Midterm II Equations & Terms Flashcards
1
Q
stoichiometric amounts
A
when compounds are mixed in the exact amounts needed for all molecules to be “used up” completely
2
Q
limiting reagent (reactant)
A
compound that is completely used up
3
Q
% yield
A
[(actual yield)/(theoretical yield)] x 100%
4
Q
Water collective properties
A
1) many different forms of ice 2) high heat capacity 3) high boiling point 4) expands upon cooling 5) dissolves many different substances
5
Q
What kind of molecule is water?
A
polar covalent (unequal electron sharing)
6
Q
hydration
A
when a solid dissolves in water, the H (partial positive charge) orients itself toward a negative charge and vice versa with the partial negative charge
7
Q
when a salt dissolves…
A
the cations and anions separate into individual ions
8
Q
solute
A
substance dissolved
9
Q
solvent
A
water (dissolved into)
10
Q
strong electrolytes…
A
conduct electrical current strongly (ionize completely!)
11
Q
weak electrolytes…
A
conduct electrical current weakly (ionize partially!)
12
Q
non-electrolytes…
A
do not conduct electrical current (no ionization!)
13
Q
what are strong electrolytes
A
substances that ionize completely when dissolved in water (ex. soluble salts, strong acids, strong bases)
14
Q
generally soluble ions
A
Li+, Na+, K+, NH4^+, NO3^-, C2H3O2^-, Cl-, Br-, I-, SO4^2-
15
Q
Cl-, Br-, I- solubility EXCEPTIONS
A
Ag+, Hg2^2+, Pb^2+ make compound INSOLUBLE
16
Q
SO4^2- solubility EXCEPTIONS
A
Sr^2+, Ba^2+, Pb^2+, Ag+, Ca^2+ make compound INSOLUBLE
17
Q
generally insoluble compounds
A
OH-, S^2-, CO3^2-, PO4^3-
18
Q
OH- and S^2- insolubility EXCEPTIONS
A
Li+, Na+, K+, NH4^+, Ca^2+, Sr^2+, Ba^2+ make compound SOLUBLE
19
Q
CO3^2- and PO4^3- insolubility EXCEPTIONS
A
Li+, Na+, K+, NH4^+ make compound SOLUBLE
20
Q
strong acid list
A
HCl, HBr, HI, HNO3, H2SO4, HClO4
21
Q
strong base list
A
NaOH, LiOH, KOH, Ca(OH)2, Ba(OH)2
22
Q
weak acid
A
any acid that dissociates (ionizes) to only a slight extent in aqueous solution (ex. HC2H3O2, acetic acid)
23
Q
non-electrolyte
A
substance that dissolves in water but doesn’t produce any ions
24
Q
Molarity (M)
A
(moles of solute)/(liters of solution)
25
standard solution
solution whose concentration is accurately known
26
dilution equation
M1V1 = M2V2 (remember, DILUTION ONLY!)
27
molecular equation (ME)
shows reactants, products, and physical states
28
complete ionic equation (CIE)
shows everything as ions, INCLUDES "spectator" ions that do not participate in reaction
29
net ionic equation (NIE)
includes only the ions that participate in the reaction
30
when diluting, remember...
AAA (always add acid)
31
Bronsted-Lowry Definition
acid=proton (H+) donor, base=proton acceptor
32
the reaction is likely acid base if...
