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General Chemistry Study Guide (CHEM115/116)

by: Cheyenne Beckelman

General Chemistry Study Guide (CHEM115/116) CHEM115/116

Marketplace > Chemistry > CHEM115/116 > General Chemistry Study Guide CHEM115 116
Cheyenne Beckelman
Lawrence University
GPA 3.92
Dr. Stork and Debbert

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This detailed study guide is a compilation of all of the material covered in any general chemistry course, as well as specifically CHEM115 and CHEM116.
Dr. Stork and Debbert
Study Guide
Chemistry, Science
50 ?




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This 62 page Study Guide was uploaded by Cheyenne Beckelman on Thursday January 29, 2015. The Study Guide belongs to CHEM115/116 at a university taught by Dr. Stork and Debbert in Fall. Since its upload, it has received 416 views.

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Date Created: 01/29/15
Chapter 1 Chemistry and Measurement 12915 248 PM A Hypothesis A tentative explanation of some natural phenomena A Law A statement or mathematical equation about a fundamental relationshipregularity in nature A Theory A tested explanation of a relationshipregularity in nature that hasn t or can t be proven Lavoisier s Law of Conservation of Mass No mass is lost during a chemical reaction so the chemicals should weight the same before after and during a reaction 0 Side Note Mass is the quantity of matter in a material whereas weight is a force of gravity on a material So weight can change in different environments elevations planets with no gravity etc whereas mass remains the same always Solid Can t be compressed fixed shape and volume Liquid Fixed volume but no fixed shape Gas No fixed volume or shape will fill any container Physical Change Chemical Change Changes form of matter but not its Changes chemical identity new chemical formulaidentity chemical is formed Physical Property Chemical Property Characteristics that can be observed Characteristics involving chemical without changing chemical identity change Mixture Chemicals that sit next to each other without reacting or having a fixed organizational structure Compound Group of chemicals with fixed structure usually reacted together Homogeneous Mixture solution A mixture that is uniform in its properties throughout given samples Heterogeneous Mixture A mixture that consists of physically distinct parts each with different properties Rules for Writing Significant Figures 1 All nonzero digits are significant For example 34 grams has two significant figures 2 Zeros that are between nonzero digits are significant For example 304 grams has three significant figures 3 Zeros written to the left of all nonzero digits are not significant For example 00034 grams has two significant figures 4 Zeros written to the right of all nonzero digits are only significant if a decimal point is written in the number For example 1000 grams has one significant figure while 10000 grams has five The zeros in the second number indicate that a value can be measured accurately to the nearest tenth of a gram while writing simply 1000 grams indicates that the measurement has been rounded to the nearest thousand grams While both mean the same thing to your calculator they don t mean the same thing to a reader 5 Numbers in scientific notation have the same number of significant figures as the portion of the number that39s before the x 10quotquot part of the number For example 430 x 105 grams has three significant figures Rules for Using Significant Figures in Calculations 1 When adding or subtracting the answer should have the same number of figures to the right of the decimal as the value with the fewest decimal places For example 34 5023 8423 Round this to 84 because 34 has only one digit to the right of the decimal 2 When multiplying or dividing the answer should have the same number of significant figures as the value with the fewest significant figures For example 1220 x 34870 425414 Round this answer to 4254 because 1220 has only four significant figures Chapter 2 Atoms Molecules and Ions 12915 248 PM Democritus Had the thought of the existence of atoms John Dalton Atomic Theory all matter is composed of small particles called atoms Joseph Thomson Discovered the Electrons using the cathode ray Milikan39s Oil Drop Experiment Calculated the ratio of the electron s mass to its electric charge Earnest Rutherford Gold foil experiment found that the atoms wasn t structure like plum pudding but that the protons were concentrated in the center of the atom James Chadwick Discovered the neutron Atomic Number No of Protons Mass Number Total no of protons and neutrons Atomic Mass Average weight of all isotopes of that element Atomic Mass Unit amu One amu is equal to onetwelfth the mass of a carbon12 atoms Nuclide An atom with a definite atomic no and mass no looks like this 9 2311Na top one is mass no bottom is atomic no Isotopes Atoms whose nuclei have the same atomic number but different mass number same no of protons different no of neutrons Inorganic Compounds Composed of elements other than carbon Hydrocarbons Compounds that only contain hydrogen and carbon Period The elements in any one horizontal row of the periodic table ie Li Be B C N O F Ne Group The elements in any one column of the periodic table ie H Li Na K Rb Cs Ft Metalloid or Semimetal An element having both metallic and nonmetallic properties ie Silicon and Germanium Anion A negatively charged ion Cation A positively charged ion Functional Group The reactive portion of a molecule that undergoes predictable reactions Molecular vs Structural Formula Molecular formulas give the exact number of different atoms of an element in a molecule whereas structural formals show how the atoms are bonded to one another in a molecule Naming Monatomic Ions Old Version add ous or ic to the stem of an element ie Mn2 is called Manganese II normally but it used to be called Manganous Fe3 is called Iron III but it used to be called Ferric Naming Binary Compounds Binary Compound A compound containing only 2 elements The Polyatomics MUST MEMORIZE TABLE 25 I Some Common Polyatomic Ions Name Formula Name Formda Mcrcuryd or mercurous Hg2 Pcrmanganatc MnO Ammonium Nilf Nitrite N02 Cyanide CN Nitrate N03 Carbonate 3032 Hydroxide 0H llydrogen carbonate or bicarbonate HCO3 Peroxide 02239 Acetate Czngo Phosphate POf Oxalate C2042 Monohydrogen phosphate PO2 quot Hypochlorite ClO Dihydrogcn phosphate H2PO4 Chlorite Clo2 Sul te sof Chloratc ClOf Sulfate 802 Perchlorate C104 Hydrogen suliite or bisul le H803quot Chromate CrOf Hydrogcn sulfate or bisulfate HSO4 Dichromate 02072 39 Thiosulfate 32032 Add an ide to the last element in the formula and use the Greek prefixes before all the elements to symbolize how many atoms there are ie 1 Mono 2 Di 3 Tri 4 Tetra 5 Penta 6 Hexa 7 Hepta 8 Octa 9 Nona 10 Deca 12 Dodeca Oxoacid An acids containing hydrogen oxygen and then other elements MUST MEMORIZE I ABLE 21 I Some Oxoanlons and Their Corresponding Oxoadds Oxoanion Oxoacid cof Carbonate ion H2C03 Carbonic acid NOf Nitrite ion HNO Nitrous acid N03 Nitrate ion HNO Nitric acid PO43 Phosphate ion H3PO4 Phosphoric acid 509 Sulfiie ion st03 Sulfurous acid 8042 Sulfare ion HZSO Sulfuric acid ClO llypochlorire ion HClO Hypochlorous acid CIOZ Chloriie ion IICIOZ Chlorrms acid ClOf Chlorate ion HC103 Chloric acid C104 Perchlorate ion HC104 Perchloric acid Hydrate A compound that contains water molecules weakly bound in its crystals Chapter 3 Calculations with Chemical Formulas and Equations 12915 248 PM 4 Steps For Balancing Equations 1 Write a Word Equation after some experience you can move straight to this stage eg sodium water gt hydrogen sodium hydroxide The arrow can be found in MicrosoftWord under Insert then Symbol 2 Put in Correct Formulas in Place of The Words eg Na H2O gt H2 NaOH Please use subscript and superscript notation from the Toolbar Add them if you don t have them 3 Balance the left hand side number of atoms with those on the right hand side This is done by adding COEFFICIENTS the big numbers eg 2Na 2H2O gt H2 2NaOH Total Na 2 total H 4 total 0 2 The space in front of the H2 of course IMPLIES 1 4 Add state symbols to show the physical state of the substance ssolid lliquid ggas aqaqueous solution eg 2Nas 2H2Ol gt H2 9 2NaOH aq Side Note You can determine the state of a substance by knowing it s bond type ie Ionically bonded chemicals are usually solid because they have