lecture notes 2 (1)

8/2/2019 Lecture Notes 2 (1)

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NORTHERN CARIBBEAN UNIVERSITY

DEPARTMENT OF BIOLOGY and CHEMISTRY

LECTURE NOTES

CHEM122: GENERAL CHEMISTRY II

Assistant Professor: Dr. Nicole White

3. States of Matter

There are 3 states of matter (solid, liquid and gas) which are classified as such because of the

interactions between the molecules found within the substance. For instance, in a solid the

molecules are found packed tightly together with only a very small amount of kinetic energy

compared to the intermolecular attraction. Solids are thus nearly incompressible and rigid and their

molecules or ions are in close contact and do not move. For a liquid its molecules possess more

kinetic energy and can move about relative to each other (called a fluid) but this kinetic energy is

not enough to overcome the intermolecular attraction.Their molecules are more tightly packed.

Liquids are relatively incompressible fluids. Gases too like liquids are fluid but they unlike liquids

and solids possess a great amount of kinetic energy, enough to overcome any intermolecular

attraction that might be present and so individual particles are able to move far away from each

other.Gases are compressible fluids.


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Each state of matter can be converted to another state by going through a phase transition a phase

is a given state of either a substance or a solution; melting freezing, boiling condensing and

sublimation deposition. A transfer of energy is required in all cases, endothermic or exothermic

processes.

Melting: change of a solid to a liquid.

Freezing: change a liquid to a solid.

Vaporization: change of a solid or liquid to a gas. Change of solid to vapor often called

sublimation.


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Condensation: change of a gas to a liquid or solid. Change of a gas to a solid often called

deposition.

Can you differentiate between the exothermic and endothermic processes?

When a substance is heated, for example ice, first there is an increase in temperature with

absorption of heat, however, at the melting point the temperature remains constant until all the ice

has been melted to water. At this stage which is seen as a flat line on the heating curve the solid and

liquid phases are in equilibrium. After complete fusion there is an increase in temperature which

levels off again once the water and steam phases are in equilibrium.

Vapour pressure, P

This is defined as the partial pressure of a vapour over the liquid measured at equilibrium at a given

temperature. At any given temperature there will be some molecules within a liquid that possess

enough kinetic energy to overcome the intermolecular attractions becoming a vapour. In a closed

system the vapour molecules cannot move away from the liquid and in time the two will interact

with the highly energetic gas molecules transferring energy to the less energetic liquid molecules

resulting in condensation.

Heat of vaporization: heatneeded for the vaporization ofa liquid.

H2O(l) H

2O(g) H =

40.7 kJHeat of fusion: heat needed for

the melting of a solid.

H2O(s) H

2O(l) H =

6.01 kJTemperature does not change

during the change from onephase to another.


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Initially the rate of vapourization is greater than that of condensation but given enough time the two

rates be the same resulting in a dynamic equilibrium and it is at this point that the vapor pressure of

the liquid can be determined. The differences in vapour pressure at different temperatures is due to

the fact that at higher temperatures more molecules will have enough kinetic energy to overcome

the intermolecular attractions and go into the vapour phase, the reverse is true for a decrease in

temperature. Various liquids have differing vapour pressures due to differences in strengths of the

intermolecular interactions that bind their molecules together the lower the vapour pressure the

stronger the intermolecular forces.

In an open system, when the vapour pressure equals the atmospheric pressure the liquid begins to

vapourize throughout resulting in bubbles and the liquid is said to be boiling.Normal boilingpointthe temperature at which the vapor pressure of a liquid is equal to atmospheric pressure (1 atm).

The Clausius-Clapeyron equation shows how the vapor pressure and temperature are related. It can

be written as:

CTR

HP

vap +

=1

ln


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R gas constant; C a constant that is characteristic of the liquid.

A straight line plot results when ln P vs. 1/T is plotted and has a slope of Hvap/R. Clausius

Clapeyron equation is true for any two pairs of points.Write the equation for each and combine to

get:

Phase diagrams

A phase diagram is a graphical way to summarize the conditions under which the different states of

a substance are stable.

The sold lines represent the conditions of temperature and pressure at which equilibrium exists

between the two phases on either side of the line. The normal melting and boiling points are those

when the pressure is 1 atmosphere. These can be found from the phase diagram by drawing a line

across at 1 atmosphere pressure. The line separating the solid and liquid phases is lightly affected

by pressure as can be seen in the figure where only a small change in temperature is observed for a

Graph of pressure-temperature

relationship; describes when 1,2,3 or

more phases are present and/or in

equilibrium with each other.

Lines indicate equilibrium state two

phases.

Triple point- Temp. and press. where all

three phases co-exist in equilibrium.

