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CHEM 120 Exam Review.docx

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Department
Chemistry
Course
CHEM 120
Professor
Carey Bissonnette
Semester
Fall

Description
CHEM 120 Exam Review *Module 1: Stoichiometry (4x1) - stoichiometry: quantitative study of the composition of compounds and mixtures and other the amount of reactants or products involved in a chemical reaction - compound: composed of two or more elements; has a fixed composition (every pure sample will always contain the same elements)  Water is always 89% oxygen by mass and 11% hydrogen - mixture: composed of two or more substances; has a variable composition  Water and ethanol can be mixed in any proportion 23 - mole: amount of substance; 6.022x10 particles - atomic mass unit: convenient for expressing the mass of a single atom or molecule  1/12 of the mas of one C; a single atom of C has a mass of 12 u - average atomic mass: weighted average of the atomic masses of the various isotopes - atomic number (Z) = # protons in nucleus - Avogadro’s constant (N )A 6.022x10 mole -1 - n = m/M = mass (g)/molar mass(g/mol) - consecutive reactions: reactions that occur sequentially; products form one reaction are consumed as reactant in subsequent reaction  A + B  C + D  C + E  F  able to ass the chemical equations together  “C” is the intermediate - simultaneous reactions: reactions that are independent and occur at same time  A + B  C + D  A + E  C + F  No product in reactant  NEVER add chemical equations together *Module 2: Aqueous Solutions and Reactions in Aqueous Solutions (4x1) - solution: homogenous mixture of one or more solutes in a solvent - homogenous: uniform composition right down to molecular level; the molecules of one substance are mixed uniformly amongst the molecules of the others - solvent: determines the phase of the solution (solid, liquid, gas); usually the most abundant component - solutes: all other components of solution - aqueous solutions have water as a solvent - (aq) = susbastnce has been dissolved in water - ionic compound: NaCl, KNO , 3H NO4etc.  Comprised of positive and negative ions arranged in regular, repeating patterns  Held in their positions by strong ionic bonding forces (positive + negative)  Solid at room temperature  Dissolves in water, pos. and neg. ions break away from sold surface and become “hydrated” (surrounded by water molecules)  Water molecules stabilize the ions in solution - molecular compounds: CO , CCl2, C H4et10 8  Stable, neutral molecules  Held together by covalent bonds (shared electrons)  Solid, liquid or gas at room temperature  Dissolves in water, molecules become hydrated  We may or may not get ions in solution; depends on whether the molecule reacts with water to produce ions (if acid/base, will produce ions) - dissociation: the separation of an entity into two or more entities + -  NaCl  Na + Cl - ionization: the generation of one or more ions  CH 3OOH + H O 2===> H O + C3 COO 3 - - an ionic compound (NaCl) produces ions in solution by dissociation - a molecular compound (HCl) products ions in solution my ionization - ionic compounds and some molecular compounds produce ions when dissolves in water; called electrolytes - solutes: nonelectrolytes (produce no ions when dissolved in water) and electrolytes (produce ions when dissolved in water) - electrolytes: strong (ionic/molecular compounds that dissociates or ionizes completely) and weak (ionic/molecular compounds that do no ionize completely – most of the molecules remain as un-ionized, neutral molecules) - molar concentration (C): number of moles per liter of solution (mol/L) - molar solubility (S): maximum number of solute per liter of solution - the concentration of a solute in a solution is limited by the solubility of the solute  C < S: solution is unsaturated; all of the solute dissolves; still more solute can be dissolved  C = S: solution is saturated; no more solute can be dissolved; addition of more solute normal causes solute to “come out of solution”  C > S: solution is super-saturated; very difficult to achieve and maintain; slightest impurity or agitation will cause solute to come out - in a precipitation reaction, ions in solution combine to form an insoluble solid (insoluble ionic compound) that precipitates from the solution Solubility Rules for Ionic Compounds Soluble: 1. salts of the alkali metals (group 1) - ex: Li, Na or K etc. X (bonded to) is soluble regardless of what X is 2. ammonium (NH ) sal4s + + - ex: (NH4)X, (NH )4X 2tc. are soluble regardless of what X is 3. nitrates (NO )3- - ex: MNO 3 M(NO ) 3 2. regardless of what M is 4. Chlorides (Cl ), bromides (Br ) and iodines (I ); except for the chlorides, 2+ + 2+ + bromides and iodines of lead (Pb ), mercury (Hg ) and (Hg 2 ) and silver (Ag ) - group 17 – halogens - ex: if X = Cl, Br or I then MX, MX 2tc. are soluble unless M = Pb, Hg or Ag 2- 5. Sulfates (SO 4 ); except for the sulfates of gr.2[calcium, strontium, barium], silver, mercury and lead - ex: M2SO 4 MSO et4. are soluble unless M = Ca, Sr, Ba or Pb, Hg, Ag Insoluble: 1. Carbonates (CO 3 2), phosphates (PO 43) and sulfides (S ); except if bonded to alkalis (rule 1) and ammonium (rule 2) 2. Hydroxides (OH ); except for hydroxides of alkalis (rule 1) - hydroxides of group 2 are slightly soluble; all others are insoluble - acid: proton (H+) donor  Considered “strong” if every acid molecule reacts with water  HCl, HBr, HI – binary acids containing one H and another element (X).  HF = weak acid (or any acid not listed)  HClO (4erchloric acid), HBrO , (p4rbromic acid) H SO (sul2uri4 acid), HNO 3 (nitric acid) – oxo acids where H is bonded directly to oxygen - base: proton (H+) acceptor  Considered “strong” if every base molecule reacts with water  Group 1 hydroxides: LiOH, NaOH, KOH, RbOH etc.  Group 2 hydroxides: Mg(OH) , Ca(O2) , Sr(OH2 etc. 2 2-  Hydride (H-) and oxide (O ) are also strong; converted to H or OH- 2ith H O 2  Base generates (OH-) ions in solutions either directly by dissociation (NaOH) or indirectly by ionization (NH reacts with water) 3  Hydroxide salts are ionic compounds containing OH- ions; produce OH- ions directly upon dissolving - acid-base reaction is a proton transfer reaction - when acid is dissolved in water, water molecules are protonated (ex. an acid produces H O (hydronium ion)) 3 - when base is dissolved in water, water molecules are de-pronated (ex. a base produces OH (hydroxide ion)) - an H 2 molecule can act as an acid or base, depending on what is dissolved; it is “amphiprotic” meaning both sides - sulfuric acid ionized in two distinct steps:  H 2O +4H O 2 HSO + H O 4- 3 +  HSO +4H O <2=> SO 42-+ H3O +  First ionization reaction goes to completion but the second does not (H SO 2 4 is a strong acid but HSO is4not) - neutralization reaction, an acid and base react to form a salt; often H O is2a product - if either acid or base is “strong”, neutralization goes essentially to completion (until limiting reactant is fully consumed) - oxidation-reduction reactions: reactions in which oxidation states change; “electron transfer reaction” Rules for Assigning Oxidation States 1. zero when in elemental form - ex: H2is zero, Na is zero 2. sum must equal total charge - ex: 2 , oxidation state is -1 3. group 1 metals (+1) and group 2 (+2) 4. F is always (-1); except when in elemental form (0) 5. H is normally (+1); except when combined with group 1 or group 2 6. O normally (-2); except when bonded to itself or bonded to fluorine 7. Cl, Br and I normally (-1); unless preceding rules dictate otherwise - oxidation state of an atom is a equal to the hypothetical charge an atom would have if the bonding electrons between pairs were given to the atom having the greatest electronegativity - Oxidation Is Loss of electrons; Reduction Is Gain of electrons  Oxidation: increase in oxidation state  Reduction: decrease in oxidation state - every electron transfer can be written as the sum of an oxidation and a reduction process, each of which is referred to as “half-reaction” - number of electrons produced in oxidation step must equal number of electrons consumed in reduction step - oxidation-reduction reactions that occur in acidic or basic solutions are difficult to balance by inspection; therefore, rules have developed a method 1. assign oxidation states 2. write half-reactions for oxidation and reduction 3. balance half-reactions separately  First with element being oxidized or reduced  Then by adding electrons to one side or another for number of electrons produced (oxidation) or consumed (reduction) 4. combine half-reactions so total electrons cancel out 5. balance net charge by adding OH- (basic solutions) or H+ (acidic) 6. balance O and H by adding H O;2check that final equation is balanced *Module 3: Gases (3x1) - gases: is most likely empty space; has a lower density and is very compressible (dentist is typically a few g/L; solids and liquids are incompressible and have densities of a few g/mL) - pressure (P): force per unit area - volume (V): provides of measure of space occupied - kelvin temperature (T): provides a measure of the average kinetic energy of the molecules in a sample  T (in K) = t(in C) + 273.15 (A) gas pressure equal (B) gas pressure greater (C) gas pressure less to barometric pressure than barometric pressure than barometric press. Pgas= P atm+ h Patm = Pgas+ h Pgas= P atm– h - increasing T: (keeping P and n constant)  V increases (gas expands when T increases) - increasing V: (keeping T and n constant)  P decreases - increasing n: (keeping T and V constant)  P increases - PV=nRT (ideal gas) – R is a universal constant - ideal gas equation is based on following assumptions:  The molecules of the gas move randomly but in straight lines, changing directions only when they collide with each other or with walls of container  When the molecules collide, kinetic energy is conserved  The distance between the molecules are much great than the sizes of the molecules themselves  The attractive and/or repulsive forces acting on a molecule are very weak except when molecules reach the same point in space or collide with walls of the container  Each molecule in the gas has its own kinetic energy but the average kinetic energy of the molecules is directly proportional to the kelvin temperature - ideal gas is realistic, provided that gas is not close to condensation point - derivations from ideal behaviour are most significant when the pressure is high (squeezed) or the temperature is very low (intramolecular) – PV=nRT works best at high T and low P - partial pressure: A (nA/n tot tot  (n An tots the “mole fraction of A”; sometimes X A  The sume of the partial pressures is equal to the total pressure (PA+P BP …c) - use of partial pressure when small amounts of gas are collected over water  Need to know VP (vapor pressure) at specific temperature  “wet gas” because some water molecules escape from surface of water to occupy and accumulate in the inverted cylinder  This method of collecting gas only works if the gas is no soluble in (doesn’t react with) water - two main assumptions of kinetic molecular theory are:  The molecules are in continuous, random motion  Average kinetic energy of the molecules is proportional to the kelvin temp. - cannot be in organized, parallel paths because the molecules would hit the irregular walls on the container and cause deflection in pathway; collision upon collision of particles lead to chaotic and randomness - for a given gas, distribution (speed) is narrow if T is small and broad if T is large; the lighter the gas, the broader the distribution of speeds - the average speed (V avg= √ (8RT/piM) and root-mean-square speed (V rms= √ (3RT/M) - molecules moving with V rms has a kinetic energy equal to average kinetic energy KE avg - for fixed T and P, lighter molecules travel faster than do heavier molecules - effusion: molecules escape from its container through an opening (ex. open tire) - diffusion: molecules of one gas mix amongst those of another (ex. perfume in air) - VavgV rms effusion and diffusion are proportional to (1√M) - (rate)x = √ M  (m/s, L/s, g/s, mol/s) y (rate)y M x - x = √ M x y M y - Graham’s Law of effusion: rate xTi= √ T1 ratexT2 T 2 - ideal gas equation is valid for gases at low pressures; derivations become increasingly significant as the pressure increases  STP: 0°C and 1bar=100kPa  Neglects: the size of molecules, intermolecular forces - real gases : V > n and P > P ideal real  Intermolecular forcecs cause the molecules to be drawn inwards and exert less pressure on the wall of container - “real” gas effects important high pressure, lower temperature (longer encounter time) - Van der Waals equation: (on data)  a – strength of intermolecular force (experimental developed)  b – measure of the sizes of molecules (experimental developed) - virial equation of state: (more accurate; preferred) *Module 4: Thermochemistry (5x1) - thermodynamics: “heat-power” concerned with inter-conversion of energy among its many different forms; predicting/understanding direction of energy transfer - physical: no change in chemical composition (heating, cooling, expansion, compression, phase change) - chemical: chemical composition is changed by chemical reaction - first law of themo.: energy is conserved - zeroth law: if 2 systems are each in thermal equilibrium with a 3 system, they are also in thermal equilibrium with each other - second law: then entropy of an isolated system cannot decrease - third law: the entropy of a perfect crystal is zero at 0K - system: part of the universe that we’re interested in  Open: energy and matter exchanged between system and surrounding  Closed: energy (not matter) exchanged between system and surrounding  Isolated: neither energy or matter is exchanged - surrounding: separated from system; real of imaginary boundaries - kinetic energy: energy an object or system has by virtue of its motion 2 (KE = ½ mv ); rotational to carry energy or vibrational bond length changing - potential energy: energy an object or system has by virtue of its position or configuration - internal energy (u): energy an object or system has by virtue of its molecular nature; the sum of kinetic energy and potential energy of all the particles in the system  Changes if: heat flows into or out of system or if the system does