Content:Page 2 6

Lesson 1: GasPage 7 8

Lesson 1: Exercise

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Lesson 2: Liquids and solids Page 14 15

Lesson 2: Exercise

Page 16 20

Lesson 3: Phase diagramPage 21 22

Lesson 3: Exercise

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Lesson 4: Solution and colligative properties Page 33 34

Lesson 4: ExercisePage 35

Periodic Table

Lesson 1: Gas1.1 Pressure Gas is a _______ fluid. It exerts _____ on a certain ____ . Pressure is measured with _______ which contains a mercury column. The unit _____ and _____ are used interchangeably. For atmospheric pressure, ___ atm = ___ mmHg = ___ torr = ___ PaExample 1 Calculation

Example 2 Calculation of the unit of pressure

1.2 Gas Laws There are three laws governing the behavior of gases, namely ______, _____, _____. The law saying that the pressure exerted by a gas is inversely proportional to the volume the gas occupies is _____.

The law saying that at constant pressure, the volume of a gas is directly proportional to the temperature (temperature unit in _____ ) of the gas is _____.

The law saying that for a gas at constant temperature and pressure the volume is directly proportional to the number of moles of gas is ______.

Example 1. Boyles Law

Example 2. Boyles Law

Example 3. Charless Law

Example 4. Avogadros Law

1.3 Ideal Gas Law The ideal gas law is a combination of the three laws. It relates to _____, _____, _____ and the ________ of a gas. The formula of the ideal gas law is _____________.

This relationship assumes that the gas behaves ______.

Standard temperature and pressure means ______ and ____

Example 1. Apply the ideal gas law

Example 2. Practice of ideal gas law

Example 3. Practice of ideal gas law

Example 4 Practice of ideal gas law

Example 5 Practice of ideal gas law

1.4 Daltons Law of partial pressure Daltons Law of partial pressure states that for a mixture of gases in a container, the total pressure is the _____ of pressures that each gas would exert if it were alone.

Example 1.

1.5 Kinetic molecular theory of gas For the KMT, Volume of gas molecules can be assumed to be ______.

Gas molecules move in random motion with _____ collisions.

Gas molecule do not exert _________ ____ on each other.

Two gases at the same temperature have the same ______ energy.

The temperature is a measure of the _______ kinetic energy of a gas, with the formula,

The expression dealing with the ______ ______ of gas particles is root mean square velocity, Example 1. Root mean velocity

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1.6 Grahams Law of Effusion ______ relates to the mixing of gases.

______ relates to the passage of a gas through an orifice into an evacuated chamber.

The relative rates of effusion of two gases are ________ proportional to the square roots of the masses of the gas particles.

The higher the molar mass, the slower the rate of effusion.

Example 1.

1.7 Real gases Real gases do not behave ideally at ____ pressure and ____ temperature. Ideal gases are the particles having no volume. But in real life, they do. Therefore, the volume available for gas to move around in a container should be (V-nb).

Ideal gases are the particles having no interaction with each other. But in real life, they have weak intermolecular forces and thus, Pobserved = P a(n/v)2.

By combing the two considerations, we have Van der Waals Equation for real gases

Example 1.

Lesson 1: Exercise

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Lesson 2: 2.1 Intermolecular force

___________ force mean forces within a molecule.

___________ force mean forces between molecules.

_____ ( molecules are locked in place; a crystalline solid has highly ordered structure

_____ ( molecules move past one another, have a definite volume, and can assume the shapes of their containers

_____ ( molecules are in constant random motion, and are compressible or expandable to fill a container.

The types of intermolecular forces determine the characteristics of the states of the substances: e.g. Melting point, boiling point, vapor pressure.

e.g. The greater the intermolecular force, the ______ the melting and boiling points.

The greater the intermolecular force, the ______ the vapor pressure.

Breaking the covalent bonds vs simply overcoming intermolecular forces.

(a) Subliming iodine crystal

(b) Dissolving HCL gas in water to make hydrochloric acid.

(c) Changing ozone, O3, to O2 gas

(d) Decomposing hydrogen peroxide, H2O2, to water and O2 Three types of forces,

(i) Dipole dipole forces: formed by the attraction between partially positive and negative charges of neighboring polar covalent molecules. (e.g. HCl.)