extremely high melting and boiling points whereas covalent molecules are usually liquid or gasses because of they re weaker bonds Molecular Formula States the exact number and type of atoms present in a molecule Empirical Formula The formula of a substance written with the smallest integer whole number subscripts Side Note For Ionic formulas the empirical formula usually is the molecular formula as well However for covalent molecules the formulas are different more often Molecular Mass The sum of the atomic masses of all the atoms in a molecule of the substance Formula Mass The sum of the atomic masses of all atoms in a formula unit of the compound Mass Percentage Mass of A in the whole mass of the whole x 100 Limiting Reactant The reactant tat is entirely consumed when a reaction goes to completion Theoretical Yield The maximum amount of product that can be obtained by a reaction from given amounts of reactants The Percentage Yield The actual yield expressed as a percentage of the theoretical yield MOLES Possible equations that have to do with moles Moles mass 9 Molar mass gmol aka mM Moles Volume L x Molarity for solutions Moles STPv 224 or ATPv 24 for gasses PV nRT 1 mole Avogadro39s Number 602 x 1023 Number of atoms in a 129 sample of Carbon12 Mole The quantity of a given substance that contains as many molecules or formula units as the number of atoms in exactly 129 of Carbon12 Molar Mass The mass of one mole of the substance gmol number given at the bottom of the element on the periodic table Side Note The amount of moles of a MOLECULE or COMPOUND is equal to the amount of moles of whichever element has a subscript of 1 Stoichiometry The coefficients and subscripts in formulas and equations You can use stoichiometric ratios in formulas and equations to calculate molesmassgas vol and solution info Chapter 4 Chemical Reactions 12915 248 PM Electrolyte A substance that dissolves in water to give an electrically conducting solution Nonelectrolyte A substance that dissolves in water to give a non conducting or very poorly conducting solution Strong Electrolyte An electrolyte that exists in solution almost entirely as ions most ionic solids that are dissolves in water ie NaCl 9 Na aq Cl39 aq Weak Electrolyte An electrolyte that dissolves in water to give a relatively small percentage f ions generally molecular substances the products react with each other again to produce the reactants Solubility Rules MUST MEMORIZE Rule DJ J Applies to li39Na39llt39Nll4 Cl Br l so2 Co2 P043 8 0 Statement Group A and ammonium compounds are soluble Acetates and nitrates are soluble Most chlorides bromides and iodides are soluble Most sulfates are soluble Most carbonates are insoluble Most phosphates are insoluble Most sul des are insoluble Most hydroxides are insoluble Exceptions AgCl ngClg PbClg AgBr HgBrg ngBrg PbBrg Ag Hgig Hgglg Pb CaSOJ SrSOJ BaSOJ A9804 143804 Pb804 Group IA carbonates NH43CO3 Group IA phosphates NH4KPO4 Group A sullides Nll43S Group A hydroxides all3 SrH3 BaH3 Side Note If something isn t soluble that means that it precipitates Use this knowledge to determine whether or not DoubleReplacement Reactions OCCUF Molecular Equation A chemical equation in which the reactants and products are written as if they were molecular substances even though they may actually exist in solution as ions ie C6OH2 aq N62CO3 aq 9 C6CO3 s 2NaOHaq Complete Ionic Equation A chemical equation in which strong electrolytes are written as separate ions in the solution ie C62aq 20Haq 2Naaq CO32aq 9 C6CO3 s 2Naaq 20H aq Spectator Ion An ion in an ionic equation that does not take part in the reaction You can cancel these ions form both side of the equation to express the essential reaction that occurs Net Ionic Equation An ionic equation from which spectator ions have been canceled only shows the ions that react in the equation Precipitate An insoluble solid compound formed during a chemical reaction in solution Exchange doublereplacement Reaction A reaction between compound that when written as a molecular equation appears to involve the exchange of parts between the two reactants ie MgClz 2AgNO3 9 2AgCl MgNO32 Side Note These are the reactions in which you predict precipitates If no precipitate is formed then no reaction takes place and so you write NR on the product side of the equation Acids and Bases AcidBase Indicator A dye used to distinguish between acidic and basic solutions by mean 5 of the color changes it undergoes in these solutions Acid A substance that produces hydrogen ions H when it dissolves in water Strong ionizes completely in water strong electrolyte Weak only partly ionizes in water weak electrolyte leave as complete formulas when reacted Strong Acids HCI HBr HI HNO3 H2504 HCO4 Split into separate ions when reacted Base A substance that produces hydroxide ions OH39 when it dissolves in water Strong Present in aqueous solution entirely as ions one of which is CH strong electrolyte Weak only partly ionized in water weak electrolyte leave as complete formulas when reacted Strong Bases Metal Hydroxides Split into separate ions when reacted Neutralization Reaction A reaction of an acid and a base that results in an ionic compound salt and possibly water Salt The ionic compound that is a product of a neutralization reaction Polyprotic Acid An acid that yields two or more acidic hydrogens per molecule ie phosphoric acid PO4339 Strong Acid Strong Base 9 Salt Water Acid Carbonate 9 C02 Acid Sulfite 9 502 Acid Sulfide 9 H25 Oxidation and Reduction The idea of Oxidation numbers is to keep track of electrons in a reaction You can determine whether electrons have been transferred from one atom to another Oxidation Number The actual charge of the atom if it exists as a monatomic ion or a hypothetical charge assigned to the atom in the substance by simple rules OxidationReduction Reaction Redox A reaction in which electrons are transferred between species or in which atoms change oxidation number TABLE 45 I Rules for Assigning Oxidation Numbers Rule Applies to Statement 1 Elements The oxidation number of an atom in an element is zero 2 Monatomic ions The oxidation number of an atom in a monatomic ion equals the charge on the ion 3 Oxygen The oxidation number of oxygen is 2 in most of its compounds An exception is O in 0 and other peroxides where the oxidation number is r 39l 4 Hydrogen The oxidation number of hydrogen is l l in most of its compounds The oxidation number of hydrogen is v l in binary compounds with a metal such as Cal3 5 Halogens The oxidation number of uorine is l in all of its compounds Each of the other halogens Cl Br I has an oxidation number of l in binary compounds except when the other element is another halogen above it in the periodic table or the other element is oxygen 6 Compounds and ions The sum ol the oxidation numbers ol the atoms in a compound is zero The sum ol the oxidation numbers of the atoms in a polyatomic ion equals the charge on the ion Oxidation The halfreaction in which there is a loss of electrons by a species or an increase of oxidation number of an atom Reduction The halfreaction in which there is a gain in electrons by a species or a decrease in the oxidation number of an atom Oxidizing Agent A species that oxidizes another species it is itself reduced Reducing Agent A species that reduces another species it is itself oxidized illutinn V n o 7 o quot n l cs39 ju ulluql FORUMl 3115 ruluumg mulling agent ugctll reduction Combination Reaction A reaction in which two substances combine to form a third substance Decomposition Reaction A reaction in which a substance decomposes into two or more other substances Displacement Reaction SingleReplacement Reaction A reaction in which an element reacts with a compound displacing another element from it Double Replacement Reaction A reaction in which two compounds react with one another and they switch their chemical partners this reaction only takes place if one of the products is a precipitate Combustion Reaction A reaction in which a substance reacts with oxygen usually with the rapid release of heat to produce a flame Molar Concentration Molarity The moles of a solute dissolves in one liter cubic decimeter of solution Molarity Moles of solute liters of solution Diluting Solutions M x Vi Mf x Vf Mmoarity Vvoume This equation can be used because the moles of solute don t change when you dilute a solution Quantitative Analysis The determination of the amount of a substance or species present in