gas and condensed matter
TRANSCRIPT
GAS AND CONDENSED MATTER
Bonding model for covalent molecular substances
Bonding for covalent molecular substances falls into two categories
1. The strong forces of attraction which holds atoms together within molecules
2. The weak forces of attraction between molecules
Forces between molecules (intermolecular forces) we will learn about the forces between
molecules or compounds are called intermolecular forces
Inter means between or among Internet, interstate, international What would Interstellar travel be?
Intermolecular Forces Forces that occur between
molecules.
Intramolecular forces What would intramolecular forces be? Forces within molecules e.g covalent, metallic or
ionic intra means within Intrastate, intranet, intracellular Most of the intermolecular forces we look at occur
between covalently bonded molecules or covalent molecular substances
Intramolecular bonds are stronger than intermolecular forces.
Overview All matter is held together by force. The force between atoms within a
molecule is a chemical or intramolecular force.
The force between molecules are a physical or intermolecular force.
These physical forces are what we overcome when a chemical changes its state (e.g. gas liquid).
What causes intermolecular forces? Molecules are made up of charged particles: positive
nuclei and negative electrons. When one molecule approaches another there is a
multitude of forces between the particles in the two molecules.
Each electron in one molecule is attracted to the nuclei in the other molecule but also repelled by the electrons
in the other molecule. The same applies for nuclei
Types of Intermolecular forces The three main types of intermolecular
forces are:1. Dipole-dipole attraction occur only btw
polar molecules2. H bonding – only with Hydrogen and
Oxygen, Fluorine and Nitrogen)3. Dispersion forces (London Dispersion
Forces)
Intermolecular Forces Forces that occur between
molecules. Dipole–dipole forces
Hydrogen bonding London dispersion forces
Dipole–Dipole Attraction
Dipole-Dipole ForcesDipole moment – molecules with polar
bonds often behave in an electric field as if they had a center of positive charge and a center of negative charge.
Molecules with dipole moments can attract each other electrostatically. They line up so that the positive and negative ends are close to each other.
Only about 1% as strong as covalent or ionic bonds.
Hydrogen BondingStrong dipole-dipole forces.Hydrogen is bound to a highly electronegative atom – nitrogen, oxygen, or fluorine.
Hydrogen Bonding in Water Blue dotted
lines are the intermolecular forces between the water molecules.
Hydrogen BondingAffects physical properties
Boiling point
London Dispersion Forces Instantaneous dipole that occurs
accidentally in a given atom induces a similar dipole in a neighboring atom.
Significant in large atoms/molecules. Occurs in all molecules, including
nonpolar ones.
London Dispersion ForcesNonpolar Molecules
London Dispersion Forces
Become stronger as the sizes of atoms or molecules increase.
Melting and Boiling Points In general, the stronger the
intermolecular forces, the higher the melting and boiling points.
Strength of Intermolecular Interactions
Hydrogen Bonding
↑ Dipole – Dipole
↑ London Dispersion Forces
The kinetic theory of matter is based on the following postulates:
1. Matter is composed of small particles called molecules
2. The particles are in constant random motionThey possess kinetic energy due to their motion
3. There are repulsive and attractive forces between particles.
They posses potential energy due to these forces
4. Average particle speed increases with temperature5. No energy is lost when the particles collide, called
elastic collision
Kinetic Molecular Theory
The kinetic energy of a particle is given by the equation:
Where:
m = particle mass in kg
v = particle velocity in m/s
KE = kg-m2/s2 = j (joule)According to postulate 4 of our kinetic theory particle velocity increases with temperature. This means as temperature increases then kinetic energy increases.
22
1KE mv
Kinetic Molecular Theory
Potential energy is the sum of the attractive and repulsive forces between particles.
Examples of these types of forces are the gravitational attractive forces between objects and the repulsive forces between the same poles of magnets.
Alternatively we can say forces between particles may be either cohesive or disruptive.
Potential Energy
Cohesive forces include dipole-dipole interactions, dispersion forces, attraction between oppositely charged ions.Cohesive forces are largely temperature independent.
e.g. magnets and gravity function the same way at different temperature.
Interparticle Forces
Disruptive forces are those forces that make particles move away from each other.
These forces result predominately from the particle
motion. Disruptive forces increase with temperature in agreement with postulate 4. We can conclude that as we increase the temperature particles will become further apart from each other.
Interparticle Forces
Tutorial 1 State and describe briefly three (3) main types of
intermolecular forces. State five (5) assumption in the kinetic molecular
theory
A Gas
Has neither a definite volume nor shape.
Uniformly fills any container.