(ii) Hydrogen Bonding: A special kind of dipole dipole formed by the hydrogen atom in a polar bond and an unshared electron pair on a nearby small electronegative atom such as N, O or F.

(iii) London Dispersion Forces (LDF): a weak force formed by instantaneous dipoles created in a molecule caused by oscillating electrons. LDF increases as the size of the atom or the molecular weight increase. (e.g. Cl2)Example 1

Example 2

Example 3

Which of the following molecules can form hydrogen bonds with other molecules of the same kind? CH3F, CH3NH2, CH3OH, CH3Br.

Example 4Explain why HOCH2CH2OH has stronger attractive forces between its molecules than

does CH3CH2OHExample 5 Propane (C3H8) is a gas at room temperature, hexane (C6H14), a liquid while candle wax (C25H52) is a solid. Explain why. 2.2 Bonding in solid

2.2.1 Molecular solid consists of atoms or molecules held together by intermolecular forces. They are soft with relatively low melting points. E.g. Ar, H2O, CO2 Structural units are ___________

The forces holding units together are __________, ____________, __________

They usually get __________ melting points and boiling points, soft and ____ electric conductivity in solid and liquid form.

They are soluble in water if ______ ; soluble in hexane if _________. Example 1

Explain why benzene boils at 80oC, toluene boils at 111oC while phenol boils at 182oC.

Example 2

2.2.2 Networkcovalent solids consist of atoms held together in large networks by covalent bonds. E.g. diamond, graphite, quartz Atoms are held in an infinite ____________ network.

The force holding the atoms is the covalent bonds and they are ________ electron-pair bonds.

They are hard, ____________ in solvent, _____ melting point / boiling point and no electric conductivity. Consider the table below,

Diamond Graphite

Molecular structure

DescriptionEach carbon atom is bonded __________ to four other carbon atoms with sp3 hybrid orbitals. The carbon atoms are arranged in layers of interconnected hexagonal rings, each one bonded to three others (with sp2 hybrid orbitals) in the layer. The layers are held together by the weak bonds

Special features Unusual hardness and very _____ melting point.Slippery and ____________

Example 1

Given that Network covalent solids (like diamond) are usually much harder and have higher melting points than molecular solids. Discuss why.

Example 2

Most Network covalent solids are hard in nature and do not conduct electricity. But why Graphite is soft and conducts electricity?

2.2.3 Ionic solids consists of ions held together by ionic bonds. The strength of an ionic bond depends greatly on the charges of the ions.

The Structural units are orderly array (_______) of ______ and ________ negative ions.

No discrete molecules in ionic solids.

The forces holding units together are the electrostatic attraction among charges on ______ and _______ ions.

The ionic solids are hard, brittle, _____ melting point and they conduct electricity in liquid form, aqueous form solution.

They are usually soluble in _____ solvent.

Example 1

Explain why MgO has a higher melting point than NaCl.

Example 2

Discuss whether ionic solid can conduct electricity in solid form.

2.2.4 Metallic solids consist entirely of metal atoms.

The model visualizes the metal as an array of positive ions immersed in a sea of __________ valence electrons. The strength of the bonding increase as the number of electrons available for bonding increases. (i.e Na Mg Al in melting point) The mobile electrons explain why metals are ________ conductors of heat and electricity and why metals are __________ and _________. They have a ____ range of hardness and melting points.

Lesson 2: Exercise

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2.3.4.5.6. Classify each substance as to the type of solid it forms and the type of intermolecular force it has: Fe C2H6 CaCl2 graphite F2

sand CH3COOH HI(g) W 7. Classify each of the following solids as metallic, network covalent, ionic, or molecular. (i) It is soluble in water, does not conduct electricity, and has a melting point (about 200oC)(ii) It is insoluble in water and conducts electricity when melted(iii) It is insoluble in water and conducts electricity. (iv) It dissolves in water, conducts electricity when present in aqueous solution, and melts above 100oC.(v) It is malleable and conducts electricity.

