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OCR AS/A Level Chemistry A > Module 2.2 Electrons, Bonding and Structure > Flashcards

Module 2.2 Electrons, Bonding and Structure Flashcards

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1
Q

The principle quantum number, n, indicates

A

The shell that the electrons occupy

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2
Q

The larger the value of n,

A

The further the shell is from the nucleus and the higher the energy level

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3
Q

Number of electrons in the first shell

A

2

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4
Q

Number of electrons in the second shell

A

8

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5
Q

Number of electrons in the third shell

A

18

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6
Q

Number of electrons in the fourth shell

A

32

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7
Q

An orbital can

A

Hold a max. of 2 electrons with opposite spins

An orbital is a region of space where electrons are most likely to be found

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8
Q

4 different types of orbital

A

s
p
d
f

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9
Q

S orbitals

A

Spherical shape

Can hold 2 electrons

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10
Q

P orbitals

A

3D dumbbell shape

Can hold 6 electrons

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11
Q

Electron energy levels

A

1s

2s
2p

3s
3p
4s
3d

4p
4d
4f

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12
Q

Filling shells and sub-shells rules

A

The lowest energy level is filled first

Each energy level must be full before the next higher energy level starts to fill

Each orbital is filled singly before pairing starts

The 4s orbital is at a slightly lower energy level than the 3d orbital. Therefore, 4s fills before 3d

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13
Q

Electron configuration of ions

A

Electrons in the highest energy levels are lost first

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14
Q

Ionic bonding

A

Metal + non-metal

Electrons are transferred from the metal atom to the non-metal atom

Oppositely charged ions formed that attract each other

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15
Q

Covalent bonding

A

Two non-metals

Electrons are shared between the atoms and are attracted to the nuclei of both bonded atoms

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16
Q

Metallic bonding

A

Occurs in metals

Electrons are shared between all atoms

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17
Q

Ionic bonds

A

Electrons are transferred from the metal atom to the non-metal atom

Oppositely charged ions form, which are bonded together by electrostatic attraction

The metal ion is positive (cation)

The non-metal ion is negative (anion)

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18
Q

Giant ionic lattices

A

Each ion is surrounded by oppositely charged ions

These ions attract each other from all directions, forming a 3D giant ionic lattice

All ionic compounds exist as giant ionic lattices in the solid state but when aqueous, the ions are free to move so they can conduct when aqueous or molten

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19
Q

Ionic compound properties: high melting and boiling points

A

Ionic compounds are solids at room temperature- a large amount of energy is needed to break the strong electrostatic bonds that hold the oppositely charged ions together in the lattice. Therefore, ionic compounds have high melting and boiling points.

The greater the charge, the stronger the electrostatic forces between the ions so more energy is required to break up the ionic lattice during melting

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20
Q

Ionic compounds: Electrical conductivity

A

In a solid ionic lattice:
The ions can’t move - can’t conduct

When aqueous or molten:
The solid lattice breaks down - the ions are free to move
Can conduct

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21
Q

Why do ionic lattices dissolve in polar substances?

A

The slight charges within the polar substances are able to attract the charged ions in the giant ionic lattice. This means that the lattice is disrupted and the ions are pulled out of it.

22
Q

Covalent bonds

A

The negatively charged shared pair of electrons is attracted to the positive charges of both nuclei

This attraction overcomes the repulsion between the two positively charged nuclei

Electrons are shared

23
Q

Carbon will always make

A

4 covalent bonds

24
Q

Nitrogen will always make

A

3 covalent bonds

25
Q

Oxygen will always make

A

2 covalent bonds

26
Q

Hydrogen will always make

A

1 covalent bond

27
Q

In a dative covalent bond,

A

one atom provides both bonding electrons from a lone pair of electrons

28
Q

For elements in groups 15-17 something weird happens from period 3. What is it?

A

After covalent bonding, one of the bonding atoms may finish up with more than eight outer shell electrons - “expansion of the octet”

29
Q

Simple covalent structures

A

The atoms within each molecule are held together by strong covalent bonds

The different molecules are held together by weak intermolecular forces (London forces)

30
Q

Simple covalent properties: Melting and boiling points

A

Low melting and boiling points:
Weak intermolecular forces
A relatively small amount of energy is needed to break them

31
Q

Simple covalent properties: Electrical conductivity

A

Don’t conduct!

No charged particles free to move

32
Q

Simple covalent properties: Solubility

A

Generally soluble in non-polar solvents:
Weak London forces are able to form between covalent molecules and the solvent
This helps the lattice break down and the substance dissolves

33
Q

Giant covalent properties: Melting and boiling points

A

High melting and boiling points:

High temperatures needed to break the strong covalent bonds within the lattice

34
Q

Giant covalent properties: Electrical conductivity

A

Can’t conduct!

No free charged particles (graphite is an exception)

35
Q

Giant covalent properties: Solubility

A

Insoluble in both polar and non-polar solvents:

The covalent bonds in the lattice are too strong to be broken by either polar or non-polar solvents

36
Q

Linear

A

2 bonded pairs around central atom

180*

37
Q

Trigonal planar

A

3 bonded pairs around central atom

120*

38
Q

Tetrahedral

A

4 bonded pairs around central atom

109.5*

39
Q

Trigonal bipyramid

A

5 bonded pairs around the central atom

90* and 120*

40
Q

Octahedral

A

6 bonded pairs around the central atom

90*

41
Q

Each lone pair reduces the bond angle by

A

2.5*

A lone pair repels more than a bonded pair

42
Q

What is the Pauling scale?

A

A scale used to measure the electronegative of an atom

43
Q

What is electronegativity?

A

The attraction of a bonded atom for the pair of electrons in a covalent bond

44
Q

Bonds between atoms with different electronegativity values result in

A

Polar bonds

45
Q

Relative strength of bond types going from least to most

A

London dispersion forces
Permanent dipole-dipole
Hydrogen bonds
Ionic and covalent bonds

46
Q

The size of London dispersion forces increases with increasing

A

numbers of electrons. The greater the number of electrons, the larger the induced dipoles and the greater the attractive forces between molecules.

47
Q

A hydrogen bond is a strong permanent dipole-permanent dipole attraction between

A

An electron deficient hydrogen atom on one molecule

and

A lone pair of electrons in a highly electronegative atom (N, O, F, Cl) on a different molecule

48
Q

Why is ice less dense than water?

A

Ice has an open lattice structure

Particles are further apart

49
Q

Why does water have a higher melting point and boiling point than expected?

A

Hydrogen bonds are strong so take more energy to break

50
Q

Why does water have a very high surface tension/is more viscous than expected?

A

Because strong hydrogen bonds hold the surface together

OCR AS/A Level Chemistry A (7 decks)

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