introduction and gases. physics - study of the properties of matter that are shared by all...
TRANSCRIPT
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Chemistry 231 Introduction and Gases
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Physics - study of the properties of matter that are shared by all substances
Chemistry - the study of the properties of the substances that make up the universe and the changes that these substances undergo
Physical Chemistry - the best of both worlds!
Physical Chemistry
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Thermodynamics – the study of energy and its transformations
Thermochemical changes – energy changes associated with chemical reactions
Thermodynamics and Thermochemistry
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Interested in the numerical values of the state variables (defined later) that quantify the systems at that point in time.
Systems can be either• macroscopic• microscopic
Studying Systems
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Described by variables such as• temperature (T)• pressure (P)• volume (V)• energy (U) • enthalpy (H)• Gibbs energy (G)
State of a System
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State Variables • system quantity whose values are fixed at constant
temperature, pressure, composition State Function
• a system property whose values depends only on the initial and final states of the system.
Path Functions • system quantity whose value is dependent on the
manner in which the transformation is carried out.
State and Path Functions
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Examples of state functions• H• G• V• T
Examples of path functions• work (w)• heat (q)
State and Path Functions (Continued)
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Metastable - the progress towards the equilibrium state is slow
Equilibrium state - state of the system is invariant with time
Equilibrium vs. Metastable
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Reversible transformation - the direction of the transformation can be reversed at any time by some infinitesimal change in the surroundings
Irreversible transformation - the system does not attain equilibrium at each step of the process
Reversible and Irreversible
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Gas - a substance that is characterised by widely separated molecules in rapid motion
Mixtures of gases are uniform. Gases will expand to fill containers.
The Definition of a Gas
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Common gases include - O2 and N2, the major components of "air"
Other common gases - F2, Cl2, H2, He, and N2O (laughing gas)
Examples of Gaseous Substances
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The pressure of a gas is best defined as the forces exerted by gas on the walls of the container
Define P = force/area The SI unit of pressure is the Pascal 1 Pa = N/m2 = (kg m/s2)/m2
The Definition of Pressure
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How do we measure gas pressure? We use an instrument called the
barometer - invented by Torricelli Gas pressure conversion factors
• 1 atm = 760 mm Hg = 760 torr• 1 atm = 101.325 kPa = 1.01325 bar• 1 bar = 1 x 105 Pa (exactly)
The Measurement of Pressure
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Experiments with a wide variety of gases revealed that four variables were sufficient to fully describe the state of a gas • Pressure (P)• Volume (V) • Temperature (T)• The amount of the gas in moles (n)
The Gas Laws
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The gas volume/pressure relationship The volume occupied by the gas is
inversely proportional to the pressure V 1/P
• note temperature and the amount of the gas are fixed
Boyle's Law
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Boyle's Law
V
1/P
V
P
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Defines the gas volume/temperature relationship.
V T (constant pressure and amount of gas)
Note T represents the temperature on the absolute (Kelvin) temperature scale
Charles and Gay-Lussac's Law
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Charles and Gay-Lussac's Law
V
t / CAbsolute Zero
(-273C = 0 K)
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Lord Kelvin – all temperature/volume plots intercepted the tc axis at -273.15°C).
Kelvin termed this absolute 0 – the temperature where the volume of an ideal gas is 0 and all thermal motion ceases!
The Kelvin temperature scale
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T (K) = [ tc (°C) + 273.15°C] K/°C• Freezing point of water: tc = 0 °C; T = 273.15
K
• Boiling point of water: tc = 100 °C; T = 373.15 K
• Room temperature: tc = 25 °C; T = 298 K
• NOTE tc = °C; T (K) = K NO DEGREE SIGN
The Temperatures Scales
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The pressure/temperature relationship For a given quantity of gas at a fixed
volume, P T, i.e., if we heat a gas cylinder, P increases!
Amonton’s Law
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The volume of a gas at constant T and P is directly proportional to the number of moles of gas
V n => n = number of moles of gas
Avogadro’s Law
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We have four relationships• V 1/P; Boyle’s law • V T; Charles’ and Gay-Lussac's law • V n; Avogadro’s law• P T; Amonton’s law
The Ideal Gas Equation of State
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Combine these relationships into a single fundamental equation of state - the ideal gas equation of state
The Ideal Gas Law
mole Katm
08206.0
314.8L
moleKJ
R
nRTPV
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An ideal gas is a gas that obeys totally the ideal gas law over its entire P-V-T range
Ideal gases – molecules have negligible intermolecular attractive forces and they occupy a negligible volume compared with the container volume
The Definition of an Ideal Gas
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Define: STP (Standard Temperature and Pressure)• Temperature - 0.00 °C = 273.15 K• Pressure - 1.000 atm• The volume occupied by 1.000 mole of an
ideal gas at STP is 22.41 L!
Standard Temperature and Pressure
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Define: SATP (Standard Ambient Temperature and Pressure)• Temperature - 25.00 °C = 273.15 K• Pressure - 1.000 bar (105 Pa)• The volume occupied by 1.000 mole of an
ideal gas at SATP is 24.78 L!
Standard Ambient Temperature and Pressure
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Let's consider two ideal gases (gas 1 and gas 2) in a container of volume V.
Partial Pressures
1
2
2 2
22 1
1
1
1
11
2
2
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The pressure exerted by gas #1 • P1 = n1 RT / V
The pressure exerted by gas #2 • P2= n2 RT / V
The total pressure of the gases • pT = nT RT / V
nT represents the total number of moles of gas present in the mixture
Partial Pressures
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P1 and P2 are the partial pressures of gas 1 and gas 2, respectively. • PT = P1 + P2 = nT (RT/V)
• PT = P1 + P2 + P3 = j PJ
• note Pj is known as the partial pressure of gas j
Partial Pressures (continued)
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Gaseous mixtures - gases exert the same pressure as if they were alone and occupied the same volume.
The partial pressure of each gas, Pi, is related to the total pressure by Pi = Xi PT
Xj is the mole fraction of gas i.• Xj= nj / nT
Dalton's Law of Partial Pressure
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In the limit of low pressures
Ideal Gas Temperature Scale
0lim p
PVT
nR
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The Isothermal Compressibility
The Isothermal Compressibility
TT P
VV
1
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The coefficient of thermal expansion
Coefficient of Thermal Expansion
PTV
V
1