Section 1
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+1, except if bonded to Alkali Metal -1
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Section 1
(50 cards)
+1, except if bonded to Alkali Metal -1
Oxidation # of Hydrogen
Bronsted-Lowry Acid
H + donor
different # of neutrons
Isotope
H2, F2, N2, O2, Cl2, Br2, I2
Diatomic Molecules
1)Assign Oxidation #s 2) Half Reaction 3) Balance Moles 2) Balance add e- to balance Oxidation #s (RED-OX = reduction on left, Oxidation on right) 4) Add H+ or OH - to balance charge 5) Add H2O to balance hydrogens 4) Balance electrons by multiplying until half reactions are equal 5) Add together & check
How to Balance a Redox Equation
End Point
In a titration, the point where the indicator changes (just after moles of solid are equal to moles of base)
Acid + Base --> Salt + Water
Neutralization Reaction (general format) (net ionic of which is always (H+) + (OH-) --> (H2O))
electrolyte
Substance that, when dissolved, is conductive
-1
Oxidation # of Halogens
0
Oxidation # of Compounds
actual yield/theoretical yield x 100%
% yield
p+
proton (symbol)
Allotrope
An element with several different forms, each with different properties (i.e. graphite & diamond)
Atomic Mass Unit
The weighted average of all the isotopes that an atom can have (in g/mol)
M1V1=M2V2
Formula used when diluting stock solution (to find amount of water or stock needed)
Arrhenius Base
Ionizes to produce OH- Ions
Solute
Substance being dissolved in a solution (lower [ ])
Monoprotic
Donates a single H+ Ion (... other prefixes also)
1) Convert to moles 2) divide by lowest moles 3) if any halfs, double all values 4) plug into Compound
How to Find an Empirical Formula Given Grams
Ionic Compounds
Compound in which there is a bond between two non-metals; when naming them you use the numerical prefixes (may NOT be reduced i.e. S2F4 may not become SF2)
Volume Metric Flask & Pipet
Tools NEEDED for dilution
charge
Oxidation # of Ions
Solvent
Substance in which something is dissolved in a solution (higher [ ])
Molecule
2 or more covalently bonded atoms
charge
Oxidation # of Polyatomic Ions
Arrhenius Acid
Ionizes to produce H+ ions
1) multiply each AMU by the percentage that represents that isotopes occurrence in nature 2) then add all the AMUs together.
How to Find a Weighted Average
Cation
positive ion
period
Horizontals on the periodic table
# protons (atom is defined by this)
Atomic #
1) Assume percentages are g (where you have 100 g) 2) convert to moles 3) divide by lowest value 4) plug into compound
How to Find an Empirical Formula Given Percentages
Metalliods
Elements on staircase on periodic table
-2, with peroxide -1
Oxidation # of Oxygen
0
Oxidation # of free elements
Bronsted-Lowry Base
H+ Acceptor
n0
neutron (symbol)
Covalently w/in themselves, Ionicly bonded w/ each other
Polyatomic Ions (bonding)
Molecular Compounds
Compound in which there is a bond between a metal and a non-metal; when naming, pay attention to ionic charges on periodic table (may be reduced)
Oxidation is Loss Reduction is Gain
OIL RIG
Molar Mass of Element/ Total Molar Mass
% Composition Formula
Solution
A homogeneous mixture with 1 phase
Ionic
All _________ compounds are electrolytes
moles of solute/ L of solution
Molarity (M)
e-
electron (symbol)
Amphoteric
a single substance that may be an acid or a base (i.e. water)
Molecular
All (though really not all, but for our purposes all) ________ compounds are NOT electrolytes
Anion
negative ion
# protons + # neutrons
Mass #
group
Verticals on the periodic table
precipitate
A solid or gas that can be formed when 2 or more aqueous reactants come together
Section 2
(50 cards)
M = square root (3RT/mm)
Kinetic Energy of an individual particle formula
Closed System
Type of system in which the energy may escape, but the mass is conserved
Thermochemistry
The heat changed in a chemical reaction.
