Chemical Bonding
Comparison of Properties
Ionic Compounds
Covalent Compounds
Metals
Properties of Matter
• Macroscopic properties of matter vary greatly
due to the type of bonding
Types of Bonding
• Metallic
• Covalent
• ionic
Metallic
Ionic
Covalent
What is a chemical bond?
• An attractive force that holds two atoms
together
• Can form by
– The attraction of positive ion to a negative ion or
– The attraction of the positive nucleus of one atom
and the negative electrons of another atom
Bond
• the interaction between two or more atoms
that allows them to form a substance different
from the independent atoms.
• involves the outer (valence) electrons of the
atoms.
• These electrons are transferred from one
atom to another or shared between them.
• Can be between
atoms of different
elements to make a
compound, like the
two hydrogen atoms
and one oxygen
atom in a water
molecule.
Sulfur
• can also be between
atoms of a single
element.
– Sulfur is an example
of an element that
has its most stable
form as a small
molecule, in this case
a ring of eight sulfur
atoms.
Chemical Bond Energy Considerations
• A chemical bond forms when it is energetically
favorable
– when the energy of the bonded atoms is less than
the energies of the separated atoms.
Bonding
• Chemical compounds are formed by the
joining of two or more atoms.
• A stable compound occurs when the total
energy of the combination has lower energy
than the separated atoms.
• The bound state implies a net attractive force
between the atoms ... a chemical bond.
Energy Changes in Bonding
• When bonds are formed, energy is released.
• Demonstrations:
– Formation of an Ionic Compound: Mg + O2
– Formation of a Molecular Compound: S + O2
Breaking Bonds
• In order to break bonds energy must be
added, usually in the form of heat, light, or
electricity.
• Demonstration: Electrolysis of water
http://www.youtube.com/watch?v=HQ9Fhd7
P_HA
• Demo: Decomposition of Nitrogen Triiodide
http://www.youtube.com/watch?v=z5vsQ8sP
gX4
• In chemical bonds, atoms can either transfer
or share their valence electrons.
Ionic Bonds
• When one or more atoms lose electrons and
other atoms gain them in order to produce a
noble gas electron configuration, the bond is
called an ionic bond.
• Bonding Animation:
• http://www.youtube.com/watch?v=QqjcCvzW
www
• Intro to bonding
• http://www.youtube.com/watch?v=31bec8Hl
7wo&NR=1&feature=endscreen
Ionic Solids
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Ionic solids are solids composed of ionic
particles (ions).
These ions are held together in a regular
array by ionic bonding.
Ionic bonding results from attractive
interactions from oppositely charged ions.
In a typical ionic solid, positively charged
ions are surrounded by negatively charged
ions and vice-versa.
The close distance between these
oppositely charged particles results in very
strong attractive forces.
The alternating pattern of positive and
negative ions continues in three
dimensions.
The regular repeating pattern is analogous
to the tiles on a floor or bricks on a wall.
called the crystal lattice.
Ionic Compounds
• Crystalline solids
(made of ions)
• High melting and
boiling points
• Conduct electricity
when melted or
dissolved in water
– Demo: Electrolytes
• Many are soluble in
water but not in nonpolar liquid
Comparison of Conductivity
Common Ionic Compounds
– NaCl - sodium chloride - table
salt
– KCl - potassium chloride present in "light" salt (mixed
with NaCl)
– CaCl2 - calcium chloride driveway salt
– NaOH - sodium hydroxide found in some surface
cleaners as well as oven and
drain cleaners
– CaCO3 - calcium carbonate found in calcium supplements
– NH4NO3 - ammonium nitrate found in some fertilizers
• Ionic compounds in the solid state are held
together by electrostatic attractions between
opposite charges.
• Sodium chloride (table salt), silver sulfide
(silver tarnish), and hydrated iron (III) oxide
(rust) are examples of ionic compounds.
Covalent (Molecular) Compounds
• Gases, liquids, or
solids (made of
molecules)
• Low melting and
boiling points
• Poor electrical
conductors in all
phases
• Many soluble in nonpolar liquids but not in
water
Molecular (Covalent) Substances
Covalent Network Solids
• Covalent because
combinations of
nonmetals
• Interconnected
• very hard and brittle
• Insoluble
• Extreme melting and
boiling points
Diamond
Covalent Bonds
• involve the sharing of a pair of valence
electrons by two atoms
• Such bonds lead to stable molecules if
they share electrons in such a way as to
create a noble gas configuration for each
atom
Covalent bonding can be visualized with the aid
of a Lewis Structure
Polar Covalent Bonds
• Covalent Bonds in which the sharing of the
electron pair is unequal
• the electrons spend more time around the
more nonmetallic atom
• In such a bond there is a charge separation
with one atom being slightly more positive
and the other more negative……. will produce
a dipole moment.
Types of Covalent bonds
• Pure Covalent (also called
non-polar covalent) bonds are
ones in which both atoms
share the electrons evenly
• By evenly, we mean that the
electrons have an equal
probability of being at a
certain radius from the nuclei
of either atom.
