Shivi YadavaHima Veeramachaneni

AP Chemistry Study GuideUnit 1 – Basic Concepts in Chemistry

Nomenclatureo Two Non-metals = Covalent

Prefix-element + prefix-element-ide Ex. P2Cl6 – diphosphorous hexachloride

o Prefixes

One MonoTwo Di

Three TriFour TetraFive PentaSix Hexa

Seven HeptaEight OctaNine NonTen Dec

o Metal + Non-metal = Ionic Name both ions

Metal ions are the sameNon-metal ions

o Binary-ide Cl-, Br-, I-, F-

Ex. NaCl – Sodium Chlorideo Polyatomic

No oxygen – “ide” Ex. Mg3S – Magnesium Sulfide

Normal number of oxygen – “ate” (4 on PT – o/e 3) Ex. MgSO4 – Magnesium Sulfate

One less Oxygen – “ite” Ex. MgSO3 – Magnesium Sulfite

Two Less Oxygen – “hypo – ite” Ex. MgSO2 – Magnesium Hyposulfite

One More Oxygen – “per-ate” Ex. MgSO5 – Magnesium Persulfate

One more Sulfur & one less Oxygen – “Thiosulfate”

Ex. MgS2O3 – Magnesium Thiosulfate Transition metal = indicate charge

o Roman Numeral o Latin (two oxidation states)

Higher = “ic” Ex. Sb(ClO)5 – Stibnic Hypochlorite

Lower = “ous” Ex. CuCl – Cuprous Chloride

o Memorize!

o Acids Name negative ion Change ending

“ate” “ic”o Ex. HNO3 – Nitric Acid

“ite” “ous”o Ex. HClO – Hypochlorous Acid

“ide” “hydro-ic”o Ex. HF – Hydrofluoric Acid

Reactionso Basic Guidelines:

All acids are aqueous unless organic Only strong acids and bases break apart 100%

Acids: HClO4, HClO3, HCl, HBr, HI, HNO3, H2SO4

Bases: LiOH, NaOH, KOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

Don’t write physical states in the ionic Hidden Reactions (clues to look for):

If one compound is aqueous that means that it is in water, and the other compound might react with that water

NH4+1 Ammonium

NH3 AmmoniaCrO4

-2 ChromateCr2O7

-2 DichromateC2O4

-2 OxalateC2H3O2

-1 AcetateOH- HydroxideMnO4

-2 PermanganateCH4 Methane

If one compound is an acid or base, then the other compound then the other compound might react with water to form an acid or base

Non-metal oxides that react with water Metal oxides that react with water

Immediately break apart because they don’t exist H2SO3H2O + SO2

H2CO3 H2O + CO2

NH4OH H2O + NH3

Ammonia reactions don’t form water Things that don’t dissolve in water:

H2O 4 gases (CO2, SO2, NH3, and H2S) Anything going against the solubility rules

Solubility Rules Soluble in water

o Alkali metal compoundso Nitrates and nitrateso Chlorates and perchlorateso Acetates (except with Ag+1)o Ammonium compoundso Chlorides, Bromides, Iodides (except with Ag+1, Hg+2, Hg2

+2, Pb+2)

o Flourides (except with Group II metals, Pb+2, Fe+3)o Sulfates, Sulfites (except with Sr+2, Ba+2, Ca+2, Pb+2, Hg2

+2, Ag+1)

o Carbonates, Phosphates, and Chromates are only soluble with alkali metals, ammonium, CaCrO4, SrCrO4

o Hyroxides are only soluble with alkali metals, ammonium, Sr+2, Ca+2, Ba+2

o Sulfides are only soluble with Group I metals, Group II metals and ammonium

o Oxides are only soluble with Group I metals and ammonium

o Synthesis Reactions: A + X -------------> AX Metals react with non-metals to produce binary salts (two elements, no

polyatomic)

Metal oxides (basic anhydrides) react with water to yield bases (metal hydroxides)

Non-metal oxides (acid anhydrides) react with water to yield acids (oxidation number of non-metal does not change – do an imaginary charge check!)

Metal oxides react with non-metal oxides to produce a polyatomic salt (Oxidation number of non-metal does not change – do an imaginary charge check!)

o Decomposition Reactions: AX -------------> A + X Acids with oxygen decompose to give non-metal oxides and water Acids

with oxygen decompose to give non-metal oxides and water (oxidation number of non-metal does not change – do an imaginary charge check!

Metallic hydroxides, or bases, decompose to give metal oxides and water Metallic carbonates decompose to give metal oxides and carbon dioxide Metallic chlorates decompose to give metal chlorides and oxygen Metallic nitrates decompose to give metal nitrites and oxygen Ammonium carbonate decomposes to give ammonia, water, and CO2

Sulfurous acid decompose to give water and sulfur dioxide Carbonic acid decomposes to give water and carbon dioxide Ammonium hydroxide decomposes to give ammonia and water Binary compounds decompose to give two elements (with energy) Hydrogen peroxide decomposes to give water and oxygen Polyatomic salts not listed above can decompose to form the metal oxide

and non-metal oxide that formed them (oxidation number of non-metal does not change – do an imaginary charge check!)

o Single Replacement Reactions: A + BX -------------> AX + B Active metals replace less active metals in ionic compounds in aqueous

solutions Active metals replace H in water to form metal hydroxides (bases) and H2 Active metals replace H in acids to form hydrogen gas and a salt Active non-metals replace less active non-metals in ionic compounds in

aqueous solutions Non-aqueous replacement reactions – reductions of metal oxides by

hydrogen or other gases: H2 + CuO → Cu + H2O (occur at high temperatures!) CO + Fe2O3 → Fe + CO2

o Double Replacement Reactions: AX + BY -------------> AY + BX Formation of a precipitate (solid) governed by the solubility rules

Formation of a gas Common gases are H2S, CO2, SO2, NH3

o Any sulfide (S-2) plus any acid forms H2S gas and a salto Any carbonate (CO3

-2) plus any acid forms CO2, HOH, and a salt

o Any sulfite (SO3-2) plus any acid forms SO2, HOH, and a salt

o Any ammonium (NH4+1) compound plus a soluble

hydroxide form NH3, HOH, and a salt Formation of a molecule – which is a compound that does not dissociate

well in water, due to its covalent nature! It stays together as a molecule! Example – H2O!

