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CAMPBELL
BIOLOGY
Outline
TENTH
EDITION
Reece • Urry • Cain • Wasserman • Minorsky • Jackson
I. Why study Chemistry
II. Atoms
A.
B.
C.
D.
2
The Chemical
Context of Life
Periodic Table
Isotopes
Electrons/orbitals
Bonding
III. Bonds
A.
B.
C.
D.
E.
F.
G.
Dr Burns
Napa Valley College
Covalent bonds
Polarity
Ionic bonds
Hydrogen bonding
Van der Waals Interactions
Shape/Structure
Chemical Rxns
© 2014 Pearson Education, Inc.
Why study Chemistry?


Chemistry is the basis for studying much of
biology
Biology follows the rules of physics and
chemistry.

Understanding chemistry is key to
understanding how biological systems work
Estuaries and wetlands can detoxify water:



Runoff from agricultural land may have high levels
of some elements such as selenium.
Plants can take up selenium and convert it to gas,
removing it from the water.
We can use these plants to remove selenium =
bioremediation
Definitions and the Basics

Matter – any substance that has mass and
takes up space

Elements are substances that can not be
broken down into simpler substances by
ordinary chemical reactions.


The smallest unit of an element is the atom. If
it is divided it will loose its unique properties



What will cross a membrane?
How is energy transferred through biological
systems?
What is a protein?
The Elements of Life

About 20–25% of the 92 elements are essential
to life

Carbon, hydrogen, oxygen, and nitrogen make
up 96% of living matter

Most of the remaining 4% consists of calcium,
phosphorus, potassium, and sulfur

Trace elements are those required by an
organism in minute quantities
© 2011 Pearson Education, Inc.
1
Table 2.1
Evolution of Tolerance to Toxic Elements
 Some elements can be toxic
 Some species can become adapted to
environments containing toxic elements
 For example, some plant communities are
adapted to serpentine
Serpentine Soil
Serpentine Plant Community
 low calcium-to-magnesium ratio
 lack of essential nutrients such as nitrogen,
potassium, and phosphorus
 high concentrations of nickel and chromium
An element’s properties depend on the
structure of its atoms
 Each element consists of unique atoms
 An atom is the smallest unit of matter that
still retains the properties of an element
Subatomic Particles
 Atoms are composed of subatomic particles
 Subatomic particles include
 Neutrons (no electrical charge)
 Protons (positive charge)
 Electrons (negative charge)
 In each neutral atom the # of electrons =
the # of protons
2
Figure 2.5
Atoms
Cloud of negative
charge (2 electrons)
 Neutrons and protons form the atomic nucleus
Electrons
Nucleus
 Electrons form a cloud around the nucleus
 Neutron mass and proton mass are almost
identical and are measured in daltons (or
atomic mass units, amus)
(a)
Electrons are found in the nucleus
1. True
2. False
50%
(b)
Atomic Number
 Atoms of the various elements differ in
number of subatomic particles
50%
 An element’s atomic number is the
number of protons in its nucleus
e
u
Tr
lse
Fa
Atomic Mass
 An element’s mass number is the sum of
protons plus neutrons in the nucleus, this
is the number at the bottom of the periodic
table entry.
Atomic Mass



Mass of a proton ≈1 dalton
Mass of a neutron ≈ 1 dalton
Mass of an electron ≈ 1/1800 dalton
 Atomic mass, the atom’s actual total
mass, can be approximated by the mass
number
3
Isotopes
Radioactive Isotopes

Isotopes are atoms of the same element
that have different numbers of neutrons

Every atom of an element has the same
number of protons. So all carbon atoms
have six protons
Radioactive Isotopes


In 1896, Henri Becquerel placed a rock on
unexposed photographic plates inside a
drawer. The rock contained uranium.

The isotopes of uranium emit energy.

After a few days the plate had an image of the
rock.

