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Chapter 8
Bonding:General Concepts
Questions to Consider
• What is meant by the term “chemical bond”?
• Why do atoms bond with each other to form compounds?
• How do atoms bond with each other to form compounds?
Copyright © Cengage Learning. All rights reserved 2
Electron Arrangement and the Periodic Table
• Electron configuration - describes the arrangement of electrons in atoms.
• The electron arrangement is the primary factor in understanding how atoms join together to form compounds.
• Valance electrons - the outermost electrons.– These are the electrons involved in chemical
bonding.
3
• For the representative elements:
– The number of valance electrons is the group number.
– The period number gives the energy level (n) of the valance shell.
• For an atom of fluorine, how many valance electrons does it have and what is the energy
level of these electrons?
• Fluorine has 7 electrons in the n=2 level
Valance Electrons
• Let’s look at fluorine more closely.
• What is the total number of electrons in fluorine?
– The atomic number is 9. It therefore has 9 protons and 9 electrons.
• If there are 7 electrons in the valance shell, (with n = 2 energy level) where are the other
two electrons?
– In the n = 1 energy level. This level holds two and only two electrons.
• Isoelectronic - they have the same electron configuration (same number of electrons)
• Nonmetallic elements tend to form negatively charged ions called anions.
• Nonmetals tend to gain electrons so they become isoelectronic with its nearest noble gas neighbor.
O
[He]2s22p4
+ 2e- O2-
[He]2s22p6 or [Ne]
Electron Configurations of Cations and Anions
Na [Ne]3s1 Na+ [Ne]
Ca [Ar]4s2 Ca2+ [Ar]
Al [Ne]3s23p1 Al3+ [Ne]
Atoms lose electrons so that cation has a noble-gas outer electron configuration.
H 1s1 H- 1s2 or [He]
F 1s22s22p5 F- 1s22s22p6 or [Ne]
O 1s22s22p4 O2- 1s22s22p6 or [Ne]
N 1s22s22p3 N3- 1s22s22p6 or [Ne]
Atoms gain electrons so that anion has a noble-gas outer electron configuration.
Of Representative Elements
+1
+2
+3 -1-2-3
Cations and Anions Of Representative Elements
Na+: [Ne] Al3+: [Ne] F-: 1s22s22p6 or [Ne]
O2-: 1s22s22p6 or [Ne] N3-: 1s22s22p6 or [Ne]
Na+, Al3+, F-, O2-, and N3- are all isoelectronic with Ne
What neutral atom is isoelectronic with H- ?
H-: 1s2 same electron configuration as He
Electron Configurations of Cations of Transition Metals
When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals.
Fe: [Ar]4s23d6
Fe2+: [Ar]4s03d6 or [Ar]3d6
Fe3+: [Ar]4s03d5 or [Ar]3d5
Mn: [Ar]4s23d5
Mn2+: [Ar]4s03d5 or [Ar]3d5
Copyright © Cengage Learning. All rights reserved 11
A Chemical Bond
• No simple, and yet complete, way to define this.• Forces that hold groups of atoms together and make
them function as a unit.• A bond will form if the energy of the aggregate is
lower than that of the separated atoms.
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The Interaction of Two
Hydrogen Atoms
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The Interaction of Two Hydrogen Atoms
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Key Ideas in Bonding
• Ionic Bonding – electrons are transferred• Covalent Bonding – electrons are shared equally by
nuclei• What about intermediate cases?
Describing Ionic Bonds
• An ionic bond is a chemical bond formed by the electrostatic attraction between positive and negative ions.
This type of bond involves the transfer of electrons from one atom (usually a metal) to another (usually a nonmetal).The number of electrons lost or gained by an atom is determined by its need to be “isoelectronic” with a noble gas.
Let’s examine the formation of NaClNa + Cl NaCl
IONIC BONDING
Sodium has alow ionization energyit readily loses this
electron .
When Sodium loses the electron, it gains the Ne configuration.
Na Na+ + e-
Chlorine has a high electron affinity.
When chlorine gains an electron, it gains the Ar configuration.
:
..
..Cl: e
..
..Cl:
Essential Features of Ionic Bonding
• Atoms with low I.E. and low E.A. tend to form positive ions.
• Atoms with high I.E. and high E.A. tend to form negative ions.
• Ion formation takes place by electron transfer.
• The ions are held together by the electrostatic force of the opposite charges.
• Reactions between metals and nonmetals (representative) tend to be ionic.
Describing Ionic Bonds
• Such noble gas configurations and the corresponding ions are particularly stable.
The atom that loses the electron becomes a cation (positive).
The atom that gains the electron becomes an anion (negative).
-1 e ])Ne([Na)s3]Ne([Na
)p3s3]Ne([Cle )p3s3]Ne([Cl 62-52
Describing Ionic Bonds
• Consider the transfer of valence electrons from a sodium atom to a chlorine atom.
