Quicksheets Gen Chem Flashcards
1
Q
how to calculate moles
A
moles = mass of sample/ molar mass
2
Q
A =
A
A = mass number = number of protons
3
Q
Z =
A
Z = atomic number = number of protons + neutrons
4
Q
Planck’s quantum theory
A
energy emitted as electromagnetic radiation from matter exists in discrete bundles called quanta
5
Q
Bohr’s model of the atom–how to calculate the energy of an electron
A
E = -Rh/n^2 Energy = -(Rydberg constant)/(principle quantum number^2) Rydberg constant = 2.18e-18J
6
Q
equation to find the electromagnetic energy of photons
A
E = hc/λ Energy = (planck’s constant)(speed of light)/(wavelength)
7
Q
Balmer series
A
n>2 to n = 2 Visible: 400-800 nm
8
Q
Lyman series correspond to
A
n >1 to n = 1 UV: 90-125 nm
9
Q
Heisenberg uncertainty principle
A
it is impossible to determine with perfect accuracy the momentum and the position of an electron simultaneously
10
Q
principle quantum number (n)
A
the larger the integer value of n, the higher the energy level and radius of the electron’s orbit; max # of electrons in energy level n is 2n^2
11
Q
azimuthal quantum number (l)
A
subshells l = 0,1,2,3 s,p,d,f max number of electrons that can exist w/in a subshell = 4l + 2
12
Q
Exceptions to the octet rule (5)
A
H(2), He(2), Li(2), Be(4), B(6) and atoms found in or beyond the 3rd period
13
Q
magnetic quantum # (ml)
A
the particular orbital w/in a sub shell where an electron is highly likely to be found at a given point in time
14
Q
spin quantum # (ms)
A
the spin of a particle is its intrinsic angular momentum and is a characteristic of the particle, like its charge
15
Q
Hund’s rule
A
orbitals are filled such that there are a max # of half-filled orbitals w/ parallel spins
16
Q
polar covalent bond
A
bonding not shared equally; pulled toward more electronegative atom
17
Q
regions of electron density: 2 Example: BeCl3 has ___ shape has ___ angle(s)
A
linear shape 180˚ angles
18
Q
regions of electron density: 3 Example: BH3 has ___ shape has ___ angle(s)
A
trigonal planar shape 120˚ angles
19
Q
regions of electron density: 4 Example: CH4 has ___ shape has ___ angle(s)
A
tetrahedral shape 109.5˚ angles
20
Q
regions of electron density: 5 Example: PCl5 has ___ shape has ___ angle(s)
A
trigonal bipyramidal shape 90, 120, 180˚ angles
21
Q
regions of electron density: 6 Example: SF6 has ___ shape has ___ angle(s)
A
octahedral shape 90˚, 180˚ angles
22
Q
lewis acid
A
accepts electrons
23
Q
lewis base
A
donates electrons
24
Q
ligands
A
coordinate bonding attaches ion/molecule with metal atom
25
describe coordinate covalent bonds
covalent bonds in which both electrons in the bond came from the same starting atom
26
describe chelation
when the central cation can be bonded to the same ligand in multiple places; it generally requires large organic ligands that can double back to form a second (or even 3rd bond) with the central cation
27
chelation therapy is used to
often used to sequester toxic metals (lead, arsenic, mercury)
28
Describe what happens in hydrogen bonding
The partial positive charge of the H atom interacts with the partial negative charge located on the electronegative atoms (FON) or nearby molecules
29
describe how dipole-dipole interactions occur
polar molecules orient themselves such that the positive region of one molecule is close to the negative region of another molecule
30
dispersion forces
the bonding electrons in covalent bonds may appear to be equally shared b/t 2 atoms, but at any particular point in time they will be located randomly through the orbital permitting the unequal sharing of electrons, causing transient polarization and counterpolarization of the electron clouds of neighboring molecules, inducing the formation of more dipoles
31
a mole is...
the amount of substance that contains the same # of particles that are found in a 12 g sample of carbon-12
32
molecular/formula weight is measured in
amu per molecule (formula unit)
33
molar mass is measured in
grams per mole
34
What happens in combustion reactions
a fuel (like a hydrocarbon) is reacted with an oxidant (like oxygen) to produce an oxide and water CH4(g) + 2O2(g) -\> CO2(g) + 2H2O (g)
35
What happens in combination reactions
two or more reactants form one product S(s) + O2(g) -\> SO2(g)
36
What happens in decomposition reactions
a compound breaks down into 2 or more substances, usually as a result of heating or electrolysis 2HgO(s) -\> 2 Hg(l) + O2(g)
37
What happens in Single-displacement rxns
an atom (or ion) of one compound is replaced by an atom of another element Zn(s) + CuSO4(aq) --\> Cu(s) + ZnSO4(aq)
38
What happens in Double-displacement rxns
aka metathesis rxns elements from 2 different compounds displace each other to form two new compounds CaCl2(aq) + 2 AgNO3 (aq) --\> Ca(NO3)2 (aq) + 2 AgCl(s)
39
Net ionic equations show...
