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Significant Figures Name methane ethane propane butane pentane hexane heptane octane nonane decane ….ene …yne Ruler A – 4.40 cm Ruler B – 4.4 cm Add 1 more decimal than the smallest measurement can measure. Type of 0 placeholder placeholder trapped precise 0.00003 3,0000 3,0003 3.0000 Significant? no no yes yes Formula CH4 C2H6 C3H8 C4H10 C5H12 C6H14 C7H16 C8H18 C9H20 C10H22 CnH2n CnH2n-2 Covalent prefixes Number Prefix Number Prefix Adding sig figs – do not count the # of sig figs! Count the # of decimal places. 1 mono- 6 hexa- 2 di- 7 hepta- % error 3 tri- 8 octa- 4 tetra- 9 nona- 5 penta- 10 deca- = I valueacc - valueexpI valueacc x 100% 106 m = 1 m 109 nm = 1 m 1012pm = 1 m micro nano n pico p Substance Liq. Water Gas at STP Density 1.00 g/mL mass of 1 mol/22.4L Average atomic mass = % mass + % mass +….. (isotope 1) (isotope 2) x x 1 mole = 6.022 x 1023 particles. (Avogadro's number) 1 12C atom = 12.00 amu (atomic mass) 1 mole 12C atom = 12.00 g (molar mass) MnO4- permanganate S2O3-2 thiosulfate CrO4-2 chromate Cr2O7-2 dichromate O2-2 peroxide ClO- hypochlorite ClO2ClO3ClO4C2O4-2 chlorite chlorate perchlorate oxalate Fe(CN)6-4 hexacyanoferrate(II)/ ferrocyanide A = mass number Z = atomic number A = number of protons + neutrons Z = number of protons A – Z = # of neutrons Isotopes= atoms with same number of protons but different numbers of neutrons Ions= electrons vary Isomer – different arrangements if atoms within a molecules PROTONS, Atomic # define the element! Fe(CN)6-3 hexacyanoferrate(III)/ ferricyanide BO3-3 borate Hg2+2 mercury (I) ion Molecular formula = the way a molecule really appears in nature Empirical formula = smallest whole number ratio of elements ex: = whole number always! = Glucose has a molecular formula (MF) of C6H12O6 Ba(NO3)2 Ba(NO2)2 + O2 8. Glucose’s empirical formula (EF) is CH2O Glucose’s molar mass (MM) is 180 g/mol Glucose’s empirical mass (EM) is 30 g/mol Diatomic molecules - BrNClHOF Other Molecular Elements o Phosphorus--one form P4 o Sulfur—often S8 o Carbon--4 allotropes, diamond and graphite, buckyballs, graphene, all covalent networks Other compounds to memorize 1. Hydrogen Peroxide – H2O2 2. Ammonia – NH3 REACTIONS TO MEMORIZE: 1. Synthesis reactions of metal oxides with water: . ex: Na2O + H2O 2NaOH CaO +H2O Ca(OH)2 2. Synthesis reactions of nonmetal oxides with water: ex: SO3 + H2O H2SO4 N2O5 + H2O 2HNO3 3. 4. Decomposition of Binary Compounds ex: 2H2O 2H2 + O2 MgS Mg + S H2O2 H2O + O2 Decomposition of Metal Carbonates: ex: CaCO3 CaO + CO2 Na2CO3 Na2O + CO2 5. Decomposition of Metal Hydroxides: ex: Ca(OH)2 CaO + H2O 2NaOH Na2O + H2O 6. Decomposition of Metal Chlorates: ex: 2KClO3 2KCl + 3O2 Mg(ClO3)2 MgCl2 + 3O2 7. Decomposition of Metal Nitrates: Decomposition of acids: certain acids can decompose to form nonmetal oxides and water. ex: H2CO3 CO2 + H2O H2SO3 SO2 + H2O 9. EF X 6 = MF; EM X 6 = MM for glucose! Solving EF problems: 1. If given percentage, drop % and use g. (assume 100 g) 2. if given mass, just leave it! 3. Convert g of each element to moles to get mole ratio. 4. Divide by smallest to get lowest whole number ratio. 5. if dividing by smallest does not give whole number ratio, then multiply all by a whole number: .2,.4,.6,.8 x 5 .25, .75 x 4 .33, .67 x 3 2NaNO3 2NaNO2 + O2 10. 11. Combustion: elements in a compound separate and combine with oxygen to make most reasonable cmpds. ex: ABC + O2 AO + BO + CO2 C2H5OH + 3 O2 2 CO2 + 3 H2O CS2 + 3 O2 CO2 + 2SO2 Single Replacement: A metal replaces a metal or a nonmetal replaces a nonmetal IF the element doing the replacing is higher on the activity series. Hydrogen replaces/is replaced by metals, though. Ex: MgCl2 + 2Na 2 NaCl + Mg Cl2 + 2 NaI 2 NaCl + I2 2 HCl + Mg MgCl2 + H2 Double replacement: Only works if 1 of following is produced: 1. Precipitate 2. Gas 3. Water If the following is made NH4OH H2SO3 H2CO3 It decomposes into… H2O + NH3 H2O + SO2 H2O + CO2 All compounds containing group 1 metals, acetates, nitrates, or ammonium are soluble. BaSO4, PbI2, and AgCl are common insoluble substances! There is no such thing as an oxide ion in solution! It will quickly become a hydroxide ion! Preparing Solutions: 1. Always use volumetric flask. 