Name:____________________________ Unit 8: Phases Part 1

1. Properties of phases

Positive and negative charges attract each other. Attractive forces between molecules are called intermolecular attractive forces. The strength of these forces determines what phase of matter a substance is in at a given temperature. Substances with weak attractive forces (London dispersion) between their molecules tend to be gases at room temperature, and substances with strong attractive forces (ionic) tend to be solids at room temperature. Phases of matter are simply stages of attraction. Gases are made of molecules with no attractive forces, allowing the molecules to spread out and fly freely past each other, taking the shape and volume of their container. Liquids are made of molecules with stronger attractive forces, allowing the molecules to flow past each other, but still stay close together and take the shape of their container. Solids are made of molecules or ions with strong attractive forces, which lock the molecules into place, forming a crystal lattice, where the particles are free to vibrate in place, but they cannot move relative to each other.

a. Particle diagrams

Using circles to represent particles, draw what each phase of matter would look like according to the description above

Solids Liquids Gases

2. Types of solids a. Ionic solids : Ionic solids are solids in which anions and cations (negatively and positively

charged atoms or groups of atoms, respectively) stick together via “electrostatic attraction.” What this basically means is that the opposite charges just like to stick to each other. When they do this, they form great big crystals in which each ion is surrounded by ions with the opposite charge. Such crystal lattices (as they’re called) are really stable, requiring lots of energy to pull apart. Examples of ionic solids include sodium chloride (table salt) and Epsom salts (magnesium sulfate heptahydrate). These

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types of solids can dissolve in water, and when they do, will conduct electricity because of there being charges that can move around.

b. Metallic solids: Metallic solids are solids in which the positively charged nuclei are held together by a bunch of valence electrons that kind of bind the whole mess together. These electrons are referred to as being “delocalized” because they don’t stay between two atoms as they do in covalent bonds – instead, they travel throughout the whole solid. Just about anything you know of that’s a metal does this kind of bonding. This is why metals conduct electricity; there are charges (electrons) that are mobile. Metallic solids AND liquids will conduct electricity.

c. Network solids : Network solids are great big crystals in which all of the atoms are stuck together using covalent bonds. Because the atoms are all locked into place, these solids usually have properties very similar to that of ionic compounds (high melting and boiling point, hard, brittle, and so forth) with the exception that they don’t conduct electricity if you melt them, nor will they dissolve into water. Typically, gemstones (such as amethyst, diamond, and ruby) and diamonds are network atomic solids.

d. Molecular solids : Molecular solids occur when covalent molecules are held together by intermolecular attractive forces. In this type of bonding, which occurs in ice and wax, the intermolecular forces between the molecules are strong enough to keep the molecules locked into place. Typically, these types of solids have much lower melting and boiling points than metallic, network atomic, or ionic solids, because the intermolecular forces holding the molecules together are much weaker than those of the bonds in the other compounds. These solids will not conduct electricity, even if they are dissolved in water, because there are no mobile charges to conduct the electricity.

Network SolidIonic Solid Metallic Solid Covalent Solid

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3. Phase change a. What is happening on a molecular level

The molecules in a solid phase can be described as being locked into place and vibrating. As the temperature increases, the particles vibrate faster, with the attractive forces still keeping them locked into place. Upon adding more energy, eventually the particles will vibrate with enough force to break away from being locked into place. This process is known as melting. The liquid particles are still attracted to each other, but have enough energy to flow past one another as they move. If you were to add more energy, these particles would move faster. If you continued to add more energy, eventually the particles would have enough energy to overcome the attraction to each other and they would spread out, turning into a gas. They would take on the volume and shape of their container and have no attraction for each other. This would occur at the boiling point. If energy were taken away from these particles, the opposite process would occur, first condensing and then freezing.

