Periodicity is the repeating patterns in physical and chemical properties of elements observed across successive periods of the periodic table.

What does the period number indicate about an element's electron configuration?

The period number gives the highest energy shell occupied by electrons in that element's atoms.

Define valence electrons

The electrons in the outermost shell of an atom that determine its chemical properties and are involved in bonding.

Define s-block elements

Elements whose highest-energy electrons occupy only s sub-shells , found in Groups 1 and 2 of the periodic table.

Elements with at least one p-electron in the outer shell belong to the .......... -block. Elements with at least one d-electron and at least one s-electron (but no f or p electrons) in the outer shell belong to the .......... -block.

Elements with at least one p-electron in the outer shell belong to the p -block. Elements with at least one d-electron and at least one s-electron (but no f or p electrons) in the outer shell belong to the d -block.

True or False? Elements in the same group of the periodic table have the same number of electron shells.

False. Elements in the same group share a similar outer-shell electron configuration , not the same number of shells. The number of shells increases down a group.

Why are hydrogen and helium not assigned to groups in the same way as other elements?

Their electron configurations are unusual and they do not fit comfortably into any group, so they are placed according to similarities in physical and chemical properties .

Define first ionisation energy

First 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.

What three factors determine the ionisation energy of an element?

Atomic radius Nuclear charge Electron shielding Together these determine the attraction between the nucleus and outer electrons.

Why does first ionisation energy decrease down a group despite increasing nuclear charge?

Increasing atomic radius and greater electron shielding by inner shells reduce the attraction between the nucleus and outer electrons, outweighing the effect of higher nuclear charge.

Identify the anomalies in the first ionisation energy trend across Period 3. IE1 of aluminium is .......... than that of magnesium because the outer electron in Al is in the .......... sub-shell, which is at a higher energy level. IE1 of sulfur is .......... than that of phosphorus because the paired electrons in sulfur's 3p sub-shell experience greater .......... .

IE1 of aluminium is lower than that of magnesium because the outer electron in Al is in the 3p sub-shell, which is at a higher energy level. IE1 of sulfur is lower than that of phosphorus because the paired electrons in sulfur's 3p sub-shell experience greater electron-electron repulsion .

True or False? A large jump between successive ionisation energies indicates that the next electron is being removed from a shell closer to the nucleus.

True. A large jump in successive ionisation energies occurs when an electron is removed from an inner shell, which is closer to the nucleus and experiences less shielding .

Define electron shielding

Electron shielding is the reduction in the effective nuclear charge felt by outer electrons, caused by repulsion from electrons in full inner shells.

How can successive ionisation energy data be used to identify the group of an element?

The position of the large jump in successive ionisation energies indicates when an inner shell is reached, revealing the number of outer-shell electrons and therefore the group number.

True or False? First ionisation energy increases steadily across a period with no exceptions.

False. There are two anomalies: IE1 drops from Group 2 to Group 13 (s to p sub-shell transition) and from Group 15 to Group 16 (paired electrons in p sub-shell increase repulsion ).

Define Metallic bonding

Metallic bonding is the strong electrostatic attraction between a lattice of positive metal ions and a surrounding 'sea' of delocalised electrons .

Define Giant covalent substance

A giant covalent substance is a substance in which atoms are joined by covalent bonds that extend indefinitely throughout the structure, forming a large lattice with no discrete molecules.

Why does diamond have an extremely high melting point and is very hard?

Each carbon atom is covalently bonded to four others in a tetrahedral arrangement, creating a 3D network of strong bonds that requires a very large amount of energy to break.

True or False? Graphite conducts electricity but diamond does not.

True. Graphite has delocalised electrons between its carbon layers that are free to carry charge, whereas in diamond all four outer electrons on every carbon are involved in covalent bonds , leaving none free.

Melting point increases from Group 1 to Group 4 (14) because metallic bonding strengthens and Group 4 has a .......... covalent structure. Melting point then drops sharply from Group 4 to Group 5 (15) because elements from Group 5 to Group 0 have .......... molecular structures with weak .......... forces.

Melting point increases from Group 1 to Group 4 (14) because metallic bonding strengthens and Group 4 has a giant covalent structure. Melting point then drops sharply from Group 4 to Group 5 (15) because elements from Group 5 to Group 0 have simple molecular structures with weak London forces.

Describe the structure of graphite and explain why its layers can slide over each other.

Each carbon atom is bonded to three others in hexagonal layers held together by weak London forces , which require little energy to overcome, allowing the layers to slide.

True or False? Metals conduct electricity in the solid state but not when molten.

False. Metals conduct electricity in both the solid and liquid states because delocalised electrons remain free to move and carry charge in both states.

What is the structure of silicon(IV) oxide and why does it have a high melting point?

Silicon(IV) oxide has a giant covalent lattice similar to diamond, with strong covalent bonds throughout the structure requiring a large amount of energy to break.