8. Give the formula of a solid containing oxygen that is:

(i) Ionic

(ii) Polar molecular

(iii) Network covalent

(iv) Nonpolar molecular

9. For each of the following pairs, choose the member with the lower boiling point. Explain your reason in each case:

(i) NO2 or SO2 (ii) NaCl or HCl (iii) NH3 or AsH3 (iv) I2 or NaI (v) HCOOH or C6H5COOHLesson 3: Vapour pressure and phase diagram3.1 Change of state

The temperature of a substance _______ when heat is input while a substance is in one phrase. The temperature of a substance _______ during a phrase change.

Heat (enthalpy) of fusion the enthalpy change associated with _______ a solid. Heat (enthalpy) of vaporization the enthalpy change associated with vaporization of a liquid.

Remind that the formula, E = ______ , is used to calculate the energy change when there is a temperature change. Example 1

Enthalpy of fusion is always smaller than the enthalpy of vaporization. Why?

(For example, Enthalpy of fusion is 6.01 kJ/mol for ice and enthalpy of vaporization is 40.67 kJ/mol for water.)

Example 2

Calculate the enthalpy change upon converting 10.0g of ice at -25oC to water vapor at 125 oC at constant pressure of 1 atm. The specific heats of ice, water, and steam are 2.09 J/goC, 4.18 J/goC and 1.84 J/goC. 3.2 Vapor Pressure Vapor pressure of a liquid is the pressure exerted by its vapor when the liquid and vapor states are in dynamic equilibrium. Two opposing processes, _________ and __________ are occurring simultaneously at the same rate. There is no observable change, but a great deal is happening on the molecular level.

Vapor pressure increase with ___________ temperature. Normal boiling point the temperature at which the vapor pressure of a liquid is ____ kPa. Clausius Clapeyron Equation:

ln(P1 / P2 ) = (vap/R )(1/T2 1/T1), where R = 8.314 J/K molExample 1 Why vapor pressures increase with increasing temperature?

Example 2

Why does it take longer to cook food at higher elevations than at sea level?

How does a pressure cooker allow the food to cook more rapidly?

Example 3

(a) For Clausius Clapeyron Equation, Acetone boils at 56.5OC at 1.0 atm. The enthalpy of vaporization of acetone is 32.0 kJ/mol. What temperature will the acetone boil at 0.750 atom?

(b) What temperature will water boil on top of Mount Everest where atmospheric pressure is 0.475 atm?

3.3 Phase diagram A graphical way to summarize the conditions under which equilibria exist between the different states of matter. The diagram allows to predict the phase of a substance that is stable at any given temperature and pressure.

Triple point point representing the temperature and pressure at which __ phases coexist. Critical temperature above which the vapor cannot be _______no matter what pressure is applied.

Critical pressure required to produce ___________ at the critical temperature

Critical point the point defined by the ________________ and ________________ .

Example 1

Referring to the Phase Diagram of CO2,

(a) What happen if the pressure increases from 1 atm to 60 atm at constant temperature of -60oC

(b) What happen if the temperature increases from -60oC to -20oC at constant 60 atm pressure.

Example 2Referring to the Phase Diagram of H2O, (a) Describe what will happen if the water is kept at 1 atm while the temperature is increasing from -10 oC to 400 oC

(b) Describe what will happen if the water is kept at 0oC while the pressure is increasing from 0 atm to 300 atm.

(c) What are point A, B, C, D represent?

(d) The slope of AB is negative. What is the significance of negative sloping?

Lesson 3: Exercise1.

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5.6.Lesson 4: Solution and colligative properties4.1 Solution composition

Solvent _________ medium

Solute substance to be _________

Solution homogeneous mixture

Formula related to solution concentration (i) Mass percent =

(ii) Mole fraction =(iii) Molarity (mol dm-3) = (iv) Molality ( mol kg -1) = Example 1

What is the mole fraction of CH3OH and H2O in a solution prepared by dissolving 1.20g of alcohol in 16.8g of water?

Example 2

A 1.13M KOH solution has a density of 1.05 g/mL. Calculate its mass precent and molality.