Equivalence Point
In a titration, the point where moles of acid are equal to moles of base (just before the indicator changes color)
0 degrees C, 1 atm
STP
Open System
Type of system in which the energy and mass may leave or enter
Specific Heat (s)
The energy required to raise 1 g of substance 1 degree C
P of a =(X of a)(total pressure)
Dalton's Law of Partial Pressures (to find partial pressure formula)
Isolated System
Type of system in which nothing is transfered (no mass or energy); ideal
Average KE = 1/2(mass)(average speed of all particles)
Average Kinetic Energy Formula
X of a = moles a/total moles
Mole Fraction
0
If heat capacity isn't mentioned, you can assume that q of cal is = ?
Surroundings
Everything in the universe that is not defined by you as part of the system
Force = mass x acceleration
Newton's Second Law
rate gas A/rate gas B = square root (mm A/ mm B)
Speed of Diffusion/Effusion formula
Pressure of H2O must be Subtracted
In ideal gas law problem, when it says "atmospheric" ...
Calorimetry
The measurement of heat changes
Temperature
Measure of the average kinetic energy of all the particles in a substance
PV=nRT
Ideal Gas Law Formula
H
Symbol for Enthalpy
PEACe 1)each molecule is a Point in space 2)Elastic collisions 3) no force of Attraction 4) no energy lost or gained to Collisions
Idea Gas Law (actual rules)
rate gas A/rate gas B = square root (mm gas B/ mm gas A)
Rate of Diffusion/Effusion formula
C=(mass)(specific heat)
Heat capacity formula
Reducing Agent
The reactant in the oxidizing reaction that forces the reduction reaction to occur
q rxn = -(q of H2O + q of cal)
q rxn = ?
q
Symbol for Total Heat absorbed or released
1 atm = 760 mmHg = 101.3 kPa
Pressure Units/Conversions
Heat Capacity (C)
The amount of energy/heat required to raise some substance 1 degree C
Oxidizing Agent
The reactant in the reduction reaction that forces the oxidation reaction to occur
Delta H or Enthalpy Change
Measure of the change in enthalpy
Temperature
Kinetic Energy is proportional to ______
0.0821 atm L/mol K
R in ideal gas law
Barometer
Instrument used to measure the pressure of atmospheric gas
purple --> pink
Colors of Reaction when (MnO4 -) --> (Mn2+)
force x distance = work done
Energy (definition)
yellow --> green
Colors of Reaction when (Cr2O7 2-) --> (Cr 3+)
Atmospheric Pressure
the weight exerted by a column of air or the pressure exerted by the Earth's atmosphere
System
What is defined by you taken from the whole universe
Exothermic
Releases/gives off heat (negative value)
8.314 J/K mol
R in instances that pertain to energy
Constant Volume
Bomb Calorimeter
q cal = CAT
q cal = ?
AH
Symbol for the heat absorbed or lost molecularly (PER MOLE)
P1V1/N1T1=P2V2/N2T2
Combined Gas Law Formula
Endothermic
Absorbs/takes in heat (positive value)
Manometer
Instrument used to measure the pressure of a not-atmospheric gas (open or closed system)
Calorimeter
A device used to measure Delta H
Constant Pressure
Coffee Cup Calorimeter
q H2O = msAT
q H2O = ?
high pressure, low temperature
Non-Ideal Gas Conditions
Section 3
(50 cards)
Positive work value; work done on system
When gas compresses ...
∆H | ∆S is spontaneous...at high temperatures
+ | +
isothermal
change that occurs at constant temperature
London dispersion forces
universal IMF for nonpolar molecules
square planar
AX4E2
octahedral
AX6
Negative work value; work done by system
When gas expands ...
State Functions
Function in which it doesn't matter HOW the changes take place, only that they do; you can use to formula Final - Initial ( i.e. energy, pressure, volume, temperature, sort of moles)
AE= AH - RTAn
Change in Energy = ? (in terms of constant pressure; for gasses)
AE = q + w
Change in Energy (AE) = ? (in terms of work)
London dispersion forces
larger molecules which have higher mass and therefore electron density have stronger...
+ ∆S
gas > liquid > solid and (aq)>(s) cause...
Hydrogen bonding
IMF that occurs with FON
T-shape
AX3E2
sublimation
phase change from solid to gas
methods of increasing rate
raising heat, adding catalyst, heighten concentration, bigger surface area
N=N.(0.5)^time/time half-life
half-life equation
endless
Significant Digits of Conversion Factors
w =-(P)(Change in V)
work = ?