Critical temp.- Temp. where substance

must always be gas, no matterwhatpressure.

Critical pressure-vapor pressure at critical temp.Critical point-point where system is at its critical pressure and

temp.

=

211

2 11lnTTR

HPP vap


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large pressure change. Also, for most substances the line tends to lean to the right indicating that

the solid is denser than the liquid. The opposite is seen as ice is less dense than water and so the

phase boundary will lean towards the left. There is a point on the diagram at which all 3 phases are

in equilibrium and this point is known as the triple point any temperature or pressure below this

point finds only the solid and gaseous phases in equilibrium. The critical point, however, is where

we have only 1 phase being stable as seen above for water only gas is observed. A supercritical

fluid is formed when a substance is heated above its critical temperature whether or not it is at its

critical pressure and has properties intermediate between those of a liquid and those of a gas.

A. Liquids

Surface tension: This is the energy required to increase the surface area of a liquid by a unit

amount. What is the theory behind this?

Molecules in the bulk of a liquid experience no net attractive force as they are pulled on all sides by

adjacent molecules. On the other hand, molecules on the surface experience a net pull toward the

interior of the liquid making the surface tend to as small a surface area as possible and a substance

does not penetrate it easily.

Viscosity: This is the resistance to flow that is exhibited by all liquids. In other words, if some

liquid were poured on the floor, the easier it is for a liquid to spread out the less viscous it is.

Consider water and mercury. Viscosity decreases as the temperature increases since increased

temperatures tend to cause increased mobility of the molecule.


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The terms explained above are dependent on the intermolecular attraction between molecules in a

substance. This force of attraction can be used to explain not only viscosity and surface tension but

also vapour pressure. Since compounds with large intermolecular forces have lower vapor

pressures, we predict that one has to go to higher temperature to make them boil.

So we can use the following principles to predict relative boiling points:

Stronger intermolecular forces mean higher boiling points.

Other things being equal:

Polar molecules boil higher than nonpolar molecules.

Hydrogen-bonded molecules boil higher than nonhydrogen-bonded molecules. Large molecules boil higher than small molecules.

Intermolecular forces

There are 2 basic types of intermolecular attractions: van der Waals forces and hydrogen bonding.

The former can be subdivided in dipole-dipole and London forces.

Hydrogen bonding: this is a weak to moderate attractive force between a hydrogen atom covalentlybonded to a very electronegative atoms, X, and a lone pair or electrons on another small,

electronegative atom, Y. Hydrogen is usually bonded to O, N or F.Strongest of the intermolecular

forces between neutral molecules.

Dipole-dipole force: attractive force between polar molecules with positive end of one molecule is

aligned with negative side of other. The strength of this intermolecular force depends on the size of

the molecular dipole moment.It is approximately 5% as strong as H-bonding forces.

London (dispersion) Forces: interactions between instantaneously formed electric dipoles on

neighboring polar or nonpolar molecules. All molecules experience this type of intermolecular

attraction. The more atoms there are in a molecule the greater the dispersion force. It is the weakest

of the intermolecular forces. Approximately 1% as strong as dipole-dipole forces (except for large

molecules where the forces can be comparable).


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The below figure shows the higher boiling points of molecules that exhibit hydrogen bonding

compared to compounds formed from elements from the same group with greater molecular

weight.

B. Solids

A nearly incompressible state of matter with a well defined shape.

Types of solids:

Crystalline a well defined arrangement of atoms; this arrangement is often seen on a macroscopic

level.

Ionic solids ionic bonds hold the solids in a regular three dimensional arrangement.

Molecular solid solids like ice are held together by intermolecular forces.

Covalent network a solid consists of atoms held together in large networks or chains by

covalent networks.

Metallic similar to covalent network except with metals. Provides high conductivity.


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Amorphous atoms are randomly arranged. No order exists in the solid. These are normally

formed when a pre-melted solid is cooled quickly so that the units cannot arrange themselves in an

ordered state and are therefore random.

The melting point of each solid is determined by the attractive forces acting within it whether they

are inter- or intra-molecular. The strength of the attractive force can also determine certain physical

properties of the solid such as electrical conductivity and hardness.

Hardness The ability of the structural units within a solid to move with respect to each

other is determined by the forces of attraction within said solid. If the attraction is strong

then the solid will be quite hard and vice versa.

Electrical conductivity in this instance the metallic solid would be the better conductor as

its electrons are delocalized whilst in the other solid types the electrons are localized. In

anionic liquid electricity is conducted via the ions.

