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Lab equipment practice


Laboratory spatula: Scoops stuff, moves it, applies it



Erlenmeyer flask: Beakers with thin necks



Test tube rack: Holds test tubes



Graduated cylinder: Holds, measures liquids



Watch glass: Holds liquid to evaporate it, holds solids for weighing, or covers beaker



Crucible: Holds hot objects

Laboratory iron rings: Holds items above work surface



Evaporation dish: Evaporates



Funnel:



Test tube: Holds materials



Ring stand (with attached titration device): Holds iron rings



Spring clamps: Holds test tubes

Click!

Wire gauze: Holds hot materials, diffuses heat



Beakers: Holds liquids for stirring, heating, etc.

Click!

Pipeclay triangle: Supports a crucible over a Bunsen burner

Click!

Test tube brush: Cleans test tubes



Florence flask: Beaker that ensures uniform heating, stirring, with thin neck



Pasteur pipet (dropper): Moves liquid; large mouth; ungraduated



Graduated pipet (dropper): Moves liquid, in measured amounts

Bunsen burner: Emits sootless open gas flame



Disposable weighing boat: Used for weighing



Wash bottle: Used for dispensing fluid



Wash bottle: Used for dispensing fluid

Click!

Crucible tongs: Picks up hot stuff



Forceps: Picks up small stuff

Click!

Beaker tongs: Picks up beakers

General terms

 * 1) Observations
 * 2) Quantitative
 * 3) Qualitative
 * 4) Hypotheses: theoretical explanation for observations
 * 5) Experiments: obtaining new observations, testing hypotheses
 * 6) Theory/Law meaning set 1
 * 7) Theory: set of tested hypotheses
 * 8) Law: theory, applied to broad range
 * 9) Theory/Law meaning set 2
 * 10) Law: what happens
 * 11) Theory: why happens

Measurement

 * 1) Precision vs accuracy:
 * 2) Precision: close to other measurement
 * 3) Accuracy: close to "true" measurement
 * 4) Types of error:
 * 5) Random/Indeterminate error: equal probability of decreasing or increasing result
 * 6) Systematic/Determinate error: consistently skewed probability; decreases accuracy, not precision
 * 7) Measurement: number & unit
 * 8) Mass: Gram (g)
 * 9) Length: Meter (m)
 * 10) Time: Second (s)
 * 11) Temperature: Kelvin (K)
 * 12) Electric current: Ampere (A)
 * 13) Amount of substance: Mole (mol)
 * 14) Luminous intensity: Candela (cd)
 * 15) Prefixes:
 * 16) Mega: 106
 * 17) Kilo: 103
 * 18) none: 101
 * 19) Deci: 10-1
 * 20) Centi: 10-2
 * 21) Milli: 10-3
 * 22) Micro: 10-6
 * 23) Nano: 10-9
 * 24) Pico: 10-12

Significant figures

 * 1) Simplify to scientific notation
 * 2) Zeroes:
 * 3) If placeholder, insignificant
 * 4) If in-number, significant
 * 5) If after number and after decimal point, significant
 * 6) Multiplication and division:
 * 7) Least number of significant digits
 * 8) Addition and subtraction:
 * 9) Round to least significant digit after

History

 * 1) 400 BCE Demokritos, Leucippus use term "atomos"
 * 2) Alchemy
 * 3) 1500s:
 * 4) Georg Bauer: Systematic metallurgy
 * 5) Paracelsus: Medicinal application of minerals
 * 6) 1600s:
 * 7) Robert Boyle: The Skeptical Chemist: Quantitative experimentation, identification of elements
 * 8) 1700s:
 * 9) Georg Stahl: Phlogiston theory
 * 10) Joseph Priestly: Discovery of oxygen
 * 11) Antoine Lavosier: Role of oxygen in combustion, law of conservation of mass, modern chemistry textbook
 * 12) 1800s:
 * 13) Joseph Proust: Law of definite composition
 * 14) John Dalton: Atomic Theory, law of multiple proportions
 * 15) Joseph Gay-Lussac: Combining volumes of gases, existence of diatomic molecules
 * 16) Amadeo Avogadro: Molar volumes of gases
 * 17) Jons Jakob Berzelius: Relative atomic masses, modern symbols for the elements
 * 18) Dmitri Mendeleyev: Periodic table
 * 19) J.J. Thomson: Discovery of the electron
 * 20) Henri Becquerel: Discovery of radioactivity
 * 21) 1900s:
 * 22) Robert Millikan: Charge and mass of the electron
 * 23) Ernest Rutherford: Existence of the nucleus, its relative size
 * 24) Meitner & Fermi: Sustained nuclear fission
 * 25) Ernest Lawrence: The cyclotron and trans-uranium elements

Dalton

 * 1) Atomic Theory (1808)
 * 2) All matter is atoms (extremely small)
 * 3) Atoms in element are identical in size, mass, and other properties; atoms of other elements differ in all
 * 4) Atoms cannot be subdivided, created, destroyed
 * 5) Atoms of diff elements combine in whole-number ratios to form chemical compounds
 * 6) In chemical reactions, atoms are rearranged
 * 7) Changes:
 * 8) none
 * 9) Atoms in element have characteristic average mass, unique to that element.
 * 10) Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions; nuclear reactions can cause these reactions
 * 11) none
 * 12) none

Thomson

 * 1) Cathode ray tube (1897)
 * 2) From classical mechanics:
 * 3) ΔX(+) ∝ e(+)/m(+)
 * 4) ΔX(-) ∝ e(-)/m(-)
 * 5) From the fact that the cathode was H2:
 * e(+)| = |e(-)| (since H is neutral; theoretically, mult electrons could have instead existed, making this incorrect)
 * 1) From the proportions and the balanced equation:
 * 2) ΔX(+)/ΔX(-) = (|e(-)|/m(-))/(|e(+)|/m(+)) = m(-)/m(+)
 * 3) Proposed "plum pudding model"
 * 4) Positive pudding
 * 5) Negative plums

Millikan

 * 1) Oil drop
 * 2) Suspended oil with electricity
 * 3) Electron mass: 9.109 x 10-31 kg
 * 4) Electron charge: 1.5924(17)×10−19 C

Rutherford

 * 1) Gold-foil experiment
 * 2) RaBr (radium bromide) source of α particles
 * 3) Experiment 1: (control)
 * 4) RaBr
 * 5) Detection screen opposite RaBr
 * 6) Results:
 * 7) 132,000 α particles/min
 * 8) Expected
 * 9) Experiment 2: (first gold-sheet experiment)
 * 10) RaBr
 * 11) Gold sheet (thinner than human hair)
 * 12) Detection screen opposite RaBr
 * 13) Results:
 * 14) 132,000 α particles/min
 * 15) Appears to confirm plum-pudding model
 * 16) Experiment 3: (backscatter)
 * 17) RaBr
 * 18) Gold sheet (thinner than human hair)
 * 19) Detection screen opposite RaBr
 * 20) Detection screen on same side as RaBr
 * 21) Results:
 * 22) 20 α particles/min on same-side detector
 * 23) Couldn't be explained as background noise
 * 24) Proves non-zero amount of backscatter (P = count rate backscattered particles / count rate of incident particles = 20/132000 = 2e-4)
 * 25) "As if you'd fired a 15" shell at tissue paper, and it came back and hit you!"
 * 26) Conclusions:
 * 27) Plum-pudding model false
 * 28) Au atoms must be mostly empty
 * 29) Very small, heavy, positively charged area
 * 30) Diameter of nucleus: 1*e-14 meters
 * 31) From the balanced nature of particle: Charge of nucelus: +Ze (atomic number * absolute value of electron charge)

General principles, 1900s

 * 1) Cathode rays have identical properties regardless of element used
 * 2) Atoms are neutral
 * 3) Positive particles must balance electrons
 * 4) Electrons are small; a 3rd particle must exist

Hierarchy

 * 1) Matter
 * 2) Mixtures
 * 3) Homogenuous (Solutions)
 * 4) Heterogenuous (Solutions)
 * 5) Pure substances
 * 6) Compounds
 * 7) Elements
 * 8) Atoms
 * 9) Electrons (is lepton)
 * 10) Protons (3 quarks)
 * 11) Neutrons (3 quarks)