2. Add solute or concentrated solution. 3. Fill to the line with deionized water. 4. Shake. McVc = MdVd Mc = molarity of concentrated solution (used to make new soln) Vc = molarity of concentrated solution (used to make new soln) Md = concentration of solution produced Vd = volume of solution produced = vol. of concentrated soln + volume of water! Leo says Ger!!!! Losing Electrons is Oxidation Gaining Electrons is Reduction! EX: 2 AgNO3 + Zn Zn(NO3)2 + 2 Ag 2 Ag+ + Zn Zn+2 + 2 Ag Silver is reduced; zinc is oxidized 2 moles of electrons are transferred (not 4!) Half reactions: 2 Ag+ + 2 e- 2 Ag Zn Zn+2 + 2e- Amount of energy to change temp of 1oC of H2O = 1 cal 4.184 J = 1 cal 1000 cal = 1 kcal = 1 Cal Half reactions are often easy! If not single replacement, use…. All Oranges Have Citrus Energy! 200 g water at 70oC 1. Separate into half reactions. Leave out spectators. 2. A = balance atoms (except H and O) 3. O = balance oxygens (add H2O for acid rxns; add 2 x OH- for bases) H = balance hydrogens (add H+ for acid rxns; add H2O for bases) 5. C = balance charge (by adding electrons) 6. E = balance electrons (by multiplying half reactions) 4. Assigning Oxidation Numbers: 1. 2. In free elements, each atom has an oxidation number of 0. For ions consisting of a single atom, the oxidation number is equal to the charge on the ion. 3. For binary ionic compounds, assign oxidation numbers (charges) as you always have! For covalent molecules, or polyatomic ions, assign oxidation numbers in the order shown below. F-O-H, closest to F, farthest from F 4. 5. 6. 7. Fluorine is always assigned first an oxidation number of –1 when in a compound. Next, assign oxygen. The oxidation number of oxygen is -2 in most compounds. Exceptions: Oxygen in peroxides is -1 Ex. H2O2 Na2O2 Then, assign hydrogen. The oxidation number of hydrogen is usually +1, but may be –1 when combined with a metal. For example: H in hydrides is -1 Ex. NaH For compounds in which both atoms cannot have the oxidation number which is equal to the charge the element commonly has, one closest to fluorine “wins”. HEAT Total amount of kinetic energy of a sample A measure of average kinetic energy of a sample q Joules T o SPECIFIC HEAT The amount of energy required to change the temp of 1 g of substance by 1oC C J/goC cal/goC J/moloC HEAT CAPACITY The amount of energy required to change the temp of a given substance by 1oC none J/oC TEMPERATURE 1 mL H2O = 1 g H2O Enthalpy = heat C, K cal/oC 100 g water at 70oC Same T Same KEavg of particles Same speed of particles Same specific heat of water More heat energy (q) Less heat energy (q) Heat of formation ∆Hf Heat of fusion ∆Hfus Heat of vaporization ∆Hvap Heat of reaction ∆Hrxn Heat of combustion ∆Hcomb Ionization energy IE Lattice Energy U Standard entropy ∆S Gibbs Free Energy of formation ∆Gf Gibbs free energy of reaction ∆Grxn Energy required to form an element from its standard state Energy required to melt a substance or RELEASED when frozen Energy required to boil a substance or RELEASED when condensed Energy required/released during a reaction (per mole) Energy required/released during a comb. reaction (per mole) Energy required to remove 1 electron from an atom Energy required to break 1 mole of ions in a crystal lattice into gaseous ions Measure of disorder compared to that of a solid crystal at 0K Measure of spontaneity; ability to do work on surroundings as compared to elements in standard state Measure of spontaneity; ability to do work on surroundings Ways to Calculate the Enthalpy of reaction, ∆Hrxn: 1. ∆Hf (products) - ∆Hf (reactants) 2. bonds broken – bonds made 3. sum of ∆Hrxn for reactions that add up to new reaction 4. measure amount of energy gained/lost by water. Divide by moles of 1 reactant reacted. Person Dalton Experiment - Mendeleev Thomson Cathode rays Millikan Rutherford Einstein Planck Bohr Oil drop experiment Gold foil experiment Photoelectric Effect Line emission spectrum of (excited) hydrogen gas De Broglie