Study Guides (248,269)
Canada (121,449)
Chemistry (166)
CHEM 120 (26)

CHEM 120 Exam Review.docx

19 Pages
Unlock Document

CHEM 120
Carey Bissonnette

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
More Less

Related notes for CHEM 120

Log In


Join OneClass

Access over 10 million pages of study
documents for 1.3 million courses.

Sign up

Join to view


By registering, I agree to the Terms and Privacy Policies
Already have an account?
Just a few more details

So we can recommend you notes for your school.

Reset Password

Please enter below the email address you registered with and we will send you a link to reset your password.

Add your courses

Get notes from the top students in your class.