3.1.3 - Bonding Flashcards
1
Q
Define ionic bonding
A
Electrostatic force of attraction between oppositely charged ions formed by electron transfer
2
Q
Ionic Bonding
Metal atoms ___ electrons to form ___ ions
A
Metal atoms lose electrons to form +ve ions
3
Q
Ionic Bonding
Non-metal atoms ____ electrons to form ___ ions
A
Non-metal atoms gain electrons to form -ve ions
4
Q
State the formula for a carbonate ion
A
CO32-
5
Q
State the formula for an ammonium ion
A
NH4+
6
Q
Name the structure of ionic crystals
A
Giant Ionic Lattice
7
Q
Name 3 physical properties of ionic compounds
A
- Conduct electricity only when they’re molten or dissolved
- High melting points
- Tend to dissolve in water
8
Q
Why can ions conduct electricity when they’re molten or dissolved?
A
∵ ions in liquid are free to move and carry a charge
9
Q
Why can’t ions conduct electricity when they’re in a solid?
A
∵ ions are in fixed position by strong ionic bonds
10
Q
Why do ionic compounds have high melting points?
A
- Giant ionic lattices
- Strong electrostatic forces of attraction between oppositely charged ions
- Takes a lot of energy to overcome these forces
11
Q
Why do ionic compounds tend to dissolve in water?
A
- Water molecules are polar
- Part of molecule has a small negative charge and other bits have small positive charges
- Charged parts pull ions away from lattice = causing it dissolve
12
Q
Ionic bonding is stronger and melting points are higher when ions are… (2x)
A
smaller and/ or have higher charges
13
Q
When do molecules form and how are they held together?
A
- Form when 2 or more atoms bond together
- Held together by strong covalent bonds
14
Q
What do single covalent bonds contain?
A
Shared pair of electrons
15
Q
Describe covalent bonding
A
- Two atoms share electrons so they’ve both got full outer shells
- Both postive nuclei are attracted electrostatically to shared electrons
16
Q
Multiple covalent bonds contain…
A
multiple shared pairs of electrons
17
Q
Draw methane, represent the covalent bonds by drawing lines
A
18
Q
Why can carbon form giant covalent structures?
A
∵ they can form 4 covalent bonds
19
Q
Describe the structure of graphite
A
- Carbon atoms are arranged in sheets of flat hexagons covalently bonded with 3 bonds each
- 4th outer electron of each carbon atom is delocalised
- Sheets of hexagons are boned together by weak van der Waal forces
20
Q
Name 5 properties of graphite
A
- Low density
- Dry lubricant/slippy
- Electrical conductor
- Insoluble in any solvent
- Very high melting point
21
Q
Explain why graphite has a low density
A
Layers are quite far apart compared to the length of covalent bonds
22
Q
Explain why graphite is an electrical conductor
A
‘Delocalised’ electrons aren’t attached to any particular carbon atoms & free to move along sheets carrying a charge
23
Q
Explain why graphite is a dry lubricant/slippy
A
Weak bonds between layers in graphite = easily broken ∴ sheets can slide over each other
24
Q
Explain why graphite has a very high melting point
A
Covalent bonds are very strong and require lots of energy to break
25
Explain why graphite is insoluble in any solvent
Covalent bonds in sheets are too strong to break
26
Describe the structure of diamond
* Each carbon atom is covalently bonding to 4 other carbon atoms (giant covalent structure)
* Tetrahedral shape
27
Name 5 properties of diamond
1. Very high melting point
2. Extremely hard
3. Good thermal conductor
4. Can't conduct electricity
5. Won't dissolve in any solvent
28
Why is diamond a good thermal conductor?
Vibrations travel easily through stiff lattice
29
Why can't diamond conduct electricity?
Outer electrons held in localised bonds
30
Why do diamond gemstones sparkle a lot?
Its structure makes it refract light a lot
31
What is dative covalent bonding (or co-ordinate bonding)?
When shared pair of electrons in covalent bond come from only one of the bonding atoms
32
Name an example of dative covalent bonding & explain how it is an example of this bonding
Ammonium ion (NH4+)
Forms when nitrogen atom in ammonia molecule donates a pair electrons to proton (H+)
33
Illustrate dative covalent bonding in an ammonium ion (NH4+)
34
Define metallic bonding
Metallic bonding is the electrostatic force of attraction between the positive metal ions and the delocalised electrons
35
Metals elements exist as...
giant metallic lattice structures
36
Describe metallic bonding
1. Outer shell electrons of metal are delocalised
1. Electrons free to move
2. Leaves positive metal ion
2. Positive metal ions attracted to delocalised negative electrons
1. Form lattice of closely packed positively ions in sea of delocalised electrons
2. This is metallic bonding
37
Name 4 properties of metals
* High melting points
* Good thermal conductors
* Good electrical conductors
* Insoluble (expect in liquid metals)
38
Why do metals have high melting points?
Strong electrostatic attraction between positive metal ions and delocalised sea of electrons
39
Why are metals good thermal conductors?
