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Flashcards in electronic_structure Deck (41):
1

Describe the relationship of wavelength, frequency and energy for electromagnetic radiation (light).

Energy is directly proportional to frequency and inversely proportional to wavelength. These relationships can be described by Plank's equation: E = hv = hc/lambda

2

In what region of the electromagnetic spectrum (in terms of wavelength) do we find the infrared, visible, and ultraviolet regions.

Infrared is considered to be above 700 nm, visible between 700-400 nm, and finally ultraviolet below 400 nm.

3

Describe the Rutherford model of the hydrogen atom circa 1910.

Based on the experimental evidence of the time (including Rutherford's goldfoil experiments), the hydrogen atom was understood to consist positively charged particle (designated as the nucleus) which is surrounded by an electron in motion at a relatively large distance from the nucleus. Evidence suggested that most of the mass (99.9%) was associated with the nucleus. From this model, it was inferred that strangely, the atom and therefore matter consisted of mainly empty space.

4

According to classical physics, what is the major flaw of the Rutherford model, which renders the model incomplete?

According to classical physics, the moving negatively charged electron should collapse into the positively charged nucleus. That is, it was not understood what kept the electron in its position away from the nucleus.

5

Describe briefly the main features of the Bohr model of the hydrogen atom.

Bohr proposed that the electron exists in discrete energy levels, n, each with defined energies.

These energy levels are said to be quantized, that is, the electron cannot exist "in between" the discrete energy levels. The electron exists predominantly in the lowest energy level (n=1), called "ground state".

If the atom is provided with an appropriate amount of energy, the electron will be "excited" to a higher energy level. An electron in higher energy level is unstable, and the electron will drop to lower energy levels and ultimately to ground state. As the electron returns to ground state, energy is lost by the atom as a photon, with an energy that exactly matches the energy difference between the higher and lower energy levels.

6

Define ground state.

In the Bohr model this is the lowest and most stable energy state for the electron. The electron exists predominantly in this state.

7

Define excited state.

An atom is in an excited state when it has absorbed energy which has allowed the electron to exist temporarily in one of the permitted higher energy states.

8

Define emission spectrum.

Discrete lines at designated wavelengths (in the UV, visible and IR) which result from electrons "falling" from higher to lower energy levels. Each element will have its own emission spectrum, which a reflection of the energy difference between excited states and ground state.

By contrast, a continuous spectrum shows a continuous set of "lines" at all wavelengths.

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9

Define absorption spectrum.

Discrete blanks in a continuous spectrum at designated wavelengths (in the UV, visible and IR) which result from electrons "jumping" from lower to higher energy levels. Each element will have its own absorption spectrum, which a reflection of the energy difference between excited states and ground state. Absorption spectrums are useful in identifying elements in stars.

10

What experimental evidence supported the concept of quantized energy states in the hydrogen atom?

An important piece of experimental evidence is the hydrogen emission spectrum. For example four discrete lines are easily discerned in the visible portion of the hydrogen emission spectrum. The wavelength of these lines correspond to energies which exactly match the expected energy difference (as calculated by Bohr) if an electron was "falling" from n=3 to n=2; n=4 to n=2; n=5 to n=2; n= 6 to n=2.

11

What are the strengths and weaknesses of the Bohr model?

The Bohr model introduces the important understanding that energy levels, where electrons are permitted to exist, are quantized.

Unfortunately, experimental data and theoretical calculations based on the model do not work when applied to multi-electron atoms - that is, all other atoms besides hydrogen. Therefore, although it is assumed that quantized energy levels exist in all atoms, a more developed model is needed, especially to describe the electronic structure in multi-electron atoms. This more developed model is of course the quantum mechanical model.

12

Define orbital.

An orbital is a mathematical solution to the Schrodinger equation based on the assumption that electrons behave as waves. This mathematical solution is related to the region of space outside the nucleus where there is a high probability of the electron existing. This mathematical solution is also related to the relative energy of the electron in the atom.

13

How many quantum numbers are needed to define an orbital? Describe each one.

Three quantum numbers are needed to define an orbital:

  1. (n) = principal quantum number- this is the most important quantum number in describing the energy state of the electron. This number also defines the shell;
  2. (l) = azumithal quantum number - this describes shape of the region of space with probability of finding the electron. Essentially it describes the shape of the boundary surface diagram.  This quantum is associated with the subshell designations s, p, d, f and therefore also defines to some degree the energy of the electron (s p d f).;
  3. (ml) = magnetic quantum number - this describes the orientation in space of the orbital.  For example a p subshell orbitals exist on the x, y, z axis (i.e. perpendicular).

14

How many orbitals exist in the hydrogen atom?

There are potentially an infinite number of orbitals. According to the quantum model, the one electron in hydrogen exists primarily in the 1s orbital. This corresponds to ground state. This one electron however can be excited to higher energy orbitals (e.g. 2s, 3p, 4d etc).

15

How many quantum numbers are needed to define an electron?

Four quantum numbers are needed to describe a single electron. In addition to n, l, ml, the fourth quantum number is called the quantum spin number (ms). The only allowed values for ms are +1/2 and -1/2. For Chem 112 students this is more easily understood as representing two possible electrons of exactly opposite spins. Electrons of different spins are often represented in orbital diagrams by arrows with the head tip either pointing upwards or downwards.

16

Define degenerate orbitals.

Degenerate orbitals are orbitals within the same subshell and are therefore of the same energy. For example all three 2p orbitals (2px,2py, 2pz) are said to be degenerate and therefore all have the same energy.

17

Define Aufbau principle

This principle predicts that in multi-electron atoms, electrons will fill the lower energy orbitals before filling the higher energy orbitals.