a material Gravimetric Analysis A type of quantitative analysis in which the amount of a species in a material is determined by converting the species to a product that can be isolated completely and weighed mass precipitation Titration A procedure for determining the amount of substance A by adding a carefully measured volume of a solution with known concentration of B until the reaction of A and B is just complete Volumetric Analysis A method of analysis based on titration molarities titration moesvoume X molarity Dilution how to You buy XL of 12M HCI You need to dilute it so that you have 2L of 2M HCI How much HCI and water do you need to mix together USE nCV 2L x 2M 4 moles of HCL 4 moles 12M x liters of 12M HCI 412 liters of 12M HCI 12M HCI 13L 2L 13L 53L You must mix 53L of water with 13L of 12M HCI Other Reaction Patterns Hydrocarbon CxHy 029 9 C02g H20g Carbonate X1CO3 Add HXz 9 Salt X1X2 C02 H20 Oxidation vs Neutralization Oxidation reactions transfer elections whereas neutralization reactions transfer protons Chapter 5 The Gaseous State 12915 248 PM Pressure The force exerted per unit area of surface Pascal Pa SI unit for pressure Barometer A device for measuring the pressure of the atmosphere mmHg millimeters of mercurytorr A unit of pressure equal to that exerted by a column of mercury 1mm high at 000 degrees Celsius Atmosphere atm Unit of pressure exactly 760 mmHg 1 atm 760 mmHg 101325 Pa 760 Torr Boyle39s Law If the temperature is kept the same then the pressure and volume of a gas are inversely proportional concept PV constant P1V1 P2V2 Charle39s Law If the moles and pressure are kept the same then the temperature and volume are directly proportion V1T2 V2T1 Combined Gas Law The volume occupied by a given amount of gas is proportional to the absolute temperature divided by the pressure PVT constant B i P i 39I 391quot Avogadro s Law The volume of one mole of any gas at standard temperature and pressure is 224 L molar volume n VSTp224 Molar Gas Volume Vm The volume of one mole of gas Standard Temperature and Pressure STP When the conditions are 0 degrees Celsius and 1 atm pressure Ideal Gas Law The law that combines all other gas laws PVnRT The R39 is the constant it varies depending on the other units used in the equa on Molar Gas Constant in Various Units Value of R 0083053 lalmquot Kmnl39y 33 I45 JI39K IHU1quotquot SANS kgnIWs FKnml 33l45kPwdnf Kwnoh IHHTZ cul KIl1ilquot Gas Density PMM dRT MM is Molar Mass d is density quot7 quot3 Pl R39I39 Ul P911 R39l39 Aim But mV equals the density d Substituting this giws PwudRT Partial Pressure The pressure exerted by a particular gas in a mixture Dalton s Law of Partial Pressures The sum of the partial pressure of all the different gases in a mixture is equal to the total pressure of the mixture Total Pressure sum of partial pressures Mole Fraction The fraction of moles of a component in the total moles of gas mixture Mole Fraction of A nAn PAP Collecting Gases over Water When gas is bubbled through water it catches some water vapor as it bubbles out the top and so the water vapor exerts some partial pressure the pressure exerted depends on the temperature of the water the gas is bubbled through TABLE 56 I Vapor Pressure of Water at Various Temperatures Pressure Temperature quot 0 mmHg 0 46 IO 92 I5 I28 I 7 I45 l I65 2 I87 23 2 I I 25 233 27 267 30 393 I 8 40 553 60 1494 80 355l 7600 Kineticmolecular Theory A gas consists of molecules in constant random motion Ek 12m x speed2 Ek is Kinetic Energy m is for mass Kinetic Theory Postulates Gases are composed of molecules whose size is negligible compared with the average distance between them so you can ignore the volume occupied by the gas molecules and focus on the volume occupied by the gas itself Molecules move randomly in straight lines in all directions and at various speeds so properties such as pressure will be the same in all directions The forces of attraction or repulsion between two molecules intermolecular forces in a gas are very weak or negligible except when they collide so a gas molecule will continue moving constantly in a straight line until it collides with another gas molecule or wall of container When molecules collide with one another the collisions are elastic so when two gas molecules collide the total kinetic energy between the two stays constant The average kinetic energy of a molecule is proportional to the absolute temperature so the higher the temperature the greater the molecular kinetic energy 3R r quot V w rms speed molecular speed umolecular speed Mmmolar mass Rmolar gas constant R831 kgm2szKmol Ttemperature Gaseous Diffusion the process in whish a gas spreads out through another gas to occupy the space uniformly Gaseous Effusion The process in which a gas flows through a hole that is so small that only one molecule can go through it at one time Graham39s Law of Effusion the rate of effusion of gas molecules from a particular hole is inversely proportional to the square root of the molecular mass of the gas at constant temperature and pressure Rule of effusion of molecules 1 3quot Side Note The volume of gas produced is directly related to the rate of effusion ratios are the same between two gasses Van der Waals Real Gas Equation V becomes V nb and P becomes I a P39 439u39 nbl HR39I39 Pn2aV2 so the complete equation is TABLE 51 I van der Waals Constants for Some Gases a b Gas L atmmol2 Lmol C03 3658 004286 CgHb 5570 006499 CgHSOH l256 0087 I0 Hc 00346 00238 Hz 02453 002651 03 L382 003 I 86 so2 6865 005679 H20 5537 003049 Thermochemistry 12915 248 PM Kinetic Energy The energy associated with the motion of an object EK12mv2 EK Kinetic Energy mmass vvolume Joule The SI unit of energy kg x m2s2 Watt a measure of the quantity of energy used per unit of time 1 watt 1 joule per second Calorie A nonSI unit of energy defined as the amount of energy required to raise the temperature of one gram of water by one degree Celsius 1 calorie 4184 Joule Potential Energy The energy associated with the position of an object how much kinetic energy could it potentialy exert Internal Energy U the sum of the kinetic and potential energies of the particles making up a substance ETOTEKEPU Law of Conservation of Energy Energy may be converted from one form to another but the total quantity of energy remains constant Thermodynamic System just System The reaction substances in the reaction Surroundings Everything that s in the vicinity of the thermodynamic system Heat q The energy that flows into or out of a system because of a difference in temperature between the thermodynamic system and its surroundings If the system absorbs heat q is positive and vice versa Thermal Equilibrium Equality between the temperature of the system and the surroundings Heat of Reaction The value of q required to return a system to the given temperature at the completion of the reaction Exothermic A Chemical reaction or a physical change in which the system emits evolves heat q is negative Endothermic A chemical reaction or a physical change in which the system absorbs heat q is positive Enthalpy H An extensive property of a substance that can be used to obtain the heat absorbed or emitted in a chemical reaction it is a state function meaning that it doesn t matter how the change occurred only the initial and final state of the system Side Note An extensive property is a property that depends on the amount of substance ie mass and volume II II lltprmlmtxl llcrcucluntxu II Enthalpy of Reaction The change in enthalpy for a reaction at a given temperature and pressure Thermochemical Equation The chemical equation for a reaction including phase labels in which the equation is given a molar interpretation and the enthalpy of reaction for these molar amounts is written directly after the equation 339lxl llll l 3l39lllml Hug394 Il ln lxl Side Note The is the k emittedabsorbed PER EQUATION meaning per number of moles specified in the equation So to find the k of heat find the moles of the specified substance and then divide the by that value Heat Capacity C The quantity of heat needed to raise the temperature of the heat required to change the temperature from initial temp to final temp the sample of substance one degree Celsius or one kelvin Specific Heat Capacity The quantity of heat required to raise the temperature of one gram of a substance by one degree Celsius or one f gt17 1 I DI 39 kelvin I x n A qthe heat required to raIse the temperature of a sample s specific heat capacity