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199 Cards in this Set
- Front
- Back
4 quantities describe the state of a gas |
1. Pressure 2. Temperature 3. Volume 4. number of moles |
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Boyle's law equation |
P1V1=P2V2 |
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Boyle's law relationships |
Increase pressure = decrease volume |
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Boyle's law constants |
Moles and temperature |
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Charles's law equation |
V1/T1 = V2/T2 |
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Charles's law relationships |
Increase temperature = increase volume |
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Charles's law constants |
Pressure and moles |
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How to solve for kelvin |
C=k- 273.15 |
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Gay lussac's law equation |
P1/T1=P2/T2 |
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Gay lussac's relationships |
Increase temperature = increase pressure |
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Gay lussac's constants |
Volume and moles |
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Combined gas law |
(P1V1/T1n1)=(P2V2/T2n2) |
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Avogadro's law equation |
V1/n1=V2/n2 |
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Avogadro's relationships |
Increase volume = increase number of moles |
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Avogadro's law constants |
Temperature and pressure |
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Ideal gas law aka universal gas law |
PV=nRT |
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Universal gas constant equation |
R=PV/nT |
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Universal gas constant number |
R= 0.08205 |
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Gas density equation |
p=PM/RT |
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Gas density relationships |
Increase p= increaseP Increase p = increase M Increase p = decrease T |
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Daltons law of partial pressures |
total pressure of a gaseous mixture is a sum of the partial pressures of each of the components gases |
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Daltons law relationships |
Increase pressure =increase concentration |
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Daltons law equation |
Ptotal=P1+P2+Pn |
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Partial pressure equals |
Ptotal x concentration of the gas |
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Concentrations of N2 O2 CO2 and H2O |
78.6% 20.9% 0.04% 0.5% |
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Relative humidity definition & relationship |
- saturation of water in air - increase air solubility =increase temperature |
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Operating roomrelative humidity |
20% to 60% |
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RH20% lower limit because? |
Decreases shelf life and increases chance of potential spark |
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Rh 60% upper limit because? |
Increases chance for mold and mildew growth and infection and less comfortable |
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Four tenets that describe ideal gases |
1. Gases consist of small particles whose volume is negligible compared to the volume of the gas 2. Gas molecules are in constant, random motion 3. Gas molecules in a sample have a range of kinetic energies, but the average kinetic energy depends on temperature 4. There are no attractive or repulsive forces between gas particles, so all collisions are elastic |
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Kinetic energy relationship in gases |
Increase temperature = increase kinetic energy |
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Graham's law of effusion equation |
Rate A /rate B =square root of (molar mass B/molar mass A) |
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Rate of effusion relationships |
Increase molecular mass = decrease rate of effusion |
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Fick's law of diffusion equation |
D =(change in P x A x S)/(d x square root of MW) |
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Rate of gas diffusion relationships |
- Increase concentration or pressure gradient = increase in diffusion -increase in solubility=increase in diffusion -increase in cross sectional area = increase in diffusion -increase in diffusion distance = decrease in diffusion -increase in molecular weight = decrease in diffusion |
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Real gases fail at which 2 tenets |
#1: at very high pressures, gas molecules represent a large percentage of the sample volume #4: real gases have attractive forces between molecules, kinetic energy is needed to overcome this attractionand pull molecules away from each other - at low temps, real gases have to use greater fraction of kinetic energy to overcome attractive forces = pressure of sample decreases |
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The Vander Waals equation |
PV = nRT |
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Van der Waals forces include |
Attraction and repulsion between atoms, molecules, and surfaces as well as other intermolecular forces |
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Van der Waals is designed to account for what |
The intermolecular forces between gas molecules and the volume of gas molecules |
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Adiabatic changes |
Occur when gases are rapidly expanded or compressed without surrounding environment |
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Adiabatic changes do what to systems energy |
No increase or decrease in systems energy |
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Adiabatic changes do what to temperature of system? |
Temperature may change die to change in kinetic energy per area |
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Energy concentration effect |
Rapid compression of gas = concentration of kinetic energy = rapid increase in temperature proportional to decrease in volume Ex-diesel engine |
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Energy dilution aka joule Thompson effect |
Rapid expansion of gas e dilution of kinetic energy = rapid decrease in temperature proportional to increase in volume Ex: frosting on bike pump |
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Which will more likely occur with compressed gas cylinders? |
Joule Thompson effect > energy concentration |
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benefits of solutions |
deliver larger more precise measured volumes better tolerated and less reactive some meds are solid and need to be dissolved as a liquid |
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homogenous mixture |
mixture that contains one or more solutes uniformly dispersed at the molecular or ionic level throughout the solution |
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solute |
the substance dissolved within the solution |
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solvent |
the substance that dissolves the solute |
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examples of solutions |
liquids, gases, and solids |
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molarity = molar concentration |
