Chapter 10 Chemical Bonding

Chapter 10 Chemical Bonding

HIV-protease

Atoms interact with other atoms to form molecules, this is chemical bondingBonding theories – are models that predict how atoms bond together to form molecules

Bonding Theories are applied to design molecules that will interfere with the activesite of HIV-protease. This delays or inhibitsthe onset of AIDS.

protease inhibitor

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CHAPTER OUTLINE Chemical Bonds Ionic Bonds and Covalent Bonds Electronegativity Bond Polarity & Electronegativity Lewis Structures Resonance Molecular Shapes Molecular Polarity

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CHEMICAL BOND

Most matter in nature is found in form of compounds: 2 or more elements held together through a chemical bond.

Elements combine together (bond) to fill their outer energy levels and achieve a stable structure (low energy).

Noble gases are un-reactive since their energy levels are complete.

The nature and type of the chemical bond is directly responsible for many physical and chemical properties of a substance: (e.g. melting point, conductivity).

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CHEMICAL BOND

When the conductivity apparatus is placed in salt solution, the bulb will light.

But when it is placed in sugar solution, the bulb does not light.

This difference in conductivity between salt and sugar is due to the different types of bonds between their atoms.

Two common types of bonding are present: ionic & covalent.

Gilbert Newton Lewis (1875 - 1946) was a famous American physical chemist known for the discovery of the covalent bond (see his Lewis dot structures and his 1916 paper "The Atom and the Molecule")

Other major contributions were his theory of Lewis acids and bases andLewis coined the term "photon" for the smallest unit of radiant energy.

Lewis is known for:Covalent bondLewis dot structuresValence bond theoryElectronic theory of acids and basesChemical thermodynamicsHeavy waterNamed photonExplained phosphorescence

The Origin of Lewis Symbols of Atoms

Drawings of cubical atoms, the corners of the cube represented possible electron positions

Lewis later cited these notes in his classic 1916 paper on chemical bonding, as being the first expression of his ideas.

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LEWIS SYMBOLS OF ATOMS

Lewis structures use Lewis symbols to show valence electrons in molecules and ions of compounds.

Lewis symbols for the first 3 periods of representative elements are shown below:

In Lewis symbols, valence electrons for each element are shown as a dot.

Lewis Bonding Theory

• atoms bond because it results in a more stable electron configuration

• atoms bond together by either transferring or sharing electrons so that all atoms obtain an outer shell with 8 electrons– Octet Rule– there are some exceptions to this rule – the

key to remember is to try to get an electron configuration like a noble gas

Lewis Symbols of Ions

• Cations have Lewis symbols without valence electrons– Lost in the cation formation

• Anions have Lewis symbols with 8 valence electrons– Electrons gained in the formation of the

anion

Li• Li+1 :F: [:F:]-1

•

•• ••

••

Ionic Bonds

• metal to nonmetal• metal loses electrons to form cation• nonmetal gains electrons to form anion• ionic bond results from + to - attraction

– larger charge = stronger attraction– smaller ion = stronger attraction

• Lewis Theory allow us to predict the correct formulas of ionic compounds

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IONIC BOND

Ionic bonds occur when electrons are transferred between two atoms.

After bonding, each atom achieves a complete shell (noble gas configuration).

Ionic bonds occur between metals and non-metals.

Metal Nonmetal

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IONIC BOND

Atoms that lose electrons (metals) form positive ions (cations).

Atoms that gain electrons (non-metals) form negative ions (anions).

The smallest particles of ionic compounds are ions (not atoms).

Cation

Anion

Using Lewis Theory to Predict Chemical Formulas of Ionic Compounds

Predict the formula of the compound that forms betweencalcium and chlorine.

Draw the Lewis dot symbolsof the elements

Ca∙∙ Cl ∙∙∙

∙ ∙∙ ∙

Transfer all the valance electronsfrom the metal to the nonmetal,adding more of each atom as yougo, until all electrons are lost from the metal atoms and all nonmetal atoms have 8 electrons

Ca∙∙ Cl ∙∙∙

∙ ∙∙ ∙Cl ∙∙∙

∙ ∙∙ ∙

Ca2+

:: Cl

:: Cl

CaCl2

Covalent Bonds• typical of molecular species• atoms bonded together to form molecules

– strong attraction

• sharing pairs of electrons to attain octets• molecules generally weakly attracted to

each other– observed physical properties of molecular

substance due to these attractions

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COVALENT BOND

Covalent bonds form when electrons are shared between two atoms.

