Electronic Structure Flashcards Preview

MCAT Gen. Chemistry - OLD > Electronic Structure > Flashcards

Flashcards in Electronic Structure Deck (38)
Loading flashcards...
1
Q

Briefly describe the Bohr Theory of the atom.

A

The Bohr Theory states:

  • Electrons can only exist in fixed orbits or energy levels located at specific distances from the nucleus
  • Any energy emitted or absorbed by an atom results from an electron jumping from one energy level to another
2
Q

In the Bohr Model, what does the hydrogen electron orbit?

A

The Bohr Model states that the hydrogen electron orbits the nucleus.

Note that all models assume that electrons orbit the nucleus. However, Bohr’s model is unique in that in most chemistry courses and on the MCAT, the Bohr model is usually restricted to the hydrogen atom only.

3
Q

In the quantum mechanical model, where does the hydrogen electron exist?

A

The electron is located somewhere in a spherical probability cloud around the nucleus, called the 1s orbital.

Note that the quantum mechanical model is the one used in most chemistry courses and on the MCAT.

4
Q

An atomic electron has not absorbed any energy. Which state is it in?

A

The atomic electron is in the ground state.

The ground state is the lowest possible energy orbital that any atomic electron may occupy.

5
Q

When a ground state hydrogen electron absorbs energy, what happens to it?

A

The hydrogen electron moves into an excited state.

For example, 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).

6
Q

What has to happen to an electron to move it from the ground state to an excited state?

A

The electron must absorb energy, typically in the form of a photon, to jump from the ground state to an excited state.

The frequency of the absorbed photon corresponds to the energy difference between the two states.

7
Q

What direction does energy flow when an atomic electron drops from the excited state back to the ground state?

A

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.

8
Q

Define:

absorption spectrum

A

The absorption spectrum of a specific material is the unique set of wavelengths of light absorbed by that substance or medium.

The absorption spectrum is typically displayed as a set of dark or “missing” lines in the spectrum, representing the absorbed wavelengths. This is shown in the third bar in the image below.

9
Q

Define:

emission spectrum

A

The emission spectrum is the unique spectrum of bright lines or bands of light emitted by a particular substance when it is electronically excited.

The emission spectrum is displayed as a set of light lines in the spectrum, representing the emitted wavelengths. This is shown by the second bar in the image below.

10
Q

How do a substance’s absorption and emission spectral lines compare to one another?

A

The absorption and emission spectral lines will overlap one another perfectly.

Both absorption and emission energy values are dependent on electrons moving between energy levels. The dark absorption line that represents jumping to a higher level should be in the exact same position as the bright emission line representing falling to a lower level; this is true because the same amount of energy is absorbed and emitted, respectively.

11
Q

What is the quantum number n called?

A

n is the principal quantum number, and is commonly referred to as the electron’s shell.

n can have any whole-number value greater than or equal to 1.

12
Q

As the principle quantum number n increases, what happens to the energy of the electron?

A

As n increases, energy increases.

Assume that the quantum number l stays constant unless told otherwise.

13
Q

What does the quantum number l represent?

A

l is the angular momentum (or azimuthal) number, and 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 MCAT will only explicitly test 0 to 3.

14
Q

In orbital theory, what do the letters s, p, d, and f indicate?

A

The letters s, p, d, and 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.

15
Q

What does the quantum number m or ml represent?

A

m or m<em>l</em> 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. For example, for an electron whose l = 1 (p subshell), m can equal -1, 0, or 1. These values correspond to the px, py, and pz orbitals, respectively.

16
Q

How many orbitals can be found in a p subshell?

A

A p subshell has three orbitals: px, py, and pz.

Remember that l = 1 for any p subshell. ml 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.

17
Q

What does the quantum number s* or ms *represent?

A

s or msis the spin quantum number, and 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.

18
Q

What is the value of l for any electron in an s orbital?

A

For any s electron, l = 0.

l, which can range from any value from 0 to n-1, determines the subshell of the electron. By definition, if l = 0 for an electron, that electron exists in an s orbital.

19
Q

What is the maximum number of electrons found in an orbital?

A

Each orbital can hold up to 2 electrons.

Note that when one orbital hold two electrons simultaneously, their spins must oppose each other.

20
Q

With 5 orbitals, how many electrons can a d subshell hold?

A

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.

21
Q

What are the geometric shapes of s, p, and d orbitals, respectively?

A
  • s = “spherical”
  • p = “peanut”
  • d = “donut”
22
Q

How many electrons exist in a filled shell with principal quantum number n?

A

There are 2n2 electrons in the filled shell.

