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Chemistry OCR(A) A Level > M3 Chapter 7 - Periodicity > Flashcards

M3 Chapter 7 - Periodicity Flashcards

1
Q

How is the periodic table arranged?

A

By chemical and physical properties and Atomic Number.

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2
Q

What does atomic number have to do in the periodic table?

A

From left to right the elements are arranged in order of increasing atomic number. Each successive element has atoms with one extra proton.

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3
Q

What is a group?

A

Elements are arranged into vertical columns called groups. Each element in a group has atoms with the same number of outer-shell electrons and similar properties.

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4
Q

What is a period?

A

The elements are arranged into horizontal rows called periods. The number of the period gives the number of the highest energy electron shell in an element’s atoms.

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5
Q

What is periodicity?

A

The repeating trend in properties across the periodic table.

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6
Q

What topics/properties influence periodicity?

A

Electron configuration, ionisation energy, structure and melting point.

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7
Q

What is the trend in electron configuration across a period?

A

Each period starts with an electron in a new highest energy shell: Across period 2 the 2s subshell fills with 2 electrons follwed by the 2p subshell with 6 electrons. Across period 3 it is the same as period 2.

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8
Q

What is the trend in electron configuration down a group?

A

Elements in each group have atoms with the same number of electrons in each subshell. This similarity in electron configuration gives elements their similar chemistry.

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9
Q

What are the four blocks of the periodic table?

A

S, P, D, F

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10
Q

What is ionisation energy?

A

Ionisation energy measures how easily an atom loses electrons to form positive ions.

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11
Q

What is the definition of 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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12
Q

Give an example of a first ionisation energy equation…

A

Na(g) –> Na+(g) + e-

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13
Q

What factors affect first ionisation energy?

A

Atomic Radius, Nuclear Charge, Electron Shielding

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14
Q

What is the effect of Atomic Radius on the ionisation energy?

A

The greater the distance between the nucleus and the outer electrons the less the nuclear attraction. Hence ionisation energy is lower if there is an increase on atomic radius.

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15
Q

What is the effect of Nuclear Charge on the ionisation energy?

A

The more protons there are in a nucleus of an atom, the greater the attraction between the nucleus and the outer electrons. Hence ionisation energy is higher if there is an increase of nuclear charge.

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16
Q

What is the effect of Electron Shielding on ionisation energy?

A

Electrons are negatively charged hence, inner shell electrons repel outer shell electrons. This repulsion is called the shielding effect. Hence an increase of electron shielding, the lower the attraction between the nucleus and the outer electrons hence ionisation energy is lower.

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17
Q

Give an example of a successive ionisation energy equation.

A

He(g) –> He+(g) + e- First Ionisation Energy
He+(g) –> He2+(g) + e- Second Ionisation Energy

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18
Q

What is the definition of 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.

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19
Q

What do successive ionisation energies show us for different energy levels in an atom?

A

In the successive ionisation energies of Fluorine, the large increase between the seventh and eighth electron shows that the eighth electron must be removed from a different shell, closer to the nucleus and with less shielding.

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20
Q

What predictions can we make from successive ionisation energies?

A
  1. The number of electrons in the outer shell
  2. The group of the element in the periodic table
  3. The identity of an element.
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21
Q

What is the trend in ionisation energy down a group?

A

The first ionisation energies decrease down a group.W

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22
Q

Why does ionisation energy decrease down the group?

A

The atomic radius increases.
There are more inner shells therefore electron shielding increases.
The nuclear attraction on outer electrons decreases.
Therefore, the overall first ionisation energy decreases because less energy is required to remove the electron.

23
Q

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

A

There is a general increase in first ionisation energy across a period.

24
Q

Why is there a general increase across the period in first ionisation energy?

A

Nuclear charge increases.
Same shell therefore shielding remains the same.
The nuclear attraction increases due to the nuclear charge increasing.
The atomic radius decreases too due to this.
The first ionisation energy increases.

