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Chemistry > Module 2.2 - Electrons, Bonding and Structure > Flashcards

Module 2.2 - Electrons, Bonding and Structure Flashcards

1
Q

How are quantum numbers used to describe the electrons in atoms?

A
  • Principal quantum number, n, indicates the shell the electron is in
  • Different shells have different principal quantum numbers
  • Larger the value of n, further the shell is from the nucleus + highest energy level
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2
Q

What phrase is ‘shell’ equivalent to?

A

Energy level

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3
Q

How do you work out the number of electrons the first 4 shells hold?

A

2n^2

E.g. 2nd shell = 2 x 2^2 = 8

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4
Q

What is the quantum number of the first shell?

A

1

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5
Q

How many electrons does the first shell hold?

A

2

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6
Q

What is the quantum number of the second shell?

A

2

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7
Q

How many electrons does the second shell hold?

A

8

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8
Q

What is the quantum number of the third shell?

A

3

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9
Q

How many electrons does the third shell hold?

A

18

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10
Q

What is the quantum number of the fourth shell?

A

4

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11
Q

How many electrons does the fourth shell hold?

A

32

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12
Q

How many electrons can an orbital contain?

A

2

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13
Q

How many s-orbitals are there in one shell?

A

1

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14
Q

What is the shape of an s orbital?

A

Spherical

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15
Q

How many p-orbitals are there in a shell, and therefore how many electrons are there in p-orbitals per electron shell (/quantum number)?

A
  • 3; px, py, pz

- 6

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16
Q

What is the shape of the p-orbital?

A

Dumbbell shaped, 8/∞

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17
Q

What shell do d-blocks start in?

A

n=3 (3rd shell)

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18
Q

How many d-orbitals are there in a shell, and therefore how many electrons can it hold?

A
  • 5

- 10

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19
Q

What shell does the f-blocks start in?

A

n=4

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20
Q

How many f-orbitals are there in each shell?

A

7 (therefore 14 electrons)

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21
Q

What method is used to show the electrons in orbitals?

A

‘Electrons in orbitals’

Up and down arrows in boxes

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22
Q

Why do the two electrons in an orbital not repel one another?

A
  • Opposite spins

- Represent the opposite spins by an up and down arrow

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23
Q

What are the 4 sub-shells?

A

s, p, d, f

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24
Q

How do the types of sub-shell change as the electron shell increases?

A

One more type is added
Shell 1: s
Shell 2: s and p etc

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25
Q

How do the energy levels of each type of orbital change?

A

s = lowest
p
d
f

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26
Q

From 1s to 4f, what is the order energy levels of each orbital from lowest to highest?

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d, 4f

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27
Q

What are the rules for the arrangement of electrons in an atom?

A
  • Electrons are added, one at a time, to ‘build up’ the atom
  • Lowest available energy level is filled first (can consider this level as being closest to the nucleus)
  • Each energy level must be filled before the next higher energy level starts to fill
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28
Q

What are the rules of filling up orbitals in the same energy level?

A
  • Each orbital in a sub-shell is filled singly before pairing starts
  • 4s orbital is at a slightly lower energy level than the 3d orbital, so fills before it
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29
Q

In an orbital, how do paired electrons move?

A

With opposite spins

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30
Q

What form is electron configuration written in?

A

nx^y
n=shell number
x=type of orbital
y=number of electrons in orbitals making up the sub-shell

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31
Q

What are the orbitals occupied and the electron configuration of:

a) Boron
b) Carbon
c) Nitrogen
d) Oxygen

A

Boron: 1s2 2s2 2px1 /// 1s2 2s2 2p1
Carbon: 1s2 2s2 2px1 2py1 /// 1s2 2s2 2p2
Nitrogen: 1s2 2s2 2px1 2py1 2pz1 /// 1s2 2s2 2p3
Oxygen: 1s2 2s2 2px2 2py1 2pz1 /// 1s2 2s2 2p4

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32
Q

Which electrons are lost to form positive ions?

A

Electrons in the highest energy level are lost first

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33
Q

What are the most stable and unreactive elements?

A

Noble gases, group 18. Already have a full outer shell of electrons (other elements react to try to get the same electron configuration as a noble gas)

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34
Q

What are the 3 main types of chemical bonding?

A
  • Metallic
  • Ionic
  • Covalent
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35
Q

What kind of materials are involved in an ionic bond?

