1. CHAPTER 7 AND 16
AP CHEMISTRY

2. LIGHT Wavelength = λ nm/ wave, m/wave, or nm or m
λf= c; c = x 108 m/s
Frequency = f
Waves/s or 1/s = c/ λ
High frequency = short wavelength
Low frequency = long wavelength
The wavelength of the radiation which produces the yellow color of a sodium vapor light is nm. What is the frequency of the radiation?
/ 109 nm / 2.998x 108 m
589.0 nm/ 1 m / s
5.090 x /s

3. PHOTON ENERGY E = hf Planck’s constant h = 6.626 x 10-34 J.s
From the last question, give the amount of energy, in Joules, that a photon would emit by an excited sodium atom? In kilojoules?
3.373 x J
3.373 x kJ
Continuous spectrum
All wavelengths
Line spectrum
Discrete energies light going through a prism
ROY G BIV

4. HYDROGEN ATOM Bohr model used hydrogen atom
Electron orbits around nucleus
Energy levels: Lyman 364 and Balmer 656
Zero energy
where a photon and electron are in completely separate energy states within an atom
Ground state
lowest amount of energy an electron has
Excited state
electron getting extra energy
E = hf = Ehi - Elo
Electrons can only give off certain energies

5. EXAMPLE
A green light of wavelength 456 nm is observed in emission spectrum of hydrogen. Calculate the energy difference, Ehi -Elo, between the two states that are responsible for this line in kJ/mole
/ 109 nm / 2.998x 108 m
456 nm / 1 m / s
6.626 x J.s / x 1014 / kJ_____
/ s / J
4.356 x kJ / 6.02 x 1023atoms
Atom / mole
262 kJ/mol

6. QUANTUM NUMBERS Most probable place an electron can be found
Ψ amplitude (height) of the electron wave
Ψ 2 directly proportional to the probability of finding the electron
Quantum numbers
Principal energy level
n = 1, 2, 3, etc.
Sublevels
s, p, d, and f
l = 0, 1, 2, 3, etc.

7. CONTINUE Orbitals or orientations m = l to -l l = 0 then m = 0
Spin
s = +1/2 , -1/2
Pauli Exclusion Principle
No two electrons can have the same four quantum numbers
Capacities
Each energy level (n) has n sublevels
Each sublevel (l) has 2l + 1 orbitals
Each orbital (m) has 2 electrons
Total number of orbitals in an energy level is n2 and total number of electrons in an energy level is 2n2
How many electrons can fit in principal energy level 3?

8. CONTINUE Orbital does not mean the electron is circling the nucleus
We do not know how the electron moves
Heisenberg principle
state that there are limitations in knowing what the position and momentum are at any given time
This uncertainty does not mean much to a large particle (ball or you) but with an electron this means a lot
Probability distribution intensity of color -electron density map
Most probable place to find a hydrogen electron is .529Ǻ from nucleus
1Ǻ = m
Radius of 1s orbital is to encompass 90% of the probability for that electron

9. ELECTRON CONFIGURATION
Electron notation
Usually a sublevel is filled to capacity before entering the next one
Do the electron configuration of the following Li
1s22s1 B
1s22s22p1 N
1s22s22p3 O
1s22s22p4
What are the electron configuration of the chlorine atom and iron atom?
What are the abbreviated electron configurations for the last two atoms?

10. QUANTUM NUMBERS Pages 296 to 298 Polyelectronic atoms
Treat electrons as if they have nuclear attraction and average repulsion from other electrons
Effective nuclear charge
Zeff = Zactual - (electron repulsion) Z = atomic number
Periodic table can predict the filling of the sublevels
Elements in groups 1 and 2 fill s
Elements in groups fill p
Transition metals in groups fill d
The two sets of 14 elements at the bottom of the table fill f
Write the electron configuration of tellurium.

11. MENDELEEV
In 1869 Mendeleev introduced an arrangement of the elements based on the fact that their chemical properties vary periodically with their atomic mass
He left gaps in his table and predicted some detailed physical and chemical properties for three elements that had not be found yet. i.e. germanium (ekasilicon) mass 73 the observed mass is 72.3
Some atoms have different electron configurations than would be predicted
Cr predicted 1s22s22p63s23p64s23d4
Actual 1s22s22p63s23p64s13d5
Cu predicted 1s22s22p63s23p64s23d9
Actual 1s22s22p63s23p64s13d10
Aufbau principle (building up) - add one electron to lowest energy orbit until it is filled
Valence electrons = outer most electrons
Inner electrons = core electrons

12. ORBITAL DIAGRAM B C Hund’s rule
Within a given sublevel, the order of filling is such that there is the maximum number of half-filled orbitals
Write the orbital diagram for silicon atom and cobalt atom
Monoatomic ion electron configuration
Ions with noble gas configuration
Elements close to a noble gas will form an ion that has the same configuration as the noble gas
19K --> K+ + 1e Ca --> Ca2+ + 2e-
17Cl-+ 1e- --> Cl S2- + 2e- --> S
Anions are formed from atoms with 1, 2, or 3 less electrons than the next noble gas
Cations are formed from atoms with 1, 2, or 3 more electrons than the preceding noble gas

