Lecture 7: Gaseous and Liquid compounds. Course lecturer : Jasmin Šutković 17 th April 201 4 . Contents. International University of Sarajevo . 1. Gaseous properties and compounds Pressure, Volume, Temperature and Amount relationship The IDEAL GAS Law Mixture of gases
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Course lecturer :
International University of Sarajevo
1. Gaseous properties and compounds
Pressure, Volume, Temperature and Amount relationship
The IDEAL GAS Law
Mixture of gases
Kinetic energy in gases
The REAL gas behavior
LIQUIDs and its kinetic energy
Intermolecular forces in liquids
Unique properties of LIQUIDS
CHANGE of states
Critical temperatures and pressures
The three common phases (or states) of matter are gas, liquid, and solid
a. Have the lowest density of the three states of matter
b. Are highly compressible
c. Completely fill any container in which they are placed
d. Their intermolecular forces are weak
e. Molecules are constantly moving independently of the other molecules present
c. Flow readily to adapt to the shape of the container
d. Sum of the intermolecular forces are between those of gases
The state of a given substance depends strongly on conditions
d. Intermolecular forces are strong
e. Molecules locked in place
The figure of the periodic table shows the locations in the periodic table of those elements that are commonly found in the gaseous, liquid, and solid states
Many of the elements and compounds are typically found as gases
• Gaseous substances include:
1. Many binary hydrides, such as the hydrogen halides
2. Hydrides of the Group-15 elements N, P, and As
3. Hydrides of the Group-14 elements C, Si, and Ge
4. Many of the simple covalent oxides of the nonmetals such as
CO, CO2, NO, NO2, SO2, SO3, and ClO2
5. Many low-molecular-mass organic compounds
6. Most of the commonly used refrigerants
• Gaseoussubstances contain covalent or polar covalent bonds and are nonpolar or slightly polar molecules.
• Boiling points of polar compounds are higher than those of nonpolar compounds of similar molecular mass.
• At the macroscopic level, in order to describe the GAS we requirefour quantities:
1. Temperature (expressed in K or Celsius (C))
2. Volume (expressed in liters (L) )
3. Amount (expressed in moles (n))
4. Pressure (given in atmospheres (atm))
• These variables are not independent — if the values of any three of these quantities are known, the fourth can be calculated.
Pressure is any force exerted from anything on any surface with which it comes in contact.
Units of pressure are derived from the units used to measure force and area
For scientific measurements, the S unit for pressure, is the newton per square meter, N/m2, called the paschal (Pa):
1 pascal (Pa) = 1 newton/meter2 (N/m2)
• Every point on Earth’s surface experiences a net pressure called atmospheric pressure.
• Pressure is due to the weight of air above that surface in the atmosphere of Earth (or that of another planet)
Air has weight.
Imagine a column ( some space) of air that is 2.5 square meter in cross-section that extends vertically above your head until the air runs out in space. At sea level that column of air weighs around 7kg . If you climb mount Everest there is much less air above you in that column, so the weight of air above you will be much less than at sea level.
This plastic bottle was sealed at approximately 4,300 maltitude, and was crushed by the increase in atmospheric pressure —at 2,700 m and 300 m— as it was brought down towards sea level.
One (1) bar is atmospheric pressure at sea level.
1atm = 760mmHg = 760torr = 1.01325x10^5 Pa = 101.325 kPa = 14.7 psi
• This relationship between pressure and volume is known as Boyle’slaw which states that at constant temperature, the volume of a fixed amount of a gas is inversely proportional to its pressure.
TEMPERATURE IS CONSTANT
At constant pressure, the volume of a given mass of an ideal gas increases or decreases by the same factor as its temperature on the absolute temperature scale.
where V is the volume of the gas; and T is the absolute temperature. The law can also be usefully expressed as follows:
Amount of gases and Volumen
V is the volume of the gas(es).
n is the amount of substance of the gas.
k is a proportionality constant.
• The relationships between the volume of a gas and its pressure, temperature, and amount are summarized in the figure below.
• The volume increases with increasing temperature or amount but decreases with increasing pressure.
Any set of relationships between a single quantity (suchas V) and several other variables (P, T, n) can be combined into a single expression that describes all the relationships simultaneously.
R = 8.3145 J/mol·K (Joule /mol·Kelvin)
Pt= P1+ P2+ P3- - - + Pi
where Pt is the total pressure and the other terms are the partial pressures of the individual gases.
•More generally, for a mixture oficomponents, the total pressure is given by
The above equation makes it clear that, at constant temperature and volume, the pressure exerted by a gas depends on only the total number of moles of gas present, whether the gas is a single chemical species or a mixture of gaseous species.
The sum of the mole fractions of all the components present must equal 1.
The kinetic molecular theory of gases explains the laws that describe the behavior of gases and it was developed during the nineteenth century by Boltzmann, Clausius, and Maxwell.
Kinetic molecular theory of gases provides a molecular explanation for the observations that led to the development of the ideal gas law !!
• Postulates 2 and 3 states that all gaseous particles behave identically, regardless of the chemical nature of their component molecules
where 2 is the average of the squares of the speeds of the particles and m is the mass of the object.