Calorimeter A device used to measure the heat absorbed or emitted during a physical or chemical change Hess39s Law For a chemical equation that can be written as the sum of two or more steps the enthalpy change for the overall equation equals the sum of the enthalpy changes for the individual steps AHreaction AH1 AHZ AH3 Standard State AH Standard thermodynamic conditions 1 atm pressure and the specifies temperature usually 25 C Allotrope One of two or more distinct forms of an element in the same physical state Reference Form The stablest form of an element under standard thermodynamic conditions Standard Enthalpy of FormationHeat of Formation AH f The enthalpy change for the formation of one mole of a substance in its standard state from its elements in their reference forms and in their standard states listed in the back of the textbook FOR AH FORMATIONS ONLY AHf sum of AHf products sum of AHf reactants Quantum Theory of the Atom 12915 248 PM The Wave Nature of Li Visible light Xrays and radios waves are all forms of electromagnetic radiation Wavelength K The distance between any two adjacent identical points of a wave Frequency Hz The number of wavelengths of that wave that pass a fixed point in one unit of time usually a second wavelength cspeed of light 300 X 108 ms vfrequency Electromagnetic Spectrum The range of frequencies or wavelengths of electromagnetic radiation III quot IIIquotquot IIIquotquot IIIquot III39 IIIquot III IIIquot IIIquot IIIquot III39 Frequency Isquot I l I I I I I I I I c F quot NW N r W FM AM Mmmd X rzm ultra ultra t I M Micmwaws Rdd dl 39 my 39 r InlIuI39cd Inlrurcd RadIo I39aws 39 Iulcl IIIlcl l l l 7 III III quot III III III III III III III III III III wanlength In 1 pm I ll pm I I lllll pm I nmI I H mm I 100mm I Him I39 lllpm I 100 quotII I mm I ll mm I I l mm quot I Visible spectrum 35H 11 Jill 50H 551 JR ISU 7U 75 Hi lllll Planck39s Quantization of Energy Planck39s Constant h A physical constant relating energy and frequency having the value 663 X 103934 Js PhotonsQuanta Particles of electromagnetic energy I hr Eenergy hPlanck s constant vfrequency Photoelectric Effect The ejection of electrons from the surface of a metal or from another material when light shines on it Continuous Spectrum A spectrum containing light of all wavelengths like a rainbow we can attain a continuous spectrum by shining light through a prism Line Spectrum A spectrum showing only certain colors or specific wavelengths of light which is what appears when a gas is heated Bohr39s Postulates Energy Level Postulate An electron can only have specific energy values in an atom called energy levels So an atom can only have specific total energy values R l n N II to w RH2179 x 1018 J Transitions Between Energy Levels An electron in an atom can change energy only by going from one energy level to another energy level aka undergo a transition llIncrg of cmiucd photon hr AI39 lf39 ll In general the energy of the emitted photon hv equals the positive energy lost by the atom AE Quantum Mechanics De Broglie Relation mmass Side Note De Broglie39s thinking was as follows since light behaved as both a particle and a wave particles of matter should show characteristics of waves under the proper circumstances Quantum MechanicsWave Mechanics The branch of physics that mathematically describes the wave properties of submicroscopic particles Heisenberg s Uncertainty Principle The momentum and position of a particle cannot be known at the same time because measuring one will change the other The product of the uncertainty in position and the uncertainty in momentum of a particle can be no smaller than Planck s constant divided by 4pi The more we know about the position of an electron the less we know about its momentumspeed And vice versa Atomic Orbital A wave function for an electron in an atom Quantum Numbers these numbers are assigned so that you can refer to a specific electronpair of electrons Principal Quantum Number n The number of a certain quantum energy level shell it can have any positive value 1 2 3 etc The smaller n is the lower the energy of the electrons in that shell and the smaller the shell is and vise versa Shells are sometimes designated by the lullmving letters Letter K I M N n l 2 3 4 Angular Momentum Quantum Number f The number of a certain subshell within a quantum energy level each subshell has a different shape and represents an area in which electrons are likely to be found it can have any integer value from 0 to n 1 iven shell 39l39he dill erenl subshells are usually denoted by letters as follows Letter p t f It I l l 2 3 4 Side Note n values that are about 4 and values that are above 3 can exist but are not in the grounded state Magnetic Quantum Number ml The quantum number that distinguishes orbitals of given n and aka that have the same energy and shape by representingshowing its different orientation in space it can be any integer from to ie For 0 5 subshell mfcan only be 0 For 1 p subshell mfcan be 1 O and 1 Spin Quantum Number ms The quantum number that refers to the two possible spins of the electron possible values are 12 and 12 only TA B L E 7 I n l l 2 2 l 3 3 l 3 2 4 4 l 4 2 4 3 I Permissible Values of Quantum Numbers for Atomic Orbitals SubsheH m Notation 0 Is 0 25 39l 0 1 2p 0 3s I 0 1 3p 2 39l 0 1 2 3d 0 49 I 0 1 4p ll2 4d 3 2 1 l 2 3 Number of Orbitals in the Subshe IKJIL J Lll39uJ DJ Electron Configurations and Periodicity 12915 248 PM Electron Configuration A particular distribution of electrons among available subshells lists the subshell symbols and after another with a superscript giving he number of electrons in that subshell Orbital Diagram A diagram to show how the orbitals of a subshell are occupied by electrons orbitals are represented by circles labeled by its subshell notation arrows inside the circle represent the electrons and their direction represent their mS value Pauli Exclusion Principle An orbital can only hold at two electrons at most and only if those electrons have opposite spins they must have different mS values Number of Maximum Number Subsh ell rbitals of Electrons s39 I 0 l 2 p I I 3 6 l 2 5 u 3 7 394 Aufbau39s Principle When writing electron configurations you must assign electrons to each subshell in a specific order ls Zs 2p 3s 3p 45 3d 4p 55 4d 5p 65 4f 5d 6p 75 5f However when writing electron diagrams fill in the shells in this order but write them with the same numbers grouped together Cr and Cu are the exceptions to this principle NobleGas Core An innershell configuration corresponding to one of the noble gases the beginning of the configuration can be substituted with one of the noble gases such as Helium Neon Argon and Krypton PseudoNobleGas Core The noblegas core together with n1d10 electrons in the configuration Valence Electron An electron in an atom outside the noblegas or pseudo noblegas core these electrons are primarily involved in chemical reactions are shown in the last orbital of a configuration if two chemical s valence shell configurations are similar then their chemical properties are probably similar as well When finding the valence electrons of an electron configuration write any incomplete shells as well as the biggest numbered shell ie 3d8 452 sblock clcmcnls p hlock clcmcnls lhlock clcmcnls 39 block clcmcnts 5f For La write 4f before 5d Hund39s Rule While drawing orbital diagrams you must place electrons into separate orbitals of the same subshell with the same spin before pairing 0 O 7 r the electrons with opposite spinned ones Iquot quot5 7 Paramagnetic Substance A substance that is weakly attracted by a magnetic field and this attraction is generally the result of unpaired electrons Diamagnetic Substance A substance that is not attracted by a magnetic field or is very slightly repelled by such a field This property generally means that the substance only has paired electrons Mendeleev made a periodic table of all the elements known at the time However he also predicted where elements discovered in the future would go Periodic Law When the elements are arranged by atomic number their physical and chemical properties vary periodically Size of BallAtomic Radius H I 0 1M li Be 3 O 0 Na Mg quot o Iill Atomic radius increases Penod Vii H Hi Will VIIIA He 0 15 ML YA 39IA 391 B C N 0 F Ne o O 0 0 0 A1 Si P S C At 0 VIII In 1m 0 o o o 0 Ge As Sc Br K 0 0 0 0 Sn TC I Xe 0 0 o o o Pb Bi Po Al Rn 0 0 o o 0 Atomic radius decreases Effective Nuclear Charge The positive charge that an