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Essentials of BiologySylvia S. Mader
Chapter 2Lecture Outline
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
2.1 The Nature of MatterMatter
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
calcium (Ca) 2%nitrogen (N) 5%
hydrogen (H) 10%oxygen
(O)65%
carbon (C)
18%
phosphorus (P) 1.1%
other elements,including sulfur 0.9%
© Tim Pannell/Corbis RF
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)
inside nucleus outside nucleus
= proton = neutron = electron
nucleus
a. 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
1
2
3
4
3
Li6.941
11
Na22.99
19
K39.10
4
Be9.012
12
Mg24.31
20
Ca40.08
5
B10.81
13
Al26.98
31
Ga69.72
6
C12.01
14
Si28.09
32
Ge72.59
7
N14.01
15
P30.97
33
As74.92
8
O16.00
16
S32.07
34
Se78.96
9
F19.00
17
Cl35.45
35
Br79.90
GroupsP
erio
ds
8
2
He4.003
10
Ne20.18
18
Ar39.95
36
Kr83.60
2 3 4 5 6 71.008
H1
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
thyroidgland
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
Electron
Firstelectron shell:can hold2 electrons
Outermostelectron shell:can hold8 electrons
Carbon (C)Atomic number = 6
Nitrogen (N)Atomic number = 7
Oxygen (O)Atomic number = 8
Hydrogen (H)Atomic number = 1
Orbital Diagrams of the First 18 Elements
2
8
8
1st Shell
2nd Shell
3rd Shell
Figure 2.6 Atoms of the six elements, CHNOPS
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
P
Phosphorus
O
Oxygen
N
Nitrogen
C
Carbon
H
Hydrogen
electron
1
H1
6
C12
7
N14
8
O16
15
P31
electronshell
nucleus
inner shell
outer(valence)shell
S
Sulfur16
S32
• 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
Chemical Bonding and Molecules
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
• When an atom loses or gains electrons, it becomes electrically charged. Why?
– Charged atoms are called ____________
– Ionic bonds are formed between oppositely charged ________________
Ionic Bonds: form between metals and nonmetals
Sodium atom (Na) Chlorine atom (Cl)
Completeouter shells
Sodium ion (Na) Chloride ion (Cl)
Sodium chloride (NaCl)Figure 2.7
(a) Hydrogen atom (H)
(c) Sodium atom (Na)
(b) Hydrogen ion (H+)
(d) Sodium ion (Na+)
1 electron
1 proton
No net electrical charge
11 electrons
11 protons
No net electrical charge
No electron
1 proton
10 electrons
11 protons
Fig. 2.03
Atoms: electrically neutral Ions: Electrically charged
Figure 2.7 Formation of sodium chloride
chlorine atom (Cl)sodium atom (Na)
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
Na Cl
Figure 2.7 continued
Cl
chlorine atom (Cl)
Na
sodium atom (Na) +
Ions – charged atoms (No. P+ ≠ 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.
Structural formula – uses straight lines H-H• One line indicates 1 pair of shared electrons.
Molecular formula – shows number of atoms involved H2
Hydrogen gas (H2)
H H
Double covalent bond: sharing 2 pairs of electrons
Oxygen gas O2 or 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…
O O
Covalent bonding in water
Full shell with 8 electrons– Slightly negative
Hydrogen atoms with unfilled shells
Full shells with 2 electrons each
+Slightly positive
+
Covalent bond(shared pairof electron)
Oxygen atom with unfilled shell
Water molecule (H2O)
Covalent Bonds: form between nonmetallic atoms
• 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
Chemical Reactions
Hydrogen gas Oxygen gas Water
Reactants Products
Reactants: on the left side of the equation – the starting
materials
Chemical Equations: symbolize chemical reactions
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
• Molecular formula for glucose
6 CO2
carbondioxide
6 H2Owater
C6H12O6
glucose+ 6 O2
oxygen+
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
C6H12O6
one molecule
indicates6 atomsof carbon
indicates6 atomsof oxygen
indicates12 atomsof hydrogen
• 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 life-supporting substance.
• Properties stem from structure of molecule
2.2 Water’s Importance to Life
• The water molecule:
– two hydrogen atoms joined to one oxygen atom by single covalent bonds
The Structure of Water
H
O
H
• 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 ( ) Why?
– oxygen atom: partially negative ( -) Why?
