Transcript UNIT 7: BONDING - St. Dominic High School

UNIT 7: BONDING
What previous knowledge will help us understand bonding?
How can we describe energy involved in a chemical bond?
How can we explain and draw ionic bonds?
How can we explain and draw covalent bonds?
What are metallic bonds and why are they good conductors?
What is the difference between bond polarity and molecule
polarity?
How are molecules geometrically arranged?
How does the VESPR theory influence the geometry of molecules?
What are the different forces that hold molecules together?
What are sigma and pi bonds?
AIM: Prior Knowledge

ENDOTHERMIC: chemical reaction that absorbs heat, producing
products with more PE than the reactants

EXOTHERMIC: chemical reaction that produces heat with less
PE than the reactants

POTENTIAL ENERGY: stored energy based on position or
composition

Why do atoms become ions? To become stable

How do atoms become ions? Gain or loss of electrons

How do metals form ions? MELPS

How do nonmetals form ions? Opposite of MELPS
BONDING
 Chemical bonds provide the glue that holds all
compounds together
 The electron structure of atoms helps explain
many aspects of chemical bonding
ENERGY AND
CHEMICAL BONDS
 Chemical bonds are the forces that holds atoms
together.
 Energy is required to overcome these attractive forces
and separate the atoms in a compound
 Breaking a chemical bond is an endothermic process
 ENDOTHERMIC: chemical reaction that absorbs energy
producing products with more potential energy than
the reactants
Ex) N2 + ENERGY  N + N
ENERGY AND
CHEMICAL BONDS
 Formation of a bond is an exothermic
process.
EXOTHERMIC: chemical reaction that releases
energy producing products with less potential
energy than the reactants
Ex) N + N  N2 + ENERGY
Bonding and stability
 When bonds are formed their products are more
stable
 The compounds have smaller amount of potential
energy
POTENTIAL ENERGY: stored energy based on position or
composition
 The bonded elements of a compound are more stable
than the individual atom or ions because the atoms
have filled their valence electron shell
Types of bonds
 There are three types of bonds:
1. Ionic
2. Covalent
3. Metallic
 They differ in the types of
elements involved.
 Also how the valence
electrons are handled
IONS
Atoms become ions
so that they can
become stable.
Atoms become ions
by either gaining or
losing electron
IONS
Metals form ions by losing electrons and
becoming a positive ion with a smaller
atomic radius. Positive ions are called
cations.
Nonmetals form ions by gaining electrons
to become a negative ion with a larger
atomic radius. Negative ions are also
called anions
IONIC BONDING
The bond that involves the transfer of one
or more electrons from a metal atom to a
nonmetal atom to form ions. The positive
ion and negative ion they attract each
other and create a bond.
IONIC BONDING
There is a large electronegativity
difference (E.D.) between a
nonmetal and a metal. The
nonmetal rips away the valence
electron from the metal atom.
Nonmetal becomes a negative
ion or anion and the metal
becomes a positive ion or cation
IONIC BONDING
Examples:
1 - EX) when Na and Cl atoms come together 
Na loses electrons and becomes Na+1
Cl gains electrons and becomes Cl-1
They attract to form ionic compound NaCl
(aka: salts or electrolytes)
Ionic bonds have the highest polarity (most unequal
