Chapter 6 - Gases

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Chapter 6 - Gases. 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

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Chapter 6 - Gases

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.
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.
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
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
Boyle’s Law
• As pressure on a gas increases the volume of the gas decreases proportionally as the temperature is held constant
• P1V1 = P2V2
Charles Law
• the volume of a gas increasesproportionally as the temperature of the gas increases, if the pressure is held Constant
• V1 = V2T1 T2
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
Combined Gas Law
• This is when all variables (T,P, and V) are changing
• P1V1 = P2V2

T1 T2

• 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
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
• Equal volumes of gases at the same temp and pressure have equal number of molecules
Molar Volume of Gases“new conversion ratio”
• 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……
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
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
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.

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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