25
What are the subshell trends in first ionisation energy?
Across a period, it was a 'general increase'. This is because there are periodic drops occuring at the same position in each period. This is caused by subshells, their energies and how orbitals fill with electrons.
26
Why is there a fall in first ionisation energy from Beryllium to Boron even though they are across a period from eachother?
This is due to Boron marking the start of filling the 2p subshell.
27
How does the new subshell being filled decrease ionisation energy?
The 2p subshell in this case has a higher energy than the 2s subshell. Therefore, in Boron the 2p electron is more easily removed than one of the 2s electrons in Beryllium. Therefore, the first ionisation energy of Boron is less than Beryllium.
28
What about Carbon vs Nitrogen? What would the first ionisation energies look like for them?
The first ionisation energy of Oxygen is less than Nitrogen.
29
What factor makes the first ionisation energy of Oxygen to be less than Nitrogen?
Electron Pairing in the 2p subshell.
30
Explain the concept of electron pairing as to why the first ionisation energy of Oxygen less than Nitrogen?
The highest energy electrons are in the 2p subshell. In oxygen, the paired electrons in one of the 2p orbitals repel one another. This makes it easier to remove an electron from an Oxygen atom than a Nitrogen atom. Hence, the first ionisation energy of Oxygen is lower than Nitrogen.
31
Give three types of bonding.
Ionic, Covalent, Metallic.
32
What is metallic bonding?
A special type of bonding specifically for metals only.
33
How does metallic bonding work?
In a solid metallic structure, each atom donates its negative outer shell electrons to a pool of electrons, which are delocalised (spread out) the entire structure & The positive ions, (cations) left behind consist of the nucleus and the inner electron shells of the metal atoms.
34
Definition of Metallic Bonding.
Metallic Bonding is the strong electrostatic force of attraction between cations and delocalised electrons.
35
Give the structure of a metallically bonded metal structure.
The cations are in a fixed shape maintaining the shape and structure of the metal. The delocalised electrons are mobile and are able to move throughout the structure, only the electrons move.
36
What shape does metallic bonding cause?
A giant metallic lattice.
37
Give the properties of metals. (3)
1. Strong metallic bonds - attraction between the cations and the delocalised electrons. 2. High electrical conductivity. 3. High melting and boiling points.
38
How can metals be electrically conductive?
Metals can conduct in solid and liquid states. This is due to the delocalised electrons being able to carry the charge through the structure due to their ability to move.
39
Why do metals have high melting and boiling points?
The strong electrostatic force of attraction between the cations and the delocalised electrons.
40
Are metals soluble?
They do not dissolve in water. Instead they could react, when there is a polar solvent and the charges in the metallic lattice combine. Eg: Sodium and Water.
41
How do non-metal element exist?
As simple covalent structures held together by weak intermolecular forces of attraction. Hence, low melting and boiling point.
42
Which non-metals exist as big boy structures? (3)
Boron, Carbon, Silicon.
43
How are these three held together?
They form a network of strong covalent bonds to form a giant covalent lattice.
44
How do Carbon and Silicon achieve such a feat?
They are both Group 14 (4) meaning their atoms have 4 electrons on the outer shell. Carbon and Silicon use these electrons to form covalent bonds to other Carbon and Silicon atoms.
45
What structural arrangement does this cause for Carbon and Silicon to have when bonded?
Tetrrahedral Structure
46
What are the bond angles of a tetrahedral structure?
109.5 due to electron pair repulsion.
47
What properties are within Giant Covalent Structures? (3)
1. High melting and boiling points. 2. Insoluble 3. Non-conductive.
48
Why are giant covalent structures with high melting and boiling points?
Due to the strong covalent bonds meaning high temperatures are required to break them.
49
Why are giant covalent structures insoluble in almost all solvents?
The strong covalent bonds are far too strong to be broken via an interaction with solvents.
50
Why are giant covalent structures unable to conduct electricity?
This is due to all four outer shell electrons being involved with covalent bonding hence none are able to carry the charge for electricity.
51
What are the exceptions to the conductivity?
Only graphene and graphite (forms of carbon) are able to conduct electricity since Carbon can form several structures in which one of the electrons is available to carry the charge.
52
What is the periodic trend in melting points across group 2?
Li, Be, B and C are all giant structures (Li and Be are giant metallic, B and C are giant covalent). Therefore, there is an overall increase across the period with all of them having relatively high melting points. However, after; N, O, F, Ne all have extremely low melting points as they are simple covalent structures meaning they have weak intermolecular forces of attraction.
53
What is the trend in melting points across group 3?
Similar to group 2, the elements of Na, Mg, Al and Si all are giant structures (Na, Mg and Al are giant metallic and Si is giant covalent). Hence, across the first four elements of group 3 there is an increase with all of the structures having a strong melting point to begin with. However, afterwards; P, S, Cl and Ar all have extremely low melting points due to being simple molecular (simple covalent) meaning they have weak intermolecular forces of attraction.

Chemistry OCR(A) A Level (8 decks)

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