Phase/state differences

 * 1) Solid
 * 2) Definite volume
 * 3) Definite shape
 * 4) Fixed particle positions
 * 5) Fixed particle distances (small)
 * 6) Molecular bonds present
 * 7) Liquid
 * 8) Definite volume
 * 9) Indefinite shape
 * 10) Unfixed particle positions
 * 11) Unfixed particle distances (small)
 * 12) Molecular bonds present
 * Gas
 * 1) Indefinite volume
 * 2) Indefinite shape
 * 3) Unfixed particle positions
 * 4) Unfixed particle distances (large)
 * 5) Molecular bonds present
 * 6) Plasma
 * 7) Indefinite volume
 * 8) Indefinite shape
 * 9) Unfixed particle positions
 * 10) Unfixed particle distances
 * 11) Molecular bonds dissasociated

Extensive vs intensive

 * 1) Extensive (dependent on amount of matter)
 * 2) Volume
 * 3) Mass
 * 4) Energy content
 * 5) Intensive (independent of amount of matter)
 * 6) Melting point
 * 7) Freezing point
 * 8) Boiling point
 * 9) Condensation point
 * 10) Density

Separation

 * 1) Mixture
 * 2) Chromatography
 * 3) Filtration
 * 4) Evaporation
 * 5) Fractional distillation (different boiling points)
 * 6) Compound (must be chemically separated)
 * 7) Electrolysis

Atomic structure

 * 1) Abnormal:
 * 2) Isotope: Abnormal neutron count
 * 3) Ion: Unbalanced electron/proton count
 * 4) Atomic mass: average natural isotope mass

Particulate composition

 * 1) Protons (1e)
 * 2) Two up quarks (+2/3e)
 * 3) One down quark (-1/3e)
 * 4) Neutrons (0)
 * 5) One up quarks (+2/3e)
 * 6) Two down quarks (-1/3e)
 * 7) Gluons hold together particles
 * 8) Electron (-1e)

Molecule vs compound

 * 1) Molecule:
 * 2) 2+ atoms joined chemically
 * 3) Compound:
 * 4) Molecule
 * 5) 2+ different elements joined chemically

Bonding types

 * 1) Two atoms close enough for atomic orbital to mix. Electronegativity values are:
 * 2) Very different: Ionic bonding.
 * 3) Similar: The atoms are classified as:
 * 4) Metals: Metallic bonding.
 * 5) Nonmetals: Covalent bonding. Electronegativity values are:
 * 6) Very close: Nonpolar covalent bonding. (Electron is equally shared; no polar region of molecule.)
 * 7) Different: Polar covalent bonding. (Electron is unequally shared; polar region of molecule.)
 * 8) Discrete/indiscreet:
 * 9) Discrete: Represented by formula
 * 10) Indiscrete: Covalent Network Structures, no formula

Nuclear reactions

 * 1) Mass defect (ΔE = Δmc2)
 * ΔE: change in energy
 * m: mass defect
 * c2: speed of light squared (really big; 8.98755179 × 1016 m2 / s2)

Fission

 * 1) Splitting heavy nucleus into two lighter nuclei (plus neutrons)
 * 2) Uranium fission: 01n + 92235U = 56142Ba + 3691U + 301n (with variation on the fission fragments' mass numbers)
 * 3) Chain reaction: (self-sustaining fission)
 * 4) Ensuring self-sustaining reaction:
 * 5) Increasing size increases likelihood neutrons will collide with additional matter
 * 6) Encasing sphere in neutron reflector (eg tungsten carbide) increases likelihood neutrons will collide with additional matter
 * 7) Fission bomb design:
 * 8) Little Boy: conventional explosive "gun" mechanism shoved two subcritical uranium 235 pieces together
 * 9) Fat Man: conventional explosive compressed plutonium core
 * 10) Fission reactor

Fusion

 * 1) Combining two lighter nuclei into a heavier nucleus
 * 2) Deuterium-tritium: 23He + 11H = 24He + 10n
 * 3) Star fusion:
 * 4) 1.3xSun mass or less: almost all energy from proton-proton fusion
 * 5) 1.3xSun mass or more: almost all energy from CNO cycle

Balancing

 * 1) Σin = Σout
 * 2) ΣAreactants = ΣAproducts
 * 3) ΣZreactants = ΣZproducts

Nuclear balancing practice
88226Ra = 24α + X

86222Rn

01n + 92235U = 53139I + 201n + X

3995Y

Decay

 * 1) Alpha decay
 * 2) Alpha particle emitted
 * 3) Alpha particle is Helium nucleus (24α = 24He)
 * 4) Limited to very large nuclei
 * 5) Example: 92238U = 24α + 90234Th
 * 6) Beta decay
 * 7) Beta particle emitted & neutron converted to proton
 * 8) Beta particle is electron (-10β = -10e)
 * 9) Example: 90234U = 91234Pa + -10e
 * 10) Gamma emission
 * 11) Gamma particle emitted during other decay; nucleus unchanged
 * 12) Gamma particle is high-energy light
 * 13) Example: 92238U = 24α + 90234Th + 200γ
 * 14) Positron emission
 * 15) Positron emitted & proton converted to neutron
 * 16) Positron is anti-electron (10e)
 * 17) Example: 1122Na = 10e + 1022Ne
 * 18) Electron capture
 * 19) Inner-orbital electron captured by nucleus
 * 20) Electron converts proton to neutron
 * 21) Example: 80201Hg + -10e = 79201Au + 200γ
 * 22) Decay reasons: return to stability (proper A and Z)
 * 23) Decay series
 * 24) Shit gets bad
 * 25) Shit gets better
 * 26) ln(N/N0)=-kt
 * 27) N0 = initial nucleides
 * 28) N = remaining nucleides
 * 29) k = rate constant
 * 30) t = elapsed time
 * 31) t0.5 = ln(2)/2 = 0.693/k
 * 32) k = rate constant
 * 33) t = elapsed time

Terms

 * 1) Cation: +
 * 2) Anion: -
 * 3) Ionic bonding: attraction between + and -
 * 4) Predicting charges:
 * 5) Alkali metals: Lose one electron, become 1+
 * 6) Alkali earth metals: Lose two electrons, become 2+
 * 7) Transition metals: no rules!!!!! (except balance electrons, yadda yadda)
 * 8) Boron group: Loses 3 electrons, become 3+
 * 9) Carbon group: Loses or gains 4 electrons
 * 10) Pnictogens: Gains 3 electrons, become 3-
 * 11) Chalcogens: Gain 2 electrons, become 2-
 * 12) Halogens: Gain 1 electron, become 1-
 * 13) Noble gases: No ions

Ionic equation balancing practice
Barium nitrate: Ba2+ + NO3-


 * 2+ + 1- (unbalanced)

Ba2+(NO3-)2
 * 2+ + 2- (balanced)

Ammonium sulfate: NH4+ | SO42-


 * 1+ | 2- (unbalanced)

(NH4+)2SO42-
 * 2+ | 2- (balanced)

Iron(III) chloride: Fe3+ | Cl-


 * 3+ | 1- (unbalanced)

Fe3+Cl3-
 * 3+ | 3- (balanced)

Aluminum sulfide: Al3+ | S2-


 * 3+ | 2- (unbalanced)

Al3+2S32-
 * 6+ | 6- (balanced)

Magnesium carbonate: Mg2+ | CO32-


 * 2+ | 2- (balanced)

Zinc hydroxide: Zn2+ | OH-


 * 2+ | 1- (unbalanced)

Zn2+ | (OH-)2
 * 2+ | 2- (balanced)

Magnesium carbonate: Mg3+ | CO43-


 * 3+ | 3- (balanced)

