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Chemistry > 2- Bonding and Structure > Flashcards

2- Bonding and Structure Flashcards

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1
Q

Properties of metals

A

high melting temperatures
good conductivity of heat and electricity
malleable
ductile

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2
Q

Metallic bond

A

the strong electrostatic attraction between the nuclei of metal cations and delocalised electrons

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3
Q

Why do metal generally have high melting temperatures

A

giant lattice structure has a lot of electrostatic forces which need a lot of energy to overcome

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4
Q

How does the size of the cation affect metal melting temperatures

A

the smaller the delocalised electron the closer the delocalised electron is to the nucleus
this increases the electrostatic forces so increases the energy required

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5
Q

Do group 1 metals have a high or low melting temperature

A

lower

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6
Q

Do group 2 metals have high or low melting temperatures

A

higher

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7
Q

Do d-block metals have high or low melting temperatures

A

high as they have more delocalised electrons

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8
Q

Why are metals good electrical conductors

A

when a potential difference is applied across the ends of the metal the delocalised electrons will be attracted so move towards the positive terminal
this movement causes an electrical charge

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9
Q

Why can metals conduct heat

A

free moving delocalised electrons pass kinetic energy along
cations closely packed together pass kinetic energy

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10
Q

Why are metals malleable and ductile

A

when stress is applied to a metal the layers of cations slide over each other because they are equal sizes
the delocalised electrons move with the cations to prevent strong forces of repulsion

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11
Q

Ionic bond

A

regular array of oppositely charged ions throughout a giant lattice confined to solid materials

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12
Q

How can you determine the strength of an ionic bond

A

calculating the amount of energy required in one mole of solid to separate the ions to infinity (in the gas phase)
When they are at an infinite distance from one another so the ions can no longer interact

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13
Q

How does ionic radius affect the strength of an ionic bond

A

As the ionic radii get larger less energy is required as there are more electron shells so more shielding which decreases the forces of attraction as this affects how closely they are packed together

As you go down group 1 and 7 the ionic radii get larger so less energy is required

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14
Q

Isoelectronic

A

same number of electrons so same electronic configuration
Na+, Mg2+, Al3+

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15
Q

Properties of ionic compounds

A

high melting temperatures
brittleness
only conduct when molten aqueous
often soluble in water

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16
Q

Why ionic compounds have high melting temperatures

A

giant lattice of oppositely charged ions so the electrostatic forces are large so require a lot of energy for ions to break free and slide past each other

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17
Q

why ionic compounds don’t conduct when solid

A

no free moving electrons and ions are not free to move

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18
Q

Why ionic compounds conduct when molten or aqueous

A

ions are mobile and will migrate to electrodes of opposite signs when a potential difference is applied

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19
Q

Why ionic compounds are soluble in water

A

energy required to break apart the lattice structure and separate ions can be supplied by the hydration of the separated ions produced

both positive and negative ions are attracted to water due to its polarity

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20
Q

Why ionic compounds are brittle and hard crystalline substances

A

when a stress is applied layers slide over each other and if the same charges are side by side they repel each other and break apart

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21
Q

Evidence for ions existing

A

electrolysis
positive ions migrate towards negative electrode
negative ions migrate towards positive electrode

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22
Q

Covalent bond

A

formed by the overlap of two atomic orbital containing a single electron each

it is the strong electrostatic attraction between nuclei of two atoms and the bonding pair of electrons

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23
Q

Sigma σ bond formation

A

end on overlap of two orbital
single bond

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24
Q

Pi π bond formation

A

a sideways overlap of p orbitals
double or triple bonds
only formed when a sigma bond is formed first