A

Metal and a non metal (electrons transferred to metal to non metal forming oppositely charged ions)

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36
Q

Using magnesium oxide as an example, show how an ionic bond can be done to give the elements involved the electron configuration of a noble gas.

A
  • Mg forms Mg2+ (1s2 2s2 2p6, same as neon)

- O forms O2- (1s2 2s2 2s6, same as neon)

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37
Q

What materials do covalent bonds form between?

A

2 non metals

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38
Q

Using hydrogen as an example, show how a covalent bond can be used to give elements the same electron configuration as a noble gas.

A

-2 H atoms covalently bond + share electrons, giving 1s2 configuration, the same as helium

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39
Q

What kind of materials do metallic bonds form between?

A

Metals (e.g. Zinc, iron, aluminium, and their alloys such as brass)

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40
Q

In a metallic structure, how many cations are the delocalised electrons shared between?

A

All of them

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41
Q

How does an ionic bond form?

A
  • Electrons are transferred from the metal to the non metal
  • Oppositely charged ions form, bonded together by electrostatic attraction
  • Metal ion is positive (cation)
  • Non metal ion is negative (anion)
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42
Q

How does an ionic bond form between sodium and oxygen, to form sodium oxide?

A
  • 2 sodium atoms (each with one electron in outer shell) give an electron to an oxygen (6 electrons in its outer shell)
  • Forms 2 Na+ ions and an O2- ion
  • 2Na –> 2Na + + 2e- (1s2 2s2 2p6 3s1 –> 1s2 2s2 2p6 [Ne])
  • O + 2e- –> O2- (1s2 2s2 2p4 –> 1s2 2s2 2p6 [Ne])
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43
Q

Describe the structure of a giant ionic lattice.

A
  • Each ion is surrounded by oppositely charged ions
  • Ions attract to each other from all directions, forming a 3D giant ionic lattice
  • All ionic compounds exist as giant ionic lattices in the solid state
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44
Q

Describe the giant ionic lattice of sodium chloride.

A
  • Sodium transfers one electron to chlorine, forming Na+ and Cl-
  • Forms a giant ionic lattice
  • Each Na+ is surrounded by 6 Cl- ions
  • Each Cl- is surrounded by 6 Na+ ions
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45
Q

Describe the ionic bonding of calcium oxide.

A
  • Calcium transfers 2 electrons to oxygen, forming Ca2+ and O2-
  • Calcium gets the same electron configuration of argon, and oxygen the same as neon
  • Ca –> Ca2+ + 2e- (oxidised)
  • O + 2e- –> O2- (reduced)
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46
Q

Describe the ionic bond of aluminium fluoride.

A
  • Aluminium loses 3 electrons forming Al3+ (same electron configuration of neon)
  • 3 fluorine atoms each gain an electron each, forming 3 x F- (same electron configuration of neon)
  • Al –> Al3+ + 3e- (oxidised)
  • 3F + 3e- –> 3F- (reduced)
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47
Q

What are the properties of ionic compounds?

A
  • High melting and boiling points
  • Conduct electricity (aqueous or molten)
  • Soluble
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48
Q

Why do ionic compounds have high melting and boiling points?

A
  • Ionic compounds are solid at room temperature (in a giant ionic lattice)
  • Lots of energy is required to break the strong electrostatic forces of attraction between the oppositely charged ions
49
Q

Explain why sodium chloride’s melting point (801ºC) is lower than that of magnesium oxide (2852ºC).

A
  • Both are ionic compounds
  • Charges on the ions in MgO (Mg2+ and O2-) are greater than the charge of the ions in NaCl (Na+ and Cl-)
  • Greater charge = stronger electrostatic forces of attraction between the ions so more energy is required to break up the ionic lattice
50
Q

Explain the electrical conductivity of a solid ionic lattice.

A
  • Ions are held in a fixed position and no ions can move

- Therefore can not conduct electricity

51
Q

Explain the electrical conductivity of molten or aqueous ionic compounds.

A
  • Solid lattice breaks down and the ions are free to move

- Therefore can conduct elecricity

52
Q

What kind of solvents do ionic compounds dissolve in?

A

Polar solvents

53
Q

Explain how ionic lattices are soluble in polar solvents.

A
  • Polar molecules break down an ionic lattice by surrounding each ion to form a solution
  • Slight charges on the polar substance are able to attract charged ions from the giant ionic lattice
  • Lattice is disrupted and ions are pulled out of it
54
Q

Explain sodium chloride’s solubility in water.

