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Chemistry > Chapter 7 > Flashcards

Chapter 7 Flashcards

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1
Q

Across period 2 what sub shells are occupied?

A

2s and 2p

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2
Q

Across period 3 what sub shells are occupied?

A

3s, 3p

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3
Q

Across period 4 what sub shells are occupied?

A

Although 3d is involved in the highest subshell, only 4s and 4p are occupied.

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4
Q

What is the trend down the group?

A

Elements in a group have the same number of electrons in each shell. This similarity in configuration gives elements in the same group similar properties.

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5
Q

Why are blocks used?

A

The elements are broken into blocks corresponding to the highest sub shell.

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6
Q

How is the atomic number of elements within a periodic table arranged?

A

From left to right elements are arranged in increasing atomic number (proton number).

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7
Q

How is the groups of elements within a periodic table arranged?

A

Elements are arranged in vertical columns called groups. Each element in that group has the same number of outer shell electrons with similar properties.

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8
Q

How is the periods of elements within a periodic table arranged?

A

They are arranged in horizontal rows, the period number gives the highest energy electron shell in an elements atoms.

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9
Q

What is periodicity?

A

The repeating trend in properties of elements across a period.

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10
Q

What does ionisation energy measure?

A

How easily an atom loses electrons to form positive ions.

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11
Q

What is meant by the first ionisation energy?

A

The energy required to remove one electron from each atom in one mole of gaseous atoms of an element forming one mole of gaseous 1+ ions. Units is KJ per mole.

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12
Q

Where is the first electron lost?

A

It is lost from the highest energy level as it has the least amount of attraction from the nucleus.

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13
Q

What are the factors affecting ionisation energy?

A

Atomic radius, nuclear charge and electron shielding.

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14
Q

How does the atomic radius affect the ionisation energy?

A

The greater the distance between the outer shell and the nucleus the less nuclear attraction. The force of attraction falls off sharply with increasing distance from the nucleus to the electrons.

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15
Q

How does nuclear charge affect the ionisation energy?

A

The greater the amount of protons in the nucleus the greater the attraction between the nucleus and the outer shell attractions

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16
Q

How does electron shielding affect the ionisation energy?

A

Electrons are negatively charged so the inner electrons repel the outer shell electrons this repulsion is called the shielding effect which reduces the attraction.

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17
Q

Describe the trend in the first ionisation energy?

A

A general increase in first ionisation energy across each period. A sharp decrease in first ionisation energy between the end of one period and the start of the next.

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18
Q

How many ionisation energies can an element have?

A

It can have as many ionisation energies as there are electrons.

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19
Q

What is the second ionisation energy?

A

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

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20
Q

Describe the second ionisation energy of helium?

A

It is greater as the two protons attracted to electrons in the 1s shell. After losing the first electron there is a greater nuclear attraction so more ionisation energy is needed to remove the second electron.

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21
Q

What does a large increase between certain ionisation energies suggest?

A

That an electron has to be removed from a different shell which is closer to the nucleus with less shielding.

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22
Q

Describe the trends in first ionisation energies down a group?

A

The atomic radius increases, more inner shells so shielding increases, nuclear attraction decreases, first ionisation energy decreases.

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23
Q

Describe the trends in first ionisation energies across a period?

A

Nuclear charge increases, same shell so have similar shielding, nuclear attraction increases with atomic radius decreases. The first ionisation energy increases.

24
Q

How do you work out the second ionisation energy?

A

You add the first ionisation energy and the second together.

25
Q

Why is there a drop in ionisation energies?

A

It is due to subshells, their energies and how orbital is filled with electrons.

26
Q

State the trend within the first ionisation energy?

A

A rise from lithium to beryllium. A fall to boron followed by a rise to carbon and nitrogen. A fall to oxygen followed by a rise to fluorine and neon.

27
Q

Compare beryllium to boron in terms of ionisation energies?

A

The fall in ionisation energy marks the filling of the 2p subshell. The 2p sub shell in boron has a higher energy than the 2s sub shell in beryllium. Therefore the 2p electron in boron is easier to remove than one in the 2s shell in beryllium. The first ionisation is lower in boron than in beryllium

28
Q

Compare nitrogen and oxygen in terms of ionisation energies?

A

In nitrogen and oxygen the highest energy levels for electron is in the 2p sub shell. In oxygen the paired electrons in one of the 2p orbital repel one another making it easier to remove an electron from an oxygen atom than a nitrogen atom. Therefore the first organisation energy of oxygen is less than the first ionisation of nitrogen.

29
Q

Is it easier to remove a spin paired electron or a half filled shell of electrons (one electron pair orbital)?

A

It is easier to remove a spin paired electron as it is less stable due to the fact that the electrons repel one another.

