Honors Chemistry - Chapter 12 & 13
Honors Chemistry - Chapter 12 & 13
Terms in this set (91)
homogeneous mixture of two or more substances in a single phase
suspension (definition and particle size)
a mixture in which particles in a solvent are so large that they settle out unless mixture is constantly stirred (particles over 1000 nm in diameter form suspensions)
particles that are intermediate in size between those in solutions and suspensions form mixtures known as colloidal dispersion on colloids
What is the following an example of? - after soil particles settle out of muddy water, water still remains cloudy
colloidal particles make up ____ _____ and water is the _____ ______
dispersed phase and dispersing medium
what is the dispersed phase and the dispersing medium?
dispersed phase = what colloidal particles make up and water is dispersing medium
What is a headlight beam in a foggy night an example of?
What is the Tyndall Effect?
when light is scattered by colloidal particles dispersed in transparent medium; a property that can be used to distinguish between solution and colloid; the phenomenon in which light is scattered by very small particles in its path
substance that dissolves in water to give a solution that conducts an electric current (ex: NaCl); when ionic compound dissolves in water the positive and negative ions separate from each other and are surrounded by water molecules
what type of molecular compounds are electrolytes?
highly polar molecular compounds (group 1 and halogens) with the exception of HF
substance that dissolves in water to give a solution that doesn't conduct electricity (ex: sugar)
what is a substance that does not conduct electricity well when dissolved water?
creating ions when there weren't any (acids)
no ions were created, but ions can move around (slats and hydroxides-bases); mobilizes ions; there are already ions present but now they can move around
property of solution, dependent on how many solute particles are dissolved in the solvency; melting point, boiling point, vapor pressure, osmotic pressure
if salt is dissolved in water, the freezing point will ____ and boiling point will ____ and ____ the vapor pressure
decrease, increase, increase
equation for molarity
M = n (solute) / volume of solution (liters); number of moles of solute in one liter of solution
equation for molality
m = n (solute) / solvent (kilograms); concentration of a solution expressed in moles of solute/kilogram of solvent
What are the factors that affect the rate of dissolution?
1. increasing surface area of solute
2. agitating a solute
3. heating a solvent
How does increasing the surface area of the solute affect the rate of dissolution?
Increasing the surface area of the solute increases the rate of dissolution; ex: sugar dissolves as sugar molecules leave crystal surface and mix with water molecules, because dissolution process at surface of solute, can be sped up if surface area of the solute is increased (ex: crushing sugar in cubes)
How does agitating a solute affect the rate of dissolution?
Agitating a solute increases the rate of dissolution; close to surface of solute, concentration of dissolved solute is high, stirring or shaking disperses the solute particles
How does heating a solvent affect the rate of dissolution?
Heating a solvent increases the rate of dissolution; temperature increases --> solvent molecules move faster --> average kinetic energy increases --> at higher temperature the collisions occur more frequently and of higher energy
There is a limit to the amount of solute that can be dissolved for
every combination of solvent with solid solute at a given temperature
when sugar is added to war ---> sugar molecules leave solid surface and do random movement in solid OR collide again with crystal and remain there --> more and more sugar added --> molecules returning to crystal at same rate as going into solution --> dynamic equilibrium between dissolution and crystallization
physical state at which opposing processes of dissolution and crystallization of solute occur at same rate
saturated solution (define and how can you tell)
solution with max amount of dissolved solute; you can tell if a solution is saturated by adding more solute - if the solute sinks to the bottom then the solution is saturated
If you want to add more solute in a saturated solution, and want the solute to dissolve, what do you have to do to the saturated solution?
add more water;dilute it
solution that contains less solute than a saturated solution under existing conditions
contains more dissolved solute than a saturated solution contains under the same conditions, may remain unchanged if left undisturbed but once crystals start to form, processes continues until equilibrium is reestablished at a lower temperature
the solubility of a substance = amount of that substance required to form a saturated solution with specific amount of solvent and specified temperature (temperature must be specified, for gasses - pressure must also be specified)
what is the general rule for solute to solvent interactions?