electron experiences from the nucleus equal to the nuclear charge but reduced by any shielding or screening from any intervening electron distribution First Ionization EnergyPotential The minimum energy needed to remove the highestenergy the outermost electron from a neutral atom in the gaseous state Side Note When removing more than one electron the ionization energy increases as more electrons are removed from an atom ie if 3 electrons are removed the first ionization energy will be less than the second which will be less than the third SizeIonization Energy 3 is 2 35536 3 333355 2 3333 Zquot 333333 Electron Affinity The energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion Basic Oxide An oxide that reacts with acids most metal oxides are basic oxides Acidic Oxide An oxide that reacts with bases most nonmetal oxides are acidic oxides Amphoteric Oxide an oxide that has both basic and acidic properties Ionic and Covalent Bonding 12915 248 PM Ionic Bonding Ionic Bond A chemical bond formed by the electrostatic attraction between positive and negative ions one or more valence electrons are transferred between atoms Cation Positive ion Anion Negative ion Lewis electrondot symbol A symbol in which the electrons in the valence shell of an atom or ion are represented by dots placed around the letter symbol of the element TABLE 9 I Lewis ElectronDot Symbols for Atoms of the Second and Third Periods IA HA quotIA VA VA VIA VllA VIIIA Period ns39 ns2 ns npl n57in ns np ns np4 ns np ns npquot Second li Bc 1 o quot39c 39l39lnrd Nd39 Mg39 39 Il 1 1 I I 413 Q I r Coulomb39s Law kconstant 899 x 109 JmC2 C is coulomb Q1electric charge of one atom Q2electric charge of other atom rdistance m The more valence electrons atoms have ie 2 or 2 instead of 1 or 1 the more energy it requires to pull them apart this explains why elements that have 1 or 1 are usually soluble whereas things with 2 or 2 charges are usually insoluble The bigger the particles are the less energy it requires to pull them apart Lattice Energy The change in energy that occurs when an ionic solid is separated into isolated ions in the gas phase ie NaCls 9 Nag CIg E639kJ Things that have small lattice energies have big radii and small charges Things that have large lattice energies have small radii and large charges BornHaber Cycle A method of determining the lattice energy of an ionic solid in which you think of the solid being formed from the elements by two different routes and add up the enthalpy changes in the alternative route Direct Nuts gtlltg mu Nd list ll Step 1 Stcp 2 V V Nalfs39 Clix Stcp 3i Slcp 4 Y Step 5 ha tg 39l39tg Nais l Nam AH 1024 kl j39lg39 Liltf39fl All 120k Nat Natty quot 3427 All 400141 Citjvz39i 24172 11341751 All 340 kl SaT39 LilTiff 39a tl H5 Nah 1113quot aCllxl AH 375 k f change for this formation reaction IS 375 kJ Iquot But the enthalpy of formation has been clclcrminctl calorimetrically and cquals ll kJ l lqualing thcsc two val ues we get 375 Ll U 4 kl Solving for 1 yields the latticc energy of 39a l Iquot 375 411kl 786 k Ionic Substance Properties High melting points because ionic solids contain large groups of cations and anions that are attracted to each other Ionic Radius A measure of the size of the spherical region around the nucleus of an ion within which the electrons are most likely to be found Isoelectronic Refers to different species having the same number and configuration of electrons Covalent Bonding Covalent Bond A chemical bond formed by the sharing of a pair of electrons between atoms an overlap of atomic orbitals so that the electrons of one atom become attracted to the nucleus of another atom 90 09 A v Bond dissociation energy Potential energy Distance between nuclei gt Minimum potential energy the lowest point is when the H2 is fully bonded the electrons are attracted to the nucleus of the other atom The graph shoots up at the left because that is when the atoms get so close to each other that their nuclei actually start repulsing each other Lewis ElectronDot Formula A structural formulas using dots to represent valence electrons On the exam you can use lines to represent the bonded pairs but you must use dots to represent the unbounded electrons Bonding Pair An electron pair shared between two atoms LoneNonbonding Pair An electron pair that remains on one atom and is not shared Coordinate Covalent Bond A bond formed when both electrons of the bond are donated by one atom A 3 B A 3 B Octet Rule The tendency of atoms in molecules to have either electrons in their valence shells only two for Hydrogen atoms Single Bond A covalent bond in which a single pair of electrons is shared by two atoms Double Bond A covalent bond in which two pairs of electrons are shared by two atoms Triple Bond A covalent bond in which three pairs of electrons are share by two atoms Polar Bond A covalent bond in which the bonding electrons spend more time near one atom than the other Electronegativity A measure of the ability of an atom in a molecule to draw bonding electrons to itself increasing clcclroncgulivily r y t Decreasing olcclroncvulivit Lr Rf Db Su Bh Hs Mt Ds Rg Uub Uul Uuq Uup Uuh Steps For Writing Lewis ElectronDot DiagramsFormulas 1 Calculate valence elections if ion included charge 2 Make skeleton structures with single bonds from central atom 3 Satisfy the octet rule on outer atoms 4 If there are extra electrons 9 add onto central atom If there aren t enough electrons 9 consider doubletriple bonds Delocalized Bonding A types of bonding in which a bonding pair of electrons is spread over a number of atoms rather than localized between two this is when you would need to write the molecule s resonance structuresdescription Resonance Description You describe the electron structure of a molecule having delocalized bonding by writing all possible electrondot formulas Exceptions to the Octet Rule More than 8 There are numerous molecules in which an atom or two have more than 8 valence electrons ie PF5 and PCI5 Less than 8 These are much less common Some examples include BECIz BF3 and other Nitrogen Oxides Formal Charge A method used to figure out which Lewis Dot Structure is correct In this method for every atom in the structure you follow this equann Formal Charge Valence electrons on free atom 12number of electrons in a bond number of lonepair electrons Bond Order Number of bonds ie 1single bond 2double bond Bond Length Distance between two nuclei Bond Energy Energy required to break a bond in a gaseous state average Example H2 F2 9 2HF AH Breaking HH bond436kJ Breaking FF bond155kJ Making HF bond 565kJ 426kJ 155kJ 2565 539kJ 6 that is the AH When doing these calculations you only need to consider the bonds that have been broken or made Molecular Geometry and Chemical Bonding Theory Electron Pairs Total Bonding Lone e2 4 4ltx 2 Ix Arrangement of Pairs Linear Tri gonal planar 39l elrahedral Molecular Geometry Linear o Tri gonal planar Bent l or Lone pair angular AX 39l39cliahcdral AX Trigonal pyramidal AX Rent or angular BCF 3F1 so CH4 H20 Example r lie F Electron Pairs Arrangement Molecular Total Bonding Lone of Pairs Geometry Example Axial atom il Trigonal oquot S 0 bipyramidal PCI l 1quot MS Cl Axial alum Seesaw f F I or distorted 9139 4 I tetrahedron 5Fquot F AX4 Lonc F pair Ti39iuonal 5 lt gt him ruinidul T h 1 ml iiibl z 39 oi it 3 2 Ax gtLone CIF3 r l39Kj pairs 39 F V P z Linear 39 I 13quot t 2 3 AX 2 Lone Xch Xle fj 39 Lone Pm 39 pair F F 6 O Octahedral SF F 3 F AX r I F F F Square F I F 6 lt 5 I gt claliedrul pyramid4 IFS I 1quot AXi F F 39 39Lone pair 39 39 Lone pair 39quot Square Fl F 4 2 planar XeF4 X AX4 Fquot F Lone pair Valence Bond Theory An approximate theory to explain the electron pair or covalent bond by quantum mechanics 1 An orbital on one atom comes to occupy a portion of the same region of space as an orbital on the other atom the two orbitals overlap 2 The total number of electrons in both orbitals is no more than two Hybrid Orbitals Orbitals used to describe bonding that are obtained by taking combinations of atomic orbitals of the isolated atoms ie when the electrons from an s orbital and a p orbital are combined to create more unpaired electrons in each orbital Bonding Orbitals Molecular orbitals that are concentrated in regions between nuclei Antibonding Orbitals Molecular orbitals having zero values in the region between two nuclei and therefore concentrated in other regions Homonuclear Diatomic Molecules Molecules composed of two like nuclei