Transcript Chapter 6 - Gases - The Lesson Locker

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 assumes the
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
increases
proportionally as the
temperature of the
gas increases, if the
pressure is held
Constant
V1 = V2
T1 T2
Boyle’s Law – inverse relationship
Charles Law – direct relationship
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
T1
= P2V2
T2
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
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
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……

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 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)

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.
.
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