Chemistry - Chapter 11

diffusion of gases

process where gas molecules spread out from an area of higher concentration to an area of lower concentration (usually there is a gradient in concentration)

what factors affect the diffusion of gases

temperature (higher temperatures increase diffusion rates bc by increasing KE, they can move faster, collide more frequently, and spread out more rapidly), concentration gradient (higher difference in particle concentration = faster net movement from high to low & a steeper gradient means molecules are densely packed in the high concentration area, which means more frequent collisions that cause them to spread out and move away faster), and size of the gas molecules (KE = 1/2mv², so if mass is small, velocity has to be big to make up for it since all gas particles have the same KE at a certain temperature, Graham’s Law, & smaller particles are less likely to collide with others which allows for a faster, straightforward path)

effusion

process where gas particles escape from a container through a tiny opening into a vacuum or another area of lower pressure that occurs without significant collisions among gas molecules

what is an ideal gas

a theoretical gas that perfectly follows the ideal gas law, which states that the pressure, volume, and temperature of a gas are related in a predictable way; they have no interactions/attractions between their particles and have no volume

what are the properties of an ideal gas

no intermolecular forces, particles have little to no volume, collisions between particle are elastic, particle movement is random and rapid, and the gas behaves according to the ideal gas law (PV=nRT)

real gases

gases that do not perfectly follow the ideal gas law due to intermolecular forces and finite volume of their particles; all gases are real gases; most gases behave ideally or near ideal conditions until extremely low temperature or extremely high temperature (where they exhibit interactions/attractions between particles and occupy significant volume); they can condense into liquids

most gases behave like ideal gases, even though all gases are “real” gases, which means they

have no significant volume and little attraction for other molecules

barometer

used to measure atmospheric pressure; the height of the Mercury column in the tube changes with atmospheric pressure (higher pressure pushes the mercury column higher, while lower pressure allows it to drop); uses pascal’s law

Pascal’s law

when pressure is applied to a confined fluid, the pressure change is transmitted equally throughout the liquid in all directions

KMT - Pascal’s law

air pressure pushes a column of Hg up into a vacuum until an equilibrium is reached between the air pressure pushing up and the weight of the mercury column pushing down

gas pressure

force per unit of area (N/m² = Pa); related to temperature, volume, and number of gas particles, resulting from collisions of gas particles with the walls of a container

how does increasing the number of particles increase pressure

As the number of gas particles increases, collisions with the walls of the container become more frequent, resulting in a higher force exerted per unit area, thus increasing the gas pressure.

how does increasing temperature increase pressure

Increasing temperature causes gas particles to move faster (higher KE), leading to more frequent and harder collisions with the container walls, which increases the force exerted per unit area and therefore raises the gas pressure.

what is a manometer

A device used to measure gas pressure, typically consisting of a U-shaped tube filled with liquid. The difference in liquid height between the two arms indicates the pressure of the gas.

closed tube manometer

find the difference of levels if not given, then that is your gas pressure bc according to the KMT, closed-tube manometers have equilibrium between the gas pressure pushing mercury up into the vacuum and the weight of the mercury column pushing down

open-tube manometer

A device used to measure gas pressure where one arm is open to the atmosphere. The difference in liquid height between the two arms reflects both the gas pressure and the atmospheric pressure.

what should you do if the mercury levels are even for an open-tube manometer

It indicates that the gas pressure is equal to atmospheric pressure, meaning there is no pressure difference.

what should you do if the mercury level is higher on the gas side, and what does it say about the relationship between the gas pressure and air pressure

The gas pressure is less than the atmospheric pressure, so you should subtract the mercury difference from the atmospheric pressure to find the gas pressure.

what should you do if the mercury level is higher on the air side, and what does it say about the relationship between the gas pressure and air pressure

The gas pressure is greater than the atmospheric pressure, so you should add the mercury difference to the atmospheric pressure to find the gas pressure.

KMT - open-tube manometer

equilibrium between atmospheric pressure, gas pressure, and the weight of the mercury column pushing down (also higher KE = higher force and collisions)

what are the four variables that allow us to completely describe a gas sample

temperature in K (a measure of the average KE of gas particles & affects how fast gas particles move and how often they collide), pressure in kPa/atm/torr/mmHg/psi (the force exerted by gas particles colliding with the walls of a container), volume in L/mL/cm³/dm³ (the space that the gas occupies & is equal to the volume of its container), and n in moles (the quantity of gas present)

state boyle’s law

Boyle's law states that at constant temperature and number of gas particles, the pressure of a gas is inversely proportional to its volume. Mathematically, this can be expressed as P1xV1 = P2xV2

KMT - Boyle’s Law

Gas pressure is caused by the force of the many collisions of gas particles with their container walls; if the volume of the container is decreased, the same number of collisions must take place, but in a small area, which means pressure must increase bc force per unit area increases (as the volume decreases, the pressure increases, and vice versa)

state Charles’s Law

at constant pressure and number of gas particles, the volume of a gas is directly proportional to its temperature in Kelvin. Mathematically, this can be expressed as V1/T1 = V2/T2

KMT - Charles’s Law

As temperature goes up, particles speed up and hit the walls of the container harder and more often, increasing the force; in order to keep pressure constant, the increased force should be spread over a greater area, which means volume must increase

state Gay-Lussac’s Law

At constant volume and number of gas particles, the pressure of a gas is directly proportional to its temperature in Kelvin. This relationship can be expressed as P1/T1 = P2/T2.

KMT - Gay-Lussac’s Law

As temperature goes up, particles speed up and hit the walls of the rigid container harder and more often, increasing the force and therefore the pressure (to force pressure up, energy has to be added through an increased temperature & bc pressure must rise to keep volume from increasing)

combined gas law

The combined gas law relates the pressure, volume, and temperature of a fixed amount of gas, combining Charles's Law, Boyle's Law, and Gay-Lussac's Law, typically expressed as P1xV1/T1 = P2xV2/T2

state Avogadro’s Law

At constant temperature and pressure, the volume of a gas is directly proportional to the number of moles of the gas, expressed as V1/N1 = V2/N2

R constants in Ideal Gas Law (PV=nRT)

0.08206 when looking at atm & 8.314 when looking at kPa

What are the two variations of the Ideal Gas Law, regarding molecular/molar mass and density

PV = mRT/M & PV = dRT/M (remember density is m/v)

State Dalton’s Law of Partial Pressure

The total pressure of a mixture of gases is equal to the sum of the partial pressures of the individual gases, expressed as PT = P1 + P2 + P3 + … (each gas has its own collisions, but the overall kinetic energy is the same)

KMT - Dalton’s Law of Partial Pressure

Gas particles are far apart and have few mutual attractions and are therefore independent of each other; each gas has its own collisions with the container walls, therefore exerting its own partial pressure

Collecting gases over water

Involves the process of gathering gases in an inverted container where water is present, requiring adjustments for the vapor pressure of water in calculations by correcting the provided pressure to only the pressure of dry gas (PT = PH20 + PDry Gas); total pressure is often room temperature (or atm)

State Graham’s Law of Effusion/Diffusion

The rate of effusion or diffusion of a gas is inversely proportional to the square root of its molecular/molar mass