CPE435 Atoms, Molecules & Chem Bonding

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Chapter 1

Atoms, Molecules &Chemical Bonding

CPE 435: PROCESS CHEMISTRY

Prepared by:Rabiatul Adawiyah Abdol Aziz



Atomic Structure

• The atom consists of positive, negative, and neutral entities(protons, electrons, and neutrons).

• Protons and neutrons are located in the nucleus of theatom, which is small. Most of the mass of the atom is due

to the nucleus.• Electrons are located outside of the nucleus. Most of the

volume of the atom is due to electrons.

2

Prentice Hall © 2003



Atomic Structure (2)3

Structure of an atom Positively charged nucleus (very dense, protons and neutrons)

and small (10-15 m)

Negatively charged electrons are in a cloud (10-10

m) aroundnucleus

Diameter is about 2 10-10 m (200 picometers (pm))[the unit angstrom (Å) is 10-10 m = 100 pm]

Particle Symbol Relative electriccharge Relative mass

Proton

Neutron

Electron

p

n

E

+1

0

-1

1

1

0.0005

Thompson Higher Education © 2007



Atomic Number and Atomic Mass4

The atomic number ( Z ) is the number of protons in theatom's nucleus

The mass number ( A) is the number of protons plusneutrons

All the atoms of a given element have the same atomicnumber Isotopes are atoms of the same element that have

different numbers of neutrons and therefore different massnumbers

The atomic mass (atomic weight ) of an element is the weighted average mass in atomic mass units (amu) of anelement’s naturally occurring isotopes




Atomic Orbitals5

Quantum mechanics: describes electron energies andlocations by a wave equation

Wave function solution of wave equation Each wave function is an orbital; ψ

A plot of ψ2 describes where electron most likely to occupy

Electron cloud has no specific boundary so we show mostprobable area




Shells & Subshells

Electrons are arranged around the nucleus of anatom, in shells. These numbers are known asprincipal quantum numbers, n.

For example, the first shell is called as quantumnumber = 1.

Each shell consists of a number of subshells,labelled as s, p, d or f .



Shells & Subshells (2)

The number of subshells in each shell equals theshell number. Hence, the first shell has onesubshell , the second shell has two subshells, etc.

shell number number of subshells

1 1 (1s)

2 2 (2s, 2p)3 3 (3s, 3p, 3d)



Shells & Subshells (3)

Each subshell contains a number of orbitals, in which theelectrons are placed.

The number of orbitals in each subshell depends on thetype of subshell:

type of subshell number of orbitalss 1

p 3

d 5 f 7



Shapes of Atomic Orbitals for Electrons9

s and p orbitals most important in organic and biologicalchemistry

s orbitals: spherical, nucleus at center p orbitals: dumbbell-shaped, nucleus at middle

d orbitals: 4 of 5 (cloverleaf-shaped) & the 5th

(elongateddumbbell-shaped), nucleus at center




Orbitals and Shells10

Orbitals are grouped in shells of increasing size andenergy

Different shells contain different numbers and kinds of orbitals

Each orbital can be occupied by two electrons



Orbitals and Shells (2)11

First shell contains one s orbital, denoted 1s, holds only two electrons

Second shell contains one s orbital (2s) and three p orbitals (2 p), eight electrons

Third shell contains one s orbital (3s), three p orbitals(3 p), and five d orbitals (3d ), 18 electrons



p-Orbitals12

In each shell there are three perpendicular p orbitals, p x , py, and p z , of equal energy

Two lobes of each p orbital are separated by region of zero electron density; a node



Electron Configurations13

Ground-state electron configuration (lowest energy arrangement) of an atom lists orbitals occupied by itselectrons.

Rules:1. Lowest-energy orbitals fill first: 1s 2s 2 p 3s 3 p

4s 3d ( Aufbau (“ build-up”) principle)2. Electrons act as if they were spinning around an axis.

Electron spin can have only two orientations, up anddown . Only two electrons can occupy an orbital, andthey must be of opposite spin ( Pauli exclusion

principle) to have unique wave equations3. If two or more empty orbitals of equal energy are

available, electrons occupy each with spins parallel untilall orbitals have one electron ( Hund's rule).