A LevelChemistryOCRFlashcards3. Periodic Table & EnergyPeriodicity

Periodicity (OCR A Level Chemistry A): Flashcards

Exam code: H432

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Define periodicity
Periodicity is the repeating patterns in physical and chemical properties of elements observed across successive periods of the periodic table.
What does the period number indicate about an element's electron configuration?
The period number gives the highest energy shell occupied by electrons in that element's atoms.
Define valence electrons
The electrons in the outermost shell of an atom that determine its chemical properties and are involved in bonding.
Define s-block elements
Elements whose highest-energy electrons occupy only s sub-shells, found in Groups 1 and 2 of the periodic table.
Elements with at least one p-electron in the outer shell belong to the ..........-block. Elements with at least one d-electron and at least one s-electron (but no f or p electrons) in the outer shell belong to the ..........-block.
Elements with at least one p-electron in the outer shell belong to the p-block. Elements with at least one d-electron and at least one s-electron (but no f or p electrons) in the outer shell belong to the d-block.
True or False?
Elements in the same group of the periodic table have the same number of electron shells.
False.
Elements in the same group share a similar outer-shell electron configuration, not the same number of shells. The number of shells increases down a group.
Why are hydrogen and helium not assigned to groups in the same way as other elements?
Their electron configurations are unusual and they do not fit comfortably into any group, so they are placed according to similarities in physical and chemical properties.
Define first ionisation energy
First 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.
What three factors determine the ionisation energy of an element?
1. Atomic radius
2. Nuclear charge
3. Electron shielding
Together these determine the attraction between the nucleus and outer electrons.
Why does first ionisation energy decrease down a group despite increasing nuclear charge?
Increasing atomic radius and greater electron shielding by inner shells reduce the attraction between the nucleus and outer electrons, outweighing the effect of higher nuclear charge.
Identify the anomalies in the first ionisation energy trend across Period 3.
1. IE1 of aluminium is .......... than that of magnesium because the outer electron in Al is in the .......... sub-shell, which is at a higher energy level.
2. IE1 of sulfur is .......... than that of phosphorus because the paired electrons in sulfur's 3p sub-shell experience greater .......... .
1. IE1 of aluminium is lower than that of magnesium because the outer electron in Al is in the 3p sub-shell, which is at a higher energy level.
2. IE1 of sulfur is lower than that of phosphorus because the paired electrons in sulfur's 3p sub-shell experience greater electron-electron repulsion.
True or False?
A large jump between successive ionisation energies indicates that the next electron is being removed from a shell closer to the nucleus.
True.
A large jump in successive ionisation energies occurs when an electron is removed from an inner shell, which is closer to the nucleus and experiences less shielding.
Define electron shielding
Electron shielding is the reduction in the effective nuclear charge felt by outer electrons, caused by repulsion from electrons in full inner shells.
How can successive ionisation energy data be used to identify the group of an element?
The position of the large jump in successive ionisation energies indicates when an inner shell is reached, revealing the number of outer-shell electrons and therefore the group number.
True or False?
First ionisation energy increases steadily across a period with no exceptions.
False.
There are two anomalies: IE1 drops from Group 2 to Group 13 (s to p sub-shell transition) and from Group 15 to Group 16 (paired electrons in p sub-shell increase repulsion).
Define Metallic bonding
Metallic bonding is the strong electrostatic attraction between a lattice of positive metal ions and a surrounding 'sea' of delocalised electrons.
Define Giant covalent substance
A giant covalent substance is a substance in which atoms are joined by covalent bonds that extend indefinitely throughout the structure, forming a large lattice with no discrete molecules.
Why does diamond have an extremely high melting point and is very hard?
Each carbon atom is covalently bonded to four others in a tetrahedral arrangement, creating a 3D network of strong bonds that requires a very large amount of energy to break.
True or False?
Graphite conducts electricity but diamond does not.
True.
Graphite has delocalised electrons between its carbon layers that are free to carry charge, whereas in diamond all four outer electrons on every carbon are involved in covalent bonds, leaving none free.
Melting point increases from Group 1 to Group 4 (14) because metallic bonding strengthens and Group 4 has a .......... covalent structure.
Melting point then drops sharply from Group 4 to Group 5 (15) because elements from Group 5 to Group 0 have .......... molecular structures with weak .......... forces.
Melting point increases from Group 1 to Group 4 (14) because metallic bonding strengthens and Group 4 has a giant covalent structure.
Melting point then drops sharply from Group 4 to Group 5 (15) because elements from Group 5 to Group 0 have simple molecular structures with weak London forces.
Describe the structure of graphite and explain why its layers can slide over each other.
Each carbon atom is bonded to three others in hexagonal layers held together by weak London forces, which require little energy to overcome, allowing the layers to slide.
True or False?
Metals conduct electricity in the solid state but not when molten.
False.
Metals conduct electricity in both the solid and liquid states because delocalised electrons remain free to move and carry charge in both states.
What is the structure of silicon(IV) oxide and why does it have a high melting point?
Silicon(IV) oxide has a giant covalent lattice similar to diamond, with strong covalent bonds throughout the structure requiring a large amount of energy to break.

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