There are different types of crystals. Some examples are shown below. With sides a, b c abd angles

, and . Side a is found opposite side and so on. Each of the below represent a unit cell which

is a repeating unit in a crystal. Generally speaking, the lattice of the crystal is comprised of points

at which an atom may be found. A scientific technique used to determine the structure of a crystal

is X-ray diffraction: powder and crystal.


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C. Gases

Gases are known for their high kinetic energies and compressibility. They exert force (P = AF in

kg/ms2 (Pa)) on objects they come into contact with and this pressure can be measured by a

barometer or manometer.

The barometer measures the pressure of the atmospheric pressure and the monometer measures the

pressure of a gas or liquid in a vessel.


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Pressure can be expressed in many units: 1 atm (atmosphere) = 101325 Pa (Pascal) = 760 mmHg or

torr. The SI units for pressure is Pascals.

Gas laws

There are various gas laws that relate pressure, volume, temperature and moles.

Boyles Law: this relates pressure and volume for a specific amount of gas at a fixed temperature.

Around 1660 Robert Boyle was experimenting with the properties of gases under conditions where

the temperature was held constants. He would take a sample of gas and measure the volume of the

gas at different pressures (without changing the temperature).

PV = constant


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The value does not remain exactly constant with volume-pressure variations. The volume of a gas

is inversely proportional to the pressure applied on it.P

V1

. This relationship,PV = constant,

deviates at high pressures.

When considering 2 sets of pressure and volume data the following equation can be applied: P1V1

= P2V2. If 3 out of the four values are known then the equation can be manipulated so that the

fourth can be determined.

Charless law: this law relates volume to temperature for a given amount of gas at a fixed

pressure. In 1787 French scientist Jacques Charles was studying the properties of gases at constant

pressure. Charless law deviates at low temperature and high pressure.

2

2

1

1,tanTV

TValsotcons

TV ==


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When the above gas laws are combined:2

22

1

11,tanT

VP

T

VPalsotcons

T

PV== .

Gay-Lussac's Law: At the beginning of the 19th century the French chemist Joseph Gay-Lussac

made several contributions to the understanding of the properties of Gases. He did studies that were

concerned the behavior of gases at constant volume. That is, how does the pressure depend on

temperature if the volume is held constant?

2

2

1

1,tanT

P

T

Palsotcons

T

P==

Gay-Lussac's Law of Combining Volumes: Gay-Lussac conducted other experiments and found

that the ratio of the volumes of gases involved in chemical reactions (at the same temperature and

pressure) were always simple fractions (fractions composed of small whole numbers).


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For example, in the reaction N2(g) + 3 H2(g) 2 NH3(g),

at constant temperature and pressure, Gay-Lussac found that the ratio of the volume of H 2(g) to the

volume of N2(g) was always 3 to 1, regardless of the initial volumes used. Likewise, the ratio of the

volume of H2(g) to the volume of NH3(g) was always 3 to 2.

Avogadros law: states that equal volumes of any two gases at the same temperature and pressure

contain the same number of molecules. One mole of gas contains the same number of molecules

6.02 1023 this volume is known as the molar gas volume, Vm = 22.4 L/mol at STP.

STP standard temperature and pressure.

The ideal gas law: An ideal gas is an idealized model for real gases that have sufficiently lowdensities. The condition of low density means that the molecules of the gas are so far apart that they

do not interact (except during collisions that are effectively elastic). From the relationship

P

TtconsVm = tan finding the volume for n moles of gas the equation becomesPV = nRT.

R gas constant= 0.0820578 Latm/Kmol = 0.0831451 Lbar/Kmol = 8.31451 J/Kmol., n number

of moles.

Gas Mixtures, Partial Pressures, and Dalton's Law: In a mixture of gases we speak of the total

pressure of the gas mixture and the "partial pressures" of the individual gases in the mixture. The

partial pressure of a gas in a mixture is defined as the pressure the gas would have (at the same

temperature and in the same volume) if all the other component gases were removed leaving only

the one gas.

Dalton's law of partial pressures says that the total pressure of a gas mixture is the sum of all the

partial pressures of all the gases in the mixture. If we number the gases as gas 1, gas 2, gas 3, and

so on then,p =p1 +p2 +p3 + . . . .

We can derive Dalton's law from the ideal g as law. Note that the n inV

nRTP=


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doesn't have any information about the identity of the gas and the equation is the same regardless of

the identity of the gas. There is no reason why n couldn't be the sum of the number moles of each

component in a gas mixture.

That is, n = n1 + n2 + n3 + . . . .

If we place this n into the ideal gas law we get,

which is Dalton's law.

