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Department
Chemistry
Course
CHEM 1061
Professor
Hyunjoo Im
Semester
Fall

Description
Class Notes SI System  Length – meter (m)  Mass – kilogram (kg)  Time – seconds (s)  Temperature – Kelvin (K)  Amount – mole (mol)  Electric Current – ampere (A) 3  Volume = m (cubic meters) o 1 mL = 1 cm 3 3 o 1L (liter) = 1 dm Prefixes: 3  Kilo = 10  Hecto = 10 2 1  Deca = 10 (da) or (dk)  Deci = 10 (d) -2  Centi = 10 -3  Milli = 10  Micro = 10 (µ) -9  Nano = 10 (n)  Pico = 10-12(p) Temperature Conversions  K = °C + 273  T F 1.8(T )C+ 32 Precision Vs. Accuracy  Precision = data collected = consistent w/ each other  Accuracy: how close measurements come close to theoretical value  You’d rather have precision Significant Figures (Sig figs)  Only for things that are measured  Counting rules: if decimal present, find first nonzero from left + count everything after; no decimal: go from left to right  When do sig figs not apply? o Conversion factors o Counting # o Fractions  Multiplying/Dividing: ans can have no more sig figs than the measured # w/ least sig figs Ex: 25.26(.0036) = 2 sig figs  Adding/Subtracting: ans can have no more decimal places than measured # w/ least places after decimal Ex: 53.00 + 65.2 = … .2  Do order of operations correctly (if switch from m/d to a/s, change # of sig figs)  *AP: 1 sig fig more or less = full credit if no idea how many sig figs, choose 3 Dimensional analysis: multiply by conversion factors; place so that units are diagonal and cancel out; usually same # sig figs as starting #; Ex: 3.5 days * (24 hr/1 day) * (3600 sec/1 hr) 3  Ex: a block of wood = 4.41 in * 4.13 in * 7.36 in. What’s volume in cm ? o 4.41 in * 4.13 in * 7.36 in = 134ish  134in * (2.54 cm/1in) = 2.20 * 10 cm 3 3 Density  Density = mass/volume  *Density of H 2 @ room temperature used to be 1.00g/mL not anymore, use density table to find using exact temperature  Ex: 125 mL Erlenmeyer Flask. Can you verify total vol. of flask from bottom to top o (Mass flask + water) – mass flask empty = mass water o Find temperature see density  find volume Moles  Allows us to work w/atoms & molecules that = too small to count  Avagadro’s # = amount of substance that contains as many elementary entities as there are atoms in exactly 12g of carbon 12. C-12 contains exactly 6.02 * 10 atoms 23  Molar mass = mass of 1 mol (6.02 * 10 ) of sample of element, molecule, or compound *round mass to hundredth’s place  Compound: made of 2 or more different elements  Molecule: made of 2 or more elements, but can be same  *not all molecules = compounds, but all compounds = molecules  Atomic Mass: average mass of existing isotopes of element including relative abundance decimal pt. # on periodic table  Mass number: represents mass of protons & neutrons for a particular isotope; must be whole #, also = # of both protons + neutrons  1 mol = 6.02 * 10 atoms, molecules, compound  1 mol of any gas @ STP (0°C, 1 atm) = 22.4L Molarity  Moles of solute per L of solution (sol’n) (both solute + solvent); measure of concentration; (M) or (mol/L)  Ex: What’s molarity of 156.5mL of sol’n that contains 36.5g Na? o 156.5mL .1565L o 36.5g Na * (1 mol Na/22.99g Na) = 1.5876…mol Na o 1.5….mol/.1565L = 10.2M Reactions (Rxns): reactants on left, products right, states of matter in parentheses, coefficients can mean atoms, molecules, moles, above arrow = heat, catalyst, electricity  *AP: AP doesn’t grade states of matter, what’s above arrow States of Matter  Solids: strong intermolecular forces (IMFs) o Densely packed difficult to compress o Fixed shapes and volume (vol) o Vibrational motion move back and forth around fixed pt.s  Liquids: fixed vol, no fixed shape o IMFs = weaker than solid gives them fluid motion, but IMFs = strong enough to hold them together  set vol o 3D motion  Gases: weakest IMFs o No fixed shape/vol take shape of container o Random motion in all directions (straight lines) Diatomics: H 2 N 2l 2r I2F “2 2cl2riff” S P 8 4 Metals  Ductile  Solids @ room temperature (temp)  Good conductors  Malleable  Shiny  React w/ acids Nonmetals  Dull  Brittle (not malleable)  Any state of matter @ room temp  Not good conductors  Tend not to react w/ acids Metalloids: characteristics of both Molarity  Dilution equation: M 1 =1M V 2 2 (initial molarity)(initial vol) = (final molarity)(final vol) o # moles of solute stay same  Ex: How many mL of 10.15 M NaOH stock sol’n needed to prepare 15.0 L of .315M NaOH? o (10.15M)V = (15.0L)(.315M) V = .4655172414L = 466mL 1 1  Ex: How many mL of 5.15 M CH OH st3ck sol’n needed to prepare 375mL of sol’n having 7.05mg of CH OH per mL sol’n? 3 o .00705g * (1 mol CH OH/32.05g CH OH) *3(1/1mL) * (1000mL/1L) = .234M o (5.15M)V =1(.234M)(375mL) o How to make: take 17.1mL of 5.15M sol’n, dilute up to 375mL w/ water Physical vs Chemical  Physical Properties: can be observed w/o changing composition of substance o Color o Odor o Solubility o Density o Mass o Conductivity o Hardness o Viscosity o Boiling/melting pts  Chemical Properties o Flammability o Reactivity o Oxidation  Physical Changes o Changes in state of matter  Chemical changes o Get new substance o Signs:  Color change  Precipitate formation  Gas formation  Temperature change  Light  Smell Stuff….  