KMT and Physical StatesIn this portion of the course we are going to focus on the states or phases of matter (solid, liquid and gas). We are going to approach this topic from the following perspective:

The physical state that a sample is found in is directly related to theKinetic Energy of the molecules in that sample.

Due to the differences in how the molecules move (or don’t move) in each phase of matter, they all behave differently. We will look at each state at a molecular level, explore the properties of each and analyze how the molecular movement changes as we move from solid to liquid and finally to gas.

There will also be some discussion on the uniqueness of water, the various “liquid” samples that exist and how we can differentiate between them. To finish up our study of physical states we will study how changes in pressure, volume, temperature and molar amount can cause changes in state.

What is KMT?KMT stands for Kinetic Molecular Theory: It is a series of assumptions that allow us to explain the observed properties and the behavior of matter in each phase on a molecular level.

For each phase we will first look at the assumptions of the KMT and then create a definition or explanation for the phase.

KMT Gases:

1. Gases consist of tiny particles that are in random ceaseless motion.

2. Collisions between molecules and the container walls are always 100% elastic. (Elastic Collisions are ones in which no energy is lost during a collision)

3. The space between gas particles is very large compared to the minute size of the particles themselves

4. There are no attractive or repulsive forces between molecules when they are in the gas phase & the volume of the particles themselves is negligible.

5. The Kinetic Energy of the particles in the gas is directly proportional to the temperature.

So what do gas molecules look like?

Imagine a glass cube with super bounce balls inside.The balls are bouncing off the walls and off each other, however, they never slow down or stop moving!

Common examples:

Bingo machine Lottery machine Cash grab booth

How do we define the Gas Phase?

A state of matter where the particles show fluidity, are in random, ceaseless motion and will distribute themselves evenly throughout the entirety of the container, thus the shape and volume of a sample in this form is determined entirely by the size and shape of its container.

Properties of Gases:

Low Density: Gases have extremely low densities, a gas is approximately 1/1,000th that of a solid or liquid of the same substance.

Compressibility: because of the low density, a gas's volume can be decreased by adding pressure ... this property is called compressibility.

Expansion or Diffusion: Because of the random ceaseless motion of the molecules in a gas state the particles will reach out to all regions of the container the substance is in. There will be an even distribution of the molecules throughout the container. A more specific explanation of this property will be covered in more detail when we get to Graham's Law of Diffusion ... a law that discusses the rates at which substances will expand in the gas phase.

Fluidity: Because the molecules are far apart and collide in a fashion similar to pool balls, the molecules are said to be fluid. By fluid I mean to refer to their ability to move past each other with relatively no resistance. Another word synonymous with Fluidity is Flowing...this paints a picture of a group of particles sliding by one another smoothly.

KMT Liquids:

The Four Assumptions of the Liquid KMT:

1. Liquids consist of tiny particles that are close together and are in constant rotational and translational motion.

2. The particles in the liquid state are held together by Intermolecular Molecular Forces (IMFs) that have a greater effect on the particles because of the lower kinetic energy (KE) of the particles.

3. Due to the closeness of the particles with respect to one another along with the attraction due to the IMF combine to provide the fluid nature of liquids.

4. Just like in gases, the KE of the particles in this state is directly proportional to the Absolute Temperature.

So what do Liquid molecules look like?

In the liquid state the particles are veryclose together (compared to gas particlesthat are very far apart) and have translational and rotational motion thatgive this state a fluid nature.

Common Examples: Lava Flow White Water

Here is an example of water flowing around the stones and boulders … this exemplifies fluidity.

Liquid: state of matter where the particles arein no set arrangement, but show signs of fluid motion. The shape of the material in this state is determined by it container since liquids always take the shape of their container. The volume of the substance in this state is fixed

Many samples of matter that we consider to be “liquid” are actually not. By definition a liquid is a sample of a pure substance with a set volume, whose molecules have translational and rotational motion and thus takes the shape of its container. So what are all the other “liquid samples? Read on…Liquids, Solutions, Colloids & Suspensions…

In this section we are going to take a look at mixtures in the liquid state. There are three types of mixtures that can exist in the liquid state. All three are dependant upon the size of the particles that are distributed throughout the liquid medium. The Classification of these liquid mixtures is primarily dependant on

the size of the particles. The size of the particles typically tells us whether the particle will stay evenly dispersed in the liquid or if it will settle out.

