Bonding
• Metallic Bonding
– In a liquid or solid state, metals readily give up
electrons
– When only other metal atoms are around,
electrons are not accepted and held, they are
free to move
– Free moving electrons called an electron sea—
reason for metallic properties of luster,
malleability, ductility, and conductivity
Determine the electronegativity difference,
bond type, and more electronegative element
with respect to the following atoms.
1.
2.
3.
4.
5.
6.
7.
8.
9.
10.
H and F
Br and Br
Al and S
Na and S
At and Cl
Ba and O
Si and O
Cr and Cl
Fe and I
K and Br
Review: bond types, bond energy
• There are 2 types of bonding: ionic, covalent (metallic not
considered)
•
ionic = stealing of electrons to form + and – ions. + and –
then attract
• Covalent = sharing of electrons
• Remember there is no clear dividing line.
What causes atoms to form molecules?
• Basically, all things that happen spontaneously are
energetically favorable
• Something must be energetically favorable about atoms
coming together as molecules (ie: by bonding atoms can
lower their energy status
An ionic bond is the chemical bond that results from the
electrostatic attraction between positive ions (cations) and
negative ions (anions) and a lattice structure results in which
the cations are surrounded by anions and the anions are
surrounded by cations.
An ionic bond is the chemical bond that results
from the electrostatic attraction between positive
ions (cations) and negative ions (anions) and a
lattice structure results in which the cations are
surrounded by anions and the anions are
surrounded by cations.
Ionic bonding
Ionic bonding involves 3 steps (3 energies)
1) loss of an electron(s) by one element,
2) gain of electron(s) by a second element,
3) attraction between positive and negative
Na
Cl
Cl–
Ionization energy
+ e–
+ Na+
e– + Na+
Electron affinity
Lattice energy
Cl–
Cl– Na+
Ionic bonding: energies
• By convention, a requirement for energy is
given a + sign (we have to put energy in) and
is called endothermic, a release of energy is
given a – sign and is called exothermic.
Ionic bonding
Ionic bonding involves 3 steps (3 energies)
1) loss of an electron(s) by one element,
2) gain of electron(s) by a second element,
3) attraction between positive and negative
Na
Cl
Cl–
Ionization energy
+ e–
+ Na+
e– + Na+
Electron affinity
Lattice energy
Cl–
+ 496
– 349
Cl– Na+ – 766
Ionic bonding: energies
• By convention, a requirement for energy is
given a + sign (we have to put energy in) and
is called endothermic, a release of energy is
given a – sign and is called exothermic.
• Problem: the sum is +147. A spontaneous
change must involve a net lowering of energy
• Solution: the lattice energy provides the
energy needed
• Note that although we represent this as a three
step process it actual occurs all at once
Review Questions
1.
A metal + non-metal gives what kind of bond?
2.
What is necessary for any stable compound to form from its elements?
3.
Define cation. Define anion.
4.
List the 3 energies involved in forming an ionic bond.
5.
What term describes a release of energy?
6.
Is breaking a bond endothermic or exothermic?
7.
Define lattice energy.
8.
Explain why metals form cations and non-metals form anions.
9.
Explain why calcium exists as Ca2+ but not as Ca3+ in ionic compounds.
10. Explain why most transition metals form a
2+ ion.
Answers
1. Ionic
2. There must be a net lowering of energy
3. Cation: a positively charged ion, Anion: a
negatively charged ion
4. Ionization energy, electron affinity, lattice
energy
5. Exothermic
6. Endothermic
7. The energy released by the imaginary
process in which isolated ions come
together to form a crystal of an ionic
compound
8. Metals form cations because they have small
IEs (and EA), non-metals form anions
because they have large EAs (and IE).
These trends are energetically favorable.
9. The first two electrons from the 4s subshell
are easily lost (they can be made up for by
the lattice energy). Losing a third electron is
not energetically favorable because of the
large third IE for Ca (the energy required to
remove the third electron can not be made
up by the lattice energy.
10. The loss of two electrons from the s subshell
accounts for the typical 2+ charge of the
transition elements.
Covalent bonding
• Just as with ionic bonds, covalent bonds
must involve a net lowering of energy
• We can explain this net lowering of energy
in two ways:
1) visualizing the combination of
attraction as two atoms approach each
other
2) drawing and combining orbital
diagrams
As atoms approach
• Recall that EA for all atoms, except the noble
gasses are negative
–
+
–
–
• In other words we have no trouble adding
electrons to atoms
–
+
• The attraction for electrons is not limited to free electrons,
but also involves electrons that are part of other atoms.
