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Unit 1 Chemistry > 2. The Periodic Table: Bonding and Structure > Flashcards

2. The Periodic Table: Bonding and Structure Flashcards

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1
Q

How do outer electrons become delocalised?

A

The OUTER electrons of the metal ATOMS are loosely held which means they can become DELOCALISED.

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2
Q

How does metallic bonding occur? What is the definition?

A

metal atoms LOSING their outer ELECTRONS to form a “common pool” of DELOCALISED electrons.

the ATTRACTION of these POSITIVELY charged ions for the delocalised ELECTRONS.

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3
Q

What can metals be considered as?

A

a giant lattice of POSITIVE ions held together by these DELOCALISED electrons

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4
Q

In order to melt or boil a metal you ..?

A

the METALLIC bonds need to be broken.

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5
Q

High melting points mean?

A

STRONG metallic bonds.

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6
Q

What is the trend for the boiling point (metallic bonding) of group one?

A

METALLIC BONDING DECREASES

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7
Q

What is the full explanation of boiling point (metallic bonding) in group one?

A

As you go down the group the boiling point DECREASES. So the metallic bonding must get WEAKER.

As you go down the group the outer ELECTRONS are FURTHER away from the nucleus (the number of energy levels has INCREASED).

The strength of the metallic bond gets WEAKER as the outer electrons are further away from the POSITIVE charge of the NUCLEUS.

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8
Q

What is the trend for boiling point (metallic bonding) in metals across a period?

A

metallic bonding INCREASES.

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9
Q

What is the full explanation of metallic bonding across a period

A

As you go across a period the metallic bonding gets STRONGER(as the boiling point is INCREASING ).

As you go across the period the number of OUTER electrons is INCREASING meaning there are more DELOCALISED electrons

so the metallic bond is STRONGER.

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10
Q

Noble gases are said to be?

A

MONATMOIC. existing as SINGLE ATOMS

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11
Q

What does monatomic mean?

A

They only exist as single atoms (noble gases)

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12
Q

Forces between monatomic noble gases are?

A

They are very low. This means the forces between are very weak

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13
Q

How do London Dispersion Forces arise?

A

The ELECTRONS in an atom may become unevenly distributed causing a TEMPORARY dipole

one side of the atom becomes slightly NEGATIVE ( δ- ) while the other side becomes slightly POSITIVE ( δ+ ).

These dipoles only exist for a fraction of a second but this has a knock on effect on the neighbouring atoms.

The δ- on one atom attracts the δ+ on a neighbouring atom. It’s this weak attraction BETWEEN the atoms that is the LONDON DISPERSION FORCE

The temporary dipole in one atom INDUCES a temporary dipole in a neighbouring atom so you get the attraction between the δ+ and δ- in the neighbouring atoms.

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14
Q

Where do LDF’s exist?

A

BETWEEN particles

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15
Q

Compared to other intermolecular forces, where to LDF’s lie

A

They are the weakest compared to the 2 others

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16
Q

As the size of atoms increase, what happens? Why?

A

The strength of LDF’s increases as the size of the atoms increase.

This is because as the atomic number increases there are more ELECTRONS.

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17
Q

What is a covalent bond?

A

A covalent bond is the electrostatic attraction between two positive nuclei and a shared pair of electrons.

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18
Q

The simplest molecular elements are?

A

diatomic (2 atoms joined together)

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18
Q

The simplest molecular elements are?

A

diatomic (2 atoms joined together)

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19
Q

What are the diatomic elements?

A

H2, N2, O2, F2, Cl2, Br2 and I2

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20
Q

Where are LDF’s located?

A

These weak forces exist between all atoms and molecules

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21
Q

What can LDF’s tell about a substance?

A

strong enough to determine the physical properties (e.g melting/boiling point) of the molecular elements.

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22
Q

What happens to the halogens when you go down a group?

A

The melting point increases.

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23
Q

Give a full explanation of what happens when you go down group 7 (halogens)

A

The melting point INCREASES as you go down the GROUP.

This is because as the ATOMIC number increases the number of ELECTRONS present in the atoms also INCREASES.

As there are more ELECTRONS present within the diatomic molecules there are stronger INTER MOLECULAR forces and so more energy is needed to break them.