• The square root of 2 is the root mean square (rms)speed (rms)
•Diffusion is the gradual mixing of gases due to the motion of their component particles even in the absence of mechanical agitation (movement) such as stirring.
The ratio of the diffusion rates of two gases is the square root of the inverse ratio of their molar masses.
If ris the diffusion rate and M is the molar mass, then
r1/r2 = M2/M1
If M1 M2, then gas #1 will diffuse more rapidly than gas #2.
•Effusion is the escape of a gas through a small (usually microscopic) opening into an evacuated space.
• Rates of effusion of gases are inversely proportional to the square root of their molar masses.
• Heavy molecules effuse through a porous material more slowly than light molecules.
•Postulates of the kinetic molecular theory of gases ignore both the volume occupied by the molecules of a gas and all interactions between molecules, whether attractive or repulsive.
• In reality, all gases have nonzero molecular volumes and the molecules of real gases interact with one another in ways that depend on the structure of the molecules and differ for each gaseous substance.
• Real gases behave differently from ideal gases at high pressures and low temperatures.
A ) In an ideal gas the interactions between molecules are not accounted
B ) in an ideal gas the actual volume taken up by the molecules of gas is not taken into account.
Molecules of an ideal gas are assumed to have zero volume; volume available to them for motion is the same as the volume of the container.
Molecules of a real gas have small but measurable volumes.
Liquefaction of gases is the condensation of gases into a liquid form.
• Both the theory and the ideal gas law predict that gases compressed to very high pressures and cooled to very low temperatures should still behave like gases.
However, as gases are compressed and cooled, they condense to form liquids.
Some characteristic properties of liquids are:
– Molecules of a liquid are packed relatively close together.
– As a result, liquids are much denser than gases
– Densities of liquids measured in units of grams per cubic centimeterr(g/cm3) or grams per milliliter (g/mL).
2. Molecular order
– Liquids exhibit short-range order because strong intermolecular attractive forces cause the molecules to pack together tightly.
– Because of the high kinetic energy of the molecules, they move rapidly with respect to one another.
– Arrangement of the molecules in a liquid is not completely random.
– Molecules in liquids are ordered because of strong intermolecular attractive forces.
– Liquids cannot be readily compressed because they have so little empty space between the component molecules.
4. Thermal expansion
– Intermolecular forces in liquids are strong enough to keep them from expanding significantly when heated.
– Volumes of liquids are somewhat fixed.
– Liquids can flow, adjusting to the shape of the container, because their molecules are free to move.
– Molecules in liquids diffuse because they are in constant motion.
There are three major types of intermolecular interactions:
1. Dipole-dipole interactions
2. London dispersion forces
3. Hydrogen bonds
The first two are described collectively as van der Waalsforces
– There are two types of dipole-dipole interactions:
1. Attractive — molecular orientations in which the positive end of one dipole is near the negative end of another (and vice versa)
2. Repulsive — molecular orientations that juxtapose the positive or negative ends of the dipoles on adjacent molecules
–The attractive intermolecular interactions are more stable than the repulsive intermolecular interactions
– Molecules that contain hydrogen atoms bonded to electronegative atoms such as O, N, F, and to a lesser extent Cl and S, tend to exhibit strong intermolecular interactions and have high boiling points.
– The large difference in electronegativity results in a large partial positive charge on hydrogen and a corresponding large partial negative charge on the O, N, or F atom.
There are three unique properties of liquids that depend intimately on the nature of intermolecular interactions
1. Surface tension
2. Capillary action
(a) A paper clip can “ﬂoat” on
water because of surface
(b) Surface tension
also allows insects such as
this water strider to “walk on
Is the ability of a liquid to flow in narrow spaces without the assistance of, and in opposition to, external forces like gravity.
The height to which the liquid rises depends on the diameter of the tube (the smaller the diameter, the higher the liquid rises) and the temperature of the liquid, but not on the angle at which the tube enters the liquid.
• Viscosity () is the resistance of a liquid to flow.
– The higher the viscosity, the slower the liquid flows through the tube and the slower the steel balls fall.
• Viscosity is expressed in units of the poise (mPa•s); the higher the number, the higherthe viscosity.
• Liquids that have strong intermolecular forces have high viscosities.
The result of this phenomenon Vapor!
1. fusion (melting) solid → liquid
2. freezing liquid → solid
3. vaporization liquid → gas
4. condensation gas → liquid
5. sublimation solid → gas
6. deposition gas → solid
ΔHsub= ΔHfus+ ΔHvap.
Fusion, vaporization, and sublimation are endothermic processes; they occur only with the absorption of heat.
Thecritical temperature of a substance is the temperature at and above which vapor of the substance cannot be liquefied, no matter how much pressure is applied.
The critical pressure of a substance is the pressure required to liquefy a gas at its critical temperature
Book chapter 10 and 11 – pages : 655-720, 731-782
Follow this lecture if you decide to read the book , what is mentioned in the powerpoint lectures ,this sould be also read in the book for more details.
Submit by 24th April - Homework 3
( Lecture 7, 8, 9 and 10)
1. why can a gas be compressed into a smaller volume? 2. why can you not squeeze a liquid into a smaller volume? 3. why can you not stir a solid? 4. what makes a gas exert pressure? 5. why does a gas completely fill its container? 6. why does a solid expand when it is heated?