chapter 8 bonding: general concepts. questions to consider what is meant by the term “chemical...

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?

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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


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.


The Interaction of Two

Hydrogen Atoms


The Interaction of Two Hydrogen Atoms


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.


• 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


• 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


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 ::

::


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.


• 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).


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


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?



What is the general trend for electronegativity across rows and down columns on the periodic table?

Explain the trend.



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


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



Which of the following bonds would be the most polar without being considered ionic?

Mg–O C–O O–O Si–O N–O



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.


Dipole Moment

57

No Net Dipole Moment (Dipoles Cancel)


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.


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.


Isoelectronic Series

• A series of ions/atoms containing the same number of electrons.

O2-, F-, Ne, Na+, Mg2+, and Al3+


Ionic Radii




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.



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


• 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.


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




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]


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]


Operational Definition of Ionic Compound

• Any compound that conducts an electric current when melted will be classified as ionic.


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.


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.


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).


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).


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


3 3CH N C( ) CH C N( ) g g

C N


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.



• 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



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.


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


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.


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


Step 3: Distribute the remaining electrons to the peripheral atoms to satisfy the octet rule.

PClCl

Cl

:: :

:

:

:

:

:

:


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

:: :

:

:

:

:

:

:

:


• Try drawing Lewis dot formulas for the following covalent compound.

SCl220 e- total

ClSCl

16 e- left

::

: :

::

::

4 e- left

Cl

C

Cl

O


• Try drawing Lewis dot formulas for the following covalent compound.

COCl224 e- total18 e- left

::

: :

::

0 e- left:

: :

Cl

C

Cl

O


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:

: :


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-



formal charge on C = 4 -2 - ½ x 6 = -1

formal charge on O = 6 -2 - ½ x 6 = +1


=1

2




-

-1 +1

C – 4 e-

O – 6 e-

2H – 2x1 e-

12 e-



HC O

H

formal charge on C = 4 -0 - ½ x 8 = 0

formal charge on O = 6 -4 - ½ x 4 = 0


=1

2




-

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


• 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.


• 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


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.


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


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


For example, the bonding in phosphorus pentafluoride, PF5, shows ten electrons surrounding the phosphorus.

: F :

::: F :

F ::

:

: F

:: PF :

::


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.


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.


VSEPR




VSEPR: Two Electron Pairs




VSEPR: Three Electron Pairs




VSEPR: Four Electron Pairs




VSEPR: Iodine Pentafluoride




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


MolecularGeometry

VSEPR

AB3 3 0trigonal planar

trigonal planar


Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


trigonal planar

AB4 4 0 tetrahedral tetrahedral


Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


trigonal planar


AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal


Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


trigonal planar


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


MolecularGeometry

VSEPR


trigonal planar

AB2E 2 1trigonal planar

bent

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

AB3E 3 1


tetrahedraltrigonal

pyramidal

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


AB3E 3 1 tetrahedraltrigonal

pyramidal

AB2E2 2 2 tetrahedral bent

H

O

H

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

AB4E 4 1trigonal

bipyramidaldistorted

tetrahedron

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

AB4E 4 1trigonal


tetrahedron

AB3E2 3 2trigonal

bipyramidalT-shaped

ClF

F

F

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR

AB5 5 0trigonal

bipyramidaltrigonal

bipyramidal

AB4E 4 1trigonal


tetrahedron

AB3E2 3 2trigonal

bipyramidalT-shaped

AB2E3 2 3trigonal

bipyramidallinear

I

I

I

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


AB5E 5 1 octahedral square pyramidal

Br

F F

FF

F

Class

# of atomsbonded to

central atom

# lonepairs on


MolecularGeometry

VSEPR


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


• 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 ::

::


• 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 :

::

: :



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:: :

::

:



Since one of the groups is a lone pair, the molecular geometry is described as bent

or angular.

O O

O:: :

::

:



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:: :

::

:


• 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 :



:Cl:

:::Cl:

:Cl

::

C

H

N

H

H :

:O

H

H :

tetrahedral

Cl:

::



:Cl:

:::Cl:

:Cl

::

C

:O

H

H :

tetrahedral

Cl:

::

H

NH H :

trigonal pyramid



:Cl:

:::Cl:

:Cl

::

C

:

O

H H :

tetrahedral

Cl:

::

trigonal pyramid bent

H

NH H :


• Five electron pairs (trigonal bipyramidal arrangement).

This structure results in both 90o and 120o bond angles.

: F :

::: F :

F :

::

: F

:: P

F ::

:


• 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.



SClF3 XeF2

F

F

FF

see-saw

:



XeF2

see-saw

S

F

F

FF

: Cl

F

F

::

F

T-shape



see-saw

S

F

F

FF

: Cl

F

F

::

F

T-shape

Xe

F

F

::

:

linear


• Six electron pairs (octahedral arrangement).

This octahedral arrangement results in 90o bond angles.

F:

::

:F

::

SF:

::

:F

::

:F:

:

:F: :



Six electron pairs also lead to other molecular geometries.

IF5 XeF4



XeF4I

FFF

:

FF

square pyramid



I

FFF

:

FF

square pyramid

XeF

F

:

FF

:

square planar

chapter 8 bonding: general concepts. questions to consider what is meant by the term “chemical...

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