ie Liz N2 02 EtC Heteronuclear Diatomic Molecules Molecules composed of two different nuclei ie CO and NO States of Matter Liquids and Solids 12915 248 PM Change of StatePhase Transition A change of a substance from one state to another Melting Solid to Liquid Freezing Liquid to Solid Vaporization Liquid to Gas Condensation Gas to Liquid Sublimation Solid directly to Gas Deposition Gas directly to Solid Vapor Pressure The partial pressure of the vapor over the liquid measured at equilibrium at a given temperature molecules with larger molecular masses have lower vapor pressures because larger dispersion forceshigher boiling pointless vapor pressure Boiling Point The temperature at which the vapor pressure of a liquid equals the pressure exerted on the liquid which is the atmospheric pressure unless the vessel containing the liquid is closed Freezing Point The temperature at which a a pure liquid changes to a crystalline solid Melting Point The temperature at which a crystalline solid changes to a liquid I Hi 51mm l l Water and steam p39l Water Ice and tllcr lcmpcmturc 39 39l mu quot lllt ul nlxlml quotII unntuui llllt l Heat of Fusion The heat needed to melt a solid the amount of heat you need to add at the first plateau point for the substance to become a liquid Heat of Vaporization The heat needed for the vaporization o a liquid the amount of heat you need to add at the second plateau point for the substance to become a gas Phase Diagram A graphical way to summarize the condition sunder which the different states of a substance occur Pressure utm39l 0 C RFC Temperature Triple Point The point on a phrase diagram representing the temperature and pressure at which three phases of a substance coexist in equilibrium Critical TemperaturePoint The temperature above which the liquid state of a substance no longer exists regardless of the pressure Critical Pressure The vapor pressure at the critical temperaturepoint MaxwellBoltzman Distribution 5 NYC Relative number of molecules 0 mm mm mm mm Sum Molecular speed ins FIGURE 525 A Maxwell s distribution of molecular speeds The dlSlllbUllUllS of speeds of H molecules are shown f0 UK and 50011 Note that the W W 1 RT speed corresponding lo the maximum m n u the curve the most probable Speed Nc Mm Mm increases Wlll l temperature r T Surface Tension The energy required to increase the surface area of a liquid by a unit amount Viscosity The resistance to flow that is exhibited by all liquids and gasses TABLE 2 Properties of Some Liquids at 20 C Molecular Weight Vapor Pressure Surface Tension Substance amu mmHg me2 Water H30 18 13 x 10 73 x ll 3 lirhuncliuxidc to 41 43 m 12 m Pcnlunc Cell 72 44 gtltI ll lb gtlt ll quot Glycerol C3H503 92 16 m 4 13 m 393 Chlnml orm CHth 11 17 X H2 27 X ll 2 urhun tetrachloride Tl l5l 87 lll 27 ll j mmnmmm cum 253 39 x Inquot 42 gtlt m Intermolecular Forces The forces of interaction between molecules which tend to be weak like Dipoledipole London Dispersion and Hydrogen Bonding Van Der Waals Forces A general term for the intermolecular forces that consist of dipoledipole and London Dispersion forces Dipoledipole Force An attractive intermolecular force resulting from the tendency of polar molecules to align themselves such that the positive end of one molecule is near the negative end of another London Dispersion Forces The weak attractive forces between molecules resulting from the small instantaneous dipoles that occur because of the varying positions of the electrons during their motion about nuclei Hydrogen Bonding A weak to moderate attractive force that exists between a hydrogen atom covalently bonded to a very electronegative atom X and a lone pair of electrons on another small electronegative atom Y Side Note The weakest intermolecular forces are the London Dispersion forces then dipoledipole forces and then hydrogen bonding Molecular Solid Covalent Simple Molecular A solid that consists of molecules held together by intermolecular forces exist as molecules these are the substances that have intramolecular and intermolecular forces Metallic Solid A solid that consists of positive cores of atoms held together by a surrounding seas of electrons metallic bonding are malleable bendable ductile can be deformed like into a thin wire without losing Viscosity kgms 10 71 24 15 58 97 20 ll m l mquot 10 m H i l l 3 toughnessstrength good conductors of electricity and are shiny because of the free electrons Ionic Solid A solid that consists of cations and anions held together by the electrical attraction of opposite charges ionic bonds are poor conductors of electricity until they are dissolved or melted because ionic solids don t contain charge carriers Covalent Network Solid A solid that consists of atoms held together in large networks or chains by covalent bonds do NOT form molecules are usually are hard and have a high melting point TABLE 115 Properties of the Different Types of Solids Hardness and Electrical Type of Solid Melting Point Brittleness Conductivity Molecular Low Soft and brittle Noneondueting Metallie Variable Variable hardness Conducting malleable lonie High to very Hard and brittle Noneondueting solid high conducting liquid Covalent network Ver high Very hard LTstlally noneondueting How do you tell if something is a covalent network or molecular Remember the Covalent Networks Diamond graphite silicon silicon oxide SiOz sandquartz The other nonmetal nonmetal molecules have a molecular structure Solutions 12915 248 PM Solute The gas or solid that is dissolved in a liquid the component in smaller amount Solvent The liquid in which the gas or solid is dissolved the component in greater amount Miscible Fluids Fluids that mix with or dissolve in each other in all proportions as opposed to immiscible Solubility The amount that dissolved in a given quantity of water at a given temperature to give a saturated solution Saturated Solution A solution that is in equilibrium with respect to a given dissolved substance Unsaturated Solution A solution not in equilibrium with respect to a given dissolved substance and in which more of the substance can dissolve Supersaturated A solution that contains more dissolved substance than a saturated solution does if a supersaturated solution is cooled to room temperature the solute will not crystallize within the solution until a crystal is added to the solution from outside Like Dissolves Likequot Substances with similar intermolecular attractions are usually soluble in one another ie polar substances are soluble only in polar substances and non polar substances are soluble only in nonpolar substances This is because when a substance dissolves it breaks bonds between the solvent molecules and between the solute molecules Substances will only dissolve if they can get this energy back form making bonds with the solvent ie oil doesn t dissolve in water because it s nonpolar and wouldn t be able to form bonds with the polar water molecules Hydration The attraction of ions for water molecules favors the dissolving of an ionic solid in water Substances with large lattice energies small ionic radii and large charges have larger energies of hydration because it takes more energy to pull the substance apart Le Chatelier39s Principle When a system in equilibrium is disturbed by a change of temperature pressure or concentration variable the system shifts in equilibrium composition in a way that tends to counteract this change of variable All gases become more soluble in a liquid at a given temperature when the partial pressure of the gas over the solution is increased Henry s Law The solubility of a gas is directly proportional to the partial pressure of the gas above the solution S kHP S solubility kH Henry s law constant P partial pressure of the gas 5 4111 s l or quot 339 fr P 9 P Colligative Properties Properties that depend on the concentration of solute molecules or ions in solution but not on the chemical identity of the solute Ways of Expressing Concentration Molarity moles of solute liters of solution Mass Percentage of Solute mass of solute mass of solution x 100 Molality moles of solute kilograms of solvent Mole Fraction moles of substance A total moles of solution VaporPressure Lowering of a solvent A colligative property equal to the vapor pressure of the pure solvent minus the vapor pressure of the solution AP vapor pressure of pure solvent P A vapor pressure of solution PA Raoult39s Law The partial pressure of solvent PA over a solution equals