SCH4U - Exam Review Package
Atomic Structure Review
1. Describe the contributions of the following scientists to our current understanding of the
atomic structure and quantum theory:
a. Bohr Planetary Model:
- Energy orbits the nucleus in shells
- Electrons travel within energy level without losing or
gaining energy
- ↑ distance from nucleus = ↑ energy
b. Heisenberg Uncertainty Principle:
- Due to its wave and particle nature, we can never know
the position and speed of an electron
- Velocity = wave nature
- Position = particle nature
c. Schrodinger Quantum Model:
- “Cloud of probability” illustrates the possibilities of
where the electron may lie as their distinct position is
unknown
d. Pauli Pauli’s Exclusion Principle:
- Electrons sharing an orbital must have opposite spins
e. Rutherford Nuclear Model: Gold Foil Experiment
- Alpha particles were fired at thin metal foils though a
small portion of it was deflected back
- Result: dense nucleus surrounded by a cloud of negative
charge
f. Thomson Plum Pudding Model:
- Cathode rays were deflected by charged objects = atoms
are mostly made of negatively charged particles
- Negative subatomic particles (electrons) were scattered
around a positive cloud within atoms
2. Based on the quantum mechanical model of the atom:
a. What is an “orbital”
i. Orbital: probability distribution maps for electrons (represented through
wave function)
b. What is the maximum electron carrying capacity of the following energy levels:
i. Equation: [x2(2)]
, n = 1 (2 electrons)
n = 2 (8 electrons)
n = 3 (18 electrons)
n = 4 (32 electrons)
n = 5 (50 electrons)
c. What orbital subshells would be found in the following energy levels:
n = 1 (s subshell)
n = 2 (p subshell)
n = 3 (d subshell)
n = 4 (f subshell)
n = 5 (g subshell)
d. How would the electrons be distributed amongst the orbital sublevels in n = 5 (ex:
number of e- per subshell)
Remember: 4 quantum numbers
- n = principal quantum number (size and energy)
- l = angular momentum quantum number (shape)
- Subshells within the main energy level
- Values (n - 1) = 4 subshells
- ml = magnetic quantum number (orientation of orbitals)
- Ranging from integers -5 to 5
- Total 25 orbitals
- ms = spin quantum number (direction of electron spin)
- Since there are 25 orbitals, there are 50 electrons
3. Electron configuration:
a. Write the ground state electron configuration of a neutral atom of nitrogen
i. Atomic number 7
ii. 1s22s22p3
b. Assign the quantum numbers to each of the electrons in your nitrogen atom
i. n=2
ii. l=1
iii. ml = -1, 0, 1
iv. ms = +½
c. Write the ground state complete electron configuration for iron, Fe.
i. Atomic number 26
ii. 1s22s22p63s23p63d64s2
d. Draw an orbital energy level diagram for titanium, Ti.
i. INCOMPELTE
Atomic Structure Review
1. Describe the contributions of the following scientists to our current understanding of the
atomic structure and quantum theory:
a. Bohr Planetary Model:
- Energy orbits the nucleus in shells
- Electrons travel within energy level without losing or
gaining energy
- ↑ distance from nucleus = ↑ energy
b. Heisenberg Uncertainty Principle:
- Due to its wave and particle nature, we can never know
the position and speed of an electron
- Velocity = wave nature
- Position = particle nature
c. Schrodinger Quantum Model:
- “Cloud of probability” illustrates the possibilities of
where the electron may lie as their distinct position is
unknown
d. Pauli Pauli’s Exclusion Principle:
- Electrons sharing an orbital must have opposite spins
e. Rutherford Nuclear Model: Gold Foil Experiment
- Alpha particles were fired at thin metal foils though a
small portion of it was deflected back
- Result: dense nucleus surrounded by a cloud of negative
charge
f. Thomson Plum Pudding Model:
- Cathode rays were deflected by charged objects = atoms
are mostly made of negatively charged particles
- Negative subatomic particles (electrons) were scattered
around a positive cloud within atoms
2. Based on the quantum mechanical model of the atom:
a. What is an “orbital”
i. Orbital: probability distribution maps for electrons (represented through
wave function)
b. What is the maximum electron carrying capacity of the following energy levels:
i. Equation: [x2(2)]
, n = 1 (2 electrons)
n = 2 (8 electrons)
n = 3 (18 electrons)
n = 4 (32 electrons)
n = 5 (50 electrons)
c. What orbital subshells would be found in the following energy levels:
n = 1 (s subshell)
n = 2 (p subshell)
n = 3 (d subshell)
n = 4 (f subshell)
n = 5 (g subshell)
d. How would the electrons be distributed amongst the orbital sublevels in n = 5 (ex:
number of e- per subshell)
Remember: 4 quantum numbers
- n = principal quantum number (size and energy)
- l = angular momentum quantum number (shape)
- Subshells within the main energy level
- Values (n - 1) = 4 subshells
- ml = magnetic quantum number (orientation of orbitals)
- Ranging from integers -5 to 5
- Total 25 orbitals
- ms = spin quantum number (direction of electron spin)
- Since there are 25 orbitals, there are 50 electrons
3. Electron configuration:
a. Write the ground state electron configuration of a neutral atom of nitrogen
i. Atomic number 7
ii. 1s22s22p3
b. Assign the quantum numbers to each of the electrons in your nitrogen atom
i. n=2
ii. l=1
iii. ml = -1, 0, 1
iv. ms = +½
c. Write the ground state complete electron configuration for iron, Fe.
i. Atomic number 26
ii. 1s22s22p63s23p63d64s2
d. Draw an orbital energy level diagram for titanium, Ti.
i. INCOMPELTE
