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CHEMISTRY A LEVEL > CHAPTER 7 - PERIODICITY > Flashcards

CHAPTER 7 - PERIODICITY Flashcards

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1
Q

How did Mendeleev arrange the periodic table

A

Order of atomic mass

Lined in groups with similar properties

Left gaps, assuming atomic mass was wrong or more elements were yet to be discovered

Predicted properties of the missing elements from group trends

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2
Q

What is a group of the periodic table (not an example)

A

Vertical columns of elements

Same number of outer shell electrons and similar properties

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3
Q

What are Periods in the periodic table

A

Horizontal rows

The number of the period gives the highest energy electron shell in the elements atom

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4
Q

What are blocks in the periodic table

A

Four distinct areas corresponding to their highest sub-shell

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5
Q

Mendeleev left gaps in his period table that were later filled when, eg. Scandium gallium and germanium were discovered. In addition, there is an entire group in the modern periodic table that Mendeleev omitted. State which group was missing and suggest why Mendeleev was unaware of it

A

Group 18 (0)

Contains unreactive noble gases, no elements from group 18 had been discovered at the time

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6
Q

What is ionisation energy

A

Measures how easily an atom loses electrons to form positive ions

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7
Q

What is first ionisation energy

A

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

eg. Na(g) –> Na+ + e-

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8
Q

What are the factors affecting ionisation energy

A

Atomic radius, Nuclear Charge, Electron Shielding

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9
Q

How does Atomic radius affect ionisation energy

A

The greater the distance between the nucleus and the outer electrons, the less the nuclear attraction

The force of attraction falls off sharply, so atomic radius has large effect (most important factor)

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10
Q

How does Nuclear charge affect Ionisation energy

A

The more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons

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11
Q

How does electron shielding affect ionisation energy

A

Electrons are negatively charged and so inner-shell electrons repel outer-shell electrons.

This repulsion is called the shielding effect, which reduces the attraction between nucleus and outer shell electrons

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12
Q

Why do second ionisation energies (or successive) require more energy to separate electrons

A

After an electron is lost, the remaining electrons are pulled closer to the nucleus

Atomic radius decreases
Nuclear ATTRACTION (not charge) increases per electron

(pg 97)

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13
Q

What does a large increase in ionisation energy suggest

A

The electron must be removed from a different shell, closer to the nucleus with no/less shielding

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14
Q

What are the two key patterns of first ionisation energies

A

General increase across each period

Sharp decrease between end of one period and the start of the next

(pg 98)

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15
Q

Why does ionisation energy decrease down a group

A

Atomic radius increases

Inner shell shielding increases

Nuclear attraction to outer electrons decreases

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16
Q

Why does ionisation energy increase across a period

A

Nuclear charge increases
Similar shielding: same shell
Nuclear attraction increases

Atomic radius decreases

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17
Q

What is the Trends in the first ionisation energy cross period 2

A

3 rises, 2 falls

Rise from Lithium - Beryllium
Fall to Boron
Rise from Carbon to Nitrogen
Fall to oxygen
Rise from Fluorine to Neon

(Pg 99)

18
Q

Why is there a decrease in first ionisation energy from beryllium to Boron

A

Filling of the 2p subshell in Boron, easier to remove than a 2s Subshell in Beryllium

(pg 100)

19
Q

Why is there a decrease in first ionisation energy from Nitrogen to Oxygen

A

Pairing of electrons in 2p sub-shell in Oxygen

Electron pair-repulsion repels other other electrons, making it easier to remove than 2p in Nitrogen

(pg 100)

20
Q

State the equations to represent the first two ionisation energies in sulfur

A

S (g) —> S+ (g) + e-
S+ (g) —> S2+ (g) + e-

21
Q

What is second ionisation energy

A

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

22
Q

Explain why successive ionisation energies increase

A

As each electron is removed, the outer shell is drawn closer to the nucleus

Nuclear attraction is greater and more energy is needed to remove the next electron

23
Q

Explain the general trend in first ionisation energy from sodium to argon

A

Nuclear charge increases from sodium to argon and outer shell.

Electrons are in the same shell with similar shielding

atomic radius decreases

results in an increase in nuclear attraction on the outer electrons and an increase in first ionisation energy.