Mixes completely with any other gas
Exerts pressure on its surroundings.
Composition of Earth’s AtmosphereCompound %(Volume) Mole Fractiona
Nitrogen 78.08 0.7808
Oxygen 20.95 0.2095
Argon 0.934 0.00934
Carbon dioxide 0.033 0.00033
Methane 2 x 10-4 2 x 10-6
Hydrogen 5 x 10-5 5 x 10-7
a. mole fraction = mol component/total mol in mixture.
Earth-like Atmosphere
A mercury barometer
The column height is proportional to the atmospheric pressure.
Atmospheric pressure results from the mass of the atmosphere and gravitational forces.
The pressure is the force per unit area.
P = F/A
1 atm = 760 mmHg
1 atm = 1.01325 E5 Pa
1 mmHg = 1 torr
Units for Expressing Pressure
Unit Value
Atmosphere 1 atm
Pascal (Pa) 1 atm = 1.01325 x 105 Pa
Kilopascal (kPa) 1 atm = 101.325 kPa
mmHg 1 atm = 760 mmHg
Torr 1 atm = 760 torr
Bar 1 atm = 1.01325 bar
mbar 1 atm = 1013.25 mbar
psi 1 atm = 14.7 psi
Pressure is equal to force/unit area SI units = Newton/meter2 = 1 Pascal (Pa) 1 standard atmosphere = 101,325 Pa
(100,000 Pa = 1 bar) 1 standard atmosphere = 1 atm =
760 mm Hg = 760 torr = 1013.25 hPa= 14.695 psi
Meteorologists often report pressure inmillibar; 1 mbar =0.001bar =0.1 kPa = 1hPa
Variables Affecting Gases Pressure (P)
Volume (V)
Temperature (T)
Number of Moles (n)
Elevation and Atmospheric Pressure
ManometerManometers are used to measure gas pressure in closed systems. For instance in a reaction vessel.
The Gas Laws and Absolute TemperatureThe relationship between the volume, pressure, temperature, and mass of a gas is called an equation of state.
We will deal here with gases that are not too dense.
Boyle’s Law: the volume of a given amount of gas is inversely proportional to the pressure as long as the temperature is constant.
Boyle Law…pressure is inversely proportional to volume (at constant T and moles, n).
The Gas Laws and Absolute Temperature
The volume is linearly proportional to the temperature, as long as the temperature is somewhat above the condensation point and the pressure is constant:
Extrapolating, the volume becomes zero at −273.15°C; this temperature is called absolute zero.
Avogadro’s Law
For a gas at constant temperature and pressure, the volume is directly proportional to the number of moles of gas (at low pressures).
V n
V1 = V2
n1 n2
The Gas Laws and Absolute TemperatureThe concept of absolute zero allows us to define a third temperature scale – the absolute, or Kelvin, scale.
This scale starts with 0 K at absolute zero, but otherwise is the same as the Celsius scale.
Therefore, the freezing point of water is 273.15 K, and the boiling point is 373.15 K.
Finally, when the volume is constant, the pressure is directly proportional to the temperature:
Combined Gas Law
Combining the gas laws the relationship P T(n/V) can be obtained.
If n (number of moles) is held constant, then PV/T = constant.
P1V1
T1
= P2V2
T2
Ideal Gas LawPV = nRT
R = universal gas constant = 0.08206 L atm K-1 mol-1
P = pressure in atm V = volume in liters n = moles T = temperature in Kelvin
Standard Temperature and Pressure (for gases)
“STP” P = 1 atmosphere T = 0C The molar volume of an ideal gas is 22.42 liters at
STP (put 1 mole, 1 atm, R, and 273 K in the ideal gas law and calculate V)
Note STP is different for other phases, e.g. solutions or enthalpies of formation.
The Ideal Gas Law
A mole (mol) is defined as the number of grams of a substance that is numerically equal to the molecular mass of the substance:
1 mol H2 has a mass of 2 g
1 mol Ne has a mass of 20 g
1 mol CO2 has a mass of 44 g
The number of moles in a certain mass of material:
• We can combine these into a general gas law:
The Ideal Gas The Ideal Gas EquationEquation
), (constant 1
TnP
V
), (constant PnTV
),(constant TPnV
• Boyle’s Law:
• Charles’s Law:
• Avogadro’s Law:
PnT
V
• R = gas constant, then
• The ideal gas equation is:
• R = 0.08206 L·atm/mol·K = 8.3145 J/mol·K• J = kPa·L = kPa·dm3 = Pa·m3
• Real Gases behave ideally at low P and high T.
