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AS Chemistry > Chapter 5: States of Matter > Flashcards

Chapter 5: States of Matter Flashcards

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1
Q

States of matter are characterized by

A

Different energies of particles

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2
Q

Kinetic theory of matter:

A

Energy of particles is directly proportional to the temperature in the Kelvin scale of the system.

State of matter at given temp and pressure is determined by strength of intermolecular force between particles

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3
Q

Properties of particles in solid, liquid, gas states

A

S: closely packed, inter particle force strong, vibrate in position, fixed shape and volume
L: more spread, force weaker, slide over each other, no fixed shape, fixed volume
G: spread out, force negligible, move freely, no fixed shape and volume

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4
Q

Explain the heating curve:

A

Solid is heated, vibrational E increases and so does temp. Line going up

Melting point. Vibrations enough energetic for molecules to move away from fixed positions to form liquid. E added at this stage used to break intermolecular force, not raise KE, so temp (line) constant

Liquid heated, KE gained, temp goes up

Boiling point. Enough energy to break all intermolecular forces and form gas. Needs more E than melting as all forces must break. KE and hence temp contstant. Gas bubbles visible throughout liquid

Gas heated under pressure, KE and temp increase

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5
Q

Basic assumptions of kinetic theory as applied to a ideal gas:

A

Ideal gas strictly obeys these:

  1. Molecules are in continuous random motion
  2. Volume occupied by molecules negligible compared to total volume occupied by gas. Ie molecules are point masses
  3. Intermolecular force of attraction is negligible
  4. Colliding molecules, perfectly elastic, no loss/gain of E
  5. Pressie of gas is due to collision of molecules against containers walls
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6
Q

Temp in Kevin is known as

A

Absolute temp

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7
Q

Absolute zero, 0K, is…

A

Point of zero kinetic energy of particles

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8
Q

0K=

A

-273°C

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9
Q

Celsius –> Kelvin ?

A

Add 273

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10
Q

How many Pa in 1 bar

A

10^5 ie 1 atm

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11
Q

Cm3 –> dm3 ??

A

Divide by 1000

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12
Q

1000cm3=xdm3

A

1dm3

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13
Q

Dm3 –> m3 ??

A

Divide by 1000

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14
Q

1m3=xdm3

A

1000dm3

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15
Q

Boyle’s law:

A

Volume occupied by fixed mass of gas varies inversely with it’s pressure at constant temp

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16
Q

Volume of gas=

A

Containers volume

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17
Q

Pressure increases as

A

Frequency or energy of collisions increases

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18
Q

Boyle’s law relationship

A

P α 1/V

PV is a constant

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19
Q

Graphs of Boyle’s law

A

V X axis, p y axis. Decreasing curve

1/V X axis, P y axis. Straight line through origin

P X axis, pv y axis. Constant (—)

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20
Q

Charles law:

A

Volume of given mass of gas directly proportional to it’s absolute temp, at constant pressure.

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21
Q

Charles law relationship

A

VαT

V/T= a constant

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22
Q

Charles law graph

A

Temp/K X axis, V y axis. Straight line through origin which is absolute 0 (0K)

23
Q

Relationship between temp and pressure

A

Increase in temp increase average KE of particles and they move faster and collide walls with more energy and frequency, raising pressure.

If volume is held constant pressure directly proportional to absolute temp
PαT
P/T= a constant

24
Q

Graph of PαT

A

Temp/K X axis, P y axis. Straight line through 0K at origin

25
Q

Charles laws:

A

VαT at const P

PαT at const V

26
Q

Combined gas law equation

A

PV/T= a constant
OR
P1V1/T1=P2V2/T2 (1-ini, 2-fin)

27
Q

Ideal gas equation:

A

PV=nRT
Derived from 3 gas laws for fixed mass of gas
Practice the derivation

28
Q

Avogadros law:

A

V α n

At const temp, and pressure, equal volumes of gases contain equal no of moles

29
Q

What is R?

