* Unit
6 Periodic Trends
and Bonding
CDO CP Chemistry
*
*There exists an attraction between
oppositely charged particles
*The greater the distance between
the particles the weaker the
attraction
*The electrons in the atom are attracted to the
protons
*Electrons closest to the nucleus feel a
stronger attraction force than electrons on
the outermost energy level
*As you move in a row from left to right on the
Periodic Table the number of protons in an
atom increases and so the attractive force on
the outermost electrons increases
*
*As you move down a column
on the periodic table the
distance between the
outermost electrons and the
nucleus is the dominant factor
determining the attractive
force
•
•
Nuclear charge – equals the number of protons in nucleus
•
Effective nuclear charge – is less than the full charge
Outer electrons – don’t experience the full attraction of the
charge as they are shielded from the nucleus and repelled
by the inner electrons
*Effective Nuclear
Charge
5
*Determining Effective
•
Equation
*
•
Nuclear Charge
Zeff = # protons - # core electrons
Trends
*
*
Increases from left to right across a periods
Stays the same in a group
6
•
•
•
Periods – Rows on the Periodic Table
•
Example: How many outer shell electrons
does P have?
Groups – Columns on the Periodic Table
Group # - equals the number of electrons in
the outer shell (valence electrons) for the
main group elements (Groups 1-0)
*Structure of the
Periodic Table
7
•
Physical properties depend ultimately on the
balance between electrostatic attractions
between oppositely charged particles (aka
columbic attraction)
*Identifying Physical
Properties
8
*
9
•
Nuclear charge – equals the number of
protons in nucleus
•
Outer electrons – don’t experience the full
attraction of the charge as they are shielded
from the nucleus and repelled by the inner
electrons
•
Effective nuclear charge – is less than the full
charge
*Effective Nuclear
Charge
10
*Determining Effective Nuclear
•
Equation
*
•
Charge
Zeff = # protons - # core electrons
Trends
*
*
Increases from left to right across a periods
Stays the same in a group
11
*Periodic Trends – The Basic
•
Idea
Force of attraction of electrons to the
nucleus depends on 2 factors – nuclear
charge and distance2
•
Across a period – number of protons
increases, increasing attraction of the
electrons. The increasing number of
valence electrons does not contribute as
much as the proton increase
•
Down a group – number of electron shells
increases, which increases distance from
the nucleus to the outer most electrons
12
*Atomic
* Atomic radius is determined by two factors.
1. The number of energy levels
2. The effective nuclear charge.
Radius
*Atomic
Radius in a Group
•
Atomic radius generally increases as you
move down a group of the periodic table.
•
Electrons are added to higher principal
energy levels and the effective nuclear
charge remains constant
*Atomic
Radius in a Period
•
As you proceed across a period, the size of successive
atoms decreases.
•
The number of energy levels remains constant but . . .
•
The effective nuclear charge increases
*Atomic
Radius for Main Group
Elements
* Example
*Choose the larger atom in each pair:
* Na or Si
* P or Sb
* Al or Cl
* Al or In
1
•
*Ionic
Ionic Radius trend is essentially the same at the
atomic radius, but now we are talking about the ions
(cation or anion)
Radius
*
If positively charged the radius decreases, because
they lose shells
*
If negatively charged the radius increases, because
they gain repulsive forces (proton, electron ratio
changes)
*
When substances have the same number of electrons
(isoelectronic), the greater the atomic number the
smaller the radius
*Example
*Choose the larger particle in each
pair
*Na or Na+
*F- or F
*Al3+ or Al
2
*Example
Predict which of the following substances has the largest
radius: P3, S2, Cl, Ar, K+, Ca2+.
3
*IONIZATION
ENERGY Ei: minimum energy Review
•
Ionization energy,
required to remove a mole of electrons
from the ground state of an atom in the
gas phase. M(g) + h  M+ + e.
•
Factor that affect IE – force of
attraction between the nucleus and the
valence electrons, stronger the force of
attraction, more energy required to pull
the electron away
*Ionization
Energy: Periodic
Ionization Energy vs atomic #
Table
*Example
* Choose the atom with the larger ionization energy in each
pair
* B or C
* O or S
* Cl or I
4
*Electronegativity
•
•
•
•
Electronegativity is the tendency of the
atoms in a molecule to attract bonding
electrons.
Electronegativity is used to predict the type
of bonding that exists between atoms in a
compound.
Electronegativity is expressed in a qualitative
measurement called the Pauling
electronegativity scale.
The scale assigns the highest value to fluorine
and the lowest to the alkali metals.
*Linus
Pauling Electronegativity
Table – in the Data Booklet
*Electronegativity
* Going across a period, electronegativity increases.
