EJC C4 2023 The Gaseous State Lecture Notes (Teachers' Copy)

2023 JC1 H2 CHEMISTRY (9729) CORE IDEA 2: STRUCTURE AND PROPERTIES Topic 4: THE GASEOUS STATE Name: ___________________________________________ Civics Group: _________ Lecture Content 1 Introduction ................................ ................................ ................................ ........................ 2 2 The Gas Laws................................ ................................ ................................ ..................... 3 2.1 Boyle’s Law ................................ ................................ ................................ ................ 3 2.2 Charles’s Law ................................ ................................ ................................ ............. 4 2.3 Avogadro’s Law ................................ ................................ ................................ .......... 6 2.4 The Ideal Gas Equation ................................ ................................ .............................. 6 2.5 Dalton’s Law of Partial Pressures ................................ ................................ ............... 8 3 The Kinetic Molecular Theory of Gases ................................ ................................ ......... 11 4 The Behaviour of Real Gases ................................ ................................ ......................... 12 4.1 Effect of Low Pressure ................................ ................................ .............................. 13 4.2 Effect of High Temperature ................................ ................................ ....................... 14 4.3 Effect of Strong Intermolecular Forces ................................ ................................ ...... 15 5 Real Life Applications ................................ ................................ ................................ ..... 17 APPENDIX ................................ ................................ ................................ ............................... 18 APPENDIX 1 ................................ ................................ ................................ ..................... 18 APPENDIX 2 ................................ ................................ ................................ ..................... 19 APPENDIX 3 ................................ ................................ ................................ ..................... 20 REFERENCES 1 Chemistry (for CIE AS & A Level) by Peter Cann & Peter Hughes 2 Chemistry3: Introducing Inorganic, Organic and Physical Chemistry by Andrew Burrows, John Holman, Andrew Parsons, Gwen Pilling & Gareth Price 3 A-Level Chemistry by E.N. Ramsden Students should be able to: (a) state the basic assumptions of the kinetic theory as applied to an ideal gas (b) explain qualitatively in terms of intermolecular forces and molecular size: (i) the conditions necessary for a gas to approach ideal behaviour (ii) the limitations of ideality at very high pressures and very low temperatures (c) state and use the general gas equation pV = nRT in calculations, including the determination of Mr (d) use Dalton’s Law to determine the partial pressures of gases in a mixture This topic is not included in the H1 Chemistry Syllabus.

T h e G a s e o u s S t a t e E u n o i a J u n i o r C o l l e g e 2 | P a g e 1 Introduction Gases tend to have the following physical properties: • Gases have much lower density than solids or liquids. (Why is density low for gases?) • Gases assume the volumes and shapes of their containers. (Why ONLY gases assume the volumes of their containers? Why liquids do not assume the volume leh?) • Gases are highly compressible and their volumes change greatly when a pressure is applied. (Why are gases highly compressible?) • Gases exert pressure equally in all directions. (Why gases exert pressure equally in all directions?) Note: ALL the above is actually because gases have MUCH weaker (as compared to liquid ah) intermolecular forces, which can be hydrogen bonding, permanent dipole - permanent dipole or instantaneous dipole-induced dipole interactions!!! Making Thinking Visible Q: Do you know what is pressure? A: Pressure refers to the force that the particles exert on a surface per unit area. Hence, pressure measured is not due to the collision between the particles. Q: Since pressure refers to the force that the particles exert on the surface per unit area, does that mean that if there are more particles hitting per unit area, the pressure will be higher? A: Yes!!! That is why the measured pressure is directly proportion al to the amount of particles, in moles (n). In addition, pressure is directly proportional to temperature (T) because the higher the temperature, the greater the kinetic energy of the particles (KE average  T). The faster the particles move, the “harder” it hits against the surface. The particles also strike the surface more frequently. Lastly, the pressure is inversely proportional to the volume of the container (V) because the greater the volume, the further the particles are spread out, the lower the frequency of collisions as the particles will take a “longer time” to hit the surface again. Q: Sir, if p of a gas is affected by n, V and T, does it mean that V of a gas is being affected by p, n and T? A: Certainly Eunoians, we will later discuss the interconnectedness of p, V, n and T in the ideal gas equation, pV = nRT. S.I. Units and Common Units of Measurement for the Properties of Gases The physical behaviour of a sample of gas can generally be described completely by four variables:

T h e G a s e o u s S t a t e E u n o i a J u n i o r C o l l e g e 3 | P a g e (a) Pressure exerted by the gas, p S.I. Unit pascal (Pa) 1 Pa = 1 N m−2 Other commonly used units atmosphere (atm) 1 atm = 101 325 Pa bar (bar) 1 bar = 105 Pa torr (torr or mm Hg) 760 torr = 760 mm Hg = 1 atm = 101 325 Pa (b) Volume occupied by the gas, V S.I. Unit cubic metre (m3) Other commonly used units cubic decimetre (dm3) 1 dm3 = 10−3 m3 cubic centimetre (cm3) 1 cm3 = 10−6 m3 litre (L) 1 L = 1 dm3 = 10−3 m3 millilitre (mL) 1 mL = 1 cm3 = 10−6 m3 (c) Amount of the gas, n S.I. Unit mole (mol) (d) Temperature of the gas, T S.I. Unit kelvin (K) Other commonly used units degree Celsius (oC) T / ºC + 273 = T / K 2 The Gas Laws Some of the first quantitative investigations were conducted in the 17th – 18th century to study the behaviour of gases. The relationship between the different macroscopic properties of gases (volume, temperature and pressure) are expressed via the different gas laws. 2.1 Boyle’s Law The volume V of a fixed mass of gas at a constant temperature is inversely proportional to the pressure p of the gas. In other words, for a fixed mass (or amount) of gas at constant temperature, when its volume increases, its pressure decreases, and vice versa. Express the following quantities in their S.I. units. (a) 2 atm (b) 500 mL (c) 22.7 dm3 (d) 30 cm3 (e) 25 ºC (a) 2 x 105 Pa (b) 500x10−6 m3 (c) 22.7x10−3 m3 (d) 30x10−6 m3 (e) 298 K Self-Check 1A LO (c) state and use the general gas equation pV = nRT in calculations, including the determination of Mr 1 where is a constant V p pV k k  =

T h e G a s e o u s S t a t e E u n o i a J u n i o r C o l l e g e 4 | P a g e Hence, for a given mass of gas under two different sets of conditions at constant temperature, Graphically, Boyle’s Law can be represented in the following forms: At constant T and n, kV p= At constant T, pV = k Therefore, pV is independent of p or V. At constant T, 1Vk p =  1Resembles graphy x  = (Resembles y = c graph) (Resembles y = mx graph) 2.2 Charles’s Law The volume V of a fixed mass of gas