show only the species that actually participate in the reaction (doesn't show spectator ions) For example, in the reaction of: Zn(s) + Cu2+(aq) + SO4^2-(aq) --\> Cu(s) + Zn^2+ (aq) + SO4^2-(aq) Spectator ion = SO4^2- Net ionic equation: Zn(s) + Cu^2+(aq) --\> Cu(s) + Zn^2+(aq)
40
What happens in Neutralization reactions
These are specific double-displacement rxns that occur when an acid reacts with a base to produce a solution of a salt (and usually water) Hcl(aq) + NaOH(aq) --\> NaCl(aq) + H2O(l)
41
Factors affecting reaction rates (4)
reactant concentrations, temperature, medium, catalysts [conc],medium, temp, catalyst
42
What do catalysts do
catalysts increase reaction rate w/o being consumed; do so by lowering the activation energy
43
Law of mass action states that
aA + bB cC + dD Kc = [C]^c[D]^d/[A]^a[B]^b Kc = equilibrium constant
44
What happens when Keq \>\>1
an equilibrium mixture contains very little of the reactants compared to the products
45
What happens when Keq \<\< 1
an equilibrium mixture contains very little of the products compared to the reactants
46
What factors affect Le Chatelier's principle?
stresses include concentration, pressure, volume, or temperature
47
A + B C + heat what will make it shift to the right in terms of concentration, volume, pressure, and temp?
Right shift = \* more A or B added \* C taken away \* if pressure is applied or volume is reduced \* If temp is reduced
48
A + B C + heat what will make it shift to the left in terms of concentration, volume, pressure, and temp?
Left shift = \* If more C is added \* If A or B is taken away \* If pressure is reduced or volume is increased \* If temp is increased
49
What does the law of conservation of energy say?
It dictates that energy can be neither created nor destroyed, but that all thermal, chemical, potential, and kinetic energies are interconvertible
50
Describe an isolated system
in an isolated system, there is no exchange of energy/matter with the environment
51
bomb calorimetry has to do with what kind of system
bomb calorimetry = a nearly isolated system
52
Describe a closed system
A closed system can exchange energy but not matter with the environment
53
describe an open system
this system can exchange both energy and matter with the environment
54
Describe an isothermal process
constant temp
55
Describe an adiabatic process
no heat exchange
56
Describe an isobaric process
constant pressure
57
describe an isovolumetric process
constant volume
58
isochoric process
constant volume
59
Endothermic processes \_\_\_heat
endothermic = absorb heat
60
Exothermic processes \_\_\_\_heat
exothermic = release
61
what is constant-volume and constant pressure calorimetry used for?
used to indicate conditions under which the heatflow is measured
62
state functions
Properties' magnitudes that depend only on the initial and final states of the system, and not on the path of the change
63
common state functions include (8)
pressure, density, temp, volume, enthalpy, internal energy, free energy, and entropy
64
Enthalpy
(∆H) = expresses heat changes at constant pressure
65
standard heat of formation
(∆H°f) = the enthalpy change that would occur if one mole of compound was formed directly from its elements in their standard states
66
standard heat of reaction
(∆H°rxn) the hypothetical enthalpy change that would occur if the rxn were carried out under standard conditions
67
calculation for the standard heat of reaction
∆H°rxn = (sum of ∆H°f of products) - (sum of ∆H°f of reactants)
68
Hess's law
enthalpies of reactions are additive
69
bond dissociation energy
an average of the energy required to break a particular type of bond in one mole of gaseous molecules
70
bond enthalpy
the standard heat of rxn that can be calculated using the values of bond dissociation energies of particular bonds
71
Entropy
(S) it's the measure of the distribution of energy (randomness) throughout a system
72
∆Suniverse can be calculated by
∆Suniverse = ∆Ssystem + ∆Ssurroundings
73
Gibbs free energy equation and implications
∆G = ∆H - T∆S -∆ G = spontaneous +∆ G = nonspontaneous ∆ G = 0 then system is in a state of equilibrium and ∆H = T∆S
74
-∆H +∆S
spontaneous at all temps
75
+∆H -∆S
nonspontaneous at all temps
76
+∆H +∆S
spontaneous only at high temps
77
-∆H -∆S
Spontaneous only at low temps
78
1 atm = ? mmHg = ? torr = ? Pa
1 atm = 760 mmHg = 760 torr = 101,325 Pa
79
STP
0 C/273 K, 1 atm
80
standard conditions
25C/298K, 1 atm, 1 M concentrations
81
Boyle's law
PV = k P1V1 = P2V2
82
Charles's law
V/T = k V1/T1 = V2/T2
83
Gay-Lussac's law
P/T = k P1/T1 = P2/T2
84
Avogadro's Principle
n/V = k n1/V1 = n2/V2
85
Combined gas law
P1V1/T1 = P2V2/T2
86
ideal gas law
PV=nRT
87
what's the effect of decreasing the volume of a sample of gas?
