Chem 112 01/07/2014
Chem 112
01/07/2014
Chem lab questions [email protected] LABS START WEEK OF JANUARY 20TH Join new course on mastering chem chem112y2014section02 (MOZILLA FIREFOX) MIDTERM Saturday, March 1 st 10am12pm Chem help center Monday – Thursday @ 11:30am1:30pm (Begins week of January 20 th) Monday and Wednesday 11:30am with Dr. Jens Mueller Tuesday 11:30 and Wednesday 12:30 with Dr. Alexandra Bartole Scott Chapter 1: Matter, measurement & problem solving 1) The scientific approach (p. 14) Hypothesis, scientific law & theory 2) Classification of matter (p.45) Matter solid, liquid & gas Solids: can be either crystalline or amorphous (eg: diamond and charcoal) Liquids: molecules are closely packed but have some ability to move around Gases: atoms have complete freedom from each other
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(p. 7) Pure Substance: made of individual atoms or molecules can be pure element or a pure compound Mixture: two or more types of atoms or molecules combined in variable proportions can be homogeneous (mixture that has uniform composition throughout) or heterogeneous (mixture that does not have uniform composition throughout)
Homogeneous eg: nitrogen & oxygen (can be mixed and will still be homogeneous) Heterogeneous eg: Wet sand
Properties and Changes of Matter (p. 810) Composition: parts or components of a sample of matter and their relative proportions Physical property: characteristics a substance displays without changing its composition (eg volatility [boiling point]) Chemical property: characteristics a substance displays only when changing its composition (eg flammability) Physical change: physical properties of sample change but composition remains unchanged Chemical change: matter is converted to a new kind, with a different composition 3) Energy changes in matter, both physical and chemical, result in either the gain or loss of energy Energy: capacity to do work Work: actions of a force applied across a distance Electrostatic force: push or pull on objects that have an electrical charge
Kinetic energy: energy of motion Potential energy: energy that is stored in matter Can be
introverted
Spontaneous processes high potential energy = less stable state 4) Temperature measure of average amount of kinetic energy (higher temperature = greater average kinetic energy) Absolute zero: 0 K & 273 C K = C + 273.15 5) SI prefix multipliers Nanotechnology length scales 1,000,000,000 times small than a meter Velocity = length/time (meters/seconds) Volume = length x length x length (m^3) Density = mass/volume (kg/m^3) % composition = mass of component/total mass x 100% (%) example: Calculating density (p.22) a man receives a ring from his fiancé. The ring has a mass of 3.15g, and displaces 0.233 cm^3 of water. Is the ring made of platinum? Density =
3.15g = 13.5g/cm^3 0.233cm^3
Chem 112 Chapter 2 atoms and elements
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Chem 112
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Early chemical discoveries Law of conservation of mass (1789: A. Lavoisier) Law of definite proportions (1797: J. Proust) Page 52 All samples of given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements Law of multiple proportions (1804: J. Dalton) Page 53 When two elements combine to form two different compounds, the relative masses of element B that combine with 1g of element A can be expressed as a ratio of small whole numbers Example: CO2: 2.67g of oxygen for 1g of carbon
Daltons Atomic Theory (1808) Page 55 1) Each chemical element is composed of minute, indestructible particles called atoms. Atoms can be neither destroyed nor be created during a chemical reaction 2) All atoms of an element are alike in mas and other properties, but the atoms of one element are different from those of all other elements
Electrons (Thomson and Mllikan) Page 5558 Electrons are particles found in all atoms Cathode rays are streams of electrons Electron has a charge of 1.60 x 10^19 C Electron has a mass of 9.1 x 10^28g Atoms must be charge neutral, so some positive charge must balance the negative charge electrons Structure of the atom: Thomson’s plum pudding atom Page 5859
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Structure of atom contains many negatively charged “electrons” These electrons are held in atom by positively charged electric field there had to be source of positive charge because atom is neutral Rutherford’s Experiment and Conclusions Page 60 Alpha particles: energetic particles that penetrate matter Atoms mostly empty space – almost all particles went straight through Atoms contain dense, positively charges particle that was small in volume compared to atom but large in mass Rutherford’s interpretation of the atom Page 60 Neutrons Page 6162 Not until 1932 that Chadwick was able to show that the missing mas came from particles that weighed nearly the same as protons but were electrically neutral Atomic mass unit = amu one amu equals 1/12 the mass of one carbon atom that contains 6 neutrons and 6 protons Elements Page 6263 Each element has unique number of protons in nucleus Number of protons in the nucleus of an atom is the atomic number Each element Isotopes Page 6465 All isotopes of element are chemically identical (undergo exact same chemical reactions)
