Define ionisation energy .

Ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous ions. It is measured under standard conditions (298 K, 100 kPa) and has units of kJ mol -1 .

Successive ionisation energies always .......... . This is because as electrons are removed, shielding decreases and the proton-to-electron ratio .......... .

Successive ionisation energies always increase . This is because as electrons are removed, shielding decreases and the proton-to-electron ratio increases .

Define first ionisation energy .

First ionisation energy ( IE 1 ) is the energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions. For example: Ca (g) → Ca + (g) + e -

True or False? First ionisation energy generally increases across a period.

True. Across a period, nuclear charge increases while shielding remains roughly constant. The outer electrons are pulled closer to the nucleus, making them harder to remove.

Why is there a dip in first ionisation energy between beryllium and boron ?

There is a dip in first ionisation energy between beryllium and boron because boron's fifth electron is in the 2p subshell , which is at a higher energy level and further from the nucleus than the 2s subshell of beryllium. It is therefore easier to remove.

What are successive ionisation energies ?

Successive ionisation energies are the energies required to remove each subsequent electron from a gaseous ion, one mole at a time. They always increase because removing electrons from an increasingly positive ion requires more energy due to greater proton-to-electron ratio and reduced shielding.

What four factors affect the size of the first ionisation energy ?

The four factors affecting the size of the first ionisation energy are: Size of the nuclear charge Distance of outer electrons from the nucleus Shielding by inner electrons Spin-pair repulsion

True or False? The dip in first ionisation energy between nitrogen and oxygen occurs because oxygen has a higher nuclear charge, making the outer electron easier to remove.

False. The dip occurs because of spin-pair repulsion in the 2p x orbital of oxygen. Two electrons share this orbital, and the repulsion between them makes one easier to remove, lowering the IE 1 slightly compared to nitrogen.

Why does first ionisation energy decrease down a group?

First ionisation energy decreases down a group because atomic radius increases as new electron shells are added, increasing the distance between the nucleus and the outer electron. Increased shielding by inner shells also reduces the effective nuclear attraction. These effects outweigh the increased nuclear charge.

Quanta are discrete packets of energy that electrons can possess when moving between energy levels. The singular form is quantum . The existence of quanta explains why line emission spectra show distinct lines rather than a continuous spectrum.

What does convergence in a line emission spectrum indicate?

Convergence in a line emission spectrum indicates that the electron is approaching a maximum energy. The point of convergence corresponds to the ionisation energy of the electron — the energy required to remove it completely from the atom.

True or False? Lines in the hydrogen line emission spectrum converge towards the higher energy end.

True. The lines in the hydrogen line emission spectrum converge (get closer together) towards the higher energy (blue) end, approaching the ionisation energy of the electron.

In a successive ionisation energy graph, large jumps in energy indicate a change of .......... ; smaller jumps indicate a change of .......... .

In a successive ionisation energy graph, large jumps in energy indicate a change of shell ; smaller jumps indicate a change of subshell .

What is a principal quantum number ?

A principal quantum number ( n ) is the number used to label each principal energy level (electron shell) around a nucleus. The lower the value of n , the closer the shell is to the nucleus and the lower its energy .

Each principal quantum shell can hold a maximum of .......... electrons, where n is the principal quantum number.

Each principal quantum shell can hold a maximum of 2 n 2 electrons, where n is the principal quantum number.

How many electrons can the third principal quantum shell ( n = 3) hold?

The third principal quantum shell ( n = 3) can hold a maximum of 18 electrons , calculated using 2 n 2 = 2 × 3 2 = 18.

True or False? A lower principal quantum number means the electron shell is further from the nucleus.

False. A lower principal quantum number ( n ) means the shell is closer to the nucleus and has lower energy. The first shell ( n = 1) is the closest to the nucleus.

Define line emission spectrum .

A line emission spectrum is a series of discrete lines of specific frequencies of light produced when electrons in excited atoms fall back to lower energy levels and emit energy. Each line corresponds to a specific electron transition between energy levels .

Define atomic orbital .

An atomic orbital is a region of space around the nucleus that can hold a maximum of two electrons with opposite spins. Orbitals exist at specific energy levels and electrons can only occupy these fixed levels.

How many electrons can each atomic orbital hold, and what is the maximum number of electrons in a p subshell ?