molar concentration = moles of solute per liter of solution |
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molar concentration equation |
M = mol of solute/L of solution |
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molar concentration depends on |
temperate of the solution, increase temperature you decrease molarity because volume increases |
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Molality = molal concentration |
moles of solute per kilogram of solvent |
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molality equation |
m = mol of solute/ kg of solvent |
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how does temperature effect molality |
temperature has no effect on molality because it does not affect mass |
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variations on percent |
1. percent by weight to volume (%w/v) 2. percent by weight (%w/w) 3. percent by volume (%v/v) |
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percent by weight to volume = %w/v (actually % by mass) |
%w/v = grams of solute/ ml solution (x 100) |
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percent by weight = %w/w |
%w/w = grams of solute/g of solution (x100) |
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percent by volume = %v/v |
%v/v = ml of solute /ml of solution (x100) |
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parts per million equation |
ppm = gram of solute /grams of solution (x10^6) |
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solubility |
amount go solute that will dissolve in a given amount of solvent at room temperature |
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what happens when you add solvent to an already saturated solution? |
excess solute will crystallize as a solid, separate as a liquid, or bubble out as a gas |
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miscible definition |
liquids that are soluble in each other in all proportions (ex. alcohol and water) |
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immiscible definition |
liquids that are not soluble in each other in all proportions (ex. oil and water) |
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like dissolves? |
like |
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polar solutes dissolve in what type of solvent |
polar solvents |
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nonpolar solutes are more soluble in what type of solvent |
nonpolar solvents |
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organic compounds are dissolved in what type of solution |
hydrochloride (HCl) |
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heat of solution (aka enthalpy of solution) |
associated energy change when a solute dissolves in a solvent |
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change in Hsoln = |
energy change associated with dissolving one mole of solute in a given solvent |
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endothermic |
energy flows into the system |
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exothermic |
energy flows out of the system |
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enthalpy (heat) |
is equal to energy when the pressure remains constant |
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effect of pressure on solubility (of gaseous solutes in liquid solvents) |
increase pressure = increase solubility |
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effects of pressure on solids and liquids |
solids and liquids are not very compressible = pressure has little effect |
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Henry's Law |
quantitative relationship between pressure and solubility |
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Henrys law equation |
S = kH(aka henrys constant) x Pgas |
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henrys law relationships |
increase pressure = increase solubility |
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relationship between temp and solubility (solids and liquids) |
solubility of solid and liquid solutes in liquid solvents increases with increasing temperature |
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effect of temp on solubility (gas and liquid) |
solubility of gaseous solutes in liquid solvents decreases with increasing temperature due to increasing vapor pressure |
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colligative properties |
physical properties/characteristics of a solution that depend on the ratio of the number of solute particle to solvent particles |
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4 common colligative properties |
1. vapor pressure 2. boiling point 3. freezing point 4. osmotic pressure |
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vapor pressure = |
pressure exerted by a vapor (gas) in contact with its liquid or solid form |
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the most energetic molecules in a liquid has sufficient what in order to do what? |
they have sufficient kinetic energy to overcome the intermolecular forces binding them into the liquid state and can escape into the gas phase |
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effect of concentration on vapor pressure |
increased solute concentration = decrease vapor pressure |
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Raoults Law |
qualitative relationship between vapor pressure and solute concentration |
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Raoults Law equation |
P = X ( mole fraction of the substance) x P^degree (vapor pressure of the substance) |
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boiling point |
the temp at which the bulk of a liquid converts to vapor at a given pressure = the temp at which the vapor pressure = atmospheric pressure |
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boiling point relationships |
increase solute concentration = increase boiling point |
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freezing point (aka melting point) |
temp at which solid state reversibly passes into liquid state |
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freezing point relationships |
increase solute concentration = decrease freezing point |
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osmosis |
diffusion of water through a semipermeable membrane |
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osmotic pressure relationship with solute concentration |
increase solute concentration = increase osmotic pressure |
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movement of water in osmosis goes from where to where |
water diffuses from area of high to area of low (water or water diffuses from area of low (solute) to high (solute) |
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tonicity |
relative concentration of solutes in the osmotic system |
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isotonic |
two solutions with equal concentration solute particles |
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hypertonic |
solution with higher concentration solute particles |
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hypotonic |
solution with lower concentration solute particle |
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osmosis and diffusion explained by |
second law of thermodynamics = entropy of the universe is constantly increasing |
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colloids |
one phase uniformly dispersed in a second phase |
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difference between colloids and solutions |
contain large particles (not molecules or ions) |