∆H | ∆S is spontaneous...never
+ | -
+ ∆S
greater # of moles of gas formed and greater volume formed cause...
+ ∆S
diatomic molecules forming mixed molecules, e.g. H2 + I2 -> 2HI, means...
triple point
temperature-pressure combination at which solid, liquid, and gas states appear
critical point
temperature-pressure point after which gas can no longer form liquid
see-saw
AX4E
fusion
melting
effects of IMF
boiling point, melting point, viscosity, vapor pressure, surface tension
solid CO2
chemical composition of dry ice
are not
For significant digits, leading zeros ____ significant
are
For significant digits, trailing zeros _____ significant
∆H | ∆S is spontaneous...at low temperatures
- | -
Adding
When ____ significant digits, round answer to least decimal place
AH rxn = (Sum of AH of formation of products) - (Sum of AH of formation of reactants)
Dirrect Method Formula
viscosity
thickness
boiling point
point at which vapor pressure=air pressure above
Multiplying
When _____ significant digits, round answer to least significant digit
vapor pressure
stronger IMF= lower... weaker IMF= higher...
Dipole-dipole forces
IMF that exists in polar molecules
AH = q/moles
Delta H (AH) = ?
trigonal bipyramidal
AX5
∆Hvap
energy needed to vaporize a mole of a liquid
square pyramidal
AX5E
An = (moles of products that are gasses) - (moles of reactants that are gasses)
Change in moles (An) =?
equilibrium
happens at lines in phase change charts
∆H | ∆S is spontaneous...always
- | +
deposition
phase change from gas to solid
endless
Significant Digits of counted things
adiabatic
change without heat transfer between the system and its surroundings
0
AH of formation for a substance in its stablest form (how it is found in nature)
bond energy
energy needed to break a bond
Section 4
(50 cards)
prop-
(organics) three carbons
non-
(organics) nine carbons
blue-green
when n=4 ->2, color=
n (first quantum number)
variable for energy of e-, goes from 1,2,3 on up
d
l=2
dec-
(organics) ten carbons
m (third quantum number)
variable for orientation of orbital (-1 through +1)
Balmer Series
spectrum of light when an electron drops to energy level n=2
Effusion
passage of gas through tiny orifice
ionic
... compounds are most conductive
specific heat
heat needed to change 1 g of substance to 1˚C
first
ln[A] vs. time is a ...-order reaction
p
l=1
hex-
(organics) six carbons
entropy (S)
degree of disorder in a system
-ous acid
if anion ends in -ite, acid name ends in...
electron affinity
energy involved in gaining an electron to become a negative ion
blue-violet
when n=5 ->2, color=
but-
(organics) four carbons
third
AP doesn't deal with ...-order reaction, don't pick it!
d orbitals
these orbitals are diagonal
oct-
(organics) eight carbons
eth-
(organics) two carbons
s orbitals
these orbitals are spherical
s
l=0
l (second quantum number)
variable for type of orbital
Boltzmann distribution
states molecules at a given temp. vary in kinetic energy along a bell-curve of molecular velocities
Diffusion
mixing of gases
zero
[A] vs. time is a ...-order reaction
hydro-ic acid
if anion ends in -ide, acid name ends in
red
when n=3 ->2, color=
g solute/g solvent x 100
mass percent
heat capacity
amount of heat needed to change a system by 1˚C
C + 273
calculation from K to C
moles solute/kg solvent
molality =
hept-
(organics) seven carbons
rate law
this MUST be determined experimentally
pent-
(organics) five carbons
not spontaneous
If K<1, then Gº>0 and reaction will be...at chemical equilibrium
f
l=3
increasing
entropy in the universe is always...
