Download Bio 102 Lecture - chapter 2 The Chemical Basis of Life

Chapter 2
The Chemical Basis of Life
1
Outline

Chemical Elements

Atoms








Periodic Table
Isotopes
Electrons and Energy
Molecules and Compounds
Chemical Bonding


Atomic Mass and Atomic Number
Ionic and Covalent
Hydrogen
Properties of Water
Acids and Bases
2
Chemical Elements

Matter:

Matter is defined as anything that has mass and
occupies space

Matter exists in three states: solid, liquid, and gas

All matter (both living and non-living) is composed of
92 naturally-occurring elements

98% of body weight of organisms are primarily
composed of six elements (carbon, hydrogen, nitrogen,
oxygen, phosphorus, and sulfur—acronym CHNOPS)
make up 98% of the body weight of organisms.
3
Atomic Structure


Atom is the smallest unit of an element
Atoms composed of subatomic particles:




Protons - positive charge; found in the nucleus
Neutrons - no charge; found in the nucleus
Electrons - negative charge; found in electron
shell
Atoms contain specific numbers of
protons, neutrons, and electrons.
4
Subatomic Particles
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= proton
= neutron
= electron
a.
b.
Subatomic Particles
Particle
Electric
Charge
Atomic Mass Unit
(AMU)
Location
Proton
+1
1
Nucleus
Neutron
0
1
Nucleus
Electron
–1
0
Electron shell
c.
5
Atomic Symbols

Each element is represented by one or two letters to give them a
unique atomic symbol e.g. H = hydrogen, Na = Sodium, C = Carbon

Atomic number = proton number = Electron number
Atomic mass or mass number = protons and neutrons
The atomic number is above the atomic symbol and the atomic
mass is below the atomic symbol


Atomic
Number
Mass
Number
6
12
6
Carbon
C
Atomic
Symbo
l
6
Periodic Table

Elements grouped in periodic table based
on characteristics



Vertical columns = groups; chemically similar
Horizontal rows = periods;
Atomic mass increases as you move
down a group or across a period.
7
Periodic Table
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
VIII
I
1
atomic number
H
atomic symbol
He
atomic mass
II
III
IV
V
VI
VII
4.003
3
4
5
6
7
8
9
10
Li
Be
B
C
N
O
F
Ne
6.941
9.012
10.81
12.01
14.01
16.00
19.00
20.18
1
12
13
14
15
16
17
18
Na
Mg
Al
Si
P
S
Cl
Ar
22.99
24.31
26.98
28.09
30.97
32.07
35.45
39.95
19
20
31
32
33
34
35
36
K
Ca
Ga
Ge
As
Se
Br
Kr
39.10
40.08
69.72
72.59
74.92
78.96
79.90
83.60
1.008
Periods
2
Groups
8
Isotopes

Isotopes:

Atoms of the same element with a differing numbers of
neutrons (and therefore have different atomic masses).
e.g. see carbon below
Some isotopes spontaneously decay
 Radioactive
 Give
 Can
off energy in the form of rays.
be used as tracers
 Mutagenic
12
6
C
Carbon 12
– Can cause cancer
13
6
C
Carbon 13
14
6
C
Carbon 14
9
Isotopes of Hydrogen
10
Electrons and Energy

Atoms normally have as many electrons as
protons. Opposite charges balance leaving atom
neutral
 Electrons revolve around nucleus in different
shells, labeled from the innermost shell as K, L,
M, N, etc.
11
12
Each shell can have a certain number of
electrons. The K-shell can have 2
Electrons, the L-shell, 8, the M-shell 18, Nshell 32.
 This is calculated by using the formula 2N²,
where N=1 for the K shell, N=2 for the L
shell, N=3 for the M shell, etc.

13
Different shells in the atom
14
Animation
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15
Arrangements of Electrons in an Atom
16
The outermost electron shell determines the reactivity of the element
17
The Octet Rule for Distribution of Electrons
8-electron configuration is stable because this
atom is having 8 valence electrons.
18
Chemical Bonding
If 3 or less electrons in the outer most shell –
Tendency to donate electrons.
 If 5 or more electrons in the outer most shell –
Tendency to receive electrons.
 A ‘chemical bond’ the force of attraction between
atoms to attain stability.


Bonds between atoms are caused by electrons in
outermost shells

The process of bond formation is called a
reaction.
19
Compounds and Molecules


Compound - when atoms of two or more different
elements bond together

CO2, H2O, C6H12O6, etc.

