Lecture V.Chemical Bondingjude.edu.sy/assets/uploads/lectures/Lecture V.Chemical Bonding.pdf ·...

Chemical Bonding

What is Chemical Bonding?

Atoms want to share electrons so that their valence shells arefull and they are chemically stable

Chemical Bonding is two or more atoms combining bysharing electrons so that a new substance isproduced that has different physical andchemical properties than its componentelements.

Chemical bonding occurs when atoms share

valence electrons

Electron Distribution in MoleculesElectron Distribution in Molecules

Gilbert. N. Lewis 1875 - 1946

• Atoms combine to achieve a more stable electronconfiguration.

•Maximum stability results when an atom isisoelectric with a nobel gas.

•Whin atoms interact to form a chemical bond, onlytheir outer regions are in contact( for this reason,when we study chemical bonding, we are concernedprimarily with the valence electrons of atoms).

•Chemists use a system of dots devised by lewisand called lewis dot symbols.

Lewis Symbols

• All noble gases except He has an s2p6 configuration.

• Octet rule: atoms tend to gain, lose, or share electronsuntil they are surrounded by 8 valence electrons (4electron pairs).

• Caution: there are many exceptions to the octet rule.

The Octet Rule

Lewis Symbols of Atoms and Ions• Also known as electron dot symbols• Use symbol of element to represent nucleus and inner

electrons• Use dots around the symbol to represent valence electrons

– put one electron on each side first, then pair• Elements in the same group have the same Lewis symbol

– Because they have the same number of valence electrons• Cations have Lewis symbols without valence electrons• Anions have Lewis symbols with 8 valence electrons

Li• Be• •B• •C• •N• •O: :F: :Ne:• •

•

• • • •

•• •• •• ••

••

Li• Li+1 :F: [:F:]-1•

•• ••

••

Writing Lewis Structures ofMolecules

• Count the total number of valence electrons from all theatoms

• Attach the atoms together with one pair of electrons– A line is often used as shorthand for a pair of electrons

that attach atoms together• Arrange the remaining electrons in pairs so that all hydrogen

atoms have 2 electrons (1 bond) and other atoms have 8electrons (combination of bonding and nonbonding)

• Occasionally atoms may violate this rule– Nonbonding pairs of electrons are also know as Lone

Pairs

Structural Formulas- LewisStructural Formulas- LewisStructuresStructures

• Valence electrons are indicated aroundthe symbol for the element

Oxygen has 6 valenceelectrons

Nitrogen has 5valence electrons

Drawing Lewis StructuresDrawing Lewis Structures• Imagine each side (top, bottom, left,

right) of the symbol of the element canhold 2 electrons for a total of 8electrons.

• Each side will hold one electron first,then will double up.

• In covalent bonding the number ofsingle electron sides (unpairedelectrons) indicates the number ofcovalent bonds the atom must haveto satisfy its octet.

• Oxygen has 6valence electrons.

• Two unpairedelectrons meansthat oxygen mustform two bonds tosatisfy its octet.

• Draw the Lewisstructure for thefollowing:– Chlorine– Phosphorus– Carbon

Lewis StructuresLewis Structures

• Atoms shareelectrons to fill theiroctets.

• A solid line indicatesa shared pair ofelectrons.

• Dots are used toindicate unsharedpairs of electrons.

Formation of a single covalent bond

Double and Triple BondsDouble and Triple Bonds• A unique characteristic of covalent

compounds is their ability to form multiplebonds between two atoms.

• Refer back to the Lewis Structures fornitrogen and oxygen.

• Nitrogen needs to share three electrons• Oxygen needs to share two electrons.

Technique for DrawingTechnique for DrawingLewis StructuresLewis Structures

1. Determine the number of valence electrons ineach atom making up the molecule

2. Add the valence electrons and divide by two3. Draw the “skeleton.” If carbon is present,

place it at the center of the molecule.4. Distribute the pairs of electrons around the

skeleton to satisfy each atoms octet.(Remember: Hydrogen only needs twoelectrons to fill its octet.)

