Atomic Theory and Structure Flashcards
From quantum numbers to electronic orbitals, use these cards to master the topic of atomic theory as tested in most introductory undergrad chemistry courses and even on the AP Chemistry exam.
Briefly describe the Bohr Theory of the atom.
The Bohr Theory states:
- Electrons can only exist in fixed orbits or energy levels.
- These energy levels are at specific distances from the nucleus.
- Any energy emitted/absorbed from/by an atom will be the result of an electron jumping from one energy level to another.
In the Bohr Model, what does the hydrogen electron orbit?
The Bohr Model states that the hydrogen electron orbits the nucleus.
Note: all models assume that electrons orbit the nucleus, but Bohr’s model is unique in that in most chemistry courses and on the AP Chemistry exam, the Bohr model is usually restricted to the hydrogen atom only.
In the quantum mechanical model, where does the hydrogen electron exist?
In a spherical probability cloud around the nucleus, called the 1s orbital.
Note: the quantum mechanical model is the one used in most chemistry courses and on the AP Chem exam.
An atomic electron has not absorbed any energy. Which state is it in?
The atomic electron is in the ground state.
The ground state is the lowest possible energy orbital that any atomic electron may occupy.
When a ground state hydrogen electron absorbs energy, what happens to it?
The hydrogen electron moves into an excited state.
Ex: a ground-state electron in hydrogen is in the 1s state. If it absorbs the right amount of energy, it can jump into the 3p state, which is excited (higher) in energy than the ground state.
What has to happen to an electron in order for it to change from the ground state to an excited state?
The electron must absorb energy, typically in the form of a photon, to go from the ground state to an excited state.
What direction does energy flow when an atomic electron drops from the excited state back to the ground state?
Energy is released from the atom.
Since the ground state is lower in energy than the excited state, the change from excited to ground is always accompanied by a release of energy from the atom.
Define:
absorption spectrum
The absorption spectrum is the unique set of wavelengths of light absorbed by a specific substance or medium.
The absorption spectrum is typically displayed as a set of dark lines (or missing lines) in the spectrum, representing the absorbed wavelengths. This is the third bar in the image.
Define:
emission spectrum
The emission spectrum is the unique spectrum of bright lines or bands of light emitted by a particular substance when it is electronically excited.
This is the second bar in the image.
How does substance absorption and emission spectral lines compare to one another?
The absorption and emission spectral lines will overlap with one another perfectly.
Both absorption and emission energy values are dependent on electrons moving between energy levels. Jumping to a higher level (dark absorption line) should be in the exact same position as jumping to a lower level (bright emission line) since it’s the exact same amount of energy absorbed and emitted, respectively.
What is the quantum number n called?
n is the principal quantum number, and is commonly referred to as the shell the electron is in.
n can have any whole number value greater than or equal to 1.
As the principle quantum number n increases, what happens to the energy?
As n increases, energy increases.
Remember: assume that the quantum number l stays constant unless told otherwise.
- What is the quantum number l called?
- What does it represent?
- l is the angular momentum (or azimuthal) number.
- It represents an electron’s subshell.
If l = 0, the electron is in an s subshell.
If l = 1, the electron is in a p subshell.
If l = 2, the electron is in a d subshell.
If l = 3, the electron is in a f subshell.
l can take any integer value from 0 to n - 1, but most chemistry courses and the AP Chem exam will only explicitly test 0 to 3.
- In orbital theory, what do s, p, d, and f indicate?
- How are these values determined?
- The letters s, p, d, f symbolize the subshells in which an electron can exist.
- The value of the quantum number l determines the subshell. s, p, d, and f subshells correspond to l = 0, 1, 2, and 3, respectively.
- What is the quantum number m or m_{l} called?
- What does it represent?
- m or m_{l} is the magnetic quantum number.
- It represents the orbital in which an electron exists.
m can hold any integer value between -l and +l, including 0.
Ex: for an electron whose l = 1 (p subshell), m can equal -1, 0, or 1. These values correspond to the p_{x}, p_{y}, and p_{z} orbitals, respectively.
How many orbitals can be found in a p subshell?
A p subshell has three orbitals: p_{x}, p_{y}, and p_{z}.
Remember: l = 1 for any p subshell. m_{l} can range from -l to l (in this case: -1, 0, or 1) in a p subshell. These values correspond to the x, y, and z orbitals.
