Essentials of
Biology
Sylvia S. Mader
Chapter 2
Lecture Outline
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2.1 The Nature of Matter
Matter



Anything that takes up space and has mass
Can exist as a solid, liquid, or a gas
Composed of elements
•
Element


•
•
cannot be broken down into another substance by
ordinary chemical means
Pure substance consisting of one type of atom
Only 92 naturally occurring elements
Four most common elements in living organisms
– CHON
Figure 2.1
Elements in
living
organisms
carbon
(C)
18%
oxygen
(O)
65%
© Tim Pannell/Corbis RF
hydrogen (H) 10%
nitrogen (N) 5%
calcium (Ca) 2%
phosphorus (P) 1.1%
other elements,
including sulfur
0.9%
Atomic structure
 Atomic theory
• Elements consist of atoms
• Atom: smallest unit of an element that has the properties of the
element
 Atomic symbols – one or two letters
• H=?
C=?
S= ?
Cl = ?
Na = ?
 Subatomic particles
• Neutrons, n
 no electrical charge, found in nucleus
• Protons, p+
 positive charge, found in nucleus
 determines the ID of an atom
• Electrons , e negative charge, outside of nucleus
 determines the chemical properties of an atom
 Mass number
• sum of protons and neutrons (electrons have nearly zero mass)
Figure 2.2 Two models of helium (He)
+
= proton
–
= neutron
inside nucleus
outside nucleus
–
+
+
+
+
nucleus
–
a.
= electron
b.
Atomic number
 All atoms of an element have this same number of
protons.
 Atoms are electrically neutral:
• How do the number of protons compare to the number of
electrons?
• Periodic table
 Elements in order of their __________________ .
 Atoms arranged in periods (rows) and groups
(columns)
 Elements’ chemical and physical characteristics
recur in predictable manner
• E.g. LiCl, NaCl, KCl; BeCl2, CaCl???
Periodic table of the elements
Location of....
• Metals?
• Nonmetals?
Figure 2.3 A portion of the periodic table
1
Groups
1
8
1
2
H
He
1.008
2
Periods
3
4
2
3
4
5
6
7
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
11
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
Isotopes
 Atoms of the same element that differ in the
number of _________________ .
 Isotopes have the same number of ________
but a different number of ___________ (different
mass numbers)
 Unstable isotopes may decay emitting radiation
 Radioactive isotopes used…
• Can be used as tracer – PET scan
• Can cause damage to cells leading to cancer
• Can be used to sterilize medical equipment
Figure 2.4 PET Scans
thyroid
gland
a.
b.
a: © Biomed Commun./Custom Medical Stock Photo; b(both):Courtesy National Institutes of Health
Figure 2.5
Chemotherapy using isotopes that emit high levels of radiation
a.
b.
a: © Natasja Weitsz/Getty Images; b: © Geoff Tompkinson/SPL/Photo Researchers, Inc.
Arrangements of Electrons
 Located outside the nucleus of an atom in
specific electron shells (energy levels)
 Each shell contains a certain number of electrons
 For atoms up through number 20
• 2 electrons fill first shell.
• 8 electrons fill each additional shell.
 Octet rule for valence shell
• Valence shell: outermost shell
• Atoms are most stable with 8 valence electrons.
 Exceptions: atoms with the only one energy level
• Atoms can give up, accept, or share electrons to have 8.
 The number of electrons in the valence shell
determines the chemical properties of an atom
Atoms of the four elements most abundant in life
First
electron shell:
can hold
2 electrons
Outermost
electron shell:
can hold
8 electrons
Electron
Hydrogen (H)
Atomic number = 1
Carbon (C)
Atomic number = 6
Nitrogen (N)
Atomic number = 7
Oxygen (O)
Atomic number = 8
Orbital Diagrams of the First 18 Elements
1st
Shell
2
2nd
Shell
8
3rd
Shell
8
Figure 2.6 Atoms of the six elements, CHNOPS
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inner shell
electron
electron
shell
H
Hydrogen
1
H
1
nucleus
P
C
N
O
Carbon
Nitrogen
outer
(valence)
shell
Oxygen
Phosphorus
6
7
8
15
C
N
O
P
12
14
16
31
S
Sulfur
16
S
32
Chemical Bonding and Molecules
• Chemical reactions:
– Atoms give up or acquire electrons in order to
complete their _____________ shells
– Result in atoms staying close together to form
molecules
– Chemical bonds hold molecules together
• Ionic Bonds
• Covalent bonds
Types of chemical bonds
 Molecule – group of atoms chemically bonded
together
• O2, H2O, C6H12O6, N2
 Compound – substance containing atoms of more
than one element
• H2O, C6H12O6
 2 types of chemical bonds
• Ionic bonds – giving up or accepting electrons
 Form between metals and nonmetals
• Covalent bonds – sharing electrons
 Form between nonmetals
Ionic Bonds: form between metals and nonmetals
• When an atom loses or
gains electrons, it
becomes electrically
charged. Why?