The resulting ions are electrostatically attracted to one another.The attraction of these oppositely charged ions for one another is the ionic bond.
ClNaClNa
e-
Electron Configurations of Ions
• As metals lose electrons to form cations and establish a “noble gas” configuration, the electrons are lost from the valence shell first.
For example, magnesium generally loses two electrons from its 3s subshell to look like neon.
)e 2( Mg Mg -2
[Ne]3s2 [Ne]
Electron Configurations of Ions
• Transition metals also lose electrons from the valence shell first, which is not the last subshell to fill according to the aufbau sequence.
For example, zinc generally loses two electrons from its 4s subshell to adopt a “pseudo”-noble gas configuration.
)e 2( Zn Zn -2
[Ar]4s23d10 [Ar]3d10
Li + F Li+ F -
The Ionic Bond
1s22s1 1s22s22p5 1s2 1s22s22p6
[He] [Ne]
Li Li+ + e-
e- + F F -
F -Li+ + Li+ F -
Covalent Bonds
When two nonmetals bond, they often share electrons since they have similar attractions for them. This sharing of valence electrons is called the covalent bond.
These atoms will share sufficient numbers of electrons in order to achieve a noble gas electron configuration (that is, eight valence electrons).
A covalent bond is a chemical bond in which two or more electrons are shared by two atoms.
Why should two atoms share electrons?
F F+
7e- 7e-
F F
8e- 8e-
F F
F F
Lewis structure of F2
lone pairslone pairs
lone pairslone pairs
single covalent bond
single covalent bond
II. Covalent Compounds
• Covalent compounds are usually formed from nonmetals.
• Molecules - compounds characterized by covalent bonding.
• not a part of a massive three dimensional crystal structure.
COVALENT BONDING
Let’s look at the formation of H2
H + H H2
• Each hydrogen has one electron in it’s valance shell.
• If it were an ionic bond it would look like this:
:H H H H
• However, both hydrogen atoms have the same tendency to gain or lose electrons.
• Both gain and loss will not occur.
• Instead, each atom gets a noble gas configuration by sharing electrons.
H:H H H
The shared electrons pair is a
Covalent Bond
Each Hydrogen atom now has two electrons around it and has a He
configuration
Features of Covalent Bonds
• Covalent bonds tend to form between atoms with similar tendency to gain or lose electrons.
• The diatomic elements have totally covalent bonds (totally equal sharing.)
H2, N2, O2, F2, Cl2, Br2, I2
:..
..F:
..
..F: :
..
..F
..
..F:
Each fluorine has eight electrons around it. Ne’s configuration.
8e-
H HO+ + OH H O HHor
2e- 2e-
Lewis structure of water
Double bond – two atoms share two pairs of electrons
single covalent bonds
O C O or O C O
8e- 8e-8e-double bonds double bonds
Triple bond – two atoms share three pairs of electrons
N N8e-8e-
N N
triple bondtriple bond
or
Lengths of Covalent Bonds
Bond Lengths
Triple bond < Double Bond < Single Bond
Copyright © Cengage Learning. All rights reserved 32
Polar Covalent Bond
• Unequal sharing of electrons between atoms in a molecule.
• Results in a charge separation in the bond (partial positive and partial negative charge).
Polar Covalent Bonds
A polar covalent bond is one in which the bonding electrons spend more time near one of the two atoms involved.
When the atoms are alike, as in the H-H bond of H2 , the bonding electrons are shared equally (a nonpolar covalent bond). When the two atoms are of different elements, the bonding electrons need not be shared equally, resulting in a “polar” bond.
Polar Covalent Bonding and Electronegativity
The Polar Covalent Bond
• Ionic bonding involves the transfer of electrons.
• Covalent bonding involves the sharing of electrons.
• Polar covalent bonding - bonds made up of unequally shared electron pairs.
1
H F FH
Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms
electron richregion
electron poorregion e- riche- poor
+ -
:..F:H :
..F H
These two electrons
are not sharedequally.
• The electrons spend more time with fluorine.
• This sets up a polar bond
• A truly covalent bond can only occur when both atoms are identical.
• Electronegativity is used to determine if a bond is polar and who gets the electrons the most.
somewhat negatively chargedsomewhat positively charged
Polar Covalent Bonds
For example, the bond between carbon and oxygen in CO2 is considered polar because the shared electrons spend more time orbiting the oxygen atoms.
The result is a partial negative charge on the oxygens (denoted )and a partial positive charge on the carbon (denoted )
C OO ::
::
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The Effect of an Electric Field on Hydrogen Fluoride Molecules
indicates a positive or negative fractional charge. or
Polar Molecules
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What is meant by the term “chemical bond?”
Why do atoms bond with each other to form molecules?
How do atoms bond with each other to form molecules?
CONCEPT CHECK!CONCEPT CHECK!
