Solubility Rules Notes Cheat Sheet

2 Solubility Rules 2.1 Soluble • Nitrates N O3−1 - All nitrates are soluble • Chlorates ClO3−1 - All chlorates are soluble • Alkali metal Cations and Ammonium cation compounds NH4+1 are all soluble • Chlorides, Bromides, and Iodides are all soluble EXCEPT Ag+1 , P b+2, and Hg +2 • Acetates - All are soluble except Ag+ • Sulfates - All are soluble except Ba+2 , P b+2, H g+2 , Ca+2, Ag+1 , and Sr+2 2.2 Insoluble • Carbonates CO3−2 - all carbonates are insoluble except alkali metals and ammonium compounds • Chromates CrO4−2 - 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 O4−3 all are insoluble except alkali metals and ammonium • Sulfites SO3−2 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 (H2 SO4 ) should be written as H + + H SO4−1 • The strong acids are: HC L, HBR, HI, H N O3 , HClO4 , and H2 SO4 • 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 → HC2 H3 O2 3 Periodic Table of Elements 4 Poly Atomic Naming • Zinc Zn+2 • Silver Ag +1 • Ammonium N H4+1 • Hydroxide OH −1 • Cyanide CN −1 • Nitrate N O3−1 • Acetate C2 H3 O2−1 • Chlorate ClO3−1 • Bromate BrO3−1 • Iodate IO3−1 • Manganate M nO3−1 • Sulfate SO4−2 • Bisulfate (Hydrogen Sulfate) HSO4−1 • Carbonate CO3−2 • Bicarbonate (Hydrogen Carbonate) HCO3−1 • Selenate SeO4−2 • Biselenate (Hydrogen Selenate)HSeO4−1 • Oxalate C2 O4−2 • Phosphate P O4−3 • Hydrogen Phosphate HP O4−2 • Dihydrogen Phosphate H2 P O4−1 • Chromate CrO4−2 Per Ate Per Ic +1 Oxygen 5 Ate Ic Most Common Ion Ite Ous -1 Oxygen Common Units, Constants and Charges 5.1 Fundamental Constants • Avogadros Number (N ) ⊲ 6.02214199 ∗ 1023mol−1 • Plancks Constant (h) ⊲ 6.62606876 ∗ 10−34J ∗ s • Speed of Light (c) ⊲ 2.99792458 ∗ 108 m/s Hypo Ite Hypo Ous -2 Oxygen 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 charge mass 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 Cl2 O Dichlorine Monoxide 7.2 Ionic • Finding Charge: N a3? Cl2+1 N acdClba (a∗b) c 7.3 7.3.1 =d Acids Polyatomic • Per...ate → Per...ic acid ⊲ HN O4 → pernitric acid • -ate → ic acid • -ite → ous acid ⊲ H + N O3 → HN O3 (Nitric Acid) ⊲ HN O2 → 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 → HCO3 8 Cations • Which cation forms a white precipate with H CL? ⊲ Ag+ (reversed proves Cl− ). • What color is a typical Manganese solution? ⊲ Pink/light purple. The precipate is dark black. • How would you test for Al+ and what would it look like? ⊲ Add Aluminom, it will make a precipate red and leave the solution clear. • Which cation forms a gel like precipate? ⊲ Aluminum. • Which cation turns deep red with K SCN ? ⊲ Iron. • How do you confirm the presence of zinc and what color is it? ⊲ Add acid, then ammonia, which results in a white/bluish precipate. • What cation turns a deep blue with ammonia? ⊲ Copper. 