The table below shows a summary of the properties of each of the three types…

Here is a molecular model for a solution…

Net Ionic Equation for Salt Water:NaCl + H2O Na+ + Cl- + H2O)

Mg(OH)2 Mg+2 + 2 (OH)-1 + H2O

The purpose of the net ionic equation is to illustrate how ionic compounds break into ions in solution and to show the total number of particles in a mixture…this will become very important in later chapters…here it is just to show that ionic compounds break into individual ions in solution.

Here is a molecular model of a colloid…

Milk, as well as other colloids, is composed of water and high numbers of colloids that are evenly distributed throughout the sample of water. The colloids can be held in place versus the particles in a suspension because their particle size is small enough that the structure of the water is able to keep them from settling out.

If you look at the table above that summarizes the three classes of the liquid mixtures, you will see that the two that we have discussed thus far have particles that remain dispersed in the liquid phase. The primary way that we can tell the difference between these two types of liquid phase mixtures is by passing a beam of light through the sample and observing the interference of the light or lack thereof.

The Tyndall Effect: when a beam of light is passed through a sample of a liquid where the particles dispersed within cannot be seen, the beam of light is scattered. You can notice a broadening of the beam as well as noticing the path that the light follows through the sample. This Tyndall Effect is noticeable when a car’s headlights cut through fog, water vapor or dust.

When a light beam is passed through a solution the particles (individual ions) are too small to scatter the beam thus the Tyndall Effect is not observed.

This is an image from a microscope that shows the complex arrangement of a number of molecules forming a mass that we refer to as a “Colloid.”

These massive collections of molecules can be viewed by a basic microscope. A well-known source/example of colloids is milk. The fat particles in the milk are colloids.

Here is a molecular model for a suspension…

In this model the particles of clay are temporarily held between the water molecules. After some time, the clay particles, due to their mass and lack of attractive forces, settle out and form a layer at the bottom of the glass. This is similar to what happens in rivers after big storms. The mass induction of water from the storm pulls soil from the land. While the river is moving quickly the soil is suspended in the water. When the storm subsides the soil will settle out and deposit on the riverbed. Below is a picture that show the sample after the clay has settled out.

The Particles in this type of liquid mixture are visible so the Tyndall Effect does not need to be observed. In most cases, Colloids are not transparent or even translucent, so the light would be completely absorbed by the mixture.

Information about Solutions…

As we have already learned, solutions are homogeneous mixtures in the liquid phase. These mixtures can be produced from mixing pure substances together. There are two components in a solution:

• The solvent is the dissolving medium (the one that makes up the majority of the mixture).

• The solute is the substance that is dissolving in the solvent (the one that comprises the minority of the solution).

When the amounts of each substance are approximately equal, then it is not worth the energy or the time to determine which is which.

When the solution is made, the solute particles are pulled apart and then evenly distributed throughout the solvent. These particles are not visible to the naked eye.

There are, however, some occurrences when the solute cannot be broken down in the solvent. These are deemed insoluble. For a quick reference you could look at the solubility rules or follow the general rule of thumb: “like dissolves like.”

When two substances are not soluble, they may show this in a variety of ways…

Liquid Solute and Liquid Solvent Interactions.

• If two liquid substances are able to freely dissolve in one another are deemed Miscible. When miscible liquids combine you end up with a homogeneous liquid mixture as shown in Figure A below.

• If two liquids are not soluble in one another, they are said to be Immiscible. When two immiscible liquids are mixed they will separate and form a distinct interface as shown in the figure below in Figure B.

Gas Solute and Liquid Solvent Interactions:

You don’t typically think about gases being dissolved in liquid solvents but it is a fairly common occurrence. Gases dissolve into liquids all the time…when water vapor molecules collide with the surface of a sample of liquid water they are pulled into the sample. Another common occurrence is the dissolution of carbon dioxide gas into pop. But how do they get the carbon dioxide to dissolve into the pop? Why does it immediately come out of solution when the cap is removed?