Sometimes particles can change directly from the solid phase to the gas phase and vice versa. Sublimation is when a solid goes directly into the gas phase, skipping the liquid phase. Solid Carbon dioxide will do this and is known as ‘dry’ ice because it doesn’t get ‘wet’, it goes directly into the gas phase. The opposite change, gas directly into solid is known as deposition and can be seen on car windshields during the winter

b. Phase change diagrams

Phase change diagrams are graphical representations of how a substances temperature changes while energy is being added or subtracted from the system. An interesting thing happens while particles are going through phase change: their temperature stays the same until it completely changes phase! Let’s say that you were melting some ice. As you added energy to the solid water (ice), its temperature would increase until it reached its melting point. If you continued adding energy, the temperature would stay the same, but the ice would start to melt. The reason for this is that as you add energy to the solid ice, it first goes to speeding up the vibrations of the solid particles. Once it reaches its melting point, the additional energy is going to go into breaking apart the particles from each other, and not to speeding up the vibrations. The same thing will happen at its boiling point, which results in two characteristic plateaus on a phase change diagram.

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Temp. (K)

Time

Temp. (K)

Time

6.___________5.___________

4.___________

3.___________

2.___________

1.___________

+ or - ΔH+ or - ΔH_____thermic

_____thermic

(particle diagram)

Solid

(particle diagram)

Liquid

(particle diagram)

Gas

Phase Changes: Label the phase changes (1-6), identify each side as either exothermic or endothermic, fill in the phase change diagrams and particle diagrams for each phase.

1)boiling 2)condensing 3)melting 4) freezing 5) sublimation 6)deposition

1) Label the phases present at each line segment above using (s), (l), and (g).

2) What is the boiling point of this substance? ________

3) What is the melting point of this substance? ________

c. Heat of phase change

In previous units we have calculated the amount of energy needed to change the temperature of water using the formula: q=mCΔT (remember that catchy song?). During a phase change, however, there is no change in temperature so this formula cannot be used to solve for the heat needed for a phase change. Instead, there are different formulas used for phase changes; q=mHf and q=mHv. Hf is the heat of fusion, the amount of energy required to melt or freeze a gram of water (334 J/g). Hv is the heat of vaporization of water; the amount of energy required to boil or condense water (2260 J/g).. ‘m’ is the mass of water involved in the phase change.

Ex: How much energy is required to melt 50.0g of water at 0oC?

q = m Hf

q = 50.0 g x 334J/g

q =16700 J

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Practice Problems

_____1) Which of the following phase changes requires heat of fusion to accomplish?a) H2O (s) H2O (g) b) H2O (g) H2O (l) c) H2O (l) H2O (g) d) H2O (s) H2O (l)

_____2) Which of the following phase changes is endothermic?a) H2O (s) H2O (l) b) H2O (g) H2O (l) c) H2O (l) H2O (s) d) H2O (g) H2O (s)

Calculate the number of joules required to (show correct numerical setup):

a) melt 20.0 g of H2O (s) at 0oC

b) boil 30.0 g of H2O (l) at 100oC

c) freeze 200.0 g of H2O (l) at 0oC

d) boil 50.0 g of H2O (g) at 100oC

Ice cubes at –12.0oC are placed in a saucepan and heated at a constant rate over a stove to

115.0oC.

Sketch a phase change diagram for the phase changes that occur between -12.0 oC and 115.0oC. Label the temperatures at which the phase changes occur. Then label each line segment with a letter (A, B, C, D, E, etc.).

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4. Gases and pressure a. Kinetic molecular theory (KMT) for gases and ideal gas laws

Before we can discuss what an ‘Ideal’ gas is, we must first discuss what is true for all gases, ideal and real. The following is true for all gases:

Gas particles have no definite volume Gas particles have no definite shape Gases are highly compressible Gases take the volume and shape of the container Gas molecules are relatively far apart from one another Gases form homogeneous mixtures with each other

The concept of the ‘Ideal gas’ is to explain the behavior of real gases. There is a list of things that we assume about gases to be true to explain their behavior. For example, we assume that ‘ideal’ gas particles have no attraction for each other, which is why they take on the volume of their container. ‘Real’ gas particles do have some attraction to each other. The reason we make these assumptions is because they are mostly true, and in doing so we can calculate a lot of information about gases. The following are the Ideal Gas Laws:

1. Gas molecules are so small that the combined size is insignificant compared to the volume occupied by the gas.

2. Gas molecules move in straight line motion until they collide with the container wall or another gas molecule

3. Any collisions between gas molecules are elastic with NO energy lost from the collisions.

4. No attractive or repulsive forces exist between gas molecules.5. The Average velocity (speed with direction) of the gas is directly proportional to the

KELVIN temperature (the higher the temperature, the faster they move).