• The ability of an atom in a molecule to attract shared electrons to itself.
• For a molecule HX, the relative electronegativities of the H and X atoms are determined by comparing the measured H–X bond energy with the “expected” H–X bond energy.
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• On the periodic table, electronegativity generally increases across a period and decreases down a group.
• The range of electronegativity values is from 4.0 for fluorine (the most electronegative) to 0.7 for cesium (the least electronegative).
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Periodic Properties• Electron Affinity
The electron affinity is the energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion.
For a chlorine atom, the first electron affinity is illustrated by:
)p3s3]Ne([Cle)p3s3]Ne([Cl 6252 Electron Affinity = -349 kJ/mol
Periodic Properties• Electron Affinity
The more negative the electron affinity, the more stable the negative ion that is formed.
Broadly speaking, the general trend goes from lower left to upper right as electron affinities become more negative.
The Pauling Electronegativity Values
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Polar Covalent BondsThe absolute value of the difference in electronegativity of two bonded atoms gives a rough measure of the polarity of the bond.
When this difference is small (less than 0.5), the bond is nonpolar.When this difference is large (greater than 0.5), the bond is considered polar.If the difference exceeds approximately 1.8, sharing of electrons is no longer possible and the bond becomes ionic.
Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-
Increasing difference in electronegativity
Classification of bonds by difference in electronegativity
Difference Bond Type
0 Covalent
2 Ionic
0 < and <2 Polar Covalent
• The greater the difference in electronegativity between two atoms, the greater the polarity of a
bond.
• Which would be more polar, a H-F bond or a H-Cl bond?
• H-F …4.0 - 2.1 = 1.9
• H-Cl… 3.0 - 2.1 = 0.9
• Therefore, the HF bond is more polar than the HCl bond.
Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H2S; andthe NN bond in H2NNH2.
Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic
H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent
N – 3.0 N – 3.0 3.0 – 3.0 = 0 Covalent
If lithium and fluorine react, which has more attraction for an electron? Why?
In a bond between fluorine and iodine, which has more attraction for an electron? Why?
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CONCEPT CHECK!CONCEPT CHECK!
What is the general trend for electronegativity across rows and down columns on the periodic table?
Explain the trend.
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CONCEPT CHECK!CONCEPT CHECK!
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Arrange the following bonds from most to least polar:
a) N–F O–F C–Fb) C–F N–O Si–Fc) Cl–Cl B–Cl S–Cla) C–F, N–F, O–Fb) Si–F, C–F, N–Oc) B–Cl, S–Cl, Cl–Cl
Copyright © Cengage Learning. All rights reserved 53
EXERCISE!EXERCISE!
Which of the following bonds would be the least polar yet still be considered polar covalent?
Mg–O C–O O–O Si–O N–O
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CONCEPT CHECK!CONCEPT CHECK!
Which of the following bonds would be the most polar without being considered ionic?
Mg–O C–O O–O Si–O N–O
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CONCEPT CHECK!CONCEPT CHECK!
Dipole Moment
• Property of a molecule whose charge distribution can be represented by a center of positive charge and a center of negative charge.
• Use an arrow to represent a dipole moment.– Point to the negative charge center with the
tail of the arrow indicating the positive center of charge.
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Dipole Moment
57
No Net Dipole Moment (Dipoles Cancel)
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Dipole Moments and Polar Molecules
H F
electron richregion
electron poorregion
= Q x rQ is the charge
r is the distance between charges
1 D = 3.36 x 10-30 C m
Dipole Moment and Molecular Geometry
Molecules that exhibit any asymmetry in the arrangement of electron pairs would have a nonzero
dipole moment. These molecules are considered polar.
H
NH H
:
Which of the following molecules have a dipole moment?H2O, CO2, SO2, and CH4
O HH
dipole momentpolar molecule
SO
O
CO O
no dipole momentnonpolar molecule
dipole momentpolar molecule
C
H
H
HH
no dipole momentnonpolar molecule
Does BF3 have a dipole moment?
Does CH2Cl2 have a dipole moment?
Stable Compounds
• Atoms in stable compounds usually have a noble gas electron configuration.
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Electron Configurations in Stable Compounds
• When two nonmetals react to form a covalent bond, they share electrons in a way that completes the valence electron configurations of both atoms.
• When a nonmetal and a representative-group metal react to form a binary ionic compound, the ions form so that the valence electron configuration of the nonmetal achieves the electron configuration of the next noble gas atom. The valence orbitals of the metal are emptied.
Copyright © Cengage Learning. All rights reserved 67
Isoelectronic Series
• A series of ions/atoms containing the same number of electrons.
O2-, F-, Ne, Na+, Mg2+, and Al3+
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Ionic Radii
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Choose an alkali metal, an alkaline earth metal, a noble gas, and a halogen so that they constitute an isoelectronic series when the metals and halogen are written as their most stable ions.