the vapor pressure of the pure solvent P A times the mole fraction of solvent X A in the solution P PEAquot Sum of both equations AP MWR BoilingPoint Elevation ATb A colligative property of a solution equal to the boiling point of the solution minus the boiling point of the pure solvent M Kb b0ng pomt elevation constant cm molal concentration aka molality FreezingPoint Depression ATf A colligative property of a solution equal to the freezing point of the pure solvent minus the freezing point of the solution A1 Osmosis The phenomenon of solvent flow through a semipermeable membrane to equalize the solute concentrations on both sides of the membrane Osmotic Pressure A colligative property of a solution equal to the pressure that when applied to the solution just stops osmosis 77 l39IR39I39 absolute temperature M molar concentration of solute R gas constant T Rates of Reaction 12915 248 PM Reaction Rate The increase in molar concentration of product of a reaction per unit time or the decrease in molar concentration of reactant per unit time AxT Variables That Affect Reaction Rates Concentration Presence of Catalyst A substance that increases the rate of reaction decreases the activation energy without being consumed in the overall reaction Temperature Surface Area of a Solid Reactant When any of these variables are increased it causes the Reaction Rate to increase because it increases the number of effective collisions between particles so they can react Calculation of the average rate The average rate of formation of 0 during the decomposition of N20 was calculated during two different time intervals When the time changes from 600 s to 1200 s the average rate is 25 X 10 quot39 molLs Later when the time changes from 4200 s to 4800 s the average rate has slowed to 5 X 10 739 molquotLs Thus the rate of the reaction decreases as the reaction proceeds Experimental Determination of Rate Concentration of 02 molL cgt c c c g g g a g g 1200 2400 lk A s 2 i 39V 39 A f f 2 3600 4800 6000 9 2N305g 4NOgtg 028 4 4i oll39 l39lllllll l39ln39 Opening for adding N 539 I v 39 r lVlcrcury manometer Insulated walcr bath 4 Flask containing NIOsl l Nogly and 03m Rate Law An equation that relates the rate of a reaction to the concentrations of reactants and catalysts raised to various powers ie Rate kN02F2 Rate Constant k A proportionality constant in the relationship between rate and concentrations ie k rate N02F2 Reaction Order regarding a specific reactant species The exponent of the concentration of that species in the rate law as determined experimentally ie order for N02 would be 1 because it s exponent in the equation is 1 Overall Order of a Reaction The sum of the orders of the reactant species in the rate law ie overall order for N02F2 is 2 RATE LAWS CAN ONLY BE FOUND BY EXPERIMENT Meaning you need data from an experiment to find the rate law of a reaction Find the exponents of each reactant find out what happens to the rate when you double the concentration triple the concentration half the concentration etc Effect on Rate of Doubling the Initial Concentration of Reactant Rate ls m Multiplied by l l l 2 2 4 m exponent Then you can solve for k by plugging in values Chemical Equilibrium 12915 248 PM Chemical Equilibrium The state reached by a reaction mixture when the rates of forward and reverse reactions have become equal The continuing forward and reverse reactions make the equilibrium dynamic Equilibrium Constant Expression An expression obtained by multiplying the concentrations of products dividing by the concentrations of reactants and raising each concentration term to a power equal to the coefficient in the chemical equation Equilibrium Constant Kc The value obtained for the equilibrium constant expression when equilibrium concentrations are substituted For the following equation aA bB cC dB This is the equilibrium constant expression ICI IDIquot IA l BIquot Law of Mass Action A relation that states that the values of the equilibriumconstant expression Kc are constant for a particular reaction at a given temperature whatever equilibrium concentrations are substituted Forward and Reverse Reaction Ratio Kc Kf Kr Equilibrium Constant Kp An equilibrium constant for a gaseous reaction in terms of partial pressures J 39llll quot39 l39l ll Iv kh39Ri x Can also turn into 9 39 An Sum of gaseous product coefficients sum of gaseous reactant coefficients Homogeneous Equilibrium An equilibrium that involves reactants and products in a single phase ie only involves gaseous reactants and products Heterogeneous Equilibrium An equilibrium involving reactants and products in more than one phase Direction of Reaction Reaction Quotient QC An expression that has the same form as the equilibriumconstant expression but whose concentration values are not necessarily those at equilibrium In order to calculate the direction of the reaction compare the reaction quotient with the equilibriumconstant expression and determine what has to happen to the quotient for it to become the equilibriumconstant expression either decreaseincrease the numeratorproducts or denominatorreactants If the expression has to decrease then more reactants have to exist so the reaction goes to the left and vise versa i Q K lhc rcuclion will go In lhc lcl39l ll Q K lhc rcuclimi will go In lhc righl I Q K the rcuclimi mixiurc is Lil cqunlibrium ICE BOX for using in Equilibrium problems I Initial Concentrations C Change in Concentrations E Equilibrium Concentrations Le Chatelier39s Principle When a system in chemical equilibrium is disturbed by a change of temperature pressure or a concentration the system shifts in equilibrium composition in a way that tends to counteract this change of variable ie pressure concentration temperature or catalyst addition Concentration Change When reactant is added or product is removed the reaction moves left to right and creates more product When product is added or reactant is removed the reaction moves right to left and creates more reactant Pressure Change If the pressure is increased by decreasing the volume of a reaction mixture the reaction shifts in the direction of fewer moles of gas Temperature Change For endothermic reactions products increase as temperature increases For exothermic reactions products increase as temperature decreases Adding a Catalyst Catalysts increase the rate of reaction but are not consumed by it A catalyst has no effect on the equilibrium composition of a reaction mixture IT MERELY SPEEDS UP THE ATTAINMENT OF EQUILIBRIUM Acids and Bases 12915 248 PM Arrhenius Theory Acids are substances that when dissolved in water increase the concentration of hydronium ions H3O Bases are substances that when dissolved in water increase the concentration of hydroxide ions OH39 BronstedLowery Theory Acids are proton donors whereas bases are protons acceptors Conjugate AcidBase Pair Two species in an acidbase reaction one acid and one base that differ by the loss or gain of a proton ie NH4 and NH3 Amphiprotic Species A species that can act as either an acid or a base it can lose or gain a proton Lewis Acid A species that can form a covalent bond by accepting an electron pair from another species Lewis Base A species that can form a covalent bond by donating an electron pair to another species r 1 H H I k H I N 2 H H I N I H H H L l clccln mpuir electronpair acceptor lnnnr Strong Acids HCI HBr HI HNO3 H2504 HCLO4 Strong Bases LiOH NaOH KOH CaOH SrOH BaOH 1A hydroxides mostly Molecular Structure and Acid Strength Larger atomic size stronger acid As you go down a column of elements the size of atom X increases so the HX bonds strength decreases Therefore the compound can dissociate more easily which makes it a stronger acid Higher electronegativity stronger acid As you go across a row of elements the electronegativity increases so the HX bond polarity increases Therefore the compound is a stronger acid Smaller negative charge stronger acid As the negative charge of compound increases there are fewer protons for the compound to donate so it is a weaker acid AcidBase Equilibria 12915 248 PM Hydrogen Ions Protons or hydronium ions H3O The Bronsted Lowry Theory Acids are proton donators and bases are proton acceptors Conjugate AcidBase Pair And acid and a base in a reaction both from the reactants and the products that differ only by 1 proton ie NH3 and NH4 Amphiprotic Substances Substances that can act as both a base and an acid aka they can both donate and accept a proton Ionization