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AP Chemistry Notes
Stephen Akiki Colchester High School
Download at http://akiscode.com/apchem
♥♠♣♦ Special Thanks to Stephen Bosley (Boser)
1 FOREWORD/DISCLAIMER
First and formost, I am going to say what everone has on their minds. No you really should not just forget about taking notes anymore in AP Chemistry class because of this packet. This packet is meant to be a review and should be used as such. However that does not mean you can use this packet as your main notes and write notes in the margins to supplement your learning. Please take into account that this entire thing was written over the course of 4 days. As such it is inevitable that I made mistakes in spelling and/or formulas. If you have any questions/comments/fixes to the text you can email me at [email protected]
- 1 FOREWORD/DISCLAIMER Contents
- 2 Solubility Rules
- 2.1 Soluble
- 2.2 Insoluble
- 2.3 Naming Rules
- 3 Periodic Table of Elements
- 4 Poly Atomic Naming
- 5 Common Units, Constants and Charges
- 5.1 Fundamental Constants
- 5.2 Charge
- 5.3 Radius
- 6 Atomic Theory
- 6.1 J.J. Thompson
- 6.2 Robert Millikan
- 6.3 Ernest Rutherford
- 6.4 Chadwick
- 6.5 John Dalton
- 7 Naming
- 7.1 Binary
- 7.2 Ionic
- 7.3 Acids
- 7.3.1 Polyatomic
- 7.3.2 Binary
- 8 Cations
- 9 Reaction Type
- 9.1 Combination (Synthesis)
- 9.2 Decomposition
- 9.2.1 Special Binary Salt Splits
- 9.3 Combustion
- 10 Blackbody Radiation
- 11 Bohr Model
- 11.1 Energy Level Formula
- 11.1.1 Energy Change during Level Jumps
- 12 Wavelength
- 13 Quantum Values
- 13.1 Quantum Value Table
- 13.2 Special cases
- 14 Periodicity
- 14.1 Electron Configuration
- 14.2 Isoelectricity
- 15 Nuclear Chemistry
- 15.1 Isotopes
- 15.2 Radiation
- 15.2.1 Alpha Radiation
- 15.2.2 Beta Radiation
- 15.2.3 Gamma Radiation
- 15.2.4 Positron Radiation
- 15.2.5 Electron Capture
- 15.3 Nuclear Equations
- 15.4 Nuclear Stability
- 15.4.1 Forces Invloved
- 15.4.2 Belt of Stability
- 15.4.3 Magic Numbers
- 15.4.4 Half-Life
- 16 Ionization and Affinity
- 16.1 Ionization Energy
- 16.2 Electron Afinity
- 17 Reactions of Metals
- 18 Chemical Bonds
- 18.1 Intramolecular
- 18.1.1 Ionic Bonding
- 18.1.2 Covalent Bonding
- 18.1.3 Metallic Bonding
- 18.2 Intermolecular
- 18.2.1 Ion-Dipole
- 18.2.2 Dipole-Dipole
- 18.2.3 Hydrogen Bond
- 18.2.4 London Dispersion/Van der Waals
- 18.2.5 Intermolecular Flowchart
- 18.3 Rule of Octet
- 19 Lewis Structures
- 19.1 Structures for Atoms
- 19.2 Structures for Ions
- 19.3 Structure for Ions of Molecules
- 19.4 Lewis Structures for Molecular Structures (Covalent)
- 19.5 Resonance Structures
- 20 Lattice Energies of Ionic Solids
- 21 Bond Lengths of Covalent Bonds
- 22 Electronegativity
- 23 Bond Enthalpy
- 24 VSEPR
- 25 Organic Chemistry
- 25.1 Polarity
