Chemical Bonding

Course: B.Sc. Microbiology/Biotechnology/Biochemistry

Sem IISub: Inorganic Chemistry

Unit 3.1

HYDROGEN BONDING

• A hydrogen atom : one electron - covalently bonded to only oneatom.

• hydrogen atom can involve itself in an additional electrostaticbond with a second atom of highly electronegative charactersuch as fluorine or oxygen. This second bond permits ahydrogen bond between two atoms or strucures.

• The strength of hydrogen bonding varies.

Hydrogen bonds connect water molecules in ordinary ice. Hydrogen bonding is also very important in proteins and nucleic acids and therefore in life processes.

Types of Hydrogen bond

• occur within one single molecule, between two like molecules, or between two unlike molecules.

• 1. Intramolecular hydrogen bonds:

- occur within one single molecule. - occurs when two functional groups of a molecule can form hydrogen bonds with

each other. - both a hydrogen donor and acceptor must be present within one molecule

For example : ethylene glycol (C2H4(OH)2) between its two hydroxyl groups

• 2. Intermolecular hydrogen bonds• occur between separate molecules in a substance. • occur between any number of like or unlike molecules as long as hydrogen donors

and acceptors are present an in positions in which they can interact.• For example, between NH3 molecules alone, between H2O molecules alone, or

between NH3 and H2O molecules.

Properties & effects of Hydrogen bond•

1. Boiling points of molecules : molecules with larger molar masses to have higher normal boiling points than molecules with smaller molar masses.

This, without taking hydrogen bonds into account, is due to greater dispersion forces. Larger molecules have more space for electron distribution and thus more possibilities for an instantanous dipole moment.

H2O, HF, and NH3 : each have higher boiling points than the same compound formed between hydrogen.

• This is because H2O, HF, and NH3 all exhibit hydrogen bonding, whereas the others do not.

•

• Furthermore, H2O has a smaller molar mass than HF but partakes in more hydrogen bonds per molecule, so its boiling point is consequently higher.

2. Viscosity

Those subtances which are capable of forming hydrogen bonds tend to have a higher viscosity than those that do not. Substances which have the possibility for multiple hydrogen bonds exhibit even higher viscosities.

Factors preventing Hydrogen bonding

• 1. Electronegativity• cannot occur without significant electronegativity differences between hydrogen and

the atom it is bonded to. • Thus, we see molecules such as PH3, which no not partake in hydrogen bonding. PH3

exhibits a trigonal pyramidal molecular geometry like that of ammmonia, but unlike NH3 it cannot hydrogen bond. This is due to the similarity in the electronegativities of phosphorous and hydrogen. Both atoms have an electronegativity of 2.1, and thus, no dipole moment occurs. This prevents the hydrogen bonding from acquiring the partial positive charge needed to hydrogen bond with the lone electron pair in another molecule.

• 2. Atom Size• The size of donors and acceptors can also effect the ability to hydrogen bond. • low ability of Cl to form hydrogen bonds. • When the radii of two atoms differ greatly or are large, their nuclei cannot achieve close

proximity when they interact, resulting in a weak interaction.

4 - VAN DER WAALS

BONDING

• weak bond.

• The intermolecular attractive forces operative between allmolecules, when they are close to one another –van dar waalsforces.

• Natural fluctuation in the electron density of all molecules andthese cause small temporary dipoles within the molecules.

• It is these temporary dipoles that attract one molecule toanother. They are called van der Waals' forces.

• The bigger a molecule is, the easier it is to polarise (to form adipole), and so the van der Waal's forces get stronger, so biggermolecules exist as liquids or solids rather than gases.

• Water molecules- attracted to each other by electrostatic forces.

• Even though the water molecule as a whole is electrically neutral, the distribution of charge in the molecule is not symmetrical and leads to a dipole moment - a microscopic separation of the positive and negative charge centers. This leads to a net attraction between such polar molecules.

• shape : influences its ability to form temporary dipoles.

• Long thin molecules can pack closer to each other thanmolecules that are more spherical.

• bigger the 'surface area' of a molecule : the greater the vander Waal's forces will be and the higher the melting andboiling points of the compound will be.

Homonuclear molecules,such as iodine, develop

temporary dipoles due tonatural fluctuations of electron

density within the molecule

Heteronuclear molecules,such as H-Cl have permanent

dipoles that attract the oppositepole in other molecules.

The dipoles can be formed as a result of unbalanced distribution

of electrons in asymettrical molecules. This is caused by the

instantaneous location of a few more electrons on one side of the

nucleus than on the other.

symmetric asymmetric

Therefore atoms or molecules containing dipoles are attracted to each other by electrostatic forces.

Ionic solids

• composed of oppositely charged ions.

• consist of positively charged cations and negatively charged anions .

• dissolved in water : seperate,

• free to move about in the water allowing the solution to conduct electrical current.

Properties of Ionic Compounds

– Have high melting and boiling temperatures.

– Are hard but brittle

• They also:

– Do NOT conduct electricity in the solid state

– They will only conduct electricity if they are melted or dissolved in water

• The physical properties of ionic compounds are very different from metals.

Lattice energy

• type of potential energy.

• Lattice energy is the energy required to break apart an ionic solid and convert its component atoms into gaseous ions.

• Value for the lattice energy – positive : an endothermic reaction.’

• “lattice energy is the reverse process, meaning it is the energy

released when gaseous ions bind to form an ionic solid.

• this process will always be exothermic, and thus the value for lattice energy will be negative.

• units - kJ/mol.”

Born-Haber Cycle

• A series of hypothetical steps and their enthalpy changes needed to convert elements to an ionic compound and devised to calculate the lattice energy.

• Using Hess’s law.

Born-Haber Cycle Steps

1. Elements (standard state) - gaseous atoms

2. form cations and anions

3. Combining gaseous anions and cations to form a solid ionic compound

Step 1: Atomisation

• The standard enthalpy change of atomisation is the ΔH required to produce one mole of gaseous atoms.

• Na(s) Na(g) ΔHoat = +109 kJmol-1

• NOTE: for diatomic gaseous elements, Cl2, ΔHoat is

equal to half the bond energy (enthalpy).

• Cl2(g) Cl(g) ΔHoat = ½ E (Cl-Cl)

ΔHoat = ½ (+242 )

ΔHoat = +121 kJmol-1

Step 2: Formation of gaseous ions

• Electron Affinity

– Enthalpy change when one mole of gaseous atoms or anions gains electrons to form a mole of negatively charged gaseous ions.

• Cl(g) + e- Cl-(g) ΔHo = -364 kJmol-1

– For most atoms = exothermic, but gaining a 2nd electron is endothermic due to the repulsion between the anion and the electron.

Becoming cations

• Ionisation energy

– Enthalpy change for one mole of a gaseous element or cation to lose electrons to form a mole of positively charged gaseous ions

• Na(g) Na+(g) + e- IE1= +494 kJmol-1

Lattice Enthalpy

• Energy required to convert one mole of the solid compound into gaseous ions.

• NaCl (s) Na+(g) + Cl-(g)

• ΔHolat = +771kJmol-1

• highly endothermic

cannot directly calculate ΔHolat , but values are

obtained indirectly through Hess’s law for the formation of the ionic compound

References

• Essentials of Physical chemistry by Bahl Arun, S Chand, 2012

• General chemistry by Ebbing Darrell D, 5th, A I T B S Publishers,

2002