Chemistry > 3.1.3 Bonding > Flashcards

3.1.3 Bonding Flashcards

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1
Q

Define covalent bonding

A

When two atoms share pairs of electrons

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2
Q

What is a dative covalent bond?

A

A dative covalent bond forms when the shared pair of electron in the covalent bond come from only one of the bonding atoms.

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3
Q

What is metallic bonding?

A

Metallic bonding is the electrostatic force of attraction between positive metal ions and delocalised electrons

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4
Q

What are the 3 main factors affecting the strength of metallic bonding?

A
  1. Nuclear Charge
  2. Number of Delocalised Electrons per Atom / Charge on Ion
  3. Size of Ion (smaller ions, stronger bond)
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5
Q

What structure do ionic structures take?

A

Giant Ionic Lattice

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6
Q

What are the properties of ionic compounds?

A
  1. High melting point and boiling point because of giant lattice of ions with strong electorstatic forces between oppositely charged ions, requires lots of energy to break.
  2. Poor conductors of electricity when solid / can conduct when molten/aqueous as ions are free to move and carry charge.
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7
Q

Explain 3 key properties of metals

A
  1. High boiling / melting points
    - strong electrostatic forces of attraction between +ive ions and delocalised electrons
  2. Good conductors of electricity
    - delocalised electrons can move through the structure and carry a charge
  3. Malleable / Ductile
    - layers of ions can slide over each other, held together by electrostatic forces
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8
Q

Describe properties of simple molecules

A
  1. Low boiling / melting points
    - due to weak intermolecular forces between molecules e.g van der waals, hydrogen bonds
  2. Poor conductivity as there aren’t any ions and electrons are localised (fixed in place)
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9
Q

Linear
b.p =
l.p =
diagram =
bond angle =
example (2) =

A

b.p = 2
l.p = 0
diagram = (google doc)
bond angle = 180
example = CO2 BeF2

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10
Q

Trigonal Planar
b.p =
l.p =
diagram =
bond angle =
example (2) =

A

b.p = 3
l.p = 0
bond angle = 120
example (2) = BF3 , AlCl3
diagram =

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11
Q

Tetrahedral
l.p =
b.p =
bond angle =
examples =
diagram =

A

l.p = 0
b.p = 4
bond angle = 109.5
examples = SiCl4 CH4
diagram = on google doc

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12
Q

Trigonal Bipyramidal
b.p =
l.p =
b.a =
examples =
diagram =

A

b.p = 5
l.p = 0
b.a = 90 and 120
examples = PCl5
diagram = on google doc

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13
Q

Octahedral
b.p =
l.p =
b.a =
examples =
diagram =

A

b.p = 6
l.p = 0
b.a = 90
examples = SF6
diagram = google doc

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14
Q

Trigonal Pyramidal
b.p =
l.p =
b.a =
examples =
diagram =

A

b.p = 3
l.p = 1
b.a = 107
examples = NCl3, PF3
diagram = google doc

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15
Q

Bent:
b.p =
l.p =
b.a =
examples =
diagram =

A

b.p = 2
l.p = 2
b.a = 104.5
examples = H20, SCl2
diagram = google doc

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16
Q

(Exam Technique) How do you explain the shape of a molecule? (5 step process)

A
  1. State the number of bonding pairs/lone pairs
  2. State that electron pairs repel and try to get as far apart as possible
  3. If there are no lone pairs, state the electron pairs repepl equally
  4. If lone pairs are present, state that the lone pairs repel more than bonding pairs
  5. State actual shape and bond angles
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17
Q

How does the presence of lone pairs affect bond angles?

A

2.5o for each lone pair

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18
Q

What shape is formed from 4 bp and 1 lp?

A

see-saw

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19
Q

What shape is formed from 3 b.p and 2 l.p?

A

T shape

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20
Q

What shape is formed from 3 l.p and 2 b.p?

A

Linear

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21
Q

What shape is made from 4 bond pairs and 2 lone pairs?

A

Square Planar

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22
Q

What is electronegativity?

A

Electronegativity is the power of an atom to attract bonded electrons in a covalent bond towards itself

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23
Q

What are the most electronegative atoms?