9 Reaction Type 9.1 Combination (Synthesis) When two or more chemicals react to form one product Example 2M g + O2 → 2M gO 2N a + S → N a2 S • Metal + Non-Metal → Metal Nonmetal (Binary Salt) • Metal Oxide + Water → Metal Hydroxide ⊲ CaO + H2 O → Ca(OH )2 ⊲ K2 O + H2 O → 2KOH • Metal Oxide + CO2 → Metal Carbonate ⊲ N a2 O + CO2 → N a2 CO3 • Nonmetallic Oxides + water → Acids (nonmetal oxides retains its oxide number) ⊲ N a2 O + SO3 → N a2 SO4 9.2 Decomposition When one chemical decomposes into 2 or more Example ∆ 2Ag2 O → 4Ag + O2 ∆ = Heat ∆ • Metal Carbonate → Metal Oxide + CO2 ∆ ⊲ CaCO3 → CaO + CO2 ∆ • Metal Hydroxide → Metal Oxide + H2 O ∆ ⊲ M g(OH )2 → M gO + H2 O ∆ • Metal Nonmetal → Metal + Nonmetal (diatomic in nature) ∆ ⊲ 2N aCl → 2N a + Cl2 ∆ • Metal Chlorates → Metal Chlorides + O2 ∆ ⊲ F e(ClO3 )2 → F eCl3 + O2 9.2.1 Special Binary Salt Splits These binary salts split into different elements (N H4 )2 CO3 → N H3 + H2 O + CO2 H2 SO3 → H2 O + SO2 H2 CO3 → H2 O + CO2 N H4 OH → N H3 + H2 O H2 O2 → H2 O + O2 9.3 Combustion Hydrocarbon + O2 → CO2 + H2 O ....⇓.... Cx Hy → double x (multiply by 2) then add 2 • C1 : meth • C2 : eth • C3 : pro • C4 : bu • C5 : pent • C6 : hex • C7 : hept • C8 : oct • C9 : non • C10: dec 10 Blackbody Radiation When an object is heated it will emmit radiant energy E = hν • E = Energy • h = Max Plancks constant (6.626 ∗ 10−34J ∗ 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: 1. Only orbits of certain radii, corresponding to certain definate energies are permitted for the electron in a hydrogen atom. 2. 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. 3. Energy is emmitted or absorbed by the e− only as the e− changes from one allowed energy state to another. 4. Flawed theory because it only works for hydrogen 11.1 Energy Level Formula En = (−2.18 ∗ 10−18J)( n12 ) • E1 : −2.18 ∗ 10−18J • E2 : −5.45 ∗ 10−19J • E3 : −2.42 ∗ 10−19J • E4 : −1.36 ∗ 10−19J • E5 : −8.72 ∗ 10−20J • E6 : −6.056 ∗ 10−20J • E∞ : 0 11.1.1 Energy Change during Level Jumps ∆E = EF − E0 • n = 3 → 2 | −3.03 ∗ 10−19J • n = 4 → 2 | −4.09 ∗ 10−19J • n = 5 → 2 | −4.578 ∗ 10−19J • n = 6 → 2 | −4.844 ∗ 10−19J 12 12.1 Wavelength De Broglie Formulas λ= h mv or λ= h p • λ = Wavelength • h = Plancks Constant (6.626 ∗ 10−34J ∗ s) • m = Mass of particle in Kg • v = Velocity of particle ( meters ) second • p = Momentum Example m = 9.11 ∗ 10−28g v = 5.97 ∗ 106 m/s −10 ∗10−34 J∗s m λ = (9.11∗106.626 −31 Kg )(5.97∗106 m/s) = 1.22 ∗ 10 13 Quantum Values 1. Principle Quantum number - (n) n = 1 (lowest) n = ∞ (at 8 or 9) Follows Bohrs En = (−2.18 ∗ 10−18 J)( n12 ) 2. Azimuthal Quantum number - (l ) l =n-1 if... • l = 0 → S shape • l = 1 → P shape • l = 2 → D shape • l = 3 → F shape Example n=3 l =2 ⇓ 3d 3. Magnetic Quantum number (orbital) - (ml ) -l and l including zero m0 = 0 m1 = −1, 0, 1 m2 = −2, −1, 0, 1, 2 4. Spin magnetic quantum number - (ms ) + 12 or - 21 13.1 n 1 2 3 4 Quantum Value Table Possible l values 0 0 1 0 1 2 0 1 2 3 13.2 Subshell 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f