1. Bio 1010
Dr. Bonnie A. Bain
Office: Science 108

2. CHAPTER 2: CHEMISTRY
Why talk about Chemistry in a Biology Class??
Many concepts in this class involve a basic understanding of Chemistry
Elements:
92 naturally occurring elements
25 of which are essential to all life

3. Figure 1.3_3
Life has
different levels
of complexity

4. Primary (96% of body weight) H, C, N2, O2
Elements found in the Human Body
Primary (96% of body weight)
H, C, N2, O2
Secondary (3.9% of body weight)
Ca, P, K, S, Na, Cl, Mg
Trace elements (less than 0.01% of body weight)
Fe, Cr, I, Si, etc.

5. Trace elements: less than 0.01%
Fig. 2-02
Carbon (C): 18.5%
Oxygen (O):
65.0%
Calcium (Ca): 1.5%
Phosphorus (P): 1.0%
Potassium (K): 0.4%
Sulfur (S): 0.3%
Sodium (Na): 0.2%
Chlorine (Cl): 0.2%
Magnesium (Mg): 0.1%
Hydrogen (H):
9.5%
Trace elements: less than 0.01%
Boron (B)
Manganese (Mn)
Chromium (Cr)
Molybdenum (Mo)
Nitrogen (N):
3.3%
Cobalt (Co)
Selenium (Se)
Copper (Cu)
Silicon (Si)
Fluorine (F)
Tin (Sn)
Iodine (I)
Vanadium (V)
Iron (Fe)
Zinc (Zn)

6. Figure 2.3

7. Fig. 2-03
GOITER

8. CHAPTER 2: CHEMISTRY Matter: Occupies space, has mass, has weight
4 states: solid, liquid, gas, plasma
composed of elements
Energy: “The capacity to do work”
Potential Energy: Stored energy (eg. battery)
Kinetic Energy: Energy of motion
(falling objects, explosions, etc.)

9. Different Forms of Energy
Chemical Energy: Energy stored in chemical bonds
Most important kind for biology
Radiant Energy: Heat and light
Electrical Energy: the flow of electrons

10. PERIODIC TABLE OF THE ELEMENTS
Figure 2.2
PERIODIC TABLE OF THE ELEMENTS

11. Bohr Atom (named after Neils Bohr)
Figure 2.5
Bohr Atom (named after Neils Bohr)

12. ATOMIC NUMBER: Number of protons
Hydrogen has 1 proton, Atomic number = 1
Helium has 2 protons, Atomic number = 2
Oxygen has 8 protons, Atomic number = 8
ATOMIC WEIGHT: Protons + neutrons
Measured in atomic mass units
Carbon has 6 protons + 6 neutrons
Atomic weight = 12
(Electrons don't weigh much, not usually counted)

13. PERIODIC TABLE OF THE ELEMENTS
Figure 2.2
PERIODIC TABLE OF THE ELEMENTS

14. 126C
Top number = Atomic Weight
Bottom number = Atomic Number

16. Electrons are arranged in shells or orbits around the nucleus

17. Electron Shells or Orbitals
Innermost shell can hold a max. of 2 electrons
Next shell holds a max. of 8 electrons, etc.
Carbon has 6 protons and 6 electrons
Innermost shell has 2 electrons
Outer shell has 4 electrons
The outer shell can potentially hold 4 more electrons
This allows Carbon to combine with other elements

18. When the outermost shell is full, the element is stable
To achieve stability, atoms can either empty their outermost electron shell or fill it to the max.
To do this, they can:
Give up an electron or
Accept an electron
Share an electron

19. Valence Number The number of extra electrons in the outer shell or
The number of deficient electrons in the outer shell
This determines the combining capacity
If an element has 1 electron in the outer shell, it can combine with another element that has 7 in its outer shell.

