3.1.3 Bonding Flashcards
1
Q
Ionic Bonding [definition]:
A
The electrostatic attraction between oppositely charged ions in a lattice
2
Q
Covalent Bonding [definition]:
A
A shared pair of e⁻s with opposite spins
- one e⁻ donated by each atom
3
Q
Sulphate formula:
A
SO₄²-
4
Q
Hydroxide formula:
A
OH−
5
Q
Nitrate formula:
A
NO₃⁻
6
Q
Giant Ionic lattice [2]:
A
- Sodium Chloride
- Magnesium oxide
7
Q
Covalent simple molecular structures [5]:
A
- Iodine
- Ice
- Carbon dioxide
- Water
- Methane
8
Q
Giant Covalent structures [4]:
A
- Diamond
- Graphite
- Silicon dioxide
- Silicon
9
Q
Ionic boiling and melting points [2]:
A
- high- because of giant lattice of ions with strong
electrostatic forces between oppositely charged ions - requires lot of energy to break
10
Q
Ionic Solubility in water:
A
Generally good
11
Q
Ionic compound conductivity (when solid):
[2]
A
- poor
- ions can’t move/ fixed in lattice
12
Q
Ionic compound conductivity (when molten):
[2]
A
- Good
- Ions can move freely
13
Q
Co-ordinate/ Dative Covalent bond [definition]:
A
A bond that contains a shared pair of electrons with both electrons supplied by one atom.
14
Q
Simple Covalent compound boiling and melting points:
A
low- because of weak intermolecular forces between molecules (specify type e.g van der
waals/hydrogen bond)
15
Q
Simple covalent compound solubility in water:
A
Generally poor
16
Q
Simple covalent compound conductivity (when solid):
[2]
A
- poor
- no ions to conduct and electrons are
localised (fixed in place)
17
Q
Simple covalent compound conductivity (when molten):
[2]:
A
- Poor
- No ions
18
Q
Giant Covalent compound boiling and melting points:
A
High- because of many strong covalent
bonds in macromolecular structure (requires a lot of energy to break many strong bonds)
19
Q
Giant covalent compound solubility in water:
A
Insoluble
20
Q
Giant covalent compound conductivity (when solid):
[2]:
A
- diamond and sand = poor because electrons are localised and can’t move
- graphite = good, has free delocalised electrons between layers to conduct
21
Q
Giant covalent compound conductivity (when molten):
A
Poor
22
Q
Metallic bonding boiling and melting points:
A
high- strong electrostatic forces between positive ions and sea of delocalised electrons
23
Q
Metallic bonding solubility in water:
A
Insoluble
24
Q
Metallic bonding conductivity (when solid):
[2]:
A
- good
- delocalised electrons can move through structure
25
Metallic bonding conductivity (when molten):
Good
26
Linear Shape [3]:
- BP:2
- LP: 0
- Bond angle: 180
27
Trigonal Planar Shape [3]:
- BP: 3
- LP: 0
- Bond angle: 120
28
Tetrahedral Shape [3]:
- BP: 4
- LP: 0
- Bond angle: 109.5
29
Bent Shape [3]:
- BP: 2
- LP: 2
- Bond angle: 104.5
30
Trigonal Pyramidal Shape [3]:
- BP: 3
- LP: 1
- Bond angle: 107
31
Trigonal Bipyramidal Shape [3]:
- BP: 5
- LP: 0
- Bond angle: 120 and 90
32
Octahedral Shape [3]:
- BP: 6
- LP: 0
- Bond angle: 90
33
Electronegativity [definition]:
Electronegativity is the relative tendency of an atom in a covalent bond to attract electrons in a covalent bond to itself
34
What are the most electronegative atoms? [3]:
- Fluorine
- Oxygen
- Nitrogen
[F, O, N]
35
Factors affecting electronegativity [2]:
- increases across a period as the number of protons increases and the atomic radius decreases cus electrons in the same shell are pulled in more
- decreases down a group cus the distance between nucleus and outer electrons increases + shielding of inner shell electrons increases
36
Purely covalent compound [definition]:
A compound containing elements of similar electronegativity and hence a small electronegativity difference will be purely covalent
37
Ionic compound [definition]:
A compound containing elements of very different electronegativity and hence a
very large electronegativity difference (> 1.7) will be ionic
38
When does a polar covalent bond form?
when the elements in the bond have different
| electronegativities. (Of around 0.3 to 1.7)
39
What is a polar covalent bond?
