4 - Electronic Structure of the Atom

To be accompanied with past exam questions

Explain what is meant by energy level

A region of definite energy within an atom that electrons can occupy

Explain what is meant by energy sub-level

A group of atomic orbitals within an atom, all of which have the same energy

Distinguish between the ground state and the excited states of the electron in a hydrogen atom

Ground state: lowest energy / n = 1 / stable //

Excited state: greater energy / n > 1 / unstable

Explain what happens when an electron moves from an excited state to its ground state

Energy is released (3)

as a photon / as light / at a specific frequency (3)

Name the instrument used to examine the line spectrum of an element

Spectrometer

State Heisenberg’s uncertainty principle

It is impossible to determine at the same time the exact position and velocity of an electron

Write the full s, p electron configuration for an atom of neon (Ne) in its ground state, including the number of electrons in each orbital

By reference to the electron configuration for neon, distinguish between an energy level and an atomic orbital

Sublevel: 2p // (Ne has) 3

Orbital: e.g. 2p_x // (Ne has) 5

Neon is chemically unreactive. Explain this property by reference to its electron configuration

Full outer energy level / full p sublevel / satisfies octet rule

Explain how the successive ionisation energy values of neon provide evidence supporting the existence of energy levels

Ninth ionisation energy significantly greater than the eighth

Draw the shape of a p orbital

Define atomic orbital

A region in space where there is a high probability of finding an electron

Explain how the line emission spectrum of hydrogen arises and provides evidence for the existence of energy levels

1. e^- revolve around the nucleus in fixed paths called orbits.

2. e^- in any orbit have a fixed amount of energy. Energy of an e^- in a particular orbit is quantised.

3. As long as an e^- is in any one particular energy level it neither gains nor loses energy.

4. When energy is supplied to an electron, it is promoted from it’s normal energy level (ground state), to a higher energy level (excited state).

5. Electrons cannot maintain this excited state and fall back down to their ground state, releasing a definite amount of energy whose value is equal to the difference in energy between the two energy levels. E₂ - E₁ = hf { h is Planck’s constant, f is the frequency of light emitted}.

6. Definite amount of energy appears as a coloured line in the line emission spectrum of the element.

Write the electronic configuration of an atom of sulphur showing the distribution of electrons in atomic orbitals in the ground state.

Hence, state how many

i) main energy levels,

ii) orbitals are occupied by electrons.

1s² 2s² 2p⁶ 3s² 3p⁴

i) 3

ii) 9

Write the electronic configuration of an atom of calcium showing the distribution of electrons in atomic orbitals in the ground state.

Hence, state how many

i) main energy levels,

ii) orbitals are occupied by electrons.

1s² 2s² 2p⁶ 3s² 3p⁶ 4s²

i) 4

ii) 10

Use the electron configuration of calcium to predict how a calcium atom is likely to behave in a chemical reaction. Explain your reasoning

Loses outer 2 electrons

to satisfy octet rule /

8 electrons in outer shell is stable

The dispositive ion M²⁺ has 28 electrons and 34 neutrons. What is

i) the atomic number,

ii) the mass number, of M?

i) 30 //

ii) 64

Who coined the name ‘electron‘?

Stoney

What were the contributions of Thomason and Millikan to our understanding of electrons?

Thomson: showed electrons were negatively charged particles /

identified electrons as subatomic particles //

measured ratio of charge to mass of electron

Millikan: measured size of charge on electron

Give the electron configuration of copper in its ground state

1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰

How are the electrons arranged in the orbitals of the highest occupied sub-level of a nitrogen atom in its ground state?

Why does an individual atom not have a definite boundary?

Heisenberg’s uncertainty principle / electron’s position and velocity not known at the same time / wave nature of electron

How is atomic radius defined?

Half the distance between the nuclei of two atoms of an element that are joined together by a single covalent bond

The bond length of the single Cl-Cl covalent bond is 0.199 nm. Predict the value in nm for the atomic radius of chlorine.

0.0995

Account for the general decrease in atomic radii of the elements across the third period of the periodic table from sodium to chlorine

Nuclear charge increasing / atomic number increases / number of protons increases

State and explain the trend in atomic radii across the second period of the periodic table of elements

State: Decrease in atomic radius

Explain: Increase in nuclear charge

Why is establishing an atomic radius for argon problematic?