Covalent bonds form between twonon-metals.

The smallest particles of covalent compounds are molecules.

Electrons shared

Single Covalent Bonds• two atoms share one pair of electrons

– 2 electrons

• one atom may have more than one single bond

F••

••

•• • F•••••••

HF••

••

•• ••

••F•••• H O

•• ••••

••

H•H• O••

• •

••

F F

Double Covalent Bond

• two atoms sharing two pairs of electrons– 4 electrons

• shorter and stronger than single bond

O••••O••

••••••

O••

• •

••O••

• •

••

O O

Triple Covalent Bond

• two atoms sharing 3 pairs of electrons– 6 electrons

• shorter and stronger than single or double bond

N••

• •

•N••

• •

•

N•••••••••• N

N N

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POLAR & NON-POLARBONDS

Two types of covalent bonds exist:

Non-polar covalent bonds occur between similar atoms.

In these bonds the electron pair is shared equally between the two protons.

Polar & Nonpolar

Electrons shared equally

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POLAR & NON-POLARBONDS

Polar covalent bonds occur between different atoms.

In these bonds the electron pair is shared unequally between the two atoms.

As a result there is a charge separation in the molecule, and partial charges on each atom.

+ H F

Dipole Moments• A dipole is a material with positively and

negatively charged ends• Polar bonds or molecules have one end

slightly positive, +; and the other slightly negative, -

– not “full” charges, come from nonsymmetrical electron distribution

• Dipole Moment, , is a measure of the size of the polarity – measured in Debyes, D

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ELECTRONEGATIVITY

Electronegativity (E.N.) is the ability of an atom involved in a covalent bond to attract the bonding electrons to itself.

Linus Pauling derived a relative Electronegativity Scale based on Bond Energies.

Cs 0.7

F 4.0

Least electronegative

Most electronegative

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ELECTRONEGATIVITY

Electronegativity increases

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BOND POLARITY &ELECTRONEGATIVITY

The more polar the

bond formed

Polarity is a measure of the inequality in the sharing of bonding electrons

The more different the

electronegativity of the elements

forming the bond

The larger the electronegativity

difference(EN)

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POLARITY &ELECTRONEGATIVITY

As difference in electronegativity

increases

Bond polarity increases

Most polar

Least polar

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POLARITY &ELECTRONEGATIVITY

Electronegativity

differenceBond Type

EN = 0 Non-polar covalent

0 < EN <1.7 Polar covalent

1.7 < EN Ionic

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H H

Hydrogen Molecule

The molecule is nonpolar covalent

Electronegativity2.1

Electronegativity2.1

POLARITY &ELECTRONEGATIVITY

EN = 0

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H Cl

Hydrogen Chloride Molecule

Electronegativity2.1

Electronegativity3.0

The molecule is polar covalent

+ -

EN = 0.9

POLARITY &ELECTRONEGATIVITY

29Sodium Chloride

Na+ Cl-

Electronegativity0.9

Electronegativity3.0

The bond is ionicNo molecule exists

EN = 2.1

POLARITY &ELECTRONEGATIVITY

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SUMMARYOF BONDING

Ionic Bond(large EN)

Covalent Bond(small to moderate EN)

Non-polar(similar electronegativities)

Polar(moderate EN)

EN > 1.7

EN = 0

0 < EN < 1.7

Bonding & Lone Pair Electrons

• Electrons that are shared by atoms are called bonding pairs

• Electrons that are not shared by atoms but belong to a particular atom are called lone pairs– also known as nonbonding pairs

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LEWIS STRUCTURES

In a Lewis structure, a shared electron pair is indicated by two dots between the atoms, or by a dash connecting them.

Unshared pairs of valence electrons (called lone pairs) are shown as belonging to individual atoms or ions.

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LEWIS STRUCTURES

Covalent molecules are best represented with electron-dot or Lewis structures.

Structures must satisfy octet rule (8 electrons around each atom).

Hydrogen is one of the few exceptions and forms a doublet (2 electrons).

