Unit Six Flashcards
Liquid
Medium packed particles, medium motion
Gas
Loosely packed particles, high motion
Temperature
A measure of the average kinetic energy particles of a substance have
Heat
Transfer of kinetic energy from a hotter object to a cooler one
Melting
Solid to liquid
Freezing
Liquid to solid
Vaporization
Liquid to gas
Condensation
Gas to liquid
Sublimation
Gas to solid
Deposition
Solid to gas
Fahrenheit to Celsius
F = 1.8 C + 32
Celsius to Fahrenheit
C = (F - 32)/1.8
Celsius and Kelvin
K = C + 273
Pressure conversions
1 atm = 14.7 psi = 101 kPa = 7.60x10^2 mm Hg (torr)
STP
Standard temperature and pressure, 1 atm and 273 K
Intramolecular forces
Hold the atoms of a molecule together
Ex. Ionic bonds, polar covalent bonds, nonpolar covalent bonds
Intermolecular forces
Hold molecules of a substance together
Ex. Dipole-dipole, hydrogen bonding, dispersion forces
Dispersion forces
Weak forces resulting from temporary shifts in electron densities, stronger for larger particles
In nonpolar molecules, bigger molecules = bigger force
Weakest type of intermolecular bonds
Dipole-dipole forces
Attraction between oppositely charged region of polar molecules
Polar covalent
Middle strength
Hydrogen bonds
Dipole-dipole attraction occurring between molecules containing hydrogen boned to fluorine, oxygen, or nitrogen
Strongest because greatest difference in electronegativity
Polar covalent
Strongest type
Are hydrogen bonds or ionic bonds stronger
Ionic because hydrogen is partial charges and ionic is full charges
Liquid density
Denser than gases
Liquid compression
Cannot compress a liquid (definite volume)
Liquid fluidity
Particles in a liquid do not have fixed locations
Liquid viscosity
A measure of resistance of a liquid to flow
Factors affecting viscosity
Stronger intermolecular forces mean greater viscosity
Larger and longer molecules will form more viscous solutions
Viscosity decreases with temperature
Surface tension
Energy required to increase the surface area of a liquid
Because intermolecular forces at the center of the liquid experience balanced attractions on all sides while molecules at the surface are only pulled inward, thus creating a taught surface
Cohesion
Attraction between identical molecules
Adhesion
Attraction between different molecules
Capillary action
What happens when the adhesion forces overcome the cohesion forces
Solid
Closely packed particles, low motion
Special properties of water due to hydrogen bonding
High melting and boiling point (takes more energy to line up or separate hydrogen bonds)
Surface tension and capillary action (strength of hydrogen bonds pulls them to the center of the liquid and towards one another and their container)
Solid water is less dense than liquid water (when cooled hydrogen bonds align in a spacious crystal pattern that makes the ice less dense than the water bc of liquid molecules)
Universal solvent (neutral, no net charge, can dissolve most things)
Crystalline solids
Atoms, ions, or molecules in an orderly, geometric structure
Molecular solid
Covalent substance, break down easily, soft, low melting points, water, sugar, dry ice
Covalent network solid
Crystal structure, carbon, silica
Ionic solid
Soluble, high melting points, salts
Metallic solid
Orderly atoms, ductile, malleable, any metal
Allotropes
Multiple different forms of the same state Ex carbon (diamond, graphite, etc)
Amorphous solids
Solid where particles are not arranged in a repeating pattern
Ex obsidian, glass rubber
Properties of gases
Gases are fluids
Gases are highly compressible
Gases completely fill containers
Gases have lower densities than liquid and solids
KMT
Kinetic molecular theory, describes the motion of particles
Gas molecules are in constant, random motion
Gas molecules are separated by huge distance relative to the size of the molecules themselves
Gas molecules have no attractive/repulsive forces
Gases are made of molecules that have mass
Molecules undergo elastic collisions
Relationship between temperature and kinetic energy
Temperature and energy of molecules are directly proportional
If temperature increases the kinetic energy increases
Boyle’s law
P1V1 = P2V2, inverse relationship
Charles law
V1/T1 = V2/T2, direct relationship
Gay-Lussac law
P1/T1 = P2/T2, direct relationship
Avogadro law
22.4 L = 1 mole gas
Combined gas law
P1V1/T1 = P2V2/T2
Ideal gas law
PV = nRT P = pressure in atm V = volume in liters n = moles R = 0.08205 if atm, 8.31 if kPa T = temperature in kelvin
Ideal gas
Gases that behave according to the assumptions of kinetic molecular theory
Assumptions made for ideal gases
No intermolecular forces
Individual molecules have no volume
Conditions where real gases behave ideally
High temperature and low pressure
Diffusion
Gas particles will travel from an area of high concentration to an area of low concentration until they are evenly distributed
Effusion
Passage of a gas through a small hole/opening in a barrier
Graham’s law of effusion
Gases that have larger molar masses move more slowly than lighter gases
Avogadro’s hypothesis
Equal volumes of gases at the same temperature and pressure contain equal number of particles
Other things about Avogadro’s hypothesis
Due mainly to the large amount of empty space between particles
From this, scientists have determined that one mole of gas = 22.4 L at STP
Dalton’s law of partial pressure
The total pressure for a mixture of gases is equal to the sum of the pressures of the individual gases in the mixture
Ptotal = P1 + P2 + Px
On a phase diagram, the top third is
Solid
On a phase diagram, the middle third is
Liquid
On a phase diagram, the bottom third is
Gas
Triple point
The temperature and pressure where all three phases of matter could exist
Where the three lines converge and meet at a single point
Critical point
Pressure and temperature above which distinct gas and liquid phases don’t exist
Density formula
D = m/v