To better understand the solubility of gases in liquids we must look into some dissolution techniques used to create the solution in the first place. In the bottling industry pop and other carbonated drinks are bottled under high pressure CO2 (5 – 10 atm). When the liquid is placed into the bottle along with the CO2, some of the gas can be dissolved in the liquid. When the drink is bottled, the CO2 pressure above the liquid is very high, which allows the CO2 to remain in solution until the cap is removed, decreasing the pressure above the liquid solution. This sudden release of a gas from a liquid is known as effervescence.

William Henry studied the impact of Pressure on the Solubility of a Gas in a Liquid and developed a law to explain it.

Henry’s Law: the solubility of a gas in a liquid is directly proportional to the partial pressure of that gas on the liquid’s surface (when temperature is constant).

As the pressure of the gas above the liquid solvent increases,the solubility of that gas also increases.

A good illustration of Immiscibility is Italian – Oil & Vinegar salad dressing. When you shake it up well there is, what seems to be a homogeneous mix of all the ingredients, but when left to sit the oil, vinegar, ect. all separate.

The Effects of Temperature on Solubility… Effects on Gas Solubility:

As the temperature of a system is increased, the KE of the gas molecules in solution also increases. Due to the greater KE, more gas molecules are able to break the attractive forces of the solvent.

Effects on Liquid and Solid Solubility:

As the temperature of a system containing a solid or a liquid solute is heated, the increased KE of the solute molecules typically yields higher solubility. Although the general trend holds true for most solids and liquids, some are conversely affected. And among those whose solubility does increase with temperature, there are drastic variations in the magnitude of the effect of temperature from substance to substance.

An example temperature change on the solubility of a solid is the difference in dissolution rates in iced tea versus hot tea. The rate at which sugar can be dissolved in the hot tea is much faster than that of the iced tea.

Why does this occur? Again refer back to the effect that increased KE has on the motion and freedom of molecules. As the temperature increases, so does the KE of the molecules, therefore, the energy of the solute and solvent are increased allowing for faster diffusion of solute particles throughout the solution.

Effects of Increased Surface Area on Rate of Diffusion: The surface area is another factor affecting the rate of solubility for solid solutes. The rate of dissolution increases with an increase in the surface area. By crushing large crystals into smaller, finer crystals the solvent has more area to attack and therefore can remove ions or molecules from the crystal surface at a much greater rate.

Effects of Solution Agitation on the Rate of Dissolution:

Agitation refers to the mixing or stirring of a solution. Agitation speeds up the rate of diffusion and dissociation by dispersing the solute particles from the areas of high concentration right around the solute crystal(s) to regions of lower concentration elsewhere in the solution.

Agitation helps the solution reach its saturation point in less time (the # of molecules that may enter the solution will not change, but the amount of time it takes to get there will be significantly shorter).

Solubility, Saturation & the Solution Process:

Solubility refers to the amount of solute that can be dissolved in a given solvent at a certain temperature.

Solubility is determined by the phase equilibrium that becomes established between dissolution and recrystalization occurring at equal rates (solution equilibrium).

Saturation is defined as the point where a maximum amount of solute has been dissolved into the solution and equilibrium has been established.

When a solution still has the ability to accept solute, it is said to be unsaturated.

Supersaturation occurs when the solvent is heated or the KE of the solution is raised and the amount of solute that enters the solution is greater than the saturation point. The solution will exist in this way for a long period of time, until it is disturbed.

Disruption may occur in any combination of the following ways…i. Decrease in Temperatureii. Increase in Pressureiii. Introduction of a seed crystal or a small crystal of the same

substance that is added to the solution and then catalyzes a chain reaction of recrystalization that will continue until the solution will

reaches its saturation point where solution equilibrium will be established.

Properties of Liquids Resulting from the KMT Assumptions… High Density – results from the closeness of the particles to one another in the sample. Liquids have approximately 90% of the density of the solid state of the same substance. Relative Incompressibility – Even under extreme increases in pressure, up to and beyond 1000 atm, liquid water only decreases by 4%. This is due to the close packing of the particles in the sample.

