Unit 5 Flashcards
Elements that are gases at room temp
H2
O2
N2
F2
Cl2
Br2
These elements make the “number seven” on the upper right-hand corner of the periodic table, with the exception of hydrogen.
(Another class of gases include the oxides of nonmetals such as COx, SOx, NOx.)
The Kinetic Molecular Theory of Gases
1) A gas consists of particles-atoms or molecules that move about randomly in straight lines at high velocities with very weak attractive forces between them.
Random motion of particles allows for easy mixing of different gases, and the random speedy motion (500 m/s = 1000 miles/hr) cause a gas to fill the entire volume of a container.
2. The size of gas particles is very small compared to the distance between each particle. Gases have low densities and most of the volume taken up by a gas is empty space, allowing for the compression of gas particles.
3.Gas particles collide with each other or the container walls without losing energy (elastic collisions). The energy is transferred from particle to another. As particles hit the container walls, they exert pressure. Increasing the force or number of collisions increases the pressure exerted by the gas on the container walls.
4. The average kinetic energy of gas particles is directly proportional to the Kelvin temperature. Gas particles have more energy and move faster as the temperature increases causing them to hit the walls of the container with more force and more often producing higher pressures.
ideal gases
-Gases that obey all of the postulates of the kinetic molecular theory are called ideal gases.
-In reality, no gas is perfectly ideal.
-Most gases can behave “ideally” when the temperature is high and pressure is low.
-When temperatures decrease and atmospheric pressure increases, gas particles get closer together and interactions between them increase causing them to behave “non-ideally.”
Pressure, collisions
The pressure a gas exerts on its container depends on the number of collisions, how forceful, and how often the particles hit the wall
Pressure equation
pressure (P) or force (F) per unit area (A)
P = F/A
-The pressure a gas exerts on its container depends on the number of collisions, how often and how forceful the particles hit the container walls. -Conditions that increase the number of collisions will increase the pressure
atmospheric pressure
-The gas particles in the atmosphere collectively exert a pressure on the earth called atmospheric pressure. -Atmospheric pressure decreases with increasing altitudes, due to gravity
-the density of air is greater near the earth’s surface and decreases with increasing altitude.
-day to day changes in weather affect atmospheric pressure.
-Atmospheric pressure can be measured using a barometer
Units for Pressure
-atmospheres (atm)
-millimeters of mercury (mmHg)
1 mmHG = 1 Torr
-Pressure units in the United States include the pounds per square inch (psi) where 1 atm = 14.7 psi.
-The SI unit of pressure is the Pascal (Pa), where 1mmHg = 133.32 Pa.
Units for Pressure conversions
1 atm = 760. mmHg
=760. Torr
= 14.7 psi
= 101, 325 Pa
= 101.325 kPa
Avogadro’s Law
V1/n1 = V2/n2
Constant P and T
V=volume
N= number of moles of a given sample of gas
-at constant temperature and pressure, gas volume is directly proportional to the number of moles of gas (amount) present.
-Gas volume is directly proportional to the number of moles of gas @ constant P and T
1 mol
=6.02 x 10^23 molecules of gas
=22.4 L
At standard temperature (0°C or 293K) and pressure (1.0 atm) or STP, one mole of a sample of gas takes up 22.4 liters of space.
STP
-Standard temperature: 0°C (273.15 K)
-standard pressure: 1.0 atm (101,325 Pa)
-when a gas is under STP conditions (0°C and 1 atm), its molar volume can be used as a conversion factor to convert between the number of moles and the volume of gas
- At STP: mole of gas = 22.4 L of gas
Molar mass
g/mol
1) check atomic mass on periodic table
2) treat it as g/mol
3) multiply molar mass by how many atoms there are
4) add them together
5) unit g/mol
Boyle’s Law
P1V1=P2V2
-constants T and n
-Gas pressure is inversely proportional to the volume of a given amount (# moles) of gas, at constant temperature.
-Changes in the volume of the container will inversely affect the amount of pressure that the gas exerts on the container
-a decrease in volume will increase the pressure that the gas exerts on the container.
-If the volume of the container increases, the pressure exerted by the gas within the container will decrease as the molecules spread out to fill the space
-alternately, if the pressure exerted on a gas is changed the volume that the gas will try to fill is inversely affected.
-So as P exerted on a gas increases, V decreases.
-Likewise, if the pressure exerted on a gas decreases, then the gas will expand (volume increases).
Charle’s Law
V1/T1 = V2/T2
Constant P and n
the volume of a given amount (# moles) of gas is directly proportional the kelvin temperature of the gas, at constant pressure.
T= kelvin temp
V=volume of the gas
-The temperature of a gas can cause the gas particles to expand when temperatures increase, and condense when the temperature decreases.
-A gaseous sample’s temperature increases as the average kinetic energy (KE), of individual molecules, increases. The temperature will decrease with decreasing KE.
-At higher temperatures, fast moving molecules tend to spread out while slower moving molecules at low temperatures will condense.
Gay-Lussac’s Law
P1/T1 = P2/T2
-constant V an n
-the pressure of a given number of moles of gas at constant volume is directly proportional to its kelvin temperature.
P = pressure exerted by gas
T= kelvin temp
Combined Gas Law
P1V1/T1 = P2V2/T2
Constant= moles of gas
Ideal Gas Law
-elates all four properties—P, V, n, and T—used to measure gases into one mathematical expression, written as
-PV = nRT
R is the ideal gas constant
All gas laws
Dalton’s Law of Partial Pressures
-the total pressure exerted by a mixture of gases is equal to the sum of the partial pressures of each gas in the mixture. The mathematical expression of Dalton’s Law is written as,
Pt= P1+P2+P3+. . .
-the effective concentration of each gas is related to the exertion of pressure by that gas on its container that we call the partial pressure.
-The partial pressure of an individual gas is equal to the total pressure of the mixture multiplied by the mole fraction of that gas.
Determine Partial Pressures
- Determine what you are given and what you need to calculate, make a plan
- Calculation of the partial pressure for each gas based on the percent composition of each gas (represents the mole fraction) given the total pressure of the mixture
- Use Dalton’s Law to check answer
At normal atmospheric pressure, we usually expel most of the N2 we inhale and only small amounts dissolve into our bloodstream, but when youre deep in the ocean where the pressure is very high, the solubility of the gases is higher too, so extra gas is being released into your blood. This causes a problem if the diver ascends too quickly because the quick release of pressure can cause gas bubbles to form in the blood vessels which can have effects like rashes, light-headedness or more seriously paralysis or death.
