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YR2 Chemistry 24-25 > M3 - Periodicity, Groups 2 and 7 > Flashcards

M3 - Periodicity, Groups 2 and 7 Flashcards

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1
Q
A
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2
Q
A

C

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3
Q

Are halogens oxidising or reducing agents?

A

Oxidising agents, as they themselves are reduced (electron acceptors)

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4
Q
A

A

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5
Q

Why are giant covalent lattices insoluble?

A

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

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6
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)
formed from carbon, silicon anf boron

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7
Q

How were the elements ordered by Mendeleev?

A

Increasing atomic mass

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8
Q
A

B

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9
Q

what is produced when Ba is added to water?

A

Barium hydroxide and hydrogen gas

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10
Q

what is the colour of iodine in an organic solvent?

A

Purple

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11
Q

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

A

Periods

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12
Q

Define First Ionisation Energy (2)

A

The energy required to remove one electron from each atom (1) in one mole of gaseous atoms to form one mole of gaseous 1+ ions (1)

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13
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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14
Q

What block is Lithium placed in and why?

A

Li is an S-block element - A s-block element has its highest energy/outermost electron in a s subshell/orbital

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15
Q

What is ionisation?

A

The removal of one or more electrons from an atom.

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16
Q

Which is the most reactive halogen and why?

A

Fluorine, as it has a small atomic radius and less shielding. Easier for fluorine to gain a electron

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17
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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18
Q

Complete the word equation;
HEAT
Calcium carbonate ->

A

Calcium oxide and carbon dioxide

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19
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. therefore lower ionisation energy

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20
Q

In period 2, explain the fall of ionisation energy 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.

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21
Q

what is the colour of aqueous chlorine?

A

Pale green

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22
Q

Why do most giant covalent lattices not conduct electricity?

A

All four outer shell electrons are involved in covalent bonding. No mobile charge carriers.

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23
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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24
Q

Explain the trend of first ionisation energy down a group

A

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

25
Q

What is Successive Ionisation Energy?

A

A measure of the energy required to remove each electron in turn

26
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.

27
Q

What are common properties of metals?

A

Strong metallic bonds
High electrical conductivity
High melting and boiling points

28
Q

What is meant by the term, Periodicity? (1)

A

The repeating trends in physical and chemical properties (1)

29
Q
A

D

30
Q

what is the colour of aqueous bromine?

A

Pale orange/brown

31
Q
A

B

32
Q

What is Disproportionation? (1)

A

The (simultaneous) oxidation and reduction of the same element (in the same redox reaction) (1)

33
Q

How does nuclear charge affect ionisation energy?

A

More protons creates a greater attraction between the nucleus and the outer electrons. therefore higher ionisation energy

34
Q

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

A

Groups

35
Q

what block are group 2 metals found in and why?

A

s block, as the highest energy/ outer electron is found within the s sub-shell / orbital

36
Q

Who created the modern periodic table?

A

Dmitri Mendeleev

37
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.

38
Q

what block are group 7 elements found in and why?

A

p block, as the highest energy/ outer electron is found within the p sub-shell / orbital

39
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.

40
Q

How are elements arranged in the Periodic table? (2)

A

Elements are arranged by:
– Increasing atomic (proton) number (1) – In groups with similar chemical properties (1)

41
Q

Define metallic bond

A

Strong electrostatic attraction between the cations and the delocalised electrons.

42
Q

Define Second Ionisation Energy.

A

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

43
Q

Why do successive ionisation energies increase?

A

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

44
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.
45
Q
A
46
Q
A

C

47
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. Strong covalent bonds.

48
Q

Why do simple covalent molecules have low melting / boiling points?

A

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

49
Q

what is meant by thermal decomposition?

A

using heat energy to break bonds and break the substance down

50
Q

what is the colour of aqueous iodine?

A

Pale brown

51
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. Stronger metallic bonds.

52
Q

what is produced when CaO is added to water?

A

Calcium hydroxide

53
Q
A

A

54
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.

55
Q

What is metallic bonding?

A

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

56
Q

Explain the trend of atomic radii across a period

A

Nuclear charge increases,
Nuclear attraction increases,
Atomic radius decreases across a period.

57
Q

What is the periodic trend in electron configuration?

A

The sub shells of n energy level fill up.

58
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.

YR2 Chemistry 24-25 (15 decks)

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