Ability to Diffuse – Similar to gases, liquids are able to diffuse and create an even mixture of two or more substances. Although both gases and liquids can diffuse, the rate of diffusion for a liquid is much slower because of how close the particles are in the liquid state.

Surface Tension – is a property that is unique to the liquid state. Surface Tension is defined as a force that tends to pull adjacent parts of a liquid’s surface together, thereby decreasing surface area to the smallest possible size.

Freezing, Melting, Boiling & Condensation – Liquid forms of pure substances have definite freezing and melting points. These points are dependent on the arrangement of particles and the attractive forces between particles. The Molar Heat of Vaporization is the amount of energy needed to change one mole of a substance from the liquid state to the gas state.

Shape is Determined by the Container – because of the fluid nature of liquids, they can flow into the shape of any container. On the other hand, the volume is set because of the forces of attraction that hold the particles in a set amount of space – a set volume.

KMT Solids:

The Three Assumptions of the Solid KMT:1. Solids consist of particles that are closely packed together in a set arrangement that

leads to a set volume and a set shape (the shape is determined by IMFs and chemical bonds).

2. Particles in the solid shape have minimal, vibrational motion due to low KE (the lowest of the three states – solid, liquid & gas).

Structure Type I – Crystal Solid. Structure Type II – Amorphous Solid.

3. The KE of the particles is directly proportional to the Absolute Temperature.

So what does a solid look like?

If you compare the lattice of sodium chloride with therack of 9 ball to the right you will see many similarities.The lattice (above) can be created using billiard balls if they were layered in a three dimensional arrangement.

Define a Solid:A Solid is a state of matter where the particles are arrangedin a rigid structure, there is only minimal, vibrationalmovement among the particles. The Volume of the substance in this state is fixed.

The common lattice shapes are provided in the table below:

Four Subcategories of Solids:There are, among crystalline and amorphous solids, four subsets. Each of these has a common set of properties. The subset that a given sample belongs to can be determined by looking at the combination of elements and the electronegativity differences between them.

A molecular solid is one that consists of atoms or molecules held together by intermolecular forces.  The physical properties of molecular solids vary greatly.  For example, ice melts at 0 oC while sucrose melts with decomposition at 184 oC. 

A covalent network is one that consists of atoms held together in large networks by covalent bonds.  One can think of a covalent network solid as a giant molecule.  In this type of solid inter- and intra-molecular forces are indistinguishable.

An ionic solid is one that consists of ions, cations and anions held together by the electrostatic attraction of opposite charges.

A metallic solid is one that consists of positive cores of atoms held together by a surrounding sea of electrons.  The delocalized electrons are from the outer shells of the metallic atoms. This type of solid also has physical properties that vary over a wide range.

Here is a summary of common properties for each type:

Type of Solid Melting Point Hardness Conductivity

Molecular Low Soft to Brittle Non-conducting

Metallic Varies Variable Hardness Malleable

Conducting

Ionic High to Very High Hard and Brittle Non-conducting Solid Conducting Liquid

Covalent Network Very High Very Hard Usually Nonconducting

Properties of Solids…

Two Distinct Classes of Solids – there are two types of solids that were hinted at above. Type I – Crystalline Solids consist of particles that are arranged in an orderly, geometric, repeating pattern. Compare this to Type II – the Amorphous Solid that consists of particles arranged randomly. Definite Shape & Volume – are determined by the crystal structure and the type of lattice shape that the substance naturally takes on. Coupled with the crystal shape are the forces of attraction that bind the atoms together to create a rigid arrangement of particles even in the amorphous solid as well as the number of particles in the arrangement.

Definite Melting Point – each substance, because of its unique combination of Periodic Properties (atomic radius, electronegativity, valence electrons and IMFs), has a unique amount of energy that must be added to the substance in the solid state to increase the KE to a point where they change to the liquid state. This energy is referred to as the Molar Heat of Fusion or the amount of energy that is needed to melt one mole of any substance at a given temperature and pressure.

