Download Ch. 13-14 Review

Ch. 13-14 Review
1. When an electron in a hydrogen atom moves from a higher to a lower energy state, the
energy difference is emitted a as quantum of energy.
2. Define the four quantum numbers (n, l, ml, ms), explain what information is given by
each, and describe the range of values each may have.
Principal quantum number (n)
 Describes the principal energy level an electron occupies
 It has values of 1,2,3,4…
Azimuthal quantum number (ℓ)
 Describes the shape of the atomic orbital
 Designates a sublevel
 Values from 0 up to and including (n-1)
 l = 0, 1, 2, 3…,(n-1)
 0 = s, 1 = p, 2 = d, 3 = f
 s = spherical, p = “peanut”, d & f = daisy and fancy
 if n = 1, then ℓ can be 0 (s orbital) = 1 sublevel
 if n = 4, then ℓ can be 0 (s), 1(p), 2(d), 3(f) = 4 sublevels
Magnetic quantum number (ml)






designates the spatial orientation of an atomic orbital in space
values range form – ℓ to + ℓ
s has 1 orbital
p has 3 orbitals
d has 5 orbitals
f has seven
Spin quantum number (ms)
 values of +1/2 and –1/2
 each orbital can hold 2 electrons with opposite spins
 Since spinning charges objects creates a magnetic field the electrons must spin in opposite
directions to minimize repulsion.
3. The space occupied by one pair of electrons is called a(n) orbital.
4. Complete the following table:
Energy level
Number of
sublevels (n)
1
1
2
2
3
3
4
4
Number of orbitals
(n2)
1
4
9
16
Maximum number
of electrons (2n2)
2
8
18
32
5. State the Pauli Exclusion Principle in your own words.
An atomic orbital can contain at most two electrons. They must have opposite spins.
6. Complete the following table
Sublevel
s
p
d
f
Number of orbitals
1
3
5
7
Maximum # of electrons
2
6
10
14
7. Write electron configurations and orbital diagrams for the following atoms:
a. Carbon
1s22s22p2
b. Iron
1s22s22p23s23p64s23d6
c. Lead
[Xe]6s24f145d106p2
8. Assign quantum numbers to the following elements
a. O : 2, 1, -1, -1/2
b. Ca : 4, 0, 0, -1/2
c. Zr : 4, 2, -1, 1/2
9. Which elements are defined by the following set of quantum numbers?
a. n = 4, l = 1, ml = 0, ms = -1/2
Bromine
b. n = 5, l = 0, ml = 0, ms = +1/2
Rubidium
c. n = 3, l = 2, ml = 0, ms = -1/2
Nickel
10. How many sublevels are there in the 3rd energy level? 3
11. How many electrons can occupy a single orbital? 2
12. When n = 5, what are the possible values for l? 0, 1, 2, 3
13. When l = 2, what are the possible values for ml? -2, -1, 0, 1, 2
14. Heisenberg stated that it was impossible to know which two things about an electron at
the same time? Momentum and Position
15. What is Hund’s rule? When electrons occupy orbitals of equal energy, one electron
enters each orbital until all orbitals contain one electron with parallel spins. Second
electrons are added and are paired so that each orbital contains two electrons with
opposite spins. Everybody gets firsts before anybody gets seconds.
16. What is the Aufbau Principle? Electrons enter orbitals of lowest energy first.
17. Write and label the wave equations. Be sure to include deBroglie’s equation, frequency,
energy, wavelength, and all constants.
υ = c/
E=hυ
= h
mv
Amplitude – wave’s height from origin the crest.
Wavelength () – the distance between crests.
Frequency (υ) – the number of wave cycles to pass a given point per unit time.
c = speed of light (3.00 x 108 m/s or 3.00 x 1010 cm/s)
E = Energy
υ = frequency
h = Planck’s constant (6.626 x 10-34 Js)
m = mass in kg
v = velocity in m/s
18. Write the spectrum from high wavelength to low wavelength. Label energy and
frequency as well.
19. Draw and label the parts of the wave.
20. Explain the differences in size between anions and cations and their atomic
counterparts. Use examples.
Anions are larger than the atoms they came from because they gain electrons,
increasing the outward pull by the electrons. Cations are smaller than the atoms they
came from since they lost an entire energy level and there are now more protons than
electrons, pulling those electrons in tighter. O vs. O2- : Both have 8 protons, but the
anion has 2 more electrons pulling outwards. Na vs. Na+: Both have 11 protons, but the
cation has 1 fewer electron pulling outwards and the electron lost was the only one in
that energy level.
21. Define isoelectronic and give an example of 5 elements that are isoelectronic.
Isoelectronic - A group of ions with the same number of electrons. The one with the highest
atomic number is the smallest in size (More protons pulling on the same # of electrons).
Na+, Mg2+, Ne, F-, O2-, N3- are isoelectronic. They all have 10 electrons. Mg2+ is the smallest
because it has 12 protons pulling on 10 electrons. (The protons win the “tug of war”) N3- is the
largest because it has 7 protons and 10 electrons. (The electrons win the “tug of war”)
22. For Cl and I, indicate which has the higher
a. Electronegativity - Chlorine
b. Electron affinity - Chlorine
c. First ionization energy – Chlorine
d. Atomic radius - Iodine
e. Ionic radius – Iodine
23. Draw simplified periodic tables that illustrate all trends you have learned. For each trend
write two sentences, in your own words, describing WHY the trends are the way they
are. Include any exceptions in your drawings and explain those as well.
You do this!