Sylvia S. Mader
Michael Windelspecht
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• This chapter should be a review.
We will go over this material very quickly.
• Read second ½ of Ch 2 for tomorrow.
• You should also read appropriate pages of CliffsAP
• Take the Sample test on page 27, grade it and turn in by Sept 2.
• Topics
2.1 Chemical Elements
2.2 Compounds and Molecules
2.3 Chemistry of Water
2.4 Acids and Bases
Mills Biology 2001

.

Elements
• An element is a substance that cannot be broken down into substances with different properties; composed of one type of atom.
• Six elements make up 95% of the body weight of organisms:
(acronym CHNOPS)
 Carbon
 Hydrogen
 Nitrogen
 Oxygen
 Phosphorus
 Sulfur

Composition of Earth’s Crust and Its Organisms
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
60
Earth’s crust organisms
40
20
0
Fe Ca K S P Si Al Mg Na O N C H
Element
© Gunter Ziesler/Peter Arnold, Inc.
5

Atoms
• An atom is the smallest part of an element that displays the property of the element.
• An element and its atom share the same name.
 Composed of subatomic particles: protons, neutrons, electrons
 Central nucleus
• Protons - positively charged
• Neutrons - no charge
 Orbiting clouds around nucleus (electron shells)
• Electrons - negatively charged, very low massnegligible in calculations

a.
c.
Particle
Proton
Neutron
Electron
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= proton
= neutron
= electron b.
+1
Subatomic Particles
Electric
Charge
Atomic Mass Unit
(AMU)
1
0
–1
1
0
Location
Nucleus
Nucleus
Electron shell
7

Atomic Number and Mass Number
• Each element is represented by one or two letters to give it a unique atomic symbol.
 H = hydrogen, Na = sodium, C = carbon
• The atomic number is equal to the number of protons in each atom of an element.
• The mass number of an atom is equal to the sum of the number of protons and neutrons in atom’s nucleus.
 The atomic mass is approximately equal to the mass number.

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
mass number atomic number
Atomic number?
Atomic mass?
atomic symbol
9

Periodic Table
• Atoms of an element are arranged horizontally by increasing atomic number in rows called periods.
• Atoms of an element arranged in vertical columns are called groups.
 Atoms within the same group share the same binding characteristics.
• Atoms shown in the periodic table are electrically neutral.
 Therefore, the atomic number tells you the number of electrons as well as the number of protons.

1
2
3
4
I
1
H
1.008
A Portion of the Periodic Table
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
atomic number atomic symbol atomic mass
3
Li
6.941
II
4
Be
9.012
III
5
B
10.81
IV
6
C
12.01
V
7
N
14.01
VI
8
O
16.00
VII
9
F
19.00
11
Na
12
Mg
22.99
24.31
13
Al
26.98
14
Si
28.09
15
P
30.97
16
S
32.07
17
Cl
35.45
19
K
39.10
20
Ca
40.08
31
Ga
69.72
32
Ge
72.59
33
As
74.92
34
Se
78.96
35
Br
79.90
VIII
2
He
4.003
10
Ne
20.18
18
Ar
39.95
36
Kr
83.60
Groups
11

2.1 Chemical Elements
Know highlighted elements
Mills Biology 2001

Isotopes
 Some isotopes spontaneously decay
• Radioactive isotopes give off energy in the form of rays and subatomic particles
• Can be helpful or harmful
12
6
C
Carbon 12
13
6
C
Carbon 13
14
6
C
Carbon 14

Uses of Low
Levels of Radiation
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Low levels of radioactive
Iodine are used to “see” the thyroid gland. a.
larynx thyroid gland trachea
In a PET scan an isotope of glucose is used to detect active tissues.
b.
a: © Biomed Commun./Custom Medical Stock Photo; b (left) : Courtesy National Institutes of Health; b: (right) © Mazzlota et al./Photo Researchers, Inc
14

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
a.
b.
a: (Peaches): © Tony Freeman/PhotoEdit; b: © Geoff Tompkinson/SPL/Photo Researchers, Inc.
15

Electrons and Energy
• Electrons are attracted to the positively charged nucleus, thus it takes energy to hold electrons in place.
• It takes energy to push them away and keep them in their own shell.
 The more distant the shell, the more energy it takes to hold in place.
• Electrons have energy due to their relative position (potential energy).
• Electrons determine chemical behavior of atoms .
16