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

PracticePractice

• Draw Lewis Structures for the followingcompounds:– Ammonia

– Ethyne- C2H2

– Carbon Dioxide

– HCN

Exceptions to the Octet RuleExceptions to the Octet Rule

• Atoms with more than an octet– SF4

• Molecules with an odd number ofelectrons– NO

– Generally short lived, unstable molecules

SO• Lewis symbol- shows the valence e-s,

represened by dots, around the symbolof the element

• Octet rule - atoms tend to lose, gain orshare e-s until they are surrounded by 8valence e-s

– Consists of full s & p sublevels– Achieve the noble gas config closest to

them in the P-table

Resonance: DelocalizedBonding

Resonance theory states that whenever amolecule or ion can be represented by two ormore plausible Lewis structures that differ onlyin the distribution of electrons, the truestructure is a composite, or hybrid, of them

Resonance structures

Molecules that Don’t Followthe Octet Rule

Molecules with an odd number of valence electronshave at least one of them unpaired and are calledfree radicals

Molecules that Don’t Followthe Octet Rule

Some molecules have incomplete octets. These areusually compounds of Be, B, and Al, generally havesome unusual bonding characteristics, and are oftenquite reactive

Molecules that Don’t Followthe Octet Rule

Some compounds have expanded valence shells,which means that the central atom has more thaneight electrons around it

Molecules that Don’t Followthe Octet Rule

An expanded valence shell may also need toaccommodate lone-pair electrons as well asbonding pairs

Odd ElectronOdd electron - if there is an odd number of

valence electrons, it is not possible togive every atom eight electrons• Let’s look at NO, nitric oxide

• It is impossible to pair all electrons as thecompound contains an ODD number ofvalence electrons

N - ON - O

Lewis Dot Symbols forRepresentative Elements

What is a Bond?

• A force that holds atoms together.

• Why?

• We will look at it in terms of energy.

• Bond energy the energy required to breaka bond.

• Why are compounds formed?

• Because it gives the system the lowestenergy.

See if you can define the following words before starting thelesson…

• Anion- negative ion• Cation-positive ion• Octet Rule- rule that states that atoms tend to gain,

lose, or share electrons so that each atom has fulloutermost energy level which is typically 8 electrons.

• Polyatomic Ion- charged group of covalently boundatoms

• Monatomic Ion- ion formed from a single atom• Molecule-neutral group of atoms united by covalent

bonds

• Unshared Pair- pair of electrons that is not involvedin bonding but instead is held exclusively by oneatom.

Principal Types of Chemical Bonds: Ionic and Covalent

• Ionic bond - a transfer of one or moreelectrons from one atom to another

• Forms attractions due to the oppositecharges of the atoms

• Covalent bond - attractive force due to thesharing of electrons between atoms

• Some bonds have characteristics of bothtypes and not easily identified as one or theother

Ionic Bonds

• metal to nonmetal, in ionic compounds

• 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 thecorrect formulas of ionic compounds

electron transfer

IONIC BONDING

Lewis Symbol

1s22s22p63s1

Ne coreimplied insymbol

Na

Lewis Symbol

1s22s22p63s1 1s22s22p63s23p5

Ne coreimplied insymbol

ClNa

Na+→

1s22s22p6 1s22s22p63s23p6

the loss or gain of electrons(dots) until

IONIC BONDING

the formation of ionic bonds is representedin terms of Lewis symbols

both species have reached an octet of electrons

ClNa Cl

Cl Cl

[Ne] 3s23p6

represents one orbital

(Pauli: 2 electrons)

ions stack together in regular crystalline structures

ionic solids typically

1. high melting andboiling points

2. brittle

3. form electrolytesolutions if theydissolve in water

electrostatic interaction

Li(s) + ½ F2(g) → LiF(s)

lattice energy

(up to few 1000 kJmol-1)

enthalpy of formation

Hess’s Law

Li+(g) + F-(g) → LiF(s)

Born-Haber Cycle

The Covalent Bond

• Covalent bonds are formed when 2 atoms share 1 ormore pairs of electrons achieving the lowestpossible energy.

• Atoms which form covalent bonds are very similar intheir tendency to gain or lose electron. In general,covalent bonds are formed between nonmetalatoms.

• Atoms may share more than one pair of electrons– Double bond –– Triple bond –– Multiple bonds are

•

• Covalent bonding between identical atoms meanselectrons are

• Covalent bonding between different atoms leads to

Covalent Bonding

• Shared electron bonds

• Due to overlap of atomic orbitals– (Valence Bond Theory)