Schrodinger Heisenberg - Discovered/Proposed Atomic Theory – 5 postulates, incorrect about all atoms of a given element identical and atoms are indivisible Designed Periodic Table; left holes for 3 missing elements and predicted their properties All matter contains electrons (and therefore protons); determined mass:charge of electron; Proposed plum pudding model of atom (p+/e- spread throughout) Determined exact mass and charge of electron Discovered positively charged, very dense nucleus photons Calculated “size” of photon – Planck’s constant Existence of energy levels/ quantized energy states of electrons De Broglie equation; wavelength of any moving object Schrodinger’s equation; calculates probability of finding electron in a given region (orbital!) within an atom by treating electron as probability wave function Uncertainty Principle; the greater the precision in measuring a small object’s location, the greater the uncertainty in measuring its velocity and vice versa (can’t know an electron’s location and velocity simultaneously) Excited electrons have gained energy and jumped to a higher energy level. They possess more energy. They fall back down to a lower energy state and must release energy in the form of 1 quantum/photon, E = hν Only certain sized photons (lines of frequency/wavelength) are emitted so each element has its own distinct line emission spectrum. This is due to the existence of quantized energy states in atoms. Pauli Exclusion Principle Aufbau Rule Hund’s Rule (Bus Rule!) Sublevel No 2 electrons in the same atom can have the same 4 quantum numbers; thus no 2 electrons can be in the same energy level/sublevel/orbital AND have same spin; 2 electrons in same orbital must have opposite spins Electrons fill up orbitals from lowest energy to highest energy (this may not be in numerical order! See aufbau box below) 2 is higher than 1; d is higher than p, etc…. If 2 equal energy orbitals are available, electrons each go to separate orbitals (with same spin) before pairing up 2p: ____ _____ _____ orbitals within Shape of orbitals Picture s s Spherical # of electrons in an orbital 2 p px, py, pz Dumbbell shaped 2 6 d dxy, dxz, dyz, dz2, dx2 –y2 4 lobed 2 10 f g 7 different f’s 9 different g’s 8 lobed/too complex !!!!! 2 2 14 18 Don’t even try AAAAH! # of electrons in sublevel 2 Quantum numbers describe where a given electron is in an atom. 1st quantum # Energy level 1, 2, 3, 4…. nd 2 quantum # sublevel s, p, d, f 3rd quantum # Orbital/orientation px, py, pz,for example 4th quantum # spin Aufbau box: (add arrows) Which one has higher Coulombic attraction? Why? 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 5g 6s 6p 6d 6f 6g 6h 7s 7p 7d 7f 7g 7h 8s 8p 8d……. 7i Example of orbital notation: C: ____ ____ _____ _____ ______ 1s 2s 2p Examples of electron configuration: Mg: (you do it!) K: 1s22s22p63s23p64s1 Mn: 1s22s22p63s23p64s23d5 Ag: 1s22s22p63s23p64s23d104p65s14d10 Pb: (you do it!) # unpaired electrons = #electrons alone in orbital C has 2 unpaired electrons K has 1, Ag has 1, Pb has 2, Mn has 5, Mg has 0 Unpaired electrons make element paramagnetic. Drawing Lewis Structures for COVALENT compounds: 1. The atom with lowest EN is the central atom. 2. Count the total number of valence electrons. MAGIC NUMBER! (adjust for ions) 3. Place one sigma bond (a pair of electrons) between each pair of bonded atoms. 4. Subtract the total number of valence electrons used for bonds from the MAGIC NUMBER. 5. Place lone pairs about each terminal atom except for hydrogen atoms. Subtract the number of lone pairs from the sum. 6. 