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Chem.200-5.Gases
General Chemistry Ch. 5 - Gases and the Kinetic-Molecular Theory
| Question | Answer |
|---|---|
| Some important industrial gases; origin and use: Methane | CH4. Origin: natural deposits. Use: domestic fuel |
| Some important industrial gases; origin and use: Ammonia | NH3. Origin: from N2 + H2. Use: Fertilizers and explosives. |
| Some important industrial gases; origin and use: Chlorine | Cl2. Origin: electrolysis of seawater. User: bleaching and disinfecting. |
| Some important industrial gases; origin and use: Oxygen | O2. Origin: liquefied air. Use: steelmaking. |
| Some important industrial gases; origin and use: Ethylene | C2H4. Origin: high-temperature decomposition of natural gas. Use: plastics. |
| While _____ behavior of gases can vary widely, their _____ behaviors are very similar to each other | Chemical; physical |
| Solid/liquid vs. gas: response to pressure changes | Solid/liquid: very resistant to changes in pressure. Gas: gas volume changes greatly with pressure |
| Solid/liquid vs. gas: response to temperature changes | Gas volumes change 50 to 100 times greater than liquids and solids when temperatures change. |
| Solid/liquid vs. gas: viscosity | Gases flow much more freely than liquids and solids. E.g. gases can leak through small holes very easily. |
| Solid/liquid vs. gas: density | Most gases have relatively low densities under normal conditions. Note: when gas is cooled, its volume decreases which means its density increases. |
| Solid/liquid vs. gas: miscibility | Gases are miscible; that is, they mix with one another in any proportion to form a solution. E.g. water and gasoline are not miscible. Gases are 100% miscible with themselves. |
| The reason gas can flow more freely is because | It has a much lower density then the other states. |
| Pressure | Defined as the force exerted per unity of surface area: P = F/A |
| Pressure (in lb/in^2; AKA psi) that atmospheric gases on earth places on the surface of earth | 14.7 PSI (pounds per square inch) |
| Why doesn’t the atmospheric pressure crush us? | The pressure on the outside of our bodies are equalized by the pressure on the inside, so there is no net pressure on the outer surface (analogous to an empty can, when it is sucked dry it gets crushed by the pressure) |
| Why doesn’t the atmospheric pressure crush only downward? | Unlike the other states, gaseous particles are moving at all directions, uniformly, at once. Thus, objects will be crushed at every direction. |
| Example of evolutionary adaptation relating to the concept of pressure | Snow leopards have paws that are very wide so that the weight of the leopard is distributed over a wider diameter. P = F/A. P decreases as the diameter of the paws increases. |
| Barometer | A device used to measure atmospheric pressure. |
| How does a standard mercury barometer work? | 1m tube is filled with Hg. It’s inverted into a dish with more Hg. As the Hg flows into the dish the portion, the top represents a vacuum. It continues to fall until the 760mmHg mark because the weight/area of Hg = atm at that point. |
| Manometer | Devices used to measure the pressure of a gas in an experiment. |
| How does a closed-end manometer work? | An evacuated tube containing Hg, with one end attached to a flask can be used to measure the gas pressure of anything inserted into the flask. Change in the height of the Hg represents the gas pressure. |
| How does an open-end manometer work? | An tube containing Hg, with one end attached to a flask and the other opened up to allow air in, can be used to compare the gas inside the flask to the atmospheric pressure outside. |
| SI unit for pressure | Pascal (Pa), which equals 1 Newton / meters^2 |
| How many Pa is one atm? | A lot. 1 atm = 101.325 kPa – 1.01325 * 10^5 Pa |
| What unit of measurement are barometer and manometers expressed in? | Millimeter of mercury (mmHg), in honor of Torricelli (founder of the barometer) this unit has been named the torr |
| How many Pa in a torr? | 1 torr = 1 mmHg = 1/760atm = 101.325/760 kPa = 133.322 Pa |
| The bar as a unit of measurement | 1 bar = 100 kPa = 1 x 10^5 Pa |
| The physical behavior of a sample of gas can be described completely by… | Four variables: pressure (P), volume (V), temperature (T), and number of moles (n). |
| Three key relationships that exist among the four gas variables | Boyle’s, Charles’, and Avogadro’s laws. Each of these gas laws expresses the effect of one variable on another, with the remaining two variables held constant. |
| The three gas laws are special cases of an all-encompassing relationship among the gas variables known as the | Ideal gas law PV = nRT |
| The unifying observation quantitatively describing gas only refers to… | Gasses that are ideal |
