General Chemistry Flashcards
Calculate moles
= mass of sample / molar mass
Avogadro’s number
6.022 x 10^23 particles
Isotopes
Same number of protons (same atomic number) but different numbers of neutrons (different mass numbers)
Planck’s quantum theory
Energy emitted as electro-magnetic radiation from matter exists in discrete bundles called quanta
Energy of an electron
E = -Rh / n^2
Electromagnetic energy of photons
E = hc / λ
Balmer series
The group of hydrogen emission lines corresponding to transitions from upper levels n > 2 to n = 2
Lyman series
The group corresponding to transitions between upper levels n > 1 to n = 1
Heisenberg uncertainty principle
It is impossible to determine with perfect accuracy the momentum and the position of an electron simultaneously
Quantum numbers
1: shell (n)
- Value = n
# 2: subshell (l)
- Value = 0 to n-1
# 3: orbital (ml)
- Value = between l and -l
#4: spin (ms)
- Value = + 1/2 or - 1/2
Principle quantum number (n)
The larger the integer value of n, the higher the energy level and radius of the electron’s orbit
The maximum number of electrons in energy level n is 2n^2
Azimuthal quantum number (l)
Refers to subshells. The four subshells correspond to = 0, 1, 2, 3 and are known as s, p, d, and f
The maximum number of electrons that can exist within a subshell is given by the equation 4l + 2`
Periodic table trends
To the right: increasing effective charge, ionization energy, electronegativity, and electron affinity - decreasing atomic radius
Up the table: increasing effective charge, ionization energy, electronegativity, and electron affinity - decreasing atomic radius
Exceptions to octet rule
Stable with fewer than 8: H (2), He (2), Li (2), Be (4), B (6)
Magnetic quantum number (ml)
This specifies the particular orbital within a subshell where an electron is highly likely to be found at a given point in time
Spin quantum number (ms)
The spin of a particle is its intrinsic angular momentum and is a characteristic of the particle, like its charge
Hund’s rule
Within a given subshell, orbitals are filled such that there are a maximum number of half-filled orbitals with parallel spins
Polar covalent bond
Bonding electron pair is not shared equally, but pulled toward more electronegative atom
Polarity of molecules
Depends on the polarity of the constituent bonds and on the shape of the molecule
- A molecule with nonpolar bonds is always nonpolar
- A molecule with polar bonds may be polar or nonpolar depending on the orientation of the bond dipoles
Linear
Regions of electron density = 2
Angle between electron pairs = 180
Trigonal planar
Regions of electron density = 3
Angle between electron pairs = 120
Tetrahedral
Regions of electron density = 4
Angle between electron pairs = 109.5
Trigonal bipyrimidal
Regions of electron density = 5
Angle between electron pairs = 90, 120, 180
Octahedral
Regions of electron density = 6
Angle between electron pairs = 90, 180
Complex ion (coordination compound)
A lewis acid-base adduct with a cation bonded to at least one electron pair donor (including water)
Ligands
Donor molecules and use coordinate covalent bonds
Chelation
A process where the central cation of a complex ion can be bonded to the same ligand multiple times
Dipole-dipole interactions
Polar molecules orient themselves such that the positive region of one molecule is close to the negative region of another molecule
Dispersion forces
The bonding electrons in covalent bonds may appear to be equally shared between two atoms, but at any particular point in time they will be located randomly throughout the orbital
- This permits unequal sharing of electrons, causing transient polarization and counter polarization of the electron clouds of neighboring molecules, inducing the formation of more dipoles
Net ionic equations
These types of equations are written showing only the species that actually participate in the reaction
Neutralization reactions
Acid reacts with a base to produce a solution of a salt (and, usually, water)
Properties of the equilibrium constant
If Keq»1, an equilibrium mixture of reactants and products will contain very little of the reactants compared to the products
If Keq«1, an equilibrium mixture will contain very little of the products compared to the reactants
If Keq is close to one, it is in equilibrium
Le Châtelier’s principle
Will shift to RIGHT: (1) more of reactants (2) if product is removed (3) if pressure applied or volume reduced (4) if temp is reduced
Will shift to LEFT: (2) more products (2) reactants are removed (3) if pressure is reduced or volume is increased (4) if temp is increased
Isolated system
No exchange of energy/matter with the environment; bomb calorimetry creates a nearly isolated system
