Chapter 6 Gases Physical Characteristics of Gases n

Chapter 6 - Gases

Physical Characteristics of Gases n 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.

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

Three Gas Laws n 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 - k. Pa

• 1 k. Pa = 1000 N/ 1 m 2 • Atmospheric pressure – pressure exerted by air particles colliding • SATP – 100 k. Pa at 25 °C • STP – 101. 3 k. Pa at 0 °C

Boyle’s Law n n As pressure on a gas increases the volume of the gas decreases proportionally as the temperature is held constant P 1 V 1 = P 2 V 2

Charles Law n n the volume of a gas increases proportionally as the temperature of the gas increases, if the pressure is held Constant V 1 = V 2 T 1 T 2

Boyle’s Law – inverse relationship Charles Law – direct relationship

Kelvin Temperature Scale n n 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 n n Formulas °C = K - 273 K= °C+273

Combined Gas Law n This is when all variables (T, P, and V) are changing ¨ P 1 V 1 T 1 = P 2 V 2 T 2

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

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

Molar Volume of Gases “new conversion ratio” n Avogadro says : § § T 1 = T 2 P 1 =P 2 V 1 = V 2 Then # particles of gas 1 = # particles of gas 2 1 mol = 6. 03 x 10 23 particles n Lets put these two ideas together…… n

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 n The two most standard temps and pressures are STP and SATP n

Molar Volume n When gases are at STP: ¨ 1 n mole of any gas = 22. 4 L/mol When gases are at SATP: ¨ 1 mole of any gas = 24. 8 L/mol

Ideal Gas Equation n Ideal Gas — is a hypothetical 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 a tiny attraction) n 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 n The smaller the molecules the closer the gas resembles an ideal gas n We assume ideal gases always. .

Equation n n PV = n. RT P = pressure (k. Pa) V = volume (L) n = moles (mol) R = universal gas constant (8. 31 k. Pa*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