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colloid exhibit what |
tyndall effect: scatter visible light passing through them |
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kinetic molecular theory of matter
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attempts to describe all states of matter and the conversion between states by 1. considering the structures of the molecules 2. how the molecules interact |
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common states of matter |
solid, liquid, and gas |
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kinetic energy by state of matter |
(least KE) solid |
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condensed states of matter = |
solids and liquids = resist compression/not easily compressed |
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Solids have |
a defined shape and volume; molecules held together by intermolecular forces |
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Liquids are fluids they have |
defined volume but no defined shape, they conform to the shape of the container in which they are placed |
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which has weaker forces liquids or solids? |
liquids, it allows the particles to flow past each other |
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solids can do what |
vibrate in place within a limited area |
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Gases = fluids = |
defined volume but no defined shape, they conform to their container (expand) in which they are placed |
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intermolecular forces present in gases? |
essentially no intermolecular forces between particles, this allows particles to flow past each other |
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melting |
conversion from solid to liquid |
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freezing |
conversions from liquid to solid |
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vaporization |
conversion of liquid to gas |
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condensation |
conversions of gas to liquid |
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deposition |
conversion of gas to solid (aka snow) |
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sublimation |
conversion of solid to gas |
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chemical bonds |
hold atoms together to form molecules |
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intermolecular forces |
determine how molecules interact with each other |
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Coulomb's Law |
particles (atoms, molecules, oppositely charged ions) are attracted to each other unless they get too close to each other then they repel each other |
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two limiting types of chemical bonds |
ionic and covalent bonds |
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octet rule |
overarching force that drives formation of chemical bonds, every atom adds, removes, or shares electrons to end up with 8 in their valence shell |
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atoms with full electron shells are |
very stable (aka noble gases) |
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elements with nearly full outer shells |
accept electrons and become anions (right side of the table) |
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elements with emptier shells |
give away electrons and become cations (left side of the table) |
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ionic compounds (aka salts) |
ions held together by ionic bonds (contain cations and anions but are electrically neutral) NOT MOLECULES! |
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salts |
ionic compounds - form highly organized crystalline lattice in solid state, almost all are solid at room temp and pressure |
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salts in liquid |
conduct electricity or dissolve in water |
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which bonds are stronger ionic or covalent? |
ionic bonds are stronger |
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covalent bonds result |
from sharing one or more pairs of electrons |
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valence bond theory |
the presence of electrons between the nuclei shield them from each other = reduces coulombic repulsive forces |
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shared electrons are able to orbit either nuclei |
= increases the electron density between the two nuclei; this stabilizes the molecule by reducing overall energy |
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decreased energy = |
increased stability when forming covalent bonds |
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there is an ideal bond length or distance where |
the energy is minimum and stability is maximum |
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energy increases if |
distance increases or decreases |
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bond dissociation energy |
amount of energy needed to break the covalent bond |
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strength of bonds vary |
stronger bonds are harder to break and are more stable |
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electronegativity |
atom's propensity for pulling electrons toward itself |
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most electronegative element
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Fluorine (F); the closer to fluorine an element is on the table the more electronegative it is |
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non polar covalent bonds |
bond between two atoms with the same electronegativity |
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non polar covalent bond electrons are |
evenly shared, and have no areas of charge, hemodiatomic molecules are always non polar (aka H2 or O2) |
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polar covalent bonds |
bond between two atoms with different electronegativity |
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increased differences in electronegativity = |
increasingly polar covalent bond |
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the more electronegative atom = |
the more pull it has, electrons spend more time here, will have partial negative charge and other atoms will have partial positive charge |
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intermolecular forces (4) |
1. dipole-dipole attraction 2. hydrogen bonding 3. London forces 4. ion-dipole interactions |
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dipole-dipole attraction |
attraction between the opposite partial charges (poles) of polar molecules |
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hydrogen bonding |
only possible when hydrogen is bonded to oxygen, nitrogen, or fluorine; H becomes positive charge and all partial negative charges are highly attracted to it |
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which bond is very important for intermolecular forces in life? |
hydrogen bonds |
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London Forces (weakest intermolecular force) |
instantaneous dipole created whenever electrons in an atom or molecule are unevenly distributed |
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what increases an atoms London forces? |
size, the larger the atom the more London force occurs |
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ion-dipole attraction |
