Chapter 7 - Periodicity Flashcards

The periodic table, Ionisation energies and Periodic trends in bonding and structure.

1
Q

Who created the modern periodic table?

A

Dmitri Mendeleev

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2
Q

How were the elements ordered by Mendeleev?

A

Increasing atomic mass

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3
Q

What is the name for the vertical columns of the periodic table?

A

Groups

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4
Q

What is the name for the horizontal rows of the periodic table?

A

Periods

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5
Q

What is the periodic trend in electron configuration?

A

The sub shells of n energy level fill up.

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6
Q

What is ionisation?

A

The removal of one or more electrons from an atom.

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7
Q

Define first ionisation energy

A

The energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.

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8
Q

How does atomic radius affect ionisation energy?

A

Greater distance between the nucleus and outer electrons, the attraction is lower.

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9
Q

How does nuclear charge affect ionisation energy?

A

More protons creates a greater attraction between the nucleus and the outer electrons.

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10
Q

How does electron shielding affect ionisation energy?

A

Inner shell electrons repel outer shell electrons, called shielding. This reduces the attraction between the nucleus and the outer electrons.

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11
Q

Why does ionisation energy decrease going down a group?

A

More electrons shells so the outer electrons are further away and there is a greater shielding effect.

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12
Q

Why does ionisation energy increase across a period?

A

The number of protons in the nucleus increases so nuclear charge increases causing atomic radius to decrease, whilst shielding stays the same.

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13
Q

Why do successive ionisation energies increase?

A

There are less electrons so the nuclear attraction on the remaining electrons will be greater.

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14
Q

Define second ionisation energy

A

The energy required to remove one electron from each ion in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions.

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15
Q

What causes the large jumps in successive ionisation energies?

A

Moving down to a closer shell, as these electrons are closer so experience a greater nuclear attraction.

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16
Q

What predictions can be made from a graph of successive ionisation energies?

A

The number of electrons in the outer shell, the group of the element in the periodic table and thus the identity of the element.

17
Q

Explain the trend of first ionisation energy down a group

A

Atomic radius increases,
More inner shells so shielding increases,
Nuclear attraction on outer electrons decreases,
First ionisation energy decreases.

18
Q

Explain the general trend of first ionisation energy across a period

A
Nuclear charge increases,
Same shell: similar shielding,
Nuclear attraction increases,
Atomic radius decreases,
First ionisation energy increases.
19
Q

In period 2, explain the fall from beryllium to boron of first ionisation energies

A

The new electron enters the 2p sub shell, which is slightly further away from the nucleus than the 2s sub shell.

20
Q

In period 2, explain the fall from nitrogen to oxygen

A

Nitrogen’s electrons in the 2p sub shell are unpaired so oxygen’s 8th electron is paired, causing repulsion and a lower ionisation energy.

21
Q

Explain the trend of atomic radii decreasing across a group

A

Nuclear charge increases,
Nuclear attraction increases,
Atomic radius decreases.

22
Q

What is metallic bonding?

A

Each atom donates an outer shell electron, which becomes delocalised. This creates cations.

23
Q

What causes metallic bonding?

A

Strong electrostatic attraction between the fixed cations and the delocalised electrons.

24
Q

What are common properties of metals?

A

Strong metallic bonds
High electrical conductivity
High melting and boiling points

25
Q

Why does the melting point and boiling point increase across the metals of a period?

A

Number of delocalised electrons per atom and charge on cation increase, so stronger electrostatic attraction.

26
Q

In period 3, why does silicon have the highest melting point?

A

Forms a giant covalent lattice, where each atom is covalently bonded to four others.

27
Q

What are the properties of giant covalent lattices?

A

High melting and boiling points
Insoluble in almost all solvents
Do not conduct (except for graphite and graphene)

28
Q

Why are giant covalent lattices insoluble?

A

Covalent bonds holding it together are too strong to be broken by interaction with solvents.

29
Q

Why do most giant covalent lattices not conduct electricity?

A

All four outer shell electrons are involved in covalent bonding.

30
Q

Why do simple molecules have low melting points?

A

Weak induced dipole-dipole forces between molecules are easy to break.