Chem 150
Unit 1 - Molecules
Organic and Biological Chemistry substances are made
of molecules. Being able to predict the structures and
interactions of molecules is important to our
understanding of Organic and Biological Chemistry
Atomic Structure
Atoms are made up of three types of subatomic particles:
• protons
• neutrons
• electrons
Two important properties of these particles are mass and
charge.
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Atomic Structure
3
•
The protons and neutrons packed
in the center of the atom in a
region called the nucleus.
•
The electrons occupy a diffuse
region surrounding the nucleus.
•
The diameter of the volume
occupied by the electrons is
≈ 10,000 times that of the nucleus.
Atomic Structure
The electrons have a specific arrangement in the the diffuse
cloud that surround the nucleus.
• The Bohr model of the atom
•
4
Electrons circle the nucleus in specific orbits, with each orbit corresponding to a
different energy level.
Atomic Structure
The electrons have a specific arrangement in the the diffuse
cloud that surround the nucleus.
• The Bohr model of the atom only works if the atom has a
single electron.
• Quantum mechanics was able to produce a better model.
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Atomic Structure
According to quantum mechanics the electrons occupy
orbitals in pairs.
• The volumes of these orbitals have specific shapes.
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Atomic Structure
The electrons in each energy level are divided into the
orbitals.
• Each orbital can hold a maximum of 2 electrons.
• The number of orbitals in each energy level, or shell,
increases with the number of the energy level.
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Atomic Structure
For the ground states of the
elements, the electrons fill
the orbitals in a predictable
way.
• In the ground state the
electrons occupy the
available orbitals that are
nearest to the center of
the atom
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Atomic Structure
The ground state distribution of the electrons is reflected in
the organization of the periodic table
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Valence Electrons
The electrons that are most important for determining the
chemical and physical properties of an atom are those
located on the surface in the highest occupied energy level.
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•
This energy level is called the valence shell.
•
The electrons in the valence shell are called the valence
electrons.
Valence Electrons
•
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For the representative elements, the number of valence
electrons can be determined from the group number for
that element.
Electron Dot Structures
The Lewis electron dot structures are a convenient way of
showing the number of valence electrons that an atom has.
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The Octet Rule
Except for helium (He), the inert gases, (He, Ne, Ar, Kr, Xe &
Rn) each have 8 electrons in their valence shell
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Max No.
allowed
Total No. Valence
in Energy
Electrons Electrons
Level
2n2
Inert
Gas
Energy
Level
Helium
(He)
1
2
2
2
Neon
(Ne)
2
10
8
8
Argon
(Ar)
3
18
8
18
Krypton
(Kr)
4
36
8
32
Xenon
(Xe)
5
54
8
50
Radon
(Rn)
6
86
8
72
The Octet Rule
The Octet Rule:
Atoms gain, lose, or share valence electrons in
order to end up with eight valence electrons.
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The Octet Rule
Chemistry is all about the different strategies that the
elements use to attain 8 electrons in their valence shell
•
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When an element has the same number of electrons as
one of the inert gases, it is said to be isoelectron with that
inert gas.
The Octet Rule
For more discussion on the octet rule, see the Chem 150
Elaboration - The Octet Rule
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Monoatomic Ions
A strategy that the representative elements use is to simply
gain or lose electrons to become isoelectronic with one of the
inert gases.
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Monoatomic Ions
Predicting the charge on the transition metals is less straight
forward than for the representative element.
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Monoatomic Ions
Predicting the charge on the transition metals is less straight
forward than for the representative element.
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Compounds
Metals and non-metals combine
to form binary ionic
compounds.
• Each monoatomic ion in a
binary ionic compound has 8
valence electrons.
2 Na
Cl2
2 Na+ + 2 Cl-
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Covalent Bonds
Review the characteristics of Metals and Nonmetals, with
respect to their approaches to adhering to the octet rule.
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Covalent Bonds
Review the characteristics of Metals and Nonmetals, with
respect to their approaches to adhering to the octet rule.
• When metals combine with nonmetals, the metal give up
their valence electrons to the nonmetal and each forms
and ion:
• An ionic compound is formed, which is held together by an
Na
Cl
Sodium Chloride (NaCl):
ionic bond.
2+
Magnesium Chloride (MgCl2):
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Cl
Mg
Cl
Covalent Bonds
When nonmetals combine with nonmetals, a different
strategy is need.
• When nonmetals combine with nonmetals, they share
valance electrons.