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25
Bond length
the distance between nuclei of two atoms that are covalently bonded
26
Bond strength
the amount of energy required to break one mole of the bond in the gaseous state
27
What happens when bond length increases
bond strength decreases as electron will be further away from nucleus reducing electrostatic attraction so less energy is required
28
Electronegativity
the ability if an atom to attract a bonding pair of electrons
29
Trends in electronegativity
decreases down a group increases across a period
30
Describe the electronegativity in non-polar bonds
they have the same electronegativity or the difference is below 0.4
31
Describe the electronegativity in polar covalent bonds
one atom has a higher electronegativity so can pull more of electron cloud towards itself difference between 0.4 and 1.8
32
Discrete (simple) molecule
electrically neutral group of two or more atoms held together by chemical bonds
33
Octet rule
a stable atom has eight electrons in its outer shell but there are exceptions
34
Exceptions to the octet rule
BeCl2 (4) BCl3 (6) PCl5 (10) SF6 (12)
35
Dative covalent bond
when both electrons in the bonding pair are supplied by only one of the atoms involved formed when an empty orbital of one atom overlaps with an orbital containing a non-bonding pair (lone pair) of another atoms represented by an arrows
36
Examples of dative covalent bonds
H3O- NH4+ Al2Cl6
37
Electron pair repulsion theory
shape of molecule or ion is caused by repulsion between electrons around the atoms electron pairs arrange themselves around the central atom so the repulsion is minimum Increasing repulsion: -bonding pair - bonding pair -lone pair - bonding pair -lone pair - lone pair
38
Normal bond ---
lies on plane
39
Dashed bond - - - -
extends backwards
40
Wedged bond
extends forwards
41
Linear shape
2 bond pairs 0 lone pairs bond angle= 180° example: BeCl2
42
Trigonal planar
3 bond pairs 0 lone pairs bond angle= 120° example: BCl3
43
Tetrahedral
4 bond pairs 0 lone pairs bond angle= 109.5 example: CH4
44
Trigonal bipyramidal
5 bond pairs 0 lone pairs bond angle= 90° and 120° between forwards and backwards bonds example: PCl5
45
Octahedral
6 bond pairs 0 lone pairs bond angle= 90° example: SF6
46
Trigonal pyramidal
3 bond pairs 1 lone pair bond angle: 107° example: NH3
47
V-shaped
2 bond pairs 2 lone pairs bond angle= 104.5° example: H2O
48
Dipole
exist when two charges of equal magnitude but opposite signs are separated shown by an arrow pointing to atom with greater electronegativity
49
How can dipoles be used to determine if a molecule os polar or non-polar
If the dipoles cancel out (go in opposite directions) then the molecule is non-polar If dipoles reinforce each other the molecule so will have an overall dipole so it is polar
50
Electron pair repulsion theory
- shape of molecule is caused by repulsion between electrons -electron pairs arrange themselves so there is minimal repulsion - increasing repulsion order: -bonding pair- bonding pair -bonding pair- lone pair -lone pair- lone pair
51
Intermolecular forces
forces holding difference molecules together that are easy to break ( broken when changing state)
52
Types of intermolecular forces
in increasing strength order -london forces (Van der Waals) - permanent dipole - hydrogen bonding
53
London forces
when a dipole is induced
54
How can london forces be made stronger
- larger molecule = more points of contact - more electrons
55
Permanent dipoles
exists between already polar molecules
56
What happens between molecules when a dipole is formed?
one flips to avoid repulsion
57
Hydrogen bonds
a bond between hydrogen and a more electronegative element
58
Between what elements can hydrogen bonds form
O-H N-H F-H
59
Why do alcohols have higher bps than alkanes?
hydrogen bonding is present in alcohols as well as LDFs while alkanes only have LDFs so alcohols require more energy to break the intermolecular bonds
60
Why do bts increase with bigger alkanes?
increasing chain length increases LDF strength as there are more electrons and more points of contact
61
Why do branched molecules have lower bts than unbranched?
less points of contact = weaker LDFs
62
Why is ice less dense than water
hydrogen bonds in ice form hexagonal rings with large open spaces but when melted the molecules come closer together increasing density
63
Why does water have an anomalous boiling temp
it is unusually high due to the two lone pairs of electrons
64
What are the conditions for dissolving?
- forces of attraction between solute and solvent particles must be stronger to overcome solvent solvent and solute solute forces -solute particles must be separated and then surrounded by solvent particles
65
Describe why NaCl can dissolve in water?
- delta negative end of water attracts Na sufficiently as electronegativity Na-O is greater than Na-Cl so Na+ ions removed from lattice in an ion-dipole interaction (exothermic) -chlorine ions will be left with Na so surrounded by water
66
Why can alcohols dissolve in water?
contain OH groups that can form hydrogen bonds with water
67
Name 4 properties of giant covalent lattices
- high melting and boiling points -do not conduct electricity - hard and strong - insoluble in water
68
Why do giant covalent structures have high melting and boiling points?
many covalent bonds require a lot of energy to overcome
69
Do giant lattices have intermolecular forces?
no
70
Why don't giant covalent structures conduct electricity?
there are no free ions or electrons to carry the charge
71
Why are giant covalent structures insoluble in water?
they are non-polar so cannot form hydrogen bonds with water
72
Why is graphite an exception?
-it conducts - it is soft and slippery - has intermolecular forces
73
Why can graphite conduct?
each carbon is bonded to three carbon atoms not four so the p-orbitals overlap creating a cloud of delocalised electrons to carry a charge
74
Why is graphite slippery?
layers can slide over each other despite weak LDFs
75
Name 4 giant covalent lattices
-diamond -graphite -graphene -silicon (IV) oxide
76
Describe the structure of diamond
- each carbon forms four sigma bonds to other carbons - 3D tetrahedral with 109.5 bond angles
77
Describe the structure of graphite
- layered - each carbon forms 3 sigma bonds -interlocking hexagonal rings

Chemistry (4 decks)

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