Lewis Electron-Dot Symbols : Lewis Electron-Dot Symbols A Lewis electron-dot symbol is a symbol in which the electrons in the valence shell of an atom or ion are represented by dots placed around the letter symbol of the element.
Note that the group number indicates the number of valence electrons. 10/11/2010 1
Lewis Electron-Dot Formulas : Lewis Electron-Dot Formulas A Lewis electron-dot formula is an illustration used to represent the transfer of electrons during the formation of an ionic bond.
The magnesium has two electrons to give, whereas the fluorines have only one “vacancy” each.
Consequently, magnesium can accommodate two fluorine atoms. 10/11/2010 2
Lewis Structures : Lewis Structures The tendency of atoms in a molecule to have eight electrons in their outer shell (two for hydrogen) is called the octet rule.
You can represent the formation of the covalent bond in H2 as follows:
This uses the Lewis dot symbols for the hydrogen atom and represents the covalent bond by a pair of dots. 3 + 10/11/2010
The Electron Probability Distribution for the H2 Molecule : The Electron Probability Distribution for the H2 Molecule 4 10/11/2010
Lewis Structures : Lewis Structures The shared electrons in H2 spend part of the time in the region around each atom.
In this sense, each atom in H2 has a helium (1s2) configuration. 5 10/11/2010
Lewis Structures : Lewis Structures The formation of a bond between H and Cl to give an HCl molecule can be represented in a similar way.
Thus, hydrogen has two valence electrons about it (as in He) and Cl has eight valence electrons about it (as in Ar) 6 10/11/2010
Lewis Structures : Lewis Structures Formulas such as these are referred to as Lewis electron-dot formulas or Lewis structures
An electron pair is either a:
bonding pair (shared between two atoms)
lone pair (electron pair that is not shared)
Hydrogen has no unbonded pairs
Chlorine has 3 unbonded pairs 7 10/11/2010
Lewis Structures : Lewis Structures Rules for obtaining Lewis electron dot formulas
Calculate the number of valence electrons for the molecule (= group # for each atom), add the charge of anion and subtract the charge of a cation
Write the skeleton structure of the molecule or ion
Put atom with lowest group number and lowest electronegativity as the central atom
Distribute electrons to atoms surrounding the central atom to satisfy the octet rule for each atom
Distribute the remaining electrons as pairs to the central atom (or atoms)
If the Central atom is deficient in electrons to complete the octet; move electron pairs from surrounding atoms to complete central atom valence electron needs 8 10/11/2010
Practice Problem : Practice Problem Write a lewis structure for CCl2F2
Step 1: Arrange Atoms (Carbon is “Central Atom” because is has the lowest group number and lowest electronegativity.
Step 2: Determine total number of valence electrons
1 x C(4) + 2 x Cl(7) + 2 x F(7) = 32
Step 3: Draw single bonds to central atom and subtract 2 e- for each single bond (4 x 2 = 8) 32 – 8 = 24 remaining
Step 4: Distribute remaining electrons in pairs around surrounding atoms 9 10/11/2010
Writing Lewis Dot Formulas : Writing Lewis Dot Formulas The Lewis electron-dot formula of a covalent compound is a simple two-dimensional representation of the positions of electrons in a molecule
Bonding electron pairs are indicated by either two dots or a dash
In addition, these formulas show the positions of lone pairs of electrons 10/11/2010 10
Writing Lewis Dot Formulas : Writing Lewis Dot Formulas The following rules allow you to write electron-dot formulas even when the central atom does not follow the octet rule.
To illustrate, draw the structure of:
Phosphorus Trichloride. 10/11/2010 11 Con’t on next slide
Writing Lewis Dot Formulas : Writing Lewis Dot Formulas Step 1: Total all valence electrons in the molecular formula. That is, total the group numbers of all the atoms in the formula.
For a polyatomic anion, add the number of negative charges to this total.
