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LOGO Lecture 7: Gaseous and Liquid compounds Course lecturer : Jasmin Šutković 17 th April 2014
LOGO
Lecture 7:Gaseous and Liquid compounds
Course lecturer : Jasmin Šutković17th April 2014
Contents International University of Sarajevo
1. Gaseous properties and compounds2. Pressure, Volume, Temperature and Amount
relationship3. The IDEAL GAS Law4. Mixture of gases 4. Kinetic energy in gases 5. The REAL gas behavior6. LIQUIDs and its kinetic energy7. Intermolecular forces in liquids8. Unique properties of LIQUIDS9. Vapor pressure10.CHANGE of states 11. Critical temperatures and pressures
The three common phases (or states) of matter are gas, liquid, and solid Gases
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 Liquids
a. Dense b. Incompressible c. Flow readily to adapt to the shape of the container d. Sum of the intermolecular forces are between those of gases
and solids
1. Gaseous Elements andCompounds
The state of a given substance depends strongly on conditions
Solids
a. Dense b. Rigid c. Incompressible d. Intermolecular forces are strong e. Molecules locked in place
Gaseous Elements andCompounds
PSE location of G,L and S elements
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
Gaseous Elements andCompounds
• Gaseous substances 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.
Gaseous Elements andCompounds
Gas Pressure
• At the macroscopic level, in order to describe the GAS we require four 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)
Units of Pressure
• 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.
Atmospheric Pressure
Athmospheric pressure example
This plastic bottle was sealed at approximately 4,300 m altitude, and was crushed by the increase in atmospheric pressure —at 2,700 m and 300 m— as it was brought down towards sea level.
Atmospheric pressure can be measured using a barometer, a closed, inverted tube filled with mercury.
The height of the mercury column is proportional to the atmospheric pressure, which is reported in units of millimeters of mercury (mmHg), also called torr.
One (1) bar is atmospheric pressure at sea level. 1atm = 760mmHg = 760torr = 1.01325x10^5 Pa = 101.325
kPa = 14.7 psi
Measurment of A.P
Pressure on a gas increases, the volume of the gas decreases ! (gas particles are forced closer together)
Pressure on a gas gas decreases, the gas volume increases ! (gas particles can now move farther apart)
Boyle carried out some experiments that determined the quantitative relationship between the pressure and volume of a gas.
2. The Relationship between Pressureand Volume
Relationship between the two quantities is described by the equation PV = constant (k).
Dividing both sides by P gives an equation that illustrates the inverse relationship between P and V:
• This relationship between pressure and volume is known as Boyle’s law which states that at constant temperature, the volume of a fixed amount of a gas is inversely proportional to its pressure.
The Relationship between Pressureand Volumen cont....
TEMPERATURE IS CONSTANT
Hot air rises and gases expand when heated.
Aleksandar Charles carried out experiments to quantify the relationship between the temperature and volume of a gas, he statet his law :
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:
Temperature and Volumen relationship
Example
Avogadro postulated that, at the same temperature and pressure, equal volumes of gases contain the same number of gaseous particles.
Avogadro’s law describes the relationship between volume and amount of gas: At constant temperature and pressure, the volume of a sample of gas is directly proportional to the number of moles of gas in the sample.
Stated mathematically:Where:
V is the volume of the gas(es). n is the amount of substance of the gas. k is a proportionality constant.
The Relationship between Amount of gases and Volumen
• 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.
In conclusion
3. Ideal Gas Law
A law relating pressure, temperature, volume, and the amount of an ideal gas.
Any set of relationships between a single quantity (such as 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)
Deriving the Ideal Gas Law
A particular set of conditions were chosen to use as a reference; 0ºC (273.15 K) and 1 atm pressure are referred to as standard temperature and pressure (STP).
The volume of 1 mol of an ideal gas at 0ºC and 1 atm pressure is 22.41L, called the standard molar volume of an ideal gas.
If the volume, temperature, and number of moles of each gas in a mixture is known, then the pressure exerted by each gas individually, which is its partial pressure, can be calculated.
Partial pressure is the pressure the gas would exert if it were the only one present (at the same temperature and volume).
This law is known as Dalton’s law of partial pressures and can be written mathematically as
Pt = P1 + P2 + P3 - - - + Pi
where Pt is the total pressure and the other terms are the partial pressures of the individual gases.
4. Mixture of gases - Partial Pressures (P)
Example
Partial Pressures cont...
For a mixture of two ideal gases, A and B, the expression for the total pressure can be written as
• More generally, for a mixture of i components, 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.
Mole Fractions of Gas Mixtures
The composition of a gas mixture can be described by the mole fractions of the gases present.
Mole fraction ( ) of any component of a mixture is the ratio of the number of moles of that component to the total number of moles of all the species present in the mixture (nt)
Mole fraction is a dimensionless quantity between 0 and 1.
If A = 1, then the sample is pure A, not a mixture.
If A = 0, then sample A is is not pure present in the mixture.
The sum of the mole fractions of all the components present must equal 1.
4. Kinetic molecular theory of gases
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 !!
Several postulates ( rules ) exists but 2 of them are essential for ideal gas behavior :
• Postulates 2 and 3 states that all gaseous particles behave identically, regardless of the chemical nature of their component molecules
At at a given temperature, all gases have the same average kinetic energy.
The average kinetic energy of the molecules of a gas is
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)
•
To sum up ...
• 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.
Result is a gas mixture with uniform composition - Graham’s law. The ratio of the diffusion rates of two gases is the square root of the
inverse ratio of their molar masses.