• The electrons are shared in pairs to form a covalent bond.
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Compounds con’d
Compounds that are formed from nonmetals are held
together by covalent bonds.
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•
A group of atoms held together by covalent bonds is called
a molecule.
•
Compounds that are composed of molecules are called
molecular compounds.
•
Unlike ionic compounds, the group of atoms in a molecular
compound stay together as a group when the compound
changes it state (solid, liquid, gas or solution)
•
The formula that describes the composition of a molecule
is called a molecular formula.
Compounds con’d
The structures of
molecules can be
represented by either
electron dot
structural formulas
or
line-bond structural
formulas.
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Compounds con’d
Compounds that are formed from nonmetals are held
together by covalent bonds.
• In some molecules, more than one pair of electrons is
shared to form double bonds (2 pairs of electrons) or triple
bonds (3 pairs of electrons).
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Oxygen (O2):
O
Nitrogen (N2):
N
O
N
O
O
N
N
Compounds, con’d
For more discussion on the formation of compounds,
see the Chem 150
Elaboration - Compounds
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Compounds, con’d
Naming binary compounds depends on type of compound
you have.
•
Binary ionic compounds with a representative metal
•
Name of metal
+ name of nonmetal with the “-ide” ending
‣
‣
•
Binary ionic compound with a transition metal
•
Name the metal with a Roman numeral to indicate the charge
+ name of nonmetal with the “-ide” ending
‣
‣
•
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NaCl Sodium chloride
MgF2 Magnesium fluoride
CuCl Copper(I) chloride
CuCl2 Copper(II) chloride
Binary molecular compound
•
Name the least electronegative element first with a prefix to indicate its number
+ name of the more electronegative element with a prefix to indicate its number
and the “-ide” ending
‣
SO3
Sulfur trioxide
Polyatomic Ions
There are some species that are held together by covalent
bond, but which are also ionic.
2-
2-
O
O
C
O
O
Carbonate ion
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O
S
O
O
Sulfate ion
Polyatomic Ions
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Structural Formulas
Structural formulas are used to explicitly show what atom is
connected to what atom in a molecule.
• Some molecules share the same molecular formula
•
These are called isomers
These all have the same molecular formula: C3H8O
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Structural Formulas
The different types of structural formulas include:
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•
Electron dot structural formula
•
Line-bond structural formula
•
Condensed structural formula
•
Skeletal structural formula
Compounds, con’d
For more discussion on structural formulas,
see the Chem 150
Elaboration - Structural Formulas
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Noncovalent Interactions
Molecules are held together by covalent bonds.
Molecules interact with other molecules through noncovalent
interactions.
• Characteristics of noncovalent interactions
• Are much weaker than covalent bonds.
• They are the interactions that hold molecules together in the solid, liquid and
•
•
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solution states.
They are easily disrupted by increasing the temperature.
The are primarily electrical in nature
Noncovalent Interactions
Charge/Charge Interaction
• Also called a salt bridge.
• Arises from permanent charges that attract or repel one
another.
• Permanent charges are also called formal charges.
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Noncovalent Interactions
Calculating the formal charge on an atom in a molecule or
polyatomic ion.
Formal Charge
•
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=
number of valence electrons
for a neutral atom
–
number of electrons around the atom
in the compound
When calculating the number of electrons around an atom
in a compound, divide the bonding electrons equally
between the two atoms participating in a bond.
Noncovalent Interactions
Calculating the formal charge on an atom in a molecule or
polyatomic ion.
H
H
N
H
H
Formal Charge on hydrogens = 1 - 1 = 0
Formal Charge on nitrogen = 5 - 4 = +1
H
H
N
H
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H
Noncovalent Interactions
Calculating the formal charge on an atom in a molecule or
polyatomic ion.
O
O
N
O
Formal Charge on singly bonded oxygens = 6 - 7 = -1
Formal Charge on doubly bonded oxygen = 6 - 6 = 0
Formal Charge on nitrogen = 5 - 4 = +1
O
O
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N
O
Noncovalent Interactions
Dipole/Dipole and Ion Dipole Interactions
• Some molecules, which are uncharged, have their valence
electrons distributed unevenly, leading to a molecule
having a positive and a negative end.
• This situation leads to a permanent dipole.
• These dipoles can interact with ions and other dipoles.