For a polyatomic cation, subtract the number of positive charges from this total. 10/11/2010 12 P 3s23p3 Cl 3s23p5
5 + 21(3x7) = 26 e- total Con’t on next slide
Writing Lewis Dot Formulas : Writing Lewis Dot Formulas Step 2: Arrange the atoms radially, with the least electronegative atom in the center. Place one pair of electrons between the central atom and each peripheral atom. 26 – 6 = 20 remaining 10/11/2010 13 Con’t on next slide
Writing Lewis Dot Formulas : Writing Lewis Dot Formulas Step 3: Distribute the remaining electrons to the peripheral atoms to satisfy the octet rule. 14 26 – (3 x 6 + 6) = 2 remaining 10/11/2010 Con’t on next slide
Writing Lewis Dot Formulas : Writing Lewis Dot Formulas Step 4: Distribute any remaining electrons (2) to the central atom. If the number of electrons on the central atom is less than the number of electrons required to complete the octet for that atom, use one or more electrons pairs from other atoms to form double or triple bonds. 15 10/11/2010
Exceptions to the Octet Rule : Exceptions to the Octet Rule Although many molecules obey the octet rule, there are exceptions where the central atom has more than eight electrons
Generally, if a nonmetal is in the third period or greater it can accommodate as many as twelve electrons, if it is the central atom
These elements have unfilled “d” subshells that can be used for bonding 16 10/11/2010
Exceptions to the Octet Rule : Exceptions to the Octet Rule For example, the bonding in phosphorus pentafluoride, PF5, shows ten electrons surrounding the phosphorus. 17 Total valence electrons
5 x 7 + 5 = 40
Distribute electrons to F atoms
5 x 6 = 30
Establish bonding pairs
5 x 2 = 10
Remaining electrons
40 – 30 – 10 = 0
Phosphorus has “0” non-bonding pairs
Since Phosphorus is in Period 3, PF5 is a “hypervalent” molecule and the phosphorus utilizes electrons from other shells to create valence shell with more than 8 electrons 10/11/2010
Exceptions to the Octet Rule : Exceptions to the Octet Rule In xenon tetrafluoride, XeF4, the xenon atom must accommodate two extra lone pairs. 18 Total valence electrons
4 x 7 + 8 = 36
Distribute electrons to F atoms
4 x 6 = 24
Establish bonding pairs
4 x 2 = 8
Remaining electrons
36 – 24 – 8 = 4
Add 2 non-bonding pairs to Xe
Xe violates “octet” rule
XeF4 is a “hypervalent” molecule and utilizes “d” orbitals to create valence shell with more than 8 electrons 10/11/2010
Delocalized Bonding: Resonance : Delocalized Bonding: Resonance The structure of ozone, O3, can be represented by two different Lewis electron-dot formulas
Experiments show, however, that both bonds are identical 19 or 10/11/2010 Ozone (O3)
Delocalized Bonding: Resonance : Delocalized Bonding: Resonance According to Resonance Theory, these two equal bonds are represented as one pair of bonding electrons spread over the region of all three atoms
This is called delocalized bonding, in which a bonding pair of electrons is spread over a number of atoms 20 10/11/2010 Ozone (O3)
Formal Charge & Lewis Structures : Formal Charge & Lewis Structures In certain instances, more than one feasible Lewis structure can be illustrated for a molecule.For example, H, C and N
The concept of “formal charge” can help discern which structure is the most likely
Formal Charge:
An atom’s formal charge is:
Total number of valence electrons
Minus all unshared electrons
Minus ½ of its shared electrons
Formal Charges must sum to actual charge of species:
Zero Charge for a Molecule
Ionic Charge for an Ion 21 10/11/2010
Formal Charge & Lewis Structures : Formal Charge & Lewis Structures When you can write several Lewis structures, choose the one having the least formal charge 22 FCH: [1 - 0 - ½(2)] = 0
FCC: [4 - 0 - ½(8)] = 0
FCN: [5 - 2 - ½(6)] = 0 FCH: [1 - 0 - ½(2)] = 0
FCC: [4 - 2 - ½(6)] = -1
FCN: [5 - 0 - ½(8)] = +1 10/11/2010 Note: HCN is a neutral molecule