The Ideal Gas The Ideal Gas EquationEquation
P
nTRV
nRTPV
Ideal Gas Law in Terms of Molecules: Avogadro’s Number
Since the gas constant is universal, the number of molecules in one mole is the same for all gases. That number is called Avogadro’s number:
The number of molecules in a gas is the number of moles times Avogadro’s number:
Ideal Gas Law in Terms of Molecules: Avogadro’s Number
Therefore we can write:
where k is called Boltzmann’s constant.
(13-4)
The Ideal Gas The Ideal Gas EquationEquation
Calculate the pressure exerted by 84.0 g of ammonia, NH3, in a 5.00 L container at 200. oC using the ideal gas law.
PV = nRTP = nRT/V n = 84.0g * 1mol/17 g T = 200 + 273P = (4.94mol)(0.08206 L atm mol-1 K-1)(473K)
(5 L)P = 38.3 atm
Tutorial 2. The pressure on a sample of an ideal gas was
increased from 715 torr to 3.55 atm at constant temperature. If the initial volume of the gas was 485. mL, what would be the final volume?
A 7.9 L sample of gas was cooled from 79°C to a temperature at which the volume of the gas was 4.3 L. Assuming the pressure remains constant, calculate the final temperature.
Calculate the pressure in atmospheres and Calculate the pressure in atmospheres and pascals of a 1.2 mol sample of methane gas in a pascals of a 1.2 mol sample of methane gas in a 3.3 L container at 25°C.3.3 L container at 25°C.
Real Gases: Deviations from Ideality
Real gases behave ideally at ordinary temperatures and pressures.
At low temperatures and high pressures real gases do not behave ideally.
The reasons for the deviations from ideality are:
1. The molecules are very close to one another, thus their volume is important.
2. The molecular interactions also become important. J. van der Waals, 1837-1923, J. van der Waals, 1837-1923,
Professor of Physics, Professor of Physics, Amsterdam. Nobel Prize 1910.Amsterdam. Nobel Prize 1910.
Real Gases:Deviations from Ideality
van der Waals’ equation accounts for the behavior of real gases at low temperatures and high pressures.
P + n a
VV nb nRT
2
2
• The van der Waals constants a and b take into account two things:
1. a accounts for intermolecular attractiona. For nonpolar gases the attractive forces are London Forcesb. For polar gases the attractive forces are dipole-dipole attractions
or hydrogen bonds.
2. b accounts for volume of gas molecules
At large volumes a and b are relatively small and van der Waal’s equation reduces to ideal gas law at high temperatures and low pressures.
The van der Waals Equation
• General form of the van der Waals equation:
Real Gases: Deviations from Ideal Real Gases: Deviations from Ideal BehaviorBehavior
2
2
V
annbV
nRTP
nRTnbVV
anP
2
2
Corrects for molecular volume
Corrects for molecular attraction
Example
Condensed matter : the three states of matter.
Some Characteristics of Gases, Liquids and Solids and the Microscopic Explanation for the Behavior
gas liquid solid
assumes the shape and volume of its container particles can move past
one another
assumes the shape of the part of the container
which it occupies particles can move/slide
past one another
retains a fixed volume and shape
rigid - particles locked into place
compressible lots of free space between particles
not easily compressible little free space between
particles
not easily compressible little free space between
particles
flows easily particles can move past
one another
flows easily particles can move/slide
past one another
does not flow easily rigid - particles cannot move/slide past one
another
Clearly, a theory used to describe the condensed states of matter must include an attraction between the particles in the substanceCondensed States
of Matter:
Liquids
Solids
.
Kinetic Theory Description of the Liquid State. Like gases, the
condensed states of matter can consist of atoms, ions, or molecules.
What separates the three states of matter is the proximity of the particles in the substance.
For the condensed states of matter the particles are close enough to interact.
Phase Changes
Triple Point Diagram of Water Regions: Each region
corresponds to one phase which is stable for any combination of P and T within its region
Lines Between Region: Lines separating the regions representing phase-transition curves
Triple Point: The triple point represents the P and T at which all 3 phases coexist in equilibrium
Critical Point: At the critical point the vapor pressure cannot be condensed to liquid no matter what pressure is applied.
Tutorial 3 Van der Waals, realized that two of the
assumptions mentioned above were questionable. He then developed the Van der Waals equation of state which predicted the formation of liquid phase. Write the equation and state two corrections that he made.
(a) Calculate the pressure exerted by 1.00 mol of CO2 in a 1.00 L vessel at 300 K, assuming that the gas behaves ideally. (b) Repeat the calculation by using the van der Waals equation.
Sketch and label the liquid region, gas region solid region and triple point in water phase diagram