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Chem 142 10272014 Instructor Dr Colleen Craig cfchemuwedu 202 Bagley Class Communications Use Canvas discussion board Email cfchemuwedu for private meeting Subject Chem 142B Syllabus Manipulating data Atomic nature of matter Stoich Behavior of gases Chem equilibrium Acidbase solubility Chapter 17 8 only solubility Evaluation Participation 5 o Aleks 10 Laboratory 15 2 midterm 40 o Finals 30 Exams 1 Friday 1017 2 Friday 1114 3 Final 1210 KNE 130 SCANTRON PHOTO ID CALC PENCIL Midterm 1 10 multiple choice questions 5pts each 50 atomic nature of matter 4 bonding lewis structure 2 stoichiometry 3 solution formation 1 3 free response 50 pts empirical and molecular formula 1Q 17 stoichiometry 2Q 15 18 Atomic Theory 10272014 Atomic Theory 10272014 Democritus Atomism 5BCE Divide an object completely atom undivided Properties of macroscopic object depends on properties of microscopic Law of conservation of mass 1780 Antoine Lavosier amp Marie Anne Paulze Mass is neither created nor destroyed Exp where mass of a closed system doesn t change after combustion Law of Definite Proportion 1790 Joseph Proust A given compound always has the same proportion of the elements regardless of where it comes from or how its made Used copper carbonate amp comparing a manmade amp natural one Dalton s atomic theory Dalton s theory of matter Elements are made of atoms Atoms of the same element are identical Compounds are formed by combinations of atoms They have the same composition Chemical reactions involve atomic reorganization but atoms don t change Only in macroscopic masses Compounds of relative amounts of atoms but not absolute amounts Assigns hydrogen as 1 and makes atomic masses based on it Atomic Theory 10272014 Assumption is wrong as hydrogen acts as H2 How does one determine absolute ratios of atoms in compounds Law of combined volumes 1808 Joseph Gay Lussac Combined pure gasses and measure volume of gases after each At constant P amp T Avogadro s constant 1811 At the same pressure and temp equal volumes of different gasses have the same number of particles Uncertain period Dalton why diatomic Why not act as monoatomic Berzelius ill make a table of atomic masses 0 as 100 Kekule set up atomic masses first then find compounds Clarity from Cannazzaro 1860 Used 2 hypothesis Compounds contain whole number of atoms Equal volume of gases under the same conditions have the same number of molecules Defined hydrogen as diatomic H2 has a mass of 2 JJ Thompson 1897 Cathode rays High electric potential across a vacuum Atomic Theory 10272014 Cathode rays can be deflected by an electric or magnetic field The properties of cathode rays are the same no matter what substance they are derived from Deduced that atoms contained charged particles that are very small making it able to be deflected Atoms are electrically neutral but contains charged particles Plum pudding model Robert Milikan 1903 Oil drop experiment Electric and gravitational field to find charge of oil drop Oil drops have an integer amount of charge the smallest value of which is the charge of an electron Ernest Rutherford 1909 Gold foil experiment Alpha particles bombarded at thin gold foil Some alpha particles bounce back Shows that gold atoms have a dense nucleus that is very small Atoms are mostly empty space Thompson 1913 Positively charged ionized neon gas affected by electric and magnetic field 3 deflection points of positively charged neon ions discovered isotopes using the mass spectrometer protons chemical identity of the atom Atomic Theory 10272014 electron ionic character neutron isotopic character Some massive definition Atomic number Z number of protons in nucleus Mass Number A sum of number of protons and neutrons in atom Average atomic mass an average of the atomic masses of the most common isotopes At room temperature all elements are solid except Gases H 2 N2 02 F2 CI2 Liquids Hg Br Formulas of elements Diatomic elements H 2 N2 02 F2 CI2 Br2 I2 Polyatomic species Sulfur S 8 Phosphorus P4 Type and number of bonds depend on both number of electrons and protons Covalent sharing of electrons non metas Ionic transfer of electrons metal nonmetals Atomic Theory 10272014 Chemical bond forms when the electrons of two atoms interact Energy is released when a bond is formed Bond because of complete octet in the valance shell of both atoms Ionic compounds have an overall charge of zero Periodicity Similar properties down a group Trend of properties across a period Electron Structure in atoms Electrons in atoms are fixed in quantized energy level Vaance electrons are in most outer shell or highest energy level Core electrons are remaining electrons Group number in A group naming valance number of electron Atoms strive for nobility Common Ions Cr2 3 Mn2 3 Atomic Theory 10272014 Fe23 Ck23 VQ3 Cu 2 Zn2 Ag Cd2 EH24 Au3 Hg2 Hg 22 Pb24 Lewis Dot structures 2D representation of the bonding pattern in a molecule only for covalent bonds