Ch10 Solids, Liquids, and Gases and Ch11 The Behavior of Gases (Gas Laws)
Wednesday 2/16/11 Chapter 10
I. The Kinetic Theory – all matter in the entire universe is in constant motion.
A. Kinetic Energy – energy of motion, one half of the mass times the volume
squared. KE = 0.5mv2
II. Solids
A. Properties – particles vibrate in a fixed place (definite shape), cannot compress (definite volume), high density (closely packed, always touching). Analogy: People in their seats in the stands at a football game.
1. Crystalline – orderly geometric repeating pattern
2. Amorphous – without shape, random arrangement of particles (glass, plastic)
B. Melting – Fusion - solid changes to liquid.
1. Increasing temperature increases energy of solid molecules until they break free into liquid phase.
2. Energy is going into the molecules (energy is being absorbed), which is an endothermic process.
C. Freezing – Solidification - liquid changes to solid
1. Decreasing temperature decreases energy of liquid molecules until they slow down and change from faster moving liquid molecules to slower moving solid molecules.
2. Energy is leaving the molecules (energy is being released), which is an exothermic process.
D. Melting point equals the freezing point, MP = FP
1. Same temperature, just moving in the opposite direction as far as energy transfer. Melting = energy in; freezing = energy out
2. Substance remains at constant temperature, in equilibrium between two phases, until phase change is complete.
3. For solid H2O (ice) the MP = FP = 0˚C and the change in heat energy required is ΔHfus=6.01kJ/mol
E. Ionic compounds have higher melting points than molecular compounds.
1. Stronger attractive forces hold ions together, due to complete opposite charges (+)(-).
2. Weaker intermolecular attractions hold polar molecules together, due to their partial charges (q+)(q-).
F. Sublimation – solid changes directly into gas (endothermic)
1. Solid can change directly into gas, skipping liquid phase, if temperature is below substances freezing point. Example: Wet laundry can freeze dry when hanging on a clothesline when the temperature is below 0 ºC
2. Deposition - Gas can move directly to a solid, skipping the liquid phase, the process is dependent on both the temperature and the pressure. (exothermic)
III. Liquids (least common state of matter is the universe)
A. Properties - slide around, change places like marbles in oil (no definite shape), in constant motion (flow, spin, vibrate), cannot compress (definite volume, always touching each other, no space between the molecules which are slightly attracted to each other due to intermolecular attractions like hydrogen bonding). Analogy: People walking in front of the stands at the football game.
1. Diffusion – spontaneous mixing of two substances caused by the random movement of the particles.
2. Surface Tension – a force of attraction that pulls the surface of a liquid together, decreasing the surface area to the smallest possible size. Example: bead of water
3. Volatile Liquids – evaporate more readily than nonvolatile (example: rubbing alcohol verses water)
B. Evaporation – at room temperature, liquid changes to gas.
1. Surface molecules in the liquid gain enough energy from the air particles above the liquid to break free from the liquid state and become gas molecules.
2. This is a cooling process as the warmer fast moving molecules leave the system as gas, and the cooler slower molecules remain as liquid.
3. Energy is going into the molecules (being absorbed), which is an endothermic process.
C. Boiling – at temperatures other than room temperature, liquid changes to gas.
1. Liquid molecules at the bottom of the container gain greater energy than the molecules above them turning into gas molecules and rising through the liquid molecules above them in the form of bubbles.
2. This process is pressure dependent; when the vapor pressure equals the external atmospheric pressure, liquids boil.
3. Energy is going into the molecules (being absorbed), which is an endothermic process.
4. Normal BP – is at standard pressure (1atm=101.3kPa=760mmHg). Example: cooking on a mountain top.
D. Condensation – gas changes to liquid
1. Gas molecules condense when they lose enough energy to slow down and become liquid molecules.
2. Energy is leaving the molecules (being released), which is an exothermic process.
E. Evaporation point equals the condensation point, EP = CP
1. Same temperature just moving in the opposite direction as far as energy transfer. evaporation = energy in; condensation = energy out
2. Substance remains at constant temperature, in equilibrium between two phases, until phase change is complete.
3. For liquid H2O (water) the BP = CP = 100˚C and the change in heat energy required is ΔHvap=40.8kJ/mol
IV. Gases
A. Properties – move randomly (elastic collisions, bounce when hit things like billiard balls, very fast, flow), can be compressed (not attracted to each other, spread out as far as possible to fill space, no definite shape or volume) Analogy: People on the field during the football game.