Sum of Formal Charges in the preferred form (0) equals molecular charge (0) FC: Total Valence e- – unshared e- – ½ shared e- Form I Form II Preferred Form - Form I (Least Formal Charge)
Formal Charge & Lewis Structures : Formal Charge & Lewis Structures 23 Ozone FCOA: [6 - 4 - ½(4)] = 0
FCOB: [6 - 2 - ½(6)] = +1
FCOC: [6 - 6 - ½(2)] = -1 FCOA: [6 - 6 - ½(2)] = -1
FCOB: [6 - 2 - ½(6)] = +1
FCOC: [6 - 4 - ½(4)] = 0 Both “Resonance” forms have the same formal charge
and thus, are identical
Note: Ozone (O3) is a neutral molecule
Sum of Formal Charges (0) equals molecular charge (0) 10/11/2010
Formal Charge & Lewis Structures : Formal Charge & Lewis Structures 10/11/2010 24 FC B = 3 – 0 -(1/2 * 6)
= 0
Even though B violates “Octet Rule”, this is the preferred form because it has “less” formal charge FC B = 3 – 0 -(1/2 * 8)
= -1
FC F = 7 – 4 - (1/2 * 4)
= +1 Boron Trifuoride
BF3 Sulfur Dioxide
SO2 FC S = 6 – 2 – (1/2 * 6) = 1 FC S = 6 – 2 – (1/2 * 8) = 0
Preferred Form (Less Formal Charge)
Resonance/Formal Charge – Nitrate Ion : Resonance/Formal Charge – Nitrate Ion 10/11/2010 25 Total Valence electrons - 3 x 6 (O) + 1 x 5 (N) + 1 (ion charge) = 24
Add 1 pair electrons between central atom and each other atom – 3 x 2 = 6
Add electrons to oxygen atoms to complete octet
Nitrogen still missing 2 electrons to complete octet
Borrow 2 electrons from one oxygen to form double bond
Formal Charge – Nitrogen: 5 – (0 + ½*8) = 5 – 4 = +1
Formal Charge – Single bonded Oxygen: 6 – (6 + ½*2) = 6 – 7 = -1 x 2 = -2
Formal Charge – Double bonded Oxygen: 6 – (4 + ½*4) = 6 – 6 = 0
Net Charge of ion is: +1 +(-2) = -1
Resonance/Formal Charge – Cyanate Ion : Resonance/Formal Charge – Cyanate Ion 10/11/2010 26 FCN = 5 – (6 + ½*1) = -2
FCC = 4 – (0 + ½*8) = 0
FCO = 6 – (2 + ½*6) = +1 FCN = 5 – (4 + ½*4) = -1
FCC = 4 – (0 + ½*8) = 0
FCO = 6 – (4 + ½*4) = 0 FCN = 5 – (2 + ½*6) = 0
FCC = 4 – (0 + ½*8) = 0
FCO = 6 – (6 + ½*2) = -1 Preferred Form:
Eliminate I – Higher formal charge on Nitrogen than Carbon & Oxygen
Positive formal charge on Oxygen, which is more electronegative than Nitrogen
Eliminate II – Forms II & III have the same magnitude of formal charges, but form III has a -1 charge on the more electronegative Oxygen atom
Forms II & III both contribute to the resonant hybrid of the Cyanate Ion,
but form III is the more important
Note: Net formal charge in form III is same as ionic charge (-1)
Formal Charge vs Oxidation No : Formal Charge vs Oxidation No “Formal Charge” is used to examine resonance hybrid structures , whereas “Oxidation Number” is used to monitor “REDOX” reactions
Formal Charge - Bonding electrons are assigned equally to the atoms as if the bonding were “Nonpolar” covalent, i.e., each atom has half the electrons making up the bond
Formal Charge = valence e- – (lone pair e- + ½ bonding e-)
Oxidation Number - Bonding electrons are assigned completely to the more electronegative atom, as if the bonding were “Ionic”
Oxidation No. = valence e- – (lone pair e- + bonding e-) 10/11/2010 27
The Valence-Shell Electron Pair Repulsion Model (VSEPR) : The Valence-Shell Electron Pair Repulsion Model (VSEPR) 10/11/2010 28 The valence-shell electron pair repulsion (VSEPR) model predicts the shapes of molecules and ions by assuming that the valence shell electron pairs are arranged as far from one another as possible.
Molecular geometry – shape of a molecule is determined by the positions of atomic nuclei relative to each other, i.e., angular arrangement
Central Atom
Place atom with “Lower Group Number” in center(N in NF3 needs more electrons to complete octet)
If atoms have same group number (SO3 or ClF3), place the atom with the “Higher Period Number” in the center.
VSEPR Model of Molecular Shapes : VSEPR Model of Molecular Shapes The following rules and figures will help discern electron pair arrangements.
Select the Central Atom (Least Electronegative Atom)
Draw the Lewis structure
Determine how many bonding electron pairs are around the central atom.