Real gases will act like Ideal gases at low pressure and high temperature. Smaller gases (H2 and He) also behave more ideally than larger gases (CO2, CH4).

b. Avagadros hypothesis

Simply put, Avogadro’s Hypothesis states that equal volumes of ALL gases, at the same temperature and pressure, have the same number of molecules. If you were to double the number of molecules, you would double the volume.

Visualizing these concepts: Using your creativity, come up with drawings to illustrate the 5 ideal gas laws and Avogadro’s hypothesis.

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Ideal Gas Law 1 Ideal Gas Law 4

Ideal Gas Law 2 Ideal Gas Law 5

Ideal Gas Law 3 Avogadro’s Hypothesis

c. Pressure, vapor pressure and boiling point

Pressure is defined as a force exerted on an area. You may be familiar with the term psi, which stands for pounds per square inch. This is what you see on your gas gauge when

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you check the pressure in your tires of your car or bicycle and measures how much force is being exerted on every square inch of your tire’s surface. The force itself is a result of the gas particles colliding with the surface of the container. There are other units to measure pressure and in chemistry we will either use atmospheres (1 atmosphere of pressure is the amount of force exerted by the atmosphere above use, about 14.7 psi), or kilopascal named after a French mathematician and physicist (1 atmosphere = 101.3 kPa).

Particles in the liquid phase, in a closed container, may evaporate (go into the gas phase while below the boiling point) and exert a pressure. We call this vapor pressure, as it is the pressure caused by the vapor and it does not depend on how much liquid is in the container.

i. Normal boiling point

Particles in the liquid phase, in a closed container, may evaporate (go into the gas phase while below the boiling point) and exert a pressure. We call this vapor pressure, as it is the pressure caused by the vapor and it does not depend on how much liquid is in the container. A substance will boil when it is heated to the point that its vapor pressure is equal to that of its environment. The normal boiling point is the temperature at which the vapor pressure is equal to standard pressure. Standard pressure has been defined as 1 atmosphere (101.3 kPa) and can be found in your reference tables under the Table A titled ‘Standard Temperature and Pressure’. If you change the pressure of the system a liquid is in, you also change its boiling point. This is how you can get LP (liquid propane) gas tanks for your grill. Propane, at standard temperature, will be in its gas phase. In order to store it in its more condensed liquid phase, the pressure needs to be increased. This is why they are stored in pressurized tanks. When you open the tank, you let some of the propane out, which is then exposed to a much lower pressure in the grill, becomes a gas which is then ignited to cook some hot dogs, hamburgers, chicken, maybe some corn…whatever your little heart desires! It’s really up to you.

Pressure cookers also make use of the impact of pressure on boiling points. A pressure cooker doesn’t allow for gas to escape as you’re cooking your meal. This causes the pressure inside to build up, which increases the temperature at which the water will boil. This higher temperature allows the food to cook faster!

Using the reference tables

Reference Table H is titles Vapor Pressure of Four Liquids and can be used to determine the boiling points of these liquids at different temperatures. Remember, the boiling point is the temperature at which the vapor pressure equals the pressure of the system.

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What are the normal boiling points for:

Propanone:____________

Ethanol:_____________

H2O:________________

Ethanoic Acid:_____________

What would the boiling point of water be at a pressure of 200 kPa?__________

What would the vapor pressure be for propanone at 50 degrees celcius?____________

Which of these 4 liquids has the strongest attractive force?_________________________

5. Gas laws: Pressure, temperature and volume a. Individual gas laws: P+T. T+V, V+P

The Gas Laws are relationships between temperature, pressure and volume of a gas. Gas law equations are used to determine what effect changing one of those variables will have on any of the others. There are three gas laws that combine to give us our ‘combined gas law’, and they are:

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1) Boyles Law : P1V1=P2V2: As pressure on a gas increases, the volume of the gas decreases. The product of the initial pressure (P1) and volume (V1) will be equal to the product of the final pressure and volume (P2V2)

2) Charles Law : V1/T1 = V2/T2: As temperature (in kelvin) of the gas increases, the volume of the gas must also increase. The quotient of the initial volume and temperature (V1/T1) is equal to the quotient of the final volume and temperature (V2/T2).