Gas density

The density of a gas can be determined by the relationship mass = moles molar mass. Density hasthe units g/cm3 or mass/volume and so manipulation of the ideal gas law gives:

RTPM

RTV

mPM

RTM

mPVnRTPV

m

m

m

=

=

==

Kinetic Molecular Theory of Gases

The kinetic molecular theory of gases states that all matter consists of extremely tiny particles that

are in constant motion. In a solid the molecules are too close and too ordered and so dont have a

lot of kinetic energy. These molecules can only vibrate. Liquid molecules have more kinetic energy


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than solids but not enough to overcome the attractive forces between molecules and so the space

between them is small. In gases, on the other hand, molecules have a large amount of kinetic

energy and so are able to escape the attraction between molecules and travel great distances. All

gas molecules have kinetic energy:

2

2

1 mvEk = . The larger the mass the smaller the average speed.

This means that heavy gases are much slower than lighter ones.

There are five fundamental concepts of kinetic molecular theory.

1. A gas is composed of molecules whose size is much smaller than the distance between them.

2. Gas molecules move randomly in straight lines at various speeds and in every possible direction.

3. Except when gas molecules collide, forces of attraction and repulsion between them are

negligible.

4. When collisions between molecules occur, they are elastic. Their kinetic energy remains

constant. Gas molecules collide (elastically) with the walls of the container to produce a pressure,

p. The molecules also collide with each other. Molecules do not lose or gain energy in their

collisions with the walls or with each other.

5. The average kinetic energy of gas molecules is proportional to the absolute temperature.


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Diffusion: This is the spread of gas molecules of one type through those of another type to occupy

the space uniformly. If we release a small amount of ammonia gas in a quiet room (no air currents

or turbulence or mixing) the ammonia will eventually spread throughout the room, due to the

motion of the molecules. It may take several minutes to spread throughout the room because as it

travels it collides with other gas molecules and changes direction constantly.

Effusion: This is the escape of gas molecules from a container through a tiny hole. The rate of

effusion is inversely proportional to the square root of its molar mass at constant temperature and

pressure. The rate at which a molecule diffuses through another gas is proportional to its average

velocity. Thus the diffusion rate is inversely proportional to the square root of the formula weight.

mMrate 1=

Effusion is dependent on the cross sectional area of the hole, the number of molecules per unit

volume, and the average molecular speed. We will not need to know how to calculate absolute

effusion rates in this course, but we will need to be able to calculate relative effusion rates. The

ratio of the effusion rate of gas A (RateA) to the effusion rate of gas B (RateB) is, from the above

discussion,

A

B

m

m

B

A

M

M

rate

rate=

4. Analytical Chemistry

Analytical chemistry deals with the separation, identification and determination of components in a

sample and consists of both qualitative and quantitative analyses. In carrying out quantitative

analysis one must follow certain steps:

1. Method selection.

2. Sample acquisition and preparation.


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3. Elimination of interference.

4. Measurement of analyte.

5. Calculation of results.

6. Determination of errors.

Method Selection: When choosing a method certain factors need to be taken into consideration,

the first and most important being the level of accuracy required. Cost is another important factor

as well as the number of samples that can be measured at any one time. Bearing all this in mind, it

is a combination of these and other factors, the best compromise depending on the needs of the

analyst, that influence the method chosen.

Sample Acquisition and Preparation: The sample chosen must be representative of the bulk

solution, sometimes referred as the stock in the case of solutions, first and foremost. Careful

records must be kept of the source and history of the sample. In some cases, the sample may need

to be ground into a suitable size, for a solid, or diluted in the case of a liquid sample. From each

selected sample multiple measurements are carried out and so it is imperative that these replicate

samples are treated identically.

Elimination of Interference: Interference is any substance present in sample that might cause an

error in the analysis of the analyte it affects an analytical signal or background.

Measuring Analyte

Calculation of Results

Determination of Errors: As stated earlier proper records need to be taken in sample preparation

and measurement and with this can the error or uncertainty associated with the final result can be

determined. These errors can be obtained from the equipment used and a good understanding of

errors is essential for any quantitative measurement and the reliability of the results reported.

Scientists try to minimize the sources of errors as best they can and ensure that no source is


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unaccounted for. This is an important area and will be dealt with in greater detail later in this

course.

Standard deviation (std. dev.)

This is a measure of how closely replicate data cluster around the mean precision. Standard

deviation can be applied to a population, , as well as to a sample of the population,s.

N

xi =2)(

and1

)(2

=

N

xxs

i

xi ith value

N number of values

meanpopulation

meansamplex

The most common applications of standard deviations are associated with addition/subtraction and

multiplication/division calculations. It is extremely important that you learn these.

222

:/

cbay

cbaynsubtractioaddition

++=

+=and 222 )()()(

:/

cc

bb

aa

yy

cbaydivisiontionmultiplica

++==

lecture notes 2 (1)

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