Exercises:

1. Write the complete or condensed electronic configurationfor the following elements:

i. Sulphur (S) (z=16);

ii. Magnesium (Mg) (z=12);

iii. Bromine (Br) (z=35);iv. Chromium (Cr) (z=24)

2. Using the Periodic Table, give the symbol(s) of:

i. an element with the ground state electronic configuration of

[Xe]6s24f 145d106p1.ii. an ion with a double positive charge (2+) with an electronic

configuration of [Ar]3d5.

iii. two elements with a groundstate configuration of ns2np3.



Quantum Mechanics

•Orbital – Each allowed combination of n, l and ml values specifiesone of the atom’s orbitals to describe the shape, size and thespatial orientation.

• An atomic orbital is specified by 3 quantum numbers:

1. Principal Quantum Number, n, related to the size of the orbital. As n becomes larger, the atom becomes largerand the electron is further from the nucleus. The higher the value of n, the higher the energy level.

Note: the total number of orbitals = n2.




Quantum Mechanics (2)

2. Orbital Angular Momentum (or, Azimuthal) Quantum

Number, l . (related to shape). This quantum number depends on

the value of n. The number of l value= the number of n value. The

values of l begin at 0 and increase to (n - 1). We usually use letters for

l (s, p, d and f for l = 0, 1, 2, and 3, respectively). Usually we refer tothe s, p, d and f -orbitals. The number of orbitals in each subshell is

2l +1 for a given l value. One s orbital (l =0), 3 p orbitals (l =1) and 5 d

orbitals (l =2) and 7 f orbitals (l =3).

3. Magnetic Quantum Number, ml related to orientation inspace. This quantum number depends on l . The magnetic quantum

number has integral values between -l and +l . The number of possible

ml values equals the number of orbitals, which is 2l +1 for a given l

value.Prentice Hall © 2003




Quantum Mechanics (3)

How many orbitals exist for n = 3?

For n = 3, l will have 3 values, i.e., 0, 1 and 2.

For l = 0, ml will have 1 value (0)

For l = 1, ml will have 3 values (-1, 0 and +1) For l = 2, ml will have 5 values, -2 through 0 to +2.

(i.e., -2, -1, 0 ,+1 and +2).

There are 9 ml values which means 9 orbitals. Inother words, n2=32=9.




Orbitals and Quantum Numbers



Molecular Orbitals

Molecular Orbital (MO) Theory - Just as electrons in atomsare found in atomic orbitals, electrons in molecules are found inmolecular orbitals.

Additive combination (bonding) MO is lower in energy Subtractive combination (antibonding) MO is higher energy

Sigma (s) bonds - Circular cross-section and are formed by head-oninteraction




Molecular Orbitals in H2 & He2

• H2 has two bonding electrons.• He2 has two bonding electrons and two antibonding

electrons.




Molecular Orbitals in Ethylene

The bonding MO is from combining p orbital lobes with the same algebraic sign

The antibonding MO is from combining lobes withopposite signs

Only bonding MO is occupied

Pi () bonds – “dumbbell” shape from sideways interaction of p orbitals




Bond Order

• Bond order = 1 for single bond.

• Bond order = 2 for double bond.• Bond order = 3 for triple bond.

• Fractional bond orders are possible.

electronsgantibondin-electronsbondingorderBond21




Bond Order (2)

• For H2

• Therefore, H2 has a single bond.

• For He2

• Therefore He2 is not a stable molecule

022orderBond21

102orderBond21




Homonuclear diatomic molecules

• Molecules composed only one type of element• For examples: Li2, Be2, B2 etc.

• AOs combine according to the following rules:• The number of MOs = number of AOs;• AOs of similar energy combine;

• As overlap increases, the energy of the MO decreases;

• Pauli exclusion principle?

• Hund’s rule?




Molecular Orbitals for Li2



Molecular Orbitals for Li2 (2)

• Each 1s orbital combines with another 1s orbital to giveone s1s and one s*

1s orbital, both of which are occupied(since Li and Be have 1s2 electron configurations).