• Allows each atom to fill valence shell withelectrons

Covalent Bonds

• Form between two or more nonmetals, inmolecular or covalent compounds

• Sharing of electrons– Equally shared---nonpolar covalent bond

– Unequally shared----polar covalent bond

• Bond order: single, double and triple

Single Covalent Bonds

Nonpolar Polar

H

F••

••

•• ••

••F•• ••

H O•• ••••

••F F

F••

••

•• • F•• ••••• H•H• O

••• •

••

Double Covalent Bonds

• two atoms sharing two pairs of electrons– 4 electrons

• shorter and stronger than single bond

O•••• O••

••••••

O O

O••

• •••

O••

• •

••

Triple Covalent Bonds

• two atoms sharing 3 pairs of electrons– 6 electrons

• shorter and stronger than single or doublebond

N•••••••••• N

N N

N••

• ••

N••

• •

•

BondsIonic Bonds

• Composed of ions thathave gained or lostelectrons to achieve afull outer shell

• Electrostatic attractiveforces

• Crystalline solids – nodiscrete molecules -formula units

• Identified by empiricalformulas

• Metal + non-metal

Covalent Bonds

• Composed of atomsthat are sharingelectrons to achieve afull outer shell

• Shared electron bonds

• Discrete molecules,forms gases, liquids,and solids

• Identified by molecularformulas

• Non-metal + non-metal

electronegativity

difference

between two atoms

involved in the bond

IONIC OR COVALENT

Covalent Bonding in Hydrogen

2. ELECTRON SHARING

1. ELECTRON FULLY TRANSFERED

IONIC BONDING

COVALENT BONDING

A + B → AB

2 Na(s) + Cl2(g) → 2 NaCl(s)

2 H2(g) + O2(g) → 2 H2O(l)

Electronegativity difference

Bond Type

Zero

Intermediate

Large

Covalent

Polar Covalent

Ionic

Co

valent C

haracter

decreases

Ion

ic Ch

aracter increases

Bonding spectrum100% covalent 100% ionic

A+ B-A B A B

Increasing ΔEN

Increasing polarity Transfer

• Bond order is the number ofshared electron pairs in a bond.

• A single bond has BO = 1, adouble bond has BO = 2, etc.

• Bond length is the distancebetween the nuclei of two atomsjoined by a covalent bond.

• Bond length depends on theparticular atoms in the bond andon the bond order.

Bond Order and Bond Length

Electronegativity

• measure of the pull an atom has on bonding electrons

• increases across period (left to right)

• decreases down group (top to bottom)

• larger difference in electronegativity, more polar thebond

– negative end toward more electronegative atom

δ+ H — F δ-

Bond Polarity

• .

0 0.4 2.0 4.0Electronegativity Difference

covalent ionic

polarnon

polar

3.0-3.0= 0.0

4.0-2.1= 1.9

3.0-0.9= 2.1

Linus Pauling

1901–1994American chemistOne of the few recipientsof two Nobel prizes. His

paper “The Nature of the Chemical Bondand the Structure of Molecules andCrystals” won him the 1954 Nobel Prize inchemistry. Won the peace prize in 1962.

ElectronegativityPauling (1901-1994)

�Electronegativity is the relative tendency of anatom in a molecule to attract a shared pair ofelectrons in a bond to itself.

�The most electronegative element is fluorine andit is given a value of 4.0.

�The higher the electronegativity value of anatom, the greater is the ability of an atom of thatelement to attract electrons to itself.

Electronegativity

Electronegativity (EN) is a measure of the ability ofan atom to attract bonding electrons to itself

The greater theelectronegativity ofan atom in a molecule,the more strongly itattracts the electronsin a covalent bond

Electronegativity

• Measure of the ability of an atom to attractshared electrons– Larger electronegativity means atom attracts more

strongly– Values 0.7 to 4.0

• Increases across period (left to right) on PeriodicTable

• Decreases down group (top to bottom) onPeriodic Table

• Larger difference in electronegativities meansmore polar bond– negative end toward more electronegative atom

Electronegativity

• Polarity is determined by difference inelectronegativity– Nonpolar covalent

– Polar covalent

– Ionic compound

•Electronegativity differences greater thanzero and less than two usually give a polarcovalent bond. (Some books use up to a 0.5difference as still being nonpolar because C-Hbonds are generally considered nonpolar.)

•Electronegativity differences of 2 or greaterare associated with ionic bonds.

Electronegativity Differenceand Bond Type

Two identical atoms have the same electronegativityand share a bonding electron pair equally. This iscalled a nonpolar covalent bond

Example: chlorine gas

All homonuclear diatomic molecules have nonpolarcovalent bonds:

H2, N2, O2, F2, Cl2, Br2, I2

Electronegativity Differenceand Bond Type

In covalent bonds between atoms with somewhatlarger electronegativity differences, electron pairs areshared unequally. This is called a polar covalentbond

Example: hydrogen chloride gas, HCl

The electrons are drawn closer to the atom of higherelectronegativity, Cl

Electronegativity Differenceand Bond Type

With still larger differences in electronegativity,electrons may be completely transferred from metalto nonmetal atoms to form ionic bonds

Example: sodium chloride, NaCl

Electronegativity Differences

Figure 11.3: Electronegativity values for selectedelements

Electronegativity