Transcript chapt02_lecture - Huntsville High School

Chapter 2
Biology is the study life and all
things are made of substances
(stuff)
Chemistry is the study of the
properties of these substances
(stuff)
So we will look at what this stuff is
made of and how they interact and
function.
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=cnXV7Ph3WPk
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2.1 Atoms
• Organisms are chemical machines
– one must know chemistry in order to
understand biology
• Any substance in the universe that has
mass and occupies space is comprised of
matter
– all matter is made up of atoms
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2.1 Atoms
• An atom is the smallest particle into which
a substance can be divided and still retain
its chemical properties.
• All atoms have the same basic structure
– at the core is a dense nucleus
comprised of two subatomic particles
• protons (positively charged)
• neutrons (no associated charge)
– orbiting the nucleus is a cloud of another
subatomic particle
• electrons (negatively charged)
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• An atom can be characterized by its proton
number or by its overall mass
Mass is constant and measured in grams
Weight is based on gravity and measured in Newtons.
– atomic number
• the number of protons in the nucleus
• atoms with the same atomic number exhibit
the same chemical properties and are
considered to belong to the same element
(Element is a substance that cannot be broken down into
any other substance by ordinary means)
– mass number
• the number of protons plus neutrons in the
nucleus
• electrons have negligible mass
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Fig. 2.2
–
–
–
+ + +
+ + +
+
–
–
–
Carbon
Nucleus contains
6 protons
6 neutrons
Hydrogen
Nucleus contains
1 proton
1 electron in orbit
around nucleus
Proton +
(Positive charge)
6 electrons in orbit
around nucleus
Neutron
(No charge)
Electron –
(Negative charge)
2.1 Atoms
• Electrons determine the chemical behavior
of atoms
– these subatomic components are the
parts of the atom that come close
enough to each other in nature to
interact
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Table 2.1
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2.1 Atoms
• Electrons are associated with energy
– electrons have energy of position, called
potential energy
– the field of energy around an atom is
arranged as levels called electron shells
• within this volume of space, orbitals are
where electrons are most likely to be found
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2.1 Atoms
• Electron shells have specific numbers of
orbitals that may be filled with electrons
– atoms that have incomplete electron orbitals
tend to be more reactive
– atoms will lose, gain, or share electrons in
order to fill completely their outermost
electron shell
– these actions are the basis of chemical
bonding
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2.1 Atoms
• As electrons move to
a lower energy level,
closer to the nucleus,
energy is released
Figure 2.3 The electrons of
atoms possess potential energy
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-
• Moving electrons to
energy levels farther
out from the nucleus
requires energy
+
Energy
level
3
Energy
level
2
Energy
level
1
++
+ +
+
+
Energy
level
1
Energy
level
2
Energy
level
3
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Fig. 2.4
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–
–
–
–
+
+
+
+
+ ++
+
+
–
–
–
–
–
(a) Helium
(b) Nitrogen
He 2 p, 2 n and 2e. 1 shell lowest energy level. N 7p, 7n, and 7e. 2 e fill first
nd
2.2 Ions and Isotopes
• Ions—atoms that have gained or lost one
or more electrons
• Isotopes—atoms that have the same
number of protons but different numbers of
neutrons
– most elements in nature exist as
mixtures of different isotopes
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Fig. 2.5
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Electron
is lost
Na
Na+
Sodium atom
11 protons
11 electrons
Sodium atom
11 protons
10 electrons
Figure 2.6 Isotopes of the element carbon
99% of all
carbon
Different
atomic
mass
Same
atomic
number
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2.2 Ions and Isotopes
• Some isotopes are unstable and break up
into particles with lower atomic numbers
– this process is known as radioactive
decay
• Radioactive isotopes have multiple uses
– Medicine: detection and treatment of
disorders
– Dating fossils
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2.2 Ions and Isotopes
• Short-lived isotopes
decay rapidly and do
not harm the body
Figure 2.7 Using a radioactive
tracer to identify cancer
• Can be used as
tracers to study how
the body functions
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2.2 Ions and Isotopes
• Dating fossils
– the rate of decay of a radioactive element is
constant
– by measuring the fraction of radioactive
elements that have decayed, scientists can
date fossils
– the older the fossil, the greater the fraction of
its radioactive atoms that have decayed
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Figure 2.8 Radioactive isotope dating
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Cosmic rays
14N
14C
Decay
14N
The most common isotope of carbon
is 12C. However, a tiny amount of 14C
exists in the atmosphere due to the
bombardment of 14N with cosmic rays.
the equilibrium ratio of 14C to 12C is
a constant, A.
The ratio of 14C to 12C in CO2 used in
photosynthesis is A.
The ratio of 14C to 12C in an animal that
eats plants is A.