- 25.2 Alkanes
- 25.3 Alkane Branch Structure Naming
- 25.3.1 Branch Structure Naming Table
- 25.4 Alkenes
- 25.5 Alkynes
- 26 Functional Groups
- 26.1 Alcohol
- 26.2 Aldehyde
- 26.3 Carboxylic Acid
- 26.4 Ester
- 26.5 Ketone
- 26.6 Ether
- 26.7 Amine
- 26.8 Amide
- 26.9 Haloalkane
- 27 Complex Ions
- 27.1 Cations
- 27.2 Anions
- 27.3 Coordination Number
- 27.4 Naming
- 27.4.1 Cations
- 27.4.2 Anions
- 28 Acidic and Basic Redox
- 29 Thermodynamics
- 29.1 Enthalpy
- 29.1.1 Stoichiometry Problems
- 29.1.2 Calorimetry
- 29.1.3 Hess Law
- 29.1.4 Standard Heat of Formation
- 29.2 Entropy
- 29.2.1 State of Matter
- 29.2.2 Number of Moles of Gasses
- 29.2.3 Pressure of Gas
- 29.3 Gibbs Law of Free Energy
- 29.3.1 ∆H, ∆S, ∆G, Relationship Table
- 30 Chemical Kinetics and Rate Laws
- 30.1 Physical State
- 30.2 Concentration
- 30.3 Temperature
- 30.4 Pressure of Gas
- 30.5 Catalysts and Inhibitors
- 30.6 Rate Laws
- 31 Reaction Mechanisms
- 32 Equilibrium
- 32.1 Types of Equilibrium
- 32.2 Equilibrium Constant Expressions
- 32.2.1 Converting Constants
- 33 Gas Laws
- 33.1 Gas Units and Conversions
- 33.2 Ideal Gas Law
- 33.3 Real Gas Law
- 33.4 Combined Gas Law
- 33.5 Daltons Law of Partial Pressures
- 33.6 Gas Collection over a Water Solution
- 34 ICE ICE (Baby)
- 35 Acids and Bases
- 35.1 Definitions of Acids and Bases
- 35.2 pH and pOH
- 35.2.1 Changing Concentrations
- 35.3 Strong Acids and Bases
- 35.3.1 Strong Acids
- 35.3.2 Strong Bases
- 35.4 Weak Acids and Bases
- 35.4.1 Ka Constant
- 35.4.2 Kb Constant
- 35.5 Common Ion Effect
- 35.6 Buffer
- 36 Equilibrium of Saturated, Soluable Salts
- 37 Kinetic Molecular Theory
- 37.1 Postulates:
- 37.2 Root Mean Square Velocity
- 37.3 Effusion and Diffusion
- 37.3.1 Effusion
- 37.3.2 Diffusion
- 37.3.3 Finding the rate
- 38 Electro Chemistry
- 38.1 Identifying Oxidation Numbers
- 38.2 Galvanic/Voltaic Cells
- 38.3 Calculating Cell Potential
- 38.3.1 Nernst Equation to Find E◦cell
- 39 Balancing Redox Reactions
2 Solubility Rules
2.1 Soluble
- Nitrates N O− 3 1 - All nitrates are soluble
- Chlorates ClO− 3 1 - All chlorates are soluble
- Alkali metal Cations and Ammonium cation compounds N H 4 +1 are all soluble
- Chlorides, Bromides, and Iodides are all soluble EXCEPT Ag+1, P b+2, and Hg+
- Acetates - All are soluble except Ag+
- Sulfates - All are soluble except Ba+2, P b+2, Hg+2, Ca+2, Ag+1, and Sr+
2.2 Insoluble
- Carbonates CO− 3 2 - all carbonates are insoluble except alkali metals and ammonium compounds
- Chromates CrO 4 − 2 - all chromates are insoluble except alkali metals, ammonium, Ca+2, and Sr+ 2
- Hydroxides OH−^1 - all hydroxides are insoluble except alkali metals, ammonium, Ba+2, Sr+2, and Ca+2^ although the last two (Sr+2^ and Ca+2) are only slightly soluble so a precipitate can form.