B. Average Kinetic Energy – of the gas particles in a system is directly proportional to the Kelvin temperature scale.
1. Temperature – the measure of the average kinetic energy of the particles in an object K = ºC + 273
2. Absolute Zero – the temperature (-273.15 ºC = 0K) at which all motion stops; has never been reached.
3. Question: What is the factor change of the average kinetic energy of a gas that changes temperature from 300K to 1200K? …of 800 K to 200K? …of 127˚C to 1227˚C?
C. Gas Pressure –
1. Atmospheric pressure – caused by the collisions of air (gas) molecules.
2. Measured with a barometer.
3. A vacuum would be empty space with NO GAS molecules so no atmospheric pressure. Outer Space
4. Increasing molecule number increases pressure, more collisions, if volume remains constant.
5. Decreasing volume will increase pressure, more collisions.
6. Standard Pressure conversions 1atm = 101.3kPa = 760mmHg
7. Vapor Pressure – gas collisions in a sealed container, liquid going to gas, gas going to liquid, reaches equilibrium when temperature remains constant, increase the temperature will increase the pressure in a sealed container.
D. Ideal Gas – based on five assumptions (example: think of an ideal car)
1. Avogadro’s Hypothesis – equal volumes of gases at same temperature & pressure contain equal numbers of particles. 1mole of gas at STP = 22.4L
2. Gas particles spread out as far as possible (volume is mostly empty space, low density, easily compressed). The volume of one mole of any gas at STP (273K and 1atm) is 22.4L.
3. Collisions between gas particles are elastic, having no net lass in total kinetic energy.
4. Motion is continuous, rapid, and random.
5. No forces of attraction between gas molecules.
6. Temperature depends on average kinetic energy. All gases at the same temperature have the same average kinetic energy.
E. Real Gases – there are no ideal gases (consider the ideal school)
1. Temperature, pressure, volume, and molecular quantities are always changing.
2. The closest real gases to the ideal conditions are the noble gases, helium and neon.
3. The more polar the gas the more it will deviate from the ideal gas behavior
V. Phase Change Diagram
A. Phase Change Concept Map – pg 348
B. Equilibrium – two opposing changes occur at equal rates in a closed system with no net change in the amount of substance in either phase.
a. At the melting point and freezing point temperature the substance moves back and forth between the solid and liquid phase until enough energy is either absorbed (becomes a liquid) or released (becomes a solid).
b. At the boiling point and condensation point temperature the substance moves back and forth between liquid and gas until enough energy is either absorbed (becomes a gas) or released (becomes a liquid).
C. Phase Change Diagram – pg 347 – graph of pressure verses temperature that shows the conditions under which the phases of a substance exist.
a. Triple Point – indicates the temperature and pressure at which all three phases of a substance exist simultaneously at equilibrium.
b. Critical Point – the temperature above which the substance could not exist as a liquid and the lowest pressure at which the substance can exist as a liquid at the critical temperature
c. Normal Temperatures – the normal boiling point, normal melting point, normal freezing point, and normal condensation point are found at Standard Pressure (1atm = 101.3kPa = 760mmHg).
Tuesday 2/22/11 Chapter 11
VI. Graham’s Law of Effusion - heavier items with the same energy move slower than
lighter items, lighter items with the same energy move faster than heavier items.
Gases with lower molar masses effuse faster than gases with higher molar masses.
A. Effusion – The passage of gas molecules under pressure through a tiny opening.
B. The rate of effusion is inversely proportional to the square root of the gases molar mass
C. Formula …NOTE: draw in the square roots over the molar masses!!! RateA = Molar MassB
Molar MassA
Example: two sealed balloons, one filled with helium and the other with air (mostly nitrogen and oxygen gas), which balloon deflates faster over time?
The helium has a molar mass of only 4g as opposed to the N2 with a molar mass of 28g and the O2 with a molar mass of 32g. The lighter gas, helium, more easily escapes through the small pores in the balloon, so it effuses faster shrinking the helium balloon more over the same time period as the balloon filled with air.
Compare the rates of effusion of Helium (gas A) & Nitrogen (gas B)…
NOTE: draw in the square roots!
Rate He = 28.0g N2 = 5.3g = 2.7 Rate of He
Rate N2 4.0g He 2.0g 1 Rate of N2
**Helium gas (smaller) effuses nearly three times faster than nitrogen gas (bigger).
VII. Dalton’s Law of Partial Pressures – at a constant volume & pressure, the total
pressure of a mixture of gases is the sum of the partial pressures of all the gases.