• Each 2s orbital combines with another 2s orbital, two give

one s2s and one s*2s orbital.• The energies of the 1s and 2s orbitals are sufficiently

different so that there is no cross-mixing of orbitals (i.e. we do not get 1s + 2s).




Molecular Orbitals for Li2 (3)

• There are a total of 6 electrons in Li2:• 2 electrons in s1s;

• 2 electrons in s*1s;

• 2 electrons in s2s; and

• 0 electrons in s*2s.

• Since the 1s AOs are completely filled, the s1s and

s*1s are filled.• Note: Generally ignore core electrons in MO

diagrams.

124orderBond21




Molecular Orbitals for Be2

• There are a total of 8 electrons in Be2:• 2 electrons in s1s;

• 2 electrons in s*1s;

• 2 electrons in s2s; and

• 2 electrons in s*2s.

• Since the bond order is zero, Be2 does not exist.

044orderBond21




Molecular Orbitals from 2p AtomicOrbitals

• The six p-orbitals (two sets of 3) must give rise to 6MOs:• s, s*, , *, , and *.

• Therefore there is a maximum of 2 bonds that can come from p-

orbitals.• The relative energies of these six orbitals can change.

• There are two ways in which two p orbitals overlap:• end-on so that the resulting MO has electron density on the axis

between nuclei (i.e. s type orbital);• sideways so that the resulting MO has electron density above and

below the axis between nuclei (i.e. type orbital).




Molecular Orbitals from 2p Atomic Orbitals (2)




Heteronuclear diatomic molecules

• Molecules composed two different elements• For examples: CO2, H2O, etc.

• Pauli exclusion principle?

• Hund’s rule?



Chemical Bonds

• Chemical bond: attractive force holding two ormore atoms together.

• Covalent bond results from sharing electrons between the atoms. Usually found between non-metals.

• Ionic bond results from the transfer of electrons

from a metal to a non-metal.• Metallic bond: attractive force holding pure

metals together.




The Octet Rule

• All noble gases except He has an s2 p6 configuration.

• Octet rule: atoms tend to gain, lose, or shareelectrons until they are surrounded by 8 valence

electrons (4 electron pairs).• Caution: there are many exceptions to the octet

rule.


http://en.wikipedia.org/wiki/Image:Octeto.png



Exceptions to the Octet Rule

• There are three classes of exceptions to the octet rule:• Molecules with an odd number of electrons;

• Molecules in which one atom has less than an octet;

• Molecules in which one atom has more than an octet.

Odd Number of Electrons

• Few examples. Generally molecules such as ClO2, NO,

and NO2 have an odd number of electrons.

N O N O




Exceptions to the Octet Rule (2)

Less than an Octet• Relatively rare.• Molecules with less than an octet are typical for compounds of Groups

1A, 2A, and 3A.• Most typical example is BF3.

• Formal charges indicate that the Lewis structure with an incompleteoctet is more important than the ones with double bonds.

More than an Octet

• This is the largest class of exceptions.

• Atoms from the 3rd period onwards can accommodate more than anoctet.

• Beyond the third period, the d -orbitals are low enough in energy toparticipate in bonding and accept the extra electron density.




Ionic Bonding (Definition)

Involves a metal and a nonmetal ion throughelectrostatic attraction.

Bond formed by the attraction between two oppositely charged ions.

The metal donates one or more electrons, forming apositively charged ion or cation with a stable electronconfiguration. These electrons then enter the non metal,causing it to form a negatively charged ion or anion

which also has a stable electron configuration. Theelectrostatic attraction between the oppositely chargedions causes them to come together and form a bond.




Ionic Bonding (Example)

Consider the reaction between sodium and chlorine:Na(s) + ½Cl2(g) NaCl(s) D H º f = -410.9 kJ

• The reaction is violently exothermic.

• NaCl is more stable than its constituent elements.• 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 Arconfiguration.

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




Energetics of Ionic Bond Formation

• The formation of Na+(g) and Cl-(g) from Na(g) and Cl(g) isendothermic.

• Why is the formation of Na(s) exothermic?