Each atomic orbital can hold a maximum of two electrons . A p subshell contains three orbitals, so it can hold a maximum of six electrons (3 x 2).

s orbitals are .......... in shape. p orbitals have a .......... shape and are oriented perpendicular to one another along the x, y and z axes.

s orbitals are spherical in shape. p orbitals have a dumbbell shape and are oriented perpendicular to one another along the x, y and z axes.

How are the three p orbitals oriented relative to one another?

The three p orbitals are oriented perpendicular to one another, lying along the x, y and z axes. This means each p orbital points at right angles (90°) to the other two.

True or False? Knowledge of the shape of d orbitals is required for Edexcel A Level Chemistry.

False. Knowledge of d orbital shapes is not required for Edexcel A Level Chemistry. Only the shapes of s and p orbitals are required.

Hund's Rule states that electrons occupy orbitals in the same subshell singly and with parallel spins before pairing up. This minimises spin-pair repulsion between electrons.

Electrons fill separate orbitals within a subshell first because electrons with the same spin .......... each other. This effect is called .......... repulsion.

Electrons fill separate orbitals within a subshell first because electrons with the same spin repel each other. This effect is called spin-pair repulsion.

What is the Pauli Exclusion Principle ?

The Pauli Exclusion Principle states that each atomic orbital can hold a maximum of two electrons, and they must have opposite spins . This applies to all orbitals across all subshells and shells.

True or False? Electrons always move to a higher-energy empty orbital rather than pair up in a lower-energy orbital.

False. Electrons fill orbitals in order of increasing energy ( Aufbau principle ) — they pair up in lower-energy orbitals rather than jump to higher-energy ones. Note: within a subshell, Hund's Rule does apply — electrons occupy each orbital singly before pairing — but they never skip to a higher subshell to avoid pairing.

What are degenerate orbitals ?

Degenerate orbitals are orbitals within the same subshell that have identical energies. For example, the p x , p y and p z orbitals in any given p subshell are all degenerate .

Why is the 4s subshell filled before the 3d subshell when writing electron configurations?

The 4s subshell is filled before the 3d subshell because the 4s orbital is slightly lower in energy than the 3d orbital. Electrons fill orbitals in order of increasing energy , so 4s is filled first.

Subshells increase in energy in the order .......... . The exception is the 3d orbital, which has slightly higher energy than the .......... orbital.

Subshells increase in energy in the order s < p < d < f . The exception is the 3d orbital, which has slightly higher energy than the 4s orbital.

What is the difference between a full and a shorthand electron configuration?

A full electron configuration lists all electrons from the 1s subshell upwards. A shorthand electron configuration replaces the inner electrons with the symbol of the nearest preceding noble gas in square brackets. For example, potassium: full = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 ; shorthand = [Ar] 4s 1 .

True or False? Transition metals lose electrons from the 3d subshell before the 4s subshell when forming positive ions.

False. Transition metals lose electrons from the 4s subshell first, not the 3d. Although the 4s is filled before the 3d when building up the configuration, it is the outer subshell and so electrons are removed from it first when forming positive ions .

What are the exceptional electron configurations of chromium and copper , and why do they occur?

Chromium is [Ar] 3d 5 4s 1 and copper is [Ar] 3d 10 4s 1 , rather than the expected [Ar] 3d 4 4s 2 and [Ar] 3d 9 4s 2 . These configurations are energetically stable because half-filled and fully-filled d subshells have lower energy.

What are s-block , p-block , d-block and f-block elements?

s-block elements have their valence electrons in an s orbital. p-block elements have their valence electrons in a p orbital. d-block elements have their valence electrons in a d orbital. f-block elements have their valence electrons in an f orbital.

All elements in the same .......... of the Periodic Table have the same number of electron shells. All elements in the same .......... have the same number of outer electrons.

All elements in the same period of the Periodic Table have the same number of electron shells. All elements in the same group have the same number of outer electrons.

Why do noble gases rarely react with other elements?

Noble gases rarely react because they have full subshells in their electronic configuration , which gives them very high ionisation energies . This stable arrangement means there is little tendency to gain, lose or share electrons.

True or False? Elements in the same group have the same number of outer electrons.

True. Elements in the same group of the Periodic Table all have the same number of outer (valence) electrons . This shared configuration explains the similar chemical properties observed within a group.

A LevelChemistryEdexcelFlashcards1. Physical ChemistryIons & Electrons

Ions & Electrons (Edexcel A Level Chemistry): Flashcards

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FrontIonisation Energy
Define ionisation energy.

