UNIT 2.7 VSEPR AND BOND HYBRIDIZATION
Chemistry is the study of matter and interactions. Chemistry overlaps with many other sciences.
Below is the contents for this sub unit.
VSEPR
Valence Shell Electron Pair Repulsion (VSEPR) Theory.
This is the theory used to predict the geometry, bond angle, interaction of molecules. The electrons in the outer shell (valence electrons) will form bonds but all the valence electrons (including those in bonds) repel each other (repulsion) and try to get as far apart as they can but still stay near the atoms they are assigned to.
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This is how the molecular geometries are determined. Make sure you review that if you haven't yet.
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There are more geometries and include those that break the octet rule. "AHHH, you'r breaking the rule you taught us." I know. Chem 1 is teaching you the rules. Chem 2 and AP Chem is giving you a better understanding of the rules and show where they can be broken.
Do a search for 'molecular geometries' and you will find charts (images) of several different types of geometries
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Linear - Review from chem 1
Trigonal planar (breaks octet rule) - Boron loves being in the middle of these. It has 3 bonds and since it has 3 valence electrons it gets a formal charge of 0 and because it is so small, it is hard to get a 4th bond or an electron pair.
Bent, Trigonal Planar, Tetrahedral - Review from chem 1
Trigonal Bipyramidal, T-Shape, See-saw, Octahedral, Square Pyramidal, Square Planar all break the octet rule. They vary in the number of atoms bound to a center atom and the number of electron pairs surrounding it.
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To check them out, an easy at home lab is to buy toothpicks and marshmallows and see how they arrange in 3D. It is hard to show on a screen.
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BOND ANGLE
So mostly we compare bond angles and the picture above is a good example. CH4, NH3, and H2O all have a tetrahedral shape if you look at the electrons but are tetrahedral, trigonal planar, and bent when you look at just the atoms.
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So a tetrahedral has an angle of about 109.5. Electron pairs repel more than an atom so on NH3 this angle decreases as the electron pushes the other atoms closer together. When you get to H2O the 2 electron pairs push the atoms even closer. This is usually a good question on AP tests as it checks your knowledge of the atom, electrons, geometries, and current theories. They usually just ask you to compare different structures and talk about the differences.
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Marshmallows and toothpicks are cheap. Talk to me during the year if you need help but this is a great way to learn how different geometries are made.
HYBRIDIZATION
So how is the tetrahedral shape even made? The orbitals we learned were s p d and f and they don't look like that.
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Hybridization. So you know what a hybrid car is, uses both gas and battery. A hybrid orbital combines the different orbital. Take the tetrahedral: it combines the s and 3 p orbitals to make 4 bonding orbitals (see picture above).
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What about a double bond like O2? The double bond uses one of the p orbitals so that isn't available. Then each O has 2 electron pair.
Answer: sp2 hybridization. This gives 3 orbital locations. One used for bonding and the other 2 for the 2 electron pair.
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What about a triple bond like N2? The triple bond uses 2 of the p orbitals so that only leaves one and then you have 1 electron pair on each N.
Answer: sp hybridization. This gives 2 orbital locations. One for bonding and the other for the electron pair. This leaves the other 2 p orbitals available for bonding.
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What about geometries that break the octet rule like the octahedral that has 6 atoms bond to a center atom?
Answer: sp3d2 hybridization. The center atom has to be at least row 3 to open up the d orbitals but they can also be used for hybridization.
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Generally. However many atoms and electron pairs are attached to the atom tell you the hybridization.
2 - sp
3 - sp2
4 - sp3
5 - sp3d
6 - sp3d2
MOLECULAR ORBITAL THEORY
Molecular Orbital Theory is a different model used to assign location to the electrons. With proper use, you can tell magnetic properties of substances. It also can be used with unpaired electrons and predict if a substance would be stable.
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In the picture above on the left side we are making H2. We see each H has 1 electron they bring in (assigned to each). In the middle there is a lower bonding orbital and a higher anti-bonding orbital. Each orbital holds 2 electrons (one up and one down). When the bonding orbital is full and the anti-bonding isn't, the molecule is stable
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On the left side we are trying to make He2. You can see the electron fill the bonding and anti-bonding orbital so this molecule is unstable. Being a noble gas we know He doesn't like bonding.
This gets more complicated as you start adding p orbitals and then the order of the bonding and anti-bonding gets complicated as you go to larger atoms.