A co-worker, Marie Curie, named this
radioactivity.
Radioisotopes in Medicine
Radioactive isotopes are unstable, and
become more stable by emitting energy and
particles

In contrast, most isotopes are stable

Radioactive isotopes are used to date fossils
and in medicine
Figure 2.7

PET scans (Positron-Emission Tomography)



Patient is injected with a compound that is labeled
with an unstable isotope
Cancer cells are growing faster and take up more
of the compound than normal cells
Abnormal tissue takes up less of the compounds
Periodic Table

The periodic table is a chart of the elements
arranged by atomic number
Cancerous
throat
tissue
4
Figure 2.9
Fig. 2.6a
Hydrogen
1H
Atomic number
2
He
4.00
Mass number
First
shell
Helium
2He
Element symbol
Electron
distribution
diagram
Lithium
3Li
Beryllium
4Be
Boron
5B
Carbon
6C
Nitrogen
7N
Oxygen
8O
Fluorine
9F
Neon
10Ne
Silicon
14Si
Phosphorus
15P
Sulfur
16S
Chlorine
17Cl
Argon
18Ar
Second
shell
Sodium
11Na
Magnesium Aluminum
12Mg
13Al
Third
shell
Figure 2.7a
2
Atomic number
He
4.003
Element symbol
Atomic mass
Answer the following questions for beryllium

The atomic number = the # of protons in an atom

Atoms have equal numbers of protons and electrons.
How many protons does beryllium (Be)
have?
1.
2.
3.
4.
5.
One
Two
Three
Four
Five
20%
1
20%
20%
2
3
20%
4
20%
5
5
How many electrons does beryllium (Be)
have?
1.
2.
3.
4.
5.
One
Two
Three
Four
Five
20%
20%
20%
2
3
20%
How many neutrons does beryllium (Be)
have?
1.
2.
3.
4.
5.
20%
One
Two
Three
Four
Five
20%
O
1
4
33%
33%
1
1.
2.
3.
4.
1.
2.
3.
4.
5.
31
32
46
47
16
20%
1
Copyright © 2009 Pearson Education, Inc.
20%
2
20%
3
electrons
neutrons
protons
none of the above
20%
4
25%
1
3
What is the mass number of an ion with 15
electrons, 16 neutrons, and a +1 charge?
20%
o
Th
20%
e
re
Fo
ur
20%
Fi
ve
Isotopes are atoms of the same element that
differ in their number of …
33%
2
Tw
5
What is the charge of an atom containing 12
protons, 11 neutrons, and 12 electrons?
1. -1
2. 0
3. +1
ne
20%
25%
25%
2
3
25%
4
The Energy Levels of Electrons

Energy is the capacity to cause change

Potential energy is the energy that matter has
because of its location or structure

The electrons of an atom differ in their amounts
of potential energy

An electron’s state of potential energy is called
its energy level, or electron shell
20%
5
© 2011 Pearson Education, Inc.
6
Figure 2.6
(a) A ball bouncing down a flight
of stairs provides an analogy
for energy levels of electrons.
Third shell (highest energy
level in this model)
Second shell (higher
energy level)
First shell (lowest energy
level)
Energy
absorbed
Energy
lost
Atomic
nucleus
(b)
Shell Model of Electrons

This is not a correct illustration of their
location, it is used to illustrate energy states
Electron Orbitals

An orbital is the three-dimensional space
where an electron is found 90% of the time

Electrons can be visualized as residing in shells
around the nucleus (but they don’t reside here).

The first shell can have up to two electrons

The second shell and third shells can have up to
eight electrons

The outer most shell is the valence shell with the
highest energy

The chemical behavior of an atom depends
mostly on the number of valence electrons
Chemical Bonds

Chemical bonds are unions between electron
structure from different atoms

Molecules are when two or more atoms join
together. They can be the same element (H2) or
different elements (H2O)

When different elements join together, the
molecule is referred to as a compound molecule
© 2011 Pearson Education, Inc.
7
Electrons and Bonding

If the valence shell is full, then the atom is
non-reactive, inert, and does not form
chemical bonds.

Incompletely filled outer orbital, then the atom
is reactive and will form chemical bonds.

How many bonds, and what type of bond it
can form depends on how many unfilled
spots in outer shell
Covalent Bonds
Covalent Bonding
 A covalent bond is the sharing of a pair of
valence electrons by two atoms

Each atom has an attractive force for the other
atoms unshared electrons, but not enough to
take it completely away

