AP Chem Ch 7-9 Flashcards
Principle QN (n)
Energy level
Integer from 1-7
Quantum numbers
Four numbers that describe an electron
- Principle QN (n)
- Angular momentum QN (l)
- Magnetic QN (ml)
- Spin QN (ms)
Angular momentum QN (l)
From 0 to n-1 When l = 0, s orgbital 1 --> p 2--> d 3 --> f
Magnetic QN (ml)
From -l to l
So when l = 0, one possibility for S orbital
When l = 1, 3 possibilities in p orbital
Spin QN
+/- 1/2
Pauli exclusion principle
No two electrons can be in the same spot
Aufbau principle
Build up – put electrons into lowest energy orbital available
Hund’s Rule
Lowest Energy confirmation generates unpaired electron – result of energy required to pair the electron in the same orbitals.
This means first do it unpaired and then pair them up
Auger emission
When a core electron is ejected, a higher energy may relax into that electron hole, resulting in an emission of energy
Exceptions to normal electron configuration
Chromium and molybendium fill up d orbital half way and have only one S electron
Copper, silver, gold all fill up the d orbital and have one in the S
Pd has no s electrons, 10 d
Homogenous magnetic field
Orient dipoles– spins align. Observed splitting of the 5s energy level with the S going to 2 different energy states. Pauli connected this to the spin in the electrons with the +1/2 and -1/2, the 2 intrinsic states of electron.
Periodic trend key factors
Number of protons, number of electrons, effective nuclear charge, down the group, across the period, shielding, distance from the nucleus
Electron affinity
The amount of energy to add an electron
Energy associated with this rxn:
X + e- –> X-
Atomic radius / size periodic trend
Increases down and to the left. Francium the biggest, helium the smallest
Ionization energy
Amount of energy required to remove an electron
M–> M+ + e-
Increases top right
Helium has the most because very hard to remove an electron
Francium the least, easier.
Increases across the period mostly, but some exceptions as move from a metal to a non metal
Electronegativity
How much an element wants an electron
Measured in Paulings, arbitrary unit for how much an element wants the electron
Increase top right. Fluorine has the most, a 4.0
Effective nuclear charge
Number of protons - inner electrons
Aluminum is +3
Chrlorine +7
Why does IE increase left to right?
Effective nuclear charge increases because number of protons increases, so while shielding stays the same because same energy level, it is harder to remove an electron against this extra positive charge from the nucleus while the atom is smaller as well.
Decreases as go down a row because it is easier to remove an electron from a bigger atom (shielding effect in effect)
Why does oxygen have lower IE than nitrogen?
Nitrogen has the p orbital half way full, so it is more stable, so it thus harder to remove that electron than from oxygen
Ionic bonding
Involves the transfer of electrons and is the positive / attractive force that results from the charged species
Electrostatic interactions
Interactions of charged particles –> forces keeping electrons around electrons –> force that makes ionic bonds ionic.
Anion
Negative ions. Non metal
Cation
Positive ions. Metal
Bond energy
Energy required to break a bond
E= (2.3110^-19 Jnm) q1 q2 / d where d is the distance between 2 nuclei
Bond energy of NaCl if the bond length is 2.76 A*
A* = A with the degrees symbol = 10^-10 m E = 2.31*10^-19 (1)(-1)/.276 nm
Covalent bonding
Share electrons. More like they’re fighting over the electrons. Nonemtals and hydrogen
As you drag a hydrogen atom closer to another, lowest energy is at the bond length
How to get delta H overall of Li (s) + 1/2 F2 (g) –> LiF(s)
Need to do a bunch of half reactions and add up the delta H’s
Li (s) –> Li (g) delta H sub = 161 kJ/mol
Li (g) –> Li+ (g) + e- delta H IE = 520
1/2 F2 (g) –> F (g) delta H = 154
F(g) + e- –> F- (g) delta H EA = -328
And lastly need a lattice energy–>
Li+ + F- –> LiF (s) delta H lattice = -1047
Add up all the delta H’s and done
Lattice energy dependent on size and magnitude of charge
Polar
Uneven distribution of electrons
Results in a dipole with a partial positive and a partial negative end
Non polar covalent bonds
When electronegativity is similar or the same
H-C or O-O or any diatomic bond is not polar
Electron domains
Areas of electron density
Bonds or lone pairs around a central atom
How many domains does a double bond count as
ONE!