Decreasing the volume of a sample of gas makes it behave less ideally b/c the individual gas particles are in closer proximity in a smaller volume, so they're more likely to engage in intermolecular interactions
88
Deviations due to pressure --as the pressure of a gas increases, what happens? (at moderately high pressure)
As the pressure of a gas increases, the particles are pushed closer and closer together. At moderately high pressure, a gas's volume is less than would be predicted by the ideal gas law due to intermolecular attraction
89
Deviations due to temp--as the temp decreases, the average velocity of the gas molecules does what? As the temp of a gas is reduced, intermolecular attraction (a) causes the gas to... At extremely low temps, what happens
temp of gas is reduced, gas has a smaller volume than would be predicted At extremely low temps, the volume of the gas particles themselves cause the gas to have a larger volume than predicted
90
Van der Waals equation of state
for deviations from ideality--when a gas doesn't closely follow the ideal gas law: (P+n^a/V^2)(V-nb) = nRT a = intermolecular attraction b = gas particles themselves
91
1 mole of gas at STP = ?
22.4 L
92
Dalton's law of partial pressures
the total pressure of a gaseous mixture is equal to the sum of the partial pressures of the individual components PT = PA + PB + PC + ... PA = PTXA XA = nA/nT (moles of A)/(total moles)
93
kinetic molecular theory of gases says that
gaseous molecular behavior based on the motion of individual molecules
94
average molecular speed equation
K = 1/2mv^2 = 3/2 KBT
95
Root mean square speed equation
urms=sqrt:(3RT/M)
96
colligative properties
physical properties dierived solely from the # of particles present and not the nature of those particles
97
freezing point depression
∆Tf = iKfm =(vant Hoff factor)(Kf)(molality)
98
Boiling point elevation
∆Tb = iKbm =(van't Hoff factor)(Kb)(molality)
99
Osmotic pressure equation
big pi = MRT =(Molarity)(0.0821)(Temp)
100
Vapor pressure lowering is known as ____ and to calculate:
Vapor pressure lowering = Raoult's law PA = XAP\*A PB = XBP\*B
101
Occurs when gas molecules distribute through a volume by random motion
diffusion
102
the flow of gas particles under pressure from one compartment to another through a small opening
effusion
103
equation for diffusion/effusion
r1/r2 = sqrt:(m2/m1)
104
Water solubility rules (4)
1. all salts containing alkali metal (Group 1) or ammonium (NH4+) cations 2. all salts containing the nitrate (NO3-) or acetate (CH3COO-) anions 3. All chlorides, bromides, and iodides (except Ag+, Pb2+, and Hg2+) 4. all salts of the sulfate ion (SO42-) except Ca2+, Sr2+, Ba2+, and Pb2+
105
Insolubility rules
5. all metal oxides except alkali metals and CaO, SrO, BaO 6. All hydroxides except alkali metals and Ca2+, Sr2+, and Ba2+ 7. All carbonates (CO32-), phosphates (PO43-), sulfides (S2-), and sulfites (SO32-) except alkali metals and ammonium
106
Calculation for the percent composition by mass
Mass of solute/mass of solution x 100%
107
mole fraction calculation
of moles of compound/total # of moles in system
108
molarity calculation
of mol of solute/L of solution
109
molality calculation
of mol of solute/kg of solvent
110
normality
of gram equivalent weights of solute/L of solution
111
solute vs. solvent vs. solution
Example: Salt = solute water = solvent salt water = solution
112
Arrhenius acid
arrhenius acid = species that produces excess H+ (protons) in an aq solution
113
Arrhenius base
arrhenius base = species that produces excess OH- (hydroxide ions)
114
Bronsted lowry acid
bronsted lowry acid donates protons
115
Bronsted Lowry base
bronsted lowry base accepts protons
116
lewis acid
lewis acid accepts electrons
117
lewis base
lewis base donates electrons
118
pH to H+ equation
pH = -log[H+]
119
POH to OH equation
pOH = -log[OH-] pOH = log(1/[OH-])
120
Kw equation
Kw = [H+][OH-] = 10^-14
121
pH+pOH =
14
122
reverse reaction where the salt ions react w/ water to give back the acid or base
hydrolysis rxn
123
oxidizing agent
causes another atom to undergo oxidation and it itself is reduced
124
reducing agent
causes another atom to be reduced and it itself is oxidized
125
A redox rxn in a galvanic cell has a \_\_\_∆G and is therefore \_\_\_\_
Galvanic cell = (-)∆G = spontaneous
126
A redox rxn in an electrolytic cell has a \_\_\_∆G and is therefore \_\_\_
Electrolytic cell = (+)∆G = nonspontaneous
127
EMF calculation
EMF = Ered,cathode- Ered,anode
128
Gibbs free energy in cell calculations
∆G = -nFEcell
