6. Gases, Liquids, and Solids Flashcards
Attractive forces that hold matter together…
-counteracts kinetic energy
-in the absence of attractive forces, the kinetic energy keeps the particles constantly moving randomly
Kinetic energy increases…
with increasing temperature
Gaseous state
-at high temperature:
molecules possess high kinetic energy; moves fast, attractive forces weak
Liquid state
-at lower temperature:
molecules move more slowly, attractive forces is stronger than gaseous state
Solid state
-at even lower temperature:
molecules no longer have velocity to move past each other
-strong attractive forces
Pressure
-force per unit area exerted against a surface
Barometer
-measure atmospheric pressure
-long glass tube that is completely filled with mercury then inverted into a pool of mercury in a dish
-measure of Mercury rise in tube= cmHg/inHg/mmHg (measure of pressure)
Evangelista Torricelli
-torr
-invented barometer
Manometer
-measure the pressure of gas in a container
-uses mercury as working fluid due to its very high density (height of liquid volume inversely proportional to liguid density); para small barometers and manometers lng
- Open-tube manometer: suited for measuring pressures equal or greater than atmospheric pressure
Pgas = Pheight + Patm - Closed-tube manometer: measuring pressure below atmospheric (vacuum)
Pgas = Ph
Boyle’s law
Robert Boyle
-for a fixed mass of gas at T=k,
⬆️ V ⬇️ P
Charles’s law
Jacques Alexandre Cesar Charles
-for a fixed mass of gas at P=k,
⬆️ V ⬇️ T
Gay-Lussac’s law
Joseph Louis Gay-Lussac
-for a fixed mass of gas at V=k,
⬆️P ⬇️T
Avogadro’s law
-equal volumes of gases at same temperature and pressure contain equal number of molecules, regardless of their identity
Standard temperature and pressure (STP)
-1 bar, 0°C
-all gases at STP or any other combination of T and P contain the same number of molecules at any given volume
Ideal gas
-a gas whose physical properties are described accurately by the ideal gas law
Ideal gas condition
-Real gases behave most like ideal gases at LOW PRESSURE (1 atm or less) and HIGH TEMPERATURE (300K or higher)
Partial pressure
-the pressure that a gas in a mixture of gases would exert if it were alone in the container
Dalton’s law of partial pressure
-total pressure is the sum of the partial pressures of each individual gas
PT = Pa + Pb +….
Pi = xi.PT
where xi: mole fraction of gas
Kinetic molecular theory
(Idealized picture of molecules of gases)
- Gases consist of particles, either atoms or molecules, constantly moving through space in straight lines, random direction, and various speeds
- The average kinetic energy of gss particles is proportional to temperature in kelvins
- Molecules collide with each other. Each time they collide, they may exchange kinetic energy but the total kinetic energy of the gas sample remains the same
- Gas particles have no volume
- There are no attractive forces between gas particles
- Molecules collide with the walls of the container, and these collisions constitute the pressure of a gas
Ideal gas
-at STP, most real gases behave in such a way that an ideal gas would
Condensation
-change of a substance from the vapor or gaseous state to the liquid state
Solidification/ Crystallization
Change from liquid to solid
London dispersion forces
-Fritz London
-extremely weak attractive forces between atoms or molecules caused by electrostatic attraction between temporary induced dipoles
-exist between all molecules
-only forces of attraction in nonpolar molecules
Dipole-dipole interactions
-attraction between positive end of a dipole of one one molecule and negative end of another dipole in the same or different molecule
Hydrogen bond
-a noncovalent force of attraction between the partial positive charge on a hydrogen atom bonded to an atom of high electronegativity (O,N,F) and a partial negative charge on a nearby O,N,F
Liquids at molecular level
-incompressible
-density of liquids are much greater tgan gases because same mass occupies a smaller volume in liquid form
Surface tension
-directly related to strength of intermolecular attraction between its molecules
-tendency of fluid surface to shrink into minimum surface area possible
Vapor pressure
-the pressure of a gas in equilibrium with its liquid form in a closed container
*at equilibrium, the rate of vaporixation equals the rate of liquefaction
Boiling point
-temperature at which the vapor pressure of a liquid is equal to the pressure of the atmosphere in contact with its surface
Normal boiling point
Temperature at which a liquid boils under a pressure of 1 atm
Factors that affect boiling point
- Intermolecular forces (esp for molecules having similar molecular weight).
⬆️ Intermolecular force, ⬆️ Boiling point - Number of sites for intermolecular interaction, (surface area).
