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Gen Chem 01 > Chapter 05: Gases > Flashcards

Chapter 05: Gases Flashcards

1
Q

Pressure

A

The force exerted per unit area by gas molecules as they strike the surfaces around them

Pressure = force/area = F/A

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2
Q

Pressure units & conversion factors

A

760 mmHg = 760 torr = 1 atm = 101,325 Pa

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3
Q

Elements that exist as gases at 25°C and 1 atm

A

1A: H

5A: N

6A: O

7A: F, Cl

8A: (all)

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4
Q

Force (formula)

A

mass × acceleration

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5
Q

SI units of force and pressure

A

Force: 1 newton (N)

Pressure: 1 pascal (Pa)

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6
Q

Newton

A

1 newton = 1 kg×m/s2

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7
Q

Pascal

A

1 pascal = 1N/m2 = 1 kg/(m×s2)

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8
Q

Barometer

A

Atmospheric pressure measurment tool

Inverted tube of Hg over open dish of Hg;
height of Hg in mm is equal to atmospheric pressure

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9
Q

Manometer

A

Instrument used to measure pressure of gas trapped in container

Gas pressure is determined by difference in liquid levels in U-shaped tube

(Pressure is high if it pushes down on liquid;
low if it cannot push down on liquid)

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10
Q

Simple Gas Laws: Four* basic properties of a gas

A

P, pressure

V, volume

T, temperature (Kelvin)

t, temperature (°C)

n, amount in moles

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11
Q

Boyle’s law

A

Inverse relationship between volume and pressure

*n is constant
*T is constant

P1V1 = P2V2

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12
Q

Charles’ law

A

Direct relationship between volume and temperature (Kelvin).

*P is constant
*n is constant

V1 = V2
T1 T2

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13
Q

Avogadro’s law

A

Direct relationship between volume and quantity (number of moles)

*P is constant
*T is constant

V1 = V2
n n

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14
Q

Ideal Gas Law

A

How a hypothetical gas behaves

PV = nRT

Where R is the ideal gas constant

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15
Q

R, ideal gas constant

A

R = 0.08206 L atm/mol K

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16
Q

Molar volume

A

The volume occupied by one mole of a substance

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17
Q

Ideal gas molar volume

A

22.414 L at STP of an ideal gas

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18
Q

STP

A

standard temperature and pressure

pressure = 1 atm at STP

temperature = 0°C or 273.15 K at STP

19
Q

Density of a gas (formula)

A

d = PM
RT

Where d = density
P = pressure in atm
M = molar mass
R = gas constant
T = temperature in Kelvin

20
Q

Molar mass of a gas (formula)

A

M = dRT
P

Where M = molar mass
d = density
R = gas constant
T = temperature in Kelvin
P = pressure in atm

21
Q

Dalton’s law

A

Partial pressures, Pn

Pn = nnRT
V

Ptotal = Px + Py + Pz

*volume and temperature are constant

22
Q

Mole fraction

A

Xn = nn
ntotal

Thus, with Dalton’s law:
Pn = XnPtotal

23
Q

Vapor pressure

A

The partial pressure of water vapor in a system

24
Q

Gas stoichiometry

A

General conceptual plan:

P, V, T / mass or volume of gas A –> moles of gas A –>
moles of gas B –> P, V, T / mass or volume of gas B

25
Q

Kinetic molecular theory

A
  1. Particle size is negligible, even those they have mass
  2. Constant motion, random directions, no overall loss of energy, just a transfer of energy (known as being elastic)
  3. Average kinetic energy is directly proportional to temperature in Kelvin
  4. Gas particles exert neither attractive nor repulsive forces
26
Q

Average kinetic energy

A

average kinetic energy = KE

( KE = 1/2×mv2 )

At same temperature, gases have same average kinetic energy

27
Q

Boyle’s law explained (KMT)

A

Decreasing volume forces gas particles into a smaller space, thus increasing collisions, and hence, pressure

*n and T are constant

28
Q

Charle’s law explained (KMT)

A

Temperature increase increases average kinetic energy

The higher the average kinetic energy, the greater the area particles will move around

Thus, volume increases

*P & n constant

29
Q

Avogadro’s law explained (KMT)

A

Increasing number of particles causes more collisions

To keep pressure constant, volume must then increase

*P and T are constant

30
Q

Dalton’s law explained (KMT)

A

Because average kinetic energy is the same (at same temperature), the total pressure of the collisions is the same

31
Q

Root mean square velocity

A

A kind of average velocity

root mean square velocity = urms

urms = (3RT/M)1/2

Where M = molar mass in kg/mol
R = gas constant in J/mol K
T = temperature in Kelvin

32
Q

R, gas constant (with joules)

A

8.314 J/mol K

*used as R in root mean square velocity

33
Q

1 Joule = ? (kg, m, s)

A

1 J = 1 kg×m2/s2

34
Q

Mean free path

A

The average distance a gas molecule travels between collisions

mean free path v as pressure ^

35
Q

Diffusion

A

Process by which gas molecules spread out
in response to concentration gradient

Heavier molecules diffuse more slowly than lighter ones

36
Q

Effusion

A

The process by which a gas escapes from a container into a vacuum through a small hole

Heavier molecules effuse more slowly than lighter ones

37
Q

Rate of effusion

A

Amount of gas that effuses in a given time

Inversely proportional to square root of molar mass of gas

38
Q

Graham’s law of effusion

A

The ratio of effusion rates of two different gases is equal to the square root of the division of the molar masses

Since rate of effusion is an inverse proportion, it can be said that rateA times the squre root of A’s molar mass is equal to rateB times the squre root of B’s molar mass.

39
Q

Real gases

A

Do not behave like ideal gases at high pressure or low temperature

Compared to ideal gases:

Low pressure: lower ratio of PV/RT
High pressure: higher ratio of PV/RT

40
Q

Effect of finite volume of gas particles

A

Size of gas particles become important at high pressure

Because size of particles take up significant portion of total gas volume

Hence, volume of the particles must be corrected – increased

41
Q

Effect of intermolecular forces

A

At low temperatures, pressure is less than an ideal gas’

Because gas atoms spend more time interacting and less colliding, thereby decreasing pressure

42
Q

van der Waals equation

A

correction for intermolecular forces (P-) × correction for particle volume (V+) = nRT

43
Q

1 atm = ? Pa

A

1 atm = 101,325 Pa

Gen Chem 01 (21 decks)

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