Bonding and structure Flashcards
1
Q
define ionic bonding
A
the strong electrostatic attraction between positively and negatively charged ions IN A LATTICE
2
Q
define giant ionic lattice
A
oppositely charged ions strongly attracted in all directions in a repeating pattern
3
Q
why are ionic compounds soluble in water
A
- dissolves in polar solvents
- cations attracted to O, anions attracted to H
4
Q
why are ionic compounds electrical conductors when aq or molten but electrical insulators when solid
A
- when aq or molten, ions are mobile and can carry a charge
- when solid, the ions are fixed in place in a lattice structure
5
Q
why are ionic compounds solid at rtp
A
strong attractions between ions, lots of energy needed to overcome attractions and break the bonds
6
Q
define covalent bonding
A
strong electrostatic force of attraction between a shared pair of electrons and the nuclei of the bonded atoms
7
Q
define dative covalent bonding
A
both electrons in a covalent bond are donated by one atom
8
Q
examples of giant covalent lattices
A
- graphite
- diamond
- silicon dioxide
9
Q
examples of giant covalent polymers
A
- polymers
- fullerenes
10
Q
examples of simple covalent molecules
A
- methane
- oxygen
- hydrogen chloride
11
Q
examples of exceptions to the octet
A
- BF3 ( hypovalent - too little e- )
- SF6 (hypervalent - too many e-)
12
Q
properties of simple covalent molecules
A
- low melting and boiling points
- insoluble in water
- don’t conduct electricity as they have no freely moving charged particles
- may react with water
13
Q
define average bond enthalpy
A
- average amount of energy needed to break a specific type of bond homolytically in a gaseous molecule
- measurement of covalent bond strength
- always +ve as bond breaking is endothermic
14
Q
define metallic bonding
A
strong electrostatic attraction between cations and delocalised e-
15
Q
structure of metals
A
metal atoms tightly packed together in lattice structure
16
Q
are metals soluble
A
- no
- more likely to lead to reactions
17
Q
why do alloys have lower melting point
A
- disrupted lattices
- weaker attraction between cations and delocalised sea of e-
18
Q
how does charge density of metal affect bond strength
A
- small, highly charged cations have greater attraction to sea of delocalised electrons
19
Q
why are electron pairs as far apart as possible in molecules
A
to minimise repulsion
20
Q
properties of metals
A
- malleable - layers can slide over each other without breaking the bond
- ductile - layers can be pulled into a wire one atom thick
- high melting + boiling point - lot of energy needed to overcome the electrostatic attraction between cations and sea of delocalised electrons
21
Q
the shape of a compound or ion is determined by
A
- no of e- pairs around central atom
- the nature of these pairs: bonding or lone pairs
22
Q
what are bonding pairs
A
- pairs of electrons that are involved in bonding
- repel each other equally
23
Q
what are lone pairs
A
- pairs of e- that aren’t involved in bonding
- repel more than bonding pairs
- more electron dense
24
Q
what does each lone pair reduce the bond angle by
A
2.5 degrees
25
normal line represents
bond in plane of the paper
26
dotted wedge represents
bond going into paper, away from you
27
bold wedge represents
bond coming out of paper towards you
28
linear molecule (2 pairs) :
1. bond pairs
2. lone pairs
3. angle
4. example
1) 2 bond pairs
2) 0 lone pairs
3) 180 degrees
4) CO2
29
linear molecule (1 pair) :
1. bond pairs
2. lone pars
3. angle
4. example
1) 1 bond pair
2) 0 lone pairs
3) 180 degrees
4) H2
30
trigonal planar molecule :
1. bond pairs
2. lone pars
3. angle
4. example
1) 3 bond pairs
2) 0 lone pairs
3) 120 degrees
4) BF3
31
tetrahedral molecule :
1. bond pairs
2. lone pars
3. angle
4. example
1) 4 bond pairs
2) 0 lone pairs
3) 109.5 degrees
4) CH4
32
pyramidal molecule :
1. bond pairs
2. lone pars
3. angle
4. example
1) 3 bond pairs
2) 1 lone pair
3) 107 degrees
4) NH3
33
non- linear molecule :
1. bond pairs
2. lone pars
3. angle
4. example
1) 2 bond pairs
2) 2 lone pairs
3) 104.5 degrees
4) H2O
34
octahedral molecule :
1. bond pairs
2. lone pars
3. angle
4. example
1) 6 bond pairs
2) 0 lone pairs
3) 90 degrees
4) SF6
35
define electronegativity
the ability of an atom to attract bonding electrons in a covalent bond
36
define polar bond
A permanent dipole within a molecule containing covalently bonded atoms with different electronegativities
37
define polar molecule
A molecule containing polar bonds with dipoles that do not cancel due to their direction
38
factors affecting electronegativity
- atomic radius
- number of shells/ shielding
- nuclear charge
39
why does electronegativity decrease down a group
- more shells, so more shielding
- atomic radius increases (outer shell further from nucleus)
- less atraction between nucleus and outer e-
40
elements in the same period have the same...
number of electron shells
41
why does electronegativity increase across a period
- nuclear charge increases
- same shielding (no. of shells)
- atomic radius decreases (outer shells closer to nucleus)
- increased attraction between nuclei and outer e-
42
what's the most electronegative element
F
43
why do **symmetrical molecules** like CCl4 not have a permanent dipole
the dipoles within the molecule cancel out
44
why do **unsymmetrical molecules** like CH3Cl have permanent dipoles and are polar
the dipoles within the molecule don't cancel out
45
what are the 3 types of intermolecular forces
- London forces (induced dipole-dipole)
- permanent dipole-dipole interactions
- Hydrogen bonding (strongest)
46
what are london forces
- weak interactions
- exist between all molecules
- when there’s an **uneven distribution of e-** causing one end of the molecule to become temporarily charged due chance and this induces a charge in a nearby molecule
47
how does the size of an atom/molecule affect london forces
- larger atoms/molecules have more e- so stronger london forces
- stronger london forces = higher melting + boiling points
48
what are permanent dipole-dipole interactions
- stronger than london forces
- exists when there's a polar covalent bond
- attraction between partially +ve and partially -ve ends of a molecule
- can induce a non-polar molecule
49
what's hydrogen bonding
- strongest type of intermolecular force
- occurs when H is bonded to F,O,N
50
what are the unusual properties of water
- ice is less dense than water
- water has a relatively high melting and boiling point
51
why is ice less dense than water
- due to H bonding, water molecules in ice are further apart than in water
52
why does water have relatively high melting and boiling points
- water has H bonding as well as london forces
- each water molecule can form 4 H bonds so a large amount of energy is needed to break them
53
what's a simple molecular lattice
the solid structure of covalently bonded molecules attracted together by intermolecular forces
54
example of simple molecular lattices
- I2
- ice
- P
- S
55
why do simple molecular lattices have low melting + boiling points
weak London forces between molecules only require little energy to break
56
why can't simple molecular lattices conduct electricity
no charged particles to carry charge throughout the structure
57
why are most simple covalent lattices not soluble
most are not polar enough to be soluble in water but anything with H bonding will dissolve