A

F
O
N
Cl

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24
Q

What scale is electronegativity measured on?

A

Pauling scale (ranging from 0 to 4)

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25
How does electronegativity change across a period?
Electronegativity will increase across a period as the number of protons increases but there is similar shielding. Furthermore, the atomic radius decreases as the electrons in the same shell are pulled in more.
26
How does electronegativity change down a gorup?
Electronegativity will decrease down a group because the distance between the nucleus and bonded electrons increase and the shielding of the inner shell electrons increases.
27
If electronegativity is similar, what type of bonding could be present?
If both are <2 in electronegativity, bonding is metallic If both are >=2 in electronegativity, bonding is non-polar covalent
28
If electronegativity is different, what type of bonding could be present?
If the difference in electronegativity is >0.5: polar covalent If difference in electronegativity >= 2: Ionic
29
What's important to note about the polarity of symmetric molecule?
It will not be polar even if the individual bonds within the molecule are polar. This individual dipoles on the bonds ''cancel out'' - there is no net dipole moment.
30
What is a dipole moment?
The overall effect of polarity of the bonds in a molecule
31
What are Van Der Waal's Forces?
These are the weakest type of intermolecular force that occurs between all molecular substances and noble gases.
32
Describe how Van Der Waal's Forces form
1) In any molecule, the electrons are moving constantly and randomly. The electron density hence fluctuates and parts of the molecule becomes more or less negative - forming temporary dipoles. 2) These temproary dipoles can induce dipoles in neighbouring molecules, which will be of the opposite sign. 3) The electrostatic forces between the induced dipoles are known as Van Der Waal's forces.
33
What are the factors affecting Van Der Waal's forces?
1) More electrons = stronger VDW forces 2) Bigger surface area = Stronger VDW forces
34
Why does the b.p of halogens down group 7 increase?
Increasing number of electrons in the bigger molecules hence stronger VDW forces form.
35
Why do long chain alkanes have a higher b.p than spherical shaped alkanes?
There is a larger surface area of contact between chained molecules than there is spherical molecules - hence there are stronger VDW forces
36
Describe permanent dipole-dipole forces (2)
1. Electrostatic forces between polar molecules 2. Stronger than VDWs so compounds have higher BP
37
What is hydrogen bonding?
The strongest type of IM force that forms between Hydrogen and F, O, N. The lone pair on these atoms attract a hydrogen atom on another molecule.
38
Use hydrogen bonding to explain why ice floats in water
1) Since ice floats in water, it must be less dense than water 2) The hydrogen bonds in ice hold the molecules further apart so density is lower whereas in water the hydrogen bonds are constantly breaking and reforming since the molecules move.
39
How do we draw hydrogen bonds?
represented by ________________________ must be longer than covalent bonds bond is between the lone pair and hydrogen label all dipoles
40
What are the 3 types of IM forces? List from weakest to strongest.
1. Van Der Waals 2. Dipole-Dipole 3. Hydrogen Bonds
41
Describe the general trend in boiling point. Explain any anomalies. Picture on google docs.
The general trend is that boiling point increases due to an increasing number of electrons leading to stronger VDW forces between molecules. The exceptions are the higher BP of H20, NH3 and HF, due to the hydrogen bonds between the molecules which are stronger than VDW forces.
42
What are the 4 types of crystal structures?
Ionic Metallic Simple Molecular Giant Covalent
43
Describe the properties of simple molecular compounds (4)
1) low m.p/b.p because of weak IM forces (VDW) 2) poor solubility in water 3) poor conducitivity of electricity in solid/when molten as there are no ions / electrons are localised 4) generally mostly gases and liquids
44
Describe properties of macromolecular compounds (5)
1) high mp/bp because of many strong covalent bonds in macromolecular structure, takes a lot of energy t obreak the many strong bond 2) insoluble in water 3) diamond and sand poor, because electrons are localised / graphite good, as free delocalised electrons between layers 4) poor conductivity when molten 5) generally solids
45
State 5 properties of metals
