Exam 1 Covers ch 9, 10, 20, 24 with book Chemistry A Molecular Approach
I am writing this without revision (because I'm lazy), so if something needs refining, there is an error, and/or there is a better way of explaining something, please mention it in the comments.
Electronegativity:
A scalar of how much an atom attracts electrons.
Covalent Bond:
A bond that exists between two nonmetals where the difference in electronegativity between the two atom is less than 0.5. In Covalent bonds the electrons are shared between the two atoms.
Polar Bond:
A bond between two atoms where their difference in electronegativity is between 0.5 and 2.0 inclusive. In polar bonds, the more electronegative atom hogs the electrons in the sharing process.
Ionic Bond:
A bond that exists between a metal and nonmetal where the difference in electronegativity is greater than 2.0. In ionic bonds, one atom has taken the electrons it needs as opposed to sharing them.
Metallic Bond:
A bond between two metals. In this type of bond electrons flow across the compound rather than get shared between atoms. This is a property that makes metals good of carrying an electrical current. Also because the atoms are not connected like X-Y, they can be molded into shapes non metallic bonds could not.
(taken from MC homework)
What are the three basic types of chemical bonds? What happens to electrons in the bonding atoms in each case?
The three types of bonds are ionic bonds, which occur between metals and nonmetals and are characterized by the transfer of electrons; covalent bonds, which occur between nonmetals and are characterized by the sharing of electrons; and metallic bonds, which occur between metals and are characterized by electrons being pooled.
A compound with be stronger than another if its radii is smaller. A good prediction for relative radii size is adding the number of electrons. So Mg2C = 12+6 = 18 and CaO = 20 + 8 = 28. Since 18 is smaller than 28 it can be predicted that Mg2C has a stronger ionic bond than CaO.
However prediction is not necessary if one knows the lattice energy between to atoms. This will be given in a table on the exam. Note higher lattice energy = higher bond strength = smaller radii = smaller bond length and vis versa.
The electronegativity of atoms will be given on the exam. Lets call the difference in electronegativity ΔE.
Covalent: 0 <= ΔE < 0.5
Polar: 0.5 <= ΔE < 2.0
Ionic: 2.0 <= ΔE < ∞
- BUILD SCELETON: Start by drawing the skeleton structure of the element. This is done by connecting atoms with single bonds. Note, the more electronegative atoms tend to be on the outside of the structure. All atoms can have 4 bonds (an octet 8 valance electrons), except:
Hydrogen (H) can only have 1
Helium (He) can only have 2
Beryllium (Be) can only have 2
Boron (B) can only have 3
Atoms with an atomic number greater than 18 have access to more electron shells and therefore can exceed 4 bonds.
Example CCl4
Cl
|
Cl - C - Cl
|
Cl
HONC, helping you build the structure. HONC is an easy way to remember how may bonds H, O, N, and C tend to have. The number of bonds is suggested by the placement of the letters, so:
H tends to have 1
O tends to have 2
N tends to have 3
C tends to have 4
- GET OCTETS: The goal now is to satisfy all atoms by getting them electrons to fill their valance shell, giving them 8 valance electrons (except for exceptions outlined above with H, He, Be, and B which want 2, 4, 6, and 6 electrons respectively).
You have a limited number of electrons to work with that we'll call glue. The number of electron glue you have is calculated by summing the valance electrons of each atom in the compound. Example, H2O = 12 + 61 = 8, CCl4 = 41 + 74 = 32.
Each single bond uses 2 electrons from the glue so the structure limits the amount of glue you have to use to get octets. Note: single bonds, double bonds, and triple bonds use 2, 4, and 6 electrons respectively. Each bond provides the number of electrons it uses to each atom it bonds, these electrons are called bonding electrons. Add lone pair electron from the glue to each atom to satisfy them until you run out. Then rearrange and implement double and triple bonds to compensate.
If it is not possible to build, change the structure.
A formal charge is calculated for each atom by
FC = # valance electrons - # lone pair electrons * 2 - ∑[for each bond](bond order)
The sum of the formal charges for each atom equals the ionic charge of the compound. If it shouldn't be ionic, the sum of formal charges should be zero.
If there are formal charges, the negative charges should be on the outside on the more electronegative atoms, and the passive formal charges on the central and lower electronegative atoms. Nature performs formal changes of zero. It is sometimes possible to rid of formal changes by expanding an octet on atoms with an atomic number greater than 18.
Since we already know about lattice energies and the values are given on the exam, this is fairly simple. Empathy of reaction is usually denoted by ΔH. It is simply calculated by the sum of lattice energy between bond broken subtracted to the sum of lattice energy between bonds joined. Because of it's nature the calculation can be further simpled to (sum of lattice energy between bonds in result) - (sum of lattice energy between bonds in source) because the bonds that don't change cancel out. So:
ΔH = ∑[each bond in result](lattice energy) - ∑[each bond in source](lattice energy)
According to VSEPR theory, what determines the geometry of a molecule?
According to VSEPR theory, the repulsion between electron groups on interior atoms of a molecule determines the geometry of the molecule.
What is a chemical bond according to valence bond theory?
According to valence bond theory a chemical bond results from the overlap of two half-filled orbitals with spin-pairing of the two valence electrons.
What is hybridization?
Hybridization is a mathematical procedure in which the standard atomic orbitals are combined to form new atomic orbitals called hybrid orbitals. Hybrid orbitals are still localized on individual atoms, but they have different shapes and energies from those of standard atomic orbitals.
Why is hybridization necessary in valence bond theory?
They are necessary in valence bond theory because they correspond more closely to the actual distribution of electrons in chemicallybonded atoms.
What is a chemical bond according to molecular orbital theory?
In molecular orbital theory, atoms will bond when the electrons in the atoms can lower their energy by occupying the molecular orbitals of the resultant molecule.
What is a bonding molecular orbital?
A bonding molecular orbital is lower in energy than the atomic orbitals from which it is formed. There is an increased electron density in the internuclear region.
What is an antibonding molecular orbital?
An antibonding molecular orbital is higher in energy than the atomic orbitals from which it is formed. There is less electron density in the internuclear region, which results in a node.
For this section you will need to memories the table on p408, also recreated bellow.
| Electron Groups | Bonding Groups | Lone Pairs | Electron Geometry | Molecular Geometry | ~ Bond Angles |
|---|---|---|---|---|---|
| 2 | 2 | 0 | Linear | Linear | 180 |
| 3 | 3 | 0 | Triangular Planar | Triangular Planar | 120 |
| 3 | 2 | 1 | Triangular Planar | Bent | <120 |
| 4 | 4 | 0 | Tetrahedral | Tetrahedral | 109.5 |
| 4 | 3 | 1 | Tetrahedral | Triangular Pyramidal | <109.5 |
| 4 | 2 | 2 | Tetrahedral | Bent | <109.5 |
| 5 | 5 | 0 | Trigonal Bipyramidal | Trigonal Bipyramidal | 120 (equatorial), 90 (axial) |
| 5 | 4 | 1 | Trigonal Bipyramidal | Seesaw | <120 (equatorial), <90 (axial) |
| 5 | 3 | 2 | Trigonal Bipyramidal | T-Shaped | <90 |
| 5 | 2 | 3 | Trigonal Bipyramidal | Linear | 180 |
| 6 | 6 | 0 | Octahedral | Octahedral | 90 |
| 6 | 5 | 1 | Octahedral | Square Pyramidal | <90 |
| 6 | 4 | 2 | Octahedral | Square Planar | 90 |