– What is the electron configuration for each species?– Determine the number of electrons for each species.– Determine the number of protons for each species.
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CONCEPT CHECK!CONCEPT CHECK!
Periodic Table Allows Us to Predict Many Properties
• Trends for:– Atomic size, ion radius, ionization energy,
electronegativity• Electron configurations• Formula prediction for ionic compounds• Covalent bond polarity ranking
Copyright © Cengage Learning. All rights reserved 71
• What are the factors that influence the stability and the structures of solid binary ionic compounds?
• How strongly the ions attract each other in the solid state is indicated by the lattice energy.
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Energy Involved in Ionic Bonding
• The transfer of an electron from a sodium atom to a chlorine atom is not in itself energetically favorable; it requires 147 kJ/mol of energy.
However, 493 kJ of energy is released when these oppositely charged ions come together.An additional 293 kJ of energy is released when the ion pairs solidify.This “lattice energy” is the negative of the energy released when gaseous ions form an ionic solid. The next slide illustrates this.
Energy Involved in Ionic Bonding
Lattice energy (E) increases as Q increases and/or
as r decreases.
cmpd lattice energyMgF2
MgO
LiF
LiCl
2957
3938
1036
853
Q= +2,-1
Q= +2,-2
r F- < r Cl-
Electrostatic (Lattice) Energy
E = kQ+Q-r
Q+ is the charge on the cation
Q- is the charge on the anionr is the distance between the ions
Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions.
Born-Haber Cycle for Determining Lattice Energy
Hoverall = H1 + H2 + H3 + H4 + H5o ooooo
Born-Haber Cycle for NaCl
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Formation of an Ionic Solid
1.Sublimation of the solid metal.• M(s) M(g) [endothermic]
2. Ionization of the metal atoms.• M(g) M+(g) + e [endothermic]
3. Dissociation of the nonmetal.• 1/2X2(g) X(g) [endothermic]
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Formation of an Ionic Solid (continued)
4 Formation of nonmetal ions in the gas phase.• X(g) + e X(g) [exothermic]
5.Formation of the solid ionic compound.• M+(g) + X(g) MX(s)
[quite exothermic]
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Operational Definition of Ionic Compound
• Any compound that conducts an electric current when melted will be classified as ionic.
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Ionic Compounds - Review
• Metals and nonmetals usually react to form ionic compounds.
• The metals are the cations and the nonmetals are the anions.
• The cations and anions arrange themselves in a regular three-dimensional repeating array called
a crystal lattice.
Models
• Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.
Copyright © Cengage Learning. All rights reserved 83
Fundamental Properties of Models
1. A model does not equal reality.2. Models are oversimplifications, and are
therefore often wrong.3. Models become more complicated and are
modified as they age.4. We must understand the underlying assumptions
in a model so that we don’t misuse it.5. When a model is wrong, we often learn much
more than when it is right.
Copyright © Cengage Learning. All rights reserved 84
Bond Energies
• To break bonds, energy must be added to the system (endothermic, energy term carries a positive sign).
• To form bonds, energy is released (exothermic, energy term carries a negative sign).
Copyright © Cengage Learning. All rights reserved 85
Bond Energies
∆H = Σn×D(bonds broken) – Σn×D(bonds formed)
D represents the bond energy per mole of bonds (always has a positive sign).
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The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy.
H2 (g) H (g) + H (g) H0 = 436.4 kJ
Cl2 (g) Cl (g)+ Cl (g) H0 = 242.7 kJ
HCl (g) H (g) + Cl (g) H0 = 431.9 kJ
O2 (g) O (g) + O (g) H0 = 498.7 kJ O O
N2 (g) N (g) + N (g) H0 = 941.4 kJ N N
Bond Energy
Bond Energies
Single bond < Double bond < Triple bond
Average bond energy in polyatomic molecules
H2O (g) H (g) + OH (g) H0 = 502 kJ
OH (g) H (g) + O (g) H0 = 427 kJ
Average OH bond energy = 502 + 427
2= 464 kJ
Bond Energies (BE) and Enthalpy changes in reactions
H0 = total energy input – total energy released= BE(reactants) – BE(products)
Imagine reaction proceeding by breaking all bonds in the reactants and then using the gaseous atoms to form all the bonds in the products.
Use bond energies to calculate the enthalpy change for:H2 (g) + F2 (g) 2HF (g)
H0 = BE(reactants) – BE(products)
Type of bonds broken
Number of bonds broken
Bond energy (kJ/mol)
Energy change (kJ)
H H 1 436.4 436.4
F F 1 156.9 156.9
Type of bonds formed
Number of bonds formed
Bond energy (kJ/mol)
Energy change (kJ)
H F 2 568.2 1136.4
H0 = 436.4 + 156.9 – 2 x 568.2 = -543.1 kJ
Bond Energy
To illustrate, let’s estimate the H for the following reaction.
