Atomic Theory and Structure

Question 1

Briefly describe the Bohr Theory of the atom.

Answer 1

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.

Question 2

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

Answer 2

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.

Question 3

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

Answer 3

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.

Question 4

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

Answer 4

The atomic electron is in the ground state.

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

Question 5

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

Answer 5

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.

Question 6

What has to happen to an electron in order for it to change from the ground state to an excited state?

Answer 6

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

Question 7

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

Answer 7

Energy flows out of the atom when an electron drops to the ground state.

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.

Question 8

Define:

absorption spectrum

Answer 8

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.

Question 9

Define:

emission spectrum

Answer 9

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.

Question 10

How does substance absorption and emission spectral lines compare to one another?

Answer 10

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.

Question 11

What is the quantum number n called?

Answer 11

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.

Question 12

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

Answer 12

As n increases, energy increases.

Remember: assume that the quantum number l stays constant unless told otherwise.

Question 13

1. What is the quantum number l called?
2. What does it represent?

Answer 13

1. l is the angular momentum (or azimuthal) number.
2. 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.

Question 14

1. In orbital theory, what do s, p, d, and f indicate?
2. How are these values determined?

Answer 14

1. The letters s, p, d, f symbolize the subshells in which an electron can exist.
2. 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.

Question 15

1. What is the quantum number m or m_l called?
2. What does it represent?

Answer 15

1. m or m_l is the magnetic quantum number.
2. 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.

Question 16

How many orbitals can be found in a p subshell?

Answer 16

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.

Question 17

1. What is the quantum number s or m_s called?
2. What does it represent?

Answer 17

1. s or m_s is the spin quantum number.
2. 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.

Question 18

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

Answer 18

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.

Question 19

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

Answer 19

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.

Question 20

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

Answer 20

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.

Question 21

What are the geometric shapes of the following orbitals?

• s orbitals
• p orbitals
• d orbitals

Answer 21

• s = ‘spherical’
• p = ‘peanut’
• d = ‘donut’

Question 22

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

Answer 22

There are 2n² electrons in the filled shell.

Ex: For the n = 2 shell: 2(2²) = 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.

Question 23

How many orbitals are there per shell with principal quantum number n?

Answer 23

A shell will have n² orbitals.

Ex: For the shell n = 2 there are 2² = 4 orbitals total. They are the 2s, 2p_x, 2p_y, and 2p_z orbitals.

Question 24

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

Answer 24

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_x²_-y², and d_z² orbitals.

Question 25

How many electrons can be found in the following subshells?

• s subshells
• p subshells
• d subshells
• f subshells

Answer 25

• 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

Question 26

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

Answer 26

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.

Question 27

Arrange the following subshells in terms of increasing energy:

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

Answer 27

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

Question 28

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

p, s, f, d

Answer 28

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.

Question 29

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

Ex: the 3d⁵ depiction of Chromium’s valence electrons.

Answer 29

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

Question 30

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

Answer 30

The full electronic structure of Calcium is:
1s²2s²2p⁶3s²3p⁶4s²

In condensed notation:
[Ar] 4s²

Since every 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.

Question 31

Define:

the Aufbau Principle

Answer 31

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.

Question 32

Define:

Hund’s Rules

Answer 32

Hund’s Rules describe the order of adding electrons to an unfilled subshell.

Hund’s Rules explain 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 spin-up, until each orbital in the subshell has one electron contained within it.

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

Question 33

Define:

the Pauli Exclusion Principle

Answer 33

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

The result of this rule is that two electrons in the same atom will never have exactly the same 4 quantum numbers (n, l, m_l, m_s).

Question 34

What types of electronic configurations lead to particularly stable atoms?

Answer 34

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 question will ask about exceptions to typical stability trends, and the answer will be due to an atom having a half-filled subshell.

Question 35

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

Answer 35

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.

Question 36

Define and explain how to calculate:

effective nuclear charge Z_eff

Answer 36

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

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

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

Question 37

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

Answer 37

Z_eff (Na) = +1.

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

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

Question 38

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

Answer 38

Z_eff (S) = +6.

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

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

Question 39

Where are the alkali metals located on the Periodic Table?

Answer 39

Alkali metals make up the first column of the Periodic Table (Group IA).

Question 40

1. What is the valence shell configuration of all alkali metals?
2. What oxidation state do they ionize to?

Answer 40

1. All alkali metals have an s¹ valence shell configuration.
2. Alkali metals are relatively electropositive, so will lose 1 valence electron and form a +1 oxidation state.

Question 41

Where are the alkaline earth metals located on the Periodic Table?

Answer 41

Alkaline earth metals make up the second column of the Periodic Table (Group IIA).

Question 42

1. What is the valence shell configuration of all alkaline earth metals?
2. What oxidation state do they ionize to?

Answer 42

1. All alkaline earth metals have an s² valence shell configuration.
2. They are relatively electropositive, so will lose 2 valence electrons and form a +2 oxidation state.

