Chemistry: The Molecular Nature Of Matter And Change Chapter 5

Front 1

Common units of Pressure

Back 1

Pascal (Pa)

kilopascal (kPa)

atmosphere (atm)

millimeters of mercury (mmHg)

torr

pounds per square inch (lb/in2 or psi)

bar

Front 2

Pascal (Pa)

kilopascal (kPa)

(at sea level and 0* C)

Back 2

1.01325 x 10^5 Pa

101.325 kPa

Front 3

Atmosphere (atm)

(at sea level and 0* C)

Back 3

1 atm

Front 4

millimeters of mercury (mmHg)

(at sea level and 0* C)

Back 4

760 mmHg

Front 5

Torr

(at sea level and 0* C)

Back 5

760 torr

Front 6

pounds per square inch

(lb/in2 or psi)

(at sea level and 0* C)

Back 6

14.7 lb/in2

Front 7

bar

(at sea level and 0* C)

Back 7

1.01325 bar

Front 8

Properties of Gases

Back 8

1. Volume (V) changes with Pressure (P)

2. Volume (V) changes with Temp (T)

3. Gases flow very freely

4. Gases have relatively low Density (D)

5. Gases form solution in any proportions

Front 9

Pressure = ?

Back 9

P = Force / Area

Front 10

when

Temperature Increases

Back 10

Volume Increases,

Density Decreases,

Pressure Increases

Front 11

when

Temperature Decreases

Back 11

Volume Decreases,

Density Increases,

Pressure Decreases

Front 12

when

Pressure Increases

(at fixed T and n)

Back 12

Volume Decreases

Front 13

1 torr = ?

Back 13

1 mm

Front 14

Boyles' Law

Back 14

At a constant temp (T), the volume (V) is inversely proportional to applied external

pressure (P).

P Increases, V Decreases

P Decreases, V Increases

Front 15

Charles' Law

Back 15

At a constant pressure (P), the volume (V) is directly proportional to its absolute temp (T) (Kelvin).

T Increases, V Increases

T Decreases, V Decreases

Front 16

Avogadro's Law

Back 16

At a fixed temp (T) and pressure (P), equal volumes (V) contain equal # particles (moles, n).

Front 17

STP

(Standard Temperature and Pressure)

Back 17

0* C

(273.15 K)

1 atm

(760 torr)

Front 18

Standard Molar Volume

(volume occupied by 1 mol of ideal gas at STP)

Back 18

22.4141 L

(22.4 L [to 4 sf])

Front 19

Ideal Gas Law

Back 19

PV = nRT

(T is always in Kelvins)

Front 20

Solving Gas Law Problems

Back 20

1. summarize gas variables and constants

2. convert units (if necessary)

3. rearrange ideal gas law and solve for

unknown

Front 21

Boyle's Law Equation

Back 21

PiVi = PfVf

(fixed n & T)

(i = initial, f = final)

Front 22

Charles' Law Equation

Back 22

Vi/Ti = Vf/Tf

(fixed n & P)

(T always in Kelvins)

(i = initial, f = final)

Front 23

Avogadro's Law Equation

Back 23

Vi/ni = Vf/nf

(fixed P and T)

(T always in Kelvins)

(i = initial, f = final)

Front 24

Ideal Gas Law

(Initial conditions to Final conditions)

Back 24

P1V1 / n1T1 = P2V2 / n2T2

(T always in Kelvins)

Front 25

R = ?

Formula

Back 25

R = PV / nT

Front 26

R = ?

(using STP)

Back 26

(.0821 L kPa/K mol)

Front 27

when

Pressure Increases

(at fixed V)

Back 27

Temperature Increases

(P1 / T1 = P2 / T2)

Front 28

R

indicates:

which may have different values based on:

Back 28

Universal gas constant

Units for Volume and Pressure

Front 29

Van der Waals Equation

Back 29

(P+ n^2 a/ V^2)(V-nb) = nRT

Front 30

Density Formula

Back 30

d = PM / RT

(M = molar mass)

Front 31

Density is directly proportional to:

Back 31

Molar Mass

Higher Molar Mass, Higher Density

Front 32

Density is inversely proportional to:

Back 32

Temperature

Temp Increases,

Volume Increases,

Density Lower

Front 33

Density from temp1 to temp2:

Back 33

d1 / d2 = T2 / T1

(inversely proportional)

Front 34

Molar Mass of a Gas Formula

Back 34

M = mRT / PV

Front 35

Dalton's Law of Partial Pressures

Back 35

In a mixture of unreacting gases, the total pressure is the sum of the partial pressures of the individual gases.

(P(total) = P1 + P2 + P3 + ...)

Front 36

Mole Fraction (X)

Back 36

Each component in a mixture contributes a fraction of the total # of moles in the mixture.

(sum of all fractions must = 1)

Front 37

Pressure of each Gas in Mixture

Back 37

P(A) = [(mole of A) / mole total] x P(total)

[P(A) = Pressure of gas 'A']

[mole total = total mole of mixture]

Front 38

mole % formula

Back 38

X(mole%A) = [P(A) / P (t)] x 100

[P(A) = pressure of element A]

[P(t) = total pressure of mixture]

Front 39

Kinetic-Molecular Theory

Postulates

Back 39

1. Particle volume (indiv. particles in empty space)

2. Particle motion (constant, straight-line motion)

3. Particle collisions (constant total Kinetic energy)

Front 40

Gas Compression vs. Liquid & Solid Compression

Back 40

The fact that liquids and solids cannot be compressed implies there is little, if any, free space between their particles.

Front 41

Kinetic Energy Equation

Back 41

Ek = ½ mass x speed^2

Front 42

Graham's Law of Effusion

Back 42

The rate of effusion of a gas is inversely proportional to the square root of its molar mass.

Front 43

Mean free path

Back 43

The average distance a particle travels between collisions at a given Temp and Pressure obtained by particle's diameter.

Front 44

Collision Frequency

Back 44

most probable speed / mean free path

(meters/second) / (meters per collision)

Front 45

Effusion and diffusion rates are _________ proportional to the square root of the molar mass because they are __________ proportional to molecular speed.

Back 45

inversely , directly

Front 46

Gay-Lussac's Equation

Back 46

Pi/Ti = Pf/Tf

(V and n are constant)

(i = initial, f = final)

Front 47

Barometer measures ________

Back 47

atm pressure

Front 48

Manometer measures ________ involved in _______.

Back 48

pressure of gas,

chemical reaction

Front 49

Diffusion is the movement of _________ through ________.

Back 49

one gas

another

Front 50

Effusion is the _______ by which gas _________ through a tiny hole in its container into ____________.

Back 50

process,

escapes,

an evacuated space.