Delocalised electrons can pass kinetic energy to each other
40
Why are metals good electrical conductors?
Delocalised electrons can move and carry current
41
Why are metals insoluble?
Strong metallic bonds
42
Name 3 factors that affect the strength of metallic bonding
1. Number of protons/strength of nuclear attraction
2. Number of delocalised electrons per atom
3. Size of ion
43
Metallic Bonding
More protons = ....
stronger bond
44
Metallic Bonding
More delocalised electrons per atom = ....
stronger bonding
45
Metallic Bonding
Smaller the ion = ...
stronger the bond
46
Explain why Mg has stronger metallic bonding than Na and a higher melting point
1. In Mg: more electrons in outer shell that are released to sea of electrons
2. Mg ion is smaller and has more than one proton
3. ∴ stronger electrostatic attraction between positive metal ions and delocalised electrons = higher energy is needed to break bonds
47
Illustrate a giant ionic lattice of sodium chloride
48
Illustrate metallic bonding in magnesium
49
What does the shape of a molecule depend on?
The number of pairs of electrons in outer shell of central atom
50
Why are the bond angles between bonding pairs reduced when lone pairs of electrons are added?
∵ they're pushed together by lone-pair repulsion
51
\_\_\_\_\_\_\_ angles are the biggest
Lone-pair/lone-pair
52
\_\_\_\_\_\_\_ angles are the second biggest
Lone-pair/bonding-pair
53
\_\_\_\_\_\_\_ angles are the smallest
Bonding-pair/bonding-pair
54
Name the shape of a molecule with 2 electron pairs (& no lone pairs)
Linear
55
Draw BeCl2
State the bond angles of the molecule
2 electron pairs (& no lone pairs)
56
Name the shape of a molecule with 3 electron pairs (& no lone pairs)
Trigonal planar
57
Draw BF3
State the bond angles of the molecule
3 electron pairs (& no lone pairs)
58
Name the shape of a molecule with 4 electron pairs (& no lone pairs)
Tetrahedral
59
Draw NH4+
State the bond angles of the molecule
4 electron pairs (& no lone pairs)
60
Name the shape of a molecule with 3 electron pairs & 1 lone pair
Trigonal Pyramidal
61
Draw PF3
State the bond angles of the molecule
3 electron pairs & 1 lone pair
62
Name the shape of a molecule with 2 electron pairs & 2 lone pairs
Bent
63
Draw H2O
State the bond angles of the molecule
2 electron pairs & 2 lone pairs
64
Name the shape of a molecule with 5 electron pairs (& no lone pairs)
Trigonal Bipyramidal
65
Draw PCl5
State the bond angles of the molecule
5 electron pairs (& no lone pairs)
66
Name the shape of a molecule with 4 electron pairs & 1 lone pair
Seesaw
67
Draw SF4
State the bond angles of the molecule
4 electron pairs & 1 lone pair
68
Name the shape of a molecule with 3 electron pairs & 2 lone pairs
T-shaped
69
Draw ClF3
State the bond angles of the molecule
3 electron pairs & 2 lone pairs
70
Name the shape of a molecule with 6 electron pairs (& no lone pairs)
Octahedral
71
SF6
State the bond angles of the molecule
6 electron pairs (& no lone pairs)
72
Name the shape of a molecule with 4 electron pairs & 2 lone pairs
Square planar
73
Draw XeF4
State the bond angles of the molecule
4 electron pairs & 2 lone pairs
74
Predict the shape of the molecule H2S (show all your steps)
75
Define Electronegativity
The power of an atom to attract a pair of electrons in a covalent bond
76
\_\_\_\_ is most electronegative element
Fluorine
77
How are polar covalent bonds created?
In a covalent bond between 2 atoms of different electronegativities, bonding electrons will be pulled towards the more electronegative atom = makes bond polar
78
Some elements (e.g. C & H) have very similar electronegativities = bond essentially \_\_\_\_
non-polar
79
In a polar bond, difference in electronegativity between 2 atoms causes a ____ \_\_\_ to form
permanent dipole
80
What is a dipole?
Difference in charge between 2 atoms caused by shift in electron density in bond
81
Greater difference in electronegativity between atoms = ...
more polar the bond
82
When are molecules with polar bonds not polar and why is this?
When polar bonds are arranged symmetrically in molecule = charges cancel out & there's no permanent dipole
83
Name 3 Intermolecular Forces
* Permanent dipole-dipole forces
* Van der Waals forces or induced dipole-dipole forces
* Hydrogen bonding
84
What type of molecules have permanent dipole-dipole forces?
Polar molecules
85
Describe how permanent dipole-dipole forces form
In a substance made from molecules with permanent dipoles = they'll be weak electrostatic forces of attraction between δ+ and δ- charges on neighbouring molecules
86
Explain why if you put a charged rod next to a jet of polar liquid (e.g water), the liquid will move towards the rod
1. ∵ polar liquids contain molecules with permanent dipoles
2. (Doesn't matter if rod is postively or negatively charged)
3. Polar molecules in liquid can turn around so the opposite charged end is attracted towards the rod
87
Where are Van der Waals forces found?
