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Solid
- substance maintains shape and volume
- rigidly packed 3-dimensional pattern
- definite shape
- definite volume
- particles packed and touching
- high density
- small compressibility
- very small thermal expansion
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Liquids
- substance assumes shape of its container
- flows readily
- indefinite shape
- definite volume
- particles touching and mobile
- high density
- small compressibility
- small thermal expansion
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Gas
- substance maintains neither shape nor volume
- indefinite shape
- indefinite volume
- particles far apart
- low density
- large compressibility
- moderate thermal expansion
- mix spontaneously and completely through diffusion
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chemical bonds
attractive forces that hold atoms together in molecules and ions together in crystals
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ionic bons
- metal + nonmetal
- complete transfer of valence electrons
- form crystalline solids
- high melting points
- water soluble
- can conduct electricity (electrolyte)
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covalent bond
- valence electrons shared by 2 atoms in a compound
- single, double, triple bonds
-
nonpolar covalent bond
- equal sharing of electrons
- occurs between same kind of atoms (H2, Cl2)
-
polar covalent bond
- unequal sharing of electrons
- occurs between atoms of different elements (H2O, SO3)
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electronegativity
- the ability of an atom to attract shared electrons to itself in a covalent bond
- electrons are attracted toward atoms of highest electronegativity
- periodic table increases left to right, decreases top to bottom
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intermolecular forces
- forces between molecules
- --dipole forces
- --hydrogen bonds
- --london dispersion forces
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dipole forces
- polar covalent bonds
- electron pairs shared unequally
- unbalanced separate center of partial positive and negative charges
- positive end attracts negative end
-
hydrogen bonds
- intermolecular force between electropositive H (bonded to N, O, F) and electronegative O, F, N
- strongest bond
-
london dispersion forces
- non polar molecules
- small transient momentary attractive forces
- weakest bond
-
kinetic energy
- proportional to kelvin temp
- temp increaes, particle velocity increases
- at fixed temp, lighter particles move faster than heavier particles
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Kinetic Molecular Theory of Gases
- 1. gases move continuously, rapidly, randomly
- 2. particles are tiny / great distances
- 3. forces between molecules are negligible
- 4. collisions are elastic (no energy loss)
- 5. average kinetic energy of gas particles is same for all gases at same temp
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atmosphere
- thin blanket of gases around earth
- 78% N2, 21% O2, 1% Ar, 4% H2O, trace CO2
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atmospheric pressure
force per unit area exerted by atmosphere
-
pressure
- force exerted per unit area
- pressure = force/area
-
pressure units
- atm: atmosphere
- mm Hg: mm mercury
- torr
- Pa: pascal (SI unit of pressure)
- psi: pounds per square inch
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standard pressure
1 atm = 760 mm Hg = 760 torr = 14.7 psi = 101,325 Pa
-
Boyle's Law
Pressure and volume inversely proportional at constant temp and qty of gas
P1 x V1 = P2 x V2 = constant k
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Charles's Law
Volume and temperature are linear
V^T^
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Gay-Lussac's Law
Pressure and temp are linear
P^T^
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Avogadro's Law
Equal volumes of gases contain equal number of molecules (or atoms)
V1/n1 = V2/n2
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Ideal Gas Law
PV = nRT
- P = pressure
- V = volume
- T = temp
- n = quanity of gas
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Dalton's Law of Partial Pressures
total pressure of gas mixture = sum of partial pressures of exerted gases
P(total) = P1 + P2 + P3
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Oxidation
- lose electron
- lose hydrogen
- gain oxygen
- increase oxidation number
-
Reduction
- gain electrons
- gain hydrogen
- lose oxygen
- decrease oxidation number
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Rules of ON
- 1. ON of any single element is 0
- 2. ON of simple ion = charge of ion (Mg2+ = +2)
- 3. ON of 1A and 2A elements when in compounds is +1, +2
- 4. ON of H is always +1
- 5. ON of O is always -2
- 6. Sum of ON of all atoms in compound = 0
- 7. Sum of ON of all atoms in polyatomic ion = charge of ion
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