type of bonding) and the most ionic character
CLUE FOR RECOGNIZING
IONIC BONDS
1. Metal and
nonmetal
2. Electronegativity
difference in greater
than 1.7 (approx)
Properties of Ionic Bonds
1. High melting and boiling point.
2. Good electrical conductor as a liquid or
when dissolved in water
3. Not good electrical conductor as solid
4. Hard substances
DRAWING LEWIS DOT
STRUCTURES FOR IONIC
BONDS
 Write the metal symbol with no dots in
brackets
 Place the charge at the top right of the
bracket
 Write the nonmetal symbol with 8 dots
around it (except H!)
 Draw brackets around the symbol and
place the charge of the ion at the top right
of the bracket
DRAWING LEWIS DOT
STRUCTURES FOR IONIC
BONDS
 Example: Draw the Lew dot structure of the
following elements Na and F
QUESTIONS:
 What is an ionic bond?
 What are the clues for recognizing an ionic
bond?
 What happened to electrons in an ionic bond?
 We will be using dot structures in order to draw
ionic bonds. Know the information from class
about the roles of metals and nonmetals in
forming ionic bonds; predict what the dot
structure of a metal will look like.
 Predict what the dot structure of a nonmetal will
look like
PRACTICE
Bonding
COVALENT/METALLIC
Bonding
ionic
transfer
m,nm
+,-
Covalent
molecular
metallic
share
M-SOME
all nm
all m
Polar covalent
Non polar covalent
AIM - What is it about the structure of
noble gases that leads to their stability?
 Noble gases (group 18) are stable and undergo few
chemical reactions – lack reactivity
 (Argon and Xenon combine with fluorine – rare)
 Why? – All have 8 valence electrons except He with 2
 Octet – configuration of 8 valence electrons
represents max # of valence electrons an atom can
have (except H and He – max 2)
 Octet Rule – states atoms generally react by gaining,
losing, or sharing electrons in order to achieve a
complete octet of 8 valence electrons – noble gas
configuration
Aim- How can we understand covalent
(molecular) bonds?
 Covalent bond – formed
when two nuclei share
electrons to achieve a
stable arrangement of
electrons – between all
nonmetals
 Diatomic – covalent bond
formed between two
nonmetals of the same
element:
Br2, I2, N2, Cl2 H2, O2, F2,
PROPERTIES
 Exist in gas, liquid, or solid state
 Good insulators
 Poor conductors
 Low melting points
 Many are soft substances
CONSTRUCTING LEWIS DOT STRUCTURE FOR
SINGLE BINARY COVALENT MOLECULAR
COMPOUNDS
1. Determine valence electrons in total (add them up for
each element in the compound
2. Divide by 2 to determine the number of pairs of
electrons in total for the compound
3. Place first pair between the two elements (use a dash
– to represent the shared pair)
4. Place remaining pairs around each elements making
sure not to violate the octet rule (Remember H can
have a max of 2 electrons)
H2
Cl2
HCl
HBr
 CONSTRUCTING LEWIS DOT STRUCTURES FOR
SINGLE TERNANRY COVALENT MOLECULAR
COMPOUNDS (more than two elements involved)
1.
Determine valence electrons in total (add them up for
each element in the compound
2.
Divide by 2 to determine the number of pairs of
electrons in total for the compound
3.
Determine the most electronegative element and
place it in the middle
4.
Place the other elements around it
5.
Start placing pairs (as dash lines) between the central
atom and the terminal atoms
6.
Place remaining around each elements making sure