Acid-base neutralization is one type – ACID PLUS BASE = WATER PLUS SALT

Hydrolysis – Reverse of an acid-base neutralization – a salt reacts with water – this will only happen with one in a trillion water molecules!

One in a trillion water molecules can break apart into H+1 and OH-1

The salt then breaks apart, and a double replacement reaction occurs, with the salt reacting with the H+1 and the OH-1

Produces an acid and a base every time! Salts are products of neutralization, but salts that undergo

hydrolysis are not neutral!o Salts of a strong acid and a weak base + H2O give an acidic

solutiono Salts of a weak acid and a strong base + H2O give a basic

solutiono Salts of a strong acid and a strong base do not undergo

hydrolysis – their solutions are neutral!o Salts of a weak acid and a weak base + H2O may give an

acidic, basic, or neutral solution – look at the strength of the acid or base produced (Ka or Kb)

o Oxidation-Reduction Reactions In a reduction/Oxidation reaction, one species is oxidized (loses

electrons) and the other species is reduced (gains electrons) The species being oxidized is called the “reducing agent” and the species

being reduced is called the “oxidizing agent” Many oxidation/reduction reactions will occur in either acidic or basic

solution, taking advantage of H+ or OH- ions, along with H2O, to aid the reduction/oxidation

These reactions are written and balanced using the half-reaction method

Acidic Solution- balancing technique Predict products Balance with

o Hydrogen= H+

o Oxygen= H2O Balance

Basic Solution Balance H w/ H+

Add OH- to each side to neutralize H+

Form H2O with H+/ OH-

Balance charged elements Ex. A solution of sodium bromide is added to an acidic solution of

potassium bromatedNaBr(aq) + KBrO3 (aq)

6H+ + Na+1 + 5Br- + K+1 + BrO3- 3Br2 + 3H2O

There are obvious signs to look for in a common redox reaction:Important Oxidizing Agents (These things are reduced!) Formed in Reaction

MnO4- (acid solution) Mn+2

MnO4- (basic solution) MnO2

MnO2 (acid solution) Mn+2

Cr2O7-2 (acid solution) Cr+3

CrO4-2 (basic solution) Cr+3

HNO3, concentrated NO2

HNO3, dilute NOH2SO4, hot concentrated SO2

Metallic Ions Metallous IonsFree Halogens Halide IonsHClO4 Cl-1

Na2O2 OH-1

H2O2 H2OPerhalates, halates, halites Halogens

Important Reducing Agents(These things are oxidized!) Formed in ReactionHalide Ions HalogensFree Metals Metal IonsMetallous Ions Metallic IonsSulfite Ions SO4

-2

Free Halogens (dilute basic solution) Hypohalite IonsFree Halogens (concentrated basic solution) Halate Ions

C2O4-2 CO2

NO2-1 NO3

-1

Sn+2 Sn+4

H2O2 O2

Chromium: dichromate to Cr3+ in acid solution; chromate to Cr(OH)3 in basic solution.

Dichromate ion can turn into chromate in basic solution, and chromate ion can turn into dichromate ion in acidic solution (this is not reduction/oxidation – the Cr still retains a +6 charge)

Oxygen: hydrogen peroxide can acts as an oxidizing agent (reduced to water) and a reducing agent (oxidized to oxygen gas).

Nitrogen: nitrate ion is an oxidizing agent only in acid solution. The reduction product is NO.

Sulfur: sulfate ion is an oxidizing agent only in acid solution. The reduction product is SO2.

o Complex Ion Reactions: Ligand= double charge and sticks onto a transitional metal Transition metal salt + ligand → complex ion If the word excess is in the problem, then it is complex!

Ex. Excess sodium cyanide solution is added to a solution of silver nitrateNaCN(aq) + AgNO3 (aq) Na+1 + CN- + Ag+1 + NO3

-1 Ag (CN)2-1

2CN- + Ag+1 Ag(CN)2-1

Aluminum salt + ligand → complex ion Beryllium salt + ligand → complex ion Both Zn+2 and Al+3 form Zn(OH)4

-2 and Al(OH)4-1 when treated with excess

hydroxide Ag+1, Cu+2, Zn+2, and Cd+2 all form complexes with NH3

Infrequently seen, but has been on the AP, and used in lab: Thiocyanate acts as a ligand and bonds to a transition metal

o A drop of potassium thiocyanate is added to a solution of iron (III) chloride:

o SCN-1 + Fe+3 → Fe(SCN)+2

Ammonia, as a ligand, gets turned into ammonium ion, and the transition metal is freed from being a complex ion

o Dilute hydrochloric acid is added to a solution of diamminesilver (I) nitrate:

o H+1 + Cl-1 + [Ag(NH3)2]+2 → AgCl + NH4+1

o Notice the destruction, rather than the formation, of a complex

Common ligands are: I-1, Br-1, F-1, OH-1, H2O, C2O4-2, NH3, SCN-1, CN-1

It is a good idea to recognize the names of these ligands as well – iodo, bromo, fluoro, hydroxy, aqua, oxalato, ammine, thiocyanato or isothiocyanato, and cyano