-ic acid
if anion ends in -ate, acid name ends in...
second
1/[A] vs. time is a ...-order reaction
violet
when n=6 ->2, color=
4.184
specific heat of water
spontaneous
If K>1, then Gº<0 and reaction will be...at chemical equilibrium
like
like dissolves...
meth-
(organics) one carbon
p orbitals
these orbitals are perpendicular
s (fourth quantum number)
variable for spin of electron (+.5 or -.5)
Section 5
(50 cards)
red/orange
color of Ca (flame test)
alkyne
(organics) triple-bonded compound
condensation
organic reaction in which two functional groups come together, resulting in the release of water
red
color of Sr (flame test)
Density
mass/volume
methoxy-
ether prefix
base
metal oxide + H20 ->
Alkali metals
Group 1 metals
ln (k1/k2) = (Ea/R)(1/T2-1/T1)
equation to find Ea from reaction rate constants at two different temperatures
phosphate, sulfide, carbonate, sulfate
generally insoluble anions (names)
green/yellow
color of Ba (flame test)
-ol
ending for alcohols
geometric isomers
two molecules with identical connectivity but different geometries
Transition metals
elements in groups 3-12
-oic acid
carboxylic acid ending
yellow
color of Na (flame test)
-al
aldehyde suffix
hydrolysis
organic reaction in which water breaks apart a molecule (splitting into two hydroxides)
Law of Multiple Proportions
atoms combine in fixed whole # ratios
Law of Conservation of Mass
mass reactants= mass products
activated complex (transition state)
peak of energy diagram
alcohol
organic w/ -OH group
mol
6.022x10^23
acid
nonmetal oxide + H2O ->
complex ions
transition metals with ammonia, hydroxide, cyanide or thiocyanate form...
diamagnetic
elements which have all electrons paired and relatively unaffected by magnetic fields
-oate
ester suffix
amine
organic w/ -NH2
blue
color of Cs (flame test)
oxide gas and water
oxoacid solution (such as HSO4-) forms...
purple
color of K (flame test)
allotrope
different form of same element
Ag+, Pb2+, Hg2+
all cations are soluble with bromide, chloride and iodide EXCEPT
insoluble
hydroxides are soluble or insoluble?
red
color of Li (flame test)
Amino-
amine prefix
Alkaline earth metals
Group 2 metals
-one
ketone suffix
Arrhenius equation
to find activation energy use the...
carbohydrates
carbon, hydrogen, oxygen compounds
0
∆Hº of pure elements=
Ag+, Pb2+, Hg2+, Sr2+, Ca2+, Ba2+
all cations are soluble with sulfate EXCEPT
k=Ae^(-Ea/RT)
Arrhenius equation
base and hydrogen gas
pure metal or metal hydride + H20 ->
alkene
(organics) double-bonded compound
Limiting reactant
reactant that's completely used up in a chemical reaction
ether
organic w/ -O-
hydrocarbons
carbon & hydrogen compounds
paramagnetic
elements which have unpaired electrons and highly affected by magnetic fields
alkane
(organics) single-bonded compound
Section 6
(50 cards)
Heisenberg Uncertainty Principle
we cannot simultaneously determine an atom's exact path or location
% error
theoretical yield-experimental yield/theoretical yieldx100
nu
frequency symbol
equivalence point
point where acid completely neutralizes base
standard solution
a solution used in titrations whose concentration is known
end point
point at which the titrated solution changes color
Molality
mol/kg of solvent, used in calculating colligative properties
22.4L
volume of gas @STP
Graham's Law
effusion of a gas is inversely proportional to the square root of the molar mass
soluble
Group 1, Ammonium, Nitrates, Acetates, Sulfates, Halides
Strong acid weak base rxn
H⁺+NH₃→NH₄
excess reactant
reactant which doesn't get used up completely in a chemical reaction
oxidation
loss of electrons, increase in oxidation #
mol Fraction
mols A/ total mols, XA
1/2mv²
Kinetic Energy per molecule
1 atm
760 mmHg, 760 torr
√3RT/M(in kg)
average speed of gas
oxidizing agent
the reactant that is being reduced, brings about oxidation
linear
AX₂, AX₂E₃
reduction agent
the reactant that is being oxidized, brings about reduction
tetrahedral
AX₄
Hund's Rule
within a sublevel, place one e⁻ per orbital before pairing them
analyte
solution in flask being titrated
titrant
buret solution used in titration
wavelength
determined by the formula h/m(in kg)v, (v=velocity)
reduction
gain of electrons, decrease in oxidation #
strong bases
Group 1 and heavier Group 2 bases
1.38x10⁻²³J/K
Boltzmann constant, used in calculating speed of gas per molecule
weak acid strong base rxn
HF+ OH⁻→H₂O
Pauli Exclusion Principle
each orbital can hold two e⁻s each w/ opposite spins
Molarity
mol/L, concentration of a solution
Pressure
force per unit area
3.0x10⁸m/s
speed of light, C
experimental yield
the actual amount of product produced in an experiment
Aufbau Principle
e⁻s fill the lowest energy orbital first, then work their way up
lambda
wavelength symbol
insoluble
Carbonates, Hydroxides, Oxides, Phosphates, Sulfides
3/2RT
Kinetic Energy per mol
√3kT/m
speed per molecule of gas
trigonal planar
AX₃
indicator
a weak acid that changes color at or near the equivalence point
% yield
experimental yield/theoretical yieldx100
strong acids
HCl, HBr, HI, HNO₃, HClO₄, H₂SO₄
E
#NAME?