There can be single, double or triple covalent
bonds
Figure 2.11-3
Hydrogen atoms (2 H)
 A single covalent bond, or single bond, is the
sharing of one pair of valence electrons
 A double covalent bond, or double bond, is
the sharing of two pairs of valence electrons
 A triple covalent bond, or triple bond, is the
sharing of three pairs of valence electrons
Hydrogen molecule (H2)
Figure 2.12a
Animation: Covalent Bonds
Name and
Molecular
Formula
Electron
Distribution
Diagram
Lewis Dot
Structure and
Structural
Formula
SpaceFilling
Model
(a) Hydrogen (H2)
8
Figure 2.12b
Figure 2.12c
Name and
Molecular
Formula
Electron
Distribution
Diagram
Name and
Molecular
Formula
SpaceFilling
Model
Lewis Dot
Structure and
Structural
Formula
(b) Oxygen (O2)
Electron
Distribution
Diagram
SpaceFilling
Model
Lewis Dot
Structure and
Structural
Formula
(c) Water (H2O)
Figure 2.12d
How many covalent bonds can an atom form?
Name and
Molecular
Formula
Electron
Distribution
Diagram
Lewis Dot
Structure and
Structural
Formula

SpaceFilling
Model


Each atom wants their outer shell filled.
Hydrogen only has one electron in its shell –
wants two, so it can form one bond.
Carbon has four electrons in outer shell,
wants eight, so it can form four bonds.
(d) Methane (CH4)
How many bonds can carbon form?
1.
2.
3.
4.
One
Two
Three
Four
25%
1
25%
25%
2
3
25%
4
How many bonds can hydrogen form?
1.
2.
3.
4.
One
Two
Three
Four
25%
1
25%
25%
2
3
25%
4
9
How many bonds can nitrogen form?
How many bonds can helium form?
1.
2.
3.
4.
None
One
Two
Three
25%
1
25%
25%
2
3
One
Two
Three
Four
25%
1
25%
2
25%
25%
3
2
3
25%
4

Some atoms have a greater pull on shared
electron than other atoms

The measure of this pull is electronegativity

When a molecule is made up of atoms with
different electronegativities it is a polar molecule

The greater the pull the more electronegative
(remember that electrons are negative)
4
Figure 2.13
–
Nonpolar bonds the atoms have same pull on
the shared electrons (H2)

25%
Electronegativity
Covalent bonds can be polar or nonpolar

25%
1
25%
Types of Covalent Bonds

One
Two
Three
Four
4
How many bonds can oxygen form?
1.
2.
3.
4.
1.
2.
3.
4.
25%
Polar bonds – the atoms don’t equally share
O
+
H
H
H2O
+
the electrons (H2O)
10
Tab 2.2
Polarity Con’t

Polar Covalent Bonding occurs with strong
electrophiles (electronegative): atoms with
nuclei that have a strong pull on electrons.
Common examples in biological molecules
include:



Water
Oxygen
Nitrogen
Sulfur (less than oxygen or nitrogen)
Alcohol
H
H
H
H
O
C
H
H
H
H
S
N
O
H
CH3
CH3
Aldehyde
Ketone
O
O
H2
C
H3C
H2
C
C
C
H2
CH3
H3C
C
C
H2
H
H
C
HC
Hydrocarbons
H2
C
H3C
H2
C
C
H2
HC
H
H2
C
C
H2
CH3
C
H2
H
CH
H
CH
C
H
C
H
Polar Functional Groups
Carboxyl

Oxygen containing:

Carboxyl = - COOH

Hydroxyl (alcohol) = - OH
Phosphates = -PO4
Carbonyl
 Ketone = - CO
 Aldehyde = - CHO




Nitrogen containing: Amino (-NH2)
Thiols - Sulfur containing compounds (-SH)
Alcohol
CH3CH2CH2OH
Ketone
Aldehyde
Ether
Hydrocarbons
CH3-O-CH2CH3
CH3CH2CH3
11
Which molecule is the most polar?
1.
2.
3.
4.
5.
Nonpolar compounds
CH3-SH
CH3-NH2
CH3-O-CH3
CH3-OH
CH3-COOH

Hydrocarbons – lots of carbons and
hydrogens bonded together
Copyright © 2009 Pearson Education, Inc.
Terminology
Ionic Bonds

Ion = atom that has gained or lost electrons,
It no longer has a balance between protons
and electrons, it is positive or negative charge

Ionic bond is an association between ions of
opposite charge: cations (positive) and anions
(negative)
 Hydrophilic (water-loving) – polar molecules
that are attracted to water
 Hydrophobic (water-fearing) – nonpolar
molecules that are pushed aside by water
Figure 2.12-1
Figure 2.12-2
+
-
Na
Cl
Na
Cl
Na
Cl
Na
Sodium atom
Cl
Chlorine atom
Na
Sodium atom
Cl
Chlorine atom
Na+
Sodium ion
(a cation)
ClChloride ion
(an anion)
Sodium chloride (NaCl)
12
Page 21
Animation: Ionic Bonds
Animation: Ionic Bonds
Right-click slide / select “Play”
© 2011 Pearson Education, Inc.
In the reaction between sodium and chloride,
which is being oxidized
1. Sodium
2. Chloride
50%
50%
Weak Chemical Bonds
 Most of the strongest bonds in organisms are
covalent bonds that form a cell’s molecules
 Weak chemical bonds are also indispensable
 Many large biological molecules are held in
their functional form by weak bonds
 The reversibility of weak bonds can be an
advantage
1
2
Copyright © 2009 Pearson Education, Inc.
13
Hydrogen Bonds