In this reaction, one C-H bond and one Cl-Cl bond must be broken.In turn, one C-Cl bond and one H-Cl bond are formed.
)g(HCl)g(ClCH)g(Cl)g(CH 324
Bond Energy
simple arithmetic yields H.
)g(HCl)g(ClCH)g(Cl)g(CH 324
)ClCl(BE)HC(BEH )ClH(BE)ClC(BE
kJ )428327240411(H
kJ 104H
Predict ∆H for the following reaction:
Given the following information: Bond Energy (kJ/mol)
C–H 413
C–N 305
C–C 347 891
∆H = –42 kJ
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3 3CH N C( ) CH C N( ) g g
C N
CONCEPT CHECK!CONCEPT CHECK!
Localized Electron Model
• A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.
Copyright © Cengage Learning. All rights reserved 94
Localized Electron Model
• Electron pairs are assumed to be localized on a particular atom or in the space between two atoms:– Lone pairs – pairs of electrons localized on an
atom– Bonding pairs – pairs of electrons found in the
space between the atoms
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Localized Electron Model
1. Description of valence electron arrangement (Lewis structure).
2. Prediction of geometry (VSEPR model).3. Description of atomic orbital types used by
atoms to share electrons or hold lone pairs.
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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.
Na..
. ..
Si .. .
:P
: .
:
.SMg. .
.Al..
:
:Cl.: Ar
:
:
::
Group I Group II Group VII Group VIIIGroup VIGroup IV Group VGroup III
Lewis symbol (Lewis structure) - a way to represent atoms (and their bonds) using the element symbol and
valence electrons as dots.
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.
As an example, let’s look at the transfer of electrons from magnesium to fluorine to form magnesium fluoride.
Lewis Electron-Dot Formulas
• Consequently, magnesium can accommodate two fluorine atoms.
::
F.:
:
:F .: Mg. .
Mg[ F ]
:
:
:
:- 2+ [ F ]
:
:
:
:-
The magnesium has two electrons to give, whereas the fluorines have only one “vacancy” each.
Drawing Lewis Structures on Molecules and Polyatomic Ions
Lewis Structure Guidelines
1.Use chemical symbols for the various elements to write the skeletal structure of the compound.– the least electronegative atom will be placed in
the central position,– hydrogen and fluorine occupy terminal positions,– carbon often forms chains of carbon-carbon
covalent bonds.
5
2.Determine the total number of valence electrons associated with each atom in the compound.– for polyatomic cations, subtract one electron for
every positive charge;– for polyatomic anions, add one electron for
every negative charge.
3.Connect the central atom to each of the surrounding atoms using electron pairs. Then give each atom an octet.– Remember, hydrogen needs only two electrons
4. Count the number of electrons you have and compare to the number you used.
• If they are the same, you are finished.• If you used more electrons than you
have add a bond for every two too many you used. Then give every atom an octet.
• If you used less electrons than you have….(see later when discuss exceptions to the octet rule)
5. Check that all atoms have the octet rule satisfied and that the total number of valance electrons are used.
Covalent Bonds
The tendency of atoms in a molecule to have eight electrons in their outer shell (two for hydrogen) is called the octet rule.
Lewis Structures
You can represent the formation of the covalent bond in H2 as follows:
H .. :H H H+This uses the Lewis dot symbols for the hydrogen atom and represents the covalent bond by a pair of dots.
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 configuration.
:H H
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).
:H
:
::ClH. . ::
Cl:+
Lewis Structures
Formulas such as these are referred to as Lewis electron-dot formulas or Lewis structures.
::H Cl:
:An electron pair is either a bonding pair (shared between two atoms) or a lone pair (an electron pair that is not shared).
bonding pair
lone pair
Coordinate Covalent Bonds
When bonds form between atoms that both donate an electron, you have:
A .. :B A B+It is, however, possible that both electrons are donated by one of the atoms. This is called a coordinate covalent bond.
A : :B A B+
Multiple Bonds
In the molecules described so far, each of the bonds has been a single bond, that is, a covalent bond in which a single pair of electrons is shared.
It is possible to share more than one pair. A double bond involves the sharing of two pairs between atoms.
CC
H
H
H
H
orC:CH
H
H
H: : ::
:
Triple bonds are covalent bonds in which three pairs of electrons are shared between atoms.
Multiple Bonds
CC orHH
::
CC HH
:::
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, we will draw the structure of PCl3, phosphorus trichloride.
3PCl
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.
3PCl5 e- (7 e-) x 3
26 e- total
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.
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.
PClCl
Cl
Writing Lewis Dot Formulas
Step 3: Distribute the remaining electrons to the peripheral atoms to satisfy the octet rule.