3) Gay-Lussac’s Law: P1/T1 = P2/T2: As the temperature of a gas increases, so does its pressure. The quotient of the initial pressure and temperature (in kelvin) (P1/T1) is equal to the quotient of the final pressure and temperature (P2/T2).

A neat way to remember how these three factors are related is by what has been trademarked as a ‘PT Cruiser’ card. Simply take an index card and write across it the letters P T V. To see how one factor will be affected by changing one of the other two, simply grasp the card with your finger covering the letter for the factor being kept constant, push the variable that is being changed in the direction in which it is changing (up for increasing, down for decreasing) and see how the other variable responds. For example: increasing pressure while temperature is kept the same. You would cover grab the card at the temperature letter, push the ‘P’ up, and see how the volume would decrease!

b. Combined Gas Law

The three individual gas laws combine to give us the combined gas law, which allows us to examine how all three variables respond to a change in a system without needing to have one of the variables kept constant (seen on the right). You can also get the other three gas laws from this combined formula. If one of the variables is kept constant, it simply drops out of the equation. For example: If temperature was kept the same, T would drop out and you would be left with: P1V1=P2V2.. An important note; ALL TEMPERATURES MUST BE IN KELVIN!

Practice: (remember temperature needs to be in kelvin).1.) What volume will 500.mL of a gas occupy if the pressure is changed from 1.00 atmosphere to

2.00 atmosphere at constant temperature?

P1 = P2=

V1= V2=

T1 = T2=

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2) A tube of hydrogen gas at a room temperature of 22.4ºC has a pressure of 88.0 KPa. What was the new temperature of the hydrogen gas when the pressure in the tube is at 101.3 KPa? (Volume is constant!)

P1 = P2=

V1= V2=

T1 = T2=

3) A gas has a volume of 700.mL at a temperature of 10.0ºC at constant pressure. What volume will the gas occupy if the temperature is raised to 50.0ºC? (Remember: change °C to K!!)

P1 = P2=

V1= V2=

T1 = T2=

4 .) A sample of gas occupies a volume of 500. mL at a pressure of 0.500 atm and a temperature of 298K. At what temperature will the gas occupy a volume of 250. mL and have a pressure of 2.50 atm?

P1 = P2=

V1= V2=

T1 = T2=

6. Honors: Gas laws; pv=nRT

The pressure and volume of a gas are proportional to the number of moles of gas and the Kelvin temperature. The equation can be derived as follows:

From the Combined Gas Law, we have

PV/T = K, K = nR (n is number of moles, R is a proportionality constant)

Since one mole of gas exerts a pressure of 1.00 atm and occupies a volume of 22.4 L at 273 K, R (the proportionality constant) can be derived as follows:

(1 atm)(22.4 L)/(1 mole)(273 K) = R13

R = 0.0821 atm-L/mol-K

This yields the IDEAL GAS LAW, which can be used to determine the pressure, volume, temperature or number of moles of gas if all of the other conditions are known, and none of the conditions have changed.

PV = nRTP= Pressure (atm) V = Volume (L) n = moles R = 0.0821 atm-L/mol-K T = Temp (K)

a. Practice with ideal gas law 1) What is the pressure exerted by 3.00 moles of gas at a temperature of 300. K in a

4.00 L container?

2) What is the volume of a sample of gas if 5.00 moles if it exerts a pressure of 0.500 atm at 200. K?

3) A sample of gas contained in a cylinder of 5.00 L exerts a pressure of 3.00 atm at 300. K. How many moles of gas are trapped in the cylinder?

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