18

Define the Pauli exclusion principle

This principle, proposed by Wolfgang Pauli, indicates that no two electrons can be described by the same four quantum numbers. In other words, each electron in an orbital must have a unique spin. Because there are only two possible spins available to an electron, this means that an orbital can only hold two electrons.

19

Explain Hund's rule

Hund's rule dictates that degenerate orbitals will first be filled with single electrons of identical spin, before electrons are paired in one or more of the degenerate orbitals. An example is shown when considering the electron configuration of nitrogen (1s22s22p3). The three electrons in the 2p orbitals are spread apart so that there is one in 2px, one in 2py and one in 2pz.

20

Define orbital diagram.

An orbital diagram shows a series of lines or boxes each representing an orbital. Electrons are represented as arrows with arrows pointing up or down, corresponding to the spin. The lines or boxes are usually labelled with the subshell designation (2s, 3p, 4d etc). Orbital diagrams are usually used to represent a subset of electrons (e.g. valence elctrons) as a means of emphasizing the electron spin.

21

Define electron configuration.

Electron configuration is a shortform notation (using the subshell designations) to represent all the electron occupied orbitals in an atom. E.g. the electron configuration for sodium is 1s22s22p63s1. This can be abbreviated even further as [Ne]3s1.

22

Define valence electrons.

Valence electrons are the outermost electrons of an atom. These electrons are the ones usually involved in covalent bonding. The number of valence electrons can be determined by noticing in which main-group column the element is found. E.g Carbon is found in main-group 4 and therefore has four valence electrons. In terms of electron configuration, these are the 2s2, 2p2 electrons.

23

Define atomic radius.

Atomic radius is the distance the valence electrons are from the nucleus.

This distance can be determined by measuring the internuclear distance between to atoms of the same element divided by 2. Given that atoms are considered to behave something like billiard balls (i.e. they behave as if they are spherical), the atomic radius is a good indication of the size of the atom.

24

State the trend for atomic radius in relation to the periodic table.

Across the periodic table the atomic radius decreases; down the periodic table the atomic radius increases.

25

Explain the trend for atomic radius.

Across the periodic table the nuclear charge increases, therefore the force of attraction between the nucleus and the valence electrons increases. Greater force of attaction pulls the valence electrons closer to the nucleus. As a result across the periodic table the atomic radius decreases.

Down the periodic table, the number of shells occupied by electrons increases and therefore the valence electrons are further from the nucleus. Therefore down the periodic table the atomic radius increases.

26

Define ionization energy.

Ionization energy is the energy required to completely remove an electron from an atom.

27

State the trend for ionization energy in relation to the periodic table.

Across the periodic table the ionization energy in general increases; down the periodic table the ionization energy decreases.

28

Explain the trend for ionization energy.

Across the periodic table the nuclear charge increases, therefore the force of attraction between the nucleus and the valence electrons increases. Greater force of attaction results in more energy required to remove an electon. As a result across the periodic table the ionization energy in general increases.

Down the periodic table, the number of shells occupied by electrons increases and therefore the valence electrons are further from the nucleus. As a result the force of attraction between the valence electrons and the nucleus decreases. Therefore down the periodic table the ionization energy decreases.

29

Define electronegativity.

Electronegativity is the attraction the nucleus has for a shared pair of electrons.

30

State the trend for electronegativity in relation to the periodic table.

Across the periodic table the electronegativity in general increases; down the periodic table the electronegativity in decreases.

31

Explain the trend for electronegativity.

Across the periodic table the nuclear charge increases, therefore the force of attraction between the nucleus and the valence electrons increases. Greater force of attaction results in a higher electronegativity. As a result across the periodic table the electronegativity in general increases.

Down the periodic table, the number of shells occupied by electrons increases and therefore the valence electrons are further from the nucleus. As a result the force of attraction between the valence electrons and the nucleus decreases. Therefore down the periodic table the electronegativity decreases.

32

Define isoelectronic series.

An isoelectronic series describes a collection of ions who all have identical electron configurations.

For example, N3-, O2-, F-, Na+, Mg2+, Al3+ all have the electron configuration 1s22s22p6. Although they have the identical electron configuration, they differ in size (in ionic radius). N3- is the largest and Al3+ the smallest because from N3- to Al3+ there is an increase in number of protons and therefore nuclear charge.

33

Explain what happens to the radius when an atom is changed to the corresponding cation.

When an atom is changed to a cation, the radius decreases. The main reason for this is a loss of an electron to form a stable cation usually results in one less shell being occupied by electrons. The electron configuration for Na is 1s22s22p63s1 and for Na+ 1s22s22p6.

34

Explain what happens to the radius when an atom is changed to the corresponding anion.

When an atom Is changed to an anion, the radius increases. The main reason for this is a gain of one or more electrons. Given that the number of protons does not change but there are more electrons and enhanced repulsion, the anion is larger than the corresponding atom.

35

In the Bohr model this is the lowest and most stable energy state for the electron. The electron exists predominantly in this state.

Ground state.

36

An atom is in this state when it has absorbed energy which has allowed the electron to exist temporarily in one of the permitted higher energy states.

Excited state.

37

Discrete lines at designated wavelengths (in the UV, visible and IR) which result from electrons "falling" from higher to lower energy levels. 

 

Emission spectrum.

38

The distance the valence electrons are from the nucleus.

 

Atomic radius.

39

The energy required to completely remove an electron from an atom.

Ionization energy.

40

The attraction the nucleus has for a shared pair of electrons.

Electronegativity.

41

Isoelectronic series describes a collection of ions who all have identical electron configurations.

 

Isoelectronic series.