Naming ionic compounds

 * 1) Cation
 * 2) Monoatomic: name of element (eg, calcium)
 * 3) Transition metals with multiple oxidation states: use Roman numerals in name (eg, II)
 * 4) Anion:
 * 5) Monoatomic: root + ide (eg, chlor-ide)
 * 6) Binary molecular compounds: (between two nonmetals)
 * 7) First element named first with element name; prefix if subscript
 * 8) Second element named second with root-ide; always prefixed
 * 9) Prefixes:
 * 10) 1 = mon(o)
 * 11) 2 = di
 * 12) 3 = tri
 * 13) 4 = tetra
 * 14) 5 = penta
 * 15) 6 = hexa
 * 16) 7 = hepta
 * 17) 8 = octa
 * 18) 9 = nona
 * 19) 10 = deka

Naming ionic compounds practice
Formula to words

H2O:

Hydrogen monoxide

P2O5:

Diphosphorus pentoxide

CO2:

Carbon dioxide

P2O5:

Carbon monoxide

P2O5:

Dinitrogen monoxide

N2O4:

Nitrogen tetroxide

SO3:

Sulfur trioxide

NO:

Nitrogen monoxide

NO2:

Nitrogen dioxide

As2O5:

Diarsenic pentoxide

PCl3:

Phosphorous trichloride

CCl4:

Carbon tetrachloride

SeF6:

Selenium hexafluoride

Words to formula

Diphosphorus pentoxide:

P2O5

Carbon dioxide:

CO2

Carbon monoxide:

P2O5

Dinitrogen monoxide:

P2O5

Silicon dioxide:

SiO2

Carbon tetrabromide:

CBr4

Sulfur dioxide:

SO2

Phosphorus pentabromide:

PBr5

Iodine trichloride:

ICl3

Nitrogen triiodide:

NI3

Dinitrogen trioxide:

N2O3

Solutions

 * 1) Matter; Can be separated by physical means:
 * 2) Yes: Pure substances; Can it be decomposed by ordinary chemical means
 * 3) Yes: Compounds (water, NaCl, sucrose)
 * 4) No: Elements (H, He, Li)
 * No: Mixtures; Is the composition uniform?
 * 1) Yes: Homogenous mixtures (air, sugar in water, stainless steel)
 * 2) No: Heterogenous mixtures (granite, wood, blood)
 * 3) Solute: dissolved substance in a solution
 * 4) Solvent: dissolving medium in a solution
 * 5) Best dissolution:
 * 6) Nonpolar solutes dissolve best in nonpolar solvents (fats, steroids, waxes, benzene, hexane, toulene)
 * 7) Polar and ionic solutes dissolve best in polar solvents (inorganic salts, sugars, water, small alcohols, acetic acid)
 * 8) Factors affecting solubility:
 * 9) Solids:
 * 10) Increases with temperature (generally)
 * 11) Increases with increasing surface area of the solid (generally)
 * 12) Increases with stirring
 * 13) Gases:
 * 14) Decreases with increases in temperature (generally)
 * 15) Increases with pressure above the solution (generally)
 * 16) Saturation:
 * 17) A solution that contains the maximum amount of solute that may be dissolved under existing conditions is saturated.
 * 18) A solution that contains less solute than a saturated solution under existing conditions is unsaturated.
 * 19) A solution that contains more dissolved solute than a saturated solution under the same conditions is supersaturated.
 * 20) Electrolytes:
 * 21) Electrolyte: A substance whose aqueous solution conducts an electric current.
 * 22) Nonelectrolyte: A substance whose aqueous solution does not conduct an electric current.
 * 23) Can be tested by running current through solution and ammeter; no current means nonelectrolyte
 * 24) Ionization:
 * 25) Ionic compounds disassociate (eg NaCl(s) = Na+(aq) + Cl-(aq))
 * 26) Ions prefer positioning to conduct a current rather than forming a solid.
 * 27) Polar water molecule pulls positive ions to the negative end (oxygen) and negative ions to the positive end (hydrogen).
 * 28) Some covalent compounds ionized: hydrogen chloride molecules, which are polar, give up their hydrogens to water, forming chloride ions (Cl-) and hydronium ions (H3O+).
 * 29) Acids:
 * 30) Strong acids are totally ionized
 * 31) Weak acids (often containing carboxyl group) rarely (5%) ionized

Solution concentration

 * 1) Mole fraction – the ratio of moles of solute to total moles of solution
 * 2) Mole fraction of A = xA = nA / (nA + nB)
 * 3) Molarity is the ratio of moles of solute to liters of solution
 * 4) Molarity = M = moles of solute / Liters of solution

Solution concentration practice
Problem: How many grams of sodium chloride are needed to prepare 1.50 liters of 0.500 M NaCl solution?


 * 1) 1.5 L * (0.5 mol / 1 L) * (58.44 g / 1 mol) = 43.8 g NaCl

Problem: What volume of stock (11.6 M) hydrochloric acid is needed to prepare 250. mL of 3.0 M HCl solution?


 * 1) MstockVstock = MdiluteVdilute
 * 2) (11.6 mol / L)(x L) = (3.0 mol / L)(0.250 L)
 * 3) x = (3.0 mol) * (0.250 L) / (11.6 L)
 * 4) x = 0.065 L

How many grams of NaOH are contained in 270.0 mL of a 0.450 M sodium hydroxide solution?

How much of each starting material would you use to prepare 2.00 L of each of the following solutions? (a) 0.250 M NaOH from solid NaOH (b) 0.160 M NaOH from 1.00 M NaOH stock solution (c) 0.150 M K2CrO4 from solid K2CrO4 (d) 0.210 M K2CrO4 from 1.75 M K2CrO4 stock solution

General

 * 1) 1 mole = 6.022e23 (6.022 * 1023) atoms = Avogadro's number
 * 2) Exactly 12 grams of carbon-12 in 1 mole of carbon-12
 * 3) Formulas:
 * 4) Empirical formula: lowest whole-number ratio of atoms in a compound (eg, CH)
 * 5) Molecular formula: actual ratio of atoms in compound; is always multiple of empirical formula (eg, C6H6 = (CH)6)
 * 6) Always empirical for ionic compounds, such as NaCl, MgCl2, Al2(SO4)3, K2CO3
 * 7) Possibly empirical for molecular compounds, such as H2O, C6H12O6, C12H22O11

Moles practice
How many grams of lithium are in 3.50 moles of lithium?

(3.50 moles Li / 1) * (6.94 g Li / 1 mol Li) = 45.1 g Li

How many moles of lithium are in 18.2 grams of lithium?

(18.2 grams Li / 1) * (1 mol Li / 6.94 g Li) = 2.62 mol Li

How many atoms of lithium are in 3.50 moles of lithium?

(3.50 moles Li / 1) * (6.022e23 atoms Li / 1 mol Li) = 2.11e24 atoms Li

How many atoms of lithium are in 18.2 grams of lithium?

(18.2 grams Li / 1) * (1 mol Li / 6.94 g Li) * (6.022e23 atoms Li / 1 mol Li) = 1.58e24 atoms Li

Calculate each of the following quantities. (a) Mass (g) of 0.59 mol of MnSO4 (b) Amount (mol) of compound in 13.3 kg of Fe(ClO4)3 (c) Number of N atoms in 78.2 mg of NH4NO2

Calculate each of the following quantities. (a) Mass (g) of 0.66 mol of KMnO4 (b) Amount (mol) of O atoms in 8.85 g of Ba(NO3)2 (c) Number of O atoms in 7.8 10-3 g of CaSO4 · 2 H2O

Formula Mass
Add up each element's atomic weight, multiplied by the number of times the element appears.

eg: Calculate the formula mass of sodium chloride.


 * 1) sodium chloride = NaCl


 * 1) Sodium atomic weight: 22.99 g Na * 1


 * 1) Chloride atomic weight: 35.45 g Cl * 1


 * 1) Sodium chloride formula mass: 22.99 g Na + 35.45 g Cl = 58.44 g NaCl

Divide each element mass in the formula by total mass and multiply by 100% to find percent composition. Add all percentages to check correctness; should equal approximately 100%.

eg: Calculate the percentage composition of sodium chloride.