High Density and Incompressibility – because of the low KE and the closeness of the particles to one another the arrangement is one where the amount of space between particles is completely minimized. Therefore, the particles cannot be pressed any closer together resulting in the incompressible nature of the solid state. Because of the same reasoning, the solid has the most mass pressed into the least volume.

Part II: Phase Equilibrium (focus on the gas liquid phase change)…

A Phase is any part of a system that has uniform composition and properties.

Equilibrium is a condition where two opposing changes are occurring at

equal rates in a closed system. The opposing changes that we are going to look at are…

Melting & Freezing Evaporation & Condensation Sublimation & Deposition

Let’s look at an example to help explain the concept of equilibrium…if we have a sample of water in a vacuum-sealed container, the water molecules at the surface will begin to break the forces of attraction and enter the gas state. As the number of particles in the gas phase increases the space above the liquid becomes filled with H2O particles behaving as an ideal gas. After a period of time some of the gas particles will collide with the surface and become absorbed into the liquid state again (giving off energy to the environment in the process).

After some time equilibrium is reached between those water particles evaporating and those condensing.

An Illustration of Our Example:

Equilibrium relationships can be established between any two phases and the opposing changes can be any between two states below (one will be endothermic and one will be exothermic when dealing with changes in state)

What an equilibrium equation will look like for changes in state…

H2O (l) + KE H2O (g) & H2O (g) H2O (l) + KE

H2O (l) + KE H2O (g)

The Double-Headed Arrow represents an equilibrium condition between the Liquid and Gaseous states of the water showing that the reaction is reversible.

An Explanation of How Environmental Factors Effect Equilibrium…

Le Chatelier’s Principle: When a system that is at equilibrium is disturbed by

change in conditions, it will adjust to reestablish the equilibrium.

Factors Effecting Phase Equilibrium:

Changes in Temperature Changes in Pressure Changes in Volume Changes in AmountAnalyzing States Under Changing Conditions:

One of the simplest ways to study the states of matter and how they are affected by changes in temperature and pressure is by studying a phase diagram

A Phase Diagram is a graph that shows the effects of changes in temperature

and pressure on the state of matter for a given substance.

A More Typical Phase Diagram

This is not a typical diagram. This diagram for water is very complicated due to its unique chemical and physical properties.

When we refer to “normal” Pressure we are referring to normal atmospheric pressure

Normal Pressure is:

1 atm or 760 mmHg or 101.325 Kpa

The triple point is the temperature and pressure at which all three phases can exist in equilibrium.

Above the critical point, molecules are Unable to liquefy. Any substance that exists beyond the CP is called a Supercritical Fluid.

Uses of the PD:The PD can be used to quickly determine the phase you can expect to find a substance @ under a given set of temperature and pressure conditions.

You can determine the MP/FP & BP/CD points for a substance under any conditions of temperature and pressure.

Heating and Cooling Curves:

Another easy tool used to show the change of state that occurs during addition of energy (heating) or loss of energy (cooling) is a Heating/Cooling Curve

Heating curves show the amount of energy needed to change a substance to a different state and the corresponding temperature at which each change occurs.

Cooling curves show the amount of energy that would have to be lost for the substance to change to a different state as well as the temperature at which the change occurs.

An Example of a Heating & Cooling Curve:

up to A In this section of the graph, the ice is warming up to 0oC. During this phase the ice is still a solid and the energy absorbed from the surrounding is used to increase the kinetic energy of the water molecules. The motion of the molecules at this point is vibrational.

A-B Here is where we see a peculiarity in the heating curve. Why do we have a flat line in the graph for a period of time? The reason the temperature is not rising in this section is that the water is undergoing a phase change from a solid to a liquid and the energy being adsorbed is being used to break the intermolecular bonds. The potential energy of the water is increasing, not the kinetic energy.

B-C Here, again, we have the water absorbing energy and increasing in temperature from 0oC to 100oC. The motion of the water molecules is now rotational and translational as opposed to simply vibrational.

C-D Once again the water is undergoing a phase change, this time from a liquid to a gas. Again, the energy is being used to weaken the bonds so that the water can enter a gaseous state. The Potential energy of the water is increasing.