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Define ionisation energy.
Ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous ions.
It is measured under standard conditions (298 K, 100 kPa) and has units of kJ mol^-1.
Successive ionisation energies always ........... This is because as electrons are removed, shielding decreases and the proton-to-electron ratio ...........
Successive ionisation energies always increase. This is because as electrons are removed, shielding decreases and the proton-to-electron ratio increases.
Define first ionisation energy.
First ionisation energy (IE₁) is the energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.
For example: Ca (g) → Ca⁺ (g) + e^-
True or False?
First ionisation energy generally increases across a period.
True.
Across a period, nuclear charge increases while shielding remains roughly constant. The outer electrons are pulled closer to the nucleus, making them harder to remove.
Why is there a dip in first ionisation energy between beryllium and boron?
There is a dip in first ionisation energy between beryllium and boron because boron's fifth electron is in the 2p subshell, which is at a higher energy level and further from the nucleus than the 2s subshell of beryllium. It is therefore easier to remove.
What are successive ionisation energies?
Successive ionisation energies are the energies required to remove each subsequent electron from a gaseous ion, one mole at a time.
They always increase because removing electrons from an increasingly positive ion requires more energy due to greater proton-to-electron ratio and reduced shielding.
What four factors affect the size of the first ionisation energy?
The four factors affecting the size of the first ionisation energy are:
1. Size of the nuclear charge
2. Distance of outer electrons from the nucleus
3. Shielding by inner electrons
4. Spin-pair repulsion
True or False?
The dip in first ionisation energy between nitrogen and oxygen occurs because oxygen has a higher nuclear charge, making the outer electron easier to remove.
False.
The dip occurs because of spin-pair repulsion in the 2p_x orbital of oxygen. Two electrons share this orbital, and the repulsion between them makes one easier to remove, lowering the IE₁ slightly compared to nitrogen.
Why does first ionisation energy decrease down a group?
First ionisation energy decreases down a group because atomic radius increases as new electron shells are added, increasing the distance between the nucleus and the outer electron. Increased shielding by inner shells also reduces the effective nuclear attraction. These effects outweigh the increased nuclear charge.
What are quanta?
Quanta are discrete packets of energy that electrons can possess when moving between energy levels.
The singular form is quantum. The existence of quanta explains why line emission spectra show distinct lines rather than a continuous spectrum.
What does convergence in a line emission spectrum indicate?
Convergence in a line emission spectrum indicates that the electron is approaching a maximum energy. The point of convergence corresponds to the ionisation energy of the electron — the energy required to remove it completely from the atom.
True or False?
Lines in the hydrogen line emission spectrum converge towards the higher energy end.
True.
The lines in the hydrogen line emission spectrum converge (get closer together) towards the higher energy (blue) end, approaching the ionisation energy of the electron.
In a successive ionisation energy graph, large jumps in energy indicate a change of .......... ; smaller jumps indicate a change of .......... .
In a successive ionisation energy graph, large jumps in energy indicate a change of shell; smaller jumps indicate a change of subshell.
What is a principal quantum number?
A principal quantum number (n) is the number used to label each principal energy level (electron shell) around a nucleus.
The lower the value of n, the closer the shell is to the nucleus and the lower its energy.
Each principal quantum shell can hold a maximum of ..........electrons, where n is the principal quantum number.
Each principal quantum shell can hold a maximum of 2n² electrons, where n is the principal quantum number.
How many electrons can the third principal quantum shell (n = 3) hold?
The third principal quantum shell (n = 3) can hold a maximum of 18 electrons, calculated using 2n² = 2 × 3² = 18.
True or False?
A lower principal quantum number means the electron shell is further from the nucleus.
False.
A lower principal quantum number (n) means the shell is closer to the nucleus and has lower energy. The first shell (n = 1) is the closest to the nucleus.
Define line emission spectrum.
A line emission spectrum is a series of discrete lines of specific frequencies of light produced when electrons in excited atoms fall back to lower energy levels and emit energy.
Each line corresponds to a specific electron transition between energy levels.
Define atomic orbital.
An atomic orbital is a region of space around the nucleus that can hold a maximum of two electrons with opposite spins.
Orbitals exist at specific energy levels and electrons can only occupy these fixed levels.
How many electrons can each atomic orbital hold, and what is the maximum number of electrons in a p subshell?