- Phosphates P O− 4 3 all are insoluble except alkali metals and ammonium
- Sulfites SO− 3 2 all are insoluble except alkali metals and ammonium
- Sulfides S−^2 all are insoluble except Alkali metals, alkali earth metals and ammonium
2.3 Naming Rules
- All strong acids and bases are soluble and should be written as the ions when completing net ionic reactions . Sulfuric acid (H 2 SO 4 ) should be written as H+^ + HSO− 41
- The strong acids are: HCL, HBR, HI, HN O 3 , HClO 4 , and H 2 SO 4
- Strong bases are any alkali metal hydroxides (LiOH, N aOH, etc) and Ca(OH) 2 , Sr(OH) 2 , Ba(OH) 2
- All acids and bases should be left in their molecular form: . Acetic acid → HC 2 H 3 O 2
3 Periodic Table of Elements
5.2 Charge
- e−^ charge = − 1. 602 ∗ 10 −^19 coulombs
- p+^ charge = 1. 602 ∗ 10 −^19 coulombs
- Atomic Mass Unit (amu) = 1. 66054 ∗ 10 −^24 . p+^ = 1.0073 amu . n◦^ = 1.0087 amu . e−^ = 5. 486 ∗ 10 −^4 amu
5.3 Radius
Angstroms (
◦ A) = 10−^10 meters
6 Atomic Theory
6.1 J.J. Thompson
- Discovered e−^ and chargemass ratio . Charge to Mass ratio: 1. 76 ∗ 108 Coulombs/Gram (Charge of e−/mass)
- Plum Pudding Model of atom
6.2 Robert Millikan
- Found charge and mass of e−
- Millikan Oil Drop: . Charge oil drops in a field and adjust field until drops levitate
6.3 Ernest Rutherford
- Discovered 3 types of radiation (Decay Particles) . Alpha particles: He2+^ size, very damaging, stoppable - α . Beta particles - e−^ size, damaging, hard to stop - β . Gamma particles - tiny, not so damaging, unstoppable - γ
- Also discovered proton and new dense nucleus model . Rutherford worked with α particles most and discredited Thompsons model of the nucleus
6.4 Chadwick
- Discovers neutron by shooting radiation at light elements and it watching it kick out a neutral particle
6.5 John Dalton
- Four Postulates . Everything made of atoms . Atoms of one element differ from those of a different element . Atoms will combine in whole number ratios . Atoms can not be created or destroyed
- Law of Constant Composition . In a compound, atom ratios are constant
7 Naming
7.1 Binary
- Smallest atomic number comes first
- Second element ends with -ide
7.1.1 Greek Prefixes
- 1-Mono
- 2-Di
- 3-Tri
- 4-Tetra
- 5-Penta
- 6-Hexa
- 7-Hepta
- 8-Octa
- 9-Nona
- 10-Deca Example Cl 2 O Dichlorine Monoxide
7.2 Ionic
- Finding Charge: N a? 3 Cl+1 2 N adc Clba (a∗b) c =^ d
7.3 Acids
7.3.1 Polyatomic
- Per...ate → Per...ic acid . HN O 4 → pernitric acid
- -ate → ic acid . H + N O 3 → HN O 3 (Nitric Acid)
- -ite → ous acid . HN O 2 → nitrous acid
- Hypo...ite → hypo...ous acid . HN O → hyponitrous acid
7.3.2 Binary
- Hydro + (stem)ic . H + Br → Hydrobromic acid . H + N → Hydronitric acid . Hydrocarbonic acid → HC . Carbonic Acid → HCO 3
- Metal Nonmetal ∆ → Metal + Nonmetal (diatomic in nature) . 2 N aCl ∆ → 2 N a + Cl 2
- Metal Chlorates ∆ → Metal Chlorides + O 2 . F e(ClO 3 ) 2 →∆ F eCl 3 + O 2
9.2.1 Special Binary Salt Splits
These binary salts split into different elements
(N H 4 ) 2 CO 3 → N H 3 + H 2 O + CO 2
H 2 SO 3 → H 2 O + SO 2 H 2 CO 3 → H 2 O + CO 2 N H 4 OH → N H 3 + H 2 O H 2 O 2 → H 2 O + O 2
9.3 Combustion
Hydrocarbon + O 2 → CO 2 + H 2 O ....⇓.... CxHy → double x (multiply by 2) then add 2
- C 1 : meth
- C 2 : eth
- C 3 : pro
- C 4 : bu
- C 5 : pent
- C 6 : hex
- C 7 : hept
- C 8 : oct
- C 9 : non
- C 10 : dec
10 Blackbody Radiation
When an object is heated it will emmit radiant energy
E = hν
- E = Energy
- h = Max Plancks constant (6. 626 ∗ 10 −^34 J ∗ s)
- ν = frequency
Photoelectric effect: Metal will give off e−s if light shines on it. Light shining on a clean sheet of metals will release e−s if ν is strong enough.
