## Glossary of Chem Chapter Five

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- Atmosphere
- the mixture of gases that surrounds the earth's surface

- Avogadro's Law
- equal volumes of gas at the same temperature contain the same number of particles

- Boyle's Law
- the volume of a given sample of gas at constant temperature varies inversely with the pressure

- Charles' Law
- the volume of a given sample of gas at contstnat pressure is directly proportional to the temperature in kelvins

- Dalton's Law of Partial Pressures
- for a mixure of gases in a container, the total pressure exerted is the sum of the pressures that each gas would exert if it were alone

- Diffusion
- the mixture of gases

- Effusion
- the passage of a gas through a tiny orfice into an evacuated chamber

- Graham's Law of Effusion
- the rate of infusion of gas is inversely proportional to the square root of the mass of the particles

- Ideal gas
- a gas that obeys the equation, PV=nRT

- Ideal Gas Law
- an equation of state for a gas, where the state of the gas is its condition at a given time; expressed by PV=nRT, where P=pressure, V= volume, n=moles of gas, R=the universal gas constant, and T=absolute temperature. This eqation expresses the behavior approached by real gases at high T and low P.

- Kinetic molecular theory
- a model that assumes that an ideal gas is composed of tiny particles (molecules) in constant motion

- Manometer
- a device used for measuring the pressure of a gas in a container

- Millimeters of mercury
- (mm Hg) a unit of pressure, also called a torr; 760 mm Hg=760 torr =101,325 Pa = 1 standard atmosphere

- Molar volume
- the volume of one mole of an ideal gas; equal to 22.4L at STP

- Mole fraction
- the ratio of the number of moles of a given component in a mixture to the total number of moles in the mixture

- Partial pressures
- the independent pressures exerted by different gases in a mixture

- Root mean square velocity
- the square root of the average of the sqaures of the individual velocities of gas particles

- Standard atmosphere
- a unit of pressure equal to 760mm of Hg

- Standard temperature and pressure
- the condition of 0 degrees celsius and 1 atm of pressure

- Torr
- another name for millimeter of mercury (mm Hg)

- Universal gas constant
- the combined proportionality constant in the ideal gas law; 0.08206 L atm/K mol

- gas described quantitatively from four factors
- volume, pressure, temperature, and amount of gas present

- absolute zero
- this temperature where volumes of gas extrapolate to zero

- pressure
- force per unit area

- Pressure is equal to force per unit area. Which of the following equations is dimensionally consistent with this definition? g = m/s2, d = g/cm3, h = height
- p = gdh

- According to the postulates of the kinetic theory of gases, the size of the molecules that make up the gas is necessary in calculations using the ideal gas law.
- false

- According to the postulates of the kinetic theory of gases, increasing the volume of a container at constant temperature will result in a decrease in the pressure the gas exerts in the container
- true

- Problems with the ideal gas law arise as
- molecular size and interactions become important at high pressure and low temperature

- Which of the following does NOT describe an ideal gas?
- strong repulsion between molecules results in ideal gas behavior

- Torricelli's original determination of atmospheric pressure
- his barometer--a filled tube with Hg and inverted it into a dish filled with Hg

- Units used for the measurement of pressure and conversions between them
- 760mm Hg= 760 torr = 1 atm = 101.3 kPa

- Determination of the pressure of an individual gas
- use a manometer; gas in flask exerts a pressure (Pgas); atomospheric gases coming into the tube exert a pressure (Patm); opposite directions; opposing forces;

a) atmosphereic pressure > gas pressure: subtract height of Hg difference from atm. pressure

b) Patm < Pgas: add h to atm. pressure

_____________________________

If the tube is closed to the atmosphere and a vacuum is established, the Hg levels in the U-tube will be equal --when a gas is introduced into the flask, the height of the Hg difference will be an actual measurement of the gas pressure

- Why mercury
- high density; workable substance

- Boyle's Law:

1) studied

2) constant

3) relationship between factors

4) formula - 1) pressure and volume

2) temperature

3) inverse

4) P1V1=P2V2

- Charles' Law:

1) studied

2) constant

3) relationship

4) formula

5) development of the Kelvin temperature scale - 1) temperature and volume

2) pressure

3) direct

4) V1/T1 = V2/T2

5) all charles' measurements comaparing T and V for various gases resulted in a straight line graph... when these lines were extrapolated to zero volume they all intersected the x-axis (temperature) at the same point: -273.15 degrees C...working with Lord Kelvin, they identified this value as absolute zero and established the relationship... K= degrees C +273.15

- Combined Gas Law
- Keep original relationship to volume for P and T and combine all three together:

p1v1/t1 = p2v2/t2

- Avogadro's Law

1) studied

2) constant

3) relationship

4) formula - 1) volume and moles/particles

2) temperature

3) direct

4) V1/N1 = V2/N2

- The derivation of the Ideal Gas Law from the 3 separate gas laws
- V=R(Tn/P)

- Review Stoichiometery
- -Calculation to relate two different substances in a chemical reaction

-Only way to relate two different substances is with mole ratio

-Use molar mass to make conversions (mass<-->mole)

- Relationship between the volume of a gas and one mole and the conditions that are necessary
- One mole = 22.4L for ANY gas at STP; set value at zero degrees celsius and 1 atm pressure

- CALCULATIONS not under STP
- * cannot use 22.4L

* use stoichiometry to determine moles of gas

*use ideal gas law equation to determine volume

(or vice versa)

- Dalton's Law of Partial Presures
- For a mixture of gases in a container, the total pressure exerted is the sum of the pressures that each gas would exert if it were alone.

- CALCULATIONS using Partial Pressures & idea of mole fraction
- mole fraction: ratio of number of moles at a single component in a miture to total number of moles in the mixture

- The statements of KMT
- 1) volume of the individual particles can assumed to be negligible

2) particles are in constant motion; the collisions of the particles with the walls of the container are the cause of the pressure exerted by the gas

3) the particles are assumed to exert no forces on each other; they are assumed neither to attract nor to repel each other

4) the average kinetic energy of a collecion of gas particles is assumed to be directly proportional to the kelvin temperature of the gas

- These statments support the gas laws

1) boyle's

2) gay-lussac's

3) charles'

4) avogadro's - 1) decrease in volume means that gas particles will hit the wall more often, thus increasing pressure

2) when the temperature of a gas increases, the speeds of its particles increase, the particles hitting the wall with greater force and frequency; since the volume remains the same, this would result in increased gas pressure

3) when the gas is heated to a higher temperature, the speeds of its molecules increase and thus they hit the walls more often and with more force; the only way to keep the pressure constant in this situation is to increase the volume in the container; this compensates for the increasd particle speeds

4) An increase in the number of gas particles at the same temperature would cause the pressure to increase if the volume were held constant; the only way to return the pressure is to increase volume