The Distribution of Electrons
• The Bohr model is a useful way to visualize electron location.
 Electrons revolve around the nucleus in energy shells (energy levels).
 For atoms with atomic numbers of 20 or less, the following rules apply:
• the first energy shell can hold up to 2 electrons
• each additional shell can hold up to 8 electrons
• each lower shell is filled first before electrons are placed in the next shell
 These rules cover most of the biologically significant elements.
17

Valence Electrons
• The outermost energy shell is called the valence shell .
• The valence shell is important because it determines many of an atom’s chemical properties.
• The octet rule states that the outermost shell is most stable when it has eight electrons.
 Exception : If an atom has only one shell, the outermost valence shell is complete when it has two electrons.
• The number of electrons in an atom’s valence shell determines whether the atom gives up, accepts, or shares electrons to acquire eight electrons in the outer shell.
– Atoms that have their valence shells filled with electrons tend to be chemically stable.
– Atoms that do not have their valence shells filled with electrons are chemically reactive.
18

H hydrogen
1H
1
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valence shell electron electron shell nucleus
C carbon
12
6
C
N nitrogen
14
7
N
O oxygen
16
8
O
P phosphorus
31
15
P
S sulfur
32
16
S
19

2.2 Compounds and Molecules
• A molecule is two or more elements bonded together.
• A compound is a molecule containing at least two different elements bonded together.
 CO
2
, H
2
O, C
6
H
12
O
6
, etc.
In Biology – usually used interchangeably (molecule and compound)
• A formula tells you the number of each kind of atom in a molecule. one molecule indicates 6 atoms of carbon indicates 12 atoms of hydrogen indicates 6 atoms of oxygen
20

• Bonds that exist between atoms in molecules contain energy.
• Bonds between atoms are caused by the interactions between electrons in outermost energy shells.
• The process of bond formation is called a chemical reaction .
21

• Covalent bonds
• Ionic bonds
• Hydrogen bonds
• An ion is an atom that has lost or gained an electron.
• An ionic bond forms when electrons are transferred from one atom to another atom and the oppositely charged ions are attracted to one.
 Example: formation of sodium chloride
• Salts are solid substances that usually separate and exist as individual ions in water.
22

Formation of Sodium Chloride
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Na sodium atom (Na)
Cl chlorine atom (Cl)
23

Na sodium atom (Na)
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Cl chlorine atom (Cl)
24

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Na Cl sodium atom (Na) chlorine atom (Cl)
+
–
Na Cl sodium ion (Na
+
) chloride ion (Cl sodium chloride (NaCl)
–
)
25

Na sodium atom (Na)
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Cl chlorine atom (Cl)
Na
+
–
Na + Cl
–
Cl a.
sodium ion (Na + ) chloride ion (Cl sodium chloride (NaCl)
–
) b.
b (middle): © Evelyn Jo Johnson; b (right): © Evelyn Jo Johnson
26

Types of Bonds: Covalent Bonds
• Covalent bonds result when two atoms share electrons so each atom has an octet of electrons in the outer shell.
 Note: In the case of hydrogen, the outer energy shell is complete when it contains 2 electrons.
• In a nonpolar covalent bond electrons are shared equally between atoms.
 Examples: hydrogen gas, oxygen gas, methane
27

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Electron Model
H H
Structural
Formula
H H
Molecular
Formula
H
2 a. Hydrogen gas
O b. Oxygen gas
O O O O
2
H
H
C
H
H
H
H C H
H
CH
4 c. Methane
28

• In a polar covalent bond electrons are shared unequally .
 Example: water
• Electronegativity is the ability of an atom to attract electrons towards itself in a chemical bond.
 In water, the oxygen atom is more electronegative than the hydrogen atoms and the bonds are polar.
Oxygen is partially negative
H
O
H
Hydrogens are partially positive
29

30

• Water is a polar molecule.

The shape of a water molecule and its polarity make hydrogen bonding possible.
• A hydrogen bond is a weak attraction between a slightly positive hydrogen atom and a slightly negative atom.
 Can occur between atoms of different molecules or within the same molecule
 A single hydrogen bond is easily broken while multiple hydrogen bonds are collectively quite strong.

Help to maintain the proper structure and function of complex molecules such as proteins and DNA
31

H
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Electron Model
2.3 Chemistry of Water
Fig. 2.9
O
H
Ball-and-stick Model

+
H

-
O hydrogen bond
H

+
O
H
104.5˚
H b. Hydrogen bonding between water molecules
Space-filling Model
Oxygen attracts the shared electrons and is partially negative.