 * 1) Sodium percentage composition: (22.99 g Na / 58.44 g NaCl) * 100% = 39.34%


 * 1) Chloride percentage composition: (35.45 g Cl / 58.44 g NaCl) * 100% = 60.66%


 * 1) Total: 39.34% + 60.66% = 100% (+/- a tiny bit is fine)

To determine empirical formula:


 * 1) Assume 100g of compound.
 * 2) Determine moles of each element per 100g.
 * 3) Divide each molesofelement by the smallest molesofelement.
 * 4) If fractional, multiply by reciprocal of smallest fraction until all are whole numbers.

Empirical formula practice
Calculate the formula mass and percentage composition of magnesium carbonate.


 * 1) magnesium carbonate = MgCO3


 * 1) Formula mass:
 * 2) 24.31 g Mg + 12.01 g C + 3(16.00 g O) = 84.32 g MgCO3
 * 3) Percentage composition:
 * 4) Magnesium: (24.31 g/84.32 g) * 100% = 28.83%
 * 5) Carbon: (12.01 g/84.32 g) * 100% = 14.24%
 * 6) Oxygen: (48 g/84.32 g) * 100% = 56.93%
 * 7) Total: 28.83% + 14.24% + 56.93% = 100%

Adipic acid contains 49.32% C, 43.84% O, and 6.85% H by mass. What is the empirical formula of adipic acid?


 * 1) Pretend 100g.
 * 2) Find mass:
 * 3) Carbon: 49.32% C * 100g = 49.32g C
 * 4) Oxygen: 43.84% C * 100g = 43.84g O
 * 5) Hydrogen: 6.85% C * 100g = 6.85g H
 * 6) Find moles:
 * 7) Carbon: 49.32g C/(12.01g C/mol C) = 4.107 mol C
 * 8) Oxygen: 43.84g O/(16.00g O/mol O) = 2.740 mol O
 * 9) Hydrogen: 6.85g H/(1.01g H/mol H) = 6.782 mol H
 * 10) Divide moles:
 * 11) Carbon: 4.107 mol C/2.740 mol O = 1.499 C/O
 * 12) Oxygen: 2.740 mol O/2.740 mol O = 1 O/O
 * 13) Hydrogen: 6.782 mol H/2.740 mol O = 2.475 H/O
 * 14) Multiply:
 * 15) Carbon: 1.499 C/O * 2 O ≈ 3
 * 16) Oxygen: 1 O/O * 2 O ≈ 2
 * 17) Hydrogen: 2.475 H/O * 2 O ≈ 5
 * 18) Consolidate: C3H5O2

The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid?


 * 1) Add masses:
 * 2) Carbon: (12.01g C/mol C) * 3 mol C = 36.03 g C
 * 3) Oxygen: (16.00g O/mol O) * 2 mol O = 32.00 g O
 * 4) Hydrogen: (1.01g H/mol H) * 5 mol H = 5.05 g H
 * 5) Total: 73.08 g
 * 6) Divide given mass by total: 146 g/73.08 g = 2
 * 7) Multiply empirical formula by number: (C3H5O2) * 2 = (C3H5O2)2 = C6H10O4

A chloride of silicon contains 79.1 mass % Cl. If the molar mass is 269 g/mol, what is the molecular formula?


 * 1) Pretend 100g.
 * 2) Find mass:
 * 3) 79.1% Cl * 100g = 79.1 g Cl
 * 4) 79.1 g / (35.453 g/mol) = 2.23 mol Cl
 * 5) 100 g (Cl + Si) - 79.1 g Cl = 20.9 g Si
 * 6) 20.9 g / (28.09 g/mol) = 0.744 mol Si
 * 7) Find ratios:
 * 8) 2.23 Cl / 0.744 = ca. 3 Cl
 * 9) 0.744 Cl / 0.744 = ca. 1 Si
 * 10) Empirical formula: SiCl3
 * 11) Molar mass: 134.449 g/mol
 * 12) 269 g/mol / 134.449 g/mol = 2
 * 13) Molecular formula: Si2Cl6

A compound contains only carbon, hydrogen, and oxygen. Combustion of 8.544 mg of the compound yields 12.81 mg CO2 and 3.50 mg H2O. The molar mass of the compound is 176.1 g/mol. What are the empirical and molecular formulas of the compound?


 * 1) 8.544 mg CxHyOz + ∞ mg O2 = 12.81 mg CO2 + 3.50 mg H2O
 * 2) Molar masses:
 * 3) CxHyOz: 176.1 g/mol
 * 4) CO2: 44.01 g/mol
 * 5) H2O: 18.016 g/mol
 * 6) Product milligrams:
 * 7) 12.81 mg CO2 = 12.01 (element molar mass) / 44.01 (compound molar mass) * 12.81 = 3.496 mg C
 * 8) 3.50 mg H2O = 2.016 (element molar mass) / 18.016 (compound molar mass) * 3.50 = 0.3917 mg H
 * O: 8.544 - (3.496 + 0.3917) = 4.656 mg O
 * 1) Moles:
 * C: 3.496 / 12.011 = 0.2911
 * H: 0.3917 / 1.008 = 0.3886
 * O: 4.656 / 16.00 = 0.2911
 * 1) Ratio:
 * C: 0.2911 / 0.2911 = 1 ; *3 = 3
 * H: 0.3886 / 0.2911 = 1.33 ; * 3 = 4
 * O: 0.2911 / 0.2911 = 1 ; * 3 = 3
 * 1) Empirical formula: C3H4O3
 * 2) Empirical formula mass: 88.06 g
 * 3) 171.6 / 88.06 = 2
 * 4) Chemical formula: C6H8O6

A chloride of silicon contains 79.1 mass % Cl. (a) What is the empirical formula of the chloride? (b) If the molar mass is 269 g/mol, what is the molecular formula?

A 0.370-mol sample of a metal oxide (M2O3) weighs 55.4 g. (a) How many moles of O are in the sample? (b) How many grams of M are in the sample? (c) What element is represented by the symbol M?

Menthol (script M = 156.3 g/mol), the strong-smelling substance in many cough drops, is a compound of carbon, hydrogen, and oxygen. When 0.1595 g menthol was burned in a combustion apparatus, 0.449 g of CO2 and 0.184 g of H2O formed. What is menthol's molecular formula?

Mass Spectrometry
https://en.wikipedia.org/wiki/Mass_spectrometry


 * 1) Equipment:
 * 2) Vacuum pump: 10-5 to 10-8
 * 3) Process:
 * 4) Sample added
 * 5) Sample vaporized
 * 6) Sample ionized
 * 7) Sample accelerated into magnetic field
 * 8) Sample path pulled up/down by electromagnet
 * 9) Effect:
 * 10) Produces spectra of masses from the molecules in a sample of material, and fragments of the molecules.
 * 11) Purpose:
 * 12) Determines elemental composition of a sample.
 * 13) Determines the masses of particles and of molecules.
 * 14) Determines potential chemical structures of molecules by analyzing the fragments.
 * 15) Determines the identity of unknown compounds by determining mass and matching to known spectra.
 * 16) Determines the isotopic composition of elements in a molecule.

Charts

 * 1) Y axis: Amount of each ion detected as compared to the most-detected ion.
 * 2) X axis: m/z (mass (unified atomic mass unit) to charge (of the ion) ratio)

Carbon dioxide (CO2) has a single molecular ion peak at 44 m/z ([CO2]+) and three fragment peaks at 12 m/z ([C]+), 16 m/z ([O]+), and 28 m/z ([CO]+).



Bromine has two isotopes 50.69% 79Br and 49.31% 81Br. This causes it to form three molecular ion peaks at 158 m/z ([79Br79Br]+) 160 m/z ([79Br81Br]+) and 162 m/z ([81Br81Br]+) and two fragment peaks at 79 m/z ([79Br]+) 81 m/z ([81Br]+).



Mass spectrometry charts practice
See: http://www.sciencegeek.net/APchemistry/APtaters/MassSpec.htm

Methylene Chloride (CH3Br). Br is 50.69% 79Br and 49.31% 81Br.





Methylene Chloride (CH2Cl2). Cl is 75.77% 35Cl and 24.23% 37Cl.





