Week 3 bonding

Prepared by:Mrs Faraziehan Senusi

PA-A11-7C

David P. White Prentice Hall ©

2003

Quantum Theory

Atomic Orbitals

Electronic Configuration

Chapter 1Atoms, Molecules & Chemical bonding

Molecular Orbitals

Bonding and Intermolecular Compounds

Introduction

Lesson Plan

At the end of this topic, the students will be able:

To describe atomic orbitals. To write electronic configurations To explain the bonding between different atoms To explain the interactions between molecules


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• Chemical bonding refers to the attractive forces that hold atoms together in compounds.

• Ionic bond results from the transfer of electrons from a metal to a nonmetal. It results from electrostatic interactions among ions. Compounds containing predominantly ionic bonding are called ionic compounds.

• Covalent bond results from sharing electrons between the atoms. Usually found between nonmetals. Those that are held together mainly by covalent bonds are called covalent compounds.

• Metallic bond: attractive force holding pure metals together.


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Chemical Bonding


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Barbara A. Gage PGCC CHM 1010

• We can classify bonds based on the kinds of atoms that are bonded together

Types of Bonds

Types of Atoms Type of BondBond

Characteristic

metals to nonmetals

Ionicelectronstransferred

nonmetals tononmetals

Covalentelectrons shared

metals tometals

Metallicelectronspooled

Barbara A. Gage PGCC CHM 1010

Types of Bonding

• All noble gases except He has an s2p6 configuration. • Atoms with one or two valence electrons more than a

closed shell are highly reactive because the extra electrons are easily removed to form positive ions.

• Atoms with one or two valence electrons fewer than a closed shell are also highly reactive because of a tendency either to gain the missing electrons and form negative ions, or to share electrons and form covalent bonds.

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


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The Octet Rule

• Between atoms of metals and nonmetals with very different electronegativity

• Bond formed by transfer of electrons• Produce charged ions all states. Conductors and

have high melting point.• Examples; NaCl, CaCl2, K2O


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Ionic Bonding

Consider the reaction between sodium and chlorine:2Na(s) + Cl2(g) 2NaCl(s) DHºf = -410.9 kJ

Ionic Bonding


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Ionic Bonding

• The reaction is violently exothermic.• We infer that the NaCl is more stable than its

constituent elements. Why?• Na has lost an electron to become Na+ and

chlorine has gained the electron to become Cl-. Note: Na+ has an Ne electron configuration and Cl- has an Ar configuration.

• That is, both Na+ and Cl- have an octet of electrons surrounding the central ion.

Ionic Bonding


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• NaCl forms a very regular structure in which each Na+ ion is surrounded by 6 Cl- ions.

• Similarly, each Cl- ion is surrounded by six Na+ ions.

• There is a regular arrangement of Na+ and Cl- in 3D.• Note that the ions are packed as closely as possible.• Note that it is not easy to find a molecular formula

to describe the ionic lattice.

Ionic Bonding


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Ionic Bonding


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Ionic Bonding

Ionic Bonding

Energetics of Ionic Bond Formation• Lattice energy: the energy required to completely

separate an ionic solid into its gaseous ions.• Lattice energy depends on the charges on the ions

and the sizes of the ions:

k is a constant (8.99 x 10 9 J·m/C2), Q1 and Q2 are the charges on the ions, and d is the distance between ions.

dQQ

El21


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Ionic Bonding

Energetics of Ionic Bond Formation• Lattice energy increases as

• The charges on the ions increase• The distance between the ions decreases.


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The energy change associated with the loss of one mole of electrons by one mole of Na atoms to form one mole of Na+ ions= 496 kJ/mol

The energy change for the gain of one mole of electrons by one mole of Cl atoms to form one mole of Cl– ions is given by the electron affinity of Cl= –349kJ/mol

The strong attractive force between ions of opposite charge draws the ions together and the energy associated with this attraction is the crystal lattice energy of NaCl, = – 789 kJ/mol.

COVALENT BONDbond formed by the sharing of electrons

• Between nonmetallic elements of more or less similar electronegativity.

• Formed by sharing electron pairs• Stable non-ionizing particles, they are not

conductors at any state• Examples; O2, CO2, C2H6, H2O, SiC

Covalent Bond

Bond Polarity and Electronegativity

• In a covalent bond, electrons are shared.• Sharing of electrons to form a covalent bond does

not imply equal sharing of those electrons.• There are some covalent bonds in which the

electrons are located closer to one atom than the other.

• Unequal sharing of electrons results in polar bonds.

when electrons are shared equally

NONPOLAR COVALENT BONDS

H2 or Cl2

Covalent bonds- Two atoms share one or more pairs of outer-shell electrons.