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All isotopes of an element have same number of protons Isotopes of element have different number of neutrons Isotopes of an element have
Mass number – atomic number = neutrons in atom (# of protons and neutrons) – (# of protons) = # of neutrons
Ions Page 68 Is atom loses or gains electrons (ionization), it is called an ion Number of protons – number of electrons = ion The Periodic Law Page 6869 Mendeleev ordered elements by increasing mass (18341907) Periodic law when elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically Put elements with similar properties in same column Used pattern to predict properties of undiscovered elements Where atomic mass order did not fit other properties, he reordered by other properties (Te and I) 1/15/2014 Periodic table Page 72
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1a = alkali metals 2a = alkali earth metals (not as reactive as alkali metals) 7a = halogens 8a = noble gases Example what are the expected ion forms of the following elements? Al^3+ S^2 Ca^2+ Cs^+
Atomic Mass Page 77 Not all atoms of an element have the same mass isotopes Average mass of an element’s atoms found in a sample must be used in calculations (average must take into account the abundance of each isotope in sample) Atomic mass the average mass Mass Spectrometry Page 79 Atoms or molecules are ionized, then accelerated down a tube Some molecules turned into fragments during ionization process These fragments can be used to help determine the structure of the molecule Their path is bent by magnetic field, separating them by mass Example Page 78 120.9038 x 0.574 +122.9042 x (100% 57.4%) = 121.7560 amu 100 100% Molar Mass: The Mole and Avogadro’s number Page 81 Mole a mole is equl to the number of carbon atoms in exactly 12g of pure carbon12 Avogadro’s constant (NA) the “number of elementary entities (atoms, molecules, ions, etc) in one mole. 1 mol = 6.022 x 10^23 atoms Mass of one mole of entities is called the molar mass of a substance (mol mass of element is numerically equal to its atomic mass (amu) but has the units g/mol) Example an aluminum sphere contains 8.55 x 10^22 Al atoms. What is the radius of the sphere in cm? The density of Al is 2.70g/cm^3. D= 2.70 g/cm^3 d= m/v v=4/3 r^3
Chem 112 Chapter 3
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Chem 112
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Types of chemical bonds Compounds are made of atoms held together by chemical bonds Bonds are forces of attraction between atoms Bonding attraction comes from attractions between protons and electrons Ionic bonds result when electrons have been transferred between atoms, resulting in oppositely charged ions that attract each other Covalent bonds result when two atoms share some of their electrons (lowest potential energy = most stable) Chemical Formulas Page 94 Chemical formula symbolic representation designed to indicate (at a minimum) the elements present and the relative number of atoms of each element. Empirical formula (e.g. CH3O) Molecule formula (e.g. C2H6O2) Structural formula (actual drawing on the bonds) Elements and compounds
Diatomic molecules H, N, O, F, CL, BR, I Polyatomic molecules P4, S8, Se8 Ionic compound and their formulas Page 100101 Metals and nonmetals No individual molecule units, instead have a 3D array of cations and anions made of formula units
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man contains polyatomic ions (several ions attached together) compounds must have no total charge (must be neutral), therefore we must balance the numbers of cations and anions in a compounds to get zero for a net charge. Page 103 ** guide to writing ionic formulas chart page 75 01/20/2014 Naming ionic compounds 1) make sure compound is ionic 2) identify and name the cation metal that forms only one type of cation metal that forms more than one type of cation 3) identify an name the anion nonmetals that form simple anions 4) combine cations and anion names metal first in formula and name cation name is the metal name nonmetal anion named by changing the ending on the nonmetal name to –ide (charge based on position on the periodic table) table 3.2 common anions ( page 107)
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Naming Binary Ionic Compounds Containing a Metal That Forms More than One Kind of Cation Metal cation name: metal name followed by a Roman numeral in parentheses to indicate its charge Transition metals can have different charges Determine charge from anion charge Nonmetal ion: same as before