A

Universal gas constant

30
Q 
When using ideal gas equation,
P must be in: 1.
V: 2.
No of moles: 3.
Temp: 4.
A
  1. Pa
  2. m3
  3. m/M
    4: K
31
Q

How do real gases behave

A

Deviated to some extent from ideal behavior

32
Q

Ideal gas defined as

A

Obeys ideal gas law PV=nRT under all conditions.

One mole of gas relationship PV/RT=1.
So graph of PV/RT against P for 1 mole is a horizontal line of intercept 1

33
Q

Real gas almost behaves like an ideal gas at

A

Low pressure

High temp

34
Q

Real gas shows greatest deviation at

A

High pressure

Low temp

35
Q

How to interpret when real gases behave most and least like ideal gases

A

Question validity of 2 assumptions from kinetic molecular theory

1) volume of gas particles negligible
2) no attractive forces between particles

36
Q

Explain why real gases can behave as ideal gases at low pressure, and not high in terms of volume

A

At low, volume of particles is about 0.05% of total so it’s valid to call it negligible

At high, space between then is reduced so percentage volume of particles increases to 20% of total - not negligible.

As a result, volume of real gas at high pres is larger than predicted from ideal gas law

37
Q

Explain why real gases can behave as ideal gases at low pressure, and not high in terms of force of attraction

A

At low pressure, particles are wide spread so forces between are negligible. But high pressure, particles are closer so attractive forces strengthen. Reduce effect of pressure of gas.

Low temp increase deviation because lower KE increases strength of forces.

38
Q

Which conditions being invalid make a real gas deviate from an ideal gas

A

High pressure and low temp molecules are close and volume occupied by individual molecules can’t be ignored compared to total. And intermolecular force of attraction is large and can’t be ignored

39
Q

Particles in liquid :

A
  • have more KE than those in solid state
  • not in fixed positions
  • in constant motion, bouncing or colliding off eachother
  • closely packed but less than solid
  • attrsctive forces less than in solid
40
Q

Melting:

A

Solid changes its state when heated. Change happens because particles gain KE and vibrate with greater amplitude.

41
Q

Melting point:

A

Temp at which energy supplied is sufficient to over come binding force

42
Q

Vaporization :

A

Liquid surface is flat, expect near walls of container. this is because particles at surface are attracted inwards and sideways by other particles.
Fastest moving surface particles escape into gas state.

43
Q

Why does temp decrease with evaporation

A

Escape of fastest particles decreases average KE of liquid so temp of liquid decreases with evaporation.

44
Q

What increases rate of evaporation

A

Higher the temp the faster rate of evaporation.

45
Q

Evaporation takes place at what temp

A

All the time

Because particles escape from surface all the time

46
Q

Evaporation in closed container

A

Escaped particles collect in space above surface and collide with eachother and walls

47
Q

Vapor pressure is

A

Vapor formed by escaping particles from particles in closed container exert a pressure

48
Q

Vapor pressure depends on:

A

Temp of liquid

Characteristics of liquid as force holding particles is different for each

49
Q

Saturated vapor pressure

A

As more and more particles escape molecules in vapor become close together and some condense back by hitting liquid surface.

When 2 opposing changes are at same time an equilibrium is reached (with molecules going from liquid to vapor as same rate as vapor to liquid) between no of particles leaving and reentering liquid when vapor pressure remains constant at given temp.

Pressure exerted by vapor at equilibrium is sat vap pres

50
Q

Boiling

A

Liquid in contact with air is heated, it’s vapor pressure increases until a temp is reached at which vapor pressure = atmospheric pressure.

51
Q

Liquid boils when …

Bp relationship with vapor pressure

A

Saturated vapor pressure = atmospheric pressure

Bp increases as vapor pressure does

52
Q

Saturated vapor pressure is a characteristic of liquid. Meaning?

A

Constant for given liquid at given temperature (+temp, +KE, able to overcome)

53
Q

Why does vapor condense back

A

Molecules in vapor come close as more molecules escape. Eventually molecules with lower KE are not able to over come attractive force of neighbouring molecules

54
Q

At equilibrium concentration of water molecules in vapor is

A

Constant

AS Chemistry (11 decks)

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