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chemistry 🫧 > Chapter 7 - Periodicity > Flashcards

Chapter 7 - Periodicity Flashcards

1
Q

What order are the elements arranged in?

A

Increasing atomic number

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2
Q

What do the groups in the periodic table tell you?

A
  • Number of electrons in the outer shell
  • Elements that have similar chemical properties
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3
Q

What do the periods in the periodic table tell you?

A

Number of the highest energy electron shell

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4
Q

What is periodicity?

A

Repeating trend in properties of the elements across a period

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5
Q

Which groups are the s-block elements?

A
  • Group 1 and 2
  • Helium
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6
Q

Which groups are the d-block elements?

A

Transition metals

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7
Q

Which groups are the p-block elements?

A

Groups 3, 4, 5, 6, 7, 8

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8
Q

What is the first ionisation energy?

A

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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9
Q

What is the first ionisation energy of Na(g) ?

A

Na(g) -> Na+ (g) + e-

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10
Q

First electron lost will be the _________ energy level and will experience the _________ attraction from the nucleus

A
  • Highest
  • Least
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11
Q

What happens to first ionisation energy when atomic radius increases? Explain why

A

Greater distance between nucleus and outer electrons
Less nuclear attraction
Lower ionisation energy

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12
Q

What happens to first ionisation energy when nuclear charge increases? Explain why

A

More protons
Greater attraction between nucleus and outer electron
Higher ionisation energy

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13
Q

What happens to first ionisation energy when there is more electron shielding Explain why

A

Shielding effect increases when number of inner shells increase
Less attraction between nucleus and outer electrons
Lower ionisation energy

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14
Q

Give the general relationship between nuclear attraction and first ionisation energy?

A

Less attraction between nucleus and outer electrons
Lower ionisation energy

Greater attraction between nucleus and outer electrons
Higher ionisation energy

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15
Q

What is the shielding effect?

A

Inner shell electrons repel outer shell electrons

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16
Q

Element has as many ionisation energies as there are ___________

A

Electrons

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17
Q

What is second ionisation energy?

A

Energy required to remove one electron from each atom in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions

18
Q

Why are second ionisation energies greater than the first?

A

After first electron is lost, remaining electron is pulled closer to nucleus, stronger nuclear attraction = more energy needed to remove electron = higher ionisation energy

19
Q

Electrons with _________ ionisation energies are ___________ to the nucleus

A

Higher
Closer

20
Q

What does the large difference/jump in ionisation energies represent?

A

Change from one shell to another

21
Q

What is the trend in first ionisation energies down a group?

A
  • Atomic radius increases
  • More inner shells so shielding increases
  • Nuclear attraction decreases
  • First ionisation energy decreases

s.n. nuclear charge increasing is outweighed by the effects of increased radius and (to a lesser extent) increased shielding

22
Q

What is the trend in first ionisation energies across a period?

A
  • Nuclear charge increases
  • Same shell numbers so similar shielding
  • Nuclear attraction increases
  • Atomic radius decreases
  • First ionisation energy increases
23
Q

Why does boron have a lower first ionisation energy than beryllium? (hint: sub-shell is the different)

A
  • Start of filling of 2p sub shell
  • 2p sub shell is higher energy than 2s sub shell
  • 2p electron easier to remove
24
Q

Why does oxygen have a lower first ionisation energy than nitrogen? (hint: sub-shell is the same)

A
  • Start of electron pairing in the p-orbitals of 2p sub shell
  • In O the paired electrons in one of the 2p orbitals repel each other
  • Easier to remove one of the paired electrons from O
25
Q

Why does aluminium have a lower first ionisation energy than magnesium? (hint: sub-shell is the different)

A
  • 3p sub-shell in Al has a higher energy level than 3s sub-shell in Mg
  • Easier to remove 3p electron from Al
26
Q

Why does sulphur have a lower first ionisation energy than phosphorus? (hint: sub-shell is the same)

A
  • P has 3 electrons in 3p sub shell which are all unpaired
  • S has 4 electron in 3p sub shell which has one pair in an orbital
  • Paired electrons repel so easier to remove an electron from S
27
Q

Describe metallic bonding?

A

Strong electrostatic attraction between cations and delocalised electrons

28
Q

Describe the electrical conductivity of metallic structures?

A

Conduct electricity in solid and liquid states since the delocalised electrons are able to carry charge

29
Q

Describe the melting and boiling points of metals?

A

High M.P. and B.P. as lots of energy needed to break strong electrostatic attraction

30
Q

Describe the solubility of metals?

A

Do not dissolve in water

31
Q

Describe giant covalent structures?

A

Boron, carbon and silicon atoms held together by covalent bonds to form giant covalent structures

32
Q

What is the shape of the structures that carbon and silicon atoms form?

A

Tetrahedral

33
Q

Describe the electrical conductivity of giant covalent structures?

A

Non conductors of electricity (except graphene and graphite)

34
Q

Why can diamond not conduct electricity?

A

All four outer shell electrons involved in bonding so no delocalised electrons to carry charge

35
Q

Why can graphene (single layer of graphite) conduct electricity?

A

Only three of the carbon outer shell electrons are involved in bonding so there are delocalised electrons to carry charge

36
Q

Describe the melting and boiling points of giant covalent structures?

A

High M.P. and B.P. as lots of energy needed to break strong covalent bonds

37
Q

Describe the solubility of giant covalent structures?

A

Insoluble in almost all solvents as covalent bonds are too strong to be broken

38
Q

What is the general trend in melting points across period 2 and 3?

A
  • Melting point increases from group 1-4
  • Sharp decrease in melting point from group 4-5
  • Melting point lower from group 5-0
39
Q

Why is there a sharp decrease in melting point from group 4-5?

A

Marks change from giant to simple molecular structures

40
Q

What structure to metals form?

A

Giant metallic lattice structures

chemistry 🫧 (21 decks)

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