Each atomic orbital can hold a maximum of two electrons. A p subshell contains three orbitals, so it can hold a maximum of six electrons (3 x 2).
s orbitals are ..........in shape. p orbitals have a ..........shape and are oriented perpendicular to one another along the x, y and z axes.
s orbitals are spherical in shape. p orbitals have a dumbbell shape and are oriented perpendicular to one another along the x, y and z axes.
How are the three p orbitals oriented relative to one another?
The three p orbitals are oriented perpendicular to one another, lying along the x, y and z axes. This means each p orbital points at right angles (90°) to the other two.
True or False?
Knowledge of the shape of d orbitals is required for Edexcel A Level Chemistry.
False.
Knowledge of d orbital shapes is not required for Edexcel A Level Chemistry. Only the shapes of s and p orbitals are required.
Define Hund's Rule.
Hund's Rule states that electrons occupy orbitals in the same subshell singly and with parallel spins before pairing up.
This minimises spin-pair repulsion between electrons.
Electrons fill separate orbitals within a subshell first because electrons with the same spin .......... each other. This effect is called .......... repulsion.
Electrons fill separate orbitals within a subshell first because electrons with the same spin repel each other. This effect is called spin-pair repulsion.
What is the Pauli Exclusion Principle?
The Pauli Exclusion Principle states that each atomic orbital can hold a maximum of two electrons, and they must have opposite spins.
This applies to all orbitals across all subshells and shells.
True or False?
Electrons always move to a higher-energy empty orbital rather than pair up in a lower-energy orbital.
False.
Electrons fill orbitals in order of increasing energy (Aufbau principle) — they pair up in lower-energy orbitals rather than jump to higher-energy ones. Note: within a subshell, Hund's Rule does apply — electrons occupy each orbital singly before pairing — but they never skip to a higher subshell to avoid pairing.
What are degenerate orbitals?
Degenerate orbitals are orbitals within the same subshell that have identical energies.
For example, the p_x, p_y and p_z orbitals in any given p subshell are all degenerate.
Why is the 4s subshell filled before the 3d subshell when writing electron configurations?
The 4s subshell is filled before the 3d subshell because the 4s orbital is slightly lower in energy than the 3d orbital. Electrons fill orbitals in order of increasing energy, so 4s is filled first.
Subshells increase in energy in the order ........... The exception is the 3d orbital, which has slightly higher energy than the ..........orbital.
Subshells increase in energy in the order s < p < d < f. The exception is the 3d orbital, which has slightly higher energy than the 4s orbital.
What is the difference between a full and a shorthand electron configuration?
A full electron configuration lists all electrons from the 1s subshell upwards. A shorthand electron configuration replaces the inner electrons with the symbol of the nearest preceding noble gas in square brackets.
For example, potassium: full = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹; shorthand = [Ar] 4s¹.
True or False?
Transition metals lose electrons from the 3d subshell before the 4s subshell when forming positive ions.
False.
Transition metals lose electrons from the 4s subshell first, not the 3d. Although the 4s is filled before the 3d when building up the configuration, it is the outer subshell and so electrons are removed from it first when forming positive ions.
What are the exceptional electron configurations of chromium and copper, and why do they occur?
Chromium is [Ar] 3d⁵ 4s¹ and copper is [Ar] 3d¹⁰ 4s¹, rather than the expected [Ar] 3d⁴ 4s² and [Ar] 3d⁹ 4s².
These configurations are energetically stable because half-filled and fully-filled d subshells have lower energy.
What are s-block, p-block, d-block and f-block elements?
s-block elements have their valence electrons in an s orbital. p-block elements have their valence electrons in a p orbital. d-block elements have their valence electrons in a d orbital. f-block elements have their valence electrons in an f orbital.
All elements in the same ..........of the Periodic Table have the same number of electron shells. All elements in the same ..........have the same number of outer electrons.
All elements in the same period of the Periodic Table have the same number of electron shells. All elements in the same group have the same number of outer electrons.
Why do noble gases rarely react with other elements?
Noble gases rarely react because they have full subshells in their electronic configuration, which gives them very high ionisation energies. This stable arrangement means there is little tendency to gain, lose or share electrons.
True or False?
Elements in the same group have the same number of outer electrons.
True.
Elements in the same group of the Periodic Table all have the same number of outer (valence) electrons. This shared configuration explains the similar chemical properties observed within a group.

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