| What is an ideal gas? | One that exhibits simple linear relationships among volume, pressure, temperature, and amount. E.g. N2, O2, H2, and the noble gases |
| Boyle’s observations | The volume occupied by gas is inversely related to its pressure |
| Boyle’s law | P = k/V (k= constant) |
| Charles’ observations | At a constant pressure, the volume occupied by a fixed amount of gas is directly proportional to its absolute (Kelvin) temperature. |
| Charles’ law | T = V/k |
| Avogadro’s Law observations | At fixed temperature and pressure, the volume occupied by a gas is directly proportional to the amount of gas |
| Avogadro’s law | n = V/k |
| Standard conditions | 0deg C, 1 atm, 22.4L |
| Gas density formula | D = m/V |
| At STP (standard temperature and pressure) does 1 mol of O2 have the same density as 1 mol of N2? | No because O2 has a greater mass than each N2 molecule and thus O2 occupies more volume and it is denser. |
| All gases are miscible only when | Thoroughly mixed, otherwise a less dense gas will lie above a denser one. |
| Rearrange PV = nRT so that you can calculate density | We know that mass divided by molar mass gives you number of moles (n) so… PV = (m/M)RT. And we know that m/V = d, so isolate m/v and change to d to get: d = (MolarMass*P/RT) |
| Why is it that heating ducts are often placed near the floor of the room? | Because the density of gas is inversely proportional to the temperature. As the warm air expands, it rises and warms the room more effectively. |
| Rearrange the ideal gas law again, this time to solve for Molar Mass. | MolarMass = dRT/P or MolarMass = (mRT/PV). Use the Dumas method to obtain the values for the variables |
| The Dumas method for determining the molar mass of a volatile liquid | Place liquid in a flask. Flask-> into a water bath with a higher temp than the liquid’s boiling point. Liquid boils into a gas; it exits the top of the flask until the it equals the atmospheric pressure. Cool and get new liquid’s mass. |
| How are all the variables obtained using the Dumas method? | V = premeasured, T = temperature of the heated water bath, P = barometric pressure, m = mass of the liquid after cooled down and re-measured. |
| Dalton’s law of partial pressures | In a mixture of unreacting gases, the total pressure is the sum of the partial pressures of the individual gases: P = P1 + P2 + P3 + … |
| Partial pressure | Each gas in a mixture exerts a partial pressure; a portion of the total pressure of the mixture, that is the same as the pressure it would exert by itself. |
| Mole fraction (X) | Each component in a mixture contributes a fraction of the total number of moles in the mixture, which is the mole fraction (X) of that component. |
| Mole percent | Multiplying the mole fraction (X) by 100 |
| How to calculate partial pressure? | P_A = (X_A)*(P_total) |
| When a gas is in contact with water, the total pressure is… | The sum of the gas pressure and the vapor pressure of water at the given temperature |
| Kinetic-molecular theory | The model that explains gas behavior in terms of particles in random motion whose volumes and interactions are negligible |
| Kinetic-Molecular Theory postulates: Particle volume | A gas consists of a large collection of individual particles. The volume of a particle is extremely small compared with the volume of the container. In essence, the model pictures gas particles as points of mass. |
| Kinetic-Molecular Theory postulates: Particle motion | Gas particles are in constant, random, straight-line motion, except when they collide with the container walls or with each other. |
| Kinetic-Molecular Theory postulates: Particle collisions | Collisions are elastic, meaning the colliding molecules exchange energy but they don’t lose any due to friction. Their total kinetic energy (E_k) is constant. |
| Average speed of particles and how temperature is related | The particles in a given gas will be moving at different temperatures but will always hover near their average, or most probable, speed. As temperature increases, the most probable speed of particle movement also increases. |
| At a given temperature, all gasses have the same | Average kinetic energy |
| The greater the number of molecules in a given container… | The more frequently they collide with the walls, and the greater the pressure is. |
| The fact that liquids and solids cannot be compressed means… | There is little, if any, free space between the molecules |
| Think of a piston in a car, how can you generate more movement out of particles? | By compressing them (increase pressure) and heating them. |
| Summarize the laws that make up the ideal gas law | Boyle’s law: V (proportional to) 1/P; Dalton’s law: P_total = P_A + P_B; Charles’ Law: V (proportional to) T; Avogadro’s law: V (proportional to) n |