D-E The kinetic energy of the now gaseous water is increasing as the temperature rises from 100oC to 120oC. The motion of the molecules is now only translational.

A Cooling Curve:

Important Points:

When a plateau appears on a heating or cooling curve the temperature stops changing. This phenomenon occurs when the substance is undergoing a change in state. When Ice melts into liquid water the temperature will remain at 0 Celsius until all the ice is gone and only the single liquid phase remains. At this point the liquid will begin to increase in temperature until it reaches the BP where it will level off again until all the liquid water is changed to gas.

Connections between Energy and State:

The state of the substance is dependent on the motion of its individual particles The motion of the particles is directly dependant on the energy of the particles. The energy of the particles is directly proportional to the temperature and

volume. The energy of the particles is inversely proportional to the pressure. If temperature increases while pressure is constant the energy of the particles

will also increase. If the temperature stays constant but the pressure is increased the energy of the

particles is decreased. Effects of Temperature Change on Phase Equilibrium.

If the change is in the form of a Temperature increase (adding energy to the system) then the endothermic change (evaporation) will be favored (in our example of the sample of water in a closed container, the forward reaction to produce more water in the vapor phase will be favored).

If the change is in the form of a Temperature decrease (removing energy from the system) then the exothermic change (condensation) will be favored (in our example this means that the amount of water vapor will be decreased and more water particles will be in the liquid state).

How are Changes in Temperature calculated?Using a Thermometer or Thermocouple to determine the final and initial temperatures.

Temperature Scales:

Celsius (0C) Kelvin (K) = 0C + 273.15

Examples of Temperatures for Water:

Scale: Absolute Zero MP BP

Celsius -273.150C 00C 1000C

Kelvin O K 273.15 K 373.15 K

Absolute Zero: a theoretical temperature where all molecular motion stops and volume approaches zero BUT mass remains. (0 K or –273.15 0C)

*This is only a theory!!! Volume cannot be eliminated while mass remains.

Effects of Volume Change on Phase Equilibrium.

If we think back on the KMT for liquids and gases we should remember that as we move from the lower energy form of matter to higher energy form of matter the volume will increase. We should also remember that as we move from low to high energy, the temperature rises.

Think on that for a minute and you will come to this: an increase in volume should result in the endothermic phase change and a decrease in volume will result in the exothermic change.

How do we measure Changes in Volume:

1. Use water displacement2. Use a rearranged density equation to solve for V…(D=M/V V=M/D)3. Calculations according to the shape and known volume equations.

Effects of Pressure Change on Phase Equilibrium.

When considering how increasing or decreasing pressure will impact phase equilibrium we need to develop an agreed upon definition for pressure:

Pressure is the measure of the force exerted on a given area

P = F/A (Newton is the standard unit of force = 1kg*m/s2)

What causes the force? The force is caused by the molecules in the states of matter colliding with the walls of the container.

The connection between pressure and state is this: as the pressure in a system increases the lower energy state will form (P up ~ T down) and as the pressure decreases the higher energy state will form (P down ~ T up)

In this section we will discuss atmospheric pressure (of the open environment) and the pressure of a gas in a closed container.

Evangelista Torricelli (1600’s):

Key Conversions:

1000 mL = 1 L

1 mL = 1 cm3

Torricelli was the first scientist to measure pressure of the atmosphere. Began his experimentation on a quest for an answer to why water could not be

pumped more than 34 feet using available pumping systems. He hypothesized that there was a connection between the height that a

substance could be pumped and its density. To study this he designed an apparatus called the barometer.

1 Pascal (pa) = 1kg/m*s2 normal air pressure = 1 kg/cm2

Torricelli's barometer demonstrated that mercury would rise to the same level no matter how tilted the tube holding the mercury became. The pressure of the mercury balances the weight of the air.

From the work that Torricelli did we know that the standard temperature and pressure of the atmosphere is 25 0C and 760 mmHg

Torricelli named the unit of graduation, 1 mmHG, a Torr.

1 Torr = 1 mmHg

The unit “atmosphere” is the standard pressure of the air at sea level

1 atm = 760 mmHg = 760 Torr