11 Bohr Model
Neils Bohr:
- Only orbits of certain radii, corresponding to certain definate energies are permitted for the electron in a hydrogen atom.
- An electron in a permitted orbit has a specific energy and is in an allowed energy state. An electron in an allowed state will not radiate energy and therefore will not spiral into the nucleus.
- Energy is emmitted or absorbed by the e−^ only as the e−^ changes from one allowed energy state to another.
- Flawed theory because it only works for hydrogen
11.1 Energy Level Formula
En = (− 2. 18 ∗ 10 −^18 J)( (^) n^12 )
- E 1 : − 2. 18 ∗ 10 −^18 J
- E 2 : − 5. 45 ∗ 10 −^19 J
- E 3 : − 2. 42 ∗ 10 −^19 J
- E 4 : − 1. 36 ∗ 10 −^19 J
- E 5 : − 8. 72 ∗ 10 −^20 J
- E 6 : − 6. 056 ∗ 10 −^20 J
- E∞: 0
11.1.1 Energy Change during Level Jumps
∆E = EF − E 0
- n = 3 → 2 | − 3. 03 ∗ 10 −^19 J
- n = 4 → 2 | − 4. 09 ∗ 10 −^19 J
- n = 5 → 2 | − 4. 578 ∗ 10 −^19 J
- n = 6 → 2 | − 4. 844 ∗ 10 −^19 J
12 Wavelength
12.1 De Broglie Formulas
λ = (^) mvh or λ = hp
- λ = Wavelength
- h = Plancks Constant (6. 626 ∗ 10 −^34 J ∗ s)
- m = Mass of particle in Kg
- v = Velocity of particle ( meterssecond )
- p = Momentum
Example m = 9. 11 ∗ 10 −^28 g v = 5. 97 ∗ 106 m/s λ = 6.^626 ∗^10
− (^34) J∗s (9. 11 ∗ 10 −^31 Kg)(5. 97 ∗ 106 m/s) = 1.^22 ∗^10
− (^10) m
14 Periodicity
14.1 Electron Configuration
14.2 Isoelectricity
Two atoms are considered isoelectric when they gain or lose electrons to become ions and have the same electron configuration as each other.
Example N a+1: 1S 2 , 2S 2 , 2P 6 N e: 1S 2 , 2S 2 , 2P 6
15 Nuclear Chemistry
Nuclear Chemistry involves changes in the nucleus of an atom.
Normal Nuclear Reactions involve electron transfer Reactions involve decay of nucleus i.e. transforming one element into another Reaction affected by factors such Affected by the type of decay and the halflife of what is decaying as pH, temp, pressure, [], etc. Reactions involve relatively small energy: Reactions deal with huge amounts of energy 400 kJ-1500kJ
15.1 Isotopes
Isotopes: Atoms of the same element that have a different number of neutrons
X − A A Z X AX
- X = Element Symbol
- A = Atomic Mass
- Z = Atomic Number
15.2 Radiation
15.2.1 Alpha Radiation
When a big nucleus ejects a He+2^ size chunk of itself.
15.2.2 Beta Radiation
When a neutrally charged particle (equal amount of p+s and e−s) ejects its e−s leaving only the p+s.