A
  • Water molecules attract the Na+ and Cl- ions
  • Ionic lattice breaks down as it dissolves and water molecules surround the ions
  • Na+ attracts delta- charges on the O atoms of the water molecules
  • Cl- attracts delta+ charges on the H atoms of the water molecules
55
Q

Describe a covalent bond.

A
  • Negatively charged shared pair of electrons are attracted to the positive nuclear charge of both the nuclei
  • This attraction overcomes the repulsion between the 2 positively charged nuclei
  • Resulting attraction is the covalent bond that holds the atoms together
  • The 2 electrons are shared
56
Q

What is the name for a covalent bond with only one shared pair of electrons?

A

Single covalent bond

57
Q

Describe the single covalent bond in hydrogen.

A

-Each hydrogen atom has 1 electron in its outer shell
-Each hydrogen atom contributes 1 electron to the covalent bond
H-H
-Each hydrogen fills its 1s sub-shell, achieving the electron configuration of helium

58
Q

How are single covalent bonds sometimes drawn?

A

By a single line e.g. H-H

59
Q

Explain the covalent bonding in an oxygen molecule.

A

-Share 2 pairs of electrons in a double covalent bond

O=O

60
Q

Explain the covalent bond of nitrogen.

A

-Share 3 pairs of electrons to form a triple bond
_
N=N

61
Q

Explain the covalent bond of carbon dioxide.

A

-2 double bonds, each between the carbon atom and an oxygen atom
O=C=O

62
Q

What is average bond enthalpy measured in?

A

kJ mol^-1

63
Q

What is a dative covalent bond also known as?

A

Coordinate bond

64
Q

What is a dative covalent bond?

A

A covalent bond in which one of the atoms supplies both of the shared electrons to the covalent bond

65
Q

How is a dative covalent bond written?

A

A–>B

Arrow shows direction in which the pairs are being donated (e.g. A is donating a pair to B)

66
Q

Describe the bonding of an ammonium ion.

A
  • Has 3 covalent bonds and one dative covalent bond
  • Formed from ammonia, NH3, and H+
  • One of the electron pairs around the nitrogen in ammonia is a lone pair
67
Q

In an ammonium ion, how can you tell which bond is the dative covalent bond?

A

You can’t, once formed a dative covalent bond is equivalent to all other covalent bonds

68
Q

Explain the bonding in an oxonium ion., H3O+.

A
  • Forms when an acid is added to water. Responsible for reactions of acids (in equations simplified to H+)
  • One of the lone pairs around the oxygen in H2O provides both electrons to form a dative covalent bond with a H+ ion
69
Q

Why can covalent bonds result in atoms that don’t follow the Octet Rule (noble gas electron configuration)?

A
  • May not be enough electrons to reach an octet

- More than 4 electrons may pair up in the bonding (expansion of the octet)

70
Q

Give one example of a covalent bond where there aren’t enough electrons to reach an octet.

A
  • Boron trifluoride (BF3)
  • Boron has 3 electrons in outer shell + each fluorine atom has 7
  • 3 covalent bonds can form
  • Each of boron’s 3 outer electrons is paired so there’s now 6 in its outer shell
  • Each of the 3 fluorine atoms now have 8 outer shell electrons, attaining the octet
  • Central boron atom does not achieve the octet
71
Q

In what elements does expansion of the octet usually occur?

A

Groups 15-17 from period 3
(As you go down the period table, there’s more outer shell electrons able to take part in bonding due to distance)
One of the atoms may have more than 8 electrons in the outer shell, breaking the Octet Rule

72
Q

Describe how expansion of the octet may occur in groups 15, 16 and 17.

A
  • 15: can form 3 or 5 covalent bonds, depending on how many electrons used in bonding
  • 16: can form 2, 4 or 6 covalent bonds, depending on how many electrons used in bonding
  • 17: can form 1, 3, 5 or 7 covalent bonds, depending on how many electrons are used in bonding
73
Q

Describe how expansion of the octet occurs in sulphur hexafluroide, SF6.

A
  • S has 6 electrons in its outer shell
  • 6 covalent bonds can form
  • Each of S’s 6 electrons is paired, so sulphur has 12 outer shell electrons (expanded the octet)
  • Each of the 6 fluorine atoms has eight electrons in its outer shell, attaining the octet
74
Q

What are the 2 types of covalent structure?

A
  • Simple molecular structure

- Giant molecular structure

75
Q

Describe the bonding in a simple molecular covalent structure.