Vinyl Chloride (CH2CHCl). Cl is 75.77% 35Cl and 24.23% 37Cl.



Stoichiometry

 * 1) When the equation is balanced it has quantitative significance:
 * 2) C2H5OH +  3O2  =  2CO2  +  3H2O
 * 3) 1 mole of ethanol reacts with 3 moles of oxygen to produce 2 moles of carbon dioxide and 3 moles of water
 * 4) Calculating masses of reactants and products
 * 5) Balance the equation.
 * 6) Convert mass or volume to moles, if necessary.
 * 7) Set up mole ratios.
 * 8) Use mole ratios to calculate moles of desired substituent.
 * 9) Convert moles to mass or volume, if necessary.

Stoichiometry practice
6.50 grams of aluminum reacts with an excess of oxygen. How many grams of aluminum oxide are formed?


 * 1) Aluminum oxide: Alx+3Oy-2
 * 2) Balance: Al2O3
 * 3) Equation: XAl + YO2 = ZAl2O3
 * 4) Balance: 4Al + 3O2 = 2Al2O3
 * 5) Al2O3 molar mass = 26.98 * 2 + 16.00 * 3 = 101.96 g/mol
 * 6) Convert:
 * 7) 6.50 g Al * (1 mol Al / 26.98 g Al) * (2 mol Al2O3 / 4 mol Al) * (101.96 g / 1 mol Al2O3)
 * 8) 6.50 / 26.98 x 2 / 4 x 101.96 = 12.3 g

Bornite (Cu3FeS3) is a copper ore used in the production of copper. When heated, the following reaction occurs.

2 Cu3FeS3(s) + 7 O2(g) = 6 Cu(s) + 2 FeO(s) + 6 SO2(g)

If 3.98 metric tons of bornite is reacted with excess O2 and the process has an 76.9% yield of copper, how much copper is produced?


 * 1) 3.98 metric tons = 3.98e6g
 * 2) 3.98e6g (2 Cu3FeS3(s)) + ∞(7 O2(g)) = x(6 Cu(s))
 * 3) Molar mass Cu3FeS3: 342.681 g/mol
 * 4) (3.98e6) * (1/342.681) * (6/2) * (63.546) * (1/1e6) * (0.769) = 1.70 metric tons Cu

Chromium(III) oxide reacts with hydrogen sulfide (H2S) gas to form chromium(III) sulfide and water. Cr2O3(s) + 3 H2S(g) → Cr2S3(s) + 3 H2O(l) (a) To produce 484 g of Cr2S3, how many moles of Cr2O3 are required? (b) How many grams of Cr2O3 are required?

Metal hydrides react with water to form hydrogen gas and the metal hydroxide. For example, SrH2(s) + 2 H2O(l) → Sr(OH)2(s) + 2 H2(g). You wish to calculate the mass (g) of hydrogen gas that can be prepared from 5.99 g of SrH2 and 4.37 g of H2O. (a) What amount (mol) of H2 can be produced from the given mass of SrH2? (b) What amount (mol) of H2 can be produced from the given mass of H2O? (c) Which is the limiting reactant? (d) How many grams of H2 can be produced?

Diborane (B2H6), a useful reactant in organic synthesis, may be prepared by the following reaction, which occurs in a nonaqueous solvent. NaBH4(s) + BF3(g) B2H6(g) + NaBF4(s)    [unbalanced]

If the reaction has a 75.0% yield of diborane, how many grams of NaBH4 are needed to make 25.0 g B2H6?

Bacterial digestion is an economical method of sewage treatment. 5CO2(g) + 55NH4+(aq) + 76O2(g) reaction arrow identifying the presence of bacteria C5H7O2N(s) + 54NO2-(aq) + 52H2O(l) + 109H+(aq) bacterial tissue

The above reaction is an intermediate step in the conversion of the nitrogen in organic compounds into nitrate ions. How much bacterial tissue is produced in a treatment plant for every 1.0 ✕ 104 kg of wastewater containing 3.0% NH4+ ions by mass?

Limiting reactant practice
What is the limiting reactant in the synthesis of ammonia (NH3) if you have 5.0 moles of N2 and 10.0 moles of H2(g) in the reaction vessel?

N2(g) + 3H2(g) = 2NH3(g)


 * 1) N2 as limiting: (N2)5.0 + (3H2(g))∞ = (2NH3(g))5/2
 * 2) H2 as limiting: (N2)∞ + (3H2(g))10/3 = (2NH3(g))10/6
 * 3) H2 is limiting.

How many milliliters of 0.383 M HCl are needed to react with 7.0 g of CaCO3? 2 HCl(aq) + CaCO3(s) → CaCl2(aq) + CO2(g) + H2O(l)

How many grams of NaH2PO4 are needed to react with 38.23 mL of 0.295 M NaOH? NaH2PO4(s) + 2 NaOH(aq) → Na3PO4(aq) + 2 H2O(l)

A sample of impure magnesium was analyzed by allowing it to react with excess HCl solution: Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g) After 1.62 g of the impure metal was treated with 0.100 L of 0.774 M HCl, 0.0125 mol HCl remained. Assuming the impurities do not react, what is the mass % of Mg in the sample?

Precipitation Reactions

 * 1) Double Replacement:
 * 2) AX + BY = AY + BX
 * 3) The ions of two compounds exchange places in an aqueous solution to form two new compounds.
 * 4) One of the compounds formed is usually a precipitate (an insoluble solid), an insoluble gas that bubbles out of solution, or a molecular compound, usually water.


 * 1) Lead(II) nitrate + potassium iodide = lead(II) iodide + potassium nitrate
 * 2) Double replacement equation: Pb(NO3)2(aq) + 2KI(aq) = PbI2(s) + 2KNO3(aq)
 * 3) Compounds as aqueous ions: Pb2+(aq) + 2 NO3-(aq) + 2 K+(aq) +2 I-(aq) = PbI2(s) + 2K+(aq) + 2 NO3-(aq)
 * 4) Eliminate spectator ions:Pb2+(aq) + 2 K+(aq) +2 I-(aq) = PbI2(s) + 2K+(aq)

Solubility rules
ON AP EXAM

All sodium, potassium, ammonium, and nitrate salts are soluble in water.

NOT ON AP EXAM

Salt chart

Gas
[16:40:28] Yixuan Liu: 1.7 atm Xe and 8.0 atm F2 gas is in a sealed metal container at constant volume. When the Xe is used up, there is 4.6 atm remaining of F2 and a solid in the container. What is the empirical formula of the solid? [16:40:43] Yixuan Liu: and then all the answer choices are XeF with a different subscript for F

Failure of the classical model

 * 1) Coulomb force: F(r) = (-e)(e)/4πε0r2 = -e2/4πε0r2
 * 2) Doesn't work with F = ma!
 * 3) e = absolute value of an electron's charge
 * 4) r = distance between two charges
 * 5) ε0 = permittivity constant of a vacuum = 8.854e-12 C2J-1m-1
 * 6) Photoelectric effect:
 * 7) Electron release:
 * 8) ν0 = threshold frequency
 * 9) If ν < ν0: no electrons released
 * 10) If ν > ν0: electrons released
 * 11) Above ν0, no additional electrons released (until a significantly higher v)
 * 12) Kinetic energy vs frequency:
 * 13) Predicted: ν has no relation on KE
 * 14) Observed: above ν0, ν has linear relationship with KE
 * 15) Kinetic energy vs intensity (I):
 * 16) Predicted: I exponentially increases KE
 * 17) Observed: I has no relation on KE
 * 18) Reason: unless individual particle has energy > ν0, doesn't emit electron
 * 19) Number of electrons emitted vs intensity (I):
 * 20) Predicted: I has no effect on #e
 * 21) Observed: I has linear relationship with #e
 * 22) Einstein plotted KE against ν for several metals
 * 23) All had same slope: 6.626e-34 Js (Planck's constant, "h")
 * 24) All had same y intercept: (6.626e-34)ν0
 * 25) KE = hν - hν0
 * 26) Both sides are energy units!
 * 27) Energy of photon proportional to frequency
 * 28) Quantized packets of light
 * 29) hν = Ei = energy of incident photon
 * 30) hν0 = φ = workfunction
 * 31) Therefore:
 * 32) KE = Ei - φ
 * 33) Ei = KE + φ

Photoelectric effect practice

 * http://www.sciencegeek.net/Activities/Ehf.html

If a beam of light with energy 4.0 eV (1 eV = 1.602e-19 J) strikes a gold surface (φ = 5.1 eV), what is the maximum kinetic energy of the ejected electrons?