7. If e- are still left at this point, assign them to the central atom. If the central atom is from the third or a higher period, it can accommodate more than four pairs of electrons. If the central atom is not yet surrounded by four electron pairs, convert one or more terminal atom lone pairs to pi bond pairs. IONIC Metal + nonmetal Metal +polyatomic ion Poly ion + nonmetal Poly ion + poly ion Strongest type of bond Electrostatic attraction between oppositely charged ions Electrons transferred, creating ions Forms crystal lattice And thus crystals Formula units Higher MP, Higher BP (must break bond to change phase) Lattice Energy Higher charge, smaller size= stronger bond Large EN difference If soluble, forms electrolytes when dissolved Electrolytic when molten, not solid High % ionic character COVALENT Nonmetal + nonmetal Within poly ion Weaker bond Electrostatic attraction between nuclei and shared electrons Electrons shared – NO IONS! A few form covalent networks (SiO2, , most form individual molecules molecules Lower MP, lower BP (only break IMFS NOT bonds to change phase) Bond energy Higher bond order, smaller atom = stronger bond Forms polar and nonpolar bonds Medium (polar) or Small (nonpolar) EN difference Nonelectrolytes, may dissolve but won’t ionize! Nonelectrolyte Low % ionic character Show resonance! Total number of sigma bonds and lone pairs on central atom Shape Atoms bonded to central atom (# of sigma bonds) Lone pairs on central atom 2 Linear 2 0 sp yes 3 Triangular Planar 3 0 sp2 yes 3 Bent 2 1 sp2 no 4 Tetrahedral 4 0 sp3 yes 4 Triangular pyramidal 3 1 sp3 no 4 Bent 2 2 sp3 no 5 Trigonal bipyramidal 5 0 dsp3 or sp3d yes 5 Unsymmetrical tetrahedron 4 1 dsp3 sp3d no Molecular Shape Picture Hybridization Symmetrical in 3d space? And bond angles (see saw) 5 T-shaped 3 2 dsp3 sp3d no 5 Linear 2 3 dsp3 sp3d yes 6 Octahedral 6 0 d2sp3 sp3d2 yes 6 Square pyramidal 5 1 d2sp3 sp3d2 no 6 Square planar 4 2 d2sp3 sp3d2 yes Example of a Lewis structure that is this shape. (Yes, draw the Lewis structure!) TYPE USES TO DETERMINE/Produce HOW Mass Spectroscopy Magnets Photoelectron Spectroscopy Beam of Light on Metal Distinguishes isotopes (separates by mass), can be used to determine % abundance of each isotope Binding Energy of electrons in different sublevels within a given atom Atomic absorption spectroscopy or Atomic Emission Spectroscopy Electrical or Heat Energy added to excite electrons The magnet bends heavier isotopes more than lighter ones “Light” beam (could be other form of ER) gives energy to electrons and removes them from the atom- this energy is measured…takes more energy to remove innermost electrons than outermost electrons Energy excites electrons and they absorb a specific wavelength, then when they fall back down, they release a specific wavelength Nucleus vibrational states altered by radio waves, gives “signature” of element Line Emission spectrum (or line absorption spectrum) which identifies the element NMR Spectroscopy (nuclear magnetic resonance) Radio waves Identity of a substance (types of (LOW Energy atoms present) right?) added to “excite” nucleus Infrared or Visible or UV IR or Visible Identity of a substance/molecule, Electrons in bonds also have spectroscopy (example: light or UV light (types of bonds present) OR quantized energies and these Using spectroscopes) to excite concentration of substance (when energies can be used to electrons in we used spectrophotometer) identify a molecule bonds Mass Spectrometry Data: …each bar represents the mass and abundance of a given isotope. What element? _______ PES data: (photoelectron spectroscopy)…each bar represents the abundance AND energy to remove each of the electrons from an atom (binding energy!). The outermost electrons take the least amount of energy to remove. 1s 3s 2p 2s When bond is formed, 2 atoms rest at a distance of minimum potential energy. This gives the typical bond length of the molecule. Boiling point – temp at which v.p equals atmospheric pressure. Normal b.p. –temp at which v.p equals 1 atm. ---------------------------------------------------------∆E = q + w ∆E = internal energy q = heat added to the system w = work done on the system So, for example if heat is added to the system AND work is done on the system, then the internal energy of the system increases and has a positive sign. If a reaction is exothermic and it expands in volume, work is done on the surroundings and so work is negative, and heat is negative, and the internal energy decreases and change in energy has a negative sign. Top of hump = activation complex/transition state/highenergy intermediate --------------------------------------------------------- Taller the “hump”, slower the reaction Activation energy is used to break bonds, create an “effective” collision Effective collision: 1. enough energy to react/break bonds 2. correct orientation A 2 step mechanism might result in double hump. During phase change, added or removed heat doesn’t change kinetic energy of particles and thus, does not change the temp. Energy is used to break IMFs or released when formed. Slow step only determines rate of reaction. (slow step has higher hump) Activation energy with catalyst How is rate law determined experimentally? 