e.g. PbCl4 Pb^2+ / Pb^4+ Lead (IV) chloride
Naming Ionic Compounds Containing Polyatomic Ions Page 111 Seldom carry suffix –ide (except OH and CN) Often end in “ite” or “ate” and have prefixed hypo or per Names of oxygen polyatomics (often called oxoanions) are related to number of oxygen atoms in ion Polyatomic ions polyatomic ions consist of two or more atoms, covalently bonded together, that carry an overall charge (often negative) Hydrated Ionic Compounds Page 112113 Hydrates are ionic ocmpounds containing a specific number of waters for eac formula unit Water of hydration often “driven off” by heating In formual, attached waters follw (CoCl2 6H2O)
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Molecular compounds Page 113 Composed of two or more nonmetals, covalent bonds, represented by molecular formula, molecule has defined beginning & end
1) make sure compounds is molecular 2) write name of first element in formula 3) write name of second element in formula with the ending –ide
e.g. NI3 Nitrogen triiodide Binary Acids Page 116 Acids are species that produce hydrogen ions when dissolved in H2O Acidity is emphasized with special names Use Greek prefix hydro Follow with name of other nonmetal using suffix “ic” Write word acid at the end of name Oxyacids Page 117 Contain 3 different elements (hydrogen, oxygen & another nonmetal) Naming 1) if polyatomic ion name ends in –ate, then change to –ic suffic 2) is ends in –ite, then change to –ous suffix 3) write word acid at end of all names
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e.g. HF (g) Hydrogen flouride NaNO3 (aq) Na+ & NO3 = oxyacid Sodium nitrate HNO2 (aq) Nitrous acid
Formula Mass & the mole concept Page 119 Formula mass: mass of a formula unit in atomic mass units. Applied to molecular and ionic compounds In case of molecules, also known as molecular mass or molecular weight = (# of atoms of 1 st elements in chemical formula x atomic mass of 1st element) + (# of atoms of 2nd element in chemical formula x atomic mass of second element) + …. e.g. calculate formula mass of glucose, C6H12O6 =(6x12.011)+(12x1.008)+(6x16) Mole concept for molecules Page 121 One mole of a compound contains 6.022 x 10^23 formula units (or molecular units for a molecular compound) Molar mass: mass of one mole of a compounds (numerically equivalent to formula mass but with unit sof g/mol) e.g. aspirin tablet contains 325g of acetylsalicylic acid (C9H804). How many acetylsalicylic molecules does it contain? Molar mass 180.16 g/mol = 1.80 x 10^3 molx 6.022 x 10^23 = 1.09 x 10^21 molecules Chemical Composition Page 123 Molecule formula tells how many moles of an atoms are in one mole of a molecule compounds
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Percent composition Page 123 Chemical formula percent composition Mass % = mass of element ‘x’ in 1 mol compound 100% mass of 1 mol of compounds e.g. Calculate mass perfect of Cl in Freon112, C2Cl4F2 203.82g/mol 4x35.45)g x 100% = 34.7g% 203.82g 1/22/2014 1/24/2014 Writing and balancing Page 142 Reactions involve chemical change & matter Provides info about reactions Formulas of reactants and products Physical states of reactants and products Relative number of reactant and products that are required can be used to determine mass of reactant and products (g) – gas, (l) – liquid e.g. 2Co2(s) + 3C(s) 4Co(s) + 3Co2(g)
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Chapter 4 Stoichiometry Page 157158 Numerical relationship between chemical amounts in a reaction Limiting reactants For reactants with multiple reactant, one might be completely used up before the other – when this reactant is used up, reaction stops Reaction that limits amount of product formed is called the limiting reaction 1/27/2014 Solutions 163164 Many chemical reactions carried out in solution because mixing facilitates the close contact between atoms, ions or molecules required for a reaction to occur Solution homogenous mixture of two or more components Solute matter which is dissolved in the solvent Dilute solutions have small amount of solute Molarity Page 165166 Concentration of a solution is given by its molarity Molarity (M) = moles of solute Volume of solution (in Solution dilutions Page 170 solutions are often stored as concentrated stock solutions
C= n litres)
V
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to make solutions of lower concentrations from these stock solutions, more solvent is added – amount of solute doesn’t change, just the volume of solution moles of solute in solution 1 = moles solute in solution 2 concentrations and volumes of the stock and new solutions are inversely proportional mol1=mol2 M1V1 = M2V2 * moles of solute are constant, only the total volume of solution changes Solution Stoichiometry Page 173 Can be used to convert between amount of reactants and/or products in a chemical reaction Concentration can be labeled with square brackets around formula [formula] e.g. [HCl(aq)]
