What we should realize here is that those
ideal bond angles aren't the same for every
known compound.
Here's the thing, we're going to say those
ideal bond angles exist only if our central
element has no lone pairs.
Once our central element starts to get lone
pairs, it's going to compress the bond angle.
It's going to make it smaller.
Here we can see an example.
Here we have methane which is CH4.
Its electronic geometry will be AX4.
It has no electron pairs around it, no lone
pairs around the central element, so its ideal
bond angle, its perfect bond angle would be
109.5.
But if we moved over to ammonia, NH3, we have
our first lone pair involved.
Lone pairs want to be as far away as everyone
else.
This is going to push the other bonds away
from it.
This causes them to compress or get smaller
and that actually makes the bond angle smaller.
Water, now we have an additional bond angle,
getting smaller, because now there are two
lone pairs pushing away.
What you're supposed to take from this is
those ideal bond angles are only if the central
element has no lone pairs.
Once the central element starts to have lone
pairs the bond is going to get smaller and
smaller.
Of course, your professor is not going to
want you to memorize every single bond angle
known to man.
All you would have to say is, you don't need
to know this exact bond angle, all you need
to know is that the electronic geometry is
AX4, so technically it's tetrahedral.
The ideal bond angle is 109.5, but because
that lone pair is there, all you'd have to
really say is, you would expect the bond angle
to be less than 109.5.
Here, since you have two lone pairs, you could
say the same exact thing again, its electronic
geometry is still AX4, ideally, it should
be 109.5, but the lone pairs being there,
make it less than 109.5.
This is what your professor would be looking
for and this is what you would have to say.
