AP Chem Ch 4-5 Flashcards
Hydration
Add water to something –> molecule bonds with water.
Heating, dehydration, reverses this.
Polar
Unequal distribution of charge. A net electronegative charge pointing in a direction
Solubility
The maximum amount of solute that dissolves into a solvent at equilibrium.
The measure of the amount of solute that can dissolve in a volume of solvent at a given set of conditions (temperature, pressure)
Solute
Thing being dissolved
Solvent
Thing being dissolved in – solute dissolved in the solvent.
Common solvent is water
Solubitility changes with temperature and stirring
Temperature up, solubility up
Stirring up, solubility up
Amount of solute up, solubility down- precipite
Solubility rules
See other flashcard deck
Super saturated
Helps make crystillation structure
Put solute into solvent, heat up and put lots in, then cool and makes crystal
Strong electrolytes
Soluble salts, strong acids and strong bases because they all three completely dissociate into ions when dissolved in water. Conduct electricity because of the ions
What do strong electrolytes do
100% dissolve in water. Carry an electrical current
Molarity
Number of mol / L
Moles of solute over the liters of the solution
If we have 11.5 g of NaOH and they molar mass is 40g/mol, what is the molarity when dissolved in 1.50 L of water?
11.5 g * 1 mol / 40 g = .288 mol
M = mol / L = .288/1.5 = .192 M
Completely dissociates in water
Example:
Say we have a solution of .50 M of Co(NO3)2
What is the concentration of each ion present
Co2+ –> M= .50
NO3 - –> M=.5*2=1.00
Example of dilution:
If we have 1 liter of 2.7 M of KMnO4, how could we make a 1.45 M solution and a .65 M solution?
Say we want .5 L solution of both 1.45 M and .65 M
M1V1=M2V2
2.7V = (1.45)(.5)
V= .27 L
Add .27 L of the 2.7 M solution and .23 L of water to make .5 L of the 1.45 M solution
2.7V= (.65)(.5)
V=.12 L
Add .12 L of the 2.7 M solution and .38 L of water to make .5 L of the .65M solution
What type of glassware is used to make series dilutions
Eurleymuer flask
Types of chemical reactions
Precipitation
Acid-base
Oxidation-reduction
What does silica gel dissolve/not dissolve in
Silica gel is not soluble in ethyl acetate
Aspirin solubility results
Soluble in ethyl acetate
Soluble in sodium bicarbonate
INSOLUBLE in HCl
Acetametaphine solubility results
Soluble in ethyl acetate, INSOLUBLE in sodium bicarbonate
Precipitation reaction
Formation of an insoluble substance
Filtrate –> solvent
Precipaite is formed
Example of precipaite reaction:
K2CrO4 (aq) + Ba(NO3)2 (aq)–> ?
K2CrO4 (aq) + Ba(NO3)2 (aq)–> 2KNO3 (aq) + BaCrO4 (s)
Types of formulas
Formula equation
Complete ionic equation
Net ionic equation
Formula equation
Overall reaction that includes everything going on
Comete ionic equation
Breaks down the aqueous materials into the ionic components
Net ionic equation
Cancels out the ions on both sides
Example of writing net ionic equation:
AgNO3 (aq) + KCl (aq)
AgNO3 (aq) + KCl (aq) –> AgCl (s) + KNO3 (aq)
Ag+ + NO3 - + K+ + Cl- –> AgCl(s) + K+ + NO3 –> complete ionic here
Cancel
Ag+ (aq) + Cl- (aq) –> AgCl(s)
Stoichiometry of precipitan reactions:
Example:
If we have 1.50 L of .1 M of AgNO3 and we add 8 grams of NaCl (.137 mol), what product will precipite and how much will form?
AgNO3 (aq) + NaCl (AQ) –> AgCl (s) + NaNO3 (AQ)
.1 M = x mol / 1.50 L.