– Charged atoms are
called ____________
– Ionic bonds are
formed between
oppositely charged
________________
Figure 2.7
Sodium atom (Na)
Chlorine atom (Cl)
Complete
outer shells
Sodium ion (Na)
Chloride ion (Cl)
Sodium chloride (NaCl)
Atoms: electrically neutral
Ions: Electrically charged
(b) Hydrogen ion (H+)
(a) Hydrogen atom (H)
1 electron
No
electron
1 proton
1 proton
No net electrical
charge
(d) Sodium ion (Na+)
(c) Sodium atom (Na)
11
electrons
11 protons
10
electrons
11 protons
No net electrical
charge
Fig. 2.03
Figure 2.7 Formation of sodium chloride
Na
sodium atom (Na)
Cl
chlorine atom (Cl)
Ionic bonding
 Forms when 2 atoms held together by the attraction between
opposite charges
 Sodium has _____electron in valence shell
• Usually _________________an electron to another atom
 Chlorine has ____electrons in valence shell
• Usually ________an electron from another atom
Figure 2.7 continued
Cl
Na
sodium atom (Na)
+
Ions – charged atoms (No. P+ ≠
chlorine atom (Cl)
No. e-)
• Sodium ion, Na+
• 1 more _____________ than electrons
• Chloride ion, Cl• 1 more _____________ than protons
• Ionic compounds often called salts
• Covalent bonding
 2 atoms ____________ electrons
 2 hydrogen atoms can share electrons to fill
their first shell – orbitals overlap.
H
H
Hydrogen gas (H2)
 Structural formula – uses straight lines H-H
• One line indicates 1 pair of shared electrons.
 Molecular formula – shows number of atoms
involved H2
Double covalent bond: sharing 2 pairs of electrons
 Oxygen gas O2 or O=O
O
O
• Triple covalent bond – sharing __?__ pairs of
electrons
 Nitrogen gas N2 or N≡N
• An atom may form bonds with more than one atom…
Covalent bonding in water
Oxygen atom with unfilled
shell
Water molecule (H2O)
Full shell with 8 electrons
– Slightly negative
Covalent
bond
(shared pair
of electron)
+
+ Slightly
positive
Full shells with 2 electrons each
Hydrogen atoms with unfilled shells
Covalent Bonds: form between nonmetallic atoms
Chemical Reactions
• Cells constantly rearrange molecules by breaking
existing chemical bonds and forming new ones
– Such changes in the chemical composition of matter
are called chemical reactions
Hydrogen gas
Oxygen gas
Reactants
Water
Products
Chemical Equations: symbolize chemical reactions
Reactants: on the left
side of the equation
– the starting materials
Products: on the right side of
the equation
– the ending materials (the stuff
produces)
Law of Conservation of Mass
– Chemical reactions do not create or destroy matter—they only
rearrange it!
• Chemical reactions
 Reactants – molecules that participate in reaction
• Shown to left of arrow
 Products – molecules formed by reactions
• Shown to right of arrow
 Equation is balanced if the same number of each type of atom
occurs on both sides of arrow.
• An overall equation for photosynthesis
6 CO2 +
carbon
dioxide
6 H2O
water
C6H12O6
glucose
+
6 O2
oxygen
• Molecular formula for glucose
one molecule
C6H12O6
indicates
6 atoms
of carbon
indicates
12 atoms
of hydrogen
indicates
6 atoms
of oxygen
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2.2 Water’s Importance to Life
• Life on Earth began in water and evolved there for 3
billion years
• The abundance of water is a major reason Earth is
habitable
– Modern life still remains tied to water
– Your cells are composed of 70%–95% water
• Water has unique properties that make it a lifesupporting substance.