After an organism dies, 14C decays, but
no additional 12C is taken into the body.
therefore, the ratio of 14C to 12C
decreases by ½ every 5,730 years,
the half-life of 14C
14C
14N
The ratio of 14C to 12C in the skeleton
after 11,460 years has been reduced
to A/4, or two half-lives ( ½ × ½ = ¼).
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2.3 Molecules
• A molecule is a group of atoms held
together by energy
• The energy holding two atoms together is
called a chemical bond
• There are 3 main types of chemical bonds
– ionic
– covalent
– hydrogen
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2.3 Molecules
• Ionic bonds involve the attraction of
opposite electrical charges
• Molecules comprised of these bonds are
often most stable as crystals
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Na
Cl
Sodium atom
Chlorine atom
Cl–
Na+
Cl–
Na+
Cl–
Na+
Cl–
Na+
Cl–
NaCl crystal
1 mm
–
+
Na+
Cl–
Sodium ion
Chloride ion
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Here are examples of ionic bonds and ionic
compounds:
NaBr - sodium bromide
KBr - potassium bromide
NaCl - sodium chloride
NaF - sodium fluoride
KI - potassium iodide
KCl - potassium chloride
CaCl2 - calcium chloride
K2O - potassium oxide
MgO - magnesium oxide
You can recognize ionic compounds because they
consist of a metal bonded to a nonmetal.
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2.3 Molecules
• Covalent bonds—form between two
atoms when they share electrons
– the number of electrons shared varies
depending on how many the atom
needs to fill its outermost electron shell
– covalent bonds are stronger than ionic
bonds
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Covalent bond
–
H
H
H
H
–
Hydrogen gas (H2)
Shared electrons
Single
covalent
bond
H
H
–
–
–
–
–
H
H
C
–
H
C
H
–
–
–
–
H
–
H
Methane gas (CH4)
Double
covalent
bond
–
–
–
–
–
–
–
–
O
O
O
–
–
O
–
–
Oxygen gas (O2)
–
–
–
–
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• Polar covalent bonds
– in polar covalent bonds, one nucleus
attracts the shared electrons more than
another nucleus
– this attraction for electrons by a nucleus
is called the atom’s electronegativity
– this results in partial charges in the
atoms that are unequally sharing
electrons; molecules with these partial
charges are called polar molecules
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These are examples of covalent bonds and
covalent compounds. Covalent compounds also
are known as molecular compounds.
PCl3
CH3CH2OH
O3 - ozone
H2 - hydrogen
H2O - water
HCl - hydrogen chloride
CH4 - methane
NH3 - ammonia
CO2 - carbon dioxide
Organic compounds, such as carbohydrates,
lipids, proteins and nucleic acids, are all examples
of molecular compounds.
You can recognize these compounds because
they consist of nonmetals bonded to each other.
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• Hydrogen bonds form between the positive
end of one polar molecule and the negative
end of another
– each atom with a partial charge acts like a
magnet to bond weakly to another polar
atom with an opposite charge
– the additive effects of many hydrogen
bonding interactions can add collective
strength to the bonds
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Water often is cited as an example of hydrogen
bonding. but the bond is usually between the
hydrogen of one water molecule and the oxygen
atoms of another water molecule, not between the
two hydrogen atoms (a common misconception).
List of molecules that exhibit hydrogen bonding:
chloroform - CHCl3 - between hydrogen of one molecule and carbon of
another molecule
ammonia - NH3 - between hydrogen of one molecule and nitrogen of
another
acetylacetone - C5H8O2- intramolecular hydrogen bonding between
hydrogen and oxygen
DNA - hydrogen bonding between base pairs
nylon - hydrogen bonds between the repeating units of the polymer
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Water molecules
–
–
O
O
+
H
H
H
–
+
+
H
H
O
–
H
H
O
H
–
O
H
H
(right): © Corbis RF
H
H
H
H
C
O
C
H
C
H
H
H
O
H
H
Covalent
bond
Hydrogen
bond
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2.3 Molecules
• van der Waals interactions are weak,
nondirectional attractive forces between
atoms that are close to each other (must
be close)
• van der Waals interactions become
significant in molecules with numerous
atoms, such as some types of proteins
(cumulative force is strong)
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2.4 Hydrogen Bonds Give Water
Unique Properties
• Water is essential for life, we are 2/3 water
and the earth is ¾ water.
– the chemistry of life is water chemistryall living things need water! – all living
things need water
• Water is a polar molecule so…
– water can form hydrogen bonds
– hydrogen bonding confers on water
many different special properties
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• Heat Storage
– water temperature changes slowly and holds
temperature well-because of the many H bonds. Take a
large input of energy to change the temp. Why your
body can maintain internal temp.