Question 43

Where are the halogens located on the Periodic Table?

Answer 43

The halogens make up the fifth column of the p block (Group VIIA).

Question 44

1. What is the valence shell configuration of all halogens?
2. What oxidation state do they ionize to?

Answer 44

1. All halogens have an s²p⁵ valence shell configuration.
2. They are quite electronegative, so they will accept one additional valence electron to take on a -1 oxidation state.

Question 45

Where are the noble gases located on the Periodic Table?

Answer 45

The noble gases make up the sixth column of the p block (Group VIIIA).

Question 46

1. What is the valence shell configuration of all noble gases?
2. What oxidation state do they ionize to?

Answer 46

1. All noble gases have an s²p⁶ valence shell configuration.
2. Trick question! Since they already have a completely filled octet, noble gases do not ionize, and they typically exist in the 0 oxidation state as free particles.

Exception: Kr and Xe, being below the 3rd row, can exceed their octet and make coordinate covalently bonded compounds such as XeF₆.

Question 47

What is the oxygen group, and where is it located on the Periodic Table?

Answer 47

The oxygen group is the group (column) below oxygen on the Periodic Table.

It includes elements such as S and Se that are chemically similar to oxygen.

Question 48

Where are the transition metals located on the Periodic Table?

Answer 48

The transition metals make up the entire d block.

Question 49

Why do transition metals have high conductivity?

Answer 49

Transition metals have high conductivity due to their unfilled d subshells.

d electrons, by their nature, are loosely bound to the atom. As such, elements with partially-filled d subshells can be thought of as nuclei floating in a sea of unattached electrons, prime conditions for electrical conductivity.

Question 50

What are the representative elements, and where are they located on the Periodic Table?

Answer 50

Representative elements are the most common elements in the solar system and the universe.

They are found in the s block and the p block of the Periodic Table.

By standard nomenclature, these are groups IA, IIA, IIIA, IVA, VA, VIA, VIIA, and VIIIA.

Question 51

What is the valence subshell for the elements in the first two columns of the periodic table?

Answer 51

The elements of the first two columns have a valence s subshell.

Group IA has an s¹ valence configuration, while IIA is s².

Note: helium also has a valence s subshell, but is typically listed on the farthest column with the noble gases, as it is chemically more similar to them than the alkaline earth metals.

Question 52

What is the valence subshell for the elements in the last six columns of the periodic table?

Answer 52

The elements of the last six columns have a valence p subshell.

Ex: Group IIIA has an s²p¹valence configuration, while VIIIA is s²p⁶.

Note: Although helium is typically listed on the farthest column with the noble gases in VIIIA, it actually has a valence s subshell.

Question 53

Describe the properties of metals in terms of their:

• position in the periodic table
• electronegativity
• preferred oxidation state

Answer 53

Metals are generally:

• found in the lower-left areas of the periodic table.
• low in electronegativity, losing electron density when bonded to nonmetals.
• found in positive oxidation states when in compounds.

Question 54

What are the main physical properties of metals?

Answer 54

Metals generally are/have:

• good conductors of heat and electricity.
• malleable, ductile, lustrous, and dense solids at room temp.
• fairly high melting and boiling points.

Question 55

Describe the properties of nonmetals in terms of:

• position in the periodic table
• electronegativity
• preferred oxidation state

Answer 55

Nonmetals are generally:

• found in the upper-right areas of the periodic table.
• high in electronegativity, gaining electron density when bonded to metals.
• found in negative oxidation states when in compounds.

Question 56

What are the main physical properties of nonmetals?

Answer 56

Nonmetals are/have:

• poor conductors of heat and electricity.
• dull and brittle if they form solids at room temperature.
• significantly lower melting and boiling points than metals (carbon is the primary exception).

Question 57

How many valence electrons does oxygen, element 8, have?

Answer 57

Oxygen has 6 valence electrons.

Oxygen is the 6th element in its row. Its valence shell configuration is 2s²2p⁴, for a total of 6 valence electrons.

Question 58

Define:

first ionization energy

Answer 58

The energy required to remove one valence electron from an atom in the gas phase.

The generic ionization energy equation is: X(g) ⇒ X⁺(g) + e^-

Question 59

Describe the general trend of ionization energy heading across a row of the Periodic Table.

Answer 59

Ionization energy increases from left to right across a row of the Periodic Table.

Other notes about ionization energy:

Atoms with fully-filled subshell will have high ionization energies.
Atoms with half-filled subshells will have higher ionization energies than their neighbors.
The alkali and alkaline earth metals have very low ionization energies.

Question 60

Which has a higher first ionization energy, Cl or Br?

Answer 60

Cl

Remember: ionization energy decreases going down a column. Br is below Cl in the halogen column.

Question 61

Describe the general trend of ionization energy heading down a column of the Periodic Table.

Answer 61

Ionization energy decreases heading down a column of the Periodic Table.