-
O

+
H H

+
Hydrogens are partially positive.
a. Water (H
2
O)
32

2.3 Chemistry of Water
• Properties of water
 Makes up 60-80% of most living cells
 Most abundant compound in living tissue
 Makes up about 66% of weight of an adult human
• About 75% of your brain is water!
• Why do we need it?
 Cohesion, adhesion, surface tension, heat transfer, solvent, transport, properties of freezing.
Mills Biology 2012

• Water molecules cling together because of hydrogen bonding.
 This association gives water many of its unique chemical properties.
• Water has a high heat capacity.
 The presence of many hydrogen bonds allow water to absorb a large amount of thermal heat without a great change in temperature.
 The temperature of water rises and falls slowly.
• Allows organisms to maintain internal temperatures
34

800
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Gas
600
400
540 calories
200
Liquid
0
Solid
80 calories freezing occurs
0 evaporation occurs
20 40 60
Temperature (
°C)
80 100 120 a. Calories lost when 1 g of liquid water freezes and calories required when
1 g of liquid water evaporates.
b. Bodies of organisms cool when their heat is used to evaporate water.
b: © Grant Taylor/Stone/Getty Images
35

• Water has a high heat of vaporization.
 Hydrogen bonds must be broken to evaporate water.
 Bodies of organisms cool when their heat is used to evaporate water.
• Water is a good solvent.
 Water is a good solvent because of its polarity.
 Polar substances dissolve readily in water.
 Hydrophilic molecules dissolve in water.
 Hydrophobic molecules do not dissolve in water.
36

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
O
H H
H
O
H
H
O
H
Na +
An ionic salt dissolves in water.
O
H
H
O
H
Cl
–
H
H
O
37

• Cohesion
 Sticks to itself
 Polarity
• Caused by H bonds
• Without these bonds water would boil at -80C (instead of +100) and freeze at –100C (instead of 0C) this would not be compatible with life
• Adhesion
 Sticks to other substances
 polarity
 “add” hesion
 Capillary action
High surface tension due to hydrogen bonding of water.

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Water evaporates, pulling the water column from the roots to the leaves.
Water molecules cling together and adhere to sides of vessels in stems.
Water enters a plant at root cells.
H
2
O
H
2
O
39

• Ice is less dense than liquid water.
 At temperatures below 4 °C, hydrogen bonds between water molecules become more rigid but also more open.
 Water expands as it reaches 0 °C and freezes.
 Ice floats on liquid water.
• Acts as an insulator on top of a frozen body of water
Ice floats, I’m glad!
40

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
ice lattice
1.0
liquid water
0.9
0 4
Temperature (ºC)
100
41

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
ice layer
Protists provide food for fish.
River otters visit ice-covered ponds.
Aquatic insects survive in air pockets.
Freshwater fish take oxygen from water
.
Common frogs and pond turtles hibernate.
42

• pH is a measure of hydrogen ion concentration in a solution.
• When water ionizes or dissociates, it releases an equal number of hydrogen
(H + ) ions and hydroxide (OH ) ions.
• Acids are substances that dissociate in water, releasing hydrogen ions.
• Bases are substances that either take up hydrogen ions (H+) or release hydroxide ions (OH−).
43

Soren Sorenson
1909 Beer Brewer
Developed pH scale
Mills Biology 2012

• The pH scale is used to indicate the acidity or basicity (alkalinity) of a solution.
 Values range from 0-14
• 0 to <7 = Acidic
• 7 = Neutral
• >7 to 14 = Basic (or alkaline)
• Actually a measure of H ion concentration
(negative log of the H ion concentration) (log is the power to which 10 must be raised to get a certain number)
 measures grams of ion/liter of solution
 pH is a shorthand measurement scale
 wide range
 10 fold change for each change in pH
• concentration of 0.1gm H ion/L = pH 1
 move decimal to left for each increase of 1 pH pH of 1 = 10 -1 pH of 6 = 10 -6
45

Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
H +
1
0
2
3
4
5
6 7 8
Acidic
9
10
11
12
13
14
OH
–
Basic
46

• A buffer is a chemical or a combination of chemicals that keeps pH within normal limits.
• Health of organisms requires maintaining the pH of body fluids within narrow limits.
 Human blood is normally pH 7.4 (slightly basic)
• If blood pH drops below 7.0, acidosis results
• If blood pH rises above 7.8, alkalosis results
• Both are life-threatening situations.
 Body has built-in mechanisms to prevent pH changes.
Read Connecting the Concepts with the Big Ideas pg 34
47

Mills Biology 2012