Basics

the first person to identify diatomic molecules such as $\ce{H2, N2, O2, Cl2}$ was Cannizzaro. all of these species are homonuclear diatomics which means they are composed to the same two atoms. this contrasts heteronuclear diatomics which are composed of 2 different atoms

atoms very close together repell and have a high PE

atoms bonded together have the lowest PE and so is the most favourable.

at the atoms get further apart, they interact less and so the PE increases as they arn’t bonded

• R is bond length and is the mean distance between the centres of 2 atoms this will usually be in the range of 100-200 pm = 1-2Å
• D is the bond dissociation energy this is the standard enthalpy change for the the bond to be broken range is around $100-500 \text{kJ mol}^{-1}$

non-covalent interactions are any atomic attractions other than covalent ,ionic or metallic bonding. intermolecular interactions are non covalent interactions between 2 or more molecules. intramolecular interactions are non-covalent interactions within a single molecule. van der Waals interactions should not be used as its too vague

What isn’t a chemical bond:

Permanent Dipoles

permanent separation of partial charge

• partially negative = red
• partially positive = blue

those with dipole moments have dipole- dipole interactions

in solids, the dipoles will be aligned with the opposite charges being next to each other

London dispersion

this is also called instantaneous induced dipole dipole interactions

these are always present in atoms and molecules

Models

Lewis model

Octet rule

in the lewis model, a covalent bond is a type of bonding associated with the sharing of 2 electrons usually between two atomic centres of a molecular entity

the Lewis model allowed for the bond order which is the measure of number of bonds between atoms, for example $\ce{Cl2}$ would have a bond order of 1, $\ce{O2}$ would have a bond order of 2, and $\ce{N2}$ would have a bond order of 3

lewis structures are represented by a classic dot and cross diagram

Problems

• cant explain stable radicals without filled octets such as NO

• fails to explain some transition metal compounds

• doesn’t account for expanded octets (hypervalent compounds) like in $\ce{PF5}$

• inaccurate for some paramagnetic molecules such as liquid water

  • paramagnetism

• it assumes pairs of electrons but they are not always in pairs

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paramagnetic compounds are weakly attracted to magnets, diatomic compounds are weakly repelled from magnets

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Pauling

← pauling found that heteronucleic diatomic bond enthalpies are much greater than homonucleic bonds and attributed this to the ionic contribution generated by differences in electronegativities

he came up with the Pauling scale, the measure of electronegativity we use

he noted that if you had 2 different diatomic molecules (say $\ce{X2 and Y2}$) then the bond enthalpy of the product fo these atoms would not fall in between the bond enthalpyes if the diatomics. if the bond enthalpy of $\ce{X2}= +100\text{kJ mol}^{-1}$ and $\ce{Y2}= +200\text{kJ mol}^{-1}$ the bond enthapy of $\ce{XY}\neq +150\text{kJ mol}^{-1}$ instead it would be greater due to the ionic contribution