| Why do heavy particles occupy the same volume as lighter particles? | E_k = 1/2*Mass x u^2 where u = speed. If a heavy object and a light object have the same kinetic energy, the heavy object must be moving more slowly. Thus higher mass = lower speed allows for volume as lighter mass. |
| H2 compared to O2: same amount of molecules, describe how they’re the same volume | H2 moves fast but its collisions yield less force. O2 is slower but its collisions yield more force. |
| Kinetic energy equation incorporating temperature | E_k(avg) = 3/2*(R/N)T where R is the gas constant and N is Avogadro’s number |
| Combine kinetic energy incorporating temperature and speed | u = sqrt((3RT/MolarMass)). Where the square root of u = the rma (root-mean-square) speed. A molecule moving at this speed has the average kinetic energy. Remember to express the MolarMass in Kg/mol |
| Root-mean-square speed of an O2 molecule | u = sqrt((2*8.314)(293)/3.2e-2)) = 478 m/s |
| Effusion | The process by which a gas escapes from its container through a tiny hole into an evacuated space |
| Graham’s law of effusion | The rate of effusion of a gas is inversely proportional to the square root of its MolarMass. Equation: rate of effusion (proportional to) 1/sqrt(MolarMass) |
| Which effuses faster, Kr or Ar? | Ar because it is lighter. Specifically, the gas with the lower molar mass effuses faster because the most probably speed of its molecules is higher. |
| Graham’s law used as a ratio | Rate_X/Rate_He = sqrt(MolarMass_He/MolarMass_X); use this ratio to find unknowns |
| Diffusion | The movement of one gas through another. Equation: Rate of Diffusion (proportional to) 1/sqrt(MolarMass) |
| Mean free path | The average distance the molecule travels between collisions at a given temperature and pressure. |
| Collision frequency | The average number of collisions per second that each molecule undergoes: (most probable speed / mean free path = collision frequency) |
| Starting from closest to earth, list layers of the atmosphere | Troposphere (airplanes fly within this layer), stratosphere, mesosphere, thermosphere, exosphere |
| Even though it may be 2000K in the exosphere, why would it still feel cold? | Because the density of the air is much lower; there is very little transfer of kinetic energy because very little collision frequency. |
| In terms of chemical composition, describe the atmosphere | It’s classified into two major regions: the homosphere and the heterosphere. |
| Homosphere | Comprised of 78% N2 and 21% O2, and 1% other gases (mostly Ar). |
| Wind | Caused by a continues cycle of warm air near the ground expanding (and density decreasing) and rising above the cool air which then drops down, warms, expands, rises, and the cycle continues |
| Heterosphere | Includes the ionosphere, containing ionic species such as O+, NO+, O2+, N2+, and free electrons. Photons break up (photodissociation) molecules and move electrons (photoionization) causing these results |
| What is the importance of the stratospheric ozone? | It is the region where O2 molecules are broken up by high energy radiation, and then reform into O3 molecules. O3 then has the capacity to absorb UV radiation which decomposes it back to O2. |
| CO2 on Earth during the era of the first life on the planet | CO2 was very high. Algae would utilize the CO2 in biosynthesis, and excrete O2 (recall: photosystem II). O2 in the atmosphere increased as a result leading the way for evolved O2-based life. |
| The point at which a gas liquefies | Condensation point |
| Why do gases begin to deviate from the ideal gas laws? | At moderately high pressure, values of PV/RT lower than ideal: due primarily to intermolecular attractions. At very high pressure: values of PV/RT greater than ideal are due primarily to molecular volume |
| Deviation from ideal gas laws: Intermolecular attractions | As pressures increase, intermolecular attractions begin to have less dominance over the attractions to surrounding molecules due to being more closely arranged with them |
| How do the intermolecular attraction issues affect gas pressure | Molecules that are heading toward the walls might be more attracted to other molecules nearby which will decrease the force and thus decrease gas pressure, thus a smaller numerator in the PV/RT ratio. |
| Deviation from ideal gas law: Molecular volume | At very high pressures, the free volume < container volume. P < V so PV/RT becomes artificially high. |
| Van der Waals equation | Johannes van der Waals realized the limitations of the ideal gas law and created his van der Waals equation which accounts for the behavior of the ideal gases. Includes Van d. w. constants: a & b. Look up formula in book. |
| Van der waals constants in ideal gases | They are zero for ideal gases because the particles don’t attract each other and have no volume |
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