Solubility Page 175 There are attractive forces between the solute species holding them together also attractive forces between the solvent molecules When we mix solute with solvent, there are attractive forces between solute species and solvent molecules If attractions between solute and solvent are strong enough – solute with dissolve Result is a solution with freemoving, charged species able to conduct electricity Electrolytes and Nonelectrolytes Page 176177 Materials that dissolve in water to form a solution that will conduct electricity are called electrolytes Materials tat dissolve in water to form a solution that will not conduct electricity are called nonelectrolytes Salt vs. sugar dissolved in water Page 177
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Acids Page 178179 Acids molecular compounds that ionize when dissolved in water Molecules are pulled apart by their attration to H2O When acids ionize, they form cations and anions Percentage of molecules that ionize caries from one acid to another Acids that ionize virtually 100% are called strong acids HCl H+ + Cl Acids that only ionize a small percentage are called week acids HF > H+ + Cl
Strong and weak electrolytes Page 177179 Strong electrolytes are compounds that siddolve completely as ions Ionic compounds = strong acids – their solutions conduct electricity well Weak electrolytes are compounds that dissolve mostly as molecules, but partially as ions Weak acids – their solutions conduct electricity, but not well When compounds containing polyatomic ion dissolve, the polyatomic ion stays together When will salt dissolve? Page 181 Predicting is a compounds will dissolve in water is not easy Best way to do it is to do some experiments and develop some rules based on those experimental results empirical method * learn table 4.1 (solubility rules)
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compounds are generally soluble in water compounds that are generally insoluble in water
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Precipitation reactions Page 184185 Reactions between aqueous solutions and ionic compounds that produce an insoluble ionic compound Insoluble product is called a precipitate e.g. 2 KI (aq) + Pb(NO3)2 (aq) PbI2 (s) + 2 KNO3 (aq) Soluble Soluble Soluble Insoluble KI = a strong electrolyte – fully dissolves in water NaCl = strong electrolyte – fully dissolves in water KCl = a strong electrolyte – fully dissolves in water NaI = a strong electrolyte – fully dissolves in water K^+ + I^ + Na^+ +Cl^ KI (aq) + NaCl (aq) no reaction No precipitate formation = no reaction ** learn procedure on page 188 e.g. 4.10 page 189
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write equation for reaction of potassium carbonate and nickel (II) chloride. Ni 2+ (aq), Cl – (aq), K + (aq), Co3 2 (aq) NiCl2 (aq) + K2Co3 (aq) NiCo3 (s) + 2KCl (aq) ** soluble = aqueous and insoluble = solid usually all NO3’s dissolve and all chlorides dissolve (with the exception of lead)
Ionic equations Page 191192 e.g. Pb(NO3)2 (aq) + 2KCl (aq) 2KNO3 (aq) + PbCl2 (s) Molecular equations Complete ionic equations Aqueous strong electrolytes are written as ions – soluble salts, strong acids, strong bases Insoluble substances, weak electrolytes, and nonelectrolytes are written as empirical formulas and as molecules respectively – solids, liquids, and gases are not dissolved Spectator ions and net ionic equation – ions that are the exact same form in the reaction and product of a reaction e.g. 4.12 Page 193 3SrCl2 (aq) + 2Li3PO4 (aq) Sr3(PO4)2 (s) + 6LiCl (aq) 3Sr2+ (aq) + 6Cl (aq) + 6Li+ (aq) + 2PO4 3 (aq) Sr3(PO4)2 (s) + 6Li+ (aq) + 6Cl (aq) ** complete ionic equation
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3Sr2+ (aq) + 2PO4 3 (aq) Sr3(PO4)2 (s) ** net ionic equation
Acids & bases Page 194195 Arrhenious definitions acids – provide H+ ions (prontons) in water H+ is donated to water molecule to form hydronium ion, H3O+ bases – inize in water to form OH ions bases like NH3, that do not containOH ions, produce OH by pulling H+ off water molecules ** common acids and bases – see table 4.2 (Page 196) Diprotic acids – they can give up 2 moles of protons whereas others give equivalent amount of protons AcidBase reactions Page 197198 Also called neutralization reactions because acid and base neutralize each other’s properties 2 HNO3 (aq) + Ca(OH)2 (aq) Ca(NO3)2 (aq) + 2 H2O (l) Net ionic equation for an acid base reaction H+ (aq) + OH (aq) H2O (l) ** as long as salt that forms is soluble in water
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e.g. 4.13 Page 199 write molecular and net ionic equation for reaction between aqueous HI and aqueous Ba(OH)2.