Standard Temperature and Pressure
0.00°C, 1 atm
strong acid strong base rxn
H⁺+OH⁻→H₂O
r₁/r₂
=√M₂/M₁
Dalton's Law
Ptotal=Pa+Pb+Pc....
Theoretical yield
amount of product produced when limiting reactant is used up
6.63x10⁻³⁴Js
Planck's constant, used to calculate energy w/frequency
Section 7
(50 cards)
seesaw
AX₄E
square planar
AX₄E₂
permanent gases
substances w/ critical temperatures below 25°C
vaporization
liquid to gas
endothermic
positive enthalpy, heat flows into system
supercritical fluid
substances above the critical temperature and pressure in which the pressure is so high that density and flowing ability of a "gas" resembles that of a liquid
1.86°C
kf of water
octahedral
AX₆
Hydrogen bonding
unusually strong dipole forces found when H is bonded to N, O, or F
Q>K
reverse rxn occurs when
bent
AX₂E, AX₂E₂
exothermic
negative enthalpy, heat flows into surroundings
LeChatelier's Principle
If a system @equilibrium is stressed, the system will shift so as to reestablish equilibrium
viscosity
resistance to flow
adhesion
molecules' tendency to stick to the container
ΔTf= kf x molality
freezing point depression formula
Cg=kPg
Henry's Law, solubility of gases is directly proportional to the partial pressure of the gas
freezing
liquid to solid
sublimation
solid to gas
system
the part of the universe one is focused upon (in thermodynamics)
melting point
point at which solid→liquid occurs
Q<K
forward rxn occurs when
rate
how fast or slow a reaction occurs, becomes slower as it reaches equilibrium
Van't Hoff factor
for colligative properties for electrolytes, the # of mols of ions/ mols of solute
single bond
1 sigma bond
Bond enthalpy
ΔH when 1 mol of bonds is broken in the gaseous state
heat of vaporization
energy required for liquid→gas
pi=(nRT)/v
osmotic pressure formula
condensation
gas to liquid
ΔTb= kb x molality
boiling point elevation formula
cohesion
molecules' tendency to stick to one another
heat capacity
heat required to raise the system 1°C
boiling point
point at which liquid→gas occurs
P₁= X₁P₁°
Raoult's Law, relations between vapor pressure and concentrations
triple bond
1 sigma bond, 2 pi bonds
0.512°C
kb of water
London dispersion forces
weakest IMFs, found in all molecules
catalyst
lowers activation energy
heat of fusion
energy required for melting to occur
work
all forms of energy except for heat
double bond
1 sigma bond, 1 pi bond
deposition
gas to solid
trigonal bipyramidal
AX₅
t-shape
AX₃E₂
square pyramidal
AX₅E
surroundings
the rest of the universe (in thermodynamics)
trigonal pyramidal
AX₃E
melting
solid to liquid
activation energy
the minimum energy that molecules must possess for collisions to be effective, Ea
Colligative properties
vapor pressure lowering, boiling point elevation, freezing point depression, osmotic pressure
Section 8
(50 cards)
permanganate
MnO₄¹⁻
Weight
amount of gravitational force exerted on an object
anode
half cell in which oxidation occurs
Buffer
a solution that resists a change in pH, contains both a weak acid and its conjugate base
chlorite
ClO₂¹⁻
no precipitate forms
if Q<Ksp
conjugate acid
the chemical formed when a base accepts a proton
sulfite
SO₃²⁻
phosphate
PO₄³⁻
nitrate
NO₃¹⁻
ammonium
NH₄¹⁺
5% rule
x can be ignored when % ionization is <5%
oxalate
C₂O₄²⁻
chloride
Cl¹⁻
Mass
a measure of resistance of an object to a change in its state of motion
96,500
Faraday's constant
Scientific Method