• Properties stem from structure of molecule
The Structure of Water
• The water molecule:
– two hydrogen atoms joined to one oxygen atom by
single covalent bonds
H
H
O
Water: a polar molecule
• The electrons of the covalent bonds are shared unequally
between oxygen and hydrogen
– unequal sharing of electrons makes water a polar molecule
– hydrogen atoms: partially positive (d ) Why?
– oxygen atom: partially negative (d -) Why?
(d )
(d )
Figure 2.9 The structure of water
(d )
The Structure of Water
• The polarity of
water results in
weak electrical
attractions between
neighboring water
molecules
()
Hydrogen bond
()
()
()
()
– These interactions
are called
hydrogen bonds
()
()
()
(b)
Figure 2.9 The structure of water
Water’s Life-Supporting Properties
•
The polarity of water molecules and the hydrogen
bonding that results explain most of water’s lifesupporting properties
– Cohesion and adhesion
– High surface tension
– High heat capacity
– High heat of vaporization
– Varying density
Water as the Solvent of Life
• A solution is a liquid consisting of two or more
substances evenly mixed
– The dissolving agent is called the solvent
– The dissolved substance is called the solute
Ion in solution
Salt crystal
Dissolving of
Sodium
Chloride (NaCl)
in Water
Salt
Electrical
attraction
Water molecules dissolve NaCl,
breaking ionic bond
Water
Water
molecules
(H2O)
Hydrogen
bonds
Edge of one
salt crystal
Ionic bond
The Cohesion of Water
• Water molecules
stick together as a
result of hydrogen
bonding
– This is called
cohesion
– Cohesion is vital
for water
transport in
plants
Microscopic tubes
• Surface tension
– is the measure of how difficult it is to stretch or break
the surface of a liquid
– Hydrogen bonds give water an unusually high
surface tension
Figure 2.13
Water Moderates Temperature
• Because of hydrogen bonding, water has a strong
resistance to temperature change
• Water can absorb and store large amounts of heat
while only changing a few degrees in temperature
– Earth’s Oceans cause temperatures to stay within
limits that permit life
• The density of ice is lower than liquid water
– This is why ice floats
Hydrogen bond
Ice
Liquid water
Stable hydrogen bonds
Hydrogen bonds
constantly break and re-form
• Water is a solvent.
 Due to polarity and H-bonding, water
dissolves many substances.
 Hydrophilic – molecules attracted to water
 Hydrophobic – molecules not attracted to
water
 Water causes NaCl to dissociate.
H+
H
O
–
+
+
O–
H H
Cl–
Na+
The salt NaCl dissociates in water.
+
• Cohesion
 Ability of water molecules to cling to each other due to
hydrogen bonding
• Adhesion
 Ability of water molecules to cling to other polar
surfaces
• Allows water to be excellent transport system in
and outside of living organisms.
• Contributes to water transport in plants
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H2 O
Water in water column
Water evaporates, pulling
the water column from
the roots to the leaves.
Water molecules cling
together and adhere to
sides of vessels in stems
and tree trunks.
Water enters a plant
at root cells.
H2 O
© Corbis RF
Figure 2.10 Cohesion and
adhesion of water
molecules
• Water has a high surface tension.
 Water molecules at the surface cling more
tightly to each other than to the air above.
 Mainly due to hydrogen bonding
Figure 2.11 Surface tension of water
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© Claude Nuridsany & Marie Perennou/Photo Researchers, Inc.
• High heat capacity
 The many hydrogen bonds linking water
molecules allow water to absorb heat without
greatly changing its temperature.
 Temperature of water rises and falls slowly.
• High heat of vaporization
 Takes a great deal of energy to break H bonds for
evaporation
 Heat is dispelled as water evaporates.
Figure 2.12 Heat of vaporization
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a.
b.
a: © The McGraw-Hill Companies, Inc./Jill Braaten, photographer; b: © SuperStock RF
• Water is less dense than ice.
 Unlike other substances, water expands as it
freezes.
 Ice floats rather than sinks.
 It makes life possible in water.
 Ice acts as an insulator.
Figure 2.13 Properties of ice
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ice layer
a. Ice
b. Pond
The Biological Significance of Ice Floating
• When water molecules get cold, they move apart,
forming ice
– A chunk of ice has fewer molecules than an equal
volume of liquid water
• Since ice floats, ponds, lakes, and even the oceans
do not freeze solid
– Marine life could not survive if bodies of water froze
solid
2.3 Acids and Bases
• Water dissociates
into an equal
number of
hydrogen ions
(H+) and
hydroxide ions
(OH-)
Figure 2.14 Dissociation
of water molecules
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reproduction or display.