1. Method of Initial Rates: Start with different concentrations of reactants and measure initial rate. a. If tripling A causes rate to increase 9fold, exponent of A is 2. 2. Measure the concentration over time. Make a graph to see what order. 1 mol gas at STP = 22.4 L 2 ways a gas may not be ideal: 1. Large size 2. Polar 5 factors that affect rate of reaction: 1. nature of reactants 2. concentration: increases # of collisions 3. phase/ surface area: increases # of collisions 4. temperature: increases # of collisions AND # of effective collisions 5. catalyst: Increases # of effective collisions by lowering activation energy Consider this reaction: A+ B D+ E Rate = k [A]x[B]y[C]z NOT to be confused with equilibrium expression: An unideal gas: 1. Tends to have inelastic collisions, and/or 2. Have a size that DOES matter 3. AND doesn’t obey the gas laws perfectly. Any gas is more real (unideal) at 1. High pressure (size of particles does matter) 2. Low temperature (particles more likely to have inelastic collisions) Any gas is more ideal at 1. Low pressure 2. High temperature 2 balloons each filled with different gases at the same volume in the same room. K= 1. Exponents (in rate law) are NOT determined by balanced equation. 2. Exponents can match slow step of mechanism. 3. Exponents can only be determined experimentally. 4. k depends on T. 5. A catalyst can be included (C perhaps) 6. Products are (almost) never included – no denominator (in most cases) Rate law Rate depends on [A] Integrated Rate law Plot needed to give a straight line Relationship of rate constant to slope of straight line Half life Half life depends on [A] He CO2 Same V, T, and P Same KE. Same number of gaseous particles Higher mass, higher density Lower mass, lower density Lower speed of particles Higher speed of particles Zero Order Rate = k NO! [A] = -kt + [A]o 1st order Rate = k[A] ln [A] = -kt + ln [A]o [A] vs. t ln [A] vs. t 2nd order Rate = k[A]2 Yes 1 = kt + 1 [A] [A]o 1/[A] vs. t Slope = -k Slope = -k Slope = k t1/2 = [A]o/2k yes t1/2 = 0.693/k NO! t1/2 = 1/k[A]o yes Yes What are IMFs? 1. NOT bonds 2. Attractions between neighboring species 3. MUCH Weaker than bonds usually Increasing m.p., b.p , Hfus, Hvap, critical point, specific heat Increasing vapor pressure, volatility, likelihood of being gas at room temp The 5 types of intermolecular forces (IMFS) listed from strongest to weakest are: 1. Ion-ion (NOT really an IMF) Example: Salt bonds 2. Ion-dipole Example: Salt dissolving in water 3. Dipole-dipole (Strongest type is hydrogen bonding) Example: Water condensing 4. Dipole-induced dipole Example: CO2 dissolving in water 5. Induced dipole-induced dipole AKA London dispersion forces AKA Van der Waals forces Example: freezing nitrogen Ways to increase solubility of a Solid Increase temp Increase surface area Agitation, stirring gas Decrease temp Increase pressure above Don’t agitate ΔHsoln (heat of solution): SUM OF - Energy required to break the water-water IMFs Energy required to break IMFS in solute Energy released when IMFS form between solute and solvent. Hydrogen bond: the ATTRACTION between a H of one molecule (that is bonded to a very EN element such as F, O, N) and a very EN element (O, F, N) of another molecule. NOT A BOND!!!!!!!!! Ways to measure the concentration of a solution: The stronger the IMFS, the lower v.p. at given T. The stronger the IMFs, the higher the normal b.p. Types of crystalline solids: Ionic –formed from ionic bonds, very strong, brittle, high MP, think NaCl Molecular – formed from the regular arrangement of IMFS between covalent molecules, low MP, not strong, think snowflakes Network – covalent bonds throughout, high MP, hard, nonconducting (usually), think SiO2 (sand/glass), C (diamond) Graphite is a unique network solid made only of carbon atoms. They are bonded together in sheets. The sheets are attracted to each other by weak IMFS. (rubs off easily) Within the sheets, there is resonance throughout, allowing graphite (a nonmetal) to conduct electricity! Molarity (M) = mol solute/L solution Molality (m) = mol solute/kg solvent Mole fraction (X) = mol solute/ mol solution % by mass = g solute/g solution % by volume = mL solute/mL solution m/v % = g solute/mL solution Colligative property: any property of a solution that changes based SOLELY on the NUMBER of solute particles dissolved. (not mass or size or IMF formed!) -freezing point depression -Vapor pressure lowering -Boiling point elevation (as a result of v.p. lowering!) WHY? Solute particles interfere with the process of evaporating, boiling, and freezing. Alloys are metal solutions. (homogeneous mixtures) Substitutional alloy: formed by 2 metals of similar radii; mp or ability to resist corrosion or conductivity might be affected. Hardness is NOT. Strong acid: Weak acid: Strong base: Weak base: Soluble ionic compound: Insoluble ionic compound: Soluble covalent compound: Insoluble covalent compound: Interstitial alloy: formed by 2 metals of different radii; adding in smaller metals tend to make the alloy harder. For both alloys, adding metals with more valence electrons increases conductivity. DO NOT MIX CONCEPTS! Periodic trends do NOT explain boiling point for example. WHY …? Does X have a larger radius than Y? a higher ionization energy than Y? a higher electronegativity than Y? a higher reactivity than Y? Does Y make a given shape? Have a given bond angle? ONLY ReasonsForce you will of use attraction Tell how many protons and electrons each species has Then, use proton-electron attraction (Coulomb’s LAW) OR electron-electron repulsion (Coulomb’s LAW) OR shielding (electrons don’t feel protons pulling on them due to intervening energy levels) VSEPR THEORY!!! Bonding pairs of electrons experience repulsion with each other so they spread out in 3d space as far as possible; Lone pair electrons induce greater repulsion and cause angles between bonding electrons to be smaller than expected. Is Y polar? Does Y dissolve in Z? Have a high freezing pt, low volatility, low vp, high critical pt, high melting pt, high Hfus, high Hvap? Does SOLUTION A Have a high b.p., high conductivity, low f.p., low v.p? Electronegativity difference (polar or nonpolar bonds?) AND Symmetrical or Unsymmetrical shape Identify intermolecular force involved Compare strengths of intermolecular force Compare [ ] of solutions Ksp problem types: always AB A+ + B1. Calculate solubility using Ksp: Do RICE solve for x 2. Calculate Ksp using solubility: Do RICE, solve for K 3. Calculate solubility within a solution of common ion: Do RICE, solve for X (should not have all 0’s on I row) 4. Calculate if precipitate forms when mix 2 solutions: - MV = MV to determine new concentration - Plug into equil expression to get Q - Compare to K 5. Calculate how much of a given substance is required to form precipitate, (or compare which precip forms 1st) - Solve equil expression for missing value! Le Chatelier’s principle: If a reaction is stressed, the reaction responds in such a way as to remove that stress. Equilibrium can be disturbed by: 1. Adding products (shifts left) 2. Adding reactants (shifts right) 3. Removing reactants (shifts left) 4. Removing products (shifts right) 5. Adding volume or decreasing partial pressure (shifts towards side with fewer gaseous or aqueous particles) 6. Increasing temperature (shifts to get to new K) a. Shifts right if endothermic b. Shifts left if exothermic Buffer: combination of (weak) conj. acid/base pair that resists change in pH A good buffer: 1. desired pH = pKa of acid 2. [conj. acid] = [conj. base] 3. [conj.acid] and [conj. base] high (increases buffering capacity) How to approach ANY titration calculation: 1. Molarity x volume = moles for both substances. Subtract to see what remains. Does a conjugate acid or base get produced? (it will in any weak titration.) Determine its moles. 