View
52
Download
0
Category
Tags:
Preview:
DESCRIPTION
Chemistry 100 Chapter 8. Chemical Bonding Basic Concepts. The Valance Electrons. When atoms interact to form chemical bonds, only the outer (valance) electrons take part. Need a tool for keeping track of valence electrons, e.g., The Lewis dot symbol 1 v.E. 7 v.E’s - PowerPoint PPT Presentation
Citation preview
Chemistry 100 Chapter 8
Chemical Bonding Basic Concepts
The Valance Electrons
When atoms interact to form chemical bonds, only the outer (valance) electrons take part.
Need a tool for keeping track of valence electrons, e.g., The Lewis dot symbol
1 v.E. 7 v.E’s
When these two elements combine to form a compound
2 Na (s) + Cl2 (g) ® 2 NaCl (s)
+ NaClNa ...Cl: :.
What’s Happening?
[Ne]3s1 [Ne]3s23p5
(g) ® Na+ (g) + e- (ionizes, loses e-) an electron configuration of [Ne]
(g) + e- ® Cl- (g) an electron configuration of [Ar]
In the crystal lattice, Na+ and Cl- ions; strong electrostatic
attractions
Na.
+ NaClNa ...Cl: :.
Cl::...
The NaCl Crystal
Ionic Bonding
Electrostatic attractions that hold ions together in an ionic compound.
The strength of interaction depends on charge magnitude and distance between them.
rqkqE 21
ionic q1 ® magnitude of charge 1q2 ® magnitude of charge 2r ® distance between the ionic centres
Stability of Ionic Compounds
The stability of ionic compounds depends on two main factors
1. The electron affinity of one of the elements2. The ionization energy of the other
Note electron affinities and ionization potentials
are gas-phase reactions? How are they related to the stability of solid
materials?
The Lattice Energy
A quantitative measure of just how strong the interaction is between the ionic centres (i.e., a measure of the strength of the ionic bond)
For the reaction KCl (s) ® K+ (g) + Cl- (g) H = 718 kJ/mol
Lattice energy (latH). The energy required to completely separate
one mole of the solid ionic compound into its gas-phase ions.
Lattice Energies of Various Ionic Compounds
Determined using a thermochemical cycle -the Born-Haber cycle (a Hess’s Law application)
Covalent Bonding
In a wide variety of molecules, the bonding atoms fulfill their valance shell requirements by sharing electrons between them.
Covalent bonds - a bond in which the electrons are shared by two atoms.
H2 ® H-H, F2 ® F-F, Cl2 ® Cl-Cl For many electron atoms (like F and Cl), we
again to worry only about the outermost (valence) electrons.
Covalent Bonding
Examples of Covalent Bonding
Let’s look at the Cl2 example. Each Cl atom has 7 valence shell
electrons 3 Lone pairs and one unpaired electron
Cl::...
Lone pairs
Unshared electron
The Cl2 Molecule
The structure we have just drawn are called Lewis structures.
The dash between the atomic centres represents bonding electrons
Redraw F2
Cl Cl:... . :
....
lone pairs (non bonding)bonding electrons
F F:... . :
....
Note both Cl2 and F2 satisfy their valence shell requirements by the formation of a single bond.
What about O2? How can we satisfy the octet rule for 2 O atoms?
O O... .
....
Valence shell requirements are satisfied by the formation of a double bond.
check out N2 ® :NºN: (triple bond) Note that the octet rule works
mainly for the second row elements. Filled valence shells can have more
than 8 electrons after Z=14 (Si). This is generally termed octet expansion.
Covalent Compounds
Compounds that contain only covalent bonds are called covalent compounds.
There are two main of covalent compounds, Molecular covalent compounds (CO2, C2H4) Network covalent compounds (SiO2, BeCl2). The network covalent compound are
characterized by an extensive “3-D” network bonding
Comparison between Ionic and Covalent Compounds
Ionic Compounds usually solids with very
high melting points conduct electricity when
molten (melted) usually quite water
soluble and they are electrolytes in aqueous solution
NaCl
Covalent Compounds usually low melting solids,
gases or liquids don’t conduct electricity
when molten aren’t very soluble in water
and are non electrolytes CCl4
The Filled Valence Shell rule
Filled Valence Shell rule Atoms participate in the formation of bonds
(either ionic or covalent) in order to satisfy their valence shell requirements.
Atoms other than H tend to form bonds until they end up being surrounded by 8 valence electrons (the noble gas configuration). Your text calls this the “octet” rule.
Electronegativity
Electronegativity is defined as the ability of an atom to attract electrons towards itself in a molecule ( (pronounced ‘chi’))
Examine the H-F covalent bond+H-F
denotes a partial “+” charge on the H atom
- denotes a partial “-“ charge on F atom
Electronegativity is related to the electron affinity and the ionization energy.
Compare the following elements. Na ® low I1, small negative E.A. ® low F ® high I1, large, negative E.A., ® high
Trends in the Values
Across a row The values generally increase as we
proceed from left to right in the periodic table.