To determine coordination number: For aqua complexes of transition metals, C.N. = 6 For others, C.N. = cation charge x 2

o Lewis Acid and Lewis Base Reactions Lewis acid reacts with a Lewis base to form an adduct:

BF3 + NH3 → F3BNH3

Phosphorus (V) oxytrichloride is added to water POCl3 + H2O → H3PO4 + Cl- + H+

Note that molecular phosphorus compounds form acids with water. PCl5 + H2O → H3PO4 + H2O + Cl- + H+ PCl3 + H2O → H3PO3 + Cl- + H+

Organic bases that have unshared pairs of electrons can react with water or other H+ suppliers:

Methylamine gas is bubbled into water: CH3NH2 + H2O → CH3NH3

+ + OH¯ Give and take electrons in order to share

Things to Practice!o Empirical Formulao Stoichiometryo Pv=nrto Dilution Formula: M1V1= M2V2

o Look at hard miscellaneous reactions sheetUnit 2 – Bonding and Molecular Structure

Intramolecular (Chemical)o Chemical properties (flammability)o Ionic

Crystalline solid (usually white) Difference in Electronegativity > 1.7 Transfer of electrons Cation (+)/ Anion (-) Strong bond high melting points Don’t conduct electricity in the solid state electrons can’t flow through

Do conduct electricity in water like dissolves like Strength is dependent on size of ions and charge of ions

Anytime you make a positive ion it becomes smaller because it loses electrons and sublevel

o Ex. CaS would be stronger than NaF because CaS has a greater charge

o Covalent Sharing electrons (2 or more non-metals) Gases Can be liquids and solids if they are large molecules Polar

0.7 < Difference in Electronegativity < 1.7 Slight charge

Non- Polar True non- polar has Difference in Electronegativity of 0 but… Difference in Electronegativity < 0.7 No charge Wax, petroleum large but no charge

o Metallic So strong because there is a sea or web of electrons By getting closer together they are able to attract electrons better

and there is a better flow Intermolecular Bonds (Physical)

o Also known as Van Dar Waals Forceso Physical properties (boiling point, melting point)o Three types of bonds

Hydrogen Bond Strongest Special type of Dipole- Dipole Bond δ+ H bonded to F, N, O

Dipole- Dipole Bond Attraction between the δ+ of one polar molecule and δ- end of

another polar molecule Ex.

London Dispersion

Bond

Weakest (almost non-existent) Frictional force can break the bonds JELLO bond Branched hydrocarbons have less London Dispersion Temporary- constantly being changed

Lattice energyo Amount of energy given off when crystal formso Energy

Positive= not-spontaneous, need energy Negative= spontaneous

o Born- Haber Cycle (Lattice Energy Problems) Ex.

Lewis Dot Diagramo Tells nothing about

shape!o Coordinates covalent

bond- double bond in which outside atom has to give up two electrons in order to share with the central atom

o Formal Charge Check # of e- that should be there - # of e- on the atom Limit formal charge as soon as possible One bond = one electron

o Resonance Chemically identical All formal charge checks are the same! One atom moves around in the Lewis Dot Diagram

o Ex.

VSEPR Theoryo Valence shell electron repulsion theory

LE Theoryo Localized Electron Theoryo Electron geometry- shape of electrons around central atomo Molecular geometry- shape of atoms around central atom

Electron Geometry Molecular Geometry s prs us

prs Hybridization Dipole Moment

Linear Linear 2 0 spTrigonal Planar Bent 2 1 sp2

Trigonal Planar 3 0 sp2 zeroTetrahedral Bent 2 2 sp3

Trigonal Pyramidal 3 1 sp3

Tetrahedral 4 0 sp3 zeroTrigonal Bypramidal Linear 2 3 sp3d zero

T- Shaped 3 2 sp3dSee-Saw 4 1 sp3dTrigonal

Bypramidal 5 0 sp3d zero

Octahedral Square Planar 4 2 sp3d2 zeroSquare Pyramidal 5 1 sp3d2

Octahedral 6 0 sp3d2 zeroo Permanent dipole moment?- cancel out charge or not?o In determining charge look at molecular geometry because these are the

only things that have a permanent chargeo Violations to the octet rule

Less than 8 electrons Be, B

More than 8 electrons Only elements that have a d sublevel 3rd period or below (p, s)

o Hybridization Theory Only central atom Blends its orbitals together to make new ones # you get = # you blend (includes unshared pairs) Sigma bonds- mixed 1st bond Pi bonds- mixed 2nd or 3rd bonds

o Just a reminder!- electrostatic bonds are physical bondso Molecular Orbital Theory

Bond Order ½ (# of bonding electrons- # of antibonding orbitals)

Unit 3 – Thermochemistry

Mattero “stuff”o Has mass and takes up space

Energyo Ability to do work or produce heato State function – does not depend on path it takeso Potential = mgho Kinetic = ½mv2

o Law of Conservation of Energy Can’t create or destroy energy Can only be converted from one form to another

Temperature (T)o Measurement of speed or randomness of particles o It is a state function

Heat (Q)o Not a state function – because it depends on the path it takes

Work (W)o Not a state function – depends on patho Work = Fdo Work = PAho Work = PΔV

When work is on the system or compressingo Work = - PΔV

When work is done by the system or compressing Specific Heat (C)

o The amount of Joules or calories needed to heat 1 gram of a substance by 1o C. Enthalpy (ΔH)

o Exothermic – release heat out of the system (-) ΔH = Hproducts – Hreactants = always negative Spontaneous

o Endothermic Absorbs heat into the system ΔH = Hproducts – Hreactants = always positive Not spontaneous- because they require constant energy

First Law of Thermodynamicso The energy in the universe is constant.