Determine the number of non-bonding electron pairs
Count a multiple bond as “one pair”
Arrange the electron pairs as far apart as possible to minimize electron repulsions
Note the number of bonding and lone pairs 10/11/2010 29
VSEPR Model of Molecular Shapes : VSEPR Model of Molecular Shapes To predict the relative positions of atoms around a given atom using the VSEPR model, you first note the arrangement of the electron pairs around that central atom.
Molecular Notation:
A – The Central Atom (Least Electronegative atom)
X – The Ligands (Bonding Pairs)
a – The Number of Ligands
E – Non-Bonding Electron Pairs
b – The Number of Non-Bonding Electron Pairs
Double & Triple Bonds count as a “single” electron pair
The Geometric arrangement is determined by:
sum (a + b) 10/11/2010 30 AXaEb
VSEPR Model of Molecular Shapes : VSEPR Model of Molecular Shapes 10/11/2010 31
VSEPR Model of Molecular Shapes : VSEPR Model of Molecular Shapes 10/11/2010 32
Arrangement of Electron Pairs About an Atom: Basic Shapes : Arrangement of Electron Pairs About an Atom: Basic Shapes 10/11/2010 33 CS2, HCN, BeF2
Arrangement of Electron Pairs About an Atom: Basic Shapes : Arrangement of Electron Pairs About an Atom: Basic Shapes 10/11/2010 34 SO3 BF3 NO3− CO32− SO2 O3 PbCl2 SnBr2
Arrangement of Electron Pairs About an Atom: Basic Shapes : Arrangement of Electron Pairs About an Atom: Basic Shapes 10/11/2010 35 CH4 SiCl4SO42- ClO4- NH3 PF3 ClO3 H3O+ H2O OF2 SCl2
Arrangement of Electron Pairs About an Atom: Basic Shapes : Arrangement of Electron Pairs About an Atom: Basic Shapes 10/11/2010 36 SF4, XeO2F2, IF4+, IO2F2- ClF3 BrF3 XeF2 I3- IF2- PF5 AsF5 SOF4
Arrangement of Electron Pairs About an Atom: Basic Shapes : Arrangement of Electron Pairs About an Atom: Basic Shapes 10/11/2010 37 SF6
IOF5 BrF5 TeF5- XeOF4 XeF4 ICl4-
Electron Pair Arrangement : Electron Pair Arrangement 10/11/2010 38
Electron Pair Arrangement : Electron Pair Arrangement 10/11/2010 39
Linear Geometry : Linear Geometry Two electron pairs (linear arrangement)
Double bonds count as a “single electron pair”
2 bonding pairs
0 non-bonding pairs
AXaEb = a + b = 2 + 0 = 2 (Linear)
Thus, according to the VSEPR model, the bonds are arranged linearly (bond angle = 180o)
Molecular shape of carbon dioxide is linear 10/11/2010 40 Carbon is central atom because it has lower group number
Trigonal Planar Geometry : Trigonal Planar Geometry Three electron pairs on Central atom
The three groups of electron pairs are arranged in a trigonal plane. Thus, the molecular shape of COCl2 is trigonal planar. The Bond angle is 120o. 10/11/2010 41 Central Atom - Carbon
3 bonding electron pairs
(double bond counts as 1 pair)
0 non-bonding electron pairs
a + b = 3 + 0 = 3
Trigonal Planar
Trigonal Planar Geometry : Trigonal Planar Geometry Effect of Double Bonds
Bond angles deviate from ideal angles when surrounding atoms and electron groups are not identical.
A double bond has greater electron density and repels two single bonds more strongly than they repel each other 10/11/2010 42
Trigonal Planar Geometry : Trigonal Planar Geometry Effect of Lone Pairs
The molecular shape is defined only by the positions of the nuclei
When one of the three electron pairs in a trigonal planar molecule is a lone (non-bonding) pair, it is held by only one nucleus
It is less confined and exerts a stronger repulsive force than a bonding pair
This results in a decrease in the angle between the bonding pairs 10/11/2010 43 The normal Trigonal Planar angle between the bonding pairs is 120o
Trigonal Planar Geometry : Trigonal Planar Geometry Three electron pairs (Effect of ‘Lone’ pairs)
(trigonal planar arrangement)
Ozone has two bonding electron pairs about the central oxygen (double bond counts as 1 pair.)
There is one non-bonding lone pair.