Characteristics dramatically different from constituent
elements
Molecule and compound is used interchangeably

In Biology molecule is used e.g. molecule of water
(H2O) molecule of glucose (C6H12O6 )
20
Compounds and Molecules
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one molecule
indicates 6 atoms
of carbon
indicates 12 atoms
of hydrogen
indicates 6 atoms
of oxygen
21
Types of Bonds: Ionic Bonding


Ionic bond - forms when electrons are
transferred from one atom to another atom.
Octet rule – atoms lose or gain electrons to fill
their outer shells and become more stable

Atoms “want” 8 electrons in outer shell



If have < 4, desire to donate electrons
If have > 4, desire to receive electrons
Consider two elements from opposite ends
of periodic table

Element from right side:



Has 7 electrons in outer shell
“Desperately wants” one more (7+1=8)
Element from left side:


Has only 1 electron in outer shell
“Desperately wants” to donate it (1-1=0=8)
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Types of Bonds: Ionic Bond Example

Sodium (Na):
 Has only 1 electron in its outermost shell
 Chlorine (Cl):
 Has 7 electrons in its outermost shell
 In a reaction between Na and Cl
 Na loses an electron and becomes a positive ion (Na+)
- CATION
 Cl gains an electron and becomes a negative ion (Cl-)
- ANION
 Attraction of oppositely charged ions holds the two
atoms together in an ionic bond.
23
Formation of Sodium Chloride
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Na
Cl
sodium atom (Na)
chlorine atom (Cl)
–
+
Na
sodium ion (Na+)
−
Na+ Cl
Cl
chloride ion (Cl−)
sodium chloride (NaCl)
a.
b.
(Crystals): © Charles M. Falco/Photo Researchers, Inc.; (Salt shaker): © Erica S. Leeds
24
25
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 (in the case of hydrogen, 2
electrons).

When atoms are horizontally closer together in the periodic table

The electrons are not permanently transferred from one atom to the other
like in NaCl

A pair of electrons from the outer shell will “time share” with one atom
and then the other

This also causes the atoms to remain together

Known as covalent bonding.
26
Covalently Bonded Molecules
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The
structural formula
of a molecule indicates a
shared pair of electrons
by a line between the
two atoms e.g. single
covalent bond (H–H),
double covalent bond
(O=O), and triple
covalent bond (N = N).
Structural
Formula
Electron Model
H
H
H
Molecular
Formula
H2
H
a. Hydrogen gas
O
O
O
O
O2
b. Oxygen gas
H
H
H
C
H
H
C
H
CH4
H
H
c. Methane
27
Nonpolar Covalent Bonds

In non-polar covalent bonds, sharing of electrons is
equal.
28
Polar Covalent Bonds

With polar covalent bonds, the sharing of
electrons is unequal.

In H2O - sharing of electrons by oxygen and hydrogen is not
equal; the oxygen atom therefore assumes a partial negative
charge.
29
Animation
Please note that due to differing
operating systems, some animations
will not appear until the presentation is
viewed in Presentation Mode (Slide
Show view). You may see blank slides
in the “Normal” or “Slide Sorter” views.
All animations will appear after viewing
in Presentation Mode and playing each
animation. Most animations will require
the latest version of the Flash Player,
which is available at
http://get.adobe.com/flashplayer.
30
Types of Bonds: Hydrogen Bonds

Water (H2O or H–O–H) is a polar molecule

H’s become slightly +, O slightly –

The H’s of water molecules are attracted to the negative parts of
the oxygen molecules and form hydrogen bond.

This bond is a weak bond because of weak attractive force
between the slightly positive charge of the hydrogen atom of one
molecule and slightly negative charge of another atom

Easily broken.

Found in water, proteins and nucleic acids
31
Hydrogen Bonds
32
Water Molecule
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Electron Model
Ball-and-stick Model
Space-filling Model
O
Oxygen attracts the shared
electrons and is partially negative.
–
O
H
O
H
H
104.5°
H
H
+
H
+
Hydrogens are partially positive.
a. Water (H2O)
+
H
–
O
H
+
hydrogen
bond
b. Hydrogen bonding between water molecules
33
The Chemistry of Water: Heat Capacity


The specific heat is the amount of heat (energy) needed
to raise a one gram of a substance by one degree
Celsius.
Water has a high heat capacity (high specific heat) due
to the hydrogen bonding.

More heat is required to raise water’s temperature than most other
liquids.
34
Properties of Water: High Heat of Vaporization

Heat of Vaporization is the quantity of heat a liquid must
absorb to be converted from the liquid to the gaseous
state. A substance which has a high heat of vaporization
requires more heat to turn from liquid to vapor.


Large numbers of hydrogen bonds must be broken to
evaporate water.
This is why sweating cools.
35
Evaporative Cooling of Animals
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Calories of Heat Energy / g
800
Gas
600
540
calories
400
200
Liquid
80
calories
Solid
0
freezing occurs
0
evaporation occurs
20 40 60 80 100 120
Temperature (°C)
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.
© Grant Taylor/Getty Images
36
Heat Content of Water at Various
Temperatures
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Calories of Heat Energy / g
800
Gas
600
540
calories
400
200
0
Liquid
80
calories
Solid
freezing occurs
0
evaporation occurs
20 40 60 80 100 120
Temperature (°C)
a. Calories lost when 1 g of liquid water freezes and
calories required when 1 g of liquid water
evaporates.
37
Properties of Water: Water as a Solvent

Solutions consist of:



A solvent (the most abundant part) and
A solute (less abundant part) that is dissolved in the solvent
Hydrophilic molecules are soluble in/ interact with/
dissolve in water. Polar compounds and Ionic
readily dissolve in water by forming hydrogen
bonds. Example: glucose, salt

A hydrophobic molecules are insoluble (do not
dissolve) in water. Example: oils or fats
38
Properties of Water: Water as a Solvent
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H
O
O
H
-

An ionic salt
dissolves in water.
Na+
H

O
H
-
- O H
H
+
+
H
O H
H
H
+
Cl–
+
H
H O
39
Properties of Water: Water as a Solvent
Ammonia (NH3)
H O
H
+

A polar molecule
dissolves in water.
-

+
N
+

H
H

H
H
O
H
+

-

-

-

O
H
H
O
H
H
40
Properties of Water: Uniqueness of Ice

Frozen water (ice) is less dense than liquid
water so ice floats in water.

Otherwise, oceans and deep lakes would fill
with ice from the bottom up

Ice acts as an insulator on top of a frozen
body of water

Melting ice draws heat from the environment
41
Water as a Transport Medium
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Water evaporates,
pulling the water
column from the
roots to the leaves.
H2O
Water molecules
cling together and
adhere to sides of
vessels in stems.
Water enters a
plant at root cells.
H2O
42
A Pond in Winter
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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.
43
Properties of Water: Cohesion & Adhesion

Cohesion – When hydrogen bonds hold water
molecules tightly together, it is called cohesion.

Adhesion – When water molecules adhere to polar
surfaces (due to hydrogen bonds), it is called
adhesion.

Cohesion and adhesion allow water be drawn many
meters up a tree in a tubular vessel
44
Properties of Water: Surface Tension

High Surface Tension
 Water molecules at surface hold more tightly than
below surface
 Allows small non-polar objects (like water strider) to sit
on top of water
45

When water ionizes or dissociates, it releases a
small but equal number of hydrogen (H+) ions
and hydroxide (OH-) ions
H–O–H
H+ + OH-
46
Acids and Bases

Acids donate hydrogen ions

Dissociate in water and release hydrogen ions (H+)

e.g. Hydrochloric acid (HCl)
HCl → H+ + Cl

Bases accepts hydrogen ions (H+) or release hydroxide
ions (OH-)

e.g. Sodium hydroxide (NaOH)

In water, it dissociates into Na+ and OH47
Acid and Base
48
pH Scale

The pH is a mathematical way of indicating the number
of H+ ions in a solution.

pH scale used to indicate acidity and alkalinity of a
solution.

Values range from 0 -14;
0 to 6 = Acidic; 7 = Neutral; 8 to 14 = Basic (or
alkaline)

Each unit in pH represents a change of 10X
pH of 4 is 10X as acidic as pH of 5
 pH of 10 is 100X more basic than pH of 8

49
The pH Scale
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
basic
acidic
H+ Ion
Concentration
10 0
10–1
10–2
10–3
10–4
10–5
10–6
10–7
10–8
10–9
10–10
10–11
10–12
10–13
10–14
pH value
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
Examples
hydrochloric acid
stomach acid, lemon
juice
vinegar, cola, beer
tomatoes
black coffee
urine
pure water
seawater
baking soda
Great Salt Lake
household ammonia
household bleach
sodium hydroxide
50
Buffers and pH

When H+ is added to pure water at pH 7, pH goes
down and water becomes acidic

When OH- is added to pure water at pH 7, pH
goes up and water becomes alkaline

Buffers are solutes in water that resist change in
pH

When H+ is added, buffer may absorb, or counter by
adding OH-

When OH- is added, buffer may absorb, or counter by
adding H+
51
Buffers in Biology

Health of organisms requires maintaining pH of
body fluids within narrow limits

Human blood normally 7.4 (slightly alkaline)

Many foods and metabolic processes add or subtract
H+ or OH- ions


Reducing blood pH to 7.0 results in acidosis

Increasing blood pH to 7.8 results in alkalosis

Both life threatening situations
Bicarbonate ion (-HCO3) in blood buffers pH to 7.4
52
Review


Chemical Elements

Atoms

Isotopes

Molecules and Compounds
Chemical Bonding

Ionic and Covalent

Hydrogen

Properties of Water

Acids and Bases
53
Chapter 2: pp. 21-35
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Na
Cl
sodium atom (Na)
chlorine atom (Cl)
–
+
Na
sodium ion (Na+)
Na+ Cl–
10th Edition
Sylvia S. Mader
Basic Chemistry
BIOLOGY
Cl
chloride ion (Cl–)
sodium chloride (NaCl)
a.
b.
(Crystals): © Charles M. Falco/Photo Researchers, Inc.; (Salt shaker): © Erica S. Leeds
PowerPoint® Lecture Slides are prepared by Dr. Isaac Barjis, Biology Instructor
Copyright © The McGraw Hill Companies Inc. Permission required for reproduction or display
54