()( )
( )
Water: a polar molecule
Figure 2.9 The structure of water
• The polarity of water results in weak electrical attractions between neighboring water molecules
– These interactions are called hydrogen bonds (b)
()
Hydrogen bond()
()()
()
()
()
()
The Structure of Water
Figure 2.9 The structure of water
• The polarity of water molecules and the hydrogen bonding that results explain most of water’s life-supporting properties
– Cohesion and adhesion
– High surface tension
– High heat capacity
– High heat of vaporization
– Varying density
Water’s Life-Supporting Properties
• A solution is a liquid consisting of two or more substances evenly mixed
Water as the Solvent of Life
– The dissolving agent is called the solvent
– The dissolved substance is called the solute
Ion in solutionSalt crystal
Dissolving of Sodium Chloride (NaCl) in Water
Salt
Water
Electricalattraction
Watermolecules(H2O)
Hydrogenbonds
Edge of onesalt crystal
Ionic bond
Water molecules dissolve NaCl,breaking ionic bond
• Water molecules stick together as a result of hydrogen bonding
– This is called cohesion
– Cohesion is vital for water transport in plants
The Cohesion of Water
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
• 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
Water Moderates Temperature
• The density of ice is lower than liquid water
– This is why ice floats
Hydrogen bond
Liquid water
Hydrogen bondsconstantly break and re-form
Ice
Stable hydrogen bonds
• 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.
++ –
The salt NaCl dissociates in water.
Na+ Cl–
++–O
H HH H
O
• 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
Figure 2.10 Cohesion and adhesion of water
molecules
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Wat
er in
wat
er c
olu
mn
Water evaporates, pullingthe water column fromthe roots to the leaves.
Water molecules clingtogether and adhere tosides of vessels in stemsand tree trunks.
H2O
Water enters a plantat root cells.
H2O
© Corbis RF
• 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 waterCopyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
© 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 vaporizationCopyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
b.
a.
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
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
a. Ice
b. Pond
ice layer
• 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
The Biological Significance of Ice Floating
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
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
OH H
OH H
OH H
OH–
OH–H+
H+
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
H – O – H
water
H+ + OH–
hydrogenion
hydroxideion
• 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+ + Cl-
Hydrochloric acid
Figure 2.15 Addition of hydrochloric acid (HCl)
OH H
OH H
OH H
OH–
Cl–
Cl–
OH–
HCl
H+
H+ H+
H+
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
• 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.
NaOH Na+ + OH-
Sodium hydroxide
Figure 2.16 Addition of sodium hydroxide (NaOH), a base
O
H H
O
H H
O
H H
OH–
OH–
OH–
OH–
NaOH
Na+
Na+
H+
H+
• 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-]
• 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
Acids, Bases, and pH
Basicsolution
Neutralsolution
Acidicsolution
Oven cleaner
Household bleach
Household ammonia
Milk of magnesia
Seawater
Human bloodPure water
Urine
Tomato juice
Grapefruit juice
Lemon juice;gastric juice
• Acidic: pH _?_ 7
H+ _?_ OH-
• Basic: pH _?_ 7
H+ _?_ OH-
• Neutral: pH _?_ 7
H+ __?__ OH-
pH ScaleThe pH scale is used to describe the acidity of a
solution
Basicsolution
Neutralsolution
Acidicsolution
Oven cleaner
Household bleach
Household ammonia
Milk of magnesia
Seawater
Human bloodPure water
Urine
Tomato juice
Grapefruit juice
Lemon juice;gastric juice
• Acidic: pH < 7
H+ > OH-
• Basic: pH > 7
H+ < OH-
• Neutral: pH = 7
H+ = OH-
pH ScaleThe pH scale is used to describe the acidity of a
solution
• 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
Figure 2.17 The pH scaleCopyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Acid
Base
Increasin
g [H
+]In
creasing
[OH
–]
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
hydrochloric acid (HCI)
stomach acid
lemon juice
Coca-Cola, beer, vinegar
tomatoes
black coffee
urine
pure water, tears
seawater
baking soda, stomach antacids
Great Salt Lake
household ammonia
bicarbonate of soda
oven cleaner
sodium hydroxide (NaOH)
normal rainwater
saliva
human blood
milk of magnesia
[H+]
[OH–]
neutral pH[H+] =
[OH–]
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 (m). 1 mm = 1000 m
• 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 m
• Base your calculations on the following hypothetical field sizes:
Low power (30x): 4.0 mm = ___m
Medium power (180x): ___mm = ___m
High power (300x): ___mm = ___m
Object viewed at high power (300x)
Remember: Field size decreases by the same factor as the magnification increases!