A
  • Atoms within each molecule are held together by strong covalent bonds
  • Different molecules are held together by weak intermolecular forces (London forces)
76
Q

Describe the simple molecular covalent structure of iodine.

A
  • Within each I2 molecule, I atoms are held together by strong covalent bonds
  • When solid I2 (simple molecular lattice) changes state, weak intermolecular forces between the I2 molecules break
77
Q

What are the properties of a simple molecular covalent structure.

A
  • Low melting and boiling points
  • Don’t conduct electricity
  • Soluble in polar substances
78
Q

Why do simple molecular covalent structures have low melting and boiling points?

A

Weak intermolecular forces so only a relatively small amount of energy needed to break them

79
Q

Why can’t simple molecular covalent structures conduct electricity?

A

There are no charged particles that can move

80
Q

Why are simple molecular covalent structures soluble in polar molecules?

A
  • Weak London forces can form between covalent molecules and these solvents
  • This helps the lattice break down and the substance to dissolve
81
Q

Give some examples of a giant covalent structure.

A
  • Diamond
  • Graphite
  • SiO2
82
Q

What are the properties of giant covalent structures?

A
  • High melting and boiling points
  • Don’t conduct electricity
  • Insoluble in polar and non polar solvents
83
Q

Why do giant covalent structures have high melting and boiling points?

A

High temperatures are needed to break the strong covalent bonds between the atoms

84
Q

Why don’t giant covalent structures conduct electricity?

A

There are no free charged particles, except for graphite

85
Q

Why are giant covalent structures insoluble in polar and non polar substances?

A

Covalent bonds in the lattice are too strong to be broken by either polar and non polar solvents

86
Q

Describe the electron repulsion theory.

A
  • All electrons have a negative charge, so electron pairs repel other electron pairs
  • Shape of the molecule is the one that allows the electrons to be as far away from each other as possible
87
Q

What is the bonding region?

A

Where a bond, e.g. double covalent bond, forms

88
Q

Give the number of bonded electron pairs around the central atom of a linear structure, and explain the bond angle using an example.

A
  • 1 or 2
  • H2, H-H (no bond angle)
  • CO2, O=C=O 180º
89
Q

Give the number of bonded electron pairs around the central atom of a trigonal planar molecule, and explain the bond angle using and example.

A
  • 3

- BF3, 120º

90
Q

Give the number of bonded electron pairs around the central atom of a tetrahedral shape, and explain the bond angle using an example.

A
  • 4

- CH4, 109.5º

91
Q

Give the number of bonded electron pairs around the central atom of a trigonal bipyramid molecule, and explain the bond angle using an example.

A
  • 5

- PCl5, 90º and 120º**

92
Q

Give the number of bonded electron pairs around the central atom of an octahedral molecule, and explain the bond angle using an example.

A
  • 6

- SF6, 90º

93
Q

Describe the order of repulsiveness between 2 different types of paired electrons, from most repulsive to least repulsive.

A

Lone pair/lone pair –> lone pair/bonded pair –> bonded pair/bonded pair

94
Q

How much does a lone pair reduce the bond angle by?

A

Around 2.5º

95
Q

Describe the shape and bond angle of a methane molecule.

A
  • CH4
  • Tetrahedral
  • 109.5º
96
Q

Describe the shape and bond angle of an ammonia molecule.

A
  • NH3
  • Pyramidal
  • 107º (one lone pair)
97
Q

Describe the shape and bond angle of water.

A
  • H2O
  • Non-linear
  • 104.5º (2 lone pairs)
98
Q

Describe the shape and bond angle of an ammonia ion.

A
  • NH4+
  • 4 electron pairs around the central N atom
  • Tetrahedral
  • -Charge will be distributed across the whole molecule (non-polar as symmetrical) shown by the square brackets with the + sign
  • 109.5º
99
Q

Where does electronegativity increase towards of the periodic table?

A

Top right (fluorine has the most electronegative atoms)

100
Q

Why is a hydrogen molecule, H2, non-polar?

A
  • The 2 bonding atoms have the same electronegativities
  • Nucleus of each atom is equally attracted to the bonding electron
  • Electrons are evenly distributed between the atoms
101
Q

Why is hydrogen chloride, HCl, a polar molecule?