X= .15 mol silver nitrate
NaCl is limiting reactant
So .137 mol of AgCl is formed, multiply by molar mass to get amount formed in grams
Arrhenius concept
Acids produce H+, bases produce OH-
Bronsted - Lowry definition
Acid is a proton donor
Base is a proton acceptor
Hydroxide with weak acid
Hydroxide is so strong as a bad that it completely reacts even with weak acids
Neutralization reactions:
What volume of .1M of HCl is required to neutralize 25 mL of a .350 M NaOH solutoon
HCl + NaOH –> H2O + NaCl (Aq)
1:1 ratio of HCl to NaOH –>
.350 M = x mol / .025 L
Get moles then need same moles of HCl and use M = mol / liter to do so
Titration
When you add a titrant –> base to acidic solution or an acid to a basic solution, observe how the pH changes as you add more of the titrant
The equivalence point is the point at which neutral pH. Endpoint when it is base if started acid or reverse if started base.
Reduction
ADDING ELECTRONS
CHARGE IS REDUCED IN REDUCTION
Oxidation
Giving away elections
The charge is increased –> less negative charge
Oxidizing agent
Facilitates oxidation, but it is reduced
Reducing agent
Facilitates reduction but it is oxidized
Rules for assigning oxidation numbers
Florine always -1
Oxygen -2
Hydrogen +1
What’s being reduced and what’s being oxidized:
Cr+ + Sn4+ –> Cr3+ + Sn2+
Cr is losing electrons, so it is being oxidized
Sn is gaining electrons, so it is being reduced
How to balance redox reactions:
Example:
Bi(OH)3 + SnO2 2- –> SnO3 2- + Bi in an acidic solution
First identify what’s being reduced and what’s being oxidized
Bi is going from 3+ to 0 –> it is being reduced
Sn is going from 2+ to 4+ –> it is being oxidized
Now write the half reactions:
3e- + Bi(OH)3 –> Bi
Now make sure the hydrogens and oxygens are balanced as well as the charge on each side
3e- + Bi(OH)3 + 3H+ –> Bi + 3H2O
Now do the same for the other half reaction:
H2O + SnO2 2- –> SnO3 2- + 2e- + 2H+
Now multiply top by 2 and bottom by 3 to get it so the electrons cancel out and write the final equation
H+ may or may not cancel
If in basic, add OH- and form water to one side and keep the OH- on the other side
Equilibrium
State where the rate of the forward reaction is equal to the rate of the reverse reaction
Lattice
The pattern for which a solid will crystallize.
Energy is required to break this lattice apart. Need the break the intra molecular forces that keep the molecule together.
Entropy
Measure of randomness – disorder and chaos occurring
Miscible substances
Substances that can mix, like alcohol and water, are miscible. Measure of solubility cannot be applied for miscible
Immiscible
Materials that do not mix together, like oil and water. Again solubility can’t be applied here
Why do some things dissolve and others don’t?
When dissolved, solvent molecules surround solute molecules. Breaks the intermolecular bonds of the individual molecules (Na and Cl ions being pulled apart by water)
Need energy to break these bonds.
Happens because entropy wants the world to move toward disorder and mixing increases the entropy, causing a drive to increase entropy and therefore free energy
Other factors to consider in solubility
Charge of the material being dissolved
Size of the molecule/compound – larger compound required more solvent per solute.
Coulomb’s Law
Opposite charges attract while alike charges repel.
F= kq1q2 / r^2
As distance between two charges increases, electrostatic forces decrease.
Force proportional to magn of both charges
Inverse square law
Force inversely prop to the square of the distance
How does Coulomb’s law apply to Chem?
Ionization energy
The protons in the nucleus attract the electrons – removing one will require some amount of energy– it also results in an increase in positive charge and will exert a stronger hold on the remaining electrons.
Ionization energy
The amount of energy required to remove electron.
First vs second vs third ionization energies
First is removal of one highest energy electron.
Second is removal of he next highest energy electron
Third is removal of the next highest energy electron
Trends with ionization energy
Group 1 wants to give up their electrons– lowest energy required to remove one
Group 17 doesn’t want to give up electrons – highest energy required to move one.
More energy required to remove inner electron than outer– easier to remove outer electrons.