• Thus, atoms are pulled toward each other
• How far they are pulled together will depend on a balance
of attraction (nucleus to electrons) and repulsion (nucleus to
nucleus and electrons to electrons)
• There are two types of covalent bonds.
In a non-polar covalent bond the electrons are shared equally between
the two atoms. (The nuclei and core electrons are indicated by the blue
spheres and the bonding electrons are indicated by the lavender dots.)
– The electrons move around the nuclei with the electrons generating
temporary positive and negative charges within the molecule.
– In a polar covalent bond the electrons are shared unequally
between the two atoms. In this situation, one atom has a greater
ability to pull the bonding electrons towards it and is said to be
more electronegative. (The green sphere represents the more
electronegative element.)
– Again, the electrons move around the nuclei with the electrons
spending the majority of the time near the more electronegative
element. This generates a partial negative charge near the more
electronegative element and a partial positive charge near the less
electronegative element.
• Molecular polarity is dependent on bond
polarity and the molecular geometry. For
small molecules: If all the regions
surrounding an atom are similar in their
electronegativities, the molecule will be
non-polar. If the regions are different, then
the molecule will be polar.
• For large molecules there may be polar and
non-polar regions.
Lewis Structures
1. The first step in drawing Lewis structures is to
determine the number of electrons to be used to
connect the atoms. This is done by simply adding
up the number of valence electrons of the atoms
in the molecule.
Consider carbon dioxide CO2
• carbon (C) has four valence electrons x 1 carbon = 4 eoxygen (O) has six valence electrons x 2 oxygens = 12 eThere are a total of 16 e- to be placed in the Lewis structure.
2. Connect the central atom to the other atoms
in the molecule with single bonds.
Carbon is the central atom, the two oxygens are bound to it
and electrons are added to fulfill the octets of the outer
atoms. (: or
means shared pair of electrons)
Some atoms can form multiple bonds (i.e.: share more than one pair of
electrons )
3. Complete the valence shell of the outer
atoms in the molecule.
4. Place any remaining electrons on the
central atom.
– there are no more available electrons
– If the valence shell of the central atom is not complete,
use a lone pair on one of the outer atoms to form a
double bond between that outer atom and the central
atom. Continue this process of making multiple bonds
between the outer atoms and the central atom until the
valence shell of the central atom is complete.
becomes
The central atom is still electron deficient,
so share another pair.
becomes
5. Double check to make sure that you have
used the correct number of electrons in
the Lewis structure and that no atom that
cannot exceed its valence shell, does not.
The best Lewis structure that can be drawn for
carbon dioxide is:
What Is VSEPR?
• The Valence Shell Electron Pair Repulsion (VSEPR)
model:
– is based on the number of regions of high electron density
around a central atom.
– can be used to predict structures of molecules or ions that
contain only non-metals by minimizing the electrostatic
repulsion between the regions of high electron density.
– can also be used to predict structures of molecules or ions that
contain multiple bonds or unpaired electrons.
– does fail in some cases.
Intermolecular forces
• Dipole—dipole
– A dipole is created by equal but opposite charges that
are separated by a short distance
– The direction of a dipole is from its positive pole to its
negative pole
used to represent
– Attraction between dipoles of polar molecules called
dipole-dipole force
– Polar molecules can induce a temporary
attraction in nonpolar molecules
– Example:
–
>
>
• Hydrogen bonding
– The intermolecular force in which a hydrogen
atom that is bonded to a highly electronegative
atom is attracted to an unshared pair of
electrons of an electronegative atom in a nearby
molecule
– Represented by dotted lines connecting the
bonded hydrogen to the unshared pair of the
electronegative atom to which it is attracted
• London dispersion forces
– The intermolecular attractions resulting from
the constant motion of electrons and the
creation of instantaneous dipoles
– Forces between all atoms and molecules
(including Noble gases)
– Increase with atomic number
Boiling Points
• The strength of the intermolecular forces is
directly correlated to the boiling points.
• The higher the bp, the more energy
required to break the intermolecular forces
• Higher bp, stronger the intermolecular
forces