Download Mr. Knittel`s Final Review Sheet I Answers

Honors Chemistry
Final Exam Review Sheet I
Name: ANSWERS (in red)
Exam Format
60 MC
4 Open Repsone
What to bring
#2 pencils
Scientific calculators
What is given
Periodic table
Names of ions and charges
Equations and constants (exact list on LCD projector)
1. Write electron configurations for the following (with or without the shortcut):
Cu = 1s22s22p63s23p64s13d10
Sodium ion = 1s22s22p6
2 2
6 2
3
P = 1s 2s 2p 3s 3p
Cl = 1s22s22p63s23p5
2 2
6 2
2
Si = 1s 2s 2p 3s 3p
Ba = [Xe]6s24f145d106p3
2. What are the three rules that are followed when completing electron configurations? Name and define
each. The Aufbau principle states that electrons will fill orbitals in order of increasing energy. Hund’s rule
states that electrons filling multiple orbitals with the same energy (for example, 2p) will fill each orbital
with a single electron before pairing. The Pauli exclusion principle states that no two electrons in the
same atom can have the same four quantum numbers.
3. For each of the rules in question 2, give 1 example of a violation of each rule. A violation of the
Aufbau principle would be the following: 1s22s22p63s23p63d5. The 3d orbital has more energy than the 4s
orbital, and therefore, should be filled AFTER the 4s is filled. A violation of Hund’s rule would be
Notice that the arrows in the 2p orbital are not spread out among all the orientations before pairing in (b).
A violation of the Pauli exclusion principle would be the following:
Notice that both arrows are pointing up, indicating that they have the same spin which would mean they
have the same four quantum numbers.
4. Write the orbital notation for the elements Na and Ar.
Na:
Ar:
NOTE: Orbital notation diagrams can be drawn from left to right as seen in the sodium example above, or
can be drawn from bottom to top as seen in the argon example above.
5. Compare and contrast the different types of orbitals. How are they similar and different? Orbitals are
3-D regions of probability for where an electron pair might be found within an atom. Each energy level
can hold a specified number of orbitals equal to n2. The four types of orbitals are s, p, d and f. The s
orbital only has one orientation in space because it is shaped like a sphere. This type of orbital is found in
all of the energy levels. The p orbital has three orientations in space and they are each shaped like a figure
8. There are no p orbitals in the first energy level. The d orbital has 5 orientations in space, most of which
look like four-leaf clovers, with one of the shapes looking like a figure 8 with a donut around its middle.
The d orbital does not exist in the first or second energy level. The f orbital has 7 orientations in space,
and the pictures are too complex to be drawn or described simply. The f orbitals do not exist in the first,
second, or third energy levels.
6. What are the 4 quantum numbers? Provide the name and symbol of each. The four quantum numbers
(in order) are: principal quantum number (n), angular momentum quantum number (l), magnetic quantum
number (m), and the spin quantum number (ms). The principal quantum number identifies the energy
level of the electron; the angular momentum quantum number specifies the type of orbital the electron is
found in; the magnetic quantum number specifies the orientation of the orbital the electron is found in;
and the spin quantum number distinguishes between the two electrons that share the previous three
quantum numbers (spin up or spin down).
7. What is the electromagnetic spectrum? Order the types of light from highest frequency to lowest. The
electromagnetic spectrum is the range of all types of electromagnetic energy (light) in order of increasing
(or decreasing) energy. From highest to lowest frequency (also highest to lowest energy), the order of the
types of light is as follows: gamma rays, x-rays, ultraviolet (UV), visible (VIBGYOR), infrared (IR),
microwaves, and radio waves.
8. How much energy is released when an electron gives releases a photon with a wavelength of 500. nm?
E = 3.98 x 10-19 J
9. If the frequency of a type of light is 4.0 MHz, what is its wavelength? λ = 75 m
10. For the 5 models of the atom: draw a picture of each and describe the progression that was made from
the original model to our current understanding of the atom. In Dalton’s atomic theory atoms were seen as
indivisible pieces of matter, and were likened to billiard balls (hard, compact, spheres). J. J. Thomson’s
work with cathode ray tubes led him to the discovery of the electron—a subatomic particle. Dalton’s
theory needed to be modified to incorporate this idea so Thomson developed the “plum pudding” model
in which negative charges were stuck into a positive blob, like plums in plum pudding. Although
Thomson didn’t find evidence of the positive charges, he reasoned they must be there to balance the
negatives. When Rutherford discovered the nucleus, Thomson’s model had to be refined further.