 * 1) KEmax = Ei - φ
 * 2) Ei = 4.0 eV
 * 3) φ = 5.1 eV
 * 4) KEmax = 4.0 eV - 5.1 eV
 * 5) KEmax = -1.1 eV = 0

Electromagnetic waves

 * 1) E(x1t) = Acos(2πx/λ - 2πνt)
 * 2) E = electric field
 * 3) x = position of the wave
 * 4) t = time
 * 5) ν = frequency
 * 6) A = amplitude
 * 7) E(x1t) = Acos(2πx/λ)
 * 8) only if time constant
 * 9) A max at intervals of λ
 * 10) E(x1t) = Acos(2πνt)
 * 11) only if position constant
 * 12) A max at intervals of int/λ
 * 13) Speed = λν m/sec
 * 14) c = 2.9979e8

Zinc experiment

 * 1) φ Zn = 6.9e-19 J
 * 2) λ UV lamp = 254 nm
 * 3) λ red laser pointer = 700 nm
 * 4) Calculate energy:
 * 5) Equation (for light):
 * 6) E = hν
 * 7) ν = c/λ
 * 8) E = hc/λ
 * 9) UV lamp: E = (6.626e-34)(2.998e8)/(254e-9) Js*m/s/m = 7.82e10-19 J
 * 10) Red laser pointer: E = (6.626e-34)(2.998e8)/(700e-9) Js*m/s/m = 2.84e10-19 J
 * 11) Ejects electrons?
 * 12) UVL: yes
 * 13) RLP: no
 * 14) Calculate photons:
 * 15) UVL: (1e-3J/s) * (photon/2.84e-19 J) * 60s = 2.1e17 photons
 * 16) It works!

Single-electron atoms

 * 1) Schroedinger equation: ĤΨ = EΨ
 * 2) Ĥ = Hamiltonian operator
 * 3) Ψ = wavefunction (description of the orbital)
 * 4) E = binding energy

Hydrogen orbital change
Wavefunctions: http://quantum.phys.cmu.edu/CQT/chaps/cqt02.pdf
 * 1) For a hydrogen atom:
 * 2) ĤΨ = ΨE
 * 3) ĤΨ = Ψ * (-1/n2) * (me4/8ε02h2)
 * 4) m = me = mass of electron
 * 5) e = charge on the e-
 * 6) ε0 = permittivity constant of a vacuum = 8.854e-12 C2J-1m-1
 * 7) h = 6.626e-34 Js
 * 8) n = principle quantum number (interger)

Rydberg constant in binding energy in general

 * 1) Rydberg constant (R):
 * 2) ĤΨ = Ψ * -RH/n2
 * 3) RH = (me4/8ε02h2) = 2.18e-18 J
 * 4) Binding energy (E): E = -RH/n2 (Always negative)
 * 5) Ionization energy (IE): IEn = -En (Always positive)

States of excitement in general

 * 1) "Excited state"
 * 2) n1 = ground
 * 3) n2 = first excited state
 * 4) n3 = second excited state
 * 5) n4 = third excited state
 * 6) En = -Z2RH / n2 (ONLY IF 1 ELECTRON; EITHER HYDROGEN OR 1+ IONS)
 * 7) He1+ 1e- atom: Z = 2
 * 8) Li2+ 1e- atom: Z = 3
 * 9) Tb64+ 1e- atom: Z = 65

Series of hydrogen orbital changes

 * 1) Series:
 * 2) nf = 1: Lyman series, UV range
 * 3) nf = 2: Balmer series, UV range
 * 4) nf = 3: Paschen series, UV range
 * 5) nf = 4: Brackett series, UV range

Photon emission in hydrogen orbital change

 * 1) Setup
 * 2) Evacuated glass tube, filled with H2
 * 3) Negative electrode, positive electrode
 * 4) Disperse emission
 * 5) Analyse wavelength
 * 6) ν = ΔE/h = (Ei - Ef)/h
 * 7) Four wavelengths:
 * 8) Red: n=3 to n=2; 656 nm; longest wavelength, lowest frequency, lowest energy
 * 9) Green: n=4 to n=2; 486 nm
 * 10) Purple: n=5 to n=2; 434 nm
 * 11) Violet: n=6 to n=2 (shortest wavelength, highest frequency, highest energy); 410 nm

Frequency in hydrogen orbital change

 * 1) General:
 * 2) Combine
 * 3) ν = (Ei - Ef)/h
 * 4) En = -RH / n2
 * 5) ν = ((-RH / ni2) - (-RH / nf2))/h
 * 6) Simplify
 * 7) ν = (-RH/h) * ((1/ni2) - (1/nf2))/h
 * 8) ν = (RH/h) * ((1/nf2) - (1/ni2))
 * 9) For nf=2:
 * 10) ν = (RH/h) * (1/4 - (1/ni2))

Rydberg constant in frequency in general

 * 1) Rydberg constant Mk. II:
 * 2) RH/h = ℜ = 3.29e15 s-1
 * 3) R∞ = 1.097373157e7 s-1

Energy levels practice in general
Calculate the wavelength of radiation emitted by a hydrogen atom when an electron makes a transition from the n=3 to the n=2 energy level. (Eundefined to Eundefined).


 * 1) Frequency:
 * 2) ν = (RH/h) * ((1/nf2) - (1/ni2))
 * 3) ν = 2.18e-18/6.626e-34 * ((1/(2^2)) - (1/(3^2))
 * 4) ν = 4.5695409e+14 s-1
 * 5) Wavelength:
 * 6) λ = c/ν = 2.998e8/4.5695409e+14 = 6.56e-7 m = 656 nm

1e- atom

 * 1) Frequency:
 * 2) For ni > nf: ν = (RHZ2/h) * ((1/nf2) - (1/ni2))
 * 3) For nf > ni: ν = (RHZ2/h) * ((1/ni2) - (1/nf2))
 * 4) Binding energy:
 * 5) Hydrogen: En = -RH/h
 * 6) All 1-electron: En = -Z2RH/h

Quantum numbers
http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch6/quantum.html
 * 1) Ψnlm(r,θ,φ)
 * n
 * 1) Name: Principle quantum number
 * 2) Describes: Total binding energy of electron (potential + kinetic)
 * 3) Describes: Shell (and number of subshells)
 * 4) Rules: Any positive integer above 1
 * l
 * 1) Name: Angular momentum quantum number
 * 2) Describes: Angular energy of the electron
 * 3) Describes: Subshell (l = 0 (s), l = 1 (p), l = 2 (d), l = 3 (f))
 * 4) Rules: Any positive integer above 0 up to l = n-1
 * 5) m (or ml)
 * 6) Name: Magnetic quantum number
 * 7) Describes: Shape of orbital / how electron behaves in magnetic field / z component of angular momentum
 * 8) Describes: Subshell description (m = -1 (x), m = 0 (z), m = +1 (y))
 * 9) Rules: Any integer above 0 from m = -l to m = l
 * 10) Alternate description:
 * 11) Ψ100(r,θ,φ) = 1s orbital
 * 12) Degenerate orbitals: for n orbitals, n2 orbitals are degenerate (have the same energy)
 * 13) For 1-electron atoms, all subshells have the same energy
 * 1) For 1-electron atoms, all subshells have the same energy

Wavefunction physical interpretation
https://youtu.be/Pj2fkkZ6Gto?t=1630
 * 1) Max Born
 * 2) [Ψnlm(r,θ,φ)]2
 * 3) Probability/Volume
 * 4) Probability of finding an electron in a given volume
 * 5) Never reaches 0 through space
 * 6) Components
 * 7) Ψnlm(r,θ,φ) = Rnl(r) x Ylm(θ,φ)
 * 8) Wavefunction = radial wavefunction x angular wavefunction
 * 9) For all s orbitals, Y is a constant
 * 10) Areas of 0 probability: Nodes
 * 11) Only occurs between orbitals
 * 12) Number of nodes = n - l - 1