• The reaction NaCl(s) Na+(g) + Cl-(g) is endothermic

(D H = +788 kJ/mol).• The formation of a crystal lattice from the ions in the gas

phase is exothermic:

Na+(g) + Cl-(g) NaCl(s) D H = -788 kJ/mol




Energetics of Ionic Bond Formation (2)

• Lattice energy: the energy required to completely separate an ionic solid into its gaseous ions.

• Lattice energy depends on the charges on the ions andthe sizes of the ions:

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

• Lattice energy increases asi. The charges on the ions increaseii. The distance between the ions decreases

d QQ E l 21




Electron Configurations of Ions of the



Electron Configurations of Ions of theRepresentative Elements

• These are derived from the electron configuration of elements with the required number of electrons addedor removed from the most accessible orbital.

• Electron configurations can predict stable ionformation:

• Mg: [Ne]3s2

• Mg+: [Ne]3s1 not stable

• Mg2+: [Ne] stable• Cl: [Ne]3s23 p5

• Cl-: [Ne]3s23 p6 = [Ar] stable




Covalent Bonding

• When two similar atoms bond, none of them wants to lose orgain an electron to form an octet.

• When similar atoms bond, they share pairs of electrons to eachobtain an octet.

• Each pair of shared electrons constitutes one chemical bond.• Example: H + H H2 has electrons on a line connecting the two

H nuclei.




Strengths of Covalent Bonds

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

• That is, for the Cl2 molecule, D(Cl-Cl) is given by D H forthe reaction: Cl2(g) 2Cl(g).

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

the bond enthalpy is a fraction of D H for the atomizationreaction: D(C-H) = ¼D H = ¼(1660 kJ) = 415 kJ.

• Bond enthalpies can either be positive or negative.






Strengths of Covalent Bonds (2)

Bond Enthalpies and the Enthalpies of Reactions • Bond enthalpies can be to calculate the enthalpy for a

chemical reaction.

• 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 bondenthalpies for bonds formed.

• For example;CH4(g) + Cl2(g) CH3Cl(g) + HCl(g) D H rxn = ?

Where D H rxn = the enthalpy for a reaction








Bond 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 getsformed.

• 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.

kJ104

Cl-HCl-CCl-ClH-C

D D D D D H rxn





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, theatoms are held closer and more tightly together.







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 electronsare located closer to one atom than the other.

• Unequal sharing of electrons results in polar bonds.

• 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).




Bond Polarity and Electronegativity (2)

• Electronegativity increases• across a period and

• down a group.• 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 polarcovalent bonds (unequal sharing of electrons);

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




Electronegativity




• There is no sharp distinction between bonding types.• The positive end (or pole) in a polar bond is represented +

and the negative pole -.





Dipole Moments• 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.• Dipole moment, m, is the magnitude of the dipole:

where Q is the magnitude of the charges.• Dipole moments are measured in debyes, D.

Qr




Metallic Bonding

• Important physical properties of pure metals: malleable,ductile, good conductors, and feel cold.

• Most metals are solids with the atoms in a close packedarrangement.

• In Cu, each atom is surrounded by 12 neighbors.• There are not enough electrons for the metal atoms to be

covalently bonded to each other.




Electron-Sea Model of Metallic Bonding

• Delocalized model is used 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 shouldincrease.

– But: group 6B metals have the highest melting points(center of the transition metals).




Prentice Hall © 2003Chapter 23






Molecular-Orbital Model for Metals

• Delocalized bonding requires the atomic orbitals on oneatom to interact with atomic orbitals on neighboringatoms.

• Example: graphite electrons are delocalized over a whole

plane, benzene molecules have electrons delocalized over aring.

• Recall: the number of molecular orbitals is equal to thenumber of atomic orbitals.

• In metals there is a very large number of orbitals.




Molecular-Orbital Model for Metals (2)

• As the number of orbitals increase, their energy spacingdecreases and they band together.

• The number of electrons do not completely fill the band of orbitals.

• Therefore, electrons can be promoted to unoccupiedenergy bands.