Rutherford placed a small, dense positive charge in the center of the atomic model surrounded by a blob
of electrons. Although Rutherford’s model is often called the planetary model, be careful not to mix it up
with Bohr – Rutherford’s model did NOT have fixed orbits for the electrons. Realizing that Rutherford’s
model would fail because there was nothing preventing the orbiting electrons from plunging into the
positive nucleus (opposites attract), Bohr modified the model further after his work with the line emission
spectrum of hydrogen. Bohr incorporated the concept of fixed energy levels where the electrons can
reside that are at particular distances from the nucleus. Finally, in the quantum model, the electrons are no
longer depicted as traveling around the nucleus in perfect circular orbits. Schrödinger’s wave equations
provided mathematical probabilities of electrons existing in 3-D regions within Bohr’s energy levels.
Dalton’s Billiard Ball Model:
Thomson’s Plum Pudding Model:
Rutherford’s Planetary Model:
Bohr’s Planetary Model with Orbits:
Quantum Model:
11. How was the nucleus of the atom discovered? The nucleus of the atom was discovered through
Rutherford’s gold foil experiment. In the experiment, alpha particles (helium nuclei) from a radioactive
source were aimed at a thin piece of gold foil. The entire apparatus was surrounded by a photographic
screen. The radioactive alpha particles would expose the photographic screen when they hit, so the
researchers could “see” where they particles were going. Since most of the alpha particles went straight,
Rutherford concluded that atoms are mostly empty space. Some particles were deflected at small angles
which led Rutherford to conclude that there was a positive charge within the gold foil because the alpha
particles were positive and the small deflection could be caused by positive-positive charge repulsion.
Surprisingly, a very few particles were deflected back at very wide angles. Rutherford concluded that this
occurred when the alpha particles were hitting some very dense matter. This dense region must be very
small since only about 1 in 8000 particles were deflected in this way. In the end, Rutherford concluded
that atoms contained a nucleus—a small, dense, positively charged are in the center of the atom.
12. What is the atomic theory? Who created it? How was it modified from the original theory? The
atomic theory was created by Dalton in 1803. It contained the following 5 postulates:
1. Elements are made of tiny particles called atoms.
2. The atoms of a given element are different from those of any other element; the atoms of different
elements can be distinguished from one another by their respective relative atomic weights.
3. All atoms of a given element are identical.
4. Atoms of one element can combine with atoms of other elements to form chemical compounds; a
given compound always has the same relative numbers of types of atoms.
5. Atoms cannot be created, divided into smaller particles, nor destroyed in the chemical process; a
chemical reaction simply changes the way atoms are grouped together.
We now know that atoms contain subatomic particles (protons, neutrons electrons) which are a
modification from postulate 5. We are also aware of the existence of isotopes which will have a
different atomic weight, but are in fact, the same element which is an amendment to postulate 3.
13. What is the Heisenberg Uncertainty Principle? Why is it thought to be true? The Heisenberg
Uncertainty Principle states that it is impossible to simultaneously know the position and velocity
(momentum) of an electron in an atom. This is thought to be true because detecting an electron is only
possible by bombarding it with another particle or a photon. Upon doing so, the particle immediately
changes position and speed.
14. What is meant by the term “block” in relation to electron configurations? How do these terms fit
together to elucidate our understanding of electron configurations? A “block” refers to sections of the
periodic table that have elements with similar endings to their electron configurations. The s block
consists of the first two columns on the periodic table, which makes sense because the s orbital can hold a
maximum of two electrons. The 6 columns at the right of the periodic table make up the p block which
can hold a maximum of six electrons (3 orbitals x 2 electrons each). The middle section of the periodic
table is referred to as the d block and contains 10 columns which corresponds to the 10 electrons the d
block can hold (5 orbitals x 2 electrons each). Finally, the f block is the two rows pulled out of the main
body of the periodic table. There are 14 columns in this section which corresponds to the maximum 14
electrons the f orbitals can hold (7 orbitals x 2 electrons each). Electron configurations can be written by
“reading” the periodic table from left to right like a book across the periods. Each period corresponds to
an energy level. Keep in mind that the d block is n-1 and the f block is n-2.