not to violate the octet rule (Remember H can have a
max of 2 electrons)
NH3
CH4
CCl4
CONSTRUCTING LEWIS DOT STRUCTURES FOR
MULTIPLE COVALENT MOLECULAR COMPOUNDS
(more than two elements involved)
1.
Determine valence electrons in total (add them up for each element in the
compound
2.
Divide by 2 to determine the number of pairs of electrons in total for the
compound
3.
Determine the most electronegative element and place it in the middle
4.
Place the other elements around it
5.
Start placing pairs (as dash lines) between the central atom and the terminal
atoms
6.
Place remaining around each elements making sure not to violate the octet
rule (Remember H can have a max of 2 electrons)
7.
If octet rule is not yet reached you can make additional pairs of electrons
into double or triple bonds until octet rule is obeyed by all elements
*can only be done with CNOPS
CO2
O2
N2
Coordinate Covalent
Bond
Sigma () Bonds
 Sigma bonds are characterized by
 Head-to-head overlap.
 Cylindrical symmetry of electron density
about the internuclear axis.
Pi () Bonds
 Pi bonds are
characterized by
 Side-to-side
overlap.
 Electron density
above and below
the internuclear
axis.
© 2009, Prentice-Hall,
Sigma and Pi Bonds
 Single bonds are always  bonds, because 
overlap is greater, resulting in a stronger
bond and more energy
 Double bonds – contain one sigma and one
pi bond
 Triple bonds – contain one sigma and two pi
bonds
 Triple bonds are the shortest and strongest
type of bond then double and single bonds
are the longest and weakest type of bond
How can we use the 1.7 rule to predict
bond polarity?
 1.7 rule is applied primarily to binary compounds.
 Determine the electronegativities of all atoms in the
bond.
 Take the difference between the bonded atoms
 If the difference is:
>1.7 then ionic
< 1.7 but not “0”, then polar covalent
=0 non polar covalent
Nonpolar
covalent
bond
Polar covalent
bond
Covalent bonds
Ionic bond
MOLECULE POLARITY:
Nonpolar/polar shapes
SNAP PAD
 SNAP : Symmetrical Nonpolar Asymmetrical Polar
 PAD: Polar Asymmetrical Dipole
 OPEN: Odd Polar Even Nonpolar
DIPOLE
MOMENT
dipole moment--separation of the charge in a molecule; product of the size
of the charge and the distance of separation
• align themselves with an electric field (figure b at right)
• align with each other as well in the absence of an electric field
• water—2 lone pairs establish a strong negative pole
• ammonia—has a lone pair which establishes a neg. pole
• note that the direction of the “arrow” indicating the dipole
moment always points to the negative pole with the cross hatch on
the arrow (looks sort of like we’re trying to make a + sign) is at the
positive pole.
Compound Polar or Nonpolar
NH3
H20
How can a molecule be both polar
and non polar?
 CX4 tetrahedral shape. They are non polar.
 The individual ligands or bonds are polar
 Conclusion – nonpolar/polar.
What Determines the Shape of a
Molecule?
 Simply put, electron
pairs, whether they be
bonding or nonbonding,
repel each other.
 By assuming the
electron pairs are
placed as far as
possible from each
other, we can predict
the shape of the
molecule.
© 2009, Prentice-Hall,
Valence Shell Electron Pair
Repulsion Theory (VSEPR)
“The best
arrangement of a
given number of
electron domains
is the one that
minimizes the
repulsions among
them.”
© 2009, Prentice-Hall,
THE VSPER MODEL AND
MOLECULAR SHAPE