15.2.3 Gamma Radiation
When a particle experiences some type of radiation (called * here) that causes the remaining nucleus to collapse. This causes gamma (γ) rays to be emitted. Gamma radiation is also caused when a positron and an electron smash into each other.
15.4 Nuclear Stability
Understanding why are some nuclides are radioactive while others are not.
15.4.1 Forces Invloved
- Electrostatic . Try to rip apart the nucleus because of like charges
- Strong Nuclear . Try to pull together the nucleus because subatomic particles naturally stick together
- The Glue . Neutrons act as the glue and more of it is required when the electrostatic force gets really strong
15.4.2 Belt of Stability
- Area A . More neutrons than protons - Beta decay → creates protons
- Area B . Less neutrons than protons - Positron emission (Smaller B) or Electron Capture (Larger B)
- Area C . Every element above 83 p+^ is radioactive and no glue can hold it together - Alpha decay
15.4.3 Magic Numbers
The Magic Numbers tend to be stable if you have either a proton or neutron in those numbers. If you have both, they are very stable.
(p+) 2 8 20 28 50 82 - (n◦) 2 8 20 28 50 82 126
- If (p+) and (n◦) even → likely stable
- If either is odd → could go either way
- If (p+) and (n◦) odd → likely unstable
15.4.4 Half-Life
The time it takes 12 the amount of a substance to decay.
Example 5g of nuclide 1 2 life of 15 years How much of the original nuclide remains after 45 years? 5 ⇓ (15 years)
⇓ (30 years)
⇓ (45 years) 0.625g
16 Ionization and Affinity
16.1 Ionization Energy
The energy needed to remove an e−^ (how easy it is to lose an e−). Needs energy (+).
16.2 Electron Afinity
How much a gaseous atom will be attracted to a free e−^ (how easy it is to gain an e−). Releases energy (-).
17 Reactions of Metals
Metal Oxides = Basic
- Metal + Water → Metal Hydroxide + H 2
- Metal + O 2 (Li or any non-Alkali metal) → Metal Oxide
- K + O 2 (Any other Alkali metal) → Metal Peroxide (O 2 − 1 ) . K + O 2 → KO 2
- Metal Oxide + H 2 O → Metal Hydroxide . N a 2 O + H 2 O → N aOH
- Metal Oxide + Acid → Salt + H 2 O . N a 2 O + HCL → N aCl + H 2 O
Nonmetal Oxides = Acidic
- Nonmetal Oxide + H 2 O → Acid . CO 2 + H 2 O → H 2 CO 3 . SO 2 + H 2 O → H 2 SO 3 . P 4 O 10 + H 2 O → H 3 P O 4
- Nonmetal Oxide + Base → Salt + H 2 O . CO 2 + N aOH → N a 2 CO 3 + H 2 O
18 Chemical Bonds
When 2 or more atoms are strongly attached (attracted) to each other.
18.1 Intramolecular
These forces act inside an atom or molecule:
18.2.5 Intermolecular Flowchart
18.3 Rule of Octet
Atoms tend to bond in such a way as to gain, lose, or share e−s in order to gain a complete valence (outer s and p).
19 Lewis Structures
19.1 Structures for Atoms
19.2 Structures for Ions
19.3 Structure for Ions of Molecules
19.4 Lewis Structures for Molecular Structures (Covalent)
- Add valence e−s from all the atoms.
- Write the symbols for the atoms. If there are more than 2 atoms, identify the central atom. Connect them with a single line which represents 2 shared e−s. Subtract the number of e−s from total found in step 1. . Central atom will be closest to Si, P or Metaloid staircase.
- Complete octets around the atoms bonded to the central atom (Hydrogen does not get more than 2).
- Place the remaining pairs around the central atom even if doing so gives more than an octet to the central atom.
- If there are not enough pairs to complete an octet in the central atom, then you ned to try using double or triple bonds.
CH 4
CH 2 Cl 2
HN O 3
CO 2
HCN