Closed system
Can exchange energy but not matter with the environment
Open system
Can exchange energy and matter
Isothermal
Temp of a system remains constant
Adiabatic
No heat exchange occurs
Isobaric
Pressure of a system remains constant
Isovolumetric (isochoric)
Volume remains constant
Endothermic
Reactions that absorb thermal energy
Exothermic
Reactions that release thermal energy
Endergonic
Reactions that are nonspontaneous
Exergonic
Reactions that are spontaneous
Enthalpy (H)
Used to express heat changes at constant pressure
Standard heat of formation (ΔHf)
The enthalpy change that would occur if one mole of a compound was formed directly from its elements in their standard states
Hess’s law
States that enthalpies of reactions are additive
- The reverse of any reaction has an enthalpy of the same magnitude as that of the forward reaction, but its sign is opposite
Bond enthalpy
Sum of bonds broken - sum of bonds formed
Entropy (S)
The measure of the distribution of energy (“randomness”) throughout a system
Gibbs free energy (G)
Combines the two factors that affect spontaneity of a reaction - changes in enthalpy, and changes in entropy
ΔG = ΔH - TΔS
Trends in G
- If ΔG is negative - spontaneous
- If ΔG is positive - nonspontaneous
- If ΔG is zero - the system is in a state of equilibrium
- ΔH = TΔS
Trends in H and S
- -ΔH and + ΔS = spontaneous at all temps
- +ΔH and -ΔS = nonspontaneous at all temps
- Both + = spontaneous only at high temps
- Both - = spontaneous only at low temps
STP vs standard conditions
STP = (0 C or 273 K, 1 atm) - generally used for gas law calculations
Standard conditions (25 C or 298 K, 1 atm, 1 M concentrations) - measuring enthalpy, entropy, Gibbs free energy, and electromotive force
Boyle’s law
PV = k or P1V1 = P2V2
Charles’s law
V / T = k or V1 / T1 = V2 / T2
Gay-Lussac’s law
P / T = k or P1 / T1 = P2 / T2
Avogadro’s principle
n / V = k or n1 / V1 = n2 / V2
Combined gas law
P1V1 / T1 = P2V2 / T2
Ideal gas law
PV = nRT
Dalton’s law of partial pressure
PT = Pa + Pb + Pc
Pa = PT(Xa)
Where Xa = na (moles of A) / nT (total moles)
Average molecular speed
K = 1/2mv^2 = 3/2KbT
Freezing point depression
ΔTf = i(Kf)m
(m) molarity
(i) van ‘t Hoff factor for ionic compounds
Boiling point elevation
ΔTb = i(Kb)m
Osmotic pressure
Π = MRT
Raoult’s law (vapor pressure lowering)
Pa = Xa(Poa)
Effusion
The flow of gas particles under pressure from one compartment to another through a small opening
Diffusion and effusion formula
r1/r2 = sq. root (m2) / (m1)
All salts containing group 1 and ammonium cations (NH4+) are
Water soluble
All salts containing nitrate (NO3-) or acetate (CH3COO-) anions
Water soluble
All chlorides, bromides, and iodides
Water-soluble with the exception of Ag+, Pb2+, and Hg2+
All salts of the sulfate ion (SO4(2-))
Water soluble with the exception of Ca2+, Sr2+, Ba2+, and Pb2+
All metal oxides are insoluble with the exception of:
Alkali metals and CaO, SrO, BaO, all of which hydrolyze to form solutions of the corresponding metal hydroxides
All hydroxides are insoluble with the exception of:
The alkali metals and Ca2+, Sr2+, and Ba2+
All carbonates, phosphates, sulfides, and sulfites are insoluble with the exception of:
Alkali metals and ammonium
Percent composition by mass
Mass of solute / mass of solution x 100
Molarity
(#) of mol of solute / liter of solution
Mole fraction
(#) of mol of compound / total # of moles in system
Molality
(#) of mol of solute / kg of solvent
Normality
(#) of gram equivalent weights of solute / liter of solution
Arrhenius
An acid is a species that produces excess H+ (protons) in an aqueous solution, and a base is a species that produces excess OH- (hydroxide ions)
Lewis
An acid is an electron pair acceptor, while a base is an electron pair donor
pH =
-log[H+] = log (1 / [H+])
pOH =
-log[OH-] = log (1 / [OH-])
Kw =
[H+][OH-] = 10^-14
pH + pOH =
14
Strong acids
HCl, HI, HBr, H2SO4, HClO3, HClO4, and HNO3
Strong bases
LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, and Ba(OH)2
Henderson-Hasselbalch equation
pH = pKa + log [conjugate base] / [weak acid]
pH = pKb + log [conjugate acid] / [weak base]
Basic to acidic compounds
NaOH, NH3, HCO3-, F-, water, H2CO3, NH4+, HSO4-, HF, HCl
Oxidizing agent
Causes another atom to undergo oxidation, and is itself reduced
Reducing agent
Causes another atom to undergo reduction, and is itself oxidized
Anode
(-)
Cathode
(+)
Galvanic cell
A redox reaction occurring here has a -ΔG
Electrolytic cell
A redox reaction occurring here has a +ΔG
Reduction potential
Each species is defined as the tendency of a species to acquire electrons and be reduced
Standard reduction potentials
emf = Ered, cathode - Ered, anode
Gibbs free energy, ΔG
ΔG = =nFEcell