Trends
* Metals have low electronegativities, non-metals high
electronegativities.
* Electronegativity decreases going down a group.
* The ability of an atom to attract electrons decreases with
increasing nuclear charge.
*Example
* Predict the order of increasing electronegativity in
each of the following groups of elements.
*C, N, O
*S, Se, O
6
•
Melting point – the temperature that give the
particles enough kinetic energy to so they
can overcome the attractive forces holding
them together, it is only a change of state
*Melting Points
28
*Melting points of the Group I
Elements
•
•
•
Melting point decreases as you go down the
group
Reason for the trend
Metallic Bonding – sea of delocalized
electrons, as the ion (+1) gets larger the
charge gets diluted over the increase surface
of the ion. As the charge density decreases
the force of attraction between the ion and
the electron is lowered therefore the melting
point goes down
29
*• Melting
Point
of
the
Group
VII
Melting Point increases as you go down the
Halogens
group
•
Reasoning
*
Halogens are diatomic – so the are
molecules which are held together by
attractive forces (more later)
*
The attractive forces increase with
increasing molar mass (molecular size)
30
*Melting points Across Period 3
•
Determine the trend by breaking into 4 main sections
*
Metals – Na, Mg, Al – increasing MP due to greater ionic charge
and increased number of delocalized e-
*
Giant Covalent – Si – Highest MP – very strong covalent bonds
between all atoms
*
Molecular Covalent – P4, S8, Cl2 – S8 highest due to increased
attractive forces, Cl2 lowest
*
Atoms – Ar – lowest weakest forces
31
*
*IONIC BONDS
*COVALENT BONDS
*METALLIC BONDS
*
•
•
•
•
Octet Rule – Elements form bonds so they have
access to 8 electrons
Elements in the main groups - form bonds such
that there are eight electrons surrounding each
atom
Obtaining an octet - is the reason for bond
formation for many compounds formed by the
main group elements
Duet Rule - H, Li and Be are octet exceptions and
form bonds to have access to only 2 electrons
33
•
Valence electrons - the electrons in the
outer shell
•
Chemical bond - a force that holds two or
more atoms together, involves only the
valence electrons
•
Bonds form – between valence electrons for
a more stable electron configuration (to
have an octet/duet)
34
*
When an atom of a nonmetal
takes one or more electrons
from an atom of a metal
so both atoms end up with
eight valence electrons
*
Neutral atoms come near each other.
Electron(s) are transferred from the Metal
atom to the Non-metal atom. They stick
together because of electrostatic forces,
like magnets.
*
*Crystalline structure.
*A regular repeating
arrangement of ions in the
solid.
*Ions are strongly bonded.
*Structure is rigid.
*High melting points- because
of strong forces between ions.
*
The
POSITIVE
CATIONS
stick to the
NEGATIVE
ANIONS, like
a magnet.
+
+
- - +
+ + - + - +
- + - +
*
*Conducting electricity is allowing
charges to move.
*In a solid, the ions are locked in place.
*Ionic solids are insulators.
*When melted, the ions can move
around.
*Melted ionic compounds conduct.
*First get them to 800ºC.
*Dissolved in water they conduct.
*
+
+
-
+
+
+
+
-
+
+
*
* Strong Repulsion breaks crystal apart.
- + - +
+ - + - + - +
*
When an atom of one
nonmetal
shares one or more
electrons
with an atom of another
nonmetal so both atoms
end up with
eight valence electrons
When one nonmetal shares one or
more electrons with an atom of
another nonmetal so both atoms end
up with eight valence electrons
*
•
Two nonmetals - share one, two or three
electron pairs
•
Bonds resulting from this sharing are
called - single, double or triple bonds
respectively
•
Multiple bonds - are frequently observed
in compounds of 2nd period elements
46
*
•
As the number of bonds between two atoms
increases, the bond length decreases.
© 2009, Prentice-Hall,
*
•
Generally – increase in the number
of bonds between the same two
atoms, the increase in the bond
strength
48
*
49
* The Lewis Structures of Covalent compound represents
elements in a compound, their valence electrons and how
they are shared, and how the elements orient themselves
around each other.
*
*
1.
Count ALL Valence electrons on all atoms in the molecule.
z
z
For an anion ion, add one electron for each negative charge.
For a cation, subtract one electron for each positive charge.
*
2. The atom with the least amount is central
atom and place the other atoms around the
central atom.
Hydrogen is never central atom
Draw a line connecting the peripheral
atoms to the central atom
Each line represent 2 electrons
Check for octet
*
3.