PClCl
Cl
:: :
:
:
:
:
:
:
Writing Lewis Dot Formulas
Step 4: Distribute any remaining electrons to the central atom. If there are fewer than eight electrons on the central atom, a multiple bond may be necessary.
PClCl
Cl
:: :
:
:
:
:
:
:
:
Writing Lewis Dot Formulas
• Try drawing Lewis dot formulas for the following covalent compound.
SCl220 e- total
ClSCl
16 e- left
::
: :
::
::
4 e- left
Cl
C
Cl
O
Writing Lewis Dot Formulas
• Try drawing Lewis dot formulas for the following covalent compound.
COCl224 e- total18 e- left
::
: :
::
0 e- left:
: :
Cl
C
Cl
O
Writing Lewis Dot Formulas
Note that the carbon has only 6 electrons.One of the oxygens must share a lone pair.
COCl224 e- total18 e- left
::
: :
::
0 e- left:
: :
Writing Lewis Dot Formulas
Note that the carbon has only 6 electrons.One of the oxygens must share a lone pair.
COCl224 e- total
Cl
C
Cl
18 e- left
::
: :
::
0 e- leftO: :
Note that the octet rule is now obeyed.
Write the Lewis structure of nitrogen trifluoride (NF3).
Step 1 – N is less electronegative than F, put N in center
F N F
F
Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)
5 + (3 x 7) = 26 valence electrons
Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
Write the Lewis structure of the carbonate ion (CO32-).
Step 1 – C is less electronegative than O, put C in center
O C O
O
Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4) -2 charge – 2e-
4 + (3 x 6) + 2 = 24 valence electrons
Step 3 – Draw single bonds between C and O atoms and complete octet on C and O atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
Step 5 - Too many electrons, form double bond and re-check # of e-
2 single bonds (2x2) = 41 double bond = 4
8 lone pairs (8x2) = 16Total = 24
Using the guidelines presented, write Lewis structures for the following:
1. H2O
2. NH3
3. CO2
4. NH4+
5. CO32-
6. N2
NNor ..N
..N
OOor ..O::
..O H - Hor H:H
Bond energy - the amount of energy required to
break a bond holding two atoms together.
triple bond > double bond > single bond
Bond length - the distance separating the nuclei of two adjacent atoms.
single bond > double bond > triple bond
Two possible skeletal structures of formaldehyde (CH2O)
H C O HH
C OH
An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.
formal charge on an atom in a Lewis structure
=1
2
total number of bonding electrons( )
total number of valence electrons in the free atom
-total number of nonbonding electrons
-
The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.
H C O HC – 4 e-
O – 6 e-
2H – 2x1 e-
12 e-
2 single bonds (2x2) = 41 double bond = 4
2 lone pairs (2x2) = 4Total = 12
formal charge on C = 4 -2 - ½ x 6 = -1
formal charge on O = 6 -2 - ½ x 6 = +1
formal charge on an atom in a Lewis structure
=1
2
total number of bonding electrons( )
total number of valence electrons in the free atom
-total number of nonbonding electrons
-
-1 +1
C – 4 e-
O – 6 e-
2H – 2x1 e-
12 e-
2 single bonds (2x2) = 41 double bond = 4
2 lone pairs (2x2) = 4Total = 12
HC O
H
formal charge on C = 4 -0 - ½ x 8 = 0
formal charge on O = 6 -4 - ½ x 4 = 0
formal charge on an atom in a Lewis structure
=1
2
total number of bonding electrons( )
total number of valence electrons in the free atom
-total number of nonbonding electrons
-
0 0
Formal Charge and Lewis Structures
1. For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present.
2. Lewis structures with large formal charges are less plausible than those with small formal charges.
3. Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms.
Which is the most likely Lewis structure for CH2O?
H C O H
-1 +1 HC O
H
0 0
Formal Charge and Lewis Structures
• In certain instances, more than one feasible Lewis structure can be illustrated for a molecule. For example,
H C N CNHor: :
The concept of “formal charge” can help discern which structure is the most likely.
Formal Charge and Lewis Structures
• The formal charge of an atom is determined by subtracting the number of electrons in its “domain” from its group number.
H C N CNHor: :
The number of electrons in an atom’s “domain” is determined by counting one electron for each bond and two electrons for each lone pair.
1 e- 4 e- 5 e- 1 e- 4 e- 5 e-
“domain” electrons
group number
I IV V I V IV
Formal Charge and Lewis Structures
The most likely structure is the one with the least number of atoms carrying formal charge. If they have the same number of atoms carrying formal charge, choose the structure with the negative formal charge on the more electronegative atom.
In this case, the structure on the left is most likely correct.
orH C N:0 0 0
CNH :formal charge
0 +1 -1
Lewis Structures and Resonance
• Write the Lewis structure at CO32- on your
paper.
• If you look at the people around you they probably put the double bond in different
places.
• Who is right?
• You all are.
A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure.