U= EK + EP

o Internal Energy = U ΔE= Q + W

o Q= positive, when flowing into systemo Q= negative, when flowing out of the systemo W= positive, when the work is done by the systemo W= negative, when the work is done to the system

If no work is being doneo ΔH = Qo ΔH = ΔE + PΔV (must be in Pa x m3)o Pressure is constant in this equation

Molar Heat of Combustion = Heat for one mole To figure out ΔH:

o Calorimetry Q=mCΔTo Hess’s Lawo Heat of Formation Table

Be careful to look at physical states Look at coefficients If it’s an element its 0!

Jouleo N x m = Pa x m3

Unit 4 – Atomic Theory and the Nucleus

Mattero Has mass and takes up spaceo Stuff you can touch

o Protons – 1 amuo Neutrons – 1 amu (1 proton + 1 electron)o Electrons – 0 amu

Energyo Lighto Electromagnetic Radiationo Ability to do worko Travels in waveso vλ = C

C = 3.00x108 m/s , v = frequency = waves/second = Hz, λ = wavelength = m FM = Megahz can move because it has more protons than AM

Kilohz Nanometers Meters = multiply by 1x10-9

Matter and Energy can indirecto Only way to move mass is with mass

Photoelectric Effecto Einstein proved that there were photons on light waves

Max Plancko E=hv h = 6.63x10-34J x sec/waves

Quantum theoryo Neils Bohro Higher frequency = more excitement = more light given offo If a lot of frequency – electrons are ionized and are given off creating electricity

Hydrogeno Hydrogen makes red, blue-green, and two violets. o Balmer’s equation – 1/λ = 1.097x107 m-1 (1/22 – 1/n2)

This predicts the wavelength of light hydrogen emits when whole numbers are inserted for n and n cannot equal 0, 1, or 2.

This is only if it is going to the second energy levelo Rydbherg’s equation – Energy = 2.18 x 10-18J x (1/n2)

This is at any level in a hydrogen atom Energylight= RH (1/nf

2 – 1/ni2)

=hv Three Complications of Bohr’s model

o Debroglie Matter travels in waves λ matter = h (constant)/mass x velocity

o Heisenberg’s Uncertainty Principle Can’t simultaneously know an electrons speed and location

o e- interact electrically and magnetically Magnetism – comes from electron spinning

Schroedingero Variables (quantum numbers)o n = principle quantum # any whole # integer ( 1 - ∞) energy level or regiono l = angular momentum quantum # (0 – n-1) sublevel (s=0, p=1, d=2, f=3)o ml = magnetic quantum # (-l – l ) orbitalo ms = spin quantum # ( + ½ , - ½ ) spin

Orbital Diagramso Pauli Exclusion Principle

An orbital can hold 2 electrons maxo Aufbau Principle

Lowest energy levels are filled firsto Hund’s Rule

Fill empty orbital’s firsto Electron Configuration

Ex. N -1s22s22p3

o Ex. 157N Top # = atomic mass number – protons + neutrons, Bottom # =

Atomic number - # of protons\o Diamagnetic – not magnetic – no unpaired electronso Paramagnetic – magnetic – unpaired electrons

To find average weight of an element = The sum of the % of certain weight x weight (amu)

Four patterns or trends in periodic tableo Shielding effect

full energy levels blocking the nucleus and its charge Shielding effect as you go P.T. Shielding effect stays the same as you go across the PT

Ex. B and C have the same shielding effect – they both have the same number of full sublevels

o Effective nuclear charge- how effective the nucleus is #protons - #electrons in full energy levels

Ex. C has more effective nuclear charge than Bo C: 5-2 = 3o B: 6-2 = 4

o Atomic Radius Across P.T. = Atomic Radius effective nuclear charge (adding protons

but not electrons in full energy levels) Down P.T.= Atomic Radius because of increasing shielding effect Exception: it gets bigger because the p sublevel extends past the atom

o Electron Affinity Energy released or required for an atom to take an electron Negative= energy released, like electrons, more stable Positive= energy absorbed, doesn’t like electrons, less stable Across P.T.= EA effective nuclear charge Down P.T. = EA shielding effect increases

o Ionization Energy Energy required to remove an electron from an element Across P.T.= IE effective nuclear charge Down P.T. = IE shielding effect increases

o Electronegativity Tendency to pull on an electron on a bond Across P.T.= EN effective nuclear charge Down P.T. = EN shielding effect increases

Four forceso Weakest Force – strongest force = gravity weak electric/magnetic strong

Radiationo Elements past 82 protons are radioactive because the neutrons can no longer

keep the protons stableo Types of Radiation

Alpha production (α) Emits Helium nuclei- 4

2He2+

Deflected by an electric field Emitted from unstable nuclei with high atomic number Outside of a body, a paper can stop Inside a body- dangerous Ex.

Beta production (β)92238

24

90234U He Th

Emit electrons Fast, penetrating Mass is basically nothing but it is fast (1/100th the speed of light) A neutron in the nucleus is converted to an electron and a proton Usually occurs when atoms have too many neutrons Ex.

Gamma Ray Production (ϒ) Photons of electromagnetic energy Pure energy Moves at speed of light Most energetic Wavelength= 10-12 meters Need more protons to stabilize neutrons Ex.

Positron Production An electron with positive charge – 0

1e It happens when u have a deficiency of neutrons for a given

number protons in nature Usually emits gamma radiation as well but since gamma radiation

has no mass or charge it is optional whether or not to write it in the nuclear equation

Ex. Electron Capture

Very rare Nucleus captures an electron It is the reverse of beta radiation Makes a neutron Releases gamma ray Ex.