Example 3

Example 4

4.2 Energies of solution formation

A major factor determining whether a solution forms is the relative strengths of ________________ between and among the ___________ and _________ particles. Like dissolves in Like polar substance (e.g. NaCl) dissolve in polar solvents (H2O) non-polar substance (e.g. C6H14) dissolve in non-polar solvents (CCl4) Explain why NaCl dissolves in H2O.

When ionic crystals NaCl are added to water, the ________ end of the water dipole is oriented toward the ________ ions, and the _________ end of the water dipole is oriented toward the _______ ions.

Interactions between solute and solvent molecules are known as ________.

When solvent is water, the interactions are referred to as _________.

For enthalpy change in forming a solution, Hsoln Hsoln = H1 + H2 + H3 H1 = input of energy to overcome attractive forces between solute molecules,

Energy is absorbed by the molecules to do so, so it should be __________ H2 = input of energy to overcome attractive forces between solvent molecules,

Energy is absorbed by the molecules to do so, so it should be __________

H3 = attractive interactions between solute and solvent,

Energy is released by the molecules to do so, so it should be __________

Note that H3 must be comparable in magnitude to H1H2 for a solution to form.

The formation of a solution can either be exothermic or endothermic.

The enthalpy of hydration depends on:

1) Distance between the ion and the dipole the closer the distance, the stronger the attraction

2) Charge on the ion the higher the charge, the stronger the attraction

3) Polarity of the solvating molecule the greater the magnitude of the dipole, the stronger the attraction

Example 1

What solvent would you choose, water of hexane C6H14 to dissolve NaCl, HF, octane and (NH4)2SO4. Example 2 The lattice energy of CaI2 is -2059kJ/mol. The enthalpy of hydration is -2163kJ/mol. What is the enthalpy of solution? Write equation to show the dissolution of CaI2 in water.

Example 3

Explain why the enthalpy of hydration of Na+ (-405 kJ/mol) is somewhat more negative than that of Cs+ (-263 kJ/mol), whereas that of Mg2+ is much more negative (-1922 kJ/mol) than that of either Na+ or Cs+.

4.3 Factors affecting solubility

A solution that is in equilibrium with undissolved solute is ________ If more solute can still be added, the solution is _________

A solution that contains a greater amount of solute than that of a saturated solution is ______________ The amount of solute needed to form a saturated solution in a given quantity of solvent (100g of H2O) at a given temperature is ________

(1) Solute solvent interactions

Non-polar ( hydrophobic; polar( hydrophilic

Pairs of liquid (such as acetone and water) that mix in all proportions are ________, whereas those that do not dissolve in one another are _________.

(2) Pressure by Henrys law : C=kP, solubility of a gas in a solvent is _______ as the pressure over the solvent increases. (C = solubility of the gas in the solution phase)

(K = Henrys Law constant)

(P = partial pressure of the gaseous solute above the solution) (3) Temperature

The solubility of most solid solutes in water ________ as the temperature of the solution increases.

The solubility of gases in water _________ with increasing the temperature.

Example 1

Determine whether or not each of the following compounds is likely to be water soluble:

CH3CH(OH)NH2 CH3(CH3)4CH2NH2 C4H9CH=CH2 NH2CH2COOH

Example 2

Example 3The solubility of CO2 at 25OC and 1 atm is 0.034M

a) what is Henrys Law Constant

b) What would the solubility of CO2 in water be at 0.038 atm and 25OC. Example 4

4.4 Colligative properties

Colligative Properties physical properties of solutions that depend on the __________ but not on the nature of solute particles.

Vapor Pressure Lowering a nonvolatile solute _______ the vapor pressure of its solution because they reduce the tendency of solvent molecules to escape. Raoults Law the partial pressure exerted by a solvent vapor above a solution, PA, equals the product of the mole fraction of the solvent in the solution, XA, times the vapor pressure of the pure solvent, PA0PA = XA PA0Example 1

Example 2

(Hint: One mole of the strong electrolyte, CuCl2, will dissolve to give 3 moles of ions)

Example 3

Example 4

(* Hint: It involves volatile solutes ( Use Ptotal = PA + PB = XA PA0 + XB PB0 ) 4.5 Boiling point elevation and freezing point depression

When there is an addition of nonvolatile solute, the vapor pressure of the solution lowers.

