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University of Waterloo
CHEM 120
Carey Bissonnette

CHEM120EXAMNOTES Module 1:Stoichiometry Review o Compound: composed of two or more elements, has fixed composition o Mixture: composed of two of more substances, has variable composition o Mole=6.022x10^23 particles (Avogadro’s number)  Based on 12 grams of carbon 12 o Atomic mass unit: 1/12 if the mass of one carbon 12 atom o Mole method:  Convert to moles (n=mm/m)  Convert between moles (ratio)  Convert from moles (m=n x mm) o Empirical formula: determine ratio of moles of one element to another, divide by smallest to get whole numbers o Molecular formula: given molar mass of the compound, mm =X (mm compound empirical),e X is the multiplication factor o Limiting reagents: calculate moles of each reactant then multiply by stoichiometry ratio to determine the limiting. Use limiting for further calculations o Percent Yield= (actual yield/theoretical)*100  Actual yield should always be less than theoretical o Consecutive reactions: a series of reactions that occur sequentially; products from one reaction are consumes as reactants for the next o Simultaneous reactions: reactions are independent and occur at the same time  Never add the chemical equations together  Complicated simultaneous reaction: introduce variables to represent difference between the masses Module 2: Aqueous solutions and reactions in aqueous solutions o Aqueous solutions have water as a solvent o Ionic compound: comprised of positive and negative ions; held in position by strong ionic forces; solids at room temperature o Molecular compound: comprised of stable, neutral molecules; solid, liquid or gas at room temperature; molecules become hydrated when dissolved in water o Dissociation: separation of an entity into two or more  Ionic compounds produce ions in solution by dissociation o Ionization: the generation of one or more ions  Molecular compounds produce ions by ionization o Strong electrolytes: ionize or dissociate completely in water o Weak electrolytes: do not ionize or dissociate completely in water o Molar concentration: # moles of solute per litre of solution o Molar solubility: maximum number of moles of solute per litre of solution CS  the solution is super-saturated Solubility rules: 1. Salts of alkali metals are soluble 2. Ammonium (NH4+) salts are soluble 3. Nitrates are soluble 4. Chlorides, bromides and iodides are soluble except with mercury and lead 5. Sulphates are soluble except with calcium, strontium, barium (group 2), mercury and lead 6. Carbonates, phosphates and sulphides are insoluble except with alkali and with ammonium 7. Hydroxides are insoluble or slightly soluble except with alkalis o Net ionic equation: break everything into ions except the solid precipitate, cancel out spectator ions o Acids: proton donor o Strong acids  Hall: hydrochloric acid  HBr: hydrobromic acid  HI: hydroiodic acid  HClO4: perchloric acid  HBrO4: perbromic acid  H2SO4: sulphuric acid (second ionization energy does not go to completion)  HNO3: nitric acid o Base: proton acceptor o Strong bases  Group 1 hydroxides: LiOH, NaOH, KOH, RbOH…  Group 2 hydroxides: Mg(OH) 2, Ca(OH)2. Sr(OH)2, Ba(OH)2 o Acid-base neutralization reactions produce water and salt o Rules for assigning oxidation states 1. The oxidation state is 0 for an atom in elemental form 2. The sum of oxidation states must equal the total charge 3. In their compounds, group 1 metals always have an oxidation state of +1, and group 2 has +2 4. F always has an oxidation state of -1 5. H normally has an oxidation state of +1 except when it combines with group 1 or 2 metal 6. O normally has an oxidation state of -2 except when bonded to itself 7. Cl, Br and I usually have an oxidation state of -1 unless the preceding riles dictate otherwise o Oxidation: loss of electrons o Reduction: gain of electrons o Balancing redox reactions in aqueous solutions  Write separate half reactions  Balance non 0,H atoms  Add electrons to each half reaction then multiply by factor so electrons equal each other and cancel out  Combine half reactions  Balance net charge in acidic by adding H+ atoms and adding OH- to basic  Balance the O and H atoms by adding water molecules  Check to make sure final equation is balanced Module 3: Gasses o Gasses: mostly empty space; low density (1-10g/L); 1000X less dense then liquid compressible o Pressure exerted by a sample of gas can be measured using a mercury manometer o P =P P =P +h P =P -h gas bar gas atm gas atm o Variable relations:  Increasing Tincreasing V PV=nRT  Increasing Vdecreasing P  Increasing nincreasing P o Derivation of ideal gas equation is based upon assumptions  Molecules of gas move randomly but in straight lines, change direction during collisions  When molecules collide, kinetic energy is converted  Distances between molecules are greater than the size of the molecules themselves (neglect point mass)  Attractive/repulsive forces are weak except when molecules collide  Each molecule in the gas has its own kinetic energy, average kinetic energy of molecules is proportional to K temperature  Deviations from