chem unit7
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States of Matter
Particle vibration fluid motion rapid, random motionRigid positions move past independent of each other
each otherFixed volume fixed volume volume of container
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Phases and Transitions
Sublimation
Condensation
EvaporationMelting
Freezing
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Intermolecular Forces
• Strongest = IONIC FORCES
• High melting points• Oppositely charged ions
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Dipole-Dipole Forces
• Dipole: contains both positively and negatively charged regions:
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Hydrogen Bonding
• Special case of dipole-dipole forces
• Causes water to have higher than expected boiling point
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Water Properties
• Surface tension
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Water Properties
• Capillary action
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Boiling PointsMolecule Boiling
Point
H2O 100°C
H2S –60.7 °C
H2Se –41°C
H2Te –2°C
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Hydrogen Bonding
• Weak bond between H and electronegative atom:
• Recall
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Hydrogen Bonding
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Hydrogen Bonding
• DNA
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Dipole-Dipole Forces
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London (Dispersion) Forces
• Weak attractions between non-polar
• Increases with size of molecule (number of electrons)
• Molecular Shape (less compact > compact)
> >
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London (Dispersion) Forces
• Weak attractions between non-polar
• “Temporary” or instantaneous dipoles
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Intermolecular Forces
IONIC >> Dipole – Dipole >> London
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Intermolecular Forces
• Which intermolecular force is expected?
CH4
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Intermolecular Forces
• Which intermolecular force is expected?
CH4
Non-polar covalent molecule
London (dispersion) forces
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Intermolecular Forces
• Which intermolecular force is expected?
Butanol
CH3CH2CH2CH2OH
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Intermolecular Forces
• Which intermolecular force is expected?
Butanol
CH3CH2CH2CH2OH
Dipole-Dipole
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Energy of Phase Changes
• Energy is required for all phase changes
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Phase Diagram (H2O)
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Phase Diagram (CO2)
Triple Point
Triple Point: where all three phases co-exist (T, p)
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Phase Diagrams (carbon)
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Properties of Matter
ForcePressure = area
h = 760 mm = 1 atm
Torricelli barometer
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Robert Boyle
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Boyle’s Law
• For a gas at constant T and nV and p are inversely proportional
pV = constant
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Charles’ Law
• At constant pressure, V/T = constant
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Charles’ Law
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Gay – Lussac’s Law
• In a constant volume:P/T = constant
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Gay-Lussac’s Law
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Combined Gas Law
• Boyle’s Law pV = constant• Charles’ Law V/T = constant• Gay-Lussac’s Law p/T = constant
Combining all three:p1V1 p2V2
So: T1 = T2
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Combined Gas Law
• A sample of gas has a volume of 400 liters when its temperature is 20°C and its pressure is 300 mm Hg. What volume will the gas occupy at STP?
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Combined Gas Law
p1V1 p2V2
T1 = T2 T in Kelvin
(300/760 mm Hg)(400 L) = (760 mm Hg) (V2)
(293 K) (273 K)
V2 = 147 Liters
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Combined Gas Law
• A sample of He gas has a volume of 250 mL at 456 torr and 25°C. At what temperature does this gas have a volume of 150 mL and 561 torr?
p1V1 p2V2
T1 = T2
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Combined Gas Law
• A sample of He gas has a volume of 250 mL at 456 torr and 25°C. At what temperature does this gas have a volume of 150 mL and 561 torr?
p1V1 p2V2 T2 = p2V2 T1
T1 = T2 p1V1
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Combined Gas Law
• A sample of He gas has a volume of 250 mL at 456 torr and 25°C. At what temperature does this gas have a volume of 150 mL and 561 torr?
p1V1 p2V2 (561/760)(0.15L)(298K)
T1 = T2 T2 = (456/760)(0.25L)
= 220 K = -53°C
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Ideal Gases
• Non-interacting
• Point particles
• Randomly moving with elastic collisions (no energy lost)
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Ideal Gases
• Avogadro’s Law:Equal volumes of gas contain the same number of molecules at the same T & p.
n = number of moles
p1V1 = constant
n T1
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Ideal Gas Law
p1V1 = constant = 0.082 L atm = R
n T1 K mole
= Universal Gas Constant
One mole of gas at STP:Volume = nRT/p = (1 mole)(0.082 Latm)
(273K)/1 atm
= 22.4 Liters
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Ideal Gas Law
• pV = nRT
• How many moles of Helium are present in a balloon that has a volume of 65 L at 20° C and 705 torr?Given Needed
V, T, p, R n n = pV/RT
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Ideal Gas Law
• pV = nRT
• How many moles of Helium are present in a balloon that has a volume of 65 L at 20° C and 705 torr?
n = pV/RT
= (705/760 atm)(65 L)
(0.082 Latm/K)(293)
= 2.5 moles He
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Ideal Gas
6.2 liters of an ideal gas are contained at 3.0 atm and 37 °C. How many moles of this gas are present?
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Ideal Gas
6.2 liters of an ideal gas are contained at 3.0 atm and 37 °C. How many moles of this gas are present?
n = pV/RT
= (3 atm)(6.2 L)(0.082 L atm/mole K)(310 K)
= 0.73 moles
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Ideal Gases and Density
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Density
• Gas density increases with molecular mass.
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Density
• What is the density of NO2 gas at 0.97 atm and 35°C?
MW = 46 g/mole Molar mass p (46 g/mole)(0.97 atm)
• Density = RT (0.082 L atm/mole K) (308 K)
= 1.767 g/L
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Gas Diffusion
• Movement of particles from region of Higher density to lower density
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Gas Diffusion
• Movement of particles from region of Higher density to lower density
Depends on density (molar mass)
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Graham’s Law proportionality
Rate of effusion inversely to square root of molar mass
Smaller molecules escape FASTER than larger molecules
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Dalton’s Law
PT = P1 + P2 + P3 + . . . .
Total pressure of a gas sample is the sun of the partial pressures.
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Dalton’s Law
• A mixture of O2, CO2 and N2 has a pT of 0.97 atm; if pO2 = 0.7 atm and pN2 = 0.12 atm, what is pCO2?
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Dalton’s Law
• A mixture of O2, CO2 and N2 has a pT of 0.97 atm; if pO2 = 0.7 atm and pN2 = 0.12 atm, what is pCO2?
• pT = pO2 + pN2 + pCO2 = 0.97 atm
• pCO2 = 0.97 atm - (0.7 atm + 0.12 atm)
= 0.15 atm
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Dalton’s Law
• The partial pressures of CH4 and O2 are 0.175 atm and 0.25 atm.
At 65°C in a volume of 2 L, how many moles of each gas are present?
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Dalton’s Law
• The partial pressures of CH4 and O2 are 0.175 atm and 0.25 atm.
At 65°C in a volume of 2 L, how many moles of each gas are present?
nCH4 = pV/RT = 0.175 atm (2L)/0.082 L atm/mole K
(338 K = 0.126 moles
nO2 = pV/RT = 0.25 atm (2L)/ 0.082 L atm/mole K (338 K)
= 0.018 moles
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Unit 7 Review
• Phases of Matter and Transitions• Intermolecular Forces• Phase Diagrams• Boyle’s, Charles’, Gay Lussac’s Laws• Ideal Gas Law pV = nRT• Graham’s Law of Diffusion• Partial Pressure