H
H
OH–
H+
O
H
H
O
Copyright © The McGraw-Hill Companies, Inc. Permission required for
reproduction or display.
H
O
H–O–H
H+
water
hydrogen
ion
+
OH–
hydroxide
ion
OH–
H
H+
• Acids
 Common examples are
lemon juice, vinegar,
tomatoes, and coffee.
 Substances that
dissociate in water,
releasing H+ ions
 Adding an acid to water
increases the number of
H+ ions.
HCl
H+ + ClHydrochloric acid
Figure 2.15 Addition of
hydrochloric acid (HCl)
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HCl
OH–
H+
H
H
O
H
Cl–
H+
O
OH–
H
H
O
H
H+
H+
Cl–
• Bases
 Common bases are
ammonia and milk of
magnesia.
 Substances that either take
up hydrogen ions or
release hydroxide ions
 Adding a base to water
either increases the
number of OH- ions or
decreases the number of
H+ ions.
Figure 2.16 Addition of
sodium hydroxide
(NaOH), a base
NaOH
H
NaOH
Na+ + OHSodium hydroxide
OH–
H
O
H
H
O
Na+
H+
OH–
OH–
Na+
H
H
O
H+
OH–
• pH
 Mathematical way to indicate number of
hydrogen ions in solution
 pH scale ranges from 0 to 14
• pH below 7 acidic – more [H+] than [OH-]
• pH above 7 basic – more [OH-] than [H+]
• pH of 7 neutral – [H+] equal to [OH-]
Acids, Bases, and pH
• Acid
 A chemical compound that donates H+ ions to
solutions
 Taste sour
• Base
 A compound that ….
• accepts H+ ions and removes them from solution
• Dissolves in water to produce hydroxide ions, OH
• Taste bitter
pH Scale
The pH scale is used to
describe the acidity of a
solution
Oven cleaner
Household bleach
Household ammonia
• Acidic: pH _?_ 7
Basic
solution
Milk of magnesia
-
H+ _?_ OH
Seawater
Human blood
Pure water
• Basic: pH _?_ 7
H+
Neutral
solution
-
_?_ OH
Tomato juice
• Neutral: pH _?_ 7
H+
Urine
-
__?__ OH
Grapefruit juice
Acidic
solution
Lemon juice;
gastric juice
pH Scale
The pH scale is used to
describe the acidity of a
solution
Oven cleaner
Household bleach
Household ammonia
• Acidic: pH < 7
Basic
solution
Milk of magnesia
-
H+ > OH
Seawater
Human blood
Pure water
• Basic: pH > 7
H+
-
Neutral
solution
< OH
Tomato juice
• Neutral: pH = 7
H+
-
= OH
Urine
Grapefruit juice
Acidic
solution
Lemon juice;
gastric juice
• Buffer
 Chemical or combination of chemicals that
keeps pH within normal limits
 Resist pH change by taking up excess H+ or
OH pH of blood is about 7.4 – maintained by
buffer
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
hydrochloric acid (HCI) 0
Coca-Cola, beer, vinegar 3
[H+]
Acid
lemon juice 2
Increasing [H+]
stomach acid 1
tomatoes 4
black coffee
normal rainwater
urine
saliva
pure water, tears
human blood
seawater
5
6
7
neutral pH
[H+] =
[OH–]
8
baking soda, stomach antacids 9
oven cleaner 13
sodium hydroxide (NaOH) 14
Base
bicarbonate of soda 12
[OH–]
Increasing [OH–]
Great Salt Lake 10
milk of magnesia
household ammonia 11
Figure 2.17 The pH scale
Self-test/Review Questions
Use these questions as a self test and then discuss your responses with
your study group/classmates—your responses will not be collected.
1. Why is carbon dioxide gas, CO2, classified as a compound
but nitrogen gas, N2, is not?
2. Which of the following are compounds? Elements?:
C6H12O6, CH4, O2, Cl2, HCl, MgCl2, Fe, Ca, Ne, NaI, I
3. What is the difference between an atom and an ion? Give
examples of each to support your response.
4. Which subatomic particle determines the identity of an
atom?
5. Which subatomic particle determines the chemical
properties of an atom?