Found between all atoms and molecules
88
Describe how Van der Waals forces form
1. Electrons in charge clouds = always moving quickly
1. At any moment, electrons in atom are likely to be more to one side than the other
2. At this moment = atom has temporary dipole
2. Dipole causes another temporary dipole in opposite direction on neighbouring atom
1. 2 dipoles are attached to each other
3. 2nd dipole causes another dipole in 3rd atom
4. Dipoles are being created and destroyed constantly ∵ electrons are constantly moving
1. Overall effect = atoms are attracted to each other
89
Describe and explain the structure of iodine at room temp
1. Iodine atoms are held together in pairs by strong covalent bonds to from I2 molecules
2. Molecules held together in molecular lattice arrangement by weak van der Waals attractions (this causes iodine to be solid at room temp.)
90
Name 3 factors that affect the strength of the Van der Waals forces
1. Size of molecules
2. Shape of molecules
3. Number of electrons
91
Explain how the size of molecules affects the strength of van der Waal forces
Larger molecules = larger electron clouds = stronger van der Waals forces
92
Explain how the shape of molecules affects the strength of van der Waal forces
Long, straight molecules lie closer than branched ones = closer together 2 molecules get = stronger the forces between them are
93
When does hydrogen bonding occur?
When hydrogen is covalently bonded to fluorine, nitrogen or oxygen
94
Hydrogen Bonding is the _____ intermolecular force
Strongest
95
Describe hydrogen bonding
1. F, N & O = very electronegative ∴ they draw bonding electrons away from hydrogen atom
2. Bond is polarised + hydrogen has high charge density = hydrogen atoms form weak bonds with lone pair of electrons on F, N or O atoms or other molecules
96
Molecules with H-bonding usually contain ____ or ____ groups
-OH or -NH groups
97
Draw hydrogen bonding occuring in water
Make sure O-H bonds are 180 degrees
98
Draw hydrogen bonding occuring in ammonia
99
Substances with h-bond have ____ boiling/melting points than similar molecules
Substances with h-bond have _higher_ boiling/melting points than similar molecules
100
Why do substances with h-bond have higher boiling/melting points than similar molecules?
∵ of extra energy needed to break h-bonds
101
Anomalously high boiling points of H2O, NH3 & HF are
caused by ___ \_\_\_\_ between molecules
Anomalously high boiling points of H2O, NH3 & HF are
caused by _hydrogen_ _bonding_ between molecules
102
Explain why ice is less dense than liquid water
1. As liquid water cools to form ice, molecules make more h-bonds & arranged themselves into regular lattice structure
2. In this structure, H2O molecules are further apart on average than molecules in liquid water
103
Explain why simple covalent compounds have lower melting/boiling points than macromolecules. (4)
1. To melt/boil simple covalent compound = just need to overcome the van der Waals forces between molecules
2. These forces are weak
3. To melt/boil macromolecules = many covalent bonds have to be broken
4. Covalent bonds = strong
104
Explain how the solubility of a substance in water depends on type of particles it contains
* Water = polar solvent
* ∴ substances that are polar or charged will dissolve in it
* Whereas non-polar or uncharged substances won't
105
Fill in the blanks
106
# Fill in the blanks
(3 examples)
107
# Fill in the blanks
(3 examples)
108
# Fill in the blanks
(3 examples)
109
110
Explain why CF4 has a bond angle of 109.5° (2)
* Around carbon = 4 bonding pairs of electrons
* ∴ these repel equally & spread as far apart as possible
111
State what is meant by macromolecular (1)
Means a giant molecule with covalent bonding
112
Predict the shape of AlCl4-. Draw a diagram of the specie to show its 3D shape. Name the shape and suggest a value for the bond angles. Explain your reasoning. (4)
113
**Application Question**
Perfume is a mixture of fragrant compounds dissolved in a volatile solvent. When applied to the skin the solvent evaporates, causing the skin to cool for a short time. After a while, the fragrance may be detected some distance away. Explain the observations. (4)
1. Solvent has low boiling point or weak intermolecular forces
2. Solvent needs energy, taken from the skin, to overcome intermolecular forces and evaporate
3. Perfume molecule slowly spreads through the room
4. By random diffusion of the perfume
114
Draw a diagram to show how 1 molecule of ammonia is attracted to 1 molecule of water. (3)
(hint h-bonding)
115
Draw NH3BCl3
116
What type of bonding does H3O+ have?
Dative Covalent Bonding
117
Fill in the blanks
118
Showing outer electrons only, draw a dot-and-cross diagram to indicate the bonding in calcium oxide (2)
Ionic Bonding
119
Explain why the boiling temperature of PH3 is greater than that of CH4 (3)
* PH3 has dipole-dipole
* between molecules
* stronger than in CH4
120
Explain why the H-F bond in HF is polar (2)
* Difference in electronegativity / F more electronegative than H
* _Bonding_ pair of _electrons attracted (drawn) towards F_ (nucleus) in the covalent bond