If r is 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.
Diffusion of Gas
• 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.
Effusion of Gas
The Behavior of Real Gases
• 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.
So…
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.
The Behavior of Real Gases
P,V and T Relationshipsin Real Gases
At low pressures, gaseous molecules are far apart
As pressure increases, intermolecular distances become smaller
Total volume occupied by gas is greater than the volume predicted by the ideal gas law, so the experimentally measured value of PV/nRT is greater than the value predicted by the ideal gas law.
Molecules of a real gas have small but measurable volumes.
V.W constants of some gases
Liquefaction of Gases
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.
6. LIQUIDS The Kinetic Molecular Description
Molecules in liquids are very close together, with essentially no empty space between them ( In solids they are more apart )
Molecules in liquids are in constant motion, and their kinetic energy (and hence their speed) depends on their temperature.
Properties of liquids can be explained using a modified version of the kinetic molecular theory.
The Kinetic Molecular Description of Liquids
Some characteristic properties of liquids are:
1. Density
– 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).
The Kinetic Molecular Description of Liquids cont...
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.
3. Compressibility
– 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.
The Kinetic Molecular Description of Liquids cont...
5. Fluidity – Liquids can flow, adjusting to the shape
of the container, because their molecules are free to move.
6. Diffusion – Molecules in liquids diffuse because they
are in constant motion.
The Kinetic Molecular Description of Liquids cont...
Fluidity and Diffusion
Intermolecular forces in liquids
intermolecular interactions, weaker than the intramolecular interactions that hold the atoms together in molecules and polyatomic ions.
Transitions between the solid, liquid, and gaseous states are due to changes in intermolecular interactions but do not affect intramolecular attractions.
Intermolecular forces are electrostatic; they arise from the interaction between positively and negatively charged species.
Intermolecular Forces cont...
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 Waals forces
Dipole-dipole interactions
– 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
Intermolecular Forces cont...
London dispersion forces
The London dispersion force is the weakest intermolecular force.
The London dispersion force is a temporary attractive force that results when the electrons in two adjacent atoms occupy positions that make the atoms form temporary dipoles.
Intermolecular Forces cont...
Hydrogen bonds
– 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.
Intermolecular Forces cont...
Example
There are three unique properties of liquids that depend intimately on the nature of intermolecular interactions
1. Surface tension 2. Capillary action 3. Viscosity
8. Unique Properties of Liquids
Surface Tension
Is the energy required to increase the surface area of a liquid by a specific amount!
Measured as energy per unit area, such as joules per square meter (J/m2)
The stronger the intermolecular forces, the higher the surface tension!
Surfactants are molecules such as soap and detergents that reduce the surface tension of polar liquids like water by disrupting the intermolecular attractions between adjacent molecules
Surface tension example...
(a) A paper clip can “float” onwater because of surfacetension.
(b) Surface tensionalso allows insects such asthis water strider to “walk onwater.”
Capillary Action
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
• 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 higher the viscosity.
• Liquids that have strong intermolecular forces have high viscosities.
9. Vapor Pressure
When the liquid is heated, its molecules obtain sufficient kinetic energy to overcome the forces holding them in the liquid and they escape into the gaseous phase.
The result of this phenomenon Vapor!
Boiling Points
As the temperature of a liquid increases, the vapor pressure of the liquid increases until it equals the external pressure, or the atmospheric pressure.
Bubbles of vapor begin to form throughout the liquid, and the liquid begins to boil.
Temperature at which a liquid boils at exactly 1 atm pressure is the normal boiling point of the liquid.
10. Changes of State
Changes of state are examples of phase changes, or phase transitions
Any of the three forms of matter (gas, liquid, solid) is converted to either of the other two
Six most common phase changes:
1. fusion (melting) solid → liquid2. freezing liquid → solid3. vaporization liquid → gas4. condensation gas → liquid5. sublimation solid → gas 6. deposition gas → solid
Energy Changes that results in phase Changes
Phase changes are always accompanied by a change in the energy of the system – we have to types of transition phases ;
1. Any transition from a more-ordered to a less-ordered state (solid to liquid, liquid to gas, or solid to gas) requires an input of energy — it is ENDOTHERMIC.
2. Any transition from a less-ordered to a more-ordered state (liquid to solid, gas to liquid, or gas to solid) releases energy — it is EXOTHERMIC.
The conversion of a solid to a liquid is called melting or fusion — energy required to melt 1 mol of the substance is its enthalpy of fusion, ΔHfus.
Energy change required to vaporize 1 mol of a substance is the enthalpy of vaporization, ΔHvap.
Energy Changes that results in phase Changes
Direct conversion of a solid to a gas, without an intervening liquid phase, is called sublimation.
Amount of energy required to sublime 1 mol of a pure solid is the enthalpy of sublimation, ΔHsub.
Enthalpy of sublimation of a substance is the sum of its enthalpies of fusion and vaporization (an application of Hess’s law):
ΔHsub = ΔHfus + ΔHvap.
Fusion, vaporization, and sublimation are endothermic processes; they occur only with the absorption of heat.
11. Critical Temperature and Pressure
Critical temperature (Tc)
The critical 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.
Critical pressure (Pc) The critical pressure of a substance is the pressure required to liquefy a gas at
its critical temperature
Critical point — combination of critical temperature and critical pressure
Readings for this lecture
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.
Reminder for Quiz 2 and Homework 3
Submit by 24th April - Homework 3
6th May Quiz 2 ( Lecture 7, 8, 9 and 10)
Homework 3 question:
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?