the products of the reaction include water and some ionic compound
33
volumetric analysis
technique for determining the amount of a certain substance by doing a titration
34
titration
carefully adding a solution of known concentration (titrant) to a solution of unknown concentration (analyte)
35
equivalence point (stoichiometric point)
when enough of the titrant has been added to react exactly with the analyte
36
indicator
a substance that's added to the analyte that changes color at or very near the equivalence point (usually a weak acid/base)
37
endpoint
the point where the indicator actually changes color
38
successful titration
1) exact reaction b/w titrant and analyte must be known 2) equivalence point must be marked accurately 3) volume of titrant required to reach equivalence point must be known
39
acid-base titration
analyte and titrant are acid/base
40
gas-evolution reaction
when 2 aqueous solutions mix and produce a new gaseous substance that bubbles out of the solution
41
oxidation-reduction reactions (redox)
reactions in which one or more electrons are transferred
42
assigning oxidation state/# rules
1) oxidation state of pure element is zero 2) oxidation state of any monoatomic ion is equal to the charge on the ion 3) sum of oxidation states on a polyatomic ion equal the charge on the ion (0, -1, +3, etc) 4) fluorine in a molecule is always -1 5) hydrogen in a molecule is usually +1 6) oxygen is usually -2
43
oxidation
an increase in oxidation state (LOSS of electrons)
44
reduction
a decrease in oxidation state (GAIN of electrons)
45
Evangelista Torricelli
invented the mercury barometer (at sea level, barometer is at 760mmHg)
46
manometer
measures pressure of sample gas, heigh difference in sides indicates pressure of gas relative to atmospheric pressure
47
pressure conversions
1 atm = 760mmHg = 760 torr = 101,325 Pa = 29.92 inHg = 14.7 psi
48
Boyle's law (Robert Boyle)
PV=k; pressure and volume are inversely proportional, P1V1=P2V2
49
Charles's law (Jacques Charles)
V=bT; volume of gas is directly proportional to the temperature, T is in Kelvin, (V1/T1)=(V2/T2)
50
Avogadro's law (Lorenzo Avogadro)
V=an; for a given gas at a constant temperature and pressure the volume is directly proportional to the number of moles of gas, (V1/n1)=(V2/n2)
51
Ideal gas law
PV=nRT, R = 0.08206 L*atm/K*mol
52
molar mass of a gas
dRT/P, d=density
53
ideal conditions for a gas
1 atm, 273.15K, 22.4L
54
Dalton's law of partial pressures
P(total)= P1+P2+P3+...+Pn
55
assumptions of ideal gas law (simplifications)
1) volume of the individual gas particles must not be important 2) forces among the particles must not be important
56
Mole fraction (X)
the ratio of the number of moles of a given component in a mixture to the total number of moles in the mixture
57
mole fraction equations
X1 = n1/n(total), X1 = P1/P(total)
58
pressure/mole fraction equation
P1=X1(P(total)), n1 = X1(n(total))
59
vapor pressure of water
pressure due to water vapor over liquid water; if calculating pressure of sample over water, make sure to subtract this from total pressure
60
Kinetic Molecular Theory (KMT)
1) particles are so small compared with the distances b/w the, that the volume of individual particles can be assumed to be negligible (0) 2) 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 3) particles are assumed to exert no forces on each other 4) average kinetic energy of a collection of gas particles is assumed to be directly proportional the the temperature in K of the gas
61
Pressure and volume relation (Boyle)
if volume DECREASES then pressure INCREASES (inversely proportional) (all else constant)
62
Pressure and temperature relation
if temperature INCREASES then pressure INCREASES (directly proportional) (all else constant)
63
Volume and temperature relation (Charles)
if volume DECREASES then temperature DECREASES (all else constant)
64
Volume and number of moles relation (Avogadro)
if number of moles INCREASES then the volume must INCREASE (all else constant)
65
Kinetic energy equation
KE = (1/2)(mv^2)
66
AVERAGE kinetic energy equation
KEavg=(2/3)(RT), if temperature increases so does energy (direct proportion)
67
the lighter a gas is...
the larger its range of velocity, and the higher its peak velocity is, see how H2 has HIGH velocity
68
as temperature increases, molecular velocity...
has a wider range/higher average
69
diffusion
describes the mixing of gases; rate of diffusion is rate of mixture
70
effusion
describes the passage of a gas through a tiny orifice into an evacuated chamber; rate of effusion is the rate of transfer into the chamber
71
Grahams law of effusion (Thomas Graham)
the rate of effusion for a gas is inversely proportional to the square root of the mass of the gas
72
effusion rate equation
(rate of effusion for gas 1= M2^1/2) / (rate of effusion for gas 2 = M1^1/2), [(mm2)^1/2]/[(mm1)^1/2]
73
van der Waals equation (Johannes Diderik van der Waals)
[P+(an^2/V^2)](V−nb)=nRT
74
correction for intermolecular forces
a(n/V)^2
75
correction for particle volume
nb (used as V-nb)
76
energy
the capacity to do work or produce heat
77