Oxygen Atom Oxygen Atom

Oxygen Molecule (O2)

http://www.usfca.edu/fac-staff/Courses/BIOL104_USF/104_Fall03_ppt/Text%20Chapter%2002/OxgnMol.swf



when electrons are shared but shared unequally

POLAR COVALENT BONDS

H2O

- water is a polar molecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

Electronegativity• Electronegativity: The ability of one atoms in a

molecule to attract electrons to itself.• Pauling set electronegativities on a scale from 0.7

(Cs) to 4.0 (F).• Electronegativity increase from left to right across

a period and decreases with increasing atomic number within each group.



Electronegativity


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Electronegativity

• Difference in electronegativity is a gauge of bond polarity:electronegativity differences around 0 result in non-

polar covalent bonds (equal or almost equal sharing of electrons);

electronegativity differences around 2 result in polar covalent bonds (unequal sharing of electrons);

electronegativity differences around 3 result in ionic bonds (transfer of electrons).


• There is no sharp distinction between bonding types.

• The positive end (or pole) in a polar bond is represented + and the negative pole -.


Consider HF:• The difference in electronegativity

leads to a polar bond.• There is more electron density on F

than on H.• Since there are two different “ends”

of the molecule, we call HF a dipole.

Prentice Hall © 2003 Chapter 8


Dipole Moments



Dipole Moments

μ = q x d% = μLiH/ μ

Strengths of Covalent Bonds

• The energy required to break a covalent bond is called the bond dissociation enthalpy, D

• When more than one bond is broken:CH4(g) C(g) + 4H(g)H = 1660 kJ

the bond enthalpy is a fraction of H for the atomization reaction:

D(C-H) = ¼H = ¼(1660 kJ) = 415 kJ.• Bond enthalpies can either be positive or negative.


Bond Enthalpies and the Enthalpies of Reactions

• We can use bond enthalpies to calculate the enthalpy for a chemical reaction.

• We recognize that in any chemical reaction bonds need to be broken and then new bonds get formed.

• The enthalpy of the reaction is given by the sum of bond enthalpies for bonds broken minus the sum of bond enthalpies for bonds formed.


Bond Enthalpies and the Enthalpies of Reactions

• Mathematically, if Hrxn is the enthalpy for a reaction, then

• We illustrate the concept with the reaction between methane, CH4, and chlorine:

CH4(g) + Cl2(g) CH3Cl(g) + HCl(g) Hrxn = ?

formed bondsbroken bonds DDHrxn


CH4(g) + Cl2(g)

CH3Cl(g) + HCl(g)

Strengths of Covalent BondsBond Enthalpies and the Enthalpies of Reactions

• In this reaction one C-H bond and one Cl-Cl bond gets broken while one C-Cl bond and one H-Cl bond gets formed.

• The overall reaction is exothermic which means than the bonds formed are stronger than the bonds broken.

• The above result is consistent with Hess’s law.

kJ 104

Cl-HCl-CCl-ClH-C

DDDDHrxn

Hess' Law states that the heat evolved or absorbed in a chemical process is the same whether the process takes place in one or in several steps.


Bond Enthalpy and Bond Length• multiple bonds are shorter than single bonds.• multiple bonds are stronger than single bonds.• As the number of bonds between atoms increases,

the atoms are held closer and more tightly together.

METALLIC BONDbond found in

metals; holds metal atoms together very strongly

Metallic Bonding• Formed between atoms of metallic elements (metal &

metal)• Could be the same metal bonded together (i.e copper

wire)• Or could be a mixture of different metals held together

(alloy)• Alloy: metal made by combining two or more metallic

elements, especially to give greater strength or to prevent corrosion

• Good electricity conductor at all states, lustrous, very high melting points, malleable & ductile


A malleable substance can be rolled or pounded into sheets. A ductile substance can be drawn into wires.

Electron-Sea Model of Metallic Bonding• We use a delocalized model for electrons in a

metal.– The metal nuclei are seen to exist in a sea of electrons.– No electrons are localized between any two metal atoms.– Therefore, the electrons can flow freely through the metal.– Without any definite bonds, the metals are easy to deform (and are

malleable and ductile).

• Problems with the electron sea model:– As the number of electrons increase, the strength of bonding should

increase and the melting point should increase.

Metallic Bonding

Ionic Bond, A Sea of Electrons


Electron-Sea Model of Metallic Bonding– group 6B metals have the highest melting points (center of

the transition metals).

Metallic Bonding

Ketelaar TriangleA plot of average electronegativity against electronegativity difference can be used to classify the bond type for binary compounds.

Dif

fere

nce

in e

lect

rone

gati

vity

Average electronegativity

Ionic bonding ~ large difference in electronegativity~ high electronegativity of one element and low for the other one element

Covalent bonding ~ small difference in electronegativity~ high electronegativity of element between nonmetals

Metallic bonding ~ small difference in electronegativity~ low electronegativity of element

For example:MgOΔX=3.44 - 1.31=2.13Xmean=2.38 MgO in ionic region

Week 3 bonding

Technology

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