20. If an element has 1 electron in the outer shell, it can combine with another element that has 7 in its outer shell.
Sodium (Na)
valence # = +1
Electron donor
Chlorine (Cl)
valence # = -1
Electron acceptor

24. Two atoms “share” electrons Example: Hydrogen is always found as H2
CHEMICAL BONDS
IONIC BOND
electron donor + electron acceptor
Na+1 + Cl-1 = NaCl
Held together by the attraction of the + and - charges
COVALENT BOND
Two atoms “share” electrons
Example: Hydrogen is always found as H2
(see polar vs non-polar)

25. CHEMICAL BONDS (cont.)
HYDROGEN BOND
Occurs when a covalently bonded hydrogen is positive and is attracted to a negatively charged atom nearby
These bonds are relatively weak, can be broken easily, usually indicated by dotted lines
Hydrogen bonds form in water, also in large protein molecules

26. H2O
(Water)

27. POLAR VS NON-POLAR MOLECULES
In a covalent bond, the atoms share the electron
NON-POLAR COVALENT BOND
Neither atom attracts the shared electron more strongly
Bonds between 2 identical atoms are always non-polar
The single covalent bond between C and H is always non-polar

29. POLAR COVALENT BOND
The sharing of the electron is unequal
This affects the chemical properties of the molecule
The “stronger” atom has a partial negative charge
The “weaker” atom has a partial positive charge
Example: water (H2O)
This structure gives water its unique properties
(universal solvent-can dissolve nearly anything)

31. Figure 2.11
WATER

32. Extra Photo 02.11x1

33. 1. Excellent solvent for polar (charged) molecules and ionic salts
PROPERTIES OF WATER
1. Excellent solvent for polar (charged) molecules and ionic salts
(ie. Many things dissolve readily in water)
Example: NaCl (salt)
Negative ends of water molecules are attracted to the Na+
Positive ends of water molecules are attracted to the Cl-
This pulls the NaCl molecule apart

35. PROPERTIES OF WATER (cont.)
Hydrophilic (water loving)
Ions and molecules which interact with water
Hydrophobic (water hating)
Non-ionized and non-polar molecules
These do not interact with water

36. PROPERTIES OF WATER (cont.)
2. Water molecules are cohesive (stick together)
This allows water to do things like fill a glass
Also gives water a high surface tension
The water molecules cling together because of the hydrogen bonds

37. Figure 2.13
Surface
tension

38. PROPERTIES OF WATER (cont.)
3. Water has a high heat of vaporization:
Water absorbs lots of heat as it warms and gives off lots of heat as it cools
Reason: hydrogen bonds
This lets us keep body temp within certain limits
Also lets lakes, rivers, oceans store vast amounts of heat, thus moderating the environment

39. PROPERTIES OF WATER (cont.)
4. Why ice floats:
When most liquids get cold, their molecules move closer together and eventually freeze solid
Water behaves differently:
As it cools down, the molecules move farther apart, eventually forming ice
A chunk of ice has fewer molecules than an equivalent amount of liquid water
It floats because it is less dense than liquid water

40. Figure 2.15

41. ACIDS, BASES, SALTS
Acid
A substance which dissociates into one
or more H+ ions and one or more negative ions
or an acid can be defined as a proton donor
HCl H+ + Cl-
Examples: hydrochloric acid, lemon juice, vinegar, tomato juice, coffee

42. ACIDS, BASES, SALTS
Base
A substance which dissociates into one or more
hydroxide ions (OH-)
or a base can be defined as a proton acceptor
NaOH Na+ + OH-
Example: Sodium hydroxide

43. ACIDS, BASES, SALTS Salts These dissociate when dissolved in water
Acid + base = a salt + water
Example:
HCl + KOH = KCl + H2O

44. pH Scale 01 2 3 4 5 6 7 8 9 10 11 12 13 14 Acidic Neutral Basic
Acidic Neutral Basic
Water is neutral (H+, OH-)
pH = 1/log [H+]
negative log of hydronium (H+) ion concentration (molarity)

46. Figure 2.17

47. CHEMICAL REACTIONS Two parts: reactant and product
Synthesis Reaction (Anabolism)
Atoms or molecules combine, form new bonds
A + B = AB
Decomposition Reaction (Catabolism) Chemical bonds break here
AB = A + B
Exchange Reaction
Part synthesis; part decomposition
AB + CD = AD + BC