The further down a column an element lies, the easier to remove. These atoms have higher n values for their valence electrons. Higher n electrons sit further from the atomic nucleus, and are therefore less strongly bound to the nucleus.

Question 62

Which has a higher first ionization energy, Si or P?

Answer 62

P

Remember: ionization energy increases from left-to-right across a column. P is to the right of Si in period 3.

Question 63

What is an atom’s second ionization energy?

Answer 63

The second ionization energy is the energy required to remove a subsequent (second) valence electron from a singly-charged ion in the gas phase.

The generic second ionization energy equation is: X⁺(g) ⇒ X²⁺(g) + e^-

Question 64

What are the relative magnitudes of any atom’s first and second ionization energies?

Answer 64

The second ionization energy is always larger in magnitude than the first ionization energy for every atom.

The removal of the first electron reduces the electron-electron repulsion energy of the molecule, allowing the positive nucleus to attract the remaining electrons more strongly, increasing the energy needed to remove subsequent electrons.

Question 65

How does atomic radius vary as atomic shell increases down a column of the Periodic Table?

Answer 65

Atomic radius increases down a column of the Periodic Table.

Each increasing shell can be thought of as another “layer” of electrons, outside the previous layer, increasing the atom’s size.

Note: The highlighted elements represent one column of the Periodic Table.

Question 66

Which has a larger radius, Cl or Br?

Answer 66

Br

Remember: atomic radius increases going down a column, and Br is below Cl in the halogens column. The valence electrons from Br are in the n=4 shell, those from Cl are in the n=3 shell.

Question 67

How does atomic radius vary as atomic number increases across a row of the Periodic Table?

Answer 67

Atomic radius decreases across a row from left to right on the Periodic Table.

This is due to effective nuclear charge. As Z_eff increases, the nucleus binds the electrons more tightly, pulling them in closer.

Note: the highlighted elements represent one full row of the Periodic Table.

Question 68

Which has a larger radius, Si or P?

Answer 68

Si

Remember: Atomic radius decreases from left-to-right across a column, and P is to the right of Si in the Period 3.

Question 69

Define:

electron affinity

Answer 69

The energy released when one valence electron is added to an atom in the gas phase.

The generic electron affinity equation is: X(g) + e^- ⇒ X^-(g)

Question 70

Describe the general trend of electron affinity heading across a row of the Periodic Table.

Answer 70

Electron affinity increases from left to right across a row of the Periodic Table.

The smaller an atom is, the closer a newly-added valence electron gets to the positively-charged nucleus, and the more energy released when that electron is added.

Question 71

Describe the general trend of electron affinity heading down a column of the Periodic Table.

Answer 71

Electron affinity decreases heading down a column of the Periodic Table.

The further down a column an element lies, the higher the value of n for its valence electrons. Higher n electrons sit further from the atomic nucleus, and so are less bound to the nucleus and release less energy when added.

Question 72

Define:

electronegativity

Answer 72

An atom’s electronegativity describes that atom’s tendency to attract electron density towards itself through a chemical bond.

Note: electronegativity only applies to atoms in a bond. There is no such concept as the electronegativity of a bare atomic species.

Question 73

Describe the general trend of electronegativity heading across a row of the Periodic Table.

Answer 73

Electronegativity increases from left to right across a row of the Periodic Table.

The smaller an atom is, the closer the electrons in its bonds get to the positively-charged nucleus, and the more strongly the electrons are attracted to the nucleus.

Question 74

Describe the general trend of electronegativity heading down a column of the Periodic Table.

Answer 74

Electronegativity decreases heading down a column of the Periodic Table.

The further down a column an element lies, the larger its radius. This puts more space between the positively-charged nucleus and the electrons in any bonds it makes, reducing the attraction between the nucleus and the electrons.

Question 75

Which element has the highest electronegativity?

Answer 75

Fluorine

Electronegativity increases the further right and the closer to the top of the Periodic Table an element is. Fluorine, which is the top-right element that isn’t a noble gas, therefore has the highest value of electronegativity.

On the Pauling scale, the most commonly-used scale for determining electronegativity values, fluorine has the highest value possible: 4.0.

Question 76

What class of elements have the highest values of electronegativity?

Answer 76

Nonmetals are the most electronegative elements.

Fluorine is the element with the highest electronegativity, and as a general rule, the closer an element is to fluorine, the higher its electronegativity.

Some other notes about electronegativity:

The halogens are the most electronegative group.
Noble gases capable of making bonds (Xe and Kr) are relatively strongly electronegative.
Metals, particularly the alkali and alkali earth metals, are generally electropositive (very low electronegativity).

Question 77

Which has a higher electronegativity, Cl or Br?

Answer 77

Cl

Remember: electronegativity decreases going down a column, and Br is below Cl in the halogens column.

Question 78

Which has a higher electronegativity, Si or P?

Answer 78

P

Remember: electronegativity increases from left-to-right across a column, and P is to the right of Si in period 3.

Question 79

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Answer 79

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