• Since the energy differences between orbitals are small, thepromotion of electrons occurs at low energy costs.








• As we move across the transition metal series, theantibonding band starts becoming filled.

• Therefore, the first half of the transition metal series haveonly bonding-bonding interactions, the second half has

bonding-antibonding interactions.• We expect the middle of the transition metal series to have

the highest melting points.

• The energy gap between bands is called the band gap.




Ketelaar Triangle

triangles used to show different compounds in varying degrees of ionic,metallic and covalent bonding.

On the right side (from ionic to covalent) should be compounds with varying difference in electronegativity, in the covalent cornercompounds with equal electronegativity such as Cl2, in the ionic corner

compounds with large electronegativity difference such as NaCl The bottom side (from metallic to covalent) is metallic bonds

with delocalized bonding and the other are covalent bonds in which theorbitals overlap in a particular direction.

The left side (from ionic to metallic) is for delocalized bonds with

varying electronegativity difference.

l l ( )



Ketelaar Triangle (2)


Intermolec lar Forces



Intermolecular Forces

• The covalent bond holding a molecule together is anintramolecular forces.

• The attraction between molecules is an intermolecular force (also called Van der Waals forces).

• Intermolecular forces are much weaker thanintramolecular forces (e.g. 16 kJ/mol vs. 431 kJ/mol forHCl).

• When a substance melts or boils the intermolecular forces

are broken (not the covalent bonds).• There are several types : ion-dipole, dipole-dipole,

hydrogen bonding and London dispersion forces.

I t l l F



Intermolecular Forces


I Di l F



Ion-Dipole Forces


• Interaction between an ion and a dipole.• For example: ionic compound dissolve in water. The

ions become separeted because the attaractions

between the ions and the oppositely charged polesof H2O molcules overcome the attraction betweenthe ions themselves.

• Strongest forces among all intermolecular forces.

I Di l F ( )



Ion-Dipole Forces (2)

--- ion-depole forces



Di l Di l F ( )



Dipole-Dipole Forces (2)

H d B di



Hydrogen Bonding


• Special case of dipole-dipole forces.• By experiments: boiling points of compounds with

H-F, H-O, and H-N bonds are abnormally high.

• Intermolecular forces are abnormally strong.• H-bonding requires H bonded to an electronegative

element (most important for compounds of F, O,and N).

– Electrons in the H-X (X = electronegative element) lie much closer to X thanH.

– H has only one electron, so in the H-X bond, the + H presents an almost bare proton to the - X.

– Therefore, H-bonds are strong.

Hydrogen Bonding



Hydrogen Bonding

E i



Exercise

Which of the folloing substances exhibits H bonding?For those that do, draw to molecules of the substance

with the H bond(s) beteen them.

i. C2

H6ii. CH3OH

iii. CH3C=O

NH2

London Dispersion Forces



London Dispersion Forces


• Weakest of all intermolecular forces.• Very weak for small particles (like H2 and He) but stronger

for larger particles (like I2 and Xe)

• It is possible for two adjacent neutral molecules to affect

each other.• The nucleus of one molecule (or atom) attracts the electrons

of the adjacent molecule (or atom).

• For an instant, the electron clouds become distorted.

• In that instant a dipole is formed (called an instantaneousdipole).

• The forces between instantaneous dipoles are calledLondon dispersion forces.

London Dispersion Forces (2)




• Polarizability is the ease with which an electron cloud can bedeformed.

• The larger the molecule (the greater the number of electrons)the more polarizable - increase as molecular weight increases

• London dispersion forces exist between all molecules -depend on the shape of the molecule.

• The strenght of the dispersion force depends on thepolarizability of the particle.


L d Di i F (3)





• The greater the surface areaavailable for contact, the greaterthe dispersion forces.

• London dispersion forces

between spherical molecules arelower than between sausage-likemolecules.

Increasing the strenghtof dispersion forces

References



References

1. David P. White (2003) Chemistry - The Central Science (9th edition). Prentice Hall.

2. John McMurry (2008) Organic Chemistry.Brooks/Cole Cengage Learning.

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CPE435 Atoms, Molecules & Chem Bonding