Place pairs of valence electrons around each peripheral
atom, except hydrogen, until octet is reached.
4.
If any electrons remain place around the central atom
until octet is reached.
5.
If central atom still does not have an octet, use a lone
pair of electrons on a neighboring atom to form a
multiple bond to the central atom.
*
* CF4
* CO2
* HCN
* ClO* NH4+
*
•
•
•
In compounds with multiple bondssometimes you can draw structures
which vary only by placement of the
double bonds
Resonance structures- are an
approximation of the true structure of
the molecule
Actually - the molecule is a superposition
of all of the resonance structures
55
*
*Ozone, O3
*CO32-
56
*
*Some structures will have less than an octet – this is
due to too few valence electrons for one atom to
form a sufficient amount of bonds
*Examples
*BeCl2
*BF3
57
*
* When the atoms in a bond are the same, the electrons are
shared equally.
* This is a nonpolar covalent bond.
* When two different atoms are connected, the atoms may not
be shared equally.
* This is a polar covalent bond.
* How do we measure how strong the atoms pull on electrons?
*
* A measure of how strongly the atoms attract
electrons in a bond.
* The bigger the electronegativity difference the
more polar the bond.
* 0.0 - 0.4 Covalent nonpolar
* 0.4 - 1.7 Covalent polar
* >1.7 Ionic
*
•
•
•
•
Ionic – Large electronegativity difference
Polar Covalent – medium difference in
electronegativity
Non-Polar Covalent – small to no difference in
electronegativity
Simply put:
* Ionic – metal and non metals
* Polar Covalent – 2 different non metals
* Non-Polar Covalent – 2 identical non metals and
C-H
60
*
*Isn’t a whole charge just a partial charge
*d+ means a partially positive
*d means a partially negative
*The Cl pulls harder on the electrons
*The electrons spend more time near the Cl
d+
H
d
Cl
*
*Determine and show if each of the following
bonds are polar
*
63
*
Valence Shell Electron Pair Repulsion Theory
Electron pairs orient themselves in order to minimize
repulsive forces
Types of e- Pairs
* Bonding pairs – form bonds
* Lone pairs – nonbonding e* Total e- pairs– bonding + lone pairs
*Lone pairs repel
more strongly than bonding
pairs!!!
*
Lone pairs reduce the bond angle between atoms
Bond Angle
*
Draw the Lewis Diagram
Tally up e- pairs on central atom (bonds + lone pairs)
* double/triple bonds = ONE pair
Shape is determined by the # of bonding pairs and lone
pairs
Know the 6 common shapes
& their bond angles!
*
*
*
X – number of bonded pairs
E – number of lone pairs
68
Attached Lone Pair Molecular
Atoms
Electrons Geometry
(X)
(E)
(AXE)
Geometry
Bond
Angle
2
0
AX2
Linear
180o
3
0
AX3
Planar
Triangular
120o
2
1
AX2E
Bent
117o
4
0
AX4
Tetrahedral
109.5o
3
1
AX3E
Pyramidal
107o
2
2
AX2E2
Bent-
105o
69
6 total
6 bond
0 lone
→ Electronic Geometry =
octahedral
SH6*
OCTAHEDRAL
90°
*
Molecules with ends
*
*Molecules with a positive and a negative
end
*Requires two things to be true
The molecule must contain polar bonds
The molecule must be symmetrical
The molecule must be completely the
same around the central atom to be
symmetrical
Must determine geometry first.
*
*HF
*H2O
*NH3
*CCl4
*CO2
*
What holds molecules to each other
*
•
•
The attractions between molecules that
hold them together in the liquid and
solid phases
Much weaker than the covalent bonds
within molecules.
*
•
They are, however, strong enough to
control physical properties such as boiling
and melting points
*
•
These intermolecular forces, which are all
electrostatic in nature, are referred to as
van der Waals forces.
• Dipole-dipole interactions
• London dispersion forces
• Hydrogen bonding
*
*Molecules that have
permanent dipoles (are
polar) are attracted to
each other.
*The positive end of one is
attracted to the negative
end of the other and viceversa.
*These forces are only
important when the
molecules are close to
each other.
*
•
Nonpolar molecules do not have
permanent dipoles. But at any given
instant, the electrons may be unevenly
distributed with in an atom or molecule.
*
•
London dispersion forces, or
dispersion forces, are attractions
between an instantaneous dipole
and an induced dipole.
*
•
•
These forces are present in all
molecules, whether they are polar or
nonpolar.
The tendency of an electron cloud to
distort in this way is called polarizability.
*
•
•
The shape of the molecule
affects the strength of
dispersion forces: long,
skinny molecules (like npentane tend to have
stronger dispersion forces
than short, fat ones (like
neopentane).