ml values 0 0 -1,0,1 0 -1,0,1 -2,-1,0,1,2 0 -1,0,1 -2,-1,0,1,2 -3,-2,-1,0,1,2,3 # of orbitals in subshell 1 1 3 1 3 5 1 3 5 7 Special cases • Chromium has 6 half-filled orbitals • Copper has one half-filled orbital and 5 filled orbitals total # of orbitals in shell 1 4 9 16 - e− in shell 2 8 18 32 - 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: 1S2 , 2S2 , 2P6 N e: 1S2 , 2S2 , 2P6 15 Nuclear Chemistry Nuclear Chemistry involves changes in the nucleus of an atom. Normal Reactions involve electron transfer Reaction affected by factors such as pH, temp, pressure, [], etc. Reactions involve relatively small energy: 400 kJ-1500kJ 15.1 Nuclear Reactions involve decay of nucleus i.e. transforming one element into another Affected by the type of decay and the halflife of what is decaying Reactions deal with huge amounts of energy Isotopes Isotopes: Atoms of the same element that have a different number of neutrons X −A A ZX A • X = Element Symbol • A = Atomic Mass • Z = Atomic Number X 15.2 15.2.1 Radiation Alpha Radiation When a big nucleus ejects a H e+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. 15.2.4 Positron Radiation When a positively charged nucleus emits its p+ leaving only the n◦ . 15.2.5 Electron Capture When an electron in orbit falls into the nucleus (positively charged) and makes it neutrally charged. 15.3 15.3.1 Nuclear Equations Radiation Table Neutron: Proton: Electron: Positron: Alpha Particle: Beta Particle: 1 0n 1 + 1p 0 − −1 e 0 − 1e 4 4 2 He or 2 α 0e− or 0 −1 −1 β Example Alpha 238 234 4 92 U →90 T h +2 He Beta 131 131 0 − 53 I → 54 Xe+ −1 e 1 0n →11 p+ −10 e− Positron 11 11 0 − 6 C →5 B+ 1 e 1 p →1 n+0 e− 1 0 1 Electron Capture 81 0 − 81 37 Rb+ −1 e →36 Kr 1 p+ 0 e− →1 n 1 −1 0 Positron-Electron Collision (Gamma) 0 0 − 0 γ e+ e → 0 1 −1 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+ ) (n◦ ) 2 2 • If (p+ ) and (n◦ ) even → likely stable • If either is odd → could go either way • If (p+ ) and (n◦ ) odd → likely unstable 8 8 20 20 28 28 50 50 82 82 126 15.4.4 Half-Life The time it takes 1 2 the amount of a substance to decay. Example 5g of nuclide life of 15 years How much of the original nuclide remains after 45 years? 5 ⇓ (15 years) 2.5 ⇓ (30 years) 1.25 ⇓ (45 years) 0.625g 1 2 16 16.1 Ionization and Affinity 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 + H2 • Metal + O2 (Li or any non-Alkali metal) → Metal Oxide • K + O2 (Any other Alkali metal) → Metal Peroxide (O2−1) ⊲ K + O2 → KO2 • Metal Oxide + H2 O → Metal Hydroxide ⊲ N a2 O + H2 O → N aOH • Metal Oxide + Acid → Salt + H2 O ⊲ N a2 O + HCL → N aCl + H2 O Nonmetal Oxides = Acidic • Nonmetal Oxide + H2 O → Acid ⊲ CO2 + H2 O → H2 CO3 ⊲ SO2 + H2 O → H2 SO3 ⊲ P4 O10 + H2 O → H3 P O4 • Nonmetal Oxide + Base → Salt + H2 O ⊲ CO2 + N aOH → N a2 CO3 + H2 O 18 Chemical Bonds When 2 or more atoms are strongly attached (attracted) to each other.