Ar-Ar bonds are unlikely to exist / do not exist

Define first ionisation energy

The energy required to removed the most loosely bound electron from an isolated atom of the element in it’s ground state

Why do first ionisation energy values show a general increase across the second period of the periodic table?

Increasing nuclear charge //

atomic radius decreasing

Why is the first ionisation energy of potassium lower than that of argon or chlorine?

Outermost electron of potassium is farther from the nucleus than argon or chlorine //

greater degree of shielding (screening) in potassium

Why is the second ionisation energy value for each element in the table bigger than its first ionisation energy value?

Requires more energy to remove electron from a positive ion than from a neutral atom /

greater attraction between the nucleus and electron /

effective nuclear charge increased

Why is the second ionisation energy of potassium very significantly bigger than its first?

Second electron comes from a new shell / first electron removed easily as it is the only electron in outer shell / second electron closer to the nucleus / second electron less shielded

Account for the decreasing trend in first ionisation energy values down Group 1 of the periodic table

Increasing atomic radius / outer electrons farther from nucleus /

extra shell //

additional screening

Why is there a general increase in first ionisation energy values across the elements of the third period?

Nuclear charge increasing //

atomic radius decreasing

The diagram represents the 10 electrons in the 4s and 3d sub-levels of a neutral atom of a certain element in its ground state.

What is the element?

Nickel (Ni)

Give two differences between an atomic orbit, as described by Bohr, and an atomic orbital

Distinguish between a 2p orbital and a 2p sublevel

2p sublevel consists of three 2p orbitals /

2p sublevel has 6 electrons but each of the 2p orbitals has 2 of these electrons

Use diagrams to distinguish between a p orbital and a p sub-level

State the maximum number of electrons that can be accommodated in a p-orbital.

2

State one piece of evidence that supports the existence of atomic energy levels

Atomic (emission, absorption, line) spectrum / ionisation energies

Identify the main energy levels involved in the electron transition that gives rise to the first (red) line of the Balmer series in the emission spectrum of the hydrogen atom

3 and //

2

Write the ground state electron configuration for a carbon atom.

How many orbitals are occupied?

1s² 2s² 2p²

4

Explain in terms of energy sublevels why the arrangement of electrons in the main energy levels in a calcium atom is 2, 8,8, 2 and not 2, 8, 10.

4s sublevel is lower in energy than the 3d / electrons fill the 4s sublevel before the 3d

Explain why the first ionisation energy of oxygen in lower than that of nitrogen despite the general increase in values across the second period of the periodic table

Nitrogen is relatively stable //

has a half-filled 2p outer sublevel

or

oxygen is less stable //

doesn’t have half filled 2p outer sublevel

{Note: the next stable state to a full energy sublevel is a half-filled energy sublevel}

The colours produced by fireworks provide evidence for energy levels in atoms of other elements.

Suggest an element that gives a blue-green colour to a firework display

Copper

[Barium acceptable]

Write the electron configuration of an atom of silicon showing the distribution of electrons in atomic orbitals in the ground state.

Hence, state how many (i) main energy levels, (ii) atomic orbitals, are occupied in the silicon atom in its ground state

i) 3 //

ii) 8

How many electrons are there in 2.3g of sodium metal, Na?

State Aufbau’s Principle

Electrons will occupy the lowest energy level available to them

State Pauli’s Exclusion Principle

No more than 2 electrons can occupy an orbital and they must have opposite spin

Why is magnesium more reactive than beryllium?

Mg has greater atomic radius / more energy levels in Mg /

more shielding (screening) effect in Mg

Mg has smaller ionisation energy /

What colour is observed in a flame test on a salt of potassium?

Lilac

What colour is observed in a flame test on a salt of sodium?

Orange

What colour is observed in a flame test on a salt of lithium?

Crimson red

What colour is observed in a flame test on a salt of barium?

Yellow / green

What colour is observed in a flame test on a salt of strontium?

Red

What colour is observed in a flame test on a salt of copper?

Blue green

Define electronegativity

The relative attraction of an atom for a shared pair of electrons in a covalent bond

Why is there an increase in electronegativity value moving from gallium to germanium in the periodic table?

Nuclear charge increasing //

atomic radius decreasing

Give one reason why electronegativity values exhibit a general increase across the second period of the periodic table

Increase in nuclear charge / decrease in atomic radius