[Ranadeep]

Hello . Im Ranadeep from India . Im in 11th standard now ..please help me in chemistry ...i find this forum very nice : )

I often confuse with the bond strength of various types of single bonds between atoms in a molecule
I know strength of covalent bond depends on the Extent of overlapping of orbitals and also the energy of overlapping orbitals
but sometimes in organic chemistry I often confuse which bonds break first ?
Most Polar (because more electro negativity difference leading to more ionic character ) or Least polar ( due to more polarisibility of anion )

If your Answer is Most Polar then how do u account for the bond enrgey of
HI < HBr < HCl < HF

and also i heard somewhere H-H bond is strongest covalent bond is this true ?
but i find HF Bond energy more than that in solomons book

and also whats this exception Bond Energy case ?
I-I < F-F < Br-Br<Cl-Cl ! how is this order ? please explain

relative overlapping of various orbitals if its possible please post a table ..SP3-Sp3 sigma covalent bond should be the strongest bond isnt it ? is it really ?

please Solve all of these asap

[cliverlong]

Why not start with some actual bond energies so you can make some comparisons?

http://www.wiredchemist.com/chemistry/data/bond_energies_lengths.html

(There may be more extensive tables elsewhere)

Now look at the values, look at what you have been asked to compare and see what are the trends and possible explanations then see if there are any anomalies

For example in comparing the bond energies of HI, HBr, HCl, HF what is changing between each example? What do you think or know could or might affect bond strength (starting from a simple model of attractions between electrons and protons within and between atoms)

Remember bond energies are averages taken across a wide range of compounds that contain those bonds.

Clive

[renge ishyo]

Generally speaking (outside of water) the more polar the bonds the harder it is to break the bonds (inside water, the logic is opposite to this and the polar bonds are easier to break than the non-polar bonds due to the ability of polar bonds to interact with water, which is where a lot of the confusion about bond strength arises).

In covalent compounds, the more the partial ionic character the stronger the bonds. For example, H₂ and Cl₂ are non polar covalent compounds . When they react (again NO water in this comparison, this is in the gas phase) to form 2HCl, the product HCl is about 80% covalent in character and 20% ionic in character. So this extra percent ionic character means the bonds in the product are stronger than the reactants by about 44 kcal/mol for the above reaction.

On the other hand, if you react H₂ and I₂ to form 2HI, the HI bond has very little partical ionic character (since the electronegativity of hydrogen and iodine are so close to one another) and as a result the products are only slighlty more polar than the reactants and far less energy is liberated in forming the products (4 kcal/mol for the above reaction). Since more energy is liberated forming HCl, you can see that the bonds in HCl are stronger than those in HI (again, outside of water) so HI < HCl.

You don't have to take my word for it, here is the explanation from the great Linus Pauling (I took some numbers from this lecture in my post, but its better to watch the real thing:) http://osulibrary.oregonstate.edu/specialcollections/coll/pauling/bond/video/1957v.1-17.html

Now if you are comparing the bond energies between compounds of a similar bond type, typically smaller "harder" molecules have stronger primary bonds with one another than larger "soft" molecules. For example, I₂ has a bond energy of 151 kj/mol while Cl₂ has a bond energy of 242 kj/mol. These are both non polar covalent compounds but the chlorine bond is closer to the nucleus of each of the two chlorine atoms than the iodine bond is and this makes the chlorine bond tighter and harder to break. In general this holds. Fluorine just happens to be a weirdo because it is so darn small. In making the bond for F₂, the electrons are forced into so small a place that they tend to repel each other weakening the bond. You see this behavior quantitatively by looking at the electron affinity for the various halogens (see: http://www.chemguide.co.uk/atoms/properties/eas.html) where Fluorine clearly falls out of line with the other three halogens (this weirdness carries over into the covalent bond strengths).

In general though, I₂<Br₂< Cl₂ when it comes to primary bonding. With secondary bonding though the order of bond strength is again opposite! And I₂ forms a solid at standard conditions, Br₂ a liquid, and Cl₂ a gas due to secondary bonding interactions such as the london forces.

As to the final question, generally bonds between two s orbitals are stronger than the bond between two p orbitals due to the "degree of overlap" between the two types of orbitals (see your text for examples of this, it is hard to draw out online but the p bonds have farther to "reach" to touch each other in comparison to the two s orbitals). Taken that fact, if you were asked which bonds were stronger between an sp2 bond and an sp3 bond you would answer sp2. The reason is that the sp2 bond has 33% s character whereas the sp3 bond has 25% s character. The more s character, the more overlap, and the stronger the bond. You apply this third concept mostly in organic chemistry when predicting bond strengths for carbon compounds.

[Ranadeep]

Thanks alot for the Reply .. its Very informative .. : )
your reason for the bond strength order  F₂<Br₂<Cl₂ is really nice : )
But i heard relative overlapping of SP₃ is greater than Sp₂ in peter sykes i guess .. so SP_{3bond should be greater than SP2}

But your reason of S Character thingy is logical and satisfactory : )

and also How do we calculate or compare the Electronegativity of various orbitals ? S orbital is most electronegative ..is that arbitrarily fixed ? like the Hydrogen in measuring the EN of various atoms

Thanks : )

[renge ishyo]

Electronegativity does not depend on the nature of the orbital, but rather the chemical species involved in the bond. Its usefullness is that it can be used to correct bond strengths to account for partial ionic character. There are a few different ways of calculating the electronegativities, but either way you end up with a table of vales that you can use. The Pauling electronegativity values for the various elements can be seen here:

http://www.tutor-homework.com/Chemistry_Help/electronegativity_table/electronegativity.html

[Ranadeep]

soo .. saying S Orbital is most electronegative is meaningless ?

[renge ishyo]

Right. But saying Fluorine is the most electronegative atom has a meaning.