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24
Q

When covalent molecular substances are melted/boiled what happens?

A

no COVALENT bonds are actually broken.

It is only the weaker intermolecular forces (LONDON DISPERSION FORCES) that are actually broken.

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25
Q

What are the two ‘special’ molecules?

A

Sulfur and Phosphorous

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26
Q

What so phosphorus and sulfur exist as?

A

DISCRETE molecules

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27
Q

What are the formulas for sulfur and phosphorous

A

P4 and S8

28
Q

As a molecule gets bigger?

A

the melting point INCREASES

29
Q

What is the full explanation of melting point increasing?

A

As the molecule get bigger the melting point INCREASES.

This is because there are more (and larger) atoms present in the molecule so there are more electrons present.

As there are more electrons there are more LDFs (London Dispersion Forces) present which means it takes more energy to separate the molecules hence the increase in melting point.

30
Q

When you melt a substance, are covalent bonds broken? Why not?

A

It’s only the weak intermolecular LDFs that are broken no covalent bonds get broken when you melt or boil a molecular substance.

31
Q

What are the 3 forms of carbon?

A

diamond, graphite and fullerene

32
Q

What is a fullerene?

A

Fullerenes exist as covalent molecules with a definite formula.

33
Q

What is a fullerene that you know of? What is its formula?

A

C60 - buckministerfullerene

34
Q

What do the structures of fullerenes look like?

A

They are a large family of ‘carbon cage’ molecules each made up of rings of 5 and 6 carbon atoms.

They can also exist as nanotubes

35
Q

In groups 5, 6 and 7 what are the forces present in the molecule?

A

Intramolecular are all strong COVALENT bonds

Intermolecular are weak LDF’s

36
Q

What are the two types of covalent elements?

A

Covalent Molecular

Covalent Network

37
Q

What are Covalent Networks? How hard is it to break them?

A

Covalent Networks are huge structures with every atom being linked to other atoms by COVALENT BONDS.

As COVALENT bonds are very STRONG it takes a large amount of energy to BREAK the bonds.

This results in covalent network substances having very HIGH melting and boiling points.

38
Q

What are the two DISTINCT forms of graphite?

A

DIAMOND and GRAPHITE

39
Q

Explain in detail the structure of diamond

A

Each C atom forms 4 covalent bonds with neighbouring ATOMS.

All 4 outer electrons are used in bonding, there is no delocalised electrons, so diamond does NOT CONDUCT electricity

Strong covalent bonds are formed in 3D, they are tetrahedral

To cut or break diamond means that you have to break strong COVALENT bonds. It needs a lot of energy. The 3D network makes diamond very HARD.

40
Q

Explain in detail the structure of graphite

A

Each C atom forms 3 covalent bonds

Only 3 outer electrons are used in bonding, each C atom has 1 outer electron, so it CAN CONDUCT electricity

The covalent bonds are formed in 2D, so in SHEETS.

Between each sheet there is weak LDF’s present

The layers of C atoms can slip and slide, LUBRICANT

To cut or break graphite means breaking LDF’s which are very weak, so it is easy to break

Graphite is very brittle

41
Q

What are the elements that have a network structure?

A

BORON and SILICON

42
Q

What are the metallic elements?

A

Lithium, Beryllium, Sodium, Magnesium, Potassium, Calcium, Aluminium

43
Q

What are the covalent network elements?

A

Boron, Carbon and Silicon

44
Q

What are the covalent molecular elements?

A

Hydrogen, Nitrogen, Carbon, Phosphorus, Oxygen, Sulfur, Fluorine, Chlorine

45
Q

What are monatomic elements?

A

Helium, Neon, Argon

46
Q

What is covalent radius?

A

half the distance between the 2 nuclei in a covalent bond

47
Q

What is covalent radius measured in?

A

picometers

48
Q

As you go across a period, what happens to the atomic size? Why?

A

It DECREASES.

as the atomic number increases the number of protons in the nucleus is increasing so the electrons are pulled in tighter causing the atoms to get smaller.

49
Q

As you go down a group, what happens to the atomic size? Why?