molecular geometry--the arrangement in space of
the atoms bonded to a central atom not necessarily
the same as the structural pair geometry lone pairs
have a different repulsion since they are
experiencing an attraction or “pull” from only one
nucleus as opposed to two nuclei. They also take up
more space around an atom as you can see on the
left.

Each lone pair or bond pair repels all other lone pairs
and bond pairs--they try to avoid each other making
as wide an angle as possible.

- works well for elements of the s and p-blocks

- VSEPR does not apply to transition element
compounds (exceptions)
MOLECULAR SHAPES
- Linear
- Pyramidal
- Bent (angular)- Tetrahedral
Intermolecular Forces
The attractions between molecules are
not nearly as strong as the intramolecular
attractions that hold compounds
together.
© 2009, Prentice-Hall,
Intermolecular Forces
They are, however, strong enough to
control physical properties such as
boiling and melting points, vapor
pressures, and viscosities.
© 2009, Prentice-Hall,
Intermolecular Forces
These intermolecular forces as a group
are referred to as van der Waals forces.
© 2009, Prentice-Hall,
Ion-Dipole Interactions
 Ion-dipole interactions (a fourth type of
force), are important in solutions of ions.
 The strength of these forces are what
make it possible for ionic substances to
dissolve in polar solvents.
© 2009, Prentice-Hall,
van der Waals Forces
 Dipole-dipole interactions
 Hydrogen bonding
 London dispersion forces
© 2009, Prentice-Hall, Inc.
Dipole-Dipole Interactions
 Molecules that have
permanent dipoles
are attracted to
each other.
 The positive end of
one is attracted to the
negative end of the
other and vice-versa.
 These forces are only
important when the
molecules are close
to each other.
© 2009, Prentice-Hall,
Dipole-Dipole Interactions
The more polar the molecule, the higher is its
boiling point.
© 2009, Prentice-Hall,
London Dispersion Forces
While the electrons in the 1s orbital of helium
would repel each other (and, therefore,
tend to stay far away from each other), it
does happen that they occasionally wind
up on the same side of the atom.
© 2009, Prentice-Hall,
London Dispersion Forces
At that instant, then, the helium atom is
polar, with an excess of electrons on the
left side and a shortage on the right side.
© 2009, Prentice-Hall,
London Dispersion Forces
Another helium nearby, then, would
have a dipole induced in it, as the
electrons on the left side of helium atom
2 repel the electrons in the cloud on
helium atom 1.
© 2009, Prentice-Hall,
London Dispersion Forces
London dispersion forces, or dispersion
forces, are attractions between an
instantaneous dipole and an induced
dipole.
© 2009, Prentice-Hall,
London Dispersion Forces
 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.
© 2009, Prentice-Hall,
Factors Affecting London Forces
 The shape of the molecule
affects the strength of
dispersion forces: long, skinny
molecules (like n-pentane
tend to have stronger
dispersion forces than short,
fat ones (like neopentane).
 This is due to the increased
surface area in n-pentane.
© 2009, Prentice-Hall,
Factors Affecting London
Forces
 The strength of dispersion forces tends to
increase with increased molecular weight.
 Larger atoms have larger electron clouds
which are easier to polarize.
© 2009, Prentice-Hall,
Which Will Have a Greater
Effect?
Dipole-Dipole Interactions or Dispersion Forces
 If two molecules are of comparable size and
shape, dipole-dipole interactions will likely
the dominating force.
 If one molecule is much larger than another,
dispersion forces will likely determine its
physical properties.
© 2009, Prentice-Hall, Inc.
How Do We Explain This?
 The nonpolar
series (SnH4 to
CH4) follow the
expected trend.
 The polar series
follows the trend
from H2Te through
H2S, but water is
quite an
anomaly.
© 2009, Prentice-Hall,
Hydrogen Bonding
 The dipole-dipole interactions
experienced when H is
bonded to N, O, or F are
unusually strong.
 We call these interactions
hydrogen bonds.
© 2009, Prentice-Hall,
Hydrogen Bonding
 Hydrogen bonding
arises in part from
the high
electronegativity of
nitrogen, oxygen,
and fluorine.
Also, when hydrogen is bonded to one of
those very electronegative elements, the
hydrogen nucleus is exposed.
© 2009, Prentice-Hall,
Summarizing Intermolecular
Forces
© 2009, Prentice-Hall, Inc.
Intermolecular Forces Affect
Many Physical Properties
The strength of
the attractions
between particles
can greatly affect
the properties of
a substance or
solution.
© 2009, Prentice-Hall,
Single Bonds
Single bonds are always  bonds, because
 overlap is greater, resulting in a stronger
bond and more energy
© 2009, Prentice-Hall,
How can we understand and recognize
metallic bonding?
 Contains positively charged metals
 Metals are arranged in a crystalline lattice structure
immersed in a:sea of mobile electrons
 Electrons are delocalized. This means no one atom
owns any electrons they belong to the whole
crystal.
M-SOME
Metals are good conductors because of
mobile ions
Sea of Mobile Electrons
Bonding
ionic
Covalent
molecular
metallic
m,nm
all nm
all m
+,-
Polar Covalent
Non Polar Covalent