Formal charge
of valence electrons - bonding electrons / 2 - # of electrons in lone pair
Example of formal charge of an oxygen with a double bond and 2 lone pairs vs an oxygen with a single bond and 3 lone pairs
With double bond --> 6 - 4/2 -4 = 0 Single bond --> 6 -2/2 -6 = -1 0 is more favorable
Resonance
Result of having more than one valid Lewis dog sfrucutre in which only difference is the placement of electrons
Chemical bonds
Force that causes atoms to behave as a unit
Sigma vs pi bond
Sigma bond is much stronger than pi because sigma has more orbitals overlapping.
Triple bond shortest and most strong because has sigma plus additional pi
Another way for delta H rxn
Sum of the energy required to break the bonds (left side) - sum of the energy released from bonds formed (right side)
Method for drawing Lewis dot structures
- Count valence electrons
- Re count valence electrons
- Arrange the atoms so that the central atom is the least electronegative
- Use the pairs of electrons to form single bonds between the atoms, the skeleton structure
- Spread remaining electrons on non central atoms trying to form octet
- Put extra electrons on central atom
- Check if central atom has octet
- Form double or triple bonds to give central atom an octet
Note that don’t need octet. Atoms at P and larger can have 10-12 electrons and be stable
Draw some Lewis dot structures!
Ok!
Equivalent resonance structure
When there are the same number of single, double, and triple bonds. The bonds are meeely rearranged. Weighted equally in trying to predict actual structure
Non equivalent resonance structures
Contain different numbers of single, double, and triple bonds (for example, four single bonds become one double and two single bonds). When these happen, it is possible one is more likely than the other. Determine which is more likely by looking at the formal charge. Want the one with the least variability (0,0,0 vs 2,-4,2)
Linear
With 2 electron domains and 2 bonds
Like CO2
Sp
Trigonal planar
3 electron domains and 3 bonds
BF3
No lone pairs
Sp2
Bent
3 electron domains and 2 bonds. For example O3. Has a lone pair sp2
Also could be 4 electron domains and 2 bonds (2 lone pairs), such as H2O sp3
Tetrahedral
4 electron domains and 4 bonds (no lone pairs)
CH4. 109.5 degrees sp3
Trigonal pyramidal
4 electron domains. 3 bonds (1 lone pair)
NH3
Sp3
Trigonal bipyramidal
5 electron domains, 5 bonds (no lone pairs)
PCl5
Sp3d
See-saw
5 electron domains but 4 bonds (1 lone pair)
SF4
Sp3d
T-shaped
5 electron domains and 3 bonds (2 lone pairs)
ClF3
Sp3d
Linear with 5 electron domains
XeF2
2 bonds and 5 domains
Sp3d
Octahedral
SF6
6 electron domains and 6 bonds
Sp3d2
Square pyramidal
6 electron domains and 5 bonds
BrF5
Square planar
6 electron domains and 4 bonds
XeF4
Hybridization
Orbitals formed between the current orbitals to allow for equal energy bonds
Molecular orbital theory
Sigma star where unpaired electrons go and sigma where bonding electrons go
If there are more or the same electrons in the sigma star than regular sigma then there isn’t a bond
Why is the first IE of sodium smaller than magnesium but the second IE of sodium is larger than the second IE of magnesium?
- First IE of Na is smaller because sodium is bigger than magnesium and it has one less proton, so less of a positive pull from the center, so easier to remove an electron
- Sodiums second electron is in the next energy level, which is even closer to the nucleus as we lose an extra orbital, so it is really hard to remove this electron
Bond order
(Number of bonding electrons - number of anti bonding electrons)/2
What does bond order signify?
Number of bonds likely to occur. In He2, the BO=0, so He2 won’t happen.
Larger bond order means stronger interaction
Paramagnetism
Attracted to an external magnetic field
UNPAIRED ELECTRON
Diamagnetism
Repelled by an external magnetic field
All electrons are paired
Relation between bond order, energy, and length
Increase bond order, increase bond energy, decrease bond length
Nitrogen triple bond
Materials with high numbers of N atoms are explosive because their decomposition results in the formation of N2, a very stable product, thus releasing large quantities of energy
What does HF molecular orbital diagram show us?
The molecular sigma orbital is closer in energy to the Florine 2p orbital, indicating that the electron will spend more time on F, supporting our theory of electronegativity and the more electronegative atom getting the electron more