2HI (aq) + Ba(OH)2 (aq) BaI2(aq) + 2H2O(l)
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Titration Page 200 Solution’s concentration can be determined by reacting it with another chemical (using stoich) Endpoint equivalence point Indicator e.g. 4.14 Page 202 H2SO4 (aq) + 2KOH (aq) K2SO4 (aq) + H2O (l)
C=n n=CV V
0.158 mol x 0.02287 L = 0.0036135 L
mols of H2SO4 = 0.0036135 / 2 mol C=n = 0.00180605 = 0.0903 V 0.02000 L
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Gas Evolution Reactions Page 203 Some reactions form a gas directly from the ion exchange: K2S(aq) + H2SO4(aq) K2SO4(aq) + H2S(g) Other reactions form a gas by the decomposition of one of the ion exchange products into a gas & water Not stable K2SO3(aq) + H2SO4(aq) K2SO4(aq) + H2SO3(aq H2SO3(aq) H2O(l) + SO2(g)
Compounds that undergo gas evolving reactions Page 205 table 4.3
e.g 4.14 Page 206 write molecular equation for reaction of aqueaous nitric acid and sodium carbonate. 2HNO3(aq) + Na2CO3(aq) 2NaNO3(aq) + CO2(g) + H2O (l) Redox reactions Page 207 To convert free element into an ion, atoms must fain or lose electrons (if one atom loses electron, another must accept them) e.g. Page 208 2Na(s) + Cl2(g) 2NaCl(s) Oxidation: Na _> Na+ + 1e reaction is written with two half reactions! Reduction: Cl2 + 2 e 2Cl
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L.E.O = lose electrons oxidize G.E.R = gain electrons reduce
Electron bookkeeping Page 209 Need a method for determining how electrons are transferred Can assign number to each element in reaction (oxidation state) that allows us to determine electron flow in reaction Oxidation stated not ion charged Oxidation states are imaginary charges assigned based on a set of rules Ion charges are real and measurable Rules for Oxidation States Page 210211 1. free elements have an oxidation state =0 e.g 2Na(s) + Cl2(g) 2NaCl(s) 2. monatomic ions have an oxidation state equal to their charge 3. a) sum of oxidation states of all atoms in neutral compound is 0 b) sum of oxidation state of all atoms in polyatomic ion equals charge on ion (e.g. NO3) written charge before number (+1 NOT 1+) 2/3/2014 Charles Law Boyles Law Avogadro’s Law
Chem 112 The Ideal Gas Law Page 244 – 245
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Ideal gas is a hypothetical gas where the molecules (or atoms) have no interaction with each other ‘R’ ideal gas constant 0.08206 atm x L mol x K PV = nRT (P1)(V1) = (P2)(V2) Boyle’s Law constant n and T Charles’s Law constant n and P Avogadro’s Law constant P and T d= m v
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Molar mass of a gas Page 251 to determine molar mass of unknown substance, one can heat a weighed sample until it becomes a gas and measure the T, P and V – then use the ideal gas law
M = m N e.g 5.8 Page 251
d=MP=m M = RT V
RTm VP
Chem 112 R = 62.364 L
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Mixture of Gases Page 252 When gases are mixed together, the molecules or atoms behave dependent of each other All gases in mixture have the same volume All completely fill the container therefore, each gas’s V = Vcontainer All gases in mixture are at the same T Therefore they have the same average kinetic energy In certain applications, the mixture can be thought of as one gas Air mixture, we can measure P, V an T of air as if it were a pure substance When working with gas mixtures – n can still be used but now refers to the total of moles Partial pressure Page 252 – 253 Pressure of single gas in a mixture of gases Dalton’s law of partial pressures sum of the partial pressures of all the gases in the mixture equals the total pressure P total = Pa + Pb P total = n total RT Pa = na RT V V
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Pa Pa = na na n total
(Pa)
e.g 5.10 Page 256 n (he) =
29.2g = 6.034 mol He M (He)
n N (o2) = 4.32 g = 0.1350 mol O2 M (O2) Ptotal=n (6.1699 mol)=(0.082058 Latmmol1K1)(298K)=12.1 atm V 12.5 L x He
= n He = 0.9781 x O2 = n n total n total
total = 6.1699 mol
total RT
O2 = 0.02188
Collecting gases Page 257 – 258 Gases can be collected by displacing water from a container Since water evaporates, water vapor is also present in the gas Partial pressure of water vapor (vapor pressure) depends on the temperature (read from table 5.4)
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e.g. 5.11 Page 259 Collecting gases over water Need to know partial pressure of water look on table Change kelvin temperature to Celsius P(O2) = P total – P (H2O) = 737.68 mmHg
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Reactions involving gases PV = nRT Page 260 When gases are at STP, use 1 mol = 22.4 L e.g. 5.12 Page 261 if 4.5L of O2 was formed at P = 745 mmHg and T = 308 K, than how many grams of Ag2O must have been decomposed? (R=62.364 LmmHgmol1K 1) n (O2) = n(Ag2O) n (O2) = 2 (0.1776 mol) = 2.01776 mol m = nM 231.74 g/mol = (82.3 g) M e.g. 5.13 Page 254
PV 245mmHg – 4.58L = 0.1776 mol RT R x 308 K
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how many grams of water form when 1.24 L of H2 at STP completely react with oxygen? 2H2 + O2 2H2O 1.24 L = 0.05533 mol 22.4 L mass H2O = n(H2O) M(H2O) n(H2O) 18.02 g/mol = 0.997 g of H2O