A
  • Non symmetrical
  • Cl is more electronegative than H
  • Cl has a greater attraction on the bonding pair of electrons than the H atom
  • Bonding atoms are held closer to the Cl atom than to the H atom
  • H has a δ+ charge (small positive charge)
  • Cl has a δ- charge (small positive charge)
102
Q

Why is tetrachloromethane, CCl4, non-polar?

A
  • Symmetrical so the dipoles cancel out

- Each CCl bond is polar

103
Q

Why is carbon dioxide non polar?

A

Although the C=O bonds are polar bonds (as oxygen is more electronegative than carbon), the molecule is symmetrical (due to 180º bond angle) so the dipoles cancel out

104
Q

Why is water a polar molecule?

A

The H-O bonds are polar bonds (as oxygen is more electronegative than hydrogen) and the molecule is not symmetrical therefore the dipoles do not cancel each other out

105
Q

What are the 2 main types of intermolecular forces?

A
  • Hydrogen bonding

- London forces

106
Q

Order all of the types of bonds by their relative strength, from strongest to weakest.

A
  • Ionic and covalent (relative strength 1000)
  • Hydrogen bonds (relative strength 50)
  • Permanent dipole-dipole forces (relative strength 10)
  • London dispersion forces (relative strength 1)
107
Q

What is a permanent dipole-dipole force?

A

Any intermolecular force containing permanent dipoles that aren’t hydrogen bonds

108
Q

What is a permanent dipole-induced dipole interaction and how do they form?

A
  • Molecules with a polar bond (permanent dipoles) bond to induced dipoles from a non polar bond
  • When a polar molecule goes towards a non polar molecule, the δ charge causes the electrons on the non polar molecule to move towards the polar molecule (for δ+) or away from the polar molecule (for δ-)
  • This causes the non-polar molecule to become slightly polar
109
Q

What is a permanent dipole-permanent dipole interaction and how do they form?

A

-Molecules with permanent dipoles are attracted to other molecules with permanent dipoles
-Oppositely charged dipoles of different molecules attract to one another
E.g. HCl: H-Cl——–H-Cl (as Cl is δ- and H is δ+)

110
Q

What are the 2 types of permanent dipole-dipole interactions?

A
  • Permanent dipole-permanent dipole interactions

- Permanent dipole-induced dipole interactions

111
Q

What is the intermolecular force between two non polar molecules?

A

London dispersion forces

112
Q

How do London dispersion forces form?

A
  • Electrons move constantly and randomly in atom’s shells. This movement unbalances the distribution of charge within the electron shells (like the density in the shells ‘wobbling’ from side to side)
  • At any moment, there’ll be an instantaneous dipole across the molecule
  • Instantaneous dipole induces a dipole in a neighbouring molecule, in turn inducing further dipoles on their neighbouring molecules
  • Small induced dipoles attract one another, causing a weak intermolecular force (London dispersion forces or instantaneous dipole-dipole forces)
113
Q

What affects the strength of the London dispersion forces?

A

Number of electrons of the atoms (as creates a larger induced dipole so grater attractive forces between molecules)

114
Q

What is the boiling points of non polar molecules?

A

Low as the weak London forces are all that need to be broken which does not require much energy

115
Q

What atoms does a hydrogen bond form between?

A

Hydrogen and the lone pair of a fluorine, nitrogen or oxygen

116
Q

What anomalous properties of water arise as a result of the hydrogen bond?

A
  • Ice floats on water
  • Higher melting and boiling point than expected
  • High surface tension of water
117
Q

Why does ice float on water?

A
  • When ice forms, water molecules arrange into an orderly pattern and hydrogen bonds form between the molecules (hydrogen bonds occur in a liquid but not as often as molecules move past each other and therefore overcome these bonds)
  • Ice has an open lattice with hydrogen bonds holding water molecules apart
  • When ice melts, the rigid hydrogen bonds collapse, allowing the H2O molecules to move closer together as hydrogen bonds are long
  • So ice is less dense than water, hence why it floats
118
Q

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

A

-The hydrogen bonds are much stronger than other intermolecular forces
-Extra strength of these bonds has to be overcome to melt/boil H2O so has a higher melting/boiling point than if there weren’t any hydrogen bonds
(-Other group 16 elements have the same structure as water but don’t have hydrogen bonds and therefore have a lower melting/boiling point)

119
Q

Why does water have a higher surface tension and viscosity than expected?

A
  • Hydrogen bonds

- Allows insects to walk along water (walking across a raft of hydrogen bonds)

Chemistry (13 decks)

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