When a bond is a polar covalent bond it has an unequal distribution of electrons in the bond and produces a charge separation, (dipole) δ+ δ- ends.
40
Symmetrical molecule [3]:
- All bonds identical
- No lone pairs
- Will not be polar even if individual bonds within the molecule are polar
(The individual dipoles on the bonds ‘cancel out’ due to the symmetrical shape of the molecule.
There is no NET dipole moment: the molecule is
NON POLAR)
41
Where do Van der Waal forces occur? [2]:
- Occur between all molecular substances and noble gases
| - They do not occur in ionic substances
42
The main factor affecting the size of Van der Waals [2]:
- The more electrons there are in the molecule the higher the chance that temporary dipoles will form
- This makes the Van der Waals stronger between the molecules and so boiling points will be greater
43
What is the weakes IMF?
Induced dipole/ Van der Waal's
44
When do ions form?
When e⁻s are transferred between elements that have a large diff in negativity (metal + non-metal)
45
How are ionic lattices structured?
Regular lattice with alternating +ive & -ve ions
46
What holds an ionic lattice together?
Very strong electrostatic forces
47
What are the properties of an ionic crystal structure? [5]:
- V high melting points
- Generally soluble in water
- Electrical insulators when solid
- Electrical conductors when molten
- Brittle
48
Why are ionic compounds generally soluble in water? [3]:
- Water is v polar
- It disrupts electrostatic attractions between ions & breaks them up
- causes the compound to dissolve
49
Why can ionic compounds NOT conduct electricity when solid? [2]:
- Ions are in a fixed position
| - So canny carry a charge
50
Why can ionic compounds conduct electricity when molten? [2]:
- Ions are free to move
| - so can carry a charge
51
Why are ionic compounds brittle? [2]:
- When moved ions no longer have alternating arrangement
| - so lattice breaks cus same charge ions repel each other
52
When do Covalent bonds occur?
Form between elements that have high electronegativity values (non-metals)
53
How many pairs of electrons are in a double bond?
2 pairs of electrons
54
How many pairs of electrons are in a triple bond?
3 pairs of electrons
55
What are the two covalent structures?
- (simple) molecular
| - Macromolecular (giant covalent)
56
What are sum examples of simple covalent molecules? [3]:
- CH₄
- H₂O
- H₂
57
What are sum examples of macromolecular covalent structures? [2]:
- Diamond
| - Graphite
58
When is a coordinate bond formed?
when BOTH electrons are donated by the same atom
59
Coordinate bond =
dative covalent
60
What is an example of a coordinate bond?
NH₄⁺
61
How do u draw a coordinate bond?
with an arrow
62
What type of covalent strucure is ice?
(simple) molecular crystal
63
Ice between 0-100 degrees Celsius (water) [2]:
- Hydrogen bonding exists between molecules
| - H-bonds hold molecules close together but allow them to move freely (liquid)
64
Ice (at 0 degrees celsius)- properties [4]:
- Molecules have less energy so H-bonds FIX molecules in position (solid)
- Ice has a 3D hexagonal structure
- Floats in water
- Relatively high melting point
65
What causes ice to be able to float in water? [2]:
- Spaces created between molecules, causing it to expand
| - So it becomes less dense (can float)
66
What causes the relatively high melting point of ice? [2]:
- Strong hydrogen bonds between molecules
| - Requires more energy to break
67
Properties of iodine as a molecular crystal structure [4]:
- Low melting/ boiling point
- Sublimes into purple gas
- Slightly soluble in water
- Shiny grey solid at room temp
68
Molecular crystal structures [2]:
- Ice (H₂O)
| - Iodine (I₂)
69
Molecular crystal structure- iodine [3]:
- I₂ is a covalently bonded molecule
- Covalent bonds very strong
- Weak Induced dipole (VDW) forces exist between molecules to give its crystal structure
70
What gives I₂ its crystal structure?
Weak Induced dipole (VDW) forces exist between molecules to give its crystal structure
71
What causes I₂'s (crystal) low melting & boiling points [2]:
- Weak VDW's forces between molecules are v easily broken
| - I₂ molecule stays intact cus diatomic
72
Why is I₂ only slightly soluble in water? [2]:
I₂ is non-polar
| - KI or I⁻ dissolve better/ helps it dissolve
73
Diamond- Structure [3]:
- Each C atom forms 4 single covalent bonds with 4 other atoms
- Forms 3D solid lattice of Carbon
- Tetrahedral arrangement (109.5)
74
Properties of diamond [4]:
- Very Hard
- Very high melting/ boiling point
- Electrical insulator
- Insoluble
75
Why is diamond hard?