Rate of Decayo Rare of decay= -ΔN/ Δt

N= number of radioactive nuclideso Rate of decay is proportional to the number of nuclides of the radioactive

material remainingo Can be written with a constant rate of decay = kNo First order process

ln (Nf /Ni ) = -kt ln Nf – ln Ni = -kt

90234

91234

10Th Pa e

1122

10

1022Na e Ne

80201

10

79201

00Hg e Au

t1/2 = .693/ko Decay Event- proportional to amount of substance

Nuclear Fission and Fusiono Fusion

Combining two like nuclei to form a more heavier and stable nucleus Proton- proton chains Release much higher energies but also take more energy

Splitting U and Pu nuclei is easier than splitting 2 H nuclei because in H you must overcome repulsion

Ex. o Fission

Splitting a heavy nucleus into two nuclei with smaller mass numbers Possible with isotopes of U and Pu Ex. Exothermic process Large, unstable, radioactive nuclei become more stable by forming

smaller, more stable nuclei Chain reaction- self sustaining fission process

Plasmao An ionized gaso Mixture of positive ions and negative ionso 4th state of mattero Neutral chargeo Conduct heat and electricity

Deuteriumo Heavy water (0.02% of all water)o Must have this for a fusion reaction

Tritium- from lithium Lithium- a common metal ΔE = ΔmC2

o ΔE= change in binding energy Binding energy- energy required to decompose the nucleus into its

components or released when the nucleus is formedo Δm= mass defect- mass decreases and some of it changes to energyo The actual mass of any atom is always less than the mathematically predicted

atom.o The mass defect was converted into pure energy when all atoms in the universe

were initially created…

Add up masses of each proton and neutron that make up nucleus Subtract actual mass of nucleus from the combined mass of nucleus from

the combined mass of the components to get mass defect Ex.

o Actual mass of a S-32 atom = 32.066 amu o Predicted mass = (16 p+)(1.0078 amu) + (16 n)(1.0087 amu)

= 32.264 amu o 32.264 amu – 32.066 amu = .198 amu lost to pure energyo .198 amu = mass defect

o When a system loses or gains energy it also gains or loses a quantity of masso Mass is always lost during a nuclear conversiono Creating a nucleus will create a mass defect

Nuclear Binding Energyo The energy required to break downo A nucleus into its components nucleons (particles)- kJ/molo 3 Steps

Determine mass defect Conversion of mass defect into energy Expressing NBE as energy per mole of atoms, or as energy pre nucleon

Unit 5 – States of Matter

Kinetic Energy = ½mv2 = 3/2RT R – constant, T – temperature Universal Gas Constant = PV/nT = R

o R= 0.0821 L x atm/mol x K, 8.314 Pa x m3/mol x K Combined Gas Law - V1P1/T1n1 = V2P2/T2n2 = k(R)

o Charles’s Law – V1/T1 = V2/T2 P and n are constanto Gay Lusaac’s Law – P1/T1 = P2/T2 V and n are constanto Boyle’s Law – V1P1=V2P2 T and n are constant

Ideal Gas Law – PV = nRT Dumas Law – MmP = DRT Avagadro’s Law – n1/V1 = n2/V2

Gas STP – 273 K – 1 atm Liquid STP – 298 K – 1 atm Dalton’s Law of Partial Pressures

o PT = Px + PY + Pz…o Pressures are collected volumes are NOTo nx/nT = Px/PT

Root Mean Square Speed - μ = Mm= kg/mole, R = 8.314 J/mole x K, T = K Diffusion/Effusion

o Diffusion – Particles moving from high to low concentrationo Effusion – A type of diffusion, but through a porous holeo Comparison of effusions into gases

Rate of Effusion x/Rate of Effusion Y = x - faster and smaller, y – slower and bigger

Vander Waals o P (V-nb) = nRT

B – Van der Waal’s constant (different for every atom, tells you size)o [Pobservation + a(n/V)2] x (V-nb) = nRT

a = Van der Waal’s constant, tells you physical attraction or charge Pobservation = real pressure

Triple Point Diagramo Shows what pressure or temperature you need to have solid, liquid, and gas. o Normal pressure = 1 atmo Supercritical fluid – Cannot be liquid or even gas it is just dense particles, like a

5th state of mattero Vapor pressure = Pressure that the vapor of that substance exerts as it vaporizeso Boiling Point – When vapor pressure of substance = atmospheric pressureo Line between solid and liquid states

Tilts to right solid is denser than liquid Tilts to left liquid is denser than solid

Heat of Vaporization – amount of heat required to boil 1 mole of liquid into gas (Joules)o Ln(Pvaporization) = -ΔHvaporization/R (1/T) + C (y-mx+b form)

Pvaporization = Pa R = 8.314 J/mole x K C = constant Steepest slope = greatest ΔHvaporization/R

Solid Stateso Crystalline Solids – repetitive lines

4 types Molecular - Ex. Ice

o Held together by weak physical bondso Weakest of the 4 solid types

o Has a low melting pointo Brittle, soft, does not conduct electricity, have no or little

charge Ionic

o Very strong, high melting points, brittle, hardo Don’t conduct solids but conduct liquids because in liquid

form the positive and negative atoms have space between them to allow electricity to flow

Metallic o Valence electrons move from one metal atom to anothero Electrostatic networko Metal atoms are strong because they’re electrons are

constantly in motion and there is more chargeo Malleable (not brittle), very low melting points and

hardness, conducts electricity Network Covalent – Ex. Diamonds

o Mostly stronger than Ionic (exception – Graphite)o Covalent bonds in many directionso High melting point, hard, typically non-conducting

o Amorphous Solids – not well defined 3-D unit structureo The strength is determined by molecular forces that hold the solid together