δ-
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δ+
δ-
δ+
δ-
δ+
Noncovalent Interactions
Determining if a molecule is polar or not:
• Determining the electron distribution in a molecule is a
very complicated calculation, there are, however, some
simple rules that can be applied to determine if a molecule
is polar or not.
Two questions are asked:
1.Does the molecule contain any polar covalent
bond?
2.If so, are these bonds arranged in a way that they
will cancel each other out?
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Noncovalent Interactions
To answer the first question, identify any polar covalent
bonds that the molecule has:
• This is done by comparing the electronegativity of the two
atoms participating in each of the bonds.
• If the one atom has a substantially higher electronegativity
that the other, the valence electrons are not shared equally
and the bond is polar.
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Noncovalent Interactions
The electronegativity values for the elements (Table 4.1 in
Raymond)
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Noncovalent Interactions
The electronegativities can be used to predict whether a
bond is ionic, polar covalent, or pure covalent:
ΔE.N. > 2
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0.5 > E.N.
Noncovalent Interactions
The electronegativities can be used to
predict whether a bond is ionic, polar
covalent, or pure covalent:
C
N
 
H O
C
O
H
N
C
O
C
F
C
C
O
O
The bond is nonpolar if C or H are bond to C H
any of the other non-metals, except N, O, C S
F or Cl; or if N, O, F or Cl are bonded to
N, O, F or Cl
O
O
F
F
In general:
• The bond is polar if N, O, F or Cl are
bonded any of the other non-metals,
including H.
•
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 
Noncovalent Interactions
Some examples of polar covalent bonds:
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Noncovalent Interactions
If a molecule contains more than one polar covalent bond,
you then need to determine the geometry and shape of the
molecule around the atoms participating in the polar covalent
bonds.
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•
This is done by counting up the groups of electrons around
each atom
•
Groups of electrons include:
•
•
•
•
A non-bonding pair of electrons
A single covalent bond
A double covalent bond (counts as a single group of electrons)
A triple covalent bond (counts as a single group of electrons)
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Noncovalent Interactions
Use arrows to represent the polar bonds and look to see if
the arrows cancel one another out
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Noncovalent Interactions
Putting is all together:
Does the molecule contain any polar
covalent bonds?
No
Nonpolar
Yes
Determine the shape of the
molecule.
Are the polar covalent bonds arranged
in space in a way that causes them to
cancel one another out?
No
Polar
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Yes
Nonpolar
Noncovalent Interactions
Putting is all together:
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Noncovalent Interactions
For more discussion on determining polar molecules,
see the Chem 150
Elaboration - Polarity
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Noncovalent Interactions
The non-covalent interactions include:
• ion/ion (salt bridge)
• dipole/dipole
• ion/dipole
• hydrogen bond
• coordinate covalent
• induced dipole (London dispersion forces)
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Noncovalent Interactions
Hydrogen Bond
• A very important noncovalent interaction in biochemistry is
the hydrogen bond.
• The hydrogen bond is an extension of the the dipole/dipole
interaction.
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Noncovalent Interactions
Requirements for hydrogen bonding:
• Donor: Hydrogen atom that is covalently bonded to a
electronegative atom. In biological molecules this is usually
either an oxygen (O) or a nitrogen (N).
• Acceptor: An electronegative atom on another molecule
that contains a non-bonded pair of electrons. In biological
molecules this is usually also an oxygen (O) or a nitrogen
(N).
Examples:
  
 

N H
donor
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O
acceptor
O H
donor
O H
acceptor
Molecular Structure of Water
Water molecules are particularly good at hydrogen bonding to
themselves.
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Noncovalent interactions
Water is particularly effective at forming Hydrogen bonds
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Noncovalent Interactions
hydrogen bonds
Ion/ion
(salt bridge)
dipole/dipole
coordinate
covalent
Ion/dipole
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Noncovalent Interactions
Induced dipole interactions
•
Ion/induced dipole and dipole/induced dipole interactions
•
In this interaction the presence of an ion or permanent dipole in one molecule
will induce a dipole in another molecule by distorting the electron cloud in that
molecule
• This interaction is always attractive.
•
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Induced dipole/induced dipole interaction.
•
To understand this interaction you have to view even nonpolar molecules as
having a fluctuating dipole.
•
When brought near another molecule, the fluctuating dipoles synchronize to
produce an attractive interaction.
•
•
•
All molecules experience this interaction.
It is also called the London dispersion force
This interaction increases in strength as the molecules get larger.
Noncovalent Interactions
Induced dipole
Interact ions
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The End