Kinetic molecular theory Page 264 – 266 Basic postulates are 1) size of a particle is very small 2) average kinetic energy of a particle is proportional to temperature (in K) 3) all collisions are elastic Can use postulates to understand the ideal gas equation (and Boyles, Charles’ & Avogadro’s Laws), as well as gas mixtures According to kinetic molecular theory – particles of different masses must have the same average KE at a given T KE = 1/2 MV2 Charles’ Law with KE Page 266 – 268 Increasing T increases their average speed, causing gas atoms/molecules to hit the wall harder and more frequently To keep P constant – the V increases Ideal vs. Real gases Page 296
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Real gases do not behave like ideal gases at high P or low T Assumptions in ideal gas law No attractions between gas moles (or atoms) Gas molecules (or atoms) do not take up space Assumptions are not valid at ow T and/or high P
Effect of P Molecular Volume Page 270 At high P amount of space occupied by molecules or atoms is significant amount of total volume Volume occupied by atoms or molecules themselves make real volume larger than idea gas law would predict Effect of Intermolecular Attractions Page 271 At low T attractions between molecules or atoms are significant Intermolecular attractions makes the real pressure less than the ideal gas law would predict Intermolecular attractions makes molar volume of a real gas smaller than predicted by ideal gas law
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Chapter 6 Thermochemistry 1. Energy & work Page 279 Energy – capacity to do work (SI units: Joules) Work – Page 280 Classification of Energy – think of energy as something an object possesses Heat (q) – one way energy can be transferred. Heat flows when two objects of different temperature are in contacts. Other way s energy transferred by work (w) Kinetic energy Energy of motions Energy that is transferred between two or more things Thermal energy Energy associated with temperature change Also a type of kinetic energy Potential energy – e.g holding a billiard ball above the ground, gravitational potential energy (obvious) or the chemical energy stored in a propane tank (not so obvious) Some forms of energy Page 280
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Electrical – Kinetic energy associated wit flow of electrical charge Heat or thermal energy – Kinetic energy associated with molecular motion Light or radiant energy – Kinetic energy associated with energy transitions in an atom or molecule Nuclear – potential energy in nucleus of atoms (stored energy) Chemical – potential energy in the attachment of atoms or because of their position
Units of energy Page 282 Amount of kinetic energy an object has is directly proportional to its mass and velocity KE = ½ mv^2 Joule (J) – amount of energy needed to move 1 kg mass a distance of 1 meter J = N x m = kg x m2/s2 Calorie (cal) – amount of energy needed to increase the temperature of one gam of water by 1 degree Celsius Kcal = energy needed to increase temperature of 1000g or water 1 degree Celsius Food calories (cal) = kcal Energy conversion factors no need to memorize 1 cal = 4.184 J 1 calorie (cal) = 1 kcal 1 kilowatthour (kWh) = 3.60 x 10^6 J Laws of Conservation of energy Page 283 Energy cannot be created or destroyed – 1st law of thermodynamics “the total energy of universe is constant”
Chem 112 Energy Flow and Conservation of Energy Page 284 Important for chemical changes System – material or process we are studying Surrounding – everything else in universe Conservation of energy Total energy change in system and surrounding must be zero (delta) energy univerise = 0 = (delta) energy system + (delta) energy surroundings
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Chem lab questions [email protected] LABS START WEEK OF JANUARY 20TH Join new course on mastering chem chem112y2014section02 (MOZILLA FIREFOX) MIDTERM Saturday, March 1 st 10am12pm Chem help center Monday – Thursday @ 11:30am1:30pm (Begins week of January 20 th) Monday and Wednesday 11:30am with Dr. Jens Mueller Tuesday 11:30 and Wednesday 12:30 with Dr. Alexandra Bartole Scott Chapter 1: Matter, measurement & problem solving 1) The scientific approach (p. 14) Hypothesis, scientific law & theory 2) Classification of matter (p.45) Matter solid, liquid & gas Solids: can be either crystalline or amorphous (eg: diamond and charcoal) Liquids: molecules are closely packed but have some ability to move around Gases: atoms have complete freedom from each other
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(p. 7) Pure Substance: made of individual atoms or molecules can be pure element or a pure compound Mixture: two or more types of atoms or molecules combined in variable proportions can be homogeneous (mixture that has uniform composition throughout) or heterogeneous (mixture that does not have uniform composition throughout)
Homogeneous eg: nitrogen & oxygen (can be mixed and will still be homogeneous) Heterogeneous eg: Wet sand