24
Q

Explain the sharp drop in first ionisation energy between neon and sodium

A

Sharp drop reflects the addition of a new shell with a resulting increase in distance and shielding.

This decreases the nuclear attraction on the outer electrons,
decreasing the first ionisation energy.

25
Q

Explain the trend in first ionisation energy is shown by helium, neon and Argon

A

Ionisation energy decreases from helium to neon to argon due to an increase in the number of Shells, so increasing atomic radius and shielding.

This causes a decrease in nuclear attraction on outer electrons

decreasing the first ionisation energy.

26
Q

Explain why aluminium has a lower first ionisation energy than magnesium

A

The 3p sub shell in aluminium has a higher energy level than the 3s sub shell in magnesium.

The 3p electron is easier to remove.

27
Q

Explain why sulphur has a lower first ionisation energy than phosphorus

A

Phosphorus has three electrons in three piece of shell one electron for each orbital.

So far has four electrons in the 3p sub shell, two electrons paired in one orbital and one electron pears in each of the other two.

The pair electrons in sulphur repel one another making it easier to remove one of these electrons than unpaired electron.

28
Q

What is Metallic bonding

A

Bonding between 2 metals

Each atom has donated all of its outer shell of negative electrons to a shared sea of delocalised electrons

Positive Cations left behind consists of the nucleus and the inner electron shells the metal atoms

29
Q

What forms as a result of metallic bonding and what allows it to form that

A

Giant Metallic lattice

Cations are fixed in position, maintaining structure and shape of the metal

Delocalised electrons are mobile and able to move throughout the structure

30
Q

What are some properties of metals

A

Strong bonds - attraction between positive ions and delocalised electrons

High electrical conductivity - Electrons can carry the charge

High MP and BP

Insoluble - leads to a reaction rather than dissolving eg. Sodium and water

31
Q

What do non-metallic elements exist as?

A

Simple covalently bonded-molecules, held together by weak intermolecular forces

32
Q

What do non-metals like Boron and carbon form

A

Giant covalent structure

Billions of atoms held together by a network of strong covalent bonds instead of small molecules

33
Q

Describe the structure of Carbon in diamond form

A

Giant covalent lattice

4 electrons in outer shell form 4 covalent bonds

all bonds 109.5 degrees due to electron pair repulsion

Tetrahedral structure around each carbon

34
Q

What are properties of giant covalent lattice structures eg. diamond

A

High MP + BP due to bond strength

Insoluble - bonds are too strong to be broken by interaction with solvents

Cannot conduct electricity due to all four outer-shell electrons involved in covalent bonding, none available to conduct electricity

(exceptions Graphene and Graphite)

35
Q

Why Can Graphene and Graphite

A

Form hexagonal structures

Using 3 out of 4 electrons

1 available to conduct electricity

Trigonal Planar - 120 degrees

36
Q

What is graphene

A

A single layer of graphite - composed of hexagonally structured carbon atoms

(pg 104)

37
Q

How is Graphite arranged

A

Stacks of parallel hexagonally arranged carbon atoms, held together by weak London forces

38
Q

Why is there are sharp decrease between group 14 (4) and 15 (5)

A

change from giant covalent to simple molecular structures

(pg 105)

39
Q

Explain what is meant by metallic bonding and why this type of bonding enables metals to conduct electricity

A

Strong electrostatic forces of attraction between cations and delocalised electrons

The delocalised electrons can move across a potential difference

40
Q

Explain how the bonding in a simple molecular lattice differs from that of a simple covalent lattice

A

Simple molecular lattice has London forces between molecules

Giant covalent lattice has covalent bonds between atoms

41
Q

Across period 4, the trend in properties is not exactly the same as across Periods 2 and 3. Suggest explanations for the following:

1) Germanium is a good conductor of electricity

2) Arsenic has a much higher melting point than nitrogen and phosphorus

A

1) the group 4 element, Germanium has a giant metallic lattice rather than a giant covalent lattice in carbon and silicon

2) The group 5 element, Arsenic has a giant covalent lattice whereas in period 2 and 3, the group 5 element has a simple molecular lattice

CHEMISTRY A LEVEL (14 decks)

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