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LEWISSTRUCTURES

Bonding electrons can be displayed by a dashed line.

Non-bonding electrons must be displayed as dots.

Polyatomic Ions

• The polyatomic ions are attracted to opposite ions by ionic bonds– Form crystal lattices

• Atoms in the polyatomic ion are held together by covalent bonds

Lewis Formulas of Molecules

• shows pattern of valence electron distribution in the molecule

• useful for understanding the bonding in many compounds

• allows us to predict shapes of molecules

• allows us to predict properties of molecules and how they will interact together

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LEWISSTRUCTURES

More complex Lewis structures can be drawn by following a stepwise method:

1. Count the number of electrons in the structure.

2. Draw a skeleton structure.- most metallic element generally central- halogens and hydrogen are generally terminal- many molecules tend to be symmetrical- in oxyacids, the acid hydrogens are attached

to an oxygen

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LEWISSTRUCTURES

More complex Lewis structures can be drawn by following a stepwise method:

3. Connect atoms by bonds (dashes or dots).

4. Distribute electrons to achieve Octet rule.

5. Form multiple bonds if necessary.

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Example 1:

Write Lewis structure for H2O

H2O = 8 electrons 2 (1) + 6 = 8Step 1:

Step 2:

H O HSkeleton structure should be

symmetrical

Step 3:

4 electrons used4 electrons remainingStep 4:

Octet rule is satisfied

Hydrogen has doublet

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Example 2:

Write Lewis structure for CO2

CO2 = 16 electrons 4 + 2(6) = 16Step 1:

Step 2:

O C OSkeleton structure should be

symmetrical

Step 3:

4 electrons used12 electrons remaining

Step 4:

Octet rule is satisfied

10 electrons used6 electrons remaining

Octet rule is NOT satisfied

Step 5:

Writing Lewis Structures forPolyatomic Ions

• the procedure is the same, the only difference is in counting the valence electrons

• for polyatomic cations, take away one electron from the total for each positive charge

• for polyatomic anions, add one electron to the total for each negative charge

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Example 3:

Write Lewis structure for CO32-

CO32- = 24 electrons 4+3(6)+2 = 24Step 1:

Step 2:

O C O

O

Step 3:

Step 4:

18 electrons remaining

12 electrons remaining6 electrons remaining0 electrons remaining

Octet rule is satisfiedOctet rule is NOT satisfied

Step 5:

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Example 4:

Determine if each of the following Lewis structures are correct or incorrect. If incorrect, rewrite the correct structure.

2(5) + 4(1) = 14

Structure is incorrect

Only 12 electrons shown

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224

Structure has 14 electrons

Octets are complete

Exceptions to the Octet Rule• H & Li, lose one electron to form cation

– Li now has electron configuration like He – H can also share or gain one electron to have

configuration like He

• Be shares 2 electrons to form two single bonds• B shares 3 electrons to form three single bonds• expanded octets for elements in Period 3 or

below

– using empty valence d orbitals

• some molecules have odd numbers of electrons– NO

:: ON

Some molecules, such as SF6 and PCl5 have more than 8 electrons around a central atom in their Lewis structure.

SF6 and PCl5 can violate the octet rule through the use of empty d orbitals:

both S and P can utilize empty d orbitals to hold pairs of electrons that help

bond halogen atoms.

Resonance

• we can often draw more than one valid Lewis structure for a molecule or ion

• in other words, no one Lewis structure can adequately describe the actual structure of the molecule

• the actual molecule will have some characteristics of all the valid Lewis structures we can draw

Resonance• Lewis structures often do not accurately represent

the electron distribution in a molecule– Lewis structures imply that O3 has a single (147 pm) and

double (121 pm) bond, but actual bond length is between, (128 pm)

• Real molecule is a hybrid of all possible Lewis structures

• Resonance stabilizes the molecule– maximum stabilization comes when resonance forms

contribute equally to the hybrid

OO

+

O OO

+

O

Resonance• we can often draw more than one valid

Lewis structure for a molecule or ion

• Real molecule is a hybrid

of all possible Lewis structuresO N

O

O·· ··

········

··

··

O N

O

O

·· ····

····

······

The three oxygens are chemically equivalent, so it makes nodifference to the ion which oxygen assumes the double bond.

represents resonance structures

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MOLECULARSHAPES

The three-dimensional shape of the molecules is an important feature in understanding their properties and interactions.