Radial probability distribution

 * 1) For s orbitals: RPD = Ψ2 * 4πr2Ψ2dr
 * 2) Probability = Probability/Volume * Volume
 * 3) Most probable : rmp = a0 = Bohr radius = 0.529 Ångstroms

p orbitals

 * 1) l = 1
 * 2) 3 orbitals
 * 3) for m = +/-1, px/py
 * 4) θ and φ (angular) dependence
 * 5) Total of 6 lobes
 * 6) Each plane: Two lobes, separated by nodal plane
 * 7) Ψ22p y =
 * 8) Highest probability: Along Y axis
 * 9) Positive Ψ: Where y is positive
 * 10) Nodal plane: xz plane (φ = 0 degrees)
 * 11) Nodes:
 * 12) Total: n - 1
 * 13) Angular: l
 * 14) Radial: n - l - 1
 * 15) As l increases, rmp decreases

Spin magnetic quantum number

 * 1) ms = +1/2 (spin up) or -1/2 (spin down)
 * 2) Independent of orbital, only describes electron
 * 3) Discovery:
 * 4) Emission spectrum of sodium
 * 5) Would expect 1 line at given frequency
 * 6) 2 lines very slightly above and below given frequency (doublet)

Pauli exclusion principle

 * 1) No two electrons in same atom can have same quantum numbers
 * 2) Distinction: orbital has 3 quantum numbers, electron 4
 * 3) Limits to 2 electrons per orbital

Multi-electron atoms
How many electrons can exist in a 2p orbital? 6. 2p has 3 complete orbitals, 2px, 2py, 2pz; each complete orbital can have 2 electrons.

Shroedinger equation for multiple electrons
Base: ĤΨ = EΨ
 * 1) H (1 electron): ĤΨ(r1θ1φ1) = EΨ(r1θ1φ1)
 * 2) He (2 electrons): ĤΨ(r1θ1φ1r2θ2φ2) = EΨ(r1θ1φ1r2θ2φ2)
 * 3) Li (3 electrons): ĤΨ(r1θ1φ1r2θ2φ2r3θ3φ3) = EΨ(r1θ1φ1r2θ2φ2r3θ3φ3)
 * 4) At high enough levels, it is mathematically impossible to solve the Shroedinger equation; instead, an approximation is used.

Hartree orbitals
Approximation: Treat the wavefunction of the many orbitals as the product of a one electron approximation for each.
 * 1) He (2 electrons):
 * 2) Ψ(r1θ1φ1r2θ2φ2) = Ψ(r1θ1φ1) * Ψ(r2θ2φ2)
 * 3) Ψ(e-#1, e-#2) = Ψ(e-#1) * Ψ(e-#2)
 * 4) Ψ100-1/2 * Ψ100+1/2
 * 5) 1s(1) * 1s(2)
 * 6) Li (3 electrons):
 * 7) Ψ(r1θ1φ1r2θ2φ2r3θ3φ3) = Ψ(r1θ1φ1) * Ψ(r2θ2φ2) * Ψ(r3θ3φ3)
 * 8) Ψ(e-#1, e-#2, e-#3) = Ψ(e-#1) * Ψ(e-#2) * Ψ(e-#3)
 * 9) Ψ100-1/2 * Ψ100+1/2 * Ψ200-1/2
 * 10) 1s(1) * 1s(2) * 2s(1)

Shorthand

 * 1) H (1 electron): 1s1
 * 2) He (2 electrons): 1s2
 * 3) Li (3 electrons): 1s22s2
 * 4) Be (4 electrons): 1s22s1
 * 5) B (5 electrons): 1s22s22p1

Multi-electron vs hydrogen atom wavefunctions
Argon: 1s22s22p63s23p6
 * 1) Similarities:
 * 2) Similar in shape
 * 3) Identical Radial Probability Distribution and nodes
 * 4) Differences:
 * 5) Each multi-electron orbital is smaller than the corresponding hydrogen atom orbital, because the nucleus is more positively charged.
 * 6) Orbital energy depends on both the shell (n) and the subshell (l) or angular momentum quantum number.
 * 7) Binding orbital energy in multi-electron atoms is lower (more negative) than in corresponding H-atom orbitals.
 * 8) In hydrogen, all orbitals for a given n have same energy level; in a multi-electron atom, orbitals increase in energy for both n and l
 * 9) For 1-electron: En = -IEn = -Z2RH/n2
 * 10) For multi-electron: En = -IEn = -(Znleff)2RH/n2
 * 11) Z != Zeffective
 * 12) Shielding and binding energy:
 * 13) Scenario: Helium (He) nucleus has 2 protons (z = 2). Electron 1 is very close to nucleus. Electron 2 is far enough out that it doesn't affect electron 1. Electron 1 is thus "no shielded" and electron 2 is "totally shielded".
 * 14) EHe = -IEHe = -(Zeff)2RH/n2 = -(+2)2RH/12 = 8.72 x 10-18 J, Zeff = 1
 * 15) EHe = -IEHe = -(Zeff)2RH/n2 = -(+1)2RH/12 = 2.18 x 10-18 J, Zeff = 2
 * 16) Reality: IEHe = 3.94 x 10-18 J
 * 17) Finding Zeff:
 * 18) Zeff = [n2(IE/RH)]1/2
 * He: [12(3.94 x 10-18 J/2.18 x 10-18 J)]1/2 = 1.34
 * 1) Shielding
 * 2) S is less shielded than P, because over RPD, average Zeff of 2p is less than Zeff of 2s (etc.)
 * 3) RPD of orbitals

Electron configurations

 * 1) Aufbau principle:
 * 2) Pauli Exclusion Principle: Only 1 spin in any given suborbital
 * 3) Hund's rule: At the same energy level, a single e- enters state before a second enters said state
 * 4) Oxygen: https://youtu.be/f7RRqxv2pzg?t=2198
 * O: 1s22s22p4
 * 1) Oml: 1s22s22px22px12px1
 * 2) Sodium:
 * Na: [Ne]3s1
 * 1) Naml: 1s22s22px22px22px23s1

Exceptions

 * 1: Half-filled d orbitals are more stable than half-filled s
 * V: [Ar]4s23d3
 * 1) Cr: [Ar]4s13d5
 * Mn: [Ar]4s23d5
 * 2: Filled d orbitals are more stable than half-filled s
 * Ni: [Ar]4s23d8
 * 1) Cu: [Ar]4s13d10
 * Zn: [Ar]4s23d10

Ion electron configurations
A d orbital with 2+ electrons has less energy (is more stable) than a s orbital. When removing the highest-energy atoms (making something ionic) an S


 * Vi: [Ar]4s23d3
 * 1) Vi (reordered from lowest to highest energy orbital): [Ar]3d34s2
 * Vi-: [Ar]4s13d3

Practice
http://www.sciencegeek.net/Chemistry/taters/Unit2ElectronNotations.htm http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch6/electronconfigpractice.html