a method of investigation involving observation and theory to test scientific hypotheses
chromate
CrO₄²⁻
bromate
BrO₃¹⁻
Bronsted-Lowry base
proton acceptors, must have an unshared pair of e⁻s
nitrite
NO₂¹⁻
Bronsted-Lowry acid
proton donors
Matter
anything occupying space and with mass
2nd law of thermodynamics
the total entropy is always increasing, all systems tend towards maximum entropy
entropy
a measure of randomness or disorder
chlorate
ClO₃²⁻
oxide
O²⁻
iodide
I¹⁻
carbonate
CO₃²⁻
acetate
C₂H₃O₂¹⁻
Law of Conservation of Mass
matter can't be created nor destroyed
salt bridge
connects the 2 half cells in a voltaic cell
voltaic cells
uses a spontaneous redox rxn to generate electrical energy, consists of 2 half cells
sulfide
S²⁻
1st law of thermodynamics
the total energy of the universe is constant, all systems tend towards minimum energy
1.0x10⁻¹⁴
Kw
spontaneity
the likelihood that a rxn will occur "by itself"
Arrhenius base
puts OH⁻ into solution
cathode
half cell in which reduction occurs
a precipitate forms
if Q>Ksp
perchlorate
ClO₄¹⁻
Arrhenius acid
puts H⁺ into solution
flouride
F¹⁻
hydroxide
OH¹⁻
conjugate base
the chemical formed when an acid donates a proton
3rd law of thermodynamics
the entropy of a pure perfectly formed crystal @0K is 0
sulfate
SO₄²⁻
Nernst Equation
Ecell= E°cell -RT/nF x lnQ
cyanide
CN¹⁻
dichromate
Cr₂O₇²⁻
Section 9
(50 cards)
Finding Empirical Formulas
1.Calculate moles of each atom in molecule 2.Divide each mole number by smallest mole number 3.If necessary, multiply every mole number to get a whole number 4.Moles of each atom is subscript in empirical formula
Kinetic Molecular Theory-FOR IDEAL GASES!!!
1.Volume of individual particles can be assumed to be zero 2.The particles are in constant motion, which causes pressure 3.Particles exert no forces on each other 4.The average kinetic energy of the particles is directly affected by temperature(K)
Naming Binary Ionic Compounds
Cation first, anion second
Dalton's Law of Partial Pressures
Ptotal=P1+P2+P3+...
Net Ionic Equation
only contains ions that change in reaction
Second-Order Half Life
1/([A]0*k)
First-Order Rate Law
r=k[A]
Law of Definite Proportion
a given compound always has exactly the same proportion of elements by mass
Integrated Zero-Order Rate Law
[A]=-kt + [A]0
First-Order Half Life
.69/k
Gamma Ray-
high-energy light
Root Mean Square Velocity
u(rms)=(3RT/M)^1/2
Acids
substances that form H+ when dissolved in water; proton donors
R=
0.0826Latm/Kmol
Zero-Order Half Life
[A]0/2k
Zero-Order Rate Law
r=k
1atm=?mmHg/Torr
760mmHg/Torr
Chemical Bonds
Force that holds atoms together
If needed, indicate charge of metal(cation) by
a Roman numeral
Osmotic Pressure
Osmotic pressure=MRT
Decrease Volume and Increase Temperature
Increase Pressure
Increase Temperature
Increase Volume
Chemical Kinetics
studies the rate at which a chemical process occurs and sheds light on its reaction mechanism
Molal FP Depression Constant
▲T=k*m(solute)
Radioactivity
Spontaneous emission of radiation
STP
0°C and 1 atm
Integrated First-Order Rate Law
ln[A]=-kt + ln[A]0
Second-Order Rate Law
r=k[A]^2
Molal BP Elevation Constant
▲T=k*m(solute)
Overall Reaction Order
n+m (these are orders of reactants)
Law of Multiple Proportions
when two elements form a series of compounds, the ratios of the masses of the second element that combine with 1g of the first element can always be reduced to whole numbers
van't Hoff Factor