So for magnesium, the first ionization energy will be bigger than the second because second is last electron in shell but the third ionization energy will be large because that goes inside the inner shell.
Photoemission spectroscopy
Shine a frequency on an atom and measure when an electron comes off.
Relative number of electrons vs energy
Less energy for outer electron shells.
2p less energy than 2s which is way less than 1s for nitrogen
Pressure
Force per area
Measured in mmHg, ATM, torr, pascals
Barometer
Glass tube with vacuum on end–> height mercury rises at ATM is 760 mmHg.
Why use mercury?
Because density – only rises 760 mm at ATM. Also a liquid at room temperature
Manometer
Also used to measure pressure but with a gas container on one end and open on the other end.
Pressure of the gas = Patm - h for gas that is less and +h for when it is more.
Pressure and volume relationship
P is proportional to 1/V
P = k/V
Pressure is inversely proportional to the volume.
P on Y axis and V on x axis produces 1/X curve.
V vs 1/P is linear.
Slope is k = PV. Constant PV
So, P1V1=P2V2
Boyles Law
If moles and temperature are constant, then P1V1=P2V2
Only true for pressures around 1 ATM
STP
0 Celsius, 273 kelvin, 1 ATM
Volume and temperature
Directly proportional
Charles Law
V= bT
V1/ T1 = V2/T2
Absolute zero
0 kelvin. No kinetic energy, no volume.
Avagadros Law
Gas at a constant t and constant p, volume directly proportional to number of moles of gas.
Molar volume of a gas at STP is always 22.42 L. Doesn’t matter the gas.
V=an
Ideal gas Law
Combine avagadros, charles, Boyle to make
PV=nRT
What will happen if water is boiled in a metal can and then the heat is turned off and the can is sealed?
Temperature and pressure down, so the can will be crushed as the water vapor molecules condense. Pressure from the atmosphere forces the pressure on the inside to collapse.
Kinetic molecular theory
- Particles are so small compared with distances between them that the volume of the individual particles of a gas can be assumed to be zero.
- The particles are in constant motion. Collisions of the particles with the walls of the container are the cause of the pressure exerted by the gas.
- Particles are assumed to exert no forces on each other. They are assumed neither to attract nor to repel each other.
- Average kinetic energy of a collection of gas particles is assumed to be directly proportional to the kelvin temperature of the gas.
How does energy relate to temperature?
KE average = 3/2 RT
Root mean square velocity
Urms = SQRT (average velocity squared) = SQRT (3RT/M)
M is the mass of a mole of gas in Kg
Use R = 8.314 J/molK
Example:
What’s the root mean square velocity of helium at STP
molar mass is 4 grams / mol
SQRT (3RT/M) = SQRT (38.314273/.004) = 1305 meters per second
Boltzmann distribution
Velocity is an average velocity
At a given temperature, the velocity varies
Higher temperature, higher average velocity and more spread out as well
Diffusion
Describes mixing of gasses
Effusion
The rate of a gas entering an evacuated chamber
Grahams law of effusion
Relates the root mean square of gases when comparing rate of effusion between two gases
Rate of effusion for gas 1 / rate of effusion for gas 2 = SQRT molar mass of 2 / SQRT molar mass 1
E1/ E2 = SQRT M2 / SQRT M1
Density relating to molar mass
Molar mass = dRT/P
Because n = m/molar mass
P= (m/molar mass)RT/V
P= dRT/ molar mass
Dalton law of partial pressure
Total pressure of a mixture of gasses is a result of the sum of the individual pressures.
P total = P1 + P2 + …
Example:
Say we have 46 L of He and 12 L of O2 both at 25 degrees Celsius and 1 ATM and they’re pumped into a 5 L tank also at 25 degrees Celsius
Find partial pressure of each gas and total pressure
Option 1. Calculate moles. Then use PV=nRT in new container to get total pressure and multiply P by mole ratio to get partial pressures.
Option 2. Use P1V1=P2V2. V2 is 5 for each gas.
Mole fraction
X= moles of 1 / total moles