O O O+ -
OOO+-
O C O
O
- -O C O
O
-
-
OCO
O
-
-
What are the resonance structures of the carbonate (CO3
2-) ion?
Delocalized Bonding: Resonance
• According to theory, one pair of bonding electrons is 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.
OO
O
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.
Lewis Structures and Exceptions to the Octet Rule
1. Incomplete Octet - less then eight electrons around an atom other than H.
• Let’s look at BF3
2. Odd Electron - if there is an odd number of valence electrons it isn’t possible to give every atom eight
electrons.
• Let’s look at NO
3. Expanded Octet - elements in 3rd period and below may have 10 and 12 electrons around it.
Exceptions to the Octet Rule
The Incomplete Octet
H HBeBe – 2e-
2H – 2x1e-
4e-
BeH2
BF3
B – 3e-
3F – 3x7e-
24e-
F B F
F
3 single bonds (3x2) = 69 lone pairs (9x2) = 18
Total = 24
Exceptions to the Octet Rule
Odd-Electron Molecules
N – 5e-
O – 6e-
11e-
NO N O
The Expanded Octet (central atom with principal quantum number n > 2)
SF6
S – 6e-
6F – 42e-
48e-
S
F
F
F
FF
F
6 single bonds (6x2) = 1218 lone pairs (18x2) = 36
Total = 48
Exceptions to the Octet Rule
For example, the bonding in phosphorus pentafluoride, PF5, shows ten electrons surrounding the phosphorus.
: F :
::: F :
F ::
:
: F
:: PF :
::
Exceptions to the Octet Rule
In xenon tetrafluoride, XeF4, the xenon atom must accommodate two extra lone pairs.
F ::
:
: F :
:
XeF :
::
: F
::
::
VSEPR Model
• VSEPR: Valence Shell Electron-Pair Repulsion.• The structure around a given atom is determined
principally by minimizing electron pair repulsions.
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Steps to Apply the VSEPR Model
1. Draw the Lewis structure for the molecule.2. Count the electron pairs and arrange them in the
way that minimizes repulsion (put the pairs as far apart as possible.
3. Determine the positions of the atoms from the way electron pairs are shared (how electrons are shared between the central atom and surrounding atoms).
4. Determine the name of the molecular structure from positions of the atoms.
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VSEPR
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VSEPR: Two Electron Pairs
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VSEPR: Three Electron Pairs
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VSEPR: Four Electron Pairs
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VSEPR: Iodine Pentafluoride
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Predicting Molecular Geometry
• The following rules and figures will help discern electron pair arrangements.
1.Draw the Lewis structure2.Determine how many electrons pairs are
around the central atom. Count a multiple bond as one pair.
3.Arrange the electrons pairs.
Arrangement of Electron Pairs About an Atom
3 pairsTrigonal planar
2 pairsLinear
4 pairsTetrahedral
5 pairsTrigonal bipyramidal
6 pairsOctahedral
Valence shell electron pair repulsion (VSEPR) model:
Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs.
AB2 2 0
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
linear linear
B B
Cl ClBe
2 atoms bonded to central atom
0 lone pairs on central atom
AB2 2 0 linear linear
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB3 3 0trigonal planar
trigonal planar
AB2 2 0 linear linear
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB3 3 0trigonal planar
trigonal planar
AB4 4 0 tetrahedral tetrahedral
AB2 2 0 linear linear
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB3 3 0trigonal planar
trigonal planar
AB4 4 0 tetrahedral tetrahedral
AB5 5 0trigonal
bipyramidaltrigonal
bipyramidal
AB2 2 0 linear linear
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB3 3 0trigonal planar
trigonal planar
AB4 4 0 tetrahedral tetrahedral
AB5 5 0trigonal
bipyramidaltrigonal
bipyramidal
AB6 6 0 octahedraloctahedral
bonding-pair vs. bondingpair repulsion
lone-pair vs. lone pairrepulsion
lone-pair vs. bondingpair repulsion> >
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB3 3 0trigonal planar
trigonal planar
AB2E 2 1trigonal planar
bent
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB3E 3 1
AB4 4 0 tetrahedral tetrahedral
tetrahedraltrigonal
pyramidal
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB4 4 0 tetrahedral tetrahedral
AB3E 3 1 tetrahedraltrigonal
pyramidal
AB2E2 2 2 tetrahedral bent
H
O
H
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB5 5 0trigonal
bipyramidaltrigonal
bipyramidal
AB4E 4 1trigonal
bipyramidaldistorted
tetrahedron
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB5 5 0trigonal
bipyramidaltrigonal
bipyramidal
AB4E 4 1trigonal
bipyramidaldistorted
tetrahedron
AB3E2 3 2trigonal
bipyramidalT-shaped
ClF
F
F
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB5 5 0trigonal
bipyramidaltrigonal
bipyramidal
AB4E 4 1trigonal
bipyramidaldistorted
tetrahedron
AB3E2 3 2trigonal
bipyramidalT-shaped
AB2E3 2 3trigonal
bipyramidallinear
I
I
I
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB6 6 0 octahedraloctahedral
AB5E 5 1 octahedral square pyramidal
Br
F F
FF
F
Class
# of atomsbonded to
central atom
# lonepairs on
central atomArrangement of electron pairs
MolecularGeometry
VSEPR
AB6 6 0 octahedraloctahedral
AB5E 5 1 octahedral square pyramidal
AB4E2 4 2 octahedral square planar
Xe
F F
FF
Predicting Molecular Geometry1. Draw Lewis structure for molecule.