This is due to the increased
surface area in n-pentane.
•
The strength of dispersion forces tends to
increase with increased molecular
weight.
•
Larger atoms have larger electron clouds
which are easier to polarize.
*
•
•
•
The dipole-dipole interactions
experienced when H is
bonded to N, O, or F are
unusually strong.
We call these interactions
hydrogen bonds.
Although it is called a bond it
is actually an intermolecular
force
*
•
Hydrogen
bonding arises
in part from
the high
electronegativi
ty of nitrogen,
oxygen, and
fluorine.
*
1. Draw the Lewis Structure
2. Determine if the molecule is polar or
nonpolar
*
•
Nonpolar molecule = dispersion forces
•
Polar molecule = dipole forces,
dispersion forces
•
Polar molecules with hydrogen directly
bonded to fluorine, oxygen or nitrogen
= hydrogen bonding, dipole forces,
dispersion forces
*
•
Usually IMF’s are described in terms of
the strongest one present
* Strongest – Hydrogen Bond
* Medium – Dipole/Dipole
* Weak – London dispersion forces
88
*
Determine the intermolecular forces for each
of the following molecules:
water
carbon tetrachloride
ammonia
carbon dioxide
phosphorus trichloride
*
90
•
Ionic – attractive forces between
oppositely charged ions, Salts are ionic
but not molecules
•
Covalent – Giant – just regular covalent
bonds but they exist between all atoms
*
91
•
Covalent – Molecular – regular covalent
bonds, forming little packages called
molecules, they are attracted to each
other by IMF
•
Metallic – electrostatic attraction
between lattice of positive ions and
delocalized electrons
*
92
Bonding
Ionic
Metallic
Giant
Covalent
Molecular
Covalent
Lattice Forces
Ionic
Ion/electron
Covalent
BP
High
High
High
Volatility
Low
Low
Low
DH vap
High
High
High
Hbond
Dipole/Dipole
London Disp.
Low
high
Low
*
Volatility – measure of the tendency to
evaporate
93
DH vap – PE change required to vaporized a
*Allotrope – different forms of an element in
the same physical state
*Carbon has 3 Allotropes
*
94
*Silicon – Giant covalent similar to diamonds
structure, has high mp and bp (but lower than
C diamond)
*Silicon Dioxide – Giant covalent structure,
each Si is bonded to 4 O in a tetrahedral
shape. Formula is SiO2 because each of the 4
O is only ½ shared by each Si. Has high mp
and bp.
*
95
*
*Solubility – how easy the solute (usually a solid)
becomes dissolved in a solvent (usually water)
*Axiom – “Like dissolves Like” – polar solutes will
dissolve in polar solvents and non polar solutes will
dissolve in non polar solvents
*Reasoning – The solute particles must be able to
interact with the solvent particles and the force of
attraction for the solute to the solvent must be
greater than for the solvent itself
96
*NaCl and Water
*Methane and NaCl
*
97
*They are what make solid and liquid
molecular compounds possible.
*The weakest are called van der Waal’s
forces - there are two kinds
*Dispersion forces
* Dipole Interactions
* depend on the number of electrons
* more electrons stronger forces
* Bigger molecules
*
*
* Depend on the number of electrons
* More electrons stronger forces
* Bigger molecules more electrons
*Fluorine is a gas
*Bromine is a liquid
*Iodine is a solid
* Occur when polar molecules are attracted to
each other.
* Slightly stronger than dispersion forces.
* Opposites attract but not completely hooked
like in ionic solids.
*
* Occur when polar molecules are attracted to
each other.
* Slightly stronger than dispersion forces.
* Opposites attract but not completely hooked
like in ionic solids.
+
d
*

d
H F
+
d

d
H F
*
d+
d+
d
d
* Are the attractive force caused by hydrogen
bonded to F, O, or N.
* F, O, and N are very electronegative so it is a
very strong dipole.
* The hydrogen partially share with the lone pair
in the molecule next to it.
* The strongest of the intermolecular forces.
*
*
d+ dH O
+
H d
*
H O
H
*
*How atoms are held together in
the solid.
*Metals hold onto there valence
electrons very weakly.
*Think of them as positive ions
floating in a sea of electrons.
*
* Electrons are free to move through the solid.
* Metals conduct electricity.
+
+ + +
+ + + +
+ + + +
*
*Hammered into shape
(bend).
*Ductile - drawn into
wires.
*
+
+ + +
+ + + +
+ + + +
*
*Electrons allow atoms to slide
by.
+ + + +
+ + + +
+ + + +