30
Q

What are semimetals/metalloids?

A

They are elements near the metal and non-metal divide and can show in between properties. The two metals in between the divide a silicon and germanium which are semimetals.

31
Q

At room temperature what state are metals at?

A

They are in solid state apart from mercury.

32
Q

What is metallic bonding?

A

It is the strong electrostatic attraction between metal cations and delocalised electrons.

33
Q

Describe the structure in metallic bonding?

A

Cations are fixed in position maintaining the shape and structure of the metal. The delocalised electrons are mobile and free to move across the whole structure meaning only the electrons can move. In a solid metal structure each atom has donated its negative outer shell electrons to a shared pool of delocalised electrons throughout the structure. The cations left behind consist of the nucleus and the inner electron shells of the metal ions.

34
Q

What is the structure described as of metals ?

A

A giant metallic lattice

35
Q

What are the properties of metals?

A

Strong metallic bonds – attraction between positive ions and delocalised electrons. High electrical conductivity and high melting and boiling point.

36
Q

Describe the periodic trend of melting points?

A

The melting point increases from group 1 to 14 (4). There is a sharp decrease in melting point from group 14 (4) and group 15 (5). The melting points are comparatively low from group 15 (5) to group 18 (0).

37
Q

What does the sharp decrease in melting point indicate?

A

A change from giant to simple molecular structures

38
Q

How are melting and boiling points different from giant structures to simple molecular structures?

A

Giant structures have strong forces so have high melting points to overcome them. For a simple molecular structures have low melting points as they have weak forces of attraction.

39
Q

Describe the electrical conductivity of metals?

A

Metals can conduct electricity in both solid and liquid state. When a voltage runs through the delocalised electrons carry the charge. This differs for ionic compounds as they can only conduct electricity in molten or aqueous states when the ions are free to move when solid they have no mobile charge carriers.

40
Q

Describe the melting and boiling points of metallic structures?

A

The melting point depends on the strength of the metallic bonds holding the atoms together in a giant metallic lattice. Most metals require large amounts of energy is to overcome the strong electrostatic attraction between the positive ions and the electrons resulting in a high melting and boiling point.

41
Q

Describe the solubility within metallic structures?

A

Metals are insoluble and don’t dissolve. Any interactions with lead to a reaction rather than dissolving.

42
Q

What are simple molecular lattice structures?

A

It’s when non metals are held together by weak intermolecular forces with strong covalent bonds between the atoms within the molecules.

43
Q

What three elements form giant covalent lattices?

A

Boron, carbon and silicon.

44
Q

What is a giant covalent lattice?

A

It’s when atoms are held together by network of strong covalent bonds to form a covalent lattice.

45
Q

Describe the structure of carbon and silicon?

A

They have four electrons in their outer shells. Carbon (in diamond form) and silicon use these for electrons to form covalent bonds with other atoms. Which result in a tetrahedral structure which bond angles of 109.5° by electron pair repulsion.

46
Q

What is graphite?

A

It’s when carbon uses three of its four outer shell electrons to form covalent bonds to other atoms with the remaining electron being released into a pool of delocalised electrons.

47
Q

What is the structure of graphite?

A

Graphic forms planar hexagonal layers with bond angles of 120° by electron repulsion theory. These layers are parallel which are held by weak London forces, despair electron is delocalised between the layer so electricity can be conducted. Between the layers that are strong covalent bonds.

48
Q

What are the physical properties of graphite?

A 
High melting and boiling point
Insoluble
Conducts electricity (delocalised electrons free to move across the whole structure)
49
Q

What are the properties of giant covalent structures dominated by?

A

The strong covalent bonds making them difficult to break down.

50
Q

Explain the melting and boiling point of covalent structures?

A

They have a high melting and boiling point as the bonds found in the covalent lattice are very strong. High temperatures are required to provide large quantities of energy needed to break these bonds.

51
Q

Explain the solubility of giant covalent structures?

A

They are insoluble in most solvents, the covalent bonds are far too strong to be broken down by the interaction with solvents.

52
Q

Explain the electrical conductivity of giant covalent structures?

A

The giant covalent lattice doesn’t conduct electricity apart from graphene in graphite. In carbon silicon all for an outer electrons are involved in covalent bonding so unable for conducting electricity where is graphene and graphite can.

53
Q

What is graphene?

A

It is a single layer of graphite composed of hexagonally composed of carbon atoms attached by strong covalent bonds.

54
Q

Describe the structure of graphene?

A

It has a giant covalent structure of carbon based on hexagonal layers with bond angles of 120° by electron pair repulsion. The spare electron is released as a delocalise electron so is a good conductor of electricity.

55
Q

What are the two allotropes of carbon?

A

Graphite and graphene

Chemistry (23 decks)

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