Photoelectric spectroscopy (PES)
https://en.wikipedia.org/wiki/Spectroscopy
 * 1) X-rays are commonly used because they have sufficient energy; UV rays may be, but are sometimes insufficient.
 * 2) Neon:
 * 3) 1s22s22p6 -> 1s22s22p5 + e- with most KE
 * 4) 1s22s22p6 -> 1s22s12p6 + e- with middle KE
 * 5) 1s22s22p6 -> 1s12s22p6 + e- with least KE
 * 6) IE = Ei + KE
 * 7) IE = Ionization energy
 * 8) Ei = Energy of photon
 * 9) KE = Kinetic energy of electron
 * 10) Neon IE (with Ei = 1254)
 * 11) 1s22s22p6 ; KE = 1232 eV ; IE2p = 22 eV (1254 - 1232)
 * 12) 1s22s22p6 ; KE = 1206 eV ; IE2p = 48 eV (1254 - 1206)
 * 13) 1s22s22p6 ; KE = 384 eV ; IE2p = 870 eV (1254 - 384)
 * 14) Each would be represented as a bar with x = KE in order 1s, 2s, 2p
 * 15) Practice: If PES produces 5 lines, which elements could it be?
 * 16) 5 lines, 5 orbitals: 1s2s2p3s3p: Al, Si, P, S, Cl, Ar
 * 17) Practice: PES would reveal how many lines in hafnium (z = 72)?
 * 18) 1s2s2p3s3p3d4s4p4d4f5s5p5d6s (or 1s2s2p3s3p4s3d4p5s4d5p6s4f5d)
 * 14

https://youtu.be/LPh2Ut7D4WA?t=746

Ionization energy

 * 1) First (smallest) ionization energy
 * 2) Implied whenever "ionization energy" is stated
 * 3) IE = -Enl
 * 4) Boron:
 * 5) B(1s22s22p1) -> B+(1s22s2) + e-
 * 6) ΔE = Ep - Er = IE = -E2p
 * 7) Second ionization energy
 * 8) FILLER
 * 9) Boron:
 * 10) B+(1s22s2) -> B2+(1s22s1) + 2e-
 * 11) ΔE = IE2 = -E2s in B+
 * 12) Third ionization energy
 * 13) FILLER
 * 14) Boron:
 * 15) B2+(1s22s1) -> B3+(1s2) + 3e-
 * 16) ΔE = IE3 = -E2s in B+
 * 17) Periodic table:
 * 18) ,
 * 19) Moving right in a row, Z (charge) while n (shell) remains constant; thus, Zeff (effective charge) increases, so IE increases
 * 20) Moving down in column, Z (charge) and n (shell) increases; increasing n overpowers increasing Z, so Zeff decreases and IE decreases
 * 21) Exceptions in Lithium - Neon row: http://staff.norman.k12.ok.us/~cyohn/index_files/ionizationenergynotes_files/image001.gif
 * 22) Beryllium (1s22s2 to Boron 1s22s22p1: Energy required to add new orbital is greater than the increased distance from the nucleus
 * 23) Nitrogen (1s22s22p3 to Boron 1s22s22p4: Oxygen has 4 electrons in a 6 electron-orbital (3 subshells of 2 each) which forces 2 electrons to pair up


 * Which element has smaller IE (and why): Al or P?
 * Al (lower Zeff)
 * P (lower Zeff)
 * Al (higher Zeff)
 * P (higher Zeff)

Electron affinity

 * 1) Electron affinity (EA or Eea) = -ΔE
 * 2) Chlorine:
 * 3) Cl + e- -> Cl- ; ΔE = -349 kJ/mol ; EA = 349 kJ/mol
 * 4) Energy is released, so negative ion is more stable than atom.
 * 5) Nitrogen:
 * 6) N + e- -> N- ; ΔE = +7 kJ/mol ; EA = -7 kJ/mol
 * 7) Energy is added, so atom is more stable than negative ion.
 * 8) Periodic table:
 * 9) Moving right in a row, EA increases.
 * 10) Moving down in a column, EA decreases.

Electronegativity

 * 1) Electronegativity (χ) is proportional to (0.5)(EA + IE)
 * 2) Periodic table:
 * 3) Upper right: high χ (electron expector)
 * 4) Bottom left: low χ (electron donor)

Atomic radius

 * 1) Atomic radius: value of r which encompasses 90% of electron density
 * 2) Atomic radius: value of r which is 0.5 of distance between two atoms in a compound
 * 3) Both are very similar values
 * 4) Periodic table:
 * 5) Moving right in a row, r decreases (because ZEff decreases).
 * 6) Moving down in a column, r increases (because new orbitals are larger).
 * 7) Ions:
 * 8) F- is larger than F (more shielding, less ZEff))
 * 9) Na+ is smaller than Na (less shielding, more ZEff))

Isoelectronic

 * 1) Neon: 1s22s22p6
 * 2) Flourine:
 * F: 1s22s22p5
 * F-: 1s22s22p6
 * 1) Sodium:
 * Na: 1s22s22p63s1
 * 1) Na+: 1s22s22p6
 * 2) Which atom is isoelectronic with Krypton (Z=36)? Selenium2- (Z=34).

Covalent bonds

 * 1) Chemical bonds: rearrangement of the nuclei and electrons of the bonded atoms results in a lower energy than separate atoms.
 * 2) Covalent bond: a pair of electrons shared between 2 atoms.
 * 3) Terms:
 * 4) Internuclear distance: r
 * 5) Energy: E = nuclear-nuclear repulsion + electron-nuclear attraction + electron-electron repulsion
 * 6) Disassociation energy: ΔEd = energy necessary to break covalent bond = bond strength = distance from EH + H to EH 2
 * 7) Energy vs interatomic distance
 * 8) Example with H:
 * 9) H+ + e- = 0 kJ/mol
 * 10) 2H (unbonded): E = 2 * -1312 kJ/mol = -2624 kJ/mol
 * 11) H2 (bonded): E = -3048 kJ/mol
 * 12) H2 is lower energy, so 2H bonds
 * 13) ΔEd = -2624 kJ/mol - (-3048 kJ/mol) = 424 kJ/mol
 * 14) N2 vs H2: N2 is stronger bond, because its ΔEd is lower than H2's, which means that the bond is more stable.

Quantitative chemical shit
Finding the stuff you've got in a pile of shit

Gravimetric analysis

 * http://www.dynamicscience.com.au/tester/solutions1/chemistry/analytical%20chem/2007nswgrv1.gif
 * https://en.wikipedia.org/wiki/Gravimetric_analysis#Procedure
 * 1) Obtain sample (analyte) with some
 * 2) Form precipitate: Add solute to form insoluble precipitate containing sample and some known (eg, AgNO3
 * 3) Isolate precipitate:
 * 4) Filter the precipitate from the solvent (eg, using sintered glass filter and vacuum pump)
 * 5) Precipitate has water removed (eg, drying/heating/etc.)
 * 6) Calculate mass of sample
 * 7) Precipitate is weighed
 * 8) Proportion of mass attributable to the precipitating agent is subtracted
 * 9) Remainder is mass of analyte

Later

 * molecular geometry https://www.google.com/search?sourceid=chrome-psyapi2&ion=1&espv=2&ie=UTF-8&q=molcular%20geometry&oq=molcular%20geometry&aqs=chrome..69i57j0l5.2521j0j7
 * http://intro.chem.okstate.edu/1314F00/Lecture/Chapter10/VSEPR.html
 * https://en.wikipedia.org/wiki/Molecular_geometry
 * http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Molecular_Geometry
 * http://www.chemmybear.com/shapes.html
 * http://2012books.lardbucket.org/books/principles-of-general-chemistry-v1.0/section_13/26009d78c259f1040f28320caaf413ba.jpg


 * Shroedinger: http://hyperphysics.phy-astr.gsu.edu/hbase/quantum/scheq.html


 * TITRATION!!!


 * Reorganize


 * General naming rules, move ionic naming rules


 * Practice photoelectric effect.


 * Gas laws


 * Solubility rules (NO3, NH4, and Group I)


 * http://ocw.mit.edu/courses/chemistry/5-111-principles-of-chemical-science-fall-2008/exams/

Dtown

 * 1) http://www.webassign.net/web/Student/Assignment-Responses/last?dep=12315235
 * 2) http://www.webassign.net/web/Student/Assignment-Responses/last?dep=12370788
 * 3) http://www.webassign.net/web/Student/Assignment-Responses/last?dep=12466037
 * 4) http://www.webassign.net/web/Student/Assignment-Responses/last?dep=12518626

Redox
http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Redox_Chemistry/Balancing_Redox_reactions http://www.webqc.org/balance.php?reaction=Ag%2BHNO3%3DAgNO3%2BH2O%2BNO2 http://www.chemteam.info/Redox/Balance-Redox-Acid.html http://www.sciencegeek.net/APchemistry/APtaters/Redox/redox1.htm