between nonmetals 3 steps 1 sum the valance electrons for all atoms and determine total number of electrons 2 use pairs of electrons to form a bond between each pair of electrons 3 Arrange remaining electrons around atoms lone pairs and or multiple bonds to satisfy octet rule for each atom or duet rule Mass spectrometer Used to find abundance of elements Find relative weight that can be used to find its atomic mass Stoichiometry 10272014 Stoichiometry 10272014 Mole unit of measure of number of units as compared to the number of carbon atoms in exactly 12g of carbon 12 1 moe 602241511 x 1023 units 1 amu 1661x 1024 g 1 Amu 1gmol Percent Composition To determine the elemental composition of a compound we can compare the weight of the compound to the weights of the elements making up the compound This pre supposes that one can actually decompose a compound Two types of decomposition processes To elements H2CO2 To simpler well characterized substancesH2O CO2 Schematic of a system Find mass of C by multiplying mass of CO2 by percentage mass of carbon Finding molecularempirical formula from percentage of elements find mass assuming 100g of sample find moles Stoichiometry 10272014 divide moles of each element by the element with the fewest moles turn it into a ratio and it becomes the empirical formula divide the empirical formula by the molecular formula Chemical Equations One way to view a chemical reaction is that it involves the reorganization of atoms from a set of initial elements andor compounds to a final set of elements andor compounds Initial elementscompounds reactants Final elementscompounds products Ex CH4 202 CO 2 2H2O Reorganization involves the breaking and forming of chemical bonds between atoms Mass is conserved All atoms present in the reactants must show up in the products A properly written chemical equation describing a reaction will be balanced That is be consistent with the conservation of mass Chemical equations provide two pieces of information The composition of the species The physical state of the species Solid s Liquid I Gas 9 Aqueous Aq Stoichiometry 10272014 States are important in thermodynamic calculations Balancing Chemical Equations Basic idea given conservation of mass the same number of atoms must be present on both reactants and products side of a equation i Step 1 Identify the most complicated or largest molecule ii Step 2 Balance the atoms in the most complicated molecule that appears in only one species on each side of the arrow iii Step 3 Balance the remaining atoms iv Step 4 Multiply by an integer to get rid of fractional coefficients Finding limiting reagents etc Aqueous solution Water based Tendency to work with common materials and under common conditions based on the world around us Pressure atmospheric Solvent water alcohols hydrocarbons H20 has a bent 105 structure A solution is a mixture of a solute being dissolved and a solvent the medium When ionic compounds dissolve in water they separate into ions that can move independently of each other NaCls Na Cl When molecular compounds dissolve in water the atoms in the molecules do not separate from each other Stoichiometry 10272014 HCH2CH2OH HOCH 2CH2OH In making an aqueous solution involving an ionic compound the interaction of charged species in the compound with the electron distribution in water results in solvation How much solute dissolves The material usually liquid we want to dissolve something in is called the solvent The material that we want to dissolve is called the solute Arrhenius How much can we measure the extent to which a solute dissolves Idea the greater the extent to which a solute dissolves the greater the concentration of ions in the solution Hypothesis the greater the ionic concentration the more electrically conductive the solution will be His bright idea can be tested by looking at the brightness of a light bulb where current passes through a solution Note BaSO4 does not give ions Strong electrolyte a large amount of ionization Weak electrolyte a few ionization Non no ions present Concentration Molarity Molarity moles of solute liters of solution Stoichiometry 10272014 Dilution Dilution is the process of making a solution of lower molarity from a solution of higher molarity Making a solution of molarity X from molarity Y where molarity X lt molarity Y First step determine the moles of solute in the final solution Second step determine the volume of the stock solution required to deliver the same amount of moles as the final solution Third step add water until we reach the volume of final solution wanted Precipitation Precipitation synthesis of ionic solid A solid precipitate forms when aqueous solutions of certain ions are mixed sometimes when we mix two solutions an insoluble solid will form the solid called a precipitate or insoluble salt is insoluble in water It is so insoluble that when its component ions find each other in solution they get locked together in large clumps driving the reaction