Gases, Liquids and Solids:• Swimmers (and all Canadians who have been

rescued by the Coast Guard!) know that the human body is slightly less dense than water but rather more dense than air.

• Liquids and solids (high density) have small molar volumes. Gases have much larger molar volumes at “normal” temperatures and pressures. Gases are “easily” compressed!

Table 1Substance ( Phase ) Experimental Molar Volumes (L.mol-1) ( 00

C and 1.00 bar)

H2O(l) 0.01802

H2O(s) 0.01965 (less dense than H2O(l))

C2H5OH(l) 0.0584

NaCl(s) 0.0270

He(g) 22.72

H2(g) 22.72

N2(g) 22.70

CO2(g) 22.56

Table 2Substance ( Phase ) Densities ( g.L-1) at 00 C and 1.00 bar

H2O(l) 999.8

H2O(s) 917.0 Icebergs!

C2H5OH(l) 789 George Street.

NaCl(s) 2165

He(g) 0.176 Party Balloons!

H2(g) 0.0887

N2(g) 1.234

CO(g) 1.234 Life is too tough??

CO2(g) 1.951 Photosynthesis.

Condensed Phases and Gases at “Normal” Temperatures and Pressures • Condensed phases (solids and liquids) have

relatively small molar volumes and “high” densities.

• Gases have relatively high molar volumes and “low” densities

• Simple Explanation – in condensed phases molecules are “touching” each other – no “empty” space.

Gases at “Normal” Temperatures and Pressures

• Gases are mostly empty space – and are thus easily compressed. This is not true at very high P and low T. (Demonstration with dry ice!)

• Gases at low pressure can be condensed if subjected to a higher (external) pressure. Gases at high pressure will expand if the external pressure is reduced (propane barbecue).

• There are many(!) pressure units.

Pressure Units

• By definition: Pressure = Force/Area• “Old” units for P: lb.in-2, mm Hg or torr• Modern or SI pressure units• P = Force/Area = N/m2 = kg.m s-2/m2 = Pascal• Standard atmospheric pressure = 101.325 kPa• 101.325 kPa = 1.01325 x 105 Pa (usual metric

abbreviations)• We often measure atmospheric pressure using a

barometer containing Hg or another liquid.

The gaseous state of three halogens (group 17)Figure 6-1

Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 6 Slide 7 of 19

Liquid PressureFigure 6-3

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liquid pressure is directly proportional to the liquid density and the height of the liquid column

P (Pa) =

AF =

AW =

Ag x m =

Ag x V x =

Ag x h x A x = g x h x

Measurement of atmospheric pressure with a mercury barometer

Figure 6-4

Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 6 Slide 9 of 19

Standard Atmospheric Pressure 1.00 atm, 101.325 kPa, 1.01325 bar, 760 torr, ~760 mm Hg

Measurement of gas pressure with an open-end manometer

Figure 6-5

Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 6 Slide 10 of 19

Relationship between gas volume and pressure – Boyle’s Law

Figure 6-6

Copyright © 2011 Pearson Canada Inc. General Chemistry: Chapter 6 Slide 11 of 19

PV = constantP 1V

6-2 Simple Gas Laws

Boyle’s Law:• The equation PV = constant is valid for a fixed

amount of a particular gas at a fixed temperature. One could take two points on the previous graph say (V1,P1) and (V2,P2) and write

• P1V1 = P2V2 = constant or just P1V1 = P2V2 • This expression can be used to predict, for

example, how the volume of a gas will change when the pressure is altered or….? We call this an initial state → final state problem.

Class Example – Boyle’s Law:• At a particular temperature and a pressure of

242 kPa a sample of argon gas Ar(g) has a volume of 3.87 L. What will be the gas volume if the pressure is reduced to 88.6 kPa? (Mention the trichotomy axiom?)