All binary molecules have a linear shape since they only contain two atoms.

More complex molecules can have various shapes (linear, bent, etc.) and need to be predicted based on their Lewis structures.

A very simple model , VSEPR (Valence Shell Electron Pair Repulsion) Theory, has been developed by chemists to predict the shape of large molecules based on their Lewis structures.

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MOLECULARSHAPES

Based on VSEPR, the electron pair groups in a molecule will repel one another and seek to minimize their repulsion by arranging themselves around the central atom as far apart as possible.

Electron pair groups can be defined as any one of the following:

bonding pairs non-bonding pairs

multiple bonds

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SUMMARY OFVSEPR SHAPES

Number of electron pair groups around central atom

MolecularShape

BondAngle

ExamplesBonding Non-bonding

2 0 Linear 180 CO2

3 0 Trigonal planar 120 BF3

2 1 Bent 120 SO2

4 0 Tetrahedral 109.5 CH4

3 1 Pyramidal 109.5 NH3

2 2 Bent 109.5 H2O

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MOLECULARSHAPES

Molecules with 2 electron pair groups around the central atom form a linear shape.

2 electron pairs around the

central atom

Shape is linear

Linear molecules have polar bonds, but are usually non-polar.

Bond angle is 180

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MOLECULARSHAPES

Molecules with 3 electron pair groups around the central atom form a trigonal planar shape.

Trigonal planar molecules have polar bonds, but are usually non-polar.

Bond angle is 1203 electron pairs

around the central atom

Shape is trigonal planar

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MOLECULARSHAPES

Molecules with 2 bonding pairs and 1 non-bonding pair groups around the central atom form a bent shape.

Bent molecules have polar bonds, and are polar.

Shape is bent

2 bonding pairs around the

central atom

1 Non-bonding pair

Bond angle is 120

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MOLECULARSHAPES

Molecules with 4 electron pairs groups around the central atom form a tetrahedral shape.

Tetrahedral molecules have polar bonds, and are usually non-polar.

4 bonding pairs around the

central atomShape is

tetrahedral Bond angle is 109.5

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MOLECULARSHAPES

Molecules with 3 bonding pairs and 1 non-bonding pair groups around the central atom form a pyramidal shape.

Pyramidal molecules have polar bonds, and are polar.

Shape is pyramidal

3 bonding pairs around the

central atom

1 Non-bonding pair

Bond angle is 109.5

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MOLECULARSHAPES

Molecules with 2 bonding pairs and 2 non-bonding pair groups around the central atom form a bent shape.

Bent molecules have polar bonds, and are polar.

Shape is bent

2 bonding pairs around the

central atom

2 Non-bonding pair

Bond angle is 109.5

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SUMMARY OFMOLECULAR SHAPES

Linear

Trigonal planar

Tetrahedral

Bent

Pyramidal

Symmetrical shapesPolar bonds

Non-polar molecules

Unsymmetrical shapes

Polar bondsPolar molecules

Polarity of Molecules

• For a molecule to be polar it must

1) have polar bonds, symmetrical shape, and

different terminal atoms

2) have polar bonds• electronegativity difference - theory• bond dipole moments – measured

3) have an unsymmetrical shape• using vector addition

• polarity effects the intermolecular forces of attraction

:: OCO

polar bonds,but nonpolar moleculebecause pulls cancel

OH H

polar bonds,and unsymmetrical

shape causes moleculeto be polar

Dipole moment is the measured polarity of a polar covalent bond. It is defined as the magnitude of charge (electrons) on the atoms and the distance between the two bonded atoms.

CH2Cl2

= 2.0 D CCl4

= 0.0 D

C

Cl

ClCl

Cl

C

Cl

ClH

H

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Adding Dipole Moments

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COMPARING PROPERTIESOF IONIC & COVALENT

COMPOUNDS

Ionic Covalent

Structural Unit Ions Molecules

Melting Point High Low

Boiling Point High Low

Solubility in H2O High Low or None

Electrical Cond. High None

Examples NaCl, AgBr H2, H2O

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THE END