States of Matter: Solid, Liquids and Phase Changes Flashcards

1
Q

crystalline solids

A
  • well-ordered matter within the solid
  • arrangement of atoms in the solid repeats itself
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2
Q

amorphous solids

A

don’t have extensive ordering of particles

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3
Q

types of crystalline solids

A
  • molecular solids
  • covalent network solids
  • metallic solids
  • ionic solids
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4
Q

molecular solids

A

molecules held together by intermolecular forces, with relatively low melting points (eg ice or benzoic acid)

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5
Q

covalent network solids

A

extended structures of atoms held together by covalent bonds with very high melting points

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6
Q

allotropes

A

different structural forms of an element

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7
Q

metallic solids

A
  • metallic bonding between atoms
  • metal atoms as cations in sea of delocalised electrons
  • high electrical conductivity
  • malleable, ductile
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8
Q

ionic solids

A

held together by electrostatic attraction between cations and anions, with high melting points

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9
Q

describe how most liquids are molecular

A

intermolecular forces keep particles close together but not strong enough to keep particles from moving past each other

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10
Q

define surface tension

A

amount of energy required to expand a liquid surface

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11
Q

how does surface tension work?

A

at top - intermolecular attractions from below mean that the net attraction is down, causing surface to contract.
in the middle - there are intermolecular attractions in all directions; no net attraction

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12
Q

how does strength of forces between particles in a liquid affect surface tension?

A

the stronger the forces between particles in a liquid, the greater the surface tension

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13
Q

define capillary action

A

the rising of a liquid in a narrow space against the pull of gravity

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14
Q

cohesive molecules

A

between molecules

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15
Q

adhesive forces

A

between molecules and container walls

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16
Q

how does water rise in a capillary tube?

A
  • strong H-bonding interactions between the water and glass (SiO2) pull the water up
  • gravity and cohesive forces pull the liquid down
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17
Q

diagram for water in a tube when adhesive>cohesive and vice versa

A
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18
Q

gas -> solid

A

deposition

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19
Q

solid -> gas

A

sublimation

20
Q

sold -> liquid

A

fusion (melting)

21
Q

liquid -> gas

A

vaporisation

22
Q

gas to liquid

A

condensation

23
Q

liquid to solid

24
Q

attractive forces in solids, liquids, gases

A

solid: many attractive forces
liquid: some attractive forces
gas: no attractive forces

25
breaking attractive forces
requires (absorbs) energy: endothermic, so ∆H is positive - adding heat increases temperature, increases KE, and overcomes attractive forces
26
forming attractive forces
releases energy: exothermic ∆H is negative
27
graph for phase changes
moving from one state of matter to another
28
how does ∆H of vaporisation relate to ∆H of fusion?
∆H(vap)>∆H(fus) for all substances
29
what type of process is a phase change?
a spontaneous process
30
spontaneous
∆G (change in free energy) is negative
31
equation for ∆G
∆G=∆H-T∆S both ∆H and ∆S determine whether a process occurs spontaneously
32
∆H =
change in enthalpy due to phase change
33
∆S =
change in entropy due to phase change
34
entropy, S
the greater the degree of randomness or disorder in a system, the greater its entropy
35
∆S =
S (final) - S (initial)
36
for a reaction to be favourable, ∆S must be
positive
37
solid -> liquid -> gas requires
heat and increases entropy
38
describe gas-liquid equilibrium and draw a graph for it
as equilibrium between Rate(vap) and Rate(cond) is achieved, gas pressure reaches a constant value
39
vapour pressure
the pressure of gas in equilibrium (co-existing) with its liquid (or solid) at a specified temperature
40
vapour pressure depends on
temperature and the strength of intermolecular forces
41
effect of boiling point on vapour pressure
- the temperature at which vapour pressure of a liquid equals the external pressure - enough kinetic energy to vaporise against external pressure; water can form vapour bubbles within
42
normal boiling point
boiling point at 1atm
43
effect of intermolecular forces on vapour pressure
greater intermolecular forces: - fewer molecules escape liquid - more KE required to overcome
44
why Hvap is greater than H fus
In the change from solid to liquid, the kinetic energy of the molecules must increase only enough to partially offset the intermolecular attractions between molecules. In the change from liquid to gas, the kinetic energy of the molecules must increase enough to overcome the intermolecular forces
45
does Xe have a lower or higher boiling point than Iodine
Xe would have a lower boiling point than iodine. Both are nonpolar with dispersion forces, but the forces between xenon atoms would be weaker than those between iodine molecules since the iodine molecules are more polarizable because of their larger size
46