1) high mp/bp due to strong attraction between +ve ions and sea of delocalised electrns 2) insoluble 3) good conductors of electricity 4) shiny 5) malleable, as layers of ions can slide over each other
46
Describe the structure of diamond (2)
1) macromolecular 2) tetrahedral arrangement of carbon atoms
47
Describe the structure of graphite
1) trigonal planar arrangement of carbon in layers 2) 3 covalent bonds per atom in each layer, 4th is delocalised
48
Describe the structure of ice (3)
1) molecular structure 2) tetrahedral arrangament 3) molecules held further apart than in liquid
49
Describe the structure of Iodine (1)
Regular arrangement of molecules held together by weak VDW forces
50
Draw the diagrams for: magnesium, iodine, ice, diamond and graphite
on google doc
51
Solid: Arrangement Movement
1. Tightly packed in a reguar arrangement 2. Vibrate in fixed positions
52
Liquid: Arrangement Movement
1. Tightly packed in a random arrangement 2. Particles move freely and have more energy than in a solid
53
Gases: Arrangement Movement
1. Spaced out and in a random arrangement 2. Particles move freely and have lots of energy The general trend is that boiling point increases due to an increasing number of electrons leading to stronger VDW forces between molecules. The exceptions are the higher BP of H20, NH3 and HF, due to the hydrogen bonds between the molecules which are stronger than VDW forces.
54
Explain why the melting point of aluminium is higher than the metling point of sodium (3)
1. Bigger charge on ions / Smaller Ions / More Protons 2. More Delocalised Electrons 3. Stronger Metallic Bond / Attraction of the Ions
55
(ESQ) Explain why CH4 has a bond angle of 109.5o
1. Around the carbon atom, there are 4 bonding pairs of electrons and no lone pairs 2. Therefore, these repel equally and spread as far apart as possible
56
What are the 2 conditions for hydrogen bonding to occur?
-attraciton between lone pair on F,O,N and H -H connected to F, O, N
57
Why is graphite soft? MS [2]
-The structure is in planes/layers -Weak bonds between planes mean they can slide
58
Explain, in terms of electronegativity, why the boiling point of H2S2 is lower than H2O2: MS [3]
-The electronegativity of S is lower than the electronegativity of O -The difference between H and S electronegativity is less -Hence, S and O have greater delta positive/negative charge, stronger bonds require more energy t obreak -There is no hydrogen bonding between the H2S2 molecules, only van der Waals forces
59
State the meaning of the term electronegativity: MS [2]
The power of an atom or nucleus to withdraw or attract a pair of electrons In a covalent bond
60
Suggest why the electronegativity of the elements increases from lithium to fluorine: MS [2]
More protons mean there is a bigger nuclear charge There is the same or similar shielding
61
Explain, in terms of its structure and bonding, why titanium has a high melting point: MS [2]
There is strong attraction Between the number of protons and delocalised electrons
62
Explain, in terms of structure and bonding, why the melting point of carbon is high: MS [3]
-Macromolecular structure is giant -Covalent bonds in the structure -Are strong and require lots of energy to break/overcom
63
Describe the structure of and bonding in graphite and explain why the melting point of graphite is very high: MS [4]
-Layers of C atoms -Are connected by covalent bonds -van der Waals forces between the layers -Strong covalent bonds are what are broken during melting
64
Explain, in terms of the intermolecular forces present in each compound, why HF has a higher boiling point than HCl: MS [3]
HF has hydrogen bonding HCl has permanent dipole-dipole bonding Hydrogen bonding is stronger
65
Describe the bonding in a metal: MS [3]
Positive ions Attract delocalised electrons Which are in a free to move ‘sea’
66
Why do diamond and graphite both have high melting points? MS [3]
Macromolecular structures Covalent bonds between atoms These are strong bonds and it require lots of energy to break them
67
Why is graphite a good conductor of electricity? MS [1]
Delocalised electrons can carry charge
68
Explain why the melting point of magnesium is higher than that of sodium: [3]
Mg2+ have a higher charge than Na+ Shorter distance between e- and ions in Mg2+ Hence stronger metallic bonding
69
Explain why the second ionisation energy of magnesium is greater than the first ionisation energy of magnesium: MS [1]
The electron is being removed from a more positive ion with a greater nuclear charge
70
Explain why nickel is ductile (can be stretched into wires): MS [1]
Layers of atoms are able to slide over each other
Chemistry
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