- What is the quantum number s or m_{s} called?
- What does it represent?
- s or m_{s} is the spin quantum number.
- It represents the spin direction of an electron.
s can have exactly one of two values, +1/2 and -1/2, corresponding to spin-up and spin-down. These two values are inherently equal in energy.
What is the value of l for any electron in an s orbital?
For any s electron, l = 0.
l can range from any value from 0 to n-1, and determines the subshell. By definition, if l = 0 for an electron, that electron exists in an s orbital.
What is the maximum number of electrons found in an orbital?
Each orbital can hold up to 2 electrons.
Note: When one orbital hold two electrons simultaneously, one must be spin-up and the other spin-down.
With 5 orbitals, how many electrons can a d subshell hold?
A d subshell holds up to 10 electrons.
Each of the 5 orbitals can have 1 spin-up electron and 1 spin-down, for a total of 2(5)=10 total.
What are the geometric shapes of the following orbitals?
- s orbitals
- p orbitals
- d orbitals
- s = ‘spherical’
- p = ‘peanut’
- d = ‘donut’
How many electrons are there in a filled shell with principal quantum number n?
There are 2n^{2} electrons in the filled shell.
Ex: For the n = 2 shell: 2(2^{2}) = 8
This shell has 4 orbitals: 2s, 2p_{x}, 2p_{y}, 2p_{z}. Each of those can hold 2 electrons, for a total of 8 in the shell.
How many orbitals are there per shell with principal quantum number n?
A shell will have n^{2} orbitals.
Ex: For the shell n = 2 there are 2^{2} = 4 orbitals total. They are the 2s, 2p_{x}, 2p_{y}, and 2p_{z} orbitals.
How many orbitals are there per subshell with azimuthal quantum number l?
A subshell will have 2(l) + 1 orbitals.
Ex: For a d subshell, l = 2, and there are 2(2) + 1 = 5 orbitals total. They are the d_{xy}, d_{xz}, d_{yz}, d_{x}^{2}_{-y}^{2}, and d_{z}^{2} orbitals.
How many electrons can be found in the following subshells?
- s subshells
- p subshells
- d subshells
- f subshells
- An s subshell holds 1x2=2 electrons
- A p subshell holds 3x2=6 electrons
- A d subshell holds 5x2=10 electrons
- An f subshell holds 7x2=14 electrons
Given two specific subshells, what determines which one will fill with electrons first?
The subshell with the lowest total energy will fill first.
Total energy can be approximated as E=n+l, where n=principal quantum # and l=azimuthal quantum #.
Ex: For a 4s subshell, n = 4 and l = 0, so n+l = 4. So a 4s subshell will fill before a 3d subshell, which has n = 3, l = 2, and n+l = 5.
Note: in the case of a tie between two subshells with the same n+l value, the one with the lower n will fill first.
Arrange the following subshells in terms of increasing energy:
4s, 6s, 3d, 2s, 4f
In order of increasing energy:
2s < 4s < 3d < 6s < 4f
The total order of all relevant subshells in the full Periodic Table is:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d
For a given value of n, please rank the following subshells in order of increasing energy:
p, s, f, d
In order of increasing energy:
s < p < d < f
These subshells differ in their value of l. For a given n, the higher the l, the higher the energy.
Explain each of the 3 terms for the spectroscopic notation for an atom’s electronic structure?
Ex: the 3d^{5} depiction of Chromium’s valence electrons.
The spectroscopic notation denotes the three most important pieces of information about a subshell: its energy level (n), subshell (l), and the total number of electrons it contains.
So Chromium’s 3d^{5} explains that there are 5 valence electrons, with n = 3 and l = 2.
Give both the full and condensed form of the spectroscopic notation for Calcium (Ca).
The full electronic structure of Calcium is:
1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}
In condensed notation:
[Ar] 4s^{2}
Since every up to 3p^{6} is completely filled, they are not chemically relevant - only valence electrons participate in chemical reactions. Therefore, they can all be abbreviated as the noble gas from the previous row, in this case Ar, which represents the element with fully-filled subshells up to 3p^{6}.
Define:
the Aufbau Principle
The Aufbau Principle describes the order in which subshells are filled with electrons as atomic number increases. Aufbau is German for ‘Building Up’.
Shells/subshells of lower energy get filled with electrons before higher energy shells/subshells.
Ex: The 1s subshell fills first, then 2s, then 2p, and so on.