13. TRANSITION METAL CATIONS
When transition metal atoms form positive ions, the outer s electrons are lost first
Mn atom 1s22s22p63s23p64s23d5
Mn2+ ion 1s22s22p63s23p63d5
Cu atom 1s22s22p63s23p64s13d10
Cu+ ion 1s22s22p63s23p63d10
Cu2+ ion 1s22s22p63s23p63d9
Give the abbreviated electron configuration of
Ti atom and Ti2+ ion
S atom and S2- ion
Ni atom and Ni2+ ion

14. PERIODIC TRENDS Atomic radius In general atomic radii
Decreases as you move left to right
Increases as you move down
Effective nuclear charge
Zeffective = Z - S
Z = number or protons
S = number of core electrons that are shielding the outer electrons from the nucleus

15. IONIC RADIUS Ionic radius
Ionic radius increases as you move down the group
Cations decrease left to right
Anions decrease left to right
Negative ions are larger than the nonmetal atoms from which they are formed
Positive ions are smaller than the metal atoms from which they are formed.
Using only the periodic table, arrange each of the following sets of atoms and ions in order of increasing size
Na, Al, Sr
O, Cl-, Se2-
V, V2+, V3+

16. IONIZATION ENERGY Energy required to remove an electron
First ionization energy
Na(g) ---> Na+(g) + 1e-
Mg(g) ---> Mg+(g) + 1e-
Second ionization energy
Mg+(g) ---> Mg2+(g) + 1e-
Increase as you move across left to right
Decrease as you move down
Inverse correlation to atomic radius
Consider three elements C,N, P. using the periodic table, predict which of the elements has
The largest atomic radius. The smallest
The largest ionization energy. The smallest

17. CONTINUED Energy required to remove an electron from a gaseous atom
E = (Z2eff/n2)(1310)kJ/mol
n = Zeff = 10.3
n = Zeff = 1.84
3s has 10 electrons repulsing the 11th electron. This reduces the pull of the electron toward it (shielding)
Shielding is not totally effective, because an outer electron could be found by the nucleus (though probability is low)
Page 318 fig
Core electrons provide effective shielding for the outer electrons
Page 321

18. ELECTRONEGATIVITY
Energy released from an atom as it acquires another electron
Page 324
Atomic radius
Decreases going across because the valence electrons are being pulled in, due to an increase of proton attraction. Shielding remains constant across the period
Electronegativity
How much an atom wants an electron
Which element has the highest electronegativity?
Fluorine

19. METALS
Metals
Metallic luster, ductile, malleable, good conductor of heat and electricity
Melting point increases with an increase in atomic number as you go into the center of the periodic table
Metals tend to lose electrons when they undergo chemical reactions
2Mg(s) + O2(g) ----> 2MgO(s) (oxidation number?)
Compounds formed from metals and nonmetals tend to be ionic
Most metal oxides are basic oxides
Metal oxide Water Metal hydroxide
MgO(s) H2O(l) > Mg(OH)2(aq)
Metal oxide Acid Salt Water
Fe2O3(s) HCl(aq) > 2FeCl3(aq) + H2O(l)

20. METALS AND NONMETALS Nonmetals Metals + nonmetals ----> salts
No luster, poor conductors
Nonmetals tend to gain electrons to become anions
Metals + nonmetals ----> salts
When compounds are made up chiefly of nonmetals they are molecular compounds
Most nonmetal oxides are acidic oxides
Nonmetal oxide water Acid
CO2(g) H2O(l) ----> H2CO3(aq)
Nonmetal oxide Base Salt Water
SO3(g) KOH(aq) ----> K2SO4(aq) + H2O(l)
Semimetals
Some properties of both, brittle, semi-conductor
Trends in metallic characteristics
The more an element shows physical and chemical properties characteristic for metals, increase right to left and top to bottom

21. THERMODYNAMICS Three thermodynamic functions
ΔH; Change in enthalpy, negative values of ΔH tend toward spontaneous
ΔS; change in entropy, positive values of ΔS tend toward spontaneous
Nature spontaneously moves to the highest probability of occurrence
Increase of entropy Ssolid < Sliquid <Sgas
Which has more entropy
Solid CO2 or gaseous CO2
N2 gas at 1 atm or N2 gas at 1.0 x 10-2 atm
N2 gas at 1.0 x 10-2 atm would be right
Predict the sign of entropy
Solid sugar is added to water to form a solution
Iodine vapor condenses on a cold surface to form crystals
ΔG; change in free energy, a reaction at constant temperature and pressure will be spontaneous if ΔG is negative
Does not mean it is fast
Means it occurs without intervention
Tells us which direction it will go