A. Partial Pressure – the pressure of one gas in a combination of gases.
B. Do sample Problem B page 367 Text (requires Table A8 on pg 859).
C. Example: A person’s breath is 18% alcohol, what is the partial pressure of the alcohol if the total pressure is 800mmHg? (not ideal b/c not at STP)
VIII. Changing the Ideal Conditions of a Gas:
A. The effect of adding or removing moles of gas … Changing Quantity
1. Quantity and Pressure rise and fall together at a constant volume –
directly proportional Examples: Tire and an Aerosol can
2. Quantity and Volume rise and fall together at a constant pressure –
directly proportional Example: Balloon
B. The effect of changing the occupied space, liters … Changing Volume
1. Volume and Pressure move in opposite directions –
inversely proportional.
2. Decrease Volume, increases pressure - Example: crowded dance floor
3. Increase Volume, decreases pressure - Example: going outside
C. The effect of changing the heat energy, Kelvins… Changing Temperature
1. Temperature and Pressure move in the same direction –
directly proportional (volume constant)
2. Temperature and Volume move in the same direction –
directly proportional (pressure constant)
IX. Gas Laws – simple mathematical relationships between volume, pressure,
temperature, and amount of a gas
A. Boyle’s Law –Pressure and volume are inversely proportional when mass and
temperature are kept constant.
1. Formula: P1V1 = P2V2 (multiplication)
2. Balloon expands at higher altitudes. Lungs deflate as we apply pressure.
3. Calculation: The Pressure on 2.50 Liters of gas changes from 100 kPa to 40 kPa. What is the new volume if the temperature remains constant?
B. Charles’ Law –Temperature and volume are directly proportional when mass
and pressure are kept constant.
1. Formula: V1 / T1 = V2 / T2 (division)
2. Note: Temperature must be in Kelvin. ** Remember that average kinetic energy is directly proportional to the Kelvin scale (x factor change). K = ºC + 273.
3. Example: Balloon in hot sun verses in air conditioning.
4. Calculation: If a sample of gas occupies 6.8 Liters at 327 ºC, what will its volume be at 27 ºC if the pressure remains constant?
C. Gay-Lussac’s Law –Temperature & pressure are directly proportional if mass
& volume are kept constant.
1. Formula: P1 / T1 = P2 / T2 (division)
2. Note: Temperature must be in Kelvin. K = ºC + 273.
3. Example: Tire will not stretch with energy increase. Aerosol can will
not expand with energy increase.
4. Calculation: A gas has a pressure of 6.58 kPa at 540 K. What will the pressure be at 210 K if the volume remains constant?
D. The Combined Gas Law – write Boyle’s, Charles’, and Gay-Lussac’s gas laws as one formula. At this point all three variables (P, T,V) can change; only mass/moles (quantity) remains constant.
1. Formula: P1V1 / T1 = P2V2 / T2
2. Calculation: If the volume of a gas-filled balloon is 30 Liters at 40 ºC and 150 kPa, what will the volume be at standard temperature and pressure?
X. The Ideal Gas Law – now quantity is added as a variable that can change. Quantity is directly proportional to pressure and volume. Quantity will be expressed in moles (n). Note: If given mass in grams, convert to moles by dividing by the molar mass.
A. The Universal Gas Law Constant: use the ideal conditions to find the constant, R, that never changes.
a. Ideal conditions of 1.0mole of gas at STP (standard temperature of 273K and standard pressure of 101.3kPa) = 22.4L
b. R = 8.31kPa*L/mol*K
C. Formula: PV=nRTAll conditions can change from the ideal!
D. Steps to solving the problems…
1) Balance the Equation.
2) List all given variables from the problem. List the gas law constant.
3) Convert pressure to kPa, temperature to Kelvin, and volume to Liters.
4) Convert mass to moles of the correct compound – use the mole ration
(coefficients) if necessary.
5) Choose the correct formula.
6) Plug (fill in the formula) and Chug (solve your algebra carefully).
7) Convert your answer into the correct units.
Example: Five innocent moles are trapped in a 2 Liter torture chamber by the arch criminal Master Cylinder. If the fiend (boo!) raises the temperature to 500 degrees Celsius (gasp!), what is the pressure?
Example: Calculate the volume of carbon dioxide produced at 200 ºC and 90 kPa, by the combustion of 75 grams of methane gas, CH4.