Due to strong Covalent bonds
76
Why does diamond have a high melting/ boiling point? [2]:
- Strong covalent bonds between molecules
| - Require lots of energy to break
77
Why is diamond a poor conductor? [2]:
- All e⁻s involved in bonding
| - so no delocalised electrons to carry a charge
78
Why's diamond insoluble? [2]:
- Due to strong c-c bonds
| - water can't break through them
79
Graphite- structure [5]:
- Each C atom forms covalent bond with 3 C atoms
- 4th C is delocalised
- Trigonal planar (120 degrees)
- Forms 2D hexagonal structure (1 atom thicc)
- Weak VDW forces between layers
80
Properties of Graphite [4]:
- Soft
- Very high melting/ boiling point
- Electrical conductor
- Insoluble
81
Why is Graphite soft? [2]:
- Weak VDW forces between layers
| - allows layers to slide over each other when force is applied to it
82
Why does graphite have a high melting point? [2]:
- Strong c-c covalent bonds
| - Require a lot of energy to break
83
Whys is graphite a good electrical conductor?
Delocalised electrons are free able to carry a charge
84
Why is graphite insoluble?
Due to very strong c-c covalent bonds
85
Where does metallic bonding happen?
Occurs between metals and alloys
86
Metallic crystal structure =
A regular lattice of cations surrounded by a sea of delocalised electrons
87
Properties of metallic bonding [5]:
- High melting points
- Conduct electricity
- Conduct heat
- Malleable
- Ductile
88
Why do metallic structures have high melting points?
Due to strong electrostatic attractions
89
Why are metallic compounds good conductors of electricity?
Delocalised electrons free to carry a charge
90
Why are metallic compounds good conductors of heat?
Delocalised electrons can quickly transfer energy
91
Why are metallic compounds malleable & ductile? [2]:
- Delocalised electrons able to move with the ion when force is applied
- They maintain the non-directional electrostatic attractions thus, maintain their shape
92
Repulsion in atoms =
Electron pairs around central atom repel each other & settle in a position to minimise repulsion
93
__=
Flat to paper
94
- - - - =
Into paper (away from u)
95
|\ =
Coming out of paper (towards u)
96
what are the combinations of electron pairs [3]:
- LP-LP
- LP-BP
- BP-BP
97
LP-LP =
Most repulsive
98
BP-BP =
Least repulsive
99
Linear shape examples [2]:
- BeCl₂
| - CO₂
100
Bent shape [example]:
H₂O
101
Trigonal planar example:
BF₃
102
Trigonal Pyramidal example:
NH₃
103
Tetrahedral shape examples [2]:
- CH₄
| - NH₄⁺
104
Trigonal Bipyramidal examples:
PCl₅
105
Octahedral shape example:
SF₆
106
Square Planar example:
ClF₄⁻
107
How to find the shape of an unknown molecule [5]:
1. Find formula
2. Identify central atom
3. Draw the molecule
4. Add all bonding electrons
5. Add 'satellite'/ valence electrons
108
electronegativity in a covalent bond =
A diff in electronegativity causes pairs of electrons in the bond to be shared unequally
109
What is the strongest intermolecular force?
Hydrogen bonding
110
when does a hydrogen bond? lol
Occurs btwn molecules that contain a H atom DIRECTLY bonded to N, O, or F
111
What are the simple molecules with hydrogen bonding? [3]:
- Hydrogen Flouride
- Water
- Ammonia
112
Groups in organic chem with hydrogen bonding [3]:
- Alcohols
- Carboxylic acids
- Amines
113
What is a hydrogen bond?
essentially its a VERY STRONG dipole
114
What is the weakest IMF force?
Induced dipole (VDW)
115
Induced dipole forces [2]:
- Occurs btwn ALL molecules
| - Strength varies ( Increases with increasing Mr cus more electrons)
116
Increased vdw = [3]:
- Increased melting/ boiling point
- this cus it would be a bigger molecule
e. g longer alkanes have higher melting/boiling point
117
Permanent dipole forces [3]:
- Occur btwn polar molecules (they have diff in electronegativity)
- Stronger than induced/vdw but weaker than h bonding
- Dipole-dipole IMF causes compounds to have higher than expected melting/ boiling points