2. Divide by volume to get molarity for substance remaining. 3. THINK….how do I get pH of this/these substances? a. –log [strong] b. If weak, RICE c. If weak acid and weak base are present, do H-Hasselbach. Equilibrium canNOT be disturbed by: 1. Adding inert gas (no effect) 2. Adding solid (no effect) The ONLY way to change K is to change the temperature K >1 products preferred at equilibrium At equivalence point of ANY titration: MaVa = MbVb, moles acid – moles base = 0 Midpoint = half moles neutralized (AND [weak acid/base] = [conj. acid/base]); pH = pKa Titrations Starting point Midpoint Equivalence point Strong acid/strong base high Not relevant Exactly 7 Vertical region Very long Strong acid/weak base not as high pH= pKa Below 7 (conjugate acid present) Not as long Strong base/ Strong acid low Not relevant Exactly 7 Very long Always pick an indicator whose pKa = pH at the equivalence point of the titration. Strong base/weak acid not as low pH = pKa Above 7 (conjugate base present) Not as long Various Ways to Describe Acid Strength Property Strong Acid Weak Acid Ka Value Too large to measure << 1 Position of the dissociation equilibrium Equilibrium concentration of H+ compared to the original concentration of HA Equilibrium concentration of HA compared to the original concentration of HA Percent Dissociation right left [H+ ] = [HA]o [H+ ] << [HA]o [HA] ~ 0 M [HA] << [HA]o [HA] ~ [HA]o 100% < 5% Strength of conjugate base Conjugate base is nonexistent How to represent in a net ionic equation H+ + A- Conjugate base is also weak, but inversely proportional in strength to the acid (Kb = Kw/Ka). HA Sign Of H0 1. Exothermic 2. Exothermic 3. Endothermic 4. Endothermic + + System Entropy Change Increasing Decreasing Increasing decreasing Galvanic Cell/ Voltaic Cell/ Electrochemical Cell Spontaneous Electrolytic Cell Nonspontaneous Ex: electrolysis water Sign of S0 Sign of G0 Sign of E0 (if a redox) K + + - - + >1 Spontaneous? (Thermodynamically favorable?) always + or - - or + Don’t know Low T + or - - or + Don’t know High T + - <1 never Chemical energy electrical energy E=+ ΔG = - anode= oxidation Cathode = reduction Electrical energy chemical energy E=- ΔG = + Anode= oxidation Cathode= reduction Ex: battery! As reaction progresses, K doesn’t change BUT When rxn is shifting right When rxn is shifting left When a reaction is at equilibrium Q approaches K Q<K Q>K Q=K G approaches 0 G=- G=+ G=0 E approaches 0 E = + E=- E=0 No ions formed usually +1 +1/-1 Increasing nuclear charge explains everything: increasing ability to attract electrons, decreasing AR, increasing EN, increasing IE, which both explain +2 -4 -3 -2 -1 Most reactive nonmetal, highest EN S block decreasing reactivity across metals, increasing reactivity among nonmetals +3 Alkali ne earth metals Alkali metals p block d block (1 behind!) 5770 Zn +2 Ag Cd +1 +2 (+2) +4 Sn Pb Transition metals * (+3) +5 +6 numeral Variable positive charge = roman numeral 89102 +7 ** Increasing # of energy levels (increased shielding) explains everything: Decreasing ability to attract electrons, Increasing AR, Decreasing IE, Decreasing EN Which explain increasing reactivity among metals, decreasing reactivity among nonmetals Sb roman numeral Roman Most reactive metal, lowest IE, lowest EN * Lanthanide series, 4f Variable positive charge = roman numeral Rare earth metals ** Actinide series, 5f Highest IE No measurable EN (too low), high IE, INERT! H 5770 * 89102 ** * ** Page 16: Look on the internet for an interesting molecule that has many atoms. Draw its Lewis structure below and 1. Give its molar mass, 2. Give its mass per molecule. Then pick 2 central atoms that have different hybridizations and label their 3. Hybridization, 4. Bond angles, 5. Shape about that atom A good website to try might be http://www.chm.bris.ac.uk/motm/motm.htm (Molecule of the Month). Which atom? 1. 2. 3. 4. 5. 4. 5. Which atom? 3.