work or has work done to it  Heat and work are mechanisms for changing the internal energy of a system - heat: energy that flows from a region of high temp. to low temp.; heat is transferred via. molecular collisions - q represents total heat transferred; (+) or (-) signs  + = heat flows into system from surroundings  - = heat flows into surroundings from system - when a pure substance is heated: (assume no chemical change)  Temperature increases; heat being used to increase kinetic energy of the molecules  A phase change occurs; heat being used to increase potential energies of molecules - Tfusfusion temperature = melting - Tvap= vaporization temperature = boiling - qfus heat of fusion - q = heat of vaporization vap - heat capacity: amount of heat required to raise temperature 1°C or 1K  Varies from substance to substance tends to increase with molecular complexity -1 -1  C = q/ΔT c(JK or J°C ), q (J)  q=mc ΔTs(pure substances – C = spesific heat capacity JK g ) -1 -1  q=ncΔT (pure substances) - C = heat capacity for heating at constant volume v - Cp= heat capacity for heating at constant pressure - work:  expansion work: work done when volume of system changes in the presence of an external pressure; external pressure causes the volume change “pressure-volume work” or “PV work”  electrical work: work done to move a charged particle (e-) form a region of high electrical potential to low  surface tension work: work done when the surface area of a liquid changes - if ΔV > 0 (expansion) then W < 0 - if ΔV < 0 (compression) then W > 0 - w= -P ΔV (P is always positive) ext ext - closed system that undergoes a change in state: ΔU = q + w (internal energy)  q = (+) endothermic – flows into system (-) exothermic flows out of system  w = (+) done on system – compression (-) done by the system – expansion - constant volume (ex. rigid, sealed vessels)  bomb calorimeter - constant atmospheric pressure (ex. “open”)  beaker - ΔU=q – P ΔV (constant V, 0)  ΔU=q v ext v  Reaction that occurs at constant volume, the heat transferred (q )vis equal to the internal energy change (ΔU) for the system -ΔH = ΔU + PΔV (constant P) q = ΔH p  Reaction that occurs at constant pressure (q )p the heat absorbed/released is equal to the enthalpy change (ΔH) for the system  Most reactions are carried out at constant pressure - qvwont equal q p - calorimeter: device that is thermally insulated form its surrounding (ex. no heat is lost to the surrounding) - qcal CΔT - qrxn-q cal - qp= q v Δn gasRT  ΔH – constant pressure heat of reaction  ΔU – constant volume.  No net consumption or net production of gas -Δ H > 0 (q p 0) = heat is absorbed, reaction is endothermic (+) - ΔH < 0 (q p 0) = heat is released, reaction is exothermic (-) - ΔH° (T) = enthalpy change for a reaction carried out at T with each substance at a pressure of 1 bar  298K and 1 bar (SATp) - formation reaction (ΔH °): reaction in which one mole of a substance is formed f form its elements in their reference form (most stable form)  elements  single compound (one mole)  ex: ½ N (g) + 3/2 H (g)  NH (g) 2 2 3  Δ H f is called the standard enthalpy of formation (T=298K) - element C exists in two different forms:  Graphite and diamond  Graphite is more stable at 298K and 1bar; graphite is the reference form of carbon - ΔH f = 0 for an element in reference form (because the formation reaction for forming an element involves absolutely no charge, ex: O  2 ) 2 - ΔH° 298 = [sum of (c/d)products] – [sum of (a/b) reactants]  Coefficients in front of ΔH f obtained from table - Hess’ Law: if a reaction can be written as the sum of other reactions, the enthalpy change is equal to the sum of various enthalpy changes involved *Module 5: Quantum Theory and H Atom (2x1) - when an electron is “confined” to a finite region of space by forces exerted on it, its total energy is restricted to certain special values  At atomic level, only certain quantities of energy are allowed; we say the energy is “quantized” - light: electromagnetic radiation that transmits energy through space or some other medium; produced when electrical charges (ex. electrons) undergo some sort of acceleration  Ex: light from a white light bulb is produced through spatial and temporal oscillations of tungsten atoms in wire filaments (vibrating back and forth) - electric field: force exerted per unit charge - at a particular instant in time: the electric field varies from point to point (wavelength λ – distance [m]) - at a particular point in space: the electric field varies in time - wavelength λ (m): distance between successful maxima - period T (s): time it takes for the electric field to return to its maximum strength -1 - frequency v (1/T in s ): # of times per second the elec
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