base (pH > 7)
Any substance that dissociates in water to increase the [H+] (and lower the pH)
The stronger an acid is, the more hydrogen ions it produces and the lower its pH
Base (pH > 7)
Substance that combines with H+ dissolved in water, and thus lowers the [H+] 43

45. Buffers Substance that resists changes in pH Act by
Releasing hydrogen ions when a base is added
Absorbing hydrogen ions when acid is added
Overall effect of keeping [H+] relatively constant 45

46. Buffers Minimize pH Changes
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Buffering range 5 pH 4 3 2 1 1X 2X 3X 4X 5X
Amount of base added 46

47. Buffers
Most biological buffers consist of a pair of molecules, one an acid and one a base
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Water
(H2O)
Carbon
dioxide
(CO2)
Carbonic
acid
(H2CO3)
Bicarbonate
ion
(HCO3–)
Hydrogen
ion
(H+) + + 47