A hydrogen bond forms when a hydrogen
Hydrogen Bonding

Weak attraction between a slightly positive
hydrogen atom and an electronegative atom
such as oxygen, nitrogen, or less commonly
sulfur.

Individually weak, but many together can be
strong.

Determines shapes of many biological
molecules including proteins and DNA
atom covalently bonded to one
electronegative atom is also attracted to
another electronegative atom

In living cells, the electronegative partners
are usually oxygen or nitrogen atoms
© 2011 Pearson Education, Inc.
Fig. 2.11a
Figure 2.14
-
+
Water (H2O)
+
-
Hydrogen bond
Ammonia (NH3)
+
+
+
Fig. 2.11b
Which of the following statements is correct about the
atoms in ammonia (NH3)?
1. The N will have a
slight positive charge
2. The N will have a
strong positive charge
3. The H will have a
slight positive charge
4. The H will have a
slight negative charge
25%
1
25%
25%
2
3
25%
4
14
Van der Waals Interactions


If electrons are distributed asymmetrically
in molecules or atoms, they can result in
“hot spots” of positive or negative charge
Van der Waals interactions are attractions
between molecules that are close together
as a result of these charges
What type of chemical bond results from an
unequal sharing of electrons between two
atoms?
1.
2.
3.
4.
Polar covalent
Nonpolar covalent
Ionic
Van der Waals
25%
1
25%
25%
2
3
25%
4
© 2011 Pearson Education, Inc.
Molecular Shape and Function
Figure 2.17
s orbital
z
Three p orbitals
Four hybrid orbitals
x

y
A molecule’s shape is usually very
important to its function
Tetrahedron
(a) Hybridization of orbitals
Space-Filling
Model

Ball-and-Stick
Model
A molecule’s shape is determined by the
positions of its atoms’ valence orbitals
Hybrid-Orbital Model
(with ball-and-stick
model superimposed)
Unbonded
Electron
pair
Water (H2O)
Methane (CH4)
(b) Molecular-shape models
© 2011 Pearson Education, Inc.
Molecular Shape and Function
Figure 2.18
Carbon
Hydrogen
Natural endorphin
Nitrogen
Sulfur
Oxygen
Morphine

Biological molecules recognize and
interact with each other with a specificity
based on molecular shape

Molecules with similar shapes can have
similar biological effects
(a) Structures of endorphin and morphine
Natural
endorphin
Brain cell
Morphine
Endorphin
receptors
(b) Binding to endorphin receptors
© 2011 Pearson Education, Inc.
15
Chemical reactions

Chemical reactions are the making and
breaking of chemical bonds

The starting molecules of a chemical
reaction are called reactants

The final molecules of a chemical reaction
are called products
Figure 2.UN02
2 H2
+
O2
Reactants
2 H2O
Reaction
Products
© 2011 Pearson Education, Inc.
Figure 2.UN07
Which of the following molecules is drawn correctly?

Photosynthesis is an important chemical
reaction

Sunlight powers the conversion of carbon
dioxide and water to glucose and oxygen
6 CO2 + 6 H20 → C6H12O6 + 6 O2
© 2011 Pearson Education, Inc.
Important Concepts

Reading for next lecture: Chapter 3

Know the vocabulary in the lecture/chapter

What are the particles of an atom, their location,
charge, and mass?

Be able to read a periodic table and determine how
many protons, neutrons, and electrons are in a
neutral atom for each element.

Important Concepts

Be able to determine how many bonds each
element can form. Understand the shell model and
valence shells, valence electrons and how they are
related to chemical bonds

What are the four major types of bonds discussed
in lecture and be able to describe them

Identify polar and nonpolar molecules
What are isotopes and radioactive isotopes, what
are their similarities and differences?
16
Important Concepts

Be able to draw two molecules hydrogen bonding
with each other

How is a molecule’s shape important in its
function, what determines the shape?
17