Substance held together by covalent bond = molecule  Ionic compound = formula unit  (NH ) SO = formula unit 4 2 4  Ex: 196.25g NiCl w2uld have how many atoms of Cl? o 196.25g NiCl * (1mol NiCl /129.59g NiCl ) * (2mol Cl/1 mol NiCl ) * (6.02 * 2 2 2 2 23 24 10 atoms Cl/1 mol Cl) = 1.8233 * 10 atoms Cl  % by mass = mass % composition: proportion of each element expressed as # of g of that element per 100g of compound  Ex: calculate mass % of each element in Ba(OH) 2 o (137.33/171.35) * 100% = 80.16% o 32/171.35 = 18.67% o 2.02/171.35 = 1.17% o *since we’re rounding all masses to 2 decimal places, % should also have 2 decimal places Solvation  Like dissolves like  Polar, ionic polar dissolves them o *sugar dissolves but doesn’t dissociate; ionics dissociate  Nonpolar dissolves nonpolar only  P. 136 solubility table MEMORIZE!!!!  Electrolyte: sol’n that contains enough ions to conduct a current  Nonelectrolyte: can’t conduct currecnt  Ion concentration of sol’n: what’s ion ion conc. Of Na and Cl in .100M sol’n of NaCl ICE tables!!, or use dimensional analysis to see how many moles of each ion there are in one formula unit Hydrates/Stoich/% Error  CaCl 26H O 2 hydrate calcium chloride hexahydrate; (-6H O) = CaC22 = anhydrate  % by mass composition: just add mass of 6H O 2 o CaCl ·62 O → 2aCl (salt) 2 6H O(g) 2  Stoichiometry: balance equations use coefficients as ratios o SiCl + 4H O → 2iO + 4HCl 2 o 2C H + 7O → 4CO + 6H O 2 6 2 2 2 o Ex: What mass of Mg needed to convert 83.6g TiCl to Ti? 4  TiCl 4 2Mg →Ti + 2MgCl 2  83.6g TiCl *4(1mol TiCl /184.68g TiCl ) * (4mol Mg/1mol TiCl ) * 4 (24.31g Mg/1mol Mg) = 21.4g Mg o Ex: If a sol’n is 65% H SO 2sul4uric acid) by mass and has a density of 1.5g/mL how many mL of sol’n are required to convert 1.00kg NH (solid 3mmonia) to (NH )4 2 ? 4  2NH +3H SO 2 (N4 ) SO 4 2 4  1000g NH * 31mol NH /17.013 NH ) * (1mol3H SO /2mol NH2) * 4 3 (98.08g H SO /1mol H SO ) * (100g sol’n/65g H SO ) * (1mL 2 4 2 4 2 4 sol’n/1.55g) = 2860mL H SO s2l’n 4  Limiting Reactant: used up 1 , determines how much product can be made o Ex: 3Mg(s) + N (g) →2Mg N (s) 3 2 o A) if 35g of Mg + 15g of N react,2how much Mg N is produc3d?2  35.00g Mg * (1mol Mg/24.31g Mg) * (1mol Mg N /3mol Mg3 *2(100.95g Mg N3/2mol Mg N ) 3 42.45g  15g N *2(1mol N /28.22g N ) * (12ol Mg N /1mol 3 )2* (100.952 Mg N /1mol Mg N ) = 54.04g 3 2 3 2  *AP: write a little note saying who’s limiting reactant o B) how much excess reactant remains?  35g Mg * (1mol Mg/24.31g Mg) * (1mol N /3mol Mg)2* (28.02g N /1mol 2 N 2 = 13.48g N 12g – 13…g = 1.553g N 2  % Yield o Actual yield (A) = produced in lab o Theoretical yield (T) = calculated o % yield = A/T * 100% o Ex: What mass of acetic acid needed to prepare 252g (actual) of ethyl acetate if expected % yield is 85%? Assume acetic acid limits  CH C3OH(a.a.) + CH CH OH(2than3l) → CH CO CH CH (eth3l 2 2 3 acetate) + H 2  .85 = 252g/xg 296g theoretical  296g e.a. * (1mol e.a./88.11g) * (1mol a.a./1mol e.a.) * (60.05g/1mol a.a.) = 202g a.a. Atomic Structure  J.J. Thomson: cathode ray tube, negatively charged particles  George Stoney: named e (electron)  Robert Millikan: determined mass of e - - +  Thomson’s model: plum pudding .e- - .e .e  Rutherford: Gold foil experiment at alpha particles - .e o Only few bounce away/back since most = empty + .e- .e- o Positively charged densely packed nucleus - +  Bohr: e in set paths w/ set amt (amount) of energy  Schrödinger: Quantum Mechanical Model - - o e in regions of probability = 90% chance of finding where e are likely to be o To place e in set energy levels (E, n) sublevels: s, p, d, f orbitals: 1, 3, 5, 7 ↑/↓ spin Sublevels Orbitals Shape Max # e - s 1 2 p 3 x3 6 d 5 10 x4 x1 f 7 Too complicated 14  Energy Diagram have s, p, d, etc lined up going down to up, ascending order, higher sublevels step up and out a bit, fill lines w/ ↑/↓ spins  Aufbau: e always fill atom from lower energy to higher energy level (Au low alphabet to high alphabet in aufbau)
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