Physical or intermolecular bonds Chemical bonds stronger

o Crystal Lattice Collection of all the unit cells 3 types

Each corner is 1/8 of an atom in an unit cell Cubic unit cell

o Lattice points are at centero 1 unit cell is 1 atomo 52% packing efficiencyo Coordination number - 6

Number of nearest neighbors surrounding atom of interest

Body centered cubico Lattice point at center of cells and cornerso One atom at center

o 68% packing efficiencyo Coordination number - 8o 2 atoms in the unit cell

Face centered cubico Lattice points at the center each face and in the cornerso Coordination number – 12o 4 atoms in unit cello 74% packing efficiency

Atomic History Ben Franklin – two charges in world William Crookes – Cathode Ray Tube/Crooks Tube J.J. Thompson

o Used Cathode Ray Tube and had 2 charged plates above and below and shot a beam of particles through it and were attracted to positive plates (electrons).

o He concluded that magnetism and electricity were relatedo Theorized positive and negative particles are part of the atom – made Plum

Pudding model – 1st model of an atom Robert Millikan – Oil Drop Experiment – discovered charge and mass of electron Eugene Goldstein

o Used cathode ray tube with holes to discover positive particleso Discovered each element has different number of protons

Ernest Rutherford o Discovered 3 different types of radiationo Gold Foil Experiment – discovered that many protons are at center of atomo Made nuclear model

James Chadwicko Discovered nucleus = protons and neutrons

Unit 6 – Kinetics Kinetics: Rate at which a chemical process occurs Reaction Rate: Speed of Reaction Differential Rate Law: How the rate depends on amount of reactants Integrated Rate Law: How to calculate the amount left or time to reach a given amount Factors that affect reaction rate:

o Temperature – higher temperature is more collision and faster reactiono Concentration of reactants- more concentration is more reactedo Catalysts- changing mechanism or how the molecules react/ makes it faster and

easier for molecules to collide

Reaction Rateo A Bo Average Rate = –ΔA/ΔT or ΔB/ΔT

Measured in M (moles / L) Average speed is different from instantaneous speed Finding on graph- find slope Not linear because it’s not constant concentration

o Rate Law- each reaction has its own equation that gives its rate as a function of reactant concentrations

Instantaneous Rate You can find the rate on a graph by finding the tangential line to

the instantaneous point and the slope of this line (y/x) Ex. 2NO2 N2O4 Rate= -k [NO2]m

o m= rate order= has to be determined experimentally= not coefficient

o k= rate constant= relates time and concentration= giveno large k rapid reaction o small k slow reaction

method of initial rates- when looking just at initial concentrations Orders of Reaction

Tells exponents on the concentrations Overall order- the sum of the order of all of the reactants Order of reactants

o 2nd order= if rate is quadrupledo 1st order= if rate is doubledo 0 order= no change in rate/ no affect

Every k value has different units 3rd order reaction – M-2s-1

2nd order reaction – s-1

1st order reaction – M/s

The Collision Theory Modelo More Collision = Higher Temperatureo Must collide with correct orientationo Activation Energy: minimum amount of energy required for reactiono Reaction Coordinate Diagram – process showing energy changes

Exothermic Initial energy is greater than final energy To determine activation energy: (initial part of bump to top of

bump) Endothermic

Initial energy is lower than final energy To determine activation energy: (from bottom of graph to top of

bump) k = Ae-Ea/RT

A = pz P – fraction of molecules w/ proper orientation z – collision frequency

ln (k) = -Ea/RT + lnA Graph – lnk vs. 1/T , slope - -Ea/RT

Lnk2/k1 = Ea/R ( 1/T1 – 1/T2)o Mechanism

Sequence of smaller reactions that describe the actual process by which reactants become products

Molecularity of process tells how many molecules are involved in the process

Reaction mechanism tells us the rate law Rate equation is based on slower step Determining legitimacy:

Sum of smaller steps must give you overall reaction Must be derived from slower step

Ex: Overall reactions: 2NO + Br2 2NOBr Step 1: NO + Br2 NOBr2 (fast) (Forward – k1 = reverse – k -1)

Step 2: NOBr2 + NO NOBr (slow) (k2)o NOBr2 is imaginary because it cancels outo Rate = k1[NO] [Br2]o Rate = k-1 [NOBr2]o k1[NO] [Br2] = k-1 [NOBr2]o [NOBr2] = k1/k-1 [NO] [Br2]o Rate = k2[NOBr2][NO]o Rate = k2k1/k-1 [NO] [Br2] [NO]o Rate = k2k1/k-1 [NO]2 [Br2]

Unit 7- Equilibrium What is equilibrium?

o Rate reverse = rate forwardo Not an equal amount of products and reactantso Reactions are reversible

K= o For pressure or concentrationo Coefficient matters because it is the exponent

Finding equilibrium is important in determining efficiency Endothermic- create more product Favors side that has less

o If it has more products it favors reactantso If it has more reactants it favors products

Pressure and moles are directly related Kp= Kc(RT)Δn

o R= .00821o Δn= change in moles of the reaction

Pure solids and liquids are never included in the Kc or Kp equationso Solids and liquids don’t change concentrations

Q= reaction quotiento Values of concentration and pressure at a certain point not at equilibriumo Instantaneouso Q is less than K favors the right side of the reactiono Q is more than K favors the left side of the reaction

Ex: ICEo NO(g) + O3(g) <-----------------> NO2(g) + O2(g)

I .55 M .55 M Ø Ø

C -x -x +x +xE .55-x .55-x x xKc = 5.00 x 10-6

Kc =

5.00 x 10-6=

5.00 x 10-6= x = .00123

o Since the Kc value is less than 10-5 we can just disregard the x values in the bottom of the fraction