law of conservation of energy
energy can be converted from one form to another but neither created nor destroyed
78
potential energy
energy due to composition or position
79
kinetic energy
energy due to the motion of the object
80
heat
the transfer of energy between two objects due to a temperature difference
81
temperature
a property reflecting the motions of particles in a substance
82
work
a force acting over a distance (W = F x d), (W = -PΔV)
83
system
the part of the universe on which we wish to focus attention
84
surroundings
everything else in the universe (other than the system)
85
exothermic
when a reaction gives off heat (E flows FROM the system)
86
endothermic
when a reaction absorbs heat (E flows INTO the system)
87
change in potential energy (ΔPE)
the energy required to break the bonds in the reactants minus the energy released when the bonds in the products are formed
88
energy of a system equation
ΔE = q + w, q=heat and w=work
89
enthalpy
H = E + PV
90
calorimeter
the device used experimentally to determine the heat associated with a chemical reaction
91
calorimetry
the science of measuring heat
92
specific heat capacity
heat capacity of a substance in grams (J/°C*g or J/K*g)
93
molar heat capacity
heat capacity of a subtance in moles (J/°C*mol or J/K*mol)
94
wavelength (λ)
the distance b/w two consecutive peaks or troughs in a wave, unit is m or nm
95
frequency (ν)
the number of waves (cycles) per second that pass a given point in space, unit is Hertz or 1/s
96
speed of light equation
c = (ν)(λ)
97
speed of light
c = 2.9979 x 10^8 m/s
98
electromagnetic spectrum (in DECREASING wavelength)
radio, microwave, infrafred, visible, ultraviolet, X-ray, Gamma ray
99
red light wavelength
625nm-750nm
100
orange light wavelength
590nm-625nm
101
yellow light wavelength
565nm-590nm
102
green light wavelength
500nm-565nm
103
blue light wavelength
450nm-500nm (lighter blue has longer wavelength)
104
violet light wavelength
380nm-450nm
105
Planck's constant
h=6.626*10^-34 J*s
106
Planck's equation
ΔE=hν (planck's constant times frequency)
107
quantum
quantized packets, only was energy can be absorbed or emitted
108
quantum equation/energy of photon equation (Einstein)
E(photon)= hc/λ, also used for ΔE for an electron
109
de Broglie equation
λ = h/mv (v is velocity)
110
1 J equals...
(1 kg*m^2)/(s^2)
111
emission spectrum
range of electromagnetic radiation of various wavelengths
112
quantum model (Niels Henrik David Bohr)
electron in a hydrogen atom moves around the nucleus only in certain allowed orbits
113
Bohr constant
E = -2.178 x 10^-18 J
114
Bohr electron transition equation
ΔE = -2.178 x 10^-18 J [(1/nf^2) - (1/ni^2)]
115
problems with Bohr model
1) only works with things that work as a circle (2D things that can be approximated as a circular thing) 2) only works on ONE electron systems
116
Schrodinger equation
Hψ=Eψ; H is the Hamiltonian operator/multiplier, ψ is the wavefunction (cartesian coordinate representation of the orbital), E is the summed energy of the atom
117
orbital
probability of finding an electron in a given area of space around the nucleus in all 3 dimensions
118
Heisenberg uncertainty principle
there is a fundamental limitation to just how precisely we can know both the position and the momentum of a particle at any given time
119
Heisenberg equation
Δx*Δ(mv) = h/4pi, or Δx*Δp = h/4pi; Δx is changing position of electron and Δmv is the changing momentum; the MORE we know about position, the LESS we know about velocity and vice versa
120
ψ^2
probability density/distribution
121
orbital size defintion
size of an orbital is the radius of the sphere that encloses 90% of the total electron probability
122
principal quantum number (n)
has integer values; related to the SIZE and ENERGY of the orbital. 1s,2p,3d,4f, etc
123
angular momentum quantum number (l)
has integer values from 0 to n-1 for each n value; is related to the SHAPE of the atomic orbital, sometimes called subshells (ex. l=0, s orbital)
124
magnetic quantum number (mₗ)
has integer values between -l and l INCLUDING zero; tells us about the relative orientation in space of the orbital
125
spin quantum number (mₛ)
either -1/2 or 1/2; indicates orientation of electron spin
126
node
area of low electron probability (n-1 gives number of nodes)
127
Pauli exclusion principle
in a given atom, no two electrons can have the same set of four quantum numbers; two electrons in same orbital must have OPPOSITE spins
128
a given orbital...
can only hold 2 electrons and the occupied orbital must contain electrons of opposite spin
129
shielding
the outermost electrons have other electrons between themselves and the nucleus leading to a net smaller attraction to the nucleus for the outermost electrons
130
electron approximation
we approximate by saying an electron is a field of charge that is the net result of the nuclear attraction and the average repulsions of all other electrons
131
types of energy in complex atoms
1) the kinetic energy of electrons moving around the nucleus 2) the potential energy of nuclear/electron attractions 3) the potential energy of electron/electron repulsions