i=moles of particles/moles of solute dissolved
Integrated Rate Law
expresses how the concentrations depend on time
In covalent bonds, prefixes are used to tell
amount of atoms present
Molarity
moles of solute/volume of soln(L)
Normality
(N) number of equivalents per liter of solution
A monoatomic cation takes name from
its element
Beta Particles-
high-speed electrons
Isotopes
atoms with the same number of protons but a different number of neutrons
1atm=?Pa
101,325 Pa
Bases
Substances that form OH- when dissolved in water; proton acceptors
Ideal Gas Law
PV=nRT
R=
8.31J/Kmol
Integrated Second-Order Rate Law
1/[A]=kt + 1/[A]0
Joule
SI unit of energy; Kg*m^2/s^2
A monoatomic anion is named by taking
its root and adding -ide
Percent Yield
Actual Yield/Theoretical Yield*100%
Molality
mol of solute/kg of solvent
Alpha Particles-
2+ charge
Enthalpy of Solution
▲Hsoln=▲H1+▲H2+▲H3+...
Section 10
(50 cards)
Volt
unit of electrical potential; J/C
Q<K
shift to right (Q?K)
w(max)
=▲G (work)
▲G
▲H-T▲S
pH=
=log[H+]
▲Ssurr
-▲H/T
Equilibrium Expression
K=[C]^l[D]^m/[A]^j[B]^k; products/reactants; solids don't count
Equilibrium constant
K
Speed of light
c=2.9979*10^8 m/s
▲Suniv
▲Ssys+▲Ssurr
Entropy
(S) the driving force for a spontaneous is an increase in entropy of the universe
Cell Potential
(Ecell) driving force of the electrons
▲H°reaction
Σn▲H°(products)-Σn▲H°(reactants)
Reaction Quotient
(Q) does the same as equilibrium expression, except it uses initial concentrations
Principal Quantum Number
(n) has values 1,2,3,...; tells energy levels
Q=K
at equilibrium (Q?K)
Theory of Relativity
E=mc^2
Le Chatelier's Principle
if a change is imposed on a system at equilibrium, the position of the equilibrium will shift in a direction that tends to reduce that change
▲G°
▲H°-T▲S°
Law of Conservation of Energy
energy can't be created nor destroyed
Angular Momentum Quantum Number
(ℓ), has values from 0 to (n-1); tells shape of atomic orbitals
Specific Heat Capacity
J/°Cg or J/Kg
Aufbau Principle
as protons are added to the nucleus, electrons are similarly added
Acid Dissociation Constant
Ka=[products]^m/[reactants]^n
▲E
q+w
Buffered Solution
a solution that resists a change in its pH
Quantum Model
electrons in a hydrogen atom move around the nucleus only in circular orbits
ℓ=0
s orbital
ℓ=4
g orbital
Nodes
where there are no electrons
Pauli Exclusion Principle
in a given atom no two electrons can have the same set of four quantum numbers
Electron Spin Quantum Number
(msubs) can only be +1/2 or -1/2
Quantum Mechanical Model
involves quantum numbers
Magnetic Quantum Number
(mℓ) has values from -ℓ to ℓ, including zero; tells orientation of the orbital relative to other orbitals
c=
λv
Work
force acting over distance
Solubility Product
(Ksp) an equilibrium expression
ℓ=2
d orbital
Galvanic Cell
a device in which chemical energy is changed to electrical energy
G
G°+RTln(Q)
Heat
the transfer of energy between two objects due to temperature difference
Cathode
where reduction occurs
Hund's Rule
the lowest energy configuration for an atom is the one having the max number of unpaired electrons allowed by the Pauli principle in a set of degenerate orbitals
ℓ=3
f orbital
Molar Heat Capacity
J/°Cmol or J/Kmol
Quantum Numbers
describe various properties of one orbital
ℓ=1
p orbital
Anode
where oxidation occurs
Hess's Law
in going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps
Q>K
shift to left (Q?K)
Section 11
(28 cards)