For example, for the n = 2 shell:

2(22) = 8

This shell has 4 orbitals: 2s, 2px, 2py, 2pz. Each of those can hold 2 electrons, for a total of 8 in the shell.

23
Q

How many orbitals exist per shell with respect to the principal quantum number n?

A

A shell will have n2 orbitals.

For example, the shell n = 2 has 22 = 4 orbitals total. They are the 2s, 2px, 2py, and 2pz orbitals.

24
Q

How many orbitals are there per subshell with azimuthal quantum number l?

A

A subshell will have 2(l) + 1 orbitals.

For example, for a d subshell (l = 2) there are 2(2) + 1 = 5 orbitals total. They are the dxy, dxz, dyz, dx2-y2, and dz2 orbitals.

25
Q

How many electrons can be found in the following subshells?

  • s subshells
  • p subshells
  • d subshells
  • f subshells
A
  • An s subshell holds 1 x 2 = 2 electrons
  • A p subshell holds 3 x 2 = 6 electrons
  • A d subshell holds 5 x 2 = 10 electrons
  • An f subshell holds 7 x 2 = 14 electrons
26
Q

Given two specific subshells, what determines which one will fill with electrons first?

A

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 #.

For example, 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. in the case of a tie between two subshells with the same n+l value, the one with the lower n will fill first.

27
Q

Arrange the following subshells in terms of increasing energy:

4s, 6s, 3d, 2s, 4f

A

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

28
Q

For a given value of n, please rank the following subshells in order of increasing energy:

p, s, f, d

A

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.

29
Q

Explain each of the 3 terms for the spectroscopic notation for an atom’s electronic structure.

For example, the 3d5 depiction of chromium’s valence electrons.

A

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 3d5 indicates that there are 5 valence electrons, with n = 3 and l = 2.

30
Q

Give both the full and condensed form of the spectroscopic notation for calcium (Ca).

A

The full spectroscopic notation of Ca is:

1s22s22p63s23p64s2

In condensed notation:

[Ar] 4s2

Since every shell up to 3p6 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 3p6.

31
Q

What does the Aufbau Principle state?

A

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. For example, the 1s subshell fills first, then 2s, then 2p, and so on.

32
Q

What does Hund’s rule state?

A

Hund’s rule describes the order of adding electrons to an unfilled subshell.

Hund’s rule explains that when electrons are added to a subshell that has more than 1 orbital (p, d, or f), each orbital first receives a single electron, each with parallel spins, until each orbital in the subshell has one electron contained within it.

Only once the subshell is half full will spin-down electrons be added, one per orbital, until the subshell is completely filled.

33
Q

What does the Pauli Exclusion Principle state?

A

The Pauli Exclusion Principle states that two electrons in the same orbital must be of different spins.

As a result of this rule, no two electrons in the same atom will ever have exactly the same 4 quantum numbers (n, l, ml, ms).

34
Q

What types of electronic configurations lead to particularly stable atoms?

A

Fully-filled and half-filled subshells make atoms particularly stable.

In particular, atoms with p3, p6, d5, and d10 valence shell configurations are especially stable.

A classic MCAT question will ask about exceptions to typical stability trends, which generally remove one electron from the s subshell to give the atom a half-filled or fully-filled d subshell.

35
Q

Use electronic structure to explain the difference in stability between atomic phosphorus (P) and sulfur (S).

A

Phosphorus’ valence shell configuration is 3p3, while sulfur’s is 3p4.The p-block electrons around phosphorus will therefore be more stable, as the p subshell is half-filled with parallel-spin electrons, a particularly stable configuration.

Sulfur will be slightly less stable, as its p subshell contains one additional electron paired in an orbital with an anti-parallel spin electron.

36
Q

Define:

effective nuclear charge (Zeff)

A

The effective nuclear charge (Zeff) is the amount of attraction that an electron in the outermost subshell has towards the positively-charged nucleus.

This number can be calculated by finding the positive charge of the nucleus and subtracting the total number of shielding electrons in the inner, fully-filled subshells.

The higher the Zeff, the more strongly the outer electrons in unfilled subshells are bound to the atom.

37
Q

What is the effective nuclear charge (Zeff) for atomic Na?

A

Zeff(Na) = +1.

Na has 11 total protons in its nucleus,and 10 total electrons in inner subshells (1s2,2s2, 2p6), so:

Zeff(Na) = +11 - 10 = +1

38
Q

What is the effective nuclear charge (Zeff) for atomic S?

A

Zeff(S) = +6.

S has 16 total protons in its nucleus and 10 total electrons in inner subshells (1s2,2s2, 2p6), so:

Zeff(S) = +16 - 10 = +6