"like dissolves like"
What happens at the molecular level when LiCl is dropped in water?
slightly charged particles of water molecules attract ions in ionic compounds and surround them; Li(+) is attracted to (-) ends of H20 and CL(-) attracted to (+) ends of H20; this attraction is strong enough to break the solute from its crystal surface
solution process with water as solvent, ions are said to be hydrated
What are the effects of pressure on solubility? (important for gasses not liquids) - there are three
1. gas in contact with surface of liquid - gas molecules can enter liquid
2. as amount of dissolved gas increases, some molecules begin to escape back to gas form
3. equilibrium established, if left undisturbed - solubility of gas in liquid is unchanged at given pressure
How does gas solubility increase?
as pressure increases/molecules collide with liquid surface more often --> rate of gas molecules entering solution increases --> rate at which they escape increases --> equilibrium is restored at higher gas solubility
solubility of a gas in a liquid is directly propertional to partial pressure of that gas on the surface of the liquid
Temperature and Gas Solubility have what kind of relationship and why?
temperature and gas have an inverse relationship; when temperature increases --> average kinetic energy increases --> more number of solute molecules can escape from attraction of solvent molecules and return to gas phase --> equilibrium is reached with fewer gas molecules in solution
enthalpy of solution:
net amount of energy absorbed as heat by solution when a specific amount of solute dissolves in a solvent
What has happened if the enthalapy of solution is negative, what has happened if it is positive?
If the enthalapy of solution is negative, energy has been released; if the enthalapy of solution is positive, energy has been absorbed
Why is the dissolution of gasses an exothermic process?
because heat causes a decrease in gas solubility
What is happening at molecular level for enthalapy of a solution with gas?
in gas state, molecules are so far apart that there are virtually no intermolecular forces of attraction between them; solute-solute interaction has little effect on enthalapy of a solution of gas
Why is energy released when a gas dissolves in liquid?
because attraction between solute gas and solvent molecules outweighs energy needed to separate solvent molecules
amount of solute in given amount of solvent/solution
Molarity and Molality are methods of expressing what?
separation of ions that occurs when an ionic compound dissolves (ex: NaCl dissolving in water)
General Solubility Guidelines (Mr. Crumm will provide on test):
1. Sodium, Pottasium, + Amonium compounds (Na+, K+, NH4+) are SOLUBLE in H20
2. Nitrates, acetates, and cholrates are SOLUBLE (No3-, CH3COO-, CLO3-)
3. Most chlorides are SOLUBLE EXCEPT those of Silver, Mercury I, and Lead (Ag, Hg2Cl2, Pb); Lead II is soluble in hot water (PbCl2)
4. Most Sulfates ((SO4) 2- charge) are SOLUBLE; EXCEPT those of barium, strontium, lead, calcium, and mercury (Ba, Sr, Pb, Ca, Hg)
5. Most carbonates ((CO3) with 2- charge), phosphates ((PO4) with 3- charge), and silicates ((SiO4) with 4- charge) are INSOLUBLE; EXCEPT those of sodium, amonium, and potassium (Na, K, NH4)
6. Most sulfides are INSOLUBLE; EXCEPT those of Ca, Sr, Na, K, NH4)
it two different soluble compounds are mixed and you get a combination of ions that forms an insoluble compound what has occurred?
double displacement and precipitation
when does precipitation occur?
when attraction between ions is greater than the attraction (between ions and surrounding water molecules);
net ionic equation;
includes only those compounds and ions that undergo a chemical change in a reaction in an aqueous solution (Crumm's test - ones that form precipitate - the solid)
ions that do not take part in a chemical reaction and are found in solution both before and after the reaction
ions are formed from solute molecules by the action of the solvent; creation of ions when there where none
how is ionization different from dissociation?
when ionic compound (ex: NaCl) dissolves, ions were already present separate from one another, when molecular compound dissolves and ionizes in a polar solvent, ions form
what provides energy needed to break covalent bonds during ionization?
energy released as head during hydration of ions provides energy needed to break covalent bonds
what determines the extent to which the solute ionizes?
bond strength determines the extent to which solute ionizes
hydronium ion, example of when it forms
H30 ion; H20 + HCl --> H30 + Cl
combination of hydrogen and any halogen (group 17 element)
What are hydrogen hallides?
Hydrogen hallides are all
1. molecular compounds with single polar covalent bons
2. all are gasses
3. are very soluble in H20
4. all are electrolytes
What are the six strong electrolytes that Mr. Crumm listed in class; Give an example of a weak electrolyte
HI, HBr, HCl, HNO3, H2SO4, HCL04 are strong electrolytes; HF is a weak electrolyte because it weakly conducts electric current
What determines whether an electrolyte is strong or weak?
whether they conduct electricity poorly or strongly which is determined by ITS ABILITY TO FORM IONS IN A SOLUTION
any compound whose dilute aqueous solutions conduct electricity well; due to presence of all or almost all of dissolved compound in forms of ions
any compound whose dilute aqueous solutions conduct electricity poorly; due to presence of small amount of dissolved compound in form of ions (ex: HF)
Why is HF a weak electrolyte?