towards products Rules All Nitrate ions are soluble Most salts containing the alkali metal ions Lit Nat Kt and ammonium are soluble Most Chloride Bromide and Iodide are soluble Exceptions are Agt Pb2 Hg22 Most sulfate ions are soluble Exceptions are BaSO4 PbSO4 Hg2SO4 and CaSO4 Most hydroxide salts are slightly soluble Most soluble are NaOH and KOH BaOH2 SrOH2 CaOH2 are marginally soluble Stoichiometry 10272014 Sulfie S2 Carbonate C032 Chromate CrO4239 Phosphate PO43 are only slightly soluble Equa ons Conventional Equation a bookkeeping of all species present and arranged for charge neutrality BaNO32aq Na2SO4aqBaSO 4s 2NaNO3aq Complete ionic Equation all aqueous species are split up into their ionic ions Ba 2NO339 2Na SO42 BaSO 4 2Na 2NO3 Net ionic equation indicates exactly the chemical change that occurs nothing more Ba2aq SO4239aqBaSO 4s Notice in the previous example that some ions appears as solvated species on both the reactants and product sides of the chemical equation Ions that appear on both the reactant and product side are referred to as spectator ions Identifying spectator ions allows for one to focus on the chemistry of interest Consistent with this focus precipitation reactions are generally written as Net Ionic Equa ons Precipitation Synthesis of an ionic solid A solid precipitate forms when aqueous solutions of certain ions are mixed Acid Base proton transfer reactions Acid donates a proton to a base forming a molecule water or another weak acid and an aqueous salt Acid proton donor base proton acceptor Stoichiometry 10272014 Bronsted Lowry Theory acidbase reaction are proton transfer process Acid is proton donor Base is proton acceptor When an acid gives its proton to water water is acting as a base Strong acids strong electrolyte undergo complete ionization HCI splits up into ions Oxyacid HNO3 H2804 Group 7 and period 3 and beyond HCI HBr HI Two chlorine oxyacids HCIO3 HCIO4 Any other acids will be weak Polyprotic acids lose only their first Ht easily Weak acids weak electrolyte undergo incomplete ionization For HF only a very few H and Ft ions exists in solution the reverse reaction dominates the chemistry They are like insoluble salts they don t dissociate very much Oxidation reduction electron transfer reaction Electron transfer from one species to another causing a change in the oxidation state of the two species Historically reactions in which oxygen atoms are transferred from one species to another were classified as oxidation reduction Oxidation adding oxygen to form an oxide Reduction removing oxygen Stoichiometry 10272014 Modern definition of Redox Chemistry eectrons transferred from one species to another Oxidation Loss of electron Reduction Gain of electron Oxidation numbers An accounting of the electrons in a chemical species The oxidation number can be interpreted as the effective charge on a species Oxidation number represents the number of electrons required to produce the charge on a species The rules for oxidation numbers Any elemental substance is O A monatomic ion is its charge Oxygen in a compound is 2 hydrogen in a compound in 1 Monatomic hydride ion Ht is 1 Big idea The oxidation numbers of the atomsions in a species must sum between the reactants and the products Balancing Half Cell equations Step 1 identify and write down the unbalanced 2 ions Step 2 Balance atoms and charges in each 2 reactions Use H2O to balance O and H to balance H assuming acidic media Use e to balance charge Step 3 Multiply each 12 reaction by integer such that the number of electrons cancels Step 4 Add and cancel Step 5 for reactions that occur in basic solution proceed as above At the end add OH to both sides for every H present combining to yield water on the H side Gas Laws 10272014 Gas Laws Different phases of matter Gasses variable shape variable volume completely independent variable Liquid variable shape constant volume independent beneath the surface limited to the volume of the liquid and the shape of the bottom of the container Solid constant shape constant volume same shape only vibrate in position Pressure is force per unit area P FA Barometer The gases in the atmosphere at sea level on a fair day no storms exert quotone atmospherequot of pressure P h g d h is height of column g is acceleration from gravity d is density Gas Laws 10272014 P can be represented as mm Hg millimeters of mercury or quottorrquot 1 atm of pressure 760mm Hg Manometer One can use a manometer quotJ tubequot connected to a gas to measure pressure The First Gas Laws Gases are relatively easy to measure and observe in a laboratory This made the physical properties of gases a popular object of study in the 17 19 centuries Boyle Charles Avogrado and others determined fundamental connections between P V n T for gases Boye s law Boyle studied the connection between P and V of gases T and n held constant PV constant When PV constant this is an ideal gas Real Gases are not ideal gases Ideal systems are approx true at low