Self-test/Review Questions
6. A carbon atom has 6 protons, and the most common isotope
of carbon has 6 neutrons. A radioactive isotope of carbon
has 8 neutrons. What are the atomic numbers and the mass
numbers of the of the stable and radioactive forms of
carbon?
7. Explain the difference between an ionic and covalent bond
in terms of what happens to the electrons in the outer shell of
the participating atoms.
8. Sodium fluoride, NaF, is often added to toothpaste to both
kill bacteria that cause cavities. It also helps to harden the
enamel of teeth thus helping it resist cavities. Is sodium
fluoride an ionic or covalent compound? How do you know?
Explain your reasoning.
9. Is carbon dioxide an ionic or covalent compound? How do
you know? Explain your reasoning.
Self-test/Review Questions (cont.)
10. Why are the following incorrect structures for the substances
below? Rewrite their structures with the correct number of chemical
bonds.
a.
Carbon dioxide gas: O—C—O
b.
Oxygen gas: O—O
c.
Nitrogen gas: N—N
11. Explain how water’s versatility as a solvent results from the fact
that water is polar molecule.
12. A bottle of Pepsi consists mostly of sugar dissolved in water,
with some carbon dioxide gas that makes fizzy and makes the pH
less than 7. Describe Pepsi using the following terms: solute,
solvent, acidic, aqueous solution
Self-test/Review Questions (cont.)
13. Which of the following are chemical changes? Physical changes? If
possible, write the balanced chemical equation for those that are a
chemical change.
a.
The alcoholic fermentation in Yeast in which yeast produce
ethanol, C2H5OH, and carbon dioxide, CO2, from the sugar
glucose, C6H12O6
b.
Water boils to form steam
c.
The healing of a cut finger
d.
Cutting a piece of wood with a saw
e.
Potassium metal, K, and chlorine gas (Cl2) combine to form
potassium chloride.
f.
The rusting of iron, Fe, to produce rust, iron (III) oxide (Fe2O3)
Self-test/Review Questions (cont.)
14. Which of these is not a subatomic particle? a) proton; b) ion; c)
neutron; d) electron
15. The outermost electron shell of every Noble Gas element
(except Helium) has ___ electrons. a) 1; b) 2; c) 4; d) 6; e) 8
16. An organic molecule is likely to contain all of these elements
except ___. a) C; b) H; c) O; d) Ne; e) N
17. The chemical bond between water molecules is a ___ bond. a)
ionic; b) polar covalent; c) nonpolar covalent; d) hydrogen
18. A solution with a pH of 7 has ___ times more H ions than a
solution of pH 9. a) 2; b) 100; c) 1000; d) 9; e) 90
19. The type of chemical bond formed when electrons are shared
between atoms is a ___ bond. a) ionic; b) covalent; c)
hydrogen
Self-test/Review Questions (cont.)
20. The type of chemical bond formed when oppositely charged
particles are attached to each other is a ___ bond. a) ionic; b)
covalent; c) hydrogen
21. Carbon has an atomic number of 6. This means it has ___. a)
six protons; b) six neutrons; c) six protons plus six neutrons; d)
six neutrons and six electrons
22. Each of the isotopes of hydrogen has ___ proton(s). a) 3; b) 1;
c) 2; d) 92; e) 1/2
23. A molecule is ___. a) a mixture of various components that can
vary; b) a combination of many atoms that will have different
ratios; c) a combination of one or more atoms that will have a
fixed ratio of its components; d) more important in a chemistry
class than in a biology class
Estimating the Size of an Object Viewed with a Microscope
•
Calculate the length and width of the following microscopic object
in both millimeters (mm) and micrometers (mm). 1 mm = 1000 mm
•
Base your calculations on the following field sizes:
Low power (40x):
4.5 mm
Medium power (100x): 1.8 mm
High power (400x):
0.45 mm
Object viewed at medium power
(100x)
Remember: Field size decreases by the same factor as
the magnification increases!
Estimating the Size of an Object Viewed with a Microscope
•
Calculate the length and width of the following microscopic object
in both millimeters and micrometers. 1 mm = 1000 mm
•
Base your calculations on the following hypothetical field sizes:
Low power (30x):
4.0 mm = ___mm
Medium power (180x): ___mm = ___mm
High power (300x):
___mm = ___mm
Object viewed at high power
(300x)
Remember: Field size decreases by the same factor as
the magnification increases!