H-F bond is much stronger than bonds between H and other halogens, when HF dissolves, some molecules inoize but reverse action, transfer of H+ ions back to F- to form HF also takes place; thus concentration of dissolved HF is greater than concentration of H30 + F ions
examples of weak electrolytes:
HF and CH3COOH
what is the difference between ionization and disassociation versus concentration and diluision
strong and weak electrolytes differ in degree of ionization/disassociation; very dilute solution w?NaCl can still be a strong electrolyte
properties that depend on the concentration of solute particles but not on their identity (concentration given sometimes in molality)
True or False: boiling point and freezing point of solution differ from pure solvents
one that has little tendency to become a gas under existing conditions
novolatile solute has a ___ boiling point and a ___ freezing point
pressure caused by molecules in gas phase that are in equilibrium with liquid phase
why is the vapor pressure of solvent containing nonvolatile solute lower than vapor pressure of pure solvent
because as solute increases --> water proportion decreases --> less H20 molecules are able to escape --> tendency of H20 molecules to enter vapor pressure decreases --> vapor pressure of solution decreases
why is vapor pressure a colligative property?
because vapor pressure lowering depends on concentration of a nonelectrolyte solute and is independent of solute identity
why does freezing point decrease and boiling point increase (w/vapor pressure)
as vapor pressure decreases --> solution liquid over a larger temperature range --> lowering of a freezing point and increasing of a boiling point
why are boiling point and freezing point colligative properties?
because they depend on the concentration of solute
molal freezing point constant (K):
freezing point depression of solvent in a 1 molal solution of nonvolatile, nonelectrolyte solute; 1.86 degrees celcius/molal
freezing point depression:
difference between freezing poitns of pure solvent and a solution of a nonelectrolyte in that solvent and is directly proportional to molal concentration of solution
How does change in vapor pressure affect change in boiling point?
more energy will be required to raisethe vapor pressure of the solution to equal the atmospheric pressure --> boiling point of solution is higher than the boiling point of the solvent
molal boiling point constant:
boiling point elevation of the solvent in a 1 molal solution of a non volatile, electrolyte solute = 0.51 degrees celcius/molal
boiling point elevation:
difference between boiling points of pure solvent and a nonelectrolyte solution of that solvent, and is dierctly proportional to molal concentration of solution
allows passage of some particles while blocking passage of others; (ex: allows H20 molecules but not sucrose molecules to pass through, sucrose molecules on solution side allow fewer water molecules leave pure H20 is greater than the reate at which they leave solution --> level of solution increases --> increases until pressure exerted by height of solution is large enough to force water molecules back through the membrane at equal rate to that at which they enter solution from pure H20 side)
movement of solvent through a semipermeable membrane from the side of lower solute concentration to side of higher solute concentration; occurs when solutions of two different concentrations are separated with a semi-permeable membrane
external pressure that must be applied to stop osmosis, (in sucrose-water example), height of solution is osmotic pressure, because osmotic pressure depends on concentration fo solute particles = colligative property
why does boiling point increase and freezing point decrease a lot more with an electrolyte?
because (ex: NaCl) produces more ions, one molecular compound breaks up into TWO ions --> produces more moles than (ex: sucrose)
changes in colligative properties caused by electrolytes will be proportional to total
molality of all dissolved particles
what is the difference between expected and actual values for Electrolyte Solutions caused by?
diff is caused by attractive forces between disassociated ions in aqueous solutions - interfere with movement
attraction between disassociated ions in aqueous solutions because ionic atmosphere that surrounds each ion, cluster of hydrated ions acts as a single unit, effective total concentration is less than expected
USE SOLUBILITY GUIDELINES FOR NET IONIC REACTION PROBLEMS
USE SOLUBILITY GUIDELINES FOR NET IONIC REACTION PROBLEMS
a nonelectrolyte is a what compound
a nonelectrolyte is an organic compound; organic compounds generally do not break up into ions (C, O,H, N) = organic compound elements
endothermic reactions dissolve more readily in
When crystals start to form, is an exothermic or endothermic process occurring?
why don't crystals form in supersaturated solutions when left undisturbed?
crystals need to be able to form around particles, in a supersaturated solution - all particles are dissolved, that's why disturbing the supersaturated solution with particles would cause crystals to form
molar fraction equation:
x with curly tail (solute) = number of moles (solute) / number of moles (solute) + number of moles (solvent)