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CHEMISTRY > CHAPTER 7 PERIODIC TABLE & ENERGY > Flashcards

CHAPTER 7 PERIODIC TABLE & ENERGY Flashcards

1
Q

What is the first ionisation energy?

A

The first ionisation energy is 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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2
Q

How do you display first ionisation energy?

A

Na(g) —> Na+(g) + e-

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3
Q

What are the three factors that affect the attraction between the nucleus and outermost electrons?

A

atomic radius, nuclear charge and electron shielding.

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4
Q

What happens across a period?

A
  • The atomic radius decreases
  • The nuclear charge increases
  • Shielding remains the same
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5
Q

What is periodicity?

A

A repeating trend in properties (of elements)

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6
Q

How is the periodic table divided into blocks?

A

Group 1-2 = s-block
Group 3-12= d-block
Group 13-18= p-block
Period 6-7= f-block

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7
Q

What are groups?

A

Elements with similar outer shell electron structures resulting in similar properties.

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8
Q

What are the electron configurations for the noble gases?

A 
Helium = 2
Neon = 10
Argon = 18
Krypton = 36
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9
Q

What happens to the first electrons lost?

A

It comes from the highest energy level and experiences the least attraction from the nucleus.

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10
Q

How many ionisation energies are there?

A

As many electrons are there eg helium has two electrons so it has 2 ionisation energies

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11
Q

Why is the second ionisation energy stronger than the first?

A

Because, after the first electron is lost, the electrons are pulled closer to the nucleus so the nuclear atrraction increases and more ionisation energy will be needed to remove the next electron.

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12
Q

What is the second ionisation energy?

A

The amount of energy required to remove an 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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13
Q

What dramatically increases the ionisation energy of an element?

A

The difference of shells. Eg the ionisation energy between a 2p and 3s would be dramatic compared to 3s and 3p.

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14
Q

What are factors effecting ionisation energies?

A
  • Nuclear charge (The greater the nuclear charge the more ionisation energy needed)
  • Shielding ( The more electron shields the lower the ionisation energy)
  • Atomic radius ( The larger the atomic radius the lower ionisation energy needed)
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15
Q

What is the difference between ionisation energy in groups and periods?

A

In periods the ionisation energy increases, in groups the ionisation energy decreases.

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16
Q

What is the trend in ionisation energy down a group?

A
  • Atomic radius increases
  • Shielding increases
  • Nuclear charge decreases
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17
Q

What are state symbols?

A

Symbols to state the state of matter the element is in eg g for gas

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18
Q

Why does Al have a lower first ionisation energy than Mg?

A

Because the paired electrons in the … subshell repel each other making it easier to remove an electron from Al than Mg therefore lowering the ionisation energy

19
Q

What happens with room temperature and metals?

A

Metals become solid.

20
Q

What is metallic bonding?

A

Metallic bonding is the strong electrostatic attraction between cations and delocalised electrons.

21
Q

What is the structure in metallic bonding?

A
  • Cations are in a fixed position maintaining the structure and shape of the metal
  • Delocalised electrons are mobile and able to move throughout the structure
22
Q

What is the structure of a metallicly bonded element called?

A

Giant metallic lattice

23
Q

What is a lattice ?

A

A lattice is a structure with a regular pattern of oppositely charged ions.

24
Q

What are the properties of metals?

A
  • Strong metallic bonds
  • High electrical conductivity
  • High melting and boiling points
25
Q

How would you explain the physical properties of metals?

A

Giant metallic lattice or metallic bonding.

26
Q

What states do metals conduct electricity in?

A

Solid and liquid.

27
Q

What are all metals called?

A

Giant metallic lattices

28
Q

What happens when a voltage is applied across a metal?

A

The delocalised electrons can move through the structure carrying a charge

29
Q

What is the melting point of metals dependent on?

A

The strength of the metallic bond

30
Q

Are metals soluble?

A

No

31
Q

What is the structure of non-metallic elements called and how is it compared to metallic structures?

A

Simple molecular lattice and they have weak intermolecular forces meaning they have low melting and boiling points.

32
Q

Which non-metals have a giant covalent lattice?

A

Boron, carbon and silicon.

33
Q

What is the structure of carbon and silicon?

A

Carbon and silicon are in group 14 and have 4 electrons on their outer shell and use these four electrons to covalently bond with other carbon or silicon atoms which results in a tetrahedral structure with a bond angle of 109.5.

34
Q

What are the properties of giant covalent structures?

A
  • High melting and boiling points. This is neccessary to provide a large quantity of energy needed to break the strong covalent bonds.
  • Insoluble because of the strength of the covalent bonds.
  • They don’t conduct electricity other than graphite and graphine
35
Q

Why can’t carbon and silicon conduct electricity?

A

Because all their electrons are used in covalent bonding.

36
Q

Why can graphite and graphene conduct electricity?

A

Because they have one electron remaining that is released into a pool of delocalised electrons. The structure is also planar hexagonal and therefore good electrical conductors.

37
Q

What is the bond angle for graphene and graphite?

A

120 degrees.

38
Q

What is graphene?

A

A single layer of graphite.

39
Q

What is the structure of graphite?

A

Graphite has parallel layers of hexagonal arranged carbon atoms bonded by weak London forces. These hexagonal layers only use 3/4 electrons that carbon has so the other is used to conduct electricity.

40
Q

What are the trends in melting points across period 2 and 3?

A

Melting point increases from group 1-14, sharp decrease between 14 and 15 and extremely low from group 15 to 18.

41
Q

Why is there a sharp decrease in melting point across period 2 and 3?

A

Because it changes from giant covalent and metallic structures to giant molecular structures.

42
Q

What is the difference in bonding between simple molecular lattices and giant covalent lattices?

A

Simple molecular lattices are bonded together by weak intermolecular forces whereas giant covalent lattices are bonded by covalent bonds

43
Q

What are the properties of graphite?

A
  • Soft and slippery
  • Conducts Electricity
  • Has a high melting point
44
Q

What is the periodic table?

A

The arrangement of elements by increasing atomic number

CHEMISTRY (7 decks)

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