[FINAL] All formulas and when to use Flashcards
Density
Q: What is the mass of a 10.00 mL sample of water at 25 degrees Celsius? The density of water at 25 degrees is 0.99707 g/cm^3.
Mass/Volume
A: 99.71 (pg 9)
Kelvin to Celius
Q: At Mt. Everest water boils at about 70 degrees Celsius. What is the temperature in Kelvin?
Kelvin= celsius +273
A: 343 K (pg 11)
Celsius to Fahrenheit
Q: What is the temperature of a human body, 36.6 degrees C in degrees F?
C= 5/9(F-3)
F=9/5(C) +32
A: 97.9 F (pg12)
Mass number
Q: How many neutrons are in Nickel 59?
Mass Number, A= # of protons + # of neutrons
A: 31
Atomic number
Atomic #, Z= number of protons which = number of electrons
Average atomic mass
Q: What is the avg. atomic mass of carbon?
M1= 12, f1=.9889, M2= 13.003355, F2= .0111
Average Atomic mass= (fractional abundance (f)1 * isotpic mass (M)1) + (fractional abundance (f)2 * isotpic mass (M)2)
A: 12.011 (pg 26)
Avogadro’s number
Q: How many atoms of Ar are contained in .345 mole of Ar?
Avogadro’s number= 6.022 *10^23 which equal 1 mole of an atom or molecule
A: 2.08 *10^23 atoms (pg 42)
mols
Q: how many mols of NH3 ( Mm= 17.03 g/mol) are there in 84.12 g of NH3?
mols= mass/ molar mass
A: 4.939 mol NH3 (pg 44)
Mass percent of element X in a compound
Q: calculate the mass percent of H in H2O
Mass percent of element X in a compound= (mass of element X/ mass of 1 mol of the compound)
A: 11.189% (pg 45)
Sample Question: How many grams of carbon dioxide (CO2, Mm= 44.010 g/mol) can be prepared from the complete combustion of 3.25 g of acetylene (C2H2, Mm = 26.038 g/mol)
2C2H2 (g) > 4CO2 (g) + 2H2O (g)
A: 11g
pg 52-53
Percent Yield
Q: A reaction theoretically can produce 7.65 g of a product, however, only 5.67 g have been obtained in the laboratory. What is the percent yield of the product in this reaction
Percent Yield= (actual yield/ theoretical yield) * 100%
A: 74.1 % (pg 62)
Sample Question: write the net ionic equation for Ca(NO3)2 (aq) + Na2CO3 (aq) > CaCO3 (s) + 2NaNO3 (aq)
A: Ca^2+ (aq) + CO3^2- (aq) > CaCO3 (s)
pg 72 and review page soluble ionic compounds on pg 67
Molarity
Q1: What is the molar concentration of .5 L solution containing 4.5 moles of C6H12O6?
Q2: How many grams of NaCl are there in 1.00 L of the solution which is .100 M?
Molarity= moles of solute (mol) / Volume of solution (L)
symbols for Molarity: C, [solute], M
A1: 9.00 M (pg 79)
A2: 5.84 g (pg 80)
mols in terms of molarity
Q: What mass of NaOH is needed to precipitate Cd^2+ ions from 20.0 mL of .100 M of Cd(NO3)2 solution?
Cd(NO3)2 + 2NaOH > Cd(OH)2 + 2NaNO3
mols = Molarity * Volume
A: .160 g NaOH (pg 85-86)
Sample Question: How many moles of methane were combusted at a constant pressure if 180 kJ of heat were released?
CH4 + 2O2 > CO2 + 2H2O; delta H = -890 kJ
A: .2 mol CH4 (pg 93)
Frequency (Hz)
The speed of light = (3*10^8)
Q: An FM radio broadcasts at 89.3 MHZ. Calculate the corresponding wavelength.
Frequency (Hz)= Speed of light (3*10^8) / wavelength (m)
A: 3.36 m (pg 97)
The energy of a photon
(Planck’s constant = 6.626* 10^-34)
Q: What is the energy of a photon of yellow light with a wavelength of 585 nm?
Energy (J) = Planck’s constant * frequency
A: 3.40 * 10^-19 (pg 100)
another formula for energy of photon if given wavelength but not frequency.
The energy of a photon= (Planck’s constant * speed of light) / wavelength (pg 100)
The energy of an electron in an H atom
Q: What is the energy of the state of an H atom in the 2nd state?
Energy of an electron = - (2.18 *10^-18) / (Prinicipal quantum number (n))^2
A: -5.45 * 10-19 (pg 103)
Sample question: What is the wavelength of light emitted when an electron in a hydrogen atom falls from n=2 to n=1 state?
A: 121 nm (pg 104)
How to calculate the formal charge (f.ch)?
Q: What is the f.ch of an oxygen atom with two lone pairs and one double bond?
f.ch = # of valence electrons - (# of lone pairs + (1/2 * # of bonding electrons)
A: 0 (pg 145)
Ideal Gas Law equation
gas constant (R)= .08206 (L * atm)/ (mol * K)
or 8.314 (J) / (mol * K)
Q: What is the volume of 1.00 mole of gas at 1.00 atm pressure at 0 degrees Celsius?
PV=nRT
A: 22.4 L (pg 187)
Combined gas law
Q: What is the temperature of 1.22 m^3 of a certain gas at 99,999 Pa if at 35 degrees Celsius and 1.00 atm pressure it takes up a volume of 1.00 m^3?
(P1 * V1) / T1 = (P2 * V2) / T2
A: 370 K or 98 degrees C (pg 189)
Sample question: what is the molar mass of a gas if a .212 g sample takes up a volume of 99.056 mL at 122 degrees C and 1.02 atm pressure?
A: 68 g/mol (pg 190)
Total pressure
(R= .082057 (L atm)/(mol K)
Q: a sample of mixed gasses contains .34 mol N2, .88 mol CO2 and 1.00 mol of O2. It takes up a volume of 56 L at 273 K. What is the total Pressure?
Total pressure= total mols * ((R * temperature) / volume)
A: .88 atm (pg 193)
Mole fraction
Q: there are 2.33 moles of gas A and .670 moles of gas B, What are the mole fractions of each?
Mole fraction= moles of a gas/ total # of moles in gas
A: Gas A= .776 Gas B=.223 (pg 194)
Partial pressure
Q: What are the partial pressures of gases A and B in a mixture if total pressure is 1.00 atm and there are 2.33 moles of gas A and .670 moles of gas B?
Partial pressure = Mole fraction * total pressure
A: Gas A=.776 atm, Gas B=.223 atm (pg 194)
Grahams Law
What is the molar mass of an unknown gas if it effuses 1.3 times faster than propane C3H8 (Mm = 44.092)
(Rate of effusion 1/ rate of effusion 2) = Square root( Molar mass 2/ molar mass 1)
A: 26 g/mol (pg 198)
heat equations
c of ice= 2.03
c of water = 4.18
calculate the amount of heat needed to heat an 18.02 g piece of ice, initially at -12.0 degrees C to water at 67.0 degrees C?
heat= specific heat capacity(c) * mass * change in temperature
(this is for when the phase is increasing in temperature)
heat= mols * change in heat
(this is for when a phase is changing at a constant tempt)
A: 11.49 kJ (pg 213)
Coulomb’s Law
Force = ( k * charge 1(positive) * Charge 2(negative)) / (distance squared)
