For each carbon, the 2s orbital hybridizes with one of the 2p orbitals to form two sp hybridized orbitals. The frontal lobes of these orbitals face away from each other forming a straight line.
The first bond consists of sp-sp orbital overlap between the two carbons. Another two bonds consist of s-sp orbital overlap between the sp hybridized orbitals of the carbons and the 1s orbitals of the hydrogens.
This leaves us with two p orbitals on each carbon that have a single carbon in them. This allows for the formation of two? Using the Lewis Structures , try to figure out the hybridization sp, sp 2 , sp 3 of the indicated atom and indicate the atom's shape. The carbon has no lone pairs and is bonded to three hydrogens so we just need three hybrid orbitals, aka sp 2. Don't forget to take into account all the lone pairs.
Every lone pair needs it own hybrid orbital. That makes three hybrid orbitals for lone pairs and the oxygen is bonded to one hydrogen which requires another sp 3 orbital. That makes 4 orbitals, aka sp 3. The carbon is bonded to two other atoms, that means it needs two hybrid orbitals, aka sp.
An easy way to figure out what hybridization an atom has is to just count the number of atoms bonded to it and the number of lone pairs. Double and triple bonds still count as being only bonded to one atom. Use this method to go over the above problems again and make sure you understand it. It's a lot easier to figure out the hybridization this way.
Introduction Carbon is a perfect example showing the value of hybrid orbitals. Carbon's ground state configuration is: According to Valence Bond Theory , carbon should form two covalent bonds, resulting in a CH 2 , because it has two unpaired electrons in its electronic configuration. That would give us the following configuration: Now that carbon has four unpaired electrons it can have four equal energy bonds.
Energy changes occurring in hybridization. Energy changes occurring in hybridization Hybridization of an s orbital with two p orbitals p x and p y results in three sp 2 hybrid orbitals that are oriented at o angle to each other Figure 3. Example: sp 2 Hybridization in Aluminum Trihydride In aluminum trihydride, one 2s orbital and two 2p orbitals hybridize to form three sp 2 orbitals that align themselves in the trigonal planar structure.
Example: sp 2 Hybridization in Ethene Similar hybridization occurs in each carbon of ethene. Energy changes occurring in hybridization Figure 1: Notice how the energy of the electrons lowers when hybridized. Example: sp Hybridization in Magnesium Hydride In magnesium hydride, the 3s orbital and one of the 3p orbitals from magnesium hybridize to form two sp orbitals. Hybridization Example: sp Hybridization in Ethyne The hybridization in ethyne is similar to the hybridization in magnesium hydride.
References John Olmsted, Gregory M. Carey Advanced Organic Chemistry Springer Wade, Jr. Aug 9, Related questions How does carbon use its "s" and "p" orbitals to form bonds in ethyne, ethene, and ethane?
Question fb1f7. Question a2. How do pi and sigma bonds relate to hybridization? What is the orbital hybridization in BrCl3? What is the orbital hybridization theory? You also need to consider various electronic effects distinct from sterics. There are also phenomena such as the anomeric effect and the gauche effect, which are often explained in terms of hyperconjugation. These are factors in determining lowest-energy conformers for given molecules, and sometimes have subtle effects on bond lengths and bond angles and I mean beyond the obvious impact on dihedral angles.
I was curious as to the statement my prof made. It was something along the lines of that bond angles can be rationalized using vdW repulsions alone.
I suspect your prof may have wanted to avoid getting into certain gray areas which are numerous in chemistry ; I've certainly seen this plenty of times with my own profs. Steric effects which is kind of a catch-all term mostly referring to Pauli electron-electron repulsion are certainly quite dominant in many cases, especially in simple and highly symmetrical structures, but they're not the whole story. Show 1 more comment. Active Oldest Votes. Electronegativity: If the electronegativity of the central atom decreases, bond angle decreases.
Good Side note: Triple bonds repel other bonding-electrons more strongly than double bonds. Improve this answer. Being a concept, hybridization has nothing to do with the actual energetics of why a molecule adopts a certain geometry. In fact, there are many cases were hybridization fails to explain reality, the water molecule being a case in point. But what it allows us is classification. It allows us to associate particular arrangements with particular geometries.
And with Chemistry, there are a ton of exceptions, so it's definitely not perfect.
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