The vapor pressure curve of the solution is shifting downward result in the following phase diagram,

Boiling point elevation ( Tb = Kb m, where Tb is the boiling point elevation

Kb is molal boiling point elevation constant

m is molality of the solution Freezing point depression ( Tf = Kf m, where Tf is the freezing point elevation

Kf is molal freezing point elevation constant

m is molality of the solution

Vant Hoff factor the ratio of experimental value to calculated value: The Vant Hoff factor is always found to be smaller value. (i.e. the measured value / observed vale is smaller than the calculated value / expected vale). It is found that Vant Hoff factor approaches whole number only in ___________ solution. In more concentrated solutions, some of the positive and negative ions are paired, decreasing the total molality of particles. This is called _______________.

Example 1 Automotive antifreeze consists of ethylene glycol (C2H6O2), a nonvolatile nonelectrolyte. Calculate the boiling point and freezing point of a 25.0 mass % solution of ethylene glycol in water. Example 2

List the following aqueous solutions in order of their expected freezing point: 0.050m CaCl2, 0.15m NaCl, 0.10m HCl, 0.050m HC2H3O2, 0.10m C12H22O11.

Example 3

When 2.50g of cortisone acetate is dissolved in 50.00g of camphor, C10H16O, (Kf = 40.0oC/m), the freezing point of the mixture is determined to be 173.44oC. That of pure camphor is 178.4OC. What is the molar mass of cortisone acetate?

Example 4

A compound contains 42.9%, 2.4%H, 16.6%N and 38.1%O. the addition of 3.16g of this compound to 75.0mL of cyclohexane (density = 0.779g/mL) gives a solution with a freezing point at 0.0OC. Find its molecular formula. (Kf = 20.2OC/m)

Example 5

A 0.00200m aqueous solution of an ionic compound Co(NH3)5(NO2)Cl freezes at 0.00732oC. How many moles of ions does 1 mole of the salt produce or being dissolved in water? 4.6 Osmotic Pressure

The pressure required to prevent osmosis by pure solvent is the _____________, , of the solution (n/V)RT = MRT

osmotic pressure of the solution

n = number of moles of solute in volume V in liters

M = molarity of solution

R = ideal gas constant

T = Temperature in Kelvin

Example 1

What osmotic pressure would a sucrose solution (C12H22O11) made of 5.00g of sucrose in 117.0 g of water exhibit at 25oC?

Example 2

Osmotic pressure is useful in determining the molecular weights. What substances are usually measured by this method?

I) Very expensive substances

II) Substances that can be prepared only in very small amounts

III) Substances of very high molecular weight that are not very solubleA) I & IIB) I & III

C) II&III

D) I & II & III

Example 3

Pepsin is an enzyme present in the human digestive tract. A solution of 0.500g sample of purified pepsin in 30.0 mL of aqueous solution exhibits an osmotic pressure of 8.92 torr, at 27.0oC. Estimate the molecular weight of pepsin.

Exercise

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To calculate enthalpy change of solution,

KF(S) ( K+ (aq) + F (aq)

1) When 1 mole ionic solid is broken down to gaseous ions

KF(S) ( K+ (g) + F (g)

Negative lattice energy = - (-821kJ)

2) When 1 mole of aqueous ions is formed from gaseous ions

K+ (g) + F (g) ( K+ (aq) + F (aq)

Enthalpy change of hydration = -819 kJ

3) When 1 mole of solute is dissolved in sufficient solvent

KF(S) ( K+ (aq) + F (aq)

Enthalpy change of hydration = +2 kJ

Applied pressure on the ____ arm of the apparatus stops net movement of solvent from the _____ side of the semipermeable membrane. This applied pressure is the osmotic pressure of the solution.

Osmosis stops when the column of a solution on the left becomes high enough to exert sufficient pressure at the membrane to counter the net movement of a solvent. At this point the solution on the ____ has become more dilute, but there still exists a difference in concentration between the two solutions.

Net movement of solvent from the pure solvent or a solution with _____ solute concentration to a solution with ______ solute concentration.

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