Chapter 6 - Gases

1. Chapter 6 - Gases

2. Physical Characteristics of Gases • Although gases have different chemical properties, gases have remarkably similar physical properties. • Gases always fill their containers (recall solids and liquids). No definite shape and volume • Gases are highly compressible: Volume decreases as pressure increases Volume increases as pressure decreases • Gases diffuse (move spontaneously throughout any available space). • Temperature affects either the volume or the pressure of a gas, or both.

3. Definition of a Gas • Therefore a definition for gas is: a substance that fills and assumesthe shape of its container, diffuses rapidly, and mixes readily with other gases.

4. Three Gas Laws • Pressure • force of colliding particles per unit area • According to the KMT gases exert pressure due to the forces exerted by gas particles colliding with themselves and the sides of the container • SI unit for pressure is kilopascals - kPa

5. 1 kPa = 1000 N/ 1 m2 • Atmospheric pressure – pressure exerted by air particles colliding • SATP – 100 kPa at 25 °C • STP – 101.3 kPa at 0 °C

6. Boyle’s Law • As pressure on a gas increases the volume of the gas decreases proportionally as the temperature is held constant • P1V1 = P2V2

7. Charles Law • the volume of a gas increasesproportionally as the temperature of the gas increases, if the pressure is held Constant • V1 = V2T1 T2

8. Boyle’s Law – inverse relationshipCharles Law – direct relationship

9. Kelvin Temperature Scale • Temperature - the average kinetic energy of the particles making up a substance • Kelvin Temp Scale: based of absolute zero — all kinetic motion stops • 273°C= 0 K 0°C = 273 K 30°C =303 K -20°C = 253 K • Formulas °C = K - 273 K= °C+273

10. Combined Gas Law • This is when all variables (T,P, and V) are changing • P1V1 = P2V2 T1 T2

11. Avogadro’s Theory and Molar volume • The kinetic molecular theory is strongly supported by experimental evidence. • The K M theory explains why gases, unlike solids and liquids, are compressible. • The K M theory explains the concept of gas pressure. • The K M theory explains Boyle’s Law — Increase volume \ decrease pressure • The KM theory explains Charles’ Law Increase volume \ increase temperature

12. History Lesson • 1808 – Joseph Guy – Lussac • “Law of Combining Volumes” • When measuring at the same temp and pressure, volumes of gas reactants and products (in chemical reactions) are always in simple whole number ratios • 1810 – AmadeoAvogadro • “Avogadro’s Theory” • Equal volumes of gases at the same temp and pressure have equal number of molecules

13. Molar Volume of Gases“new conversion ratio” • Avogadro says : • T1 = T2 • P1 =P2 • V1 = V2 • Then # particles of gas 1 = # particles of gas 2 • 1 mol = 6.03 x 10 23 particles • Lets put these two ideas together……

14. Therefore for all gases at a specific temp and pressure there must be a certain volume that contains exactly 1 mole of particles - molar volume • The two most standard temps and pressures are STP and SATP

15. Molar Volume • When gases are at STP: • 1 mole of any gas = 22.4 L/mol • When gases are at SATP: • 1 mole of any gas = 24.8 L/mol

16. Ideal Gas Equation • Ideal Gas — is ahypothetical gas that obeys all the gas laws perfectly under all conditions. It is composed of particles with no attraction to each other. (Real gas particles do have atiny attraction) • The further apart the gas particles are, the faster they are moving the less attractive force they have and behave the most like ideal gases • The smaller the molecules the closer the gas resembles an ideal gas • We assume ideal gases always. .

17. Equation • PV = nRT • P= pressure (kPa) • V = volume (L) n = moles (mol) R = universal gas constant (8.31 kPa*L ) Mol * K T = temperature (K) • Sometimes the n must be converted to mass after the equation is completed. If this is necessary, use a conversion