A

As you go down a group the atomic size INCREASES.

This is because as you down the group you are adding extra shells of electrons so the atoms are bigger.

Sodium (2, 8, 1) is smaller than potassium (2, 8, 8, 1) as there is an extra shell of electrons in potassium.

50
Q

What is density?

A

mass per volume

51
Q

The greater the mass of an atom?

A

the greater the density

52
Q

The larger the atom?

A

The lower the density.

53
Q

When does atomic size only have an effect?

A

on solids and liquids i.e. when the atoms are CLOSE TOGETHER.

54
Q

Explain group 1’s density and atomic size

A

Group 1 has a LOW DENSITY, because they have a LARGE ATOMIC SIZE

55
Q

When you go across a period, what happens to the density in metals? Why?

A

the density INCREASES.

this is because the relative atomic mass is INCREASING and the covalent radius is DECREASING.

56
Q

What happens to the density when you go down a metal group? Why?

A

The density INCREASES

The RAM and SIZE of the metals INCREASES.

It increases, because the increase in RAM has an effect on the density too

57
Q

What happens to the first ionisation energy when you go across a period?

A

ionisation energy INCREASES

58
Q

Why does an increase in first ionisation energy happen when you go across a period?

A

This is because as the ATOMIC number increases extra PROTONS are added to the nucleus. So as the nuclear charge INCREASES this pulls the ELECTRON shells TIGHTER to the nucleus.

As the electrons are held MORE tightly to the nucleus it takes MORE energy to remove them. Therefore the first IONISATION energy INCREASES.

59
Q

What happens to the 1st ionisation energy when you go DOWN a group?

A

the first ionisation energy DECREASES

60
Q

Why when you go down a group, the first ionisation energy decreases?

A
  1. The nuclear charge increases so more protons.
  2. Extra energy level, outer electrons further away from nucleus
  3. The extra layers of electrons shield the outer energy levels from the pull of the nucleus.
  4. This means LESS energy is needed to pull off an electron from the outer energy level so the first ionisation energy DECREASES.
61
Q

What is the definition of first ionisation energy?

A

The First Ionisation Energy of an element is the energy required to remove 1 mole of electrons from 1 mole of atoms in the GAS state.

62
Q

What are the units of ionisation energy?

A

kilojoules per mole (kJmol-1)

63
Q

Where can you find the general equation for ionisation energies in the data book?

A

At the top of page 11

64
Q

What is the definition of 2nd ionisation energy?

A

The Second Ionisation Energy is the energy required to remove 1 mole of electrons from 1 mole of positive ions (+1 charge) in the gas state.

65
Q

What is the definition of the 3rd ionisation energy?

A

The THIRD ionisation energy is when a further mole of electrons is removed from ions with a 2+ charge.

66
Q

Explain why the second ionisation energy of an element is always greater than the first ionisation energy.

A
  1. There are less electrons present after the first has been removed
  2. The number of protons hasn’t changed so they are now pulling in a smaller number of electrons so they are pulled in tighter.
  3. As the electrons are now held more tightly is takes more energy to pull another one off hence why the 2nd ionisation energy must be higher.
67
Q

Explain why the second ionisation energy of K is much greater than the second ionisation energy of Mg.

A

The electron arrangement for the elements are K(2,8,8,1) and Mg (2,8,2)

The electron arrangements for the ions are K+ (2,8,8) and Mg+ (2,8,1).

For the potassium the second electron would be removed from a full outer shell.

Remember a full outer shell is a stable electron arrangement so the K+ doesn’t want to lose this 2nd electron so the 2nd ionisation energy is high.

For the Mg+ the second electron is removed from a partially filed shell so it’s a lot easier to remove hence the lower 2nd ionisation energy value.

68
Q

How do you calculate the energy required for Mg(g) → Mg2+(g) + 2e- ?

A

First, you write the first and second ionisation energy equations

Mg(g) → Mg+(g) + e- 738 kJmol-1
Mg+(g) → Mg2+(g) + e- 1451kJmol-1

Then you add them together, the two Mg+ would cancel each other out.

Mg(g) → Mg2+(g) + 2e- 2189kJmol-1

Unit 1 Chemistry (3 decks)

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