Properties and Changes of Matter (p. 810) Composition: parts or components of a sample of matter and their relative proportions Physical property: characteristics a substance displays without changing its composition (eg volatility [boiling point]) Chemical property: characteristics a substance displays only when changing its composition (eg flammability) Physical change: physical properties of sample change but composition remains unchanged Chemical change: matter is converted to a new kind, with a different composition 3) Energy changes in matter, both physical and chemical, result in either the gain or loss of energy Energy: capacity to do work Work: actions of a force applied across a distance Electrostatic force: push or pull on objects that have an electrical charge
Kinetic energy: energy of motion Potential energy: energy that is stored in matter Can be
introverted
Spontaneous processes high potential energy = less stable state 4) Temperature measure of average amount of kinetic energy (higher temperature = greater average kinetic energy) Absolute zero: 0 K & 273 C K = C + 273.15 5) SI prefix multipliers Nanotechnology length scales 1,000,000,000 times small than a meter Velocity = length/time (meters/seconds) Volume = length x length x length (m^3) Density = mass/volume (kg/m^3) % composition = mass of component/total mass x 100% (%) example: Calculating density (p.22) a man receives a ring from his fiancé. The ring has a mass of 3.15g, and displaces 0.233 cm^3 of water. Is the ring made of platinum? Density =
3.15g = 13.5g/cm^3 0.233cm^3
Chem 112 Chapter 2 atoms and elements
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Chem 112
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Early chemical discoveries Law of conservation of mass (1789: A. Lavoisier) Law of definite proportions (1797: J. Proust) Page 52 All samples of given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements Law of multiple proportions (1804: J. Dalton) Page 53 When two elements combine to form two different compounds, the relative masses of element B that combine with 1g of element A can be expressed as a ratio of small whole numbers Example: CO2: 2.67g of oxygen for 1g of carbon
Daltons Atomic Theory (1808) Page 55 1) Each chemical element is composed of minute, indestructible particles called atoms. Atoms can be neither destroyed nor be created during a chemical reaction 2) All atoms of an element are alike in mas and other properties, but the atoms of one element are different from those of all other elements
Electrons (Thomson and Mllikan) Page 5558 Electrons are particles found in all atoms Cathode rays are streams of electrons Electron has a charge of 1.60 x 10^19 C Electron has a mass of 9.1 x 10^28g Atoms must be charge neutral, so some positive charge must balance the negative charge electrons Structure of the atom: Thomson’s plum pudding atom Page 5859
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Structure of atom contains many negatively charged “electrons” These electrons are held in atom by positively charged electric field there had to be source of positive charge because atom is neutral Rutherford’s Experiment and Conclusions Page 60 Alpha particles: energetic particles that penetrate matter Atoms mostly empty space – almost all particles went straight through Atoms contain dense, positively charges particle that was small in volume compared to atom but large in mass Rutherford’s interpretation of the atom Page 60 Neutrons Page 6162 Not until 1932 that Chadwick was able to show that the missing mas came from particles that weighed nearly the same as protons but were electrically neutral Atomic mass unit = amu one amu equals 1/12 the mass of one carbon atom that contains 6 neutrons and 6 protons Elements Page 6263 Each element has unique number of protons in nucleus Number of protons in the nucleus of an atom is the atomic number Each element Isotopes Page 6465 All isotopes of element are chemically identical (undergo exact same chemical reactions)
Chem 112
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All isotopes of an element have same number of protons Isotopes of element have different number of neutrons Isotopes of an element have
Mass number – atomic number = neutrons in atom (# of protons and neutrons) – (# of protons) = # of neutrons
Ions Page 68 Is atom loses or gains electrons (ionization), it is called an ion Number of protons – number of electrons = ion The Periodic Law Page 6869 Mendeleev ordered elements by increasing mass (18341907) Periodic law when elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically Put elements with similar properties in same column Used pattern to predict properties of undiscovered elements Where atomic mass order did not fit other properties, he reordered by other properties (Te and I) 1/15/2014 Periodic table Page 72
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1a = alkali metals 2a = alkali earth metals (not as reactive as alkali metals) 7a = halogens 8a = noble gases Example what are the expected ion forms of the following elements? Al^3+ S^2 Ca^2+ Cs^+
