Ch.4-Electronic Structure of Atoms

Question 1

What is an energy level?

Answer 1

An energy level is a region of definite energy within the atom that electrons can occupy. Electrons in atoms occupy levels that are outside the nucleus.

Question 2

How can you determine how many electrons an energy level can hold?

Answer 2

To determine how many electrons an energy level can hold, we use the formula 2n^2 (n representing the number of the energy level).

Question 3

What is a spectroscope?

Answer 3

A spectroscope can be used to analyse the light emitted by elements.

Question 4

What is the emission spectrum (of an element)?

Answer 4

The emission spectrum (line spectrum) is electromagnetic radiation that is characteristic of an element. It is seen in sodium street lamps and particular colours in fireworks.

Question 5

What are absorption spectra?

Answer 5

The absorption spectrum of an element is the spectrum that is observed after white light has been passed through it. An absorption spectrum of an element consists of a series of dark lines on a coloured background.

Question 6

What is the Balmer series?

Answer 6

The Balmer series contains all of the lines in the visible region of the hydrogen spectrum.

Question 7

What is an atomic absorption spectrometer?

Answer 7

An atomic absorption spectrometer can be used to measure the amount of an element in a sample from the amount of light absorbed by free atoms of the element. It can be used in analysis of heavy metals and to estimate the amount of lead in a blood sample.

Question 8

What is Heisenburg’s uncertainty principle?

Answer 8

The Heisenburg uncertainty principle states that it is impossible to determine, at the same time, the exact position and velocity of an electron. In order to determine position or velocity of an electron within an atom, it must be illuminated with radiation. However, radiation interferes with the electron by changing the position and momentum of the electron.

Question 9

What was Bohr’s theory of atoms?

Answer 9

The Bohr theory only explained the line spectrum of hydrogen. The energy of an electron in the hydrogen atom is quantised (restricted to certain values). Electrons have energy values that are restricted by energy levels. When an electron moves from a higher energy level to a lower energy level, a definite amount of energy is emitted.

Question 10

What are the Lyman and Paschen series?

Answer 10

The Lyman series contains the lines in the ultraviolet region of the hydrogen spectrum, it occurs if electron jump back to the 1st energy level. The Paschen series contains the lines in the infrared region of the hydrogen spectrum, it occurs if an electron jumps back to the 3rd energy level.

Question 11

What is the amount of energy release when an electron moves from a higher energy level to a lower one?

Answer 11

The amount of energy is equal to the difference between the 2 levels:
E2 - E1 = hf, where h=Plank’s constant (6.63X10^-34 Js) and f=frequency of the light emitted.

Question 12

What happens when an atom absorbs an amount of energy equal to the energy difference of 2 levels?

Answer 12

The electron will move from the lower energy level to the higher energy level (E2-E1=hf).

Question 13

What is ground state and excited state?

Answer 13

Ground state is the lowest energy state while the excited state is the higher energy state.

Question 14

What did Louis de Broglie say?

Answer 14

In 1923, Louis de Broglie stated that electrons have the properties of waves as well as particles.

Question 15

What is an atomic orbital?

Answer 15

An atomic orbital is a region in space where the probability of finding an electron is relatively high.

Question 16

What is an energy sublevel?

Answer 16

An energy sublevel is a subdivision of an energy level containing one or more atomic orbitals, all of which have the same energy.

Question 17

State the rules used in assigning electrons to the various sublevels and orbitals.

Answer 17

The Aufbau principle states that electrons will occupy the lowest energy sublevel available. Hund’s rule states that electrons tend to occupy orbitals of equal energy singly where possible. Pauli’s principle states that not more than 2 electrons can occupy an orbital at the one time.

Question 18

How many electrons can occupy an orbital at one time?

Answer 18

Not more than 2 electrons can occupy an orbital at one time.

Question 19

Why is a knowledge of the electronic configuration of atoms important to chemists?

Answer 19

The outer electronic configuration of an element largely determines its chemical properties. Elements with similar outer electronic configurations, such as the alkali metals, have similar chemical properties.

Question 20

How many electrons does each group have on their outermost energy level?

Answer 20

Each group has the same number of electrons on its outermost energy level as its group number, group II has 2 electrons. Except for helium, which has 2 electrons on its outermost energy level as opposed to 8 like the rest of the noble gases.