When combining two equations in order to find the K for the total reaction:o we have to multiply the K values togethero if we flip one of the reactions take reciprocal of the K value for that reactiono if we multiply by a number take the K value of that reaction to that exponent

(number you multiplied by) Le Chatelier’s Principle

o Lower the concentration or remove some of the substance It will shift to produce more of that substance

o Increase the concentration or add some of the substance it will shift to produce less of that substance

o increasing volume=decreasing the pressure and vice versao if you increase the pressure

then the system will shift so that the least number of gas molecules will form

because the more gas molecules there are the more collisions there are (which increases the pressure)

o if you decrease the pressure then the system will shift so that the most number of gas molecules will

form because the more gas molecules there are the more collisions

there are (which increases the pressure)o increase temperature

endothermic reaction will be favored because it will take in some of the excess heat

o decrease temperature exothermic reaction will be favored

because it will produce the heat that was lost Catalyst will increase speed but has no effect on K

Unit 8- Acids and Bases Acids

o Has H+ (Arrhenius Acid)o Donate a proton (H+) (Bronsted-Lowry Acid)o Has a low pHo Usually aqueouso Reacts with a baseo Reacts with OH-

o Electron pair acceptor(Lewis acid) Bases

o Has OH- or donate OH- (Arrhenius base)o Accepting a proton or H+ (Bronstend-Lowry Base)o Reacts with an acido High pHo Electron pair donor(Lewis Base)

Ex:o NH3 + H2PO4

- <-----> NH4+ + HPO4

-2 o Base acid conjugate acid conjugate baseo The base is accepting the protono The acid is donating the protono Bronsted-Lowry Acid and Base Reaction

Base always turns into acid and acid always turns into base Water can be an acid or base Strong acids/ bases break apart 100% so they do not have a Ka or Kb value Weak acids and bases don’t break apart 100% so they need Ka or Kb value which is still

concentration of products over reactantso If K is small then the acid or base is weako K measures how well the acid or base dissolves

pH= measure of H+ in the solutiono pH= -log (H+)o H+= 10-pH

o Same thing for OH- (pOH) pH scale

o 1-6 acid

o 7 H2Oo 8-14 base

pH + pOH = 14 Kw = Ka (Kb)

o Kw= 1x10-14

pKa= -log Ka pKb= -log Kb

pH= pKa + log Henderson-Hasselbach when the pKa is low then it is a strong acid and it breaks apart easily (same for base) % dissociation= how much of the acid or base dissolves

o Weak acids or usually 3 to 5% dissociation The more oxygens the greater the pull, the stronger the acid Hydrolysis: substance reacts with water the greater the charge on the metal ion, the stronger the acidity (transition metals

because they have open d sublevels) Buffers

o Buffer is a blockero Slows the reactiono Cushions blow as weak acid/ base dissolveso Must be weak acid or base and its conjugate solid (paired ionic compound)o Ex:

HC3H5O3(aq) <-------> H+ + C3H5O3-

I .12 M Ø .1 MC -x +x +xE .12 – x x .1 + x

Ka=x= 1.65 x 10-4

pH= 3.78

Second method= pH= pKa + logo Look at buffer lab for more examples with more complex problems

Kspo For things that do not dissolve well in watero Not for things that dissolve 100%o It is to see how much will redissolveo Do normal ice but disregard reactant because it is a solid

o Ex: PbCl2 ----> Pb+2 + 2Cl-

I Ø ØC +x +2xE x 2xKsp= [x][2x]2

Ksp= x[4x2]Ksp=2.5 x 10-5

X= .0184

Unit 9- Thermodynamics Spontaneous process= occurs without outside intervention; naturally without much

effort Enthalpy = ΔH= change in Heat for a reaction

o Endothermic= positiveo Exothermic= negative

Entropy= state of disorder or randomness (S)o This increases with heat

ΔSuniv = ΔSsys + ΔSsurr

o ΔSuniv= positive Disorder will increase USUALLY spontaneous

o ΔSuniv= negative Disorder will decrease USUALLY non spontaneous

o Sign of ΔSsurr is dependent on the heat flow of the reactiono Magnitude of ΔSsurr depends on temperature

High temperature= more randomness Low temperature= less randomness

o ΔSsurr= -ΔH/T J/K Negative because when we have a negative ΔH then the negatives cancel

out & the ΔSsurr becomes positive showing that it is spontaneouso States

Going from gas to liquid to solid is decreasing disorder (ΔS = negative) Going from solid to liquid to gas is increasing disorder (ΔS = positive) Decreasing the moles of gas is decreasing disorder (ΔS = negative)

Gibbs’ Free Energy

o ΔG˚= ΔH˚ - TΔS˚ use signs in this equation to predict spontaneity This determines spontaneity best- because it includes everything ΔG= negative= spontaneous ΔG=Ø= equilibrium ΔG= positive= non-spontaneous ΔH= - , ΔS= + then ΔG= negative spontaneous ΔH= + , ΔS= - then ΔG= positive non-spontaneous

o ΔSuniv= -ΔG/T ΔS˚= ∑ΔS˚products - ∑ΔS˚reactants

ΔH˚= ∑ΔH˚products - ∑ΔH˚reactants

ΔG˚= ∑ΔG˚products - ∑ΔG˚reactants

G = G˚ + RT ln Po G = new

G = ΔG˚ + RT ln Q ΔG˚= -RT lnK

o at standard state

Unit 10- Electrochemistry interchange between electrical and chemical energy Galvanic/ Voltaic- spontaneous generates current spontaneously

o For example batteries Two kinds

Wet Dry (most common)

Electrolytic Cell- not spontaneous outside current must be appliedo Used to generate elements from compounds