occurs between ions and polar molecules |
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ion-dipole attraction strength depends on |
the strength of the charges of the poles and the size of the ion
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ion-dipole attraction allows |
ionic solids to dissolve in water, NaCl for example |
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polar compounds dissolve in |
polar compounds |
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non polar compounds are insoluble in |
polar compounds |
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non polar compounds dissolve easily in |
nonpolar compounds |
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ionic compounds dissolve easily in |
polar solvents |
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surface tension |
molecules near or on the surface are subject to surrounding forces and forces from below = unbalanced force, which creates "skin" on the surface of the liquid |
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substances with increased intermolecular attractions have |
increased surface tensions |
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if somethings density/weight/force is less than the surface tension of water |
it will not sink in water |
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water is very cohesive liquid = |
water molecules have a strong attraction for other water molecules, causes water to form spherical drops |
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surface tensions does what in the lungs |
|
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capillary rise |
caused by surface tension = tendency of a fluid to rise in a narrow diameter tube |
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LaPlace's Law |
describes relationship between wall tension to pressure and the radius in cylinders and spheres |
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LaPlace's Law equations |
cylinders = T = Pxr or P = T/r spheres = 2T = Pr pr P = 2T/r |
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LaPlace's Law relationships |
increased tension = increased pressure increased radius = increased tension increased radius = decreased pressure |
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Cylinder and sphere examples |
cylinders = blood vessels and aortic aneurysms spheres = saccular aneurysms and alveoli and cardiac chambers |
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Law of Laplace of hollow viscera |
P = 2T/R where P is pressure required to keep the viscera expanded, and T is surface tension and R is radius; surfactant decreases surface tension |
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surfactants |
reduce the surface tension of a liquid and increase the ability of the liquid to function as a solvent |
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surfactants are commonly known as |
detergents or soaps, have a polar head and non polar tail; forming monolayers, bilayers, and micelles |
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monolayers |
polar head goes into water and non polar tail sticks out of or sits on surface of water |
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surfactant molecules get between the water molecules and disrupt their hydrogen bonding = |
decreased surface tension of water, aka water is less cohesive |
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surfactant in lungs creates |
monolayer in the alveoli to keep from collapsing and to equilibrate the pressure needed to inflate alveoli of diff sizes |
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bilayers |
tails of surfactant molecules dissolve into each other to form a double layer |
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bilayer set up |
non polar tails form middle of the bilayer and polar heads form outside surfaces of bilayer; preventing polar compounds from penetrating the bilayer without carrier proteins |
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Micelles |
nonpolar tails dissolve in each other forming a spherical structure; how soaps and detergent work; creates non polar microenvironment in water |
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viscosity = |
measure of a fluids resistance to flow |
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fluids with high viscosity |
have increased resistance to flow |
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viscosity increases |
as intermolecular forces increase |
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vaporization requires |
energy |
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increased temp = |
increased vapor pressure |
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volatility = |
tendency of a liquid to evaporate |
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increased volatility = |
increased vapor pressure |
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dynamic equilibrium = |
system comprised of at least 2 states when the populations of the states are constant; in a closed system rates of vaporization and condensation are equal; equilibrium is achieved |
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substances with greater intermolecular forces have |
lower vapor pressure |
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molecules with increased intermolecular forces are |
more tightly bond in the liquid state, cannot escape not enough kinetic energy |
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volatile liquids = |
liquids that have a high vapor pressure at room temp (all our volatile anesthetics) |
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if pressurized anesthetic gases flow through a closed container holding a liquid volatile anesthetic |
the anesthetic gases will mix with the gaseous portion of the volatile anesthetic and exit the vaporizer |
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the temp of the vaporizer and the pressure of the anesthetic gas determine |
how much volatile anesthetic in its gas phase exits the vaporizer (this is also determined by selecting a setting on the vaporizer); each agent has different vapor pressure so must be calibrated for that specific agent or can under or over dose |
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boiling point = |
the temp at which the bulk of a liquid converts to vapor at a given pressure; temp at which the vapor pressure is equal to atmospheric pressure |
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the temp will never rise above boiling point |
the rest of additional heat energy transforms the liquid into the gas |
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the heat energy is transformed into molecules as |
kinetic energy = molecules are able to escape liquid phase and move into gas |
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increased intermolecular force = |
increased boiling point |
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boiling point increases as |
pressure increases (how pressure cooker work) |
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melting point = |
temp at which a solid state reversibly passes into liquid state (also known as freezing point) |
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melting point temp is the temp at which average kinetic energy is |
sufficient to overcome intermolecular forces that hold the molecules in the solid state |
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increased intermolecular forces do what to melting point |
increased melting point |