Question 21

How do coloured lines on the Balmer series appear?

Answer 21

All of the lines on the Balmer series are due to electrons dropping from higher levels to the n=2 energy level. From 7 to 2, is invisible, 6 to 2, violet, 5 to 2, indigo, 4 to 2, green, 3 to 2, red.

Question 22

How can the light spectra of elements other than hydrogen be explained?

Answer 22

Electrons receive energy from a flame to move to higher energy levels. They are then unstable and rapidly fall down to lower levels. Each time they do this they emit a colour whose energy is equal to the energy difference between the 2 levels. This colour appears on the line spectrum of that element. The line spectrum is unique because the spacings between energy levels in that element are also unique.

Question 23

What are the limitations of the Bohr theory?

Answer 23

The Bohr theory only worked well for hydrogen, it did not take into account that electrons have the properties of particles as well as waves, it did not allow for Heisenburg’s uncertainty principle, it did not explain the discovery of sublevels or the existence of orbitals.

Question 24

What is atomic radius?

Answer 24

The atomic radius of an element is half the distance between the nuclei of 2 atoms of the element that are joined together by a single covalent bond.

Question 25

What happens to the size of atomic radius going down a group?

Answer 25

Atomic radius increases on going down a group, as the number of energy levels occupied increases. The bigger the radius, the greater the screening effect, due to the presence of more inner energy level electrons.

Question 26

What happens to the size of atomic radius going across a period?

Answer 26

Atomic radius decreases going across a group because of the increasing nuclear charge, meaning greater attractive force on the outer shell electrons. Within the same energy, the screening effect is the same as the number of electrons on the inner levels remains the same.

Question 27

What is the screening effect?

Answer 27

The screening effect occurs when electrons in the inner shells help to weaken the force of attraction between the nucleus and the electron in the outer shell.

Question 28

What is the first ionisation energy of an element?

Answer 28

The first ionisation energy of an element is the minimum energy in kilojoules required to remove the most loosely bound electron from each isolated atom in a mole of the element in its ground state.

Question 29

How do first ionisation energies change going across a period?

Answer 29

First ionisation energies generally increase going across a period, this is due to, the increase in nuclear charge caused by the increase in the number of protons in successive elements. The decrease in atomic radius means that electrons are held more tightly, making them harder to remove.

Question 30

How do first ionisation energies change going down a period?

Answer 30

The value of first ionisation energies decrease going down a group because of the increase in atomic radius, making it easier to remove an electron from an atom. The screening effect of inner energy levels makes the nuclear charge much less effective on the outermost electron.

Question 31

Why aren’t the electronic configuration of copper and chromium as expected?

Answer 31

The unexpected configurations are due to the extra stability in the case of chromium of a structure with half filled 3d and 4s sublevels, and in the case of copper of a structure with a full 3d and a half-filled 4s sublevel.

Question 32

What exceptions are observed for the trend of first ionisation energies going across a period?

Answer 32

Beryllium has a higher first ionisation energy than boron. The loosest electron in boron is more easily removed than beryllium’s because of beryllium’s more stable structure, it has a full outer sublevel. Also, oxygen has 2 electrons in the 2px orbital, while 2py and 2pz have 1. The 2 electrons repel each other, making the outermost electron easier to remove. The nitrogen atom has a half full outer p sublevel, making it more stable.

Question 33

What is the second ionisation energy of an element?

Answer 33

The second ionisation energy of an element is the minimum energy required to remove the most loosely bound electron from each singly charged positive ion in a mole of these ions.

Question 34

What evidence is there for the existence of sublevels?

Answer 34

Successive ionisation energies increase as more electrons are removed. Large jumps in the ionisation energy reveal where electrons are being removed from the next principle energy level. For example there are large jumps between the 2nd, 3rd, 10th and 11th ionisation energies for magnesium.

Question 35

Name the test for the evidence of Bohr’s theory? Explain.

Answer 35

The flame test. Dip a damp wooden splint into a sample of salt. Place it in the bunsen burner flame and note the colour the salt burns. Lithium turns crimson. Sodium turns yellow. Potassium turns lilac. Barium turn green. Copper turns blue-green. Strontium turns red.

Question 36

What is electronegativity?

Answer 36

Electronegativity is the relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond. Nuclear charge and atomic radius affects electronegativity.