⬆️ Surface area ⬆️ Boiling point - Molecular shape
(i.e. pentane (linear shape) vs 2,2-dimethylpropane (spherical shape); pentane higher b.pt. since larger surface area)
Carbon
-five crystalline form
(i.e. diamond and graphite, buckyball, nanotubes, soot; carbon atoms are packed differently)
Buckyballs
-60 carbon atoms named buckminsterfullerene “buckyball”
-named after Buckminster Fuller for his invention geodesic domes (similar structure as buckyball)
Fullerenes
-cage like structures containing 72, 80, and even larger atoms
Nanotube
-new variations of fullerene
-cross section of each tube is only 10^(-9) m
-industrial interest due to optical and electronic property
Soot
-solidifies directly out of carbon vapor
-amorphous solid
Types of Solids
1. Ionic
- ions in a a crystal lattice
-high melting point
-NaCl, K2SO4
Types of Solids
2. Molecular
-molecules in a crystal lattice shape
-low melting point
-ice, aspirin
Types of Solids
3. Polymeric
-giant molecules
-can be crystalline, semicrystalline, amorphous
-low melting point or cannot be melted
-soft or hard
-rubber, plastic, proteins
Types of Solids
4. Network
-very large number of atoms connected by covalent bonds
-very hard
-very high melting point or cannot be melted
-diamond, quartz
Types of Solids
5. Amorphous
-randomly arranged atoms or molecules
-mostly soft, can be made to flow, no melting point
-soot, tar, glass
Phase change
-change from one physical state to another
Heat of fusion/ melting
-heat necessary to melt 1.0g of any solid
Specific heat
-heat required to raise the temperature of 1g of fluud sample
Heat of vaporization
-amount of heat necessary to vaporize 1.0 g of liquid at its normal boiling point
Sublimation
-transition from solid state directly to vapor state
-solids usually sublime only at reduced pressures (less than 1 atm)
Phase diagram
-Pressure (y) against temperature (x)
-solid liquid vapor
Triple point
All three phases coexist
For water:
Specific heat of ice
0.48 cal/g °C
Heat of fusion of ice
80 cal/g
Specific heat of liquid water
1.0 cal/g °C
Heat of vaporization
540 cal/g
Specific heat of steam
0.48 cal/g °C
Gas vs Vapor
•Gas- substance that is normally in gaseous state at ordinary temperatures and pressures
•Vapor- gaseous form of any substance that is a liquid or a solid at normal temperature.
Atmospheric pressure
-pressure exerted by Earth’s atmosphere
-depends on locations, temperature, and condition
-atmosphere is much denser near the surface of the Earth than at high altitudes;
denser air, greater pressure it exerts
Absolute zero
-Lowest attainable temperature (theoretically) (-273.15°C)
Absolute temperature scale/
Kelvin temperature scale
-William Thomson, Lord Kelvin
-absolute zero as starting point
Molar mass of unknown substance
-bulb of known volume is filled with unknown gas
-temperature and pressure
-total mass of bulb plus gas
-mass of empty bulb
-difference of total mass and mass empty = mass of gas
-density of gas = mass of gas/ volume of bulb
Molar mass = (density.R.T) / P
Scuba diving and gas laws
•Application of Boyle’s Law
33 ft of seawater = 1 atm pressure
⬆️ Pressure ⬆️ Depth
-divers must ascend slowly to adjust to decreasing pressure
-otherwise, rupture membranes of lungs due to air expansion; OR air embolism (air forced into capillaries, blocking normal blood flow) resulting to lost of consciousness
•Application of partial pressure
-body functions best when
PO2 = 0.20 atm
-when diver submerged, pressure of water on diver is greater than atmospheric pressure
-special valve adjusts pressure air breathed from scuba tank to ensure air pressure = water pressure all times (adjusts oxygen content in air breathed in terms of %volume) -otherwise, body cavities like lungs and sinuses will collpase
Gas pressure
-result of collissions between molecules and the walls of their containers
-depends on frequency of collission per unit area and how hard molecules hit the wall
⬆️Temperature ⬆️KE of molecules
⬆️Speed of molecules
Maxwell speed distribution
-at given temperature, distribution curve tells us the number of molecules moving at certain speed
-peak of each curve
MOST PROBABLE SPEED: speed of the largest number of molecules
⬆️ Temp, ⬆️ Most probable speed
(As temp increases, the peak shifts toward right and the curve flattens out, indicating larger number of molecules are moving at greater speed)
Root-mean-square speed
-average molecular speed at any temperature
•Earth, unlike Jupiter, doesn’t have appreciable amount of Hydrogen/ Helium
-weaker gravitational attraction, smaller escape velocity required
-average speed of He exceeds Ni and O2, so trace amount of He found in atmosphere
Diffusion
-gradual mixing of molecules of one gas with molecules of another by virtue of their kinetic properties
-region of higher concentration to one of lower concentration