These groups have a:
Trigonal Planar arrangement
AXaEb (a + b = 2 + 1 = 3)
Since one of the groups is a lone pair, the molecular geometry is described as bent or angular. 10/11/2010 44 SO3
BF3
NO3-
CO32-
Tetrahedral Geometry : Tetrahedral Geometry Four electron pairs
(Tetrahedral Arrangement)
Four electron pairs about the central atom lead to three different molecular geometries.
a + b = 4 + 0 a + b = 3 + 1 a + b = 2 + 2
= 4 = 4 = 4 10/11/2010 45
Tetrahedral Geometry : Tetrahedral Geometry Molecular Geometries produced by variable non-bonding electron pairs 10/11/2010 46 CH4, SiCl4, SO42-, ClO4- PF3, ClO3-, H3O+ OF2, SCl2
Trigonal Bipyramidal : Trigonal Bipyramidal Five electron pairs
(trigonal bipyramidal arrangement)
This structure results in both 90o and 120o bond angles. 10/11/2010 47 ASF5 SOF4
Trigonal Bipyramidal : Trigonal Bipyramidal Other molecular geometries are possible when one or more of the electron pairs is a lone pair. 10/11/2010 48 XeO2F2 IF4+ IOF2- ClF3 BrF3 XeF2 I3- IF2-
Octahedral Geometry : Octahedral Geometry Six electron pairs
(Octahedral arrangement)
This octahedral arrangement results in:
90o bond angles 10/11/2010 49 SF6 IOF5
Other Geometries : Other Geometries Six electron pairs
(octahedral arrangement) 10/11/2010 50 square pyramidal square planar BrF5 TeF5- XeOF4 XeF4 ICl4- Iodine violates octet rule
Iodine is sp3d2 hybridized
Iodine uses d orbitals Noble gases not
always inert
Xenon forms
6 electron domains
Sample Problem : Sample Problem In the ICl4– ion, the electron pairs are arranged around the central iodine atom in the shape of
a. a tetrahedron.
b. a trigonal bipyramid.
c. a square plane.
d. an octahedron.
e. a trigonal pyramid.
Ans: a 10/11/2010 51 AXaEb
a + b = 4 + 0 = 4 (AX4 – Tetrahedral) AX4
Dipole Moment andMolecular Geometry : Dipole Moment andMolecular Geometry 10/11/2010 52 The dipole moment is a measure of the degree of charge separation in a molecule.
The polarity of individual bonds within a molecule can be viewed as vector quantities.
Thus, molecules that are perfectly symmetric have a zero dipole moment. These molecules are considered nonpolar.
Dipole Moment andMolecular Geometry : Dipole Moment andMolecular Geometry 10/11/2010 53 However, molecules that exhibit any asymmetry in the arrangement of electron pairs would have a nonzero dipole moment. These molecules are considered polar NH3 PF3 ClO3 H3O+
Dipole Moment andMolecular Geometry : Dipole Moment andMolecular Geometry 10/11/2010 54
Sample Problem : Sample Problem The nitrogen atom would be expected to have the positive end of the dipole in the species
a. NH4+
b. Ca3N2
c. HCN
d. AlN
e. NO+
Ans: e 10/11/2010 55 N is more electronegative than H N is more EN than Ca N is more EN than C N is more EN than Al O is more EN than Nitrogen
Practice Problem : Practice Problem Which of the following molecules is polar?
a. BF3 b. CBr4 c. CO2
d. NO2 e. SF6
Ans: d 10/11/2010 56 The Lewis structures for BF3, CBr4, CO2, and SF6 do not have any non-bonding electrons on the central atom
The Lewis structure for NO2 shows one double bond and a lone non-bonding electron on the Nitrogen
The VESPR Molecular Geometry for NO2 is AX2E1 (a + b = 2 + 1 = 3) - Trigonal Planar
Formal Charge on N is 5 – 1 -1/2(6) = +1
NO2 molecule is polar
Practice Problem : Practice Problem Which of the following compounds is nonpolar?
a. XeF2 b. HCl c. SO2
d. H2S e. N20
Ans: a
HCL is ionic and very polar
SO2 has AX2E1 Trigonal Planar Bent geometry with a dipole moment (Polar)
H2S has AX2E2 Tetrahedral Bent geometry and with a dipole moment (polar)
N2O has AX2E0 linear with asymmetric geometry.Since oxygen is more EN than N, the molecule is polar
XeF2 has AX2E3 Trigonal Bypyramidal Geometry, but linear molecular geometry (nonpolar) 10/11/2010 57