2. Count number of lone pairs on the central atom and number of atoms bonded to the central atom.
3. Use VSEPR to predict the geometry of the molecule.
What are the molecular geometries of SO2 and SF4?
SO O
AB2E
bent
S
F
F
F F
AB4E
distortedtetrahedron
Predicting Molecular Geometry
• Two electron pairs (linear arrangement).
You have two double bonds, or two electron groups about the carbon atom.
Thus, according to the VSEPR model, the bonds are arranged linearly, and the
molecular shape of carbon dioxide is linear. Bond angle is 180o.
C OO ::
::
Predicting Molecular Geometry
• Three electron pairs (trigonal planar arrangement).
The three groups of electron pairs are arranged in a trigonal plane. Thus, the molecular shape of COCl2 is trigonal
planar. Bond angle is 120o.
Cl
C
:
::
O
Cl :
::
: :
Predicting Molecular Geometry
• Three electron pairs (trigonal planar arrangement).
Ozone has three electron groups about the central oxygen. One group is a lone pair.
These groups have a trigonal planar arrangement.
O O
O:: :
::
:
Predicting Molecular Geometry
• Three electron pairs (trigonal planar arrangement).
Since one of the groups is a lone pair, the molecular geometry is described as bent
or angular.
O O
O:: :
::
:
Predicting Molecular Geometry
• Three electron pairs (trigonal planar arrangement).
Note that the electron pair arrangement includes the lone pairs, but the molecular
geometry refers to the spatial arrangement of just the atoms.
O O
O:: :
::
:
Predicting Molecular Geometry
• Four electron pairs (tetrahedral arrangement).
Four electron pairs about the central atom lead to three different molecular
geometries.
:Cl:
:::Cl:
:Cl
:: C Cl:
::
H
N
H
H :
:O
H
H :
Predicting Molecular Geometry
• Four electron pairs (tetrahedral arrangement).
:Cl:
:::Cl:
:Cl
::
C
H
N
H
H :
:O
H
H :
tetrahedral
Cl:
::
Predicting Molecular Geometry
• Four electron pairs (tetrahedral arrangement).
:Cl:
:::Cl:
:Cl
::
C
:O
H
H :
tetrahedral
Cl:
::
H
NH H :
trigonal pyramid
Predicting Molecular Geometry
• Four electron pairs (tetrahedral arrangement).
:Cl:
:::Cl:
:Cl
::
C
:
O
H H :
tetrahedral
Cl:
::
trigonal pyramid bent
H
NH H :
Predicting Molecular Geometry
• Five electron pairs (trigonal bipyramidal arrangement).
This structure results in both 90o and 120o bond angles.
: F :
::: F :
F :
::
: F
:: P
F ::
:
Predicting Molecular Geometry
• Other molecular geometries are possible when one or more of the electron pairs is a lone pair.
SF4 ClF3 XeF2
Let’s try their Lewis structures.
Predicting Molecular Geometry
• Other molecular geometries are possible when one or more of the electron pairs is a lone pair.
SClF3 XeF2
F
F
FF
see-saw
:
Predicting Molecular Geometry
• Other molecular geometries are possible when one or more of the electron pairs is a lone pair.
XeF2
see-saw
S
F
F
FF
: Cl
F
F
::
F
T-shape
Predicting Molecular Geometry
• Other molecular geometries are possible when one or more of the electron pairs is a lone pair.
see-saw
S
F
F
FF
: Cl
F
F
::
F
T-shape
Xe
F
F
::
:
linear
Predicting Molecular Geometry
• Six electron pairs (octahedral arrangement).
This octahedral arrangement results in 90o bond angles.
F:
::
:F
::
SF:
::
:F
::
:F:
:
:F: :
Predicting Molecular Geometry
• Six electron pairs (octahedral arrangement).
Six electron pairs also lead to other molecular geometries.
IF5 XeF4
Predicting Molecular Geometry
• Six electron pairs (octahedral arrangement).
XeF4I
FFF
:
FF
square pyramid
Predicting Molecular Geometry
• Six electron pairs (octahedral arrangement).
I
FFF
:
FF
square pyramid
XeF
F
:
FF
:
square planar