22. SECOND LAW OF THERMODYNAMICS
Any spontaneous process will increase entropy in the universe. Entropy will always increase
Example
CH4(g) + 2O2(g) ----> CO2(g) + 2H2O(g)
Factors affecting it:
Energy factor: at 25 oC, 1 atm, exothermic reactions are spontaneous (H < 0)
Randomness factor; when other things are equal system tends to move from an ordered structure to a random one

23. ENTROPY ΔS = Sproduct - Sreactants
This is done the same way as enthalpy
Measure of change in order
As the number of moles increase, the ΔS will be positive (usually)
2SO3(g) ----> 2SO2(g) + O2(g): ΔS = 187 J
ΔS = [2(248) + 1(205)] - [2(257)]
N2(g) + 3H2(g) ----> 2NH3(g); what is the ΔS?
Reactions where ΔSo is positive tend to be spontaneous at high temperatures

24. WATER H2O(l) ----> H2O(g) Δ Ssys is positive, entropy increases
Δ Ssurr determined by the flow of heat. Heat is going into the system, randomness decreases, Δ Ssurr is negative
When Δ Ssys = +, and Δ Ssurr = -. Δ Suniv will be temperature dependent
Above 100 oC, spontaneous the direction it is written
Below 100 oC, occurs opposite direction
Sign of Δ Ssurr depends on the direction of the heat flow - lowest energy
Magnitude of ΔSsurr depends on the temperature
ΔSsurr = +J/K exothermic
ΔSsurr = -J/K endothermic
ΔSsurr = -ΔH (kJ need to change to J)
T(oC need to change to K)

25. PROBLEM
In the metallurgy of antimony, the pure metal is recovered via different reactions, depending on the composition of the ore. For example
Sb2S3(s) + 2Fe(s) -----> 2Sb(s) + 3FeS(s) ΔH = -125kJ
Sb4O6(s) + 6C(s) > 4Sb(s) + 6CO(g) Δ H = 778 kJ
Δ Ssurr = - Δ H (kJ need to change to J)
T (oC need to change to K)
_ x 105J
298 K
419 J/K ΔS positive so exothermic
_ 7.78 x 105J
-2.61 x 103J/K ΔS negative so endothermic
Page 778

26. FREE ENERGY
ΔG = ΔH - T ΔS
At a constant temperature and pressure a process will be spontaneous if ΔG is negative, ΔG (free energy) is decreasing. Energy not used in breaking down or building up
ΔSo = ΔSsys , (o) all substances are in their standard state
Page 780 tables 16.4 and 16.5
What temperature is this process spontaneous at 1 atm?
Br2(l) -----> Br2(g)
What is the normal boiling point? (Hint at the boiling point Br2(l) and Br2(g) are at equilibrium so ΔGo = 0)
333K above this number ΔSo controls it, below this number ΔHo controls it.

27. THIRD LAW OF THERMODYNAMICS
At zero Kelvin the entropy of a perfect crystal would be zero (unattainable ideal)
ΔS found at 298 K and 1 atm
The more complex a molecule is, the higher ΔS value is
ΔG = change in free energy occurs when reactants are in standard states forming products in standard state
As the ΔG increases in negative value the further right it will go to reach equilibrium
2SO2(g) + O2(g) -----> 2SO3(g)
Carried out at 25 oC and 1 atm. Calculate ΔH, ΔS, and ΔG
ΔHf (kJ/mol) ΔS (J/K.mol) SO2(g)
SO3(g)
O2(g)

28. ΔG = ΔH - T ΔS
ΔG is dependent on pressure, concentration, and temperature
ΔG < 0 spontaneous
ΔG > 0 nonspontaneous
ΔG = 0 equilibrium
Effect of ΔH and ΔS on spontaneity
If ΔH > 0, ΔS < 0, ΔG > 0 at all temperatures it will be nonspontaneous
If ΔH < 0, ΔS > 0, ΔG < 0 at all temperatures it will be spontaneous
If ΔH > 0, ΔS > 0, then ΔG > 0 at low temperatures, and ΔG < 0 at high temperatures so it will be spontaneous at high temperatures
If ΔH < 0, ΔS < 0, then ΔG < 0 at low temperatures, and ΔG > 0 at high temperatures so it will be nonspontaneous at high temperatures

29. SECOND METHOD TO CALCULATE ΔG
Hess’s law
Using the following data (25oC)
Cdiamond(s) + O2(g) ----> CO2(g) ΔG = -397 kJ
Cgraphite(s) + O2(g) ----> CO2(g) ΔG = -394 kJ
Calculate ΔG for the reaction
Cdiamond(s) > Cgraphite(s)
Third method to calculate ΔGf (formation)
ΔGf = ∑ np ΔGf (product) - ∑nrΔGf (reactant)
Methanol is a high-octane fuel. Calculate ΔG for the reaction
2CH3OH(g) + 3O2(g) ----> 2CO2(g) + 4H2O(g)
ΔGf (kJ/mol) ΔGf (kJ/mol)
CH3OH(g) CO2(g) O2(g) H2O(g)