of Water H20 9 Haq OH39aq Ionic Product of Water Kw The Kc for water It is ALWAYS equal to 1 x 1014 The effect of temperature on Kw Selfionization of water is an endothermic process So if the temperature increases then the concentration of hydroxide and hydronium ions will increase pH og H pOH log OH39 pH pOH 14 Example Problem Given pH 9 9 pOH 14 pOH 5 Kw H OH39 103914 Example Problem Given H 10393 10393 x OH39 103914 OH39 103911 AcidDissociation Constant Ka concentrations of aqueous products concentrations of aqueous reactants BaseIonization Constant Kb The equilibrium constant for the ionization of a weak base same as Kb but for bases Hydrolysis The reaction of an ion with water to produce the conjugate acid and hydroxide ion or the conjugate base and hydronium ion For example e 21o e Ll ii il my i i titl Sodium ion Na 39 is unreactive with water but the cyanide ion CN reacts to pro duce Chi and Oil C39 titi i llsOh 39 llCNt uq ll ttq From the Bronsted lo39ry point of View the 7K ion acts as a base because it accepts a proton from H30 You can also see however that H ion is a product so you expect the solution to have a basic pH This explains why solutions ol 39aCi39 are basic Predicting Whether A Alt Solution is Acidic Basic or Neutral A salt of a strong base and a strong acid 9 no hydrolysable ions so is neutral A salt of a strong base and a weak acid 9 anions hydrolyze so is acidic A salt of a weak base and a strong acid 9 cations hydrolyze so is basic A salt of a weak base and a weak acid 9 both ions hydrolyze so acidity or basicity depends on relative acidbase strengths compare Ka and Kb if the Ka is bigger then its acidic and vice versa CommonIon Effect The shift in an ionic equilibrium caused by the addition of a solute that provides an ion that takes part in the equilibrium Buffers A solution characterized by the ability to resist changes in pH when limited amounts of acid or base are added to it The pH of a Buffer Use the concentration of H Solubility and ComplexIon Equilibria 12915 248 PM Solubility Product Constant Ksp The equilibrium constant for the solubility equilibrium of a slightly soluble or nearly insoluble ionic compound Same as equilibrium constant but is used for aqueous solutions instead of gasses CommonIon Effect When one ion in an aqueous solution is increased it will react with the other ion to create a precipitate and return the system to equilibrium Ion Product QC The product of ion concentrations in a solution each concentration raised to a power equal to the number of ions in the formula of the ionic compound Same as Q vs Kc for equilibrium constant Predicting Precipitation If Qc gt KSp then precipitation will occur If Qc lt KSp then precipitation will not occur Thermodynamics 12915 248 PM Internal Energy U The sum of the kinetic and potential energies of the system State Function A property of a system that depends only on its present state which is completely determined by variables such as temperature and pressure AU Uf Ui Work The energy exchange that results when a force F moves an object through a distance D W F x D First Law of Thermodynamics The change in internal energy of a system AU equals q w q is the heart ABSORBED BY the system and w is work done ON the system W PAV Enthalpy H The quantity U PV Spontaneous Process A physical or chemical change that occurs by itself it may require activation energy but not work Nonspontaneous Process A physical or chemical change that doesn t occur by itself and requires work to be done to it Entropy S A measure of how dispersed the energy of a system is among the different possible ways that system can contain energy measure of disorder Second Law of Thermodynamics The total entropy of a system and its surroundings always increases for a spontaneous process AS entropy created qT For a spontaneous process at a given temperature T the change in entropy of the system is greater than the heat divided by the absolute temperature qT Spontaneous AS gt qT Third Law of Thermodynamics A substance that is perfectly crystalline at UK has an entropy of zero Standard EntropyAbsolute Entropy S The entropy value for the standard state of the species latm 1M and 25 C AS Z nS products Z mS reactants Disorder increases AS is positive when 1 Temperature increases 2 Solid dissolves in a solvent 3 Solid 9 Liquid 9 Gas 4 Less gas molecules 9 more gas molecules Gibbs Free Energy G A thermodynamic quantity defined by the equation G H TS When AG is negative the reaction is spontaneous AG AH TAS RTan 39nFEocell AG RTInQ AGof products AGOf reactants Standard Free Energy of Formation AG f The freeenergy change that occurs when 1 mol of substance is formed from its elements in their stablest states at 1 atm and at a specified temperature usually 25 C The more negative AG is the MORE spontaneous a reaction will be AG lt 10kJ 9 spontaneous reaction almost entirely products AG gt 10k 9 nonspontaneous reaction almost entirely reactants 10kJ lt AG lt 10k 9 equilibrium reaction equal amounts of products and reactants Electrochemistry 12915 248 PM VoltaicGalvanic Cells Electrochemical Cell A system consisting of electrodes that dip into an electrolyte and in which a chemical reaction either uses or generates an electric current Voltaic galvanic Cell An electrochemical cell in which a spontaneous reaction generates an electric current Electrolytic Cell An electrochemical cell in which an electric current drives an otherwise nonspontaneous reaction Balancing Redox Reactions Step 1 Identify which substances oxidation numbers change Step 2 Split equation into two half reactions Step 3 Balance each half reaction first things that change oxidation number then 0 s and H s by adding H and H20 then balance electrical charge by adding e Step 4 For bases only add the same number of OH as H on both sides of the half equation Let OH and H form water on one side Step 5 Combine half reactions and cancel out whatever you can Construction HalfCell The portion of an electrochemical cell in which a halfreaction takes place Salt Bridge A tube of an electrolyte in a gel that is connected to the two halfcells of a voltaic cell the salt bridge allows the flow of ions but prevents the mixing of the different solutions that would allow direct reaction of the cell reactants Anode The electrode at which oxidation occurs species gains electrons Cathode The electrode at which reaction occurs species loses electrons Good way to remember 9 AN OX amp RED CAT Cell Reaction The net reaction that occurs in the voltaic cell Notation 7 U il If i 21quot Center double line represents salt bride Cell Potential Potential Difference The difference in electric potential electrical pressure between two points Measured in Volts Volt V The SI unit of potential difference Joules Coulombs x Volts Faraday Constant F 96485 x 104 The magnitude of charge on one mole of electrons Cell Potential EceuElectromotive Force EMF The maximum potential difference between the electrodes of a voltaic cell Inux HI u39l W maXImum work attainable n moles of electrons FFaraday Constant Standard Cell Potential E cequot The EMF of a voltaic cell operating under standardstate conditions solutes are both 1M gas pressures are 1 atm and temperature has specified value usually 25 C This is what you calculate using the Table of Standard Electrode Potentials Standard Electrode Potential E The electrode potential when the concentrations of solutes are 1M gas pressures are 1 atm and temperature has a specified value usually 25 C Use the Table of Standard Electrode Potentials to find the Voltage of a reaction and to figure out which way it goes Nernst Equation An equation relating the cell potential to its standard cell potential and the reaction quotient ECell EoCell 39 RTquotFan ECequot cell potential under nonstandard conditions V E Cequot Cell potential under standard conditions R Gas constant 831 coltcoulombmo K T Temperature K n Moles of electrons exchanged in reaction F Faraday s constant 96500 coulombsmol Q Reaction quotient like the one used for Rate of Reactions with initial concentrations instead of equilibrium concentrations Electrolysis The process of producing a chemical change in an electrolytic cell nonspontaneous must provide electrical energy to make reactions occur For example using a DC supply Electroplating When you use an electrolysis reaction to coat an electrode with a metal that is dissolved in solution 12915 248 PM


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