Down a group The values generally decrease as we
descend the group. Transition metals
Essentially constant values
Plot of Values
Electronegativity and Bond Type
Can we use the electronegativity values to help us deduce the type of bonding in compounds?
values bond type
0.0 < < 0.5 non-polar covalent0.5 1.9 polar covalent
2.0 3.3 Ionic bond
An Outline for Drawing Lewis Structures
Predict arrangement of atoms (i.e., predict the skeletal arrangement of the molecule or ion). · The H is always a terminal atom,
bonded to ONE OTHER ATOM ONLY. A halogen atom is usually a terminal atom.
· Note that the central atom usually has the least negative electron affinity.
Count total number of valence shell electrons (include ionic charges).
Place 1 pair electrons (sigma bond, ) between each pair of bonded atoms (i.e., the central atom and each one of the terminal atoms).
Place remaining electrons around the terminal atoms to satisfy the filled valence shell rule. (lone pairs).
All remaining electrons are assigned to the central atom. Atoms in the 3rd or higher row can have more than eight electrons around them. If a central atom does not have a filled
valence shell, use a lone pair of electrons from a terminal atom to make a pi () bond.
Formal Charges
Definition: formal charge on atom = number of valence electrons – number of non-bonding - ½ the number of bonding electrons.
Formal charge in a Lewis Structure is a bookkeeping “device” keeps track of the electrons “associated”
with certain atoms in the molecule vs. the valence e-‘s in the isolated atom!
How does it work?
Rules for Formal Charges
Neutral molecules ® S formal charges = 0 Ions ® S formal charges = charge of ion For molecules where the possibility of multiple
Lewis Structures with different formal charges exist Neutral molecule - choose the structure with
the fewest formal charges. Structures with large formal charges are less
likely than ones with small formal charges Two Lewis Structures with similar formal
charge distribution ® negative formal charges on more electronegative atom
Resonance Structures
Examine the NO3- anion.
The structures differ in the location of the N=O double bond.
They are said to be resonance structures. The actual structure of the molecule is a
combination of three resonance structures (the resonance hybrid).
N
O
O O N
O
OO
Experimental Evidence for Resonance.
The resonance structures for benzene C6H6
We would expect to find two different bond lengths in benzene (C=C and C-C bonds). C= C ® bond length = 133 pm = 0.133 nm C-C ® bond length = 0.154 nm
Experimentally, all benzene carbon-carbon bond lengths are equivalent at 0.140 nm
Exceptions to the Filled Valence Shell Rule
Be compounds BeH2, BeCl2, Boron and Al compounds BF3, AlCl3, BCl3 BF3 is stable The B central atom has a tendency
to pick up an unshared e- pair from another compound
BF3 + NH3 ® BF3NH3 the B-N bond is an example of a coordinate
covalent bond, or a “dative” bond ® i.e. a bond in which one of the atoms donates both bonding electrons.
Odd e- molecules
These molecules have uneven numbers of electrons \ no way that they can form octets.
Examples NO and NO2. These species have an odd
number of electrons.
N OO.
..
......:N O
.....
:
Look at the dimerization reaction of NO2
.
2 NO2 (g) ⇄ N2O4 (g) Keq = 210
N NO
OO
O: ..: : :
: : : :..
Valence Shells having more than 8 Electrons (Expanded Octets)
A central atom having more than 8 valance shell electrons is possible with atomic number 14 and above.
Cl
Cl ClPCl Cl
Reason - elements in this category can use the energetically low-lying d orbitals to
accommodate extra electrons
Look at HClO3
High formal charge on the electronegative Cl atom (f.c.(Cl) = 7-2-1/2 (6) = +2)
This resonance structure would make a very small contribution to the overall resonance hybrid.
Cl O
O
O H......:
..: :
..
..
With the possibility of using the low lying d-orbitals on the Cl atom to accommodate extra electron pairs, we may write other Lewis structures
Note: the final three structures reduce the formal charges
Cl O
O
O H......:
..: :
..
.. Cl O
O
O H......
..: :
..
.. Cl O
O
O H......
: :....:
Cl O
O
O H......
: :....
Bond Energies and Thermochemistry
Look at the energy required to break 1 mole of gaseous diatomic molecules into their constituent gaseous atoms.
H2 (g) ® H (g) + H (g) H° = 436.4 kJCl2 (g) ® Cl (g) + Cl (g) H° = 242 kJ
These enthalpy changes are called bond dissociation energies. In the above examples, the enthalpy changes are designated D (H-H) and D (Cl-Cl).
For Polyatomic Molecules.
CO2 (g) ® C (g) + 2 O (g) H = 745 kJ Denote the H of this reaction D(C=O) What about dissociating methane into C + 4
H’s?CH4 (g) ® C(g) + 4 H (g) H° = 1650 kJ
Note 4 C-H bonds in CH4 \ D (C-H) = 412 kJ/mol
H2O (g) ® 2 H (g) + O (g) H° = 929 kJ/mol H2O It takes more energy to break the first O-H
bond.H2O (g) ® H (g) + OH (g) H° = 502 kJ/mol H2OHO (g) ® H (g) + O (g) H = 427 kJ/mol H2O
Note: we realize that all chemical reactions involve the breaking and reforming of chemical bonds. Break bonds ® add energy. Make bonds ® energy is released.
rxnH° S D(bonds broken) - S D(bonds formed)
These are close but not quite exact. Why? The bond energies we use are averaged bond
energies, i.e., This is a good approximate for equations
involving diatomic species. We can only use the above procedure for GAS
PHASE REACTIONS ONLY.
Recommended