• Ice Formation
– few hydrogen bonds break at low temperatures
• water becomes less dense as it freezes because
hydrogen bonds stabilize and hold water molecules
farther apart (more air space so it floats)
• High Heat of Vaporization
– at high temperatures, hydrogen bonds can be broken
• water requires a lot of energy to vaporize because of
all the hydrogen bonds that must be broken (you
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sweat and cool off it releases heat)
Figure 2.9 Ice formation
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Water molecules
Stable hydrogen bonds
Unstable hydrogen bonds
Water molecules
(a) Liquid water
(b) Ice
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2.4 Hydrogen Bonds Give Water
Unique Properties
• Cohesion (because it is polar)
– Attraction of water molecules
to other water molecules
– Example: surface tension
• Adhesion
Figure 2.10 Cohesion
Water strider
– Attraction of water molecules
to other polar molecules
– Example: capillary action
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• Water is highly polar
– in solution, water molecules tend to form
the maximum number of hydrogen bonds
• hydrophilic molecules are attracted to
water and dissolve easily in it
–these molecules are also polar and
can form hydrogen bonds
• hydrophobic molecules are repelled by
water and do not dissolve
–these molecules are nonpolar and do
not form hydrogen bonds
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Table 2.2
2.5 Water Ionizes
• Covalent bonds within a water molecule
sometimes break spontaneously
H 2O
OH–
hydroxide
ion
+
H+
hydrogen
ion
• This process of spontaneous ion formation is
called ionization
– It is not common because of the strength
of covalent bonds
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pH
• A convenient way to express the hydrogen
ion concentration of a solution
pH = _ log [H+]
• The pH scale is logarithmic
– A difference of one unit represents a
ten-fold change in H+ concentration
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H+ Ion
Concentration
Examples of
Solutions
H+
100
Figure 2.12
The pH scale
pH Value
Hydrochloric acid
0
10-1
1
10-2
2
Stomach acid
Lemon juice
10-3
3
Vinegar, cola, beer
10-4
4
Tomatoes
10-5
5
Black coffee
Normal rain water
10-6
6
Urine
Saliva
10-7
7
10-8
8
Pure water
Blood
Seawater
10-9
9
Baking soda
10-10
10
10-11
11
Great Salt Lake
Milk of magnesia
Household ammonia
10-12
12
Household bleach
10-13
13
Oven cleaner
10-14
14
Sodium hydroxide
H+
+
+
H+ H H
+
OH– H
OH–
OH–
H+
OH–
H+
OH–
H+ H+
OH–
OH– H+
H+ OH–
–
OH– OH
OH–
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(top): © The McGraw-Hill Companies, Inc./Bob Coyle, photographer; (middle): © The McGraw-Hill Companies, Inc./Jacques
Cornell photographer; (bottom): © The McGraw-Hill Companies, Inc./Jill Braaten photographer
• Pure water has a pH of 7
– there are equal amounts of [H+] relative
to [OH–]
• Acid—any substance that dissociates in
water and increases the [H+]
– acidic solutions have pH values below 7
• Base—any substance that combines with
[H+] when dissolved in water
– basic solutions have pH values above 7
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• The pH in most living cells and their
environments is fairly close to 7
– proteins involved in metabolism are
sensitive to any pH changes
• Acids and bases are routinely encountered
by living organisms
– from metabolic activities (i.e., chemical
reactions)
– from dietary intake and processing
• Organisms use buffers to minimize pH
disturbances
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Fig. 2.13
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9
8
7
pH
6
Buffering
range
5
4
3
2
1
0
0
1
4
2
3
Amount of base added
5
2.5 Water Ionizes
• Buffer—a chemical substance that takes
up or releases hydrogen ions
– buffers don’t remove the acid or the
base affecting pH but minimize their
effect on it
– most buffers are pairs of substances,
one an acid and one a base
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Figure 2.13 Buffers minimize changes in pH
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9
8
7
pH
6
Buffering
range
5
4
3
2
1
0
0
1
4
2
3
Amount of base added
5
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Inquiry & Analysis
Using Radioactive Decay to Date
the Iceman
• In the graph, what is the
dependent variable?
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Radioactive Decay of Carbon Isotope 14C
• What proportion of the 14C
present in Ötzi’s body when he
died is still there today?
• How old are the remains of the
Iceman Ötzi?
Proportion of 14C remaining
1/
1
1.000
1/
2
0.500
1/
4
0.250
1/
1/
16
1/
8
0.125
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0
0.062
0.031
0.000
0
5,730
11,460 17,190 22,920
Time (years)
28,650
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Fig. 2.14
1 Holding one’s breath
1
causes levels of CO2 to
rise in the blood.
2 CO2 combines with water to
form carbonic acid (H2CO3).
3 Carbonic acid dissociates to form
–
bicarbonate ion (HCO3 and H+.
CO2
CO2 + H2O
2
H2CO3
–
3 HCO3 + H+
4 pH sensors detect a drop in pH due to an increase
in H+ concentration in the blood and send a signal
to the brain, forcing the person to breathe.
pH sensors
4
Page 47
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O
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O
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