pressures Charles s Law Charles studied the connection between T and V of gases Note P and n held constant VT constant amp T in Kelvin Avogadro s Law Avogadro studied the connection between V and n of gases P and T held constant Gas Laws 10272014 Volumes of gases that react do so in small whole number ratios Vn constant Ideal Gas Law PV nRT Common R Values 0082057 L atm mol 1 k 1 83145 J mol 1 K1 What is ideality Recall that the molecules in a gaseous substance are very far apart So we can make the assumptions Molecules of an ideal gas do not attract or repel one another The volume of an ideal gas molecule is negligible with respect to the container Especially an ideal gas is a collection of non interacting point particles When would you expect ideality to fail High P molecules get too close start interacting Low T same thing Partial Pressure For a mixture of ideal gases in a container Total pressure the sum of the individual gas pressures Gas Laws 10272014 Mole Fraction and Partial Pressure X nintotal Mole fraction X ratio of the number of moles of a component in a mixture to the total number of moles in the mixture X Pi Ptotal Daton s Law Pi xPtotal Kinetic Molecular Theory KMT The gas laws of Boyle Charles and Avogadro are empirical meaning they are based on observation of macroscopic property These laws offer a general description of behavior based on many experiments The empirical gas laws can tell you what happens to an ideal gas but not why KMT is a theoretical molecular level model of ideal gases which can be used to predict the macroscopic behavior of gaseous system KTM postulates Gas particles are so small that their volume is negligible Gas particles are in constant random motion This motion is associated with an average kinetic energy that is directly proportion to the kelvin temperature of the gas Gas molecules constantly collide with each other and with the container walls The collisions of the particles with the container walls are the cause of the pressure exerted by the gas Collisions are elastic The main ideas you should take from KMT are that we can describe temperature and pressure from a molecular perspective Pressure arises from molecules banging into the container walls Temperature is directly related to the kinetic energy of the gas molecules The more KE they have the greater their temperature Chemical Equilibrium 10272014 Rate of the reverse reaction equals the rate of the forward reaction KbKm CmCb Keq It7tI391aiaftstl tli I jIl limit2 Fittilis NHI 39 39ia I39quot 39quotH39 I i iitattt 1iiI39I1IE39 Ee ttnnt5 iI39iuiiiE1i39iiilr1 quotaJ39i39glE1ll gE liF urnm l 5 It1IsIIiiJJIl if rquoti39rt1IlIII ltl linE3 the reaction does not stop at equilibrium The macroscopic concentrations stop changing but the identities of the reactants and the products are constantly in flux since the concentrations of reactants and products are constant equilibrium we can use these concentrations to define a quantity referred to as the equilibrium constant Recall we are considering chemical reactions that occur in solutions with numerous species present When discussing gases we saw that ideal behavior was recovered at low pressure The same is true for species in solution Ideal behavior is recovered at nfinite dilution The activity of a substance is a measure of the non ideality of its properties The activity of a component of a reaction system can be described as the effective concentration or pressure of that component defined relative to a standard state Standard state at a given temperature is For gases 1 atm For solutions 1 M For pure liquids the state of the liquid itself For pure solids the state of the solid itself ai cico activity of i concentration of i concentration at ideal state Chemical Equilibrium 10272014 ai PiPo O p g pR K 2 IlfaE r equilibrium constant has no units mt each component of K is unit less At a given temperature K is a constant value no matter what the initial concentrations were Equilibrium does not tell you how quickly equilibrium is attained The steps involved in attaining equilibrium If equilibrium will be attained Partial Pressure g Molarity aq Different Types of Equilibrium Keq or K refers to any equil Constant KC species are measured in molL Kp species are measured in pressure Ka refers to acid dissociation Kb refers to base hydrolysis Ksp refers to sparingly soluble salt equilibria Kl refers to formation of complex ions They All Have The Same Form Small K 1030 Chemical Equilibrium 10272014 Intermediate K 104 104 Large K 103 Properties of equilibrium constant K for a reaction written in reverse is the reciprocal of the original reactionKneWKorigina If you multiply a chemical equation by a factor n the new K is simply the old powered by the factor KC KpRT de cb KpRI deltangas Notice all RT39dequota 9aS is doing is providing a conversion factor between concentration 1 and pressure units ONLY LOOK AT GAS MOLECULES Number of solids does not affect the number of gaseous molecules