Transcript Chapter Two Atoms & The Periodic Table

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Gas vs. Vapor
 Vapor is a term referring to the gas state of
something that is liquid/solid at room temperature
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Few elements exist as gases at room temp.
 H,N,O,F, Cl, and the noble gases
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Molecular compounds with LOW molar masses
tend to be gases at room temp (i.e. CO2, HCl)
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Gases expand and take shape of
container
 No definite shape or volume
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Gases are compressible
Gases have MUCH lower densities
 Density is also HIGHLY variable depending
on temperature and pressure
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Gases form homogeneous mixtures with
one another
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1)
2)
3)
4)
Four basic assumptions:
Gases are tiny particles with large amounts
of space in between (volume of gas particle
is negligible).
Gases in constant, random, straight-line
motion (all collisions perfectly elastic)
No attractive/repulsive forces on each other
Average kinetic energy is proportional to the
absolute temperature
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As molecular weight (mass) increases, speed
decreases of gas particles
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Diffusion: mixing of gases as the result of
random motion and collisions
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Effusion: escape of gas molecules from
container to a region of vacuum
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Pressure = Force/Area
 SI unit for force is Newtons (N)
 SI unit for pressure is pascals (Pa)
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Gases exert pressure on everything they
touch
 Internal pressure & External pressure
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Atmospheric Pressure: pressure exerted by
Earth’s atmosphere
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Atmospheres (atm)
Pascals (Pa) and kilopascals (kPa)
Bar (bar)
Millimeters of Mercury (mmHg)
Torr (torr)
Pounds per Square Inch (psi)
1.00 atm = 101,325 Pa = 101.325 kPa = 1.01325 bar =
760. mmHg = 760. torr = 14.7 psi
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Barometer: measures atmospheric pressure
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Manometer:
measures pressures
other than
atmospheric
pressure
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Convert the
following:
 12.4 psi  kPa
 509 bar  atm
 98,320. Pa  mmHg
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Assuming atmospheric
pressure is 765.0 mmHg,
what is the pressure of
the gas in the bulb?
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Gases and their physical state can be
described with the following four parameters:
 Pressure (P)
 Volume (V)
 Temperature (T)
 Number of moles (n)
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Relationship between Pressure & Volume
 Constant # of moles and temperature
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P1V1 = P2V2 = k
(k is a constant)
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As pressure increases, volume decreases (inverse)
 If pressure doubles, volume was halved
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Ex: What will be the pressure of 21 mL of a gas if
its pressure was 731 mmHg at 43 mL?
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Relationship between Volume & Temperature
 Constant pressure & number of moles
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V1 = V2
T1 T2
As temperature increases, so
does volume (direct relationship)
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Based on absolute (Kelvin) scale, so temp MUST be
in Kelvin! ( °C K, add 273.15)
 Ex: If a 1.45 L balloon is cooled from 25.0°C to 15.0°C,
what is the new volume?
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Relationship between Volume & # of moles
 Constant pressure and temperature

V1 = V2
n 1 n2
As # of moles increases, volume
increases (direct relationship)
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Avogadro proposed equal volumes of different
gases contained the same amount of particles at
the same temperature and pressure
 1 mole = 22.41 L @ STP (1 atm and 0°C)
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Relationship between Volume, Pressure,
Temperature, & # of moles
 Combination of Boyle’s & Charles’s Law
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P1V1 = P2V2
n1 T1 n2T2
*Remember, temp must
be in Kelvin!
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Ex: If a child releases a 6.25-L balloon where the temperature is
28.50°C and the air pressure is 757.2 mmHg, what will be the volume
of the balloon when it has risen to a height where the temperature is
-34.35°C and the air pressure is 366.4 mmHg?
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At what temperature will a gas sample
occupy 100.0 L if it originally occupies 76.1 L
at 89.5°C? Assume constant P.
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If the pressure of a container is 1.50 atm at a
volume off 37.3 mL, what will be the new
volume if the pressure is changed to 1.18
atm?
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Includes all 4 gas variables
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PV = nRT
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P = atm or kPa
V = liters
n = moles
T = Kelvin
R = 0.08206 atm*L/mol*K or 8.314 kPa*L/mol*K
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One in which a sample of gas’s pressurevolume-temperature behavior is predicted by
the ideal gas law equation
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Most of the time the differences between an
ideal gas and a real gas are negligible
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What is the volume of 5.12 moles of an ideal
gas at 32.0°C and 749.5 torr?
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Consider the reaction of zinc with
hydrochloric acid to produce zinc chloride
and hydrogen gas.
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How many grams of zinc are necessary to
produce 11.5 L of hydrogen gas at 825 torr
and 25.0°C?
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Ex: Calculate the density of CO2 at room
temperature (25.0°C) and 1.00 atm.
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Gas Laws and KMT assume ideal gas behavior
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Most gases exhibit ideal (or nearly ideal)
behavior under ordinary conditions
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Conditions of high pressure and low
temperature constitute a deviation of a real
gas from an ideal gas
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In a mixture of gases, each gas exerts a
pressure as though it were by itself
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Each gas exerts its own pressure, regardless
of the presence of any other gas
 Partial Pressure
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Ptot = P1 + P2 + P3 + … Pn
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Another way of showing the quantity of a gas
in a mixture of gases
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Mole Fraction = moles of gas X
total # of moles
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Remember:
 Mole fractions less than 1
 Sum of mole fractions equals 1
 Mole fractions are unitless
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Since n and P are proportional at a set T and V,
mole fractions can also be found be dividing
partial pressure by total pressure
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Mole Fraction = Partial Pressure
Total Pressure
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Ex: Three gases occupy a volume of 1.55 L.
Gas 1 has a pressure of 1.30 atm; Gas 2 has a
pressure of 3.10 atm; Gas 3 has a pressure of
2.70 atm. What is the total pressure? What is
the mole fraction for each gas?
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Ex: A 1.00 L vessel contains 0.215 moles of N2
gas and 0.0118 moles of H2 gas at 25.5°C.
Determine the partial pressure of each gas
and the total pressure in the vessel.
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What is the mole fraction of CO2 in a mixture
of 0.756 moles of N2, 0.189 moles of O2, and
0.0132 moles of CO2? If the total pressure in
the container is 212.4 atm, what is the
pressure of CO2?
In the lab, you react aluminum with excess
hydrochloric acid. As the reaction goes to
completion, you observe that 19.12 mL of gas has
been collected in the buret. The leveling bulb has
been held in such a way where the liquid level is
equal to that of the buret. If the pressure in the
room is 0.9894 atm, the temperature is 22.5°C, and
the vapor pressure of water at this temperature is
0.02687 atm, what is the number of moles produced
of hydrogen gas alone?
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Stoichiometry where reactants and products
are all gases
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Ex: How many liters of ammonia gas can be
produced from 2.38 L of nitrogen gas reacted
with excess hydrogen gas?
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Recall: at a set temperature and pressure
(typically STP), the volumes of 1 mole of any
gas are equal.
 Use mole ratio
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Ex: N2(g) + 3H2(g)  2NH3(g)
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Must use ideal gas law if conditions are NOT
at STP
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Determine the mass of NaN3 needed for an
air bag to produce 100.0 L of N2 gas at 85.0°C
and 1.00 atm according to the equation:
2NaN3(s)  2Na(s) + 3N2(g)