Download History of Atomic theory

SCH 4UI - History of Atomic Theory Review
Name: _________________________________
Marking Scheme:
A.
Date: _______________________________
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D
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Matching. Match each scientist with their contribution to Atomic Theory. Note that there are more
“contributions” than scientists. No “contribution” may be used more than once.
Scientists
Contribution
1. ____ Bohr
2. ____ Alchemists
3. ____ Thomson
4. ____ deBroglie
5. ____ Hiesenberg
6. ____ Rutherford
7. ____ Dalton
8. ____ Planck
9. ____ Democritus
A. His gold foil experiments proved that the atom is composed of a very
tiny, positively charged nucleus which contains most of the mass. Most
of the atom is filled with empty space occupied by electrons.
B. This is the first proposed model of the atom. Matter is composed of indestructible, indivisible atoms. Different elements are composed of different particles.
C. His model postulated that electrons were embedded in a positive sphere
like raisins in a bun. He stated that each element is characterized by the
number of electrons in the atom.
D. He used spectral emission lines to propose a theory that stated that electrons travel in orbits that correspond to specific energy levels.
E. Used wave equations to determine the energy states of matter. His theories led to the development of the secondary quantum number.
F. He determined that particles of matter (electrons) also exhibit wave
characteristics.
G. His uncertainty principle stated that we (the observer) can never exactly
know both the position and momentum of a particle (electron).
H. This group of scientist/philosophers generated the basics of modern
chemistry by developing lab equipment, chemical procedures and the
identification and use of many current substances.
I. He determined that particles emit energy in distinct “packets” of energy
known as quanta.
J. He determined that atoms cannot be created or destroyed (The Law of
Conservation of Matter). He also stated that atoms combine in specific
ratios when forming compounds.
10. ____ Schrodinger
B.
Modified True/False. Indicate whether the sentence or statement is true or false. If false, change the identified
word or phrase to make the sentence or statement true.
____ 1.
Bohr's atomic theory states that electrons may only possess specific amounts of energy.
____________________
____ 2.
The quantum mechanical model of the atom is capable of determining the exact location of an electron in
an atom. ____________________
____ 3.
All alkali earth metals have valence electrons in s orbitals when they are in their ground state.
____________________
____ 4.
Schrodinger's equation describes the electron as a wave. ____________________
____ 5.
The 2p orbitals in oxygen have three unpaired electrons. ____________________
____ 6.
The Pauli exclusion principle requires that two electrons in the same orbital have the same spin.
____________________
____ 7.
The electron configuration for sulfur, S is shown below. _________________________
3p
3s
2p
2s
1s
____ 8.
The valence p orbitals in phosphorus, P, are half-filled. ____________________
____ 9.
Hund's rule states that you must fill electrons into the lowest energy levels first. ____________________
____10.
A photon of light is equal to one quantum of energy. ____________________
C.
Multiple Choice. Identify the letter of the choice that best completes the statement or answers the question.
____ 11.
The 3p atomic orbital has the shape of
a. a sphere
d. two perpendicular dumb-bells
b. a torus
e. an egg
c. a dumb-bell
____ 12.
Rutherford's gold foil experiment showed that the atom is mostly empty space because
a. some of the alpha particles were reflected right back
b. some of the alpha particles were deflected
c. most of the alpha particles went straight through the foil
d. all of the alpha particles went straight through the foil
e. all of the alpha particles were deflected
____ 13.
The lines in the line spectrum of an atom results from
a. energy absorbed by electrons dropping back down to a lower energy level
b. energy absorbed by electrons jumping to a higher energy level
c. energy released by electrons jumping to a higher energy level
d. energy released by electrons dropping back down to a lower energy level
e. none of the above
____ 14.
Why do energy levels exist in atoms?
a. electrons are negatively charged
b. electrons are attracted to certain numbers of neutrons
c. electrons are able to possess any range of energy
d. electrons will only display certain colours
e. electrons are only able to possess quanta of energy
____ 15.
What did Heisenberg contribute to the quantum mechanical model of the atom?
a. the uncertainty principle
b. concept of quanta of energy
c. the idea that every mass has a wave with which it is associated
d. the wave equation
e. a relationship between energy and mass
____ 16.
What did Schrodinger contribute to the quantum mechanical model of the atom?
a. the uncertainty principle
b. concept of quanta of energy
c. the idea that every mass has a wave with which it is associated
d. the wave equation
e. a relationship between energy and mass
____ 17.
"A region of space in which there is a high probability of finding an electron" is the definition of
a. Orbital
d. photon
b. absorption spectrum
e. dipole
c. Quantum
____ 18.
"A packet of energy that can be absorbed or released by an electron" is a description of
a. Orbital
d. photon
b. absorption spectrum
e. dipole
c. Quantum
____
a.
b.
19.
Which of the following is the electron configuration for the valence shell of oxygen?
d.
e.
c.
____ 20.
Which of the following is the electron configuration for magnesium?
2 2
8
a. 1s 2s 2p
d. 1s32s32p33s2
3 3
4 2
b. 1s 2s 2p 3s
e. 1s22s22p63s2
2 2
7 1
c. 1s 2s 2p 3s
____ 21.
What experimental evidence led Bohr to believe that electrons can possess only specific amounts of
energy?
a. most alpha particles went straight through the gold foil
b. some alpha particles were deflected by the gold foil
c. the line spectra produced by excited atoms
d. atoms are electrically neutral
e. none of the above
____ 22.
What led Rutherford to believe that atoms contain a positive nucleus?
a. most alpha particles went straight through the gold foil
b. some alpha particles were deflected by the gold foil
c. the line spectra of excited atoms
d. atoms are electrically neutral
e. none of the above
____ 23.
Unlike Bohr's model of the atom, the quantum mechanical model of the atom treats the electron like
a. a tiny particle
d. a wave
b. a proton
e. a photon
c. a positive particle
____
s
24.
Which atoms could have the valence electron configuration shown below?
p
a. N3b. O2c. Ar
d. Cl1e. all of the above
____ 25.
a. 1s22s22p4
b. 1s21p6
c. 1s22s22p5
Which of the following is the electron configuration for fluoride, F1-?
d. 1s22s22p6
e. 1s22s22p63s1
D.
Short Answer
1.
What is the difference between an "orbit" as described in the Bohr-Rutherford model of the atom and an "orbital"
as described in the quantum mechanical model of the atom? (3)
2.
Write the long form and short hand electron configurations for gold and calcium. (4)
3.
Draw energy level diagrams for Scandium, Sc and Europium, Eu . (4)
4.
Sulfur can have the valence of +6, as in the sulfur atom in sulfate (SO42-), how is this possible? (3)
5.
What is the maximum number of electrons in the third energy level? (2)
6.
Propose an energy level diagram for the valence shell of Fe3+ and explain it. (3)
7.
Why is it possible for antimony to have both +3 and +5 charges? Explain using an energy level diagram. (3)
8.
What is the difference between paramagnetism and ferromagnetism? (3)
9.
Palladium and gold are examples of atoms with anomalous electron configurations. What are the predicted
electron configurations for these atoms? What are the actual electron configurations for these atoms? Why are the
actual electron configurations different from the predicted electron configurations? (5)