• Polar covalent bonds are ones
in which the electrons have a
higher probability of being in
the proximity of one of the
atoms
• Determined by
Electronegativity Difference
Electronegativity
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•
the periodic property that indicates
the strength of the attraction an
atom has for the electrons it shares in
a bond.
Atoms with high electronegativities
tend to hold tightly to their electrons
or to form negative ions.
– These elements are found to the
upper right on the periodic table.
•
Atoms with low electronegativities
tend to have a lower attraction for
their electrons and may form positive
ions.
– These elements are found to the
lower left on the periodic table.
Pure covalent or Non-polar
covalent bond
• Electronegativity difference of 0.3 or less in
between the two atoms.
• A pure covalent bond can form between two
atoms of the same element (such as in
diatomic oxygen molecule)
• or atoms of different elements that have
similar electronegativies (such as in the
carbon and hydrogen atom in methane).
Polar Covalent Bond
• A is a pair of electrons shared between two atoms
with significantly different electronegativities (from
0.3 to 1.7 difference).
• These bonds tend to form between highly
electronegative non-metals and other non-metals,
such as the bond between hydrogen and oxygen in
water.
Ionic Bonds
• In compounds that have elements with very
different electronegativities (greater than 1.7
difference), the electrons can be considered to
have been transferred to form ions.
• Many of the properties of a compound, such
as solubility and boiling point, depend, in part,
on the degree of the polarity of its bonds.
Examples to Determine Bond Character
• Using electronegativity in the prediction of the
polarity of a chemical bond.
• sodium bonded to chlorine
– Difference between the electronegativities of Na(0.9) and
Cl(3.0) are so great that they form an ionic bond.
• The hydrogen molecule (2 H atoms bonded to each
other)
• zero electronegativity difference, form a non-polar
covalent bond.
Bond Character
•
Nonpolar-Covalent bonds (H2)
– Electrons are equally shared
– Electronegativity difference of 0 to 0.3
•
Polar-Covalent bonds (HCl)
– Electrons are unequally shared
– Electronegativity difference between .3 and 1.7
•
Ionic Bonds (NaCl)
– Electrons are transferred
– Electronegativity difference of more than 1.7
Diatomic Molecules
• hydrogen gas H2
• the halogens:
– chlorine Cl2
– fluorine F2
– bromine Br2
– iodine I2
• Nitrogen N2
• Oxygen O2
Pneumonic Device to remember the diatomic
molecules: Professor BrINClHOF
Metals and Metallic Bonding
• Typical Properties of Metals
– Malleable
– Ductile
– Good Conductors of Heat and Electricity
– Generally high melting and boiling points
Metallic Bonds
• The properties of metals suggest that their
atoms possess strong bonds
• yet the ease of conduction of heat and
electricity suggest that electrons can move
freely in all directions in a metal
• The general observations give rise to a picture
of "positive ions in a sea of electrons" to
describe metallic bonding.
Metal Properties
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Malleable and Ductile
Strong and Durable
Good conductors of heat and electricity.
Their strength indicates that the atoms are difficult to
separate… strong bonds
• but malleability and ductility suggest that the atoms are
relatively easy to move in various directions.
• The electrical conductivity suggests that it is easy to move
electrons in any direction in these materials.
• The thermal conductivity also involves the motion of
electrons. All of these properties suggest the nature of the
metallic bonds between atoms.
Metallic Bonding
the Electron Sea Model
• Explained by the Electron Sea Model
• the atoms in a metallic solid contribute their
valence electrons to form a “sea” of electrons
that surrounds metallic cations.
• delocalized electrons are not held by any
specific atom and can move easily throughout
the solid.
• A metallic bond is the attraction between
these electrons and the metallic cation.
• A mixture of elements that has metallic
properties is called an alloy.
• Two types of alloys
– An interstitial alloy is one in which the small holes
in a metallic crystal are filled by other smaller
atoms.
– A substitutional alloy is one in which atoms of the
original metal are replaced by other atoms of
similar size.
Ionic Compounds
Covalent Compounds
Metallic Compounds
-Formed from a combination of metals and
nonmetals.
-Electron transfer from the cation to the anion.
-Opposite charged ions attract each other.
-Formed from a combination of
nonmetals.
-Electron sharing between
atoms.
-Formed from a combination of
metals
-“sea of electrons”;
electrons can move among
atoms
Solids at room temperature
Can be solid, liquid, or gas at
room temperature.
Solids at room temperature
High melting points
Low melting points
Various melting points
Dissolve well in water
Do not dissolve in water (Sugar
is an exception)
Do not dissolve in water.
Conduct electricity only when dissolved in
water; electrolytes
Do not conduct electricity; non
electrolytes
Conduct electricity in solid form.
Brittle, hard
Soft
Metallic compounds range in
hardness. Group 1 and 2 metals
are soft; transition metals are
hard. Metals are malleable,
ductile, and have luster.