ideal gas behaviour are most significant with high pressure and low temperature o Calculating gas variables before/after use… 𝑃1 1 𝑃2 2 =  n and R are constant so they are 𝑇1 𝑇2 cancelled out o Ideal gas equation derives relationship with molar mass and density… 𝑀𝑀 = 𝑑𝑅𝑇 𝑃 o Fixed volumes of two reactants, treat as limiting reagent  Multiply given volume by stoichiometry ratio o The sum of partial pressure is equal to the total pressure  Calculate by dividing moles of desired substance by the total moles, then multiplying by total pressure o Collecting gas over water  Wet gas is yielded meaning water vapour is mixed within the compound  Use table to determine pressure of water at given temperature then subtract to determine pressure of substance  o Kinetic molecular theory  Molecules are in continuous random motion 3  The average kinetic energy is proportional to kelvin𝑣𝑔 = 2𝑅𝑇 temperature  Distribution of a gas is narrow if T is small  For fixed temperature, the lighter the gas, the broader the distribution 8𝑅𝑇 𝑇 3𝑅𝑇 𝑇 𝑉𝑎𝑣𝑔 = 𝜋𝑀𝑀 = 𝑀𝑀 𝑉𝑟𝑚𝑠 = 𝑀𝑀 = 𝑀𝑀 o Effusion: molecules escape from its container through 𝑟𝑎𝑡𝑒𝑥 = 𝑀𝑀𝑦 opening 𝑟𝑎𝑡𝑒𝑦 𝑀𝑀𝑥 o Diffusion: molecules of one gas mix with another o Real gasses: deviations from ideal gas behaviour are small at STP but very significant at high pressures and low temperatures  Ideal gas model neglects sizes of the molecules and intermolecular forces  Van der Waals equation accounts for… (Equation is not accurate) -a: measure of the strength of intermolecular forces (modifies pressure) -b: measure of the size of the molecules (modifies volume) o Viral equation of state: contains correction terms to account for deviations from ideal gas behaviour 𝑃𝑉 𝐵 𝐶 𝑅𝑇 = 1 +𝑉 + 2… 𝑉 Module 4: Thermochemistry o Physical change: heating, cooling, expansion, compression, phase change o Chemical change: changed by chemical reaction o Types of systems  Open: energy and matter can be exchanged between system and surroundings  Closed: only energy is exchanged between system and surroundings  Isolated: neither energy nor matter is exchanged between system and surroundings o Internal energy: energy of an object or system that is molecular in nature; sum of the kinetic and potential energies of all the particles in a system  Heat flows in or out of the system (heat out=neg)  System does work or has work done on it (expansion=neg) o Heat: energy is transition; flows from area of high temperature to area of low temperature  2 objects with the same temperature= thermal equilibrium  Increasing temperature=increasing average kinetic energy  Phase changes occur at constant temperature; heat is used to overcome attractive forces that cause the molecule to aggregate (increases the potential energy of the molecules) o Heat capacity: amount of heat required to raise the temperature of an object or substance by 1degree Celsius or 𝑞 = 𝑚𝑐Δ𝑇 1 kelvin Or 𝑞 = 𝑛𝑐Δ𝑇  Depends whether it occurs at constant heat or constant pressure  Pay attention to the units of C to know which equation to use  Always divide into separate entities to determine rationality behind the positive or negative value of q 𝑞1 + 𝑞2 = 0  Assume no heat is loss so use equation o Work: energy in transition; many types  Expansion work: volume of the system changes in 𝑤 = −𝑃 𝑒𝑥𝑡ΔV the presence of an external pressure  Minus sign in the work equation is from the fact that the pressure exerted by the surroundings opposed an increase in volume. o First law of thermodynamics: Energy is conserved, cannot be created not destroyed, just converted or transferred  Governed by heat and work Δ = 𝑞 + 𝑤  U has not interpretation, but the change in U does  When heat is absorbed, the molecular motions are affected in a random way  When work is done, molecules are moved in a particular way o Constant volume heat of reaction qv= o Constant pressure heat of reaction qp= o Enthalpy  o Most reactions are carries out at constant pressure Δ = 𝑈 + 𝑃𝑉 o Convert constant volume heats of reaction into constant Δ = Δ + Δ𝑛 𝑔𝑎𝑠RT pressure heats of reaction; qv does not equal qp o Bomb calorimeter: insulated from surroundings; treat as constant volume equation o Thermochemical equations: chemical equation that includes . Right side of the equation for exergonic, left side for endergonic  Reverse reaction=change the sign of value  Multiply reaction by “n”=multiply value by “n”  Add reactions together= add the values together o Formation reactions: reaction to create 1 mole of a substance that is formed from its elements.  is 0 for an element in its reference form (graphite is elemental form for carbon)  = ∑ − ∑ Module 5: Overview of quantum theory o Wavelength: distance between successive maxima o Period: time it takes for the electric field to return to its maximum strength λv = 𝑐 o Frequency: # of times per second the electric field C=2.998*10^8m/s reaches its maximum value o Light is classified according to wavelength, visible light is between 400- 750nm Increasing frequency  Gamma rays, x-rays, UV rays, visible, infrared, microwave, radio Increasing o Blackbody radiation: a heated solid pro
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