Anode= electrons move away from it Cathode = electrons move toward it Most reactions are single replacement reactions Salt bridge

o Positive ion of chemical in salt bridge Is going to flow into container with cathode to balance out the negative

chargeo Negative ion of chemical in salt bridge

Is going to flow into container with anode to balance out positive charge

Oxidation/ Reduction Reactionso Oxidizing agent gains electrons at cathodeo Reducing agent loses electrons at anodeo Oxidation = lose electrons at anodeo Reduction= gain electrons at cathode

Ε= electric potential= voltage potential Current= measurement how many e- you get during a period of time Voltage= force put on the electricity Voltage will not change as we manipulate equations When trying to figure out Etotal

o Flip reaction that is not the most spontaneouso Keep reaction that is most spontaneous

More spontaneous= gaining electrons Ex: Cu+2 + 2e- Cu(s) Eo = -0.34 V

Zn (s) Zn2+ + 2e- Eo = 0.76 VEototal= 1.11 V

To balance half reactions use Jauss’ short cut:The Half-Reaction Method of Balancing

Oxidation/Reduction ReactionsLet’s use the following reaction as a sample problem:Potassium permanganate solution is added to an acidic solution of oxalic acid.

1. Write an unbalanced ionic equation for the reaction – only show reactive species, and not spectator ions.

H2C2O4 + MnO4-1 → CO2 + Mn+2

2. Assign oxidation numbers to the elements that are being oxidized and reduced, and what new oxidation number they become - these can be difficult to identify!

C = +3 → +4 Mn = +7 → +2C2O4

-2 → CO2 MnO4-1 → Mn+2

3. Divide the reaction into oxidation and reduction half-reactions and balance these half-reactions one at a time.

a. First balance elements that aren’t hydrogen and oxygen. b. Balance oxygens by adding waters to the appropriate side of the equation.c. Balance hydrogens by adding H+1 to the appropriate side of the equation.d. Balance electric charge by adding electrons (e-) to the appropriate side of the equation.e. If the reaction is in basic solution, add OH-1 molecules to both sides of the equation,

equal to the number of H+1 ions in the equation. Simplify the equation by combining the H=1 ions and OH-1 ions into water, and reduce the equation to its simplest terms.OXIDATION REACTIONH2C2O4 → CO2

H2C2O4 → 2 CO2 to balance the carbonH2C2O4 → 2 CO2 + 2e- carbon went from +3 to +4H2C2O4 → 2 CO2 + 2e- + 2H+1 both sides should have the

same charge, and the hydrogens and oxygens should be balanced

REDUCTION REACTIONMnO4

-1 → Mn+2

MnO4-1 + 5e- → Mn+2 manganese went from +7 to

+2MnO4

-1 + 5e- + 8H+1 → Mn+2 both sides should have the same charge – adding 8H+1 to both sides gives each side a +2 charge

MnO4-1 + 5e- + 8H+1 → Mn+2 + 4 H2O Balance oxygen and hydrogen by

adding 4 waters to the right

4. Combine these half-reactions so that electrons are neither created nor destroyed – meaning, you have the same number of electrons on the right and left hand sides.2 x (MnO4

-1 + 8 H+1 + 5 e- → Mn+2 + 4 H2O) the top reaction needs to be multiplied by 2, the bottom

5 x (H2C2O4 → 2 CO2 + 2 e- + 2 H+1) by 5, to balance the transfer of electrons

2 MnO4-1 + 16 H+1 + 5 H2C2O4 → 10 CO2 + 2 Mn+2 + 8 H2O + 10 H+1

5. Balance the remainder of the equation by inspection, if necessary.2 MnO4

-1 + 16 H+1 + 5 H2C2O4 → 10 CO2 + 2 Mn+2 + 8 H2O + 10 H+1 In this case, we simply reduce the number of H+1 on each side!

2 MnO4-1 + 6 H+1 + 5 H2C2O4 → 10 CO2 + 2 Mn+2 + 8 H2O

+

E=EMF= potential difference= voltage (V)= work/ charge= W/q= J/Co Electromotive forceo Voltage = amount of work on an electron

E= - W/q Amp= C/s Voltage = positive= spontaneous W= qE q= nF

o n= # of moles of electronso F= Faraday’s constant= 96485 C/ 1 mol of e-

Wmax= ΔG ΔG= -nFE Two things that determine spontaneity of a battery

o Determined by concentration of elementso Natural tendency

V= work/q Nernst Equations

o E˚= (RT/nF) (ln K)o E= E˚-(RT/nF) (lnQ)

Q= not at equilibriumo E = E˚ - (0.0591/ n) log Q

At 298 K In joules

ΔG changes from balancing but the voltage does not

Solubility Rules1. All alkali metal compounds are soluble in water. All nitrates and nitrites are soluble in

water. All chlorates and perchlorates are soluble. All acetates are soluble in water, except when they contain Ag+1. All (NH4)+1 compounds are soluble.

2. All chlorides, bromides, and iodides are soluble in water, except those containing Ag+, Hg2

+2, Hg+2, and Pb+2.3. All fluorides are soluble in water, except those that contain group II metals and Pb+2 and

Fe+3.4. All sulfates and sulfites are soluble in water, except those containing Sr+2, Ba+2, Hg2 +2,

Ag+1, Ca+2, and Pb+2.5. All carbonates, phosphates, and chromates are insoluble in water, except those of the

alkali metals and (NH4)+1. CaCrO4 and SrCrO4 are exceptions.

6. All hydroxides are insoluble, except those of the alkali metals, and ones containing NH4

+1, Sr+2, Ca+2, and Ba+2. 7. All sulfides are insoluble, except those of group I and group II metals and NH4

+1. 8. All oxides are insoluble, except those of group I and NH4

+1.