Atomic Mass Page 77 Not all atoms of an element have the same mass isotopes Average mass of an element’s atoms found in a sample must be used in calculations (average must take into account the abundance of each isotope in sample) Atomic mass the average mass Mass Spectrometry Page 79 Atoms or molecules are ionized, then accelerated down a tube Some molecules turned into fragments during ionization process These fragments can be used to help determine the structure of the molecule Their path is bent by magnetic field, separating them by mass Example Page 78 120.9038 x 0.574 +122.9042 x (100% 57.4%) = 121.7560 amu 100 100% Molar Mass: The Mole and Avogadro’s number Page 81 Mole a mole is equl to the number of carbon atoms in exactly 12g of pure carbon12 Avogadro’s constant (NA) the “number of elementary entities (atoms, molecules, ions, etc) in one mole. 1 mol = 6.022 x 10^23 atoms Mass of one mole of entities is called the molar mass of a substance (mol mass of element is numerically equal to its atomic mass (amu) but has the units g/mol) Example an aluminum sphere contains 8.55 x 10^22 Al atoms. What is the radius of the sphere in cm? The density of Al is 2.70g/cm^3. D= 2.70 g/cm^3 d= m/v v=4/3 r^3
Chem 112 Chapter 3
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Chem 112
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Types of chemical bonds Compounds are made of atoms held together by chemical bonds Bonds are forces of attraction between atoms Bonding attraction comes from attractions between protons and electrons Ionic bonds result when electrons have been transferred between atoms, resulting in oppositely charged ions that attract each other Covalent bonds result when two atoms share some of their electrons (lowest potential energy = most stable) Chemical Formulas Page 94 Chemical formula symbolic representation designed to indicate (at a minimum) the elements present and the relative number of atoms of each element. Empirical formula (e.g. CH3O) Molecule formula (e.g. C2H6O2) Structural formula (actual drawing on the bonds) Elements and compounds
Diatomic molecules H, N, O, F, CL, BR, I Polyatomic molecules P4, S8, Se8 Ionic compound and their formulas Page 100101 Metals and nonmetals No individual molecule units, instead have a 3D array of cations and anions made of formula units
Chem 112
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man contains polyatomic ions (several ions attached together) compounds must have no total charge (must be neutral), therefore we must balance the numbers of cations and anions in a compounds to get zero for a net charge. Page 103 ** guide to writing ionic formulas chart page 75 01/20/2014 Naming ionic compounds 1) make sure compound is ionic 2) identify and name the cation metal that forms only one type of cation metal that forms more than one type of cation 3) identify an name the anion nonmetals that form simple anions 4) combine cations and anion names metal first in formula and name cation name is the metal name nonmetal anion named by changing the ending on the nonmetal name to –ide (charge based on position on the periodic table) table 3.2 common anions ( page 107)
Chem 112
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Naming Binary Ionic Compounds Containing a Metal That Forms More than One Kind of Cation Metal cation name: metal name followed by a Roman numeral in parentheses to indicate its charge Transition metals can have different charges Determine charge from anion charge Nonmetal ion: same as before
e.g. PbCl4 Pb^2+ / Pb^4+ Lead (IV) chloride
Naming Ionic Compounds Containing Polyatomic Ions Page 111 Seldom carry suffix –ide (except OH and CN) Often end in “ite” or “ate” and have prefixed hypo or per Names of oxygen polyatomics (often called oxoanions) are related to number of oxygen atoms in ion Polyatomic ions polyatomic ions consist of two or more atoms, covalently bonded together, that carry an overall charge (often negative) Hydrated Ionic Compounds Page 112113 Hydrates are ionic ocmpounds containing a specific number of waters for eac formula unit Water of hydration often “driven off” by heating In formual, attached waters follw (CoCl2 6H2O)
Chem 112
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Molecular compounds Page 113 Composed of two or more nonmetals, covalent bonds, represented by molecular formula, molecule has defined beginning & end
1) make sure compounds is molecular 2) write name of first element in formula 3) write name of second element in formula with the ending –ide
e.g. NI3 Nitrogen triiodide Binary Acids Page 116 Acids are species that produce hydrogen ions when dissolved in H2O Acidity is emphasized with special names Use Greek prefix hydro Follow with name of other nonmetal using suffix “ic” Write word acid at the end of name Oxyacids Page 117 Contain 3 different elements (hydrogen, oxygen & another nonmetal) Naming 1) if polyatomic ion name ends in –ate, then change to –ic suffic 2) is ends in –ite, then change to –ous suffix 3) write word acid at end of all names
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e.g. HF (g) Hydrogen flouride NaNO3 (aq) Na+ & NO3 = oxyacid Sodium nitrate HNO2 (aq) Nitrous acid
Formula Mass & the mole concept Page 119 Formula mass: mass of a formula unit in atomic mass units. Applied to molecular and ionic compounds In case of molecules, also known as molecular mass or molecular weight = (# of atoms of 1 st elements in chemical formula x atomic mass of 1st element) + (# of atoms of 2nd element in chemical formula x atomic mass of second element) + …. e.g. calculate formula mass of glucose, C6H12O6 =(6x12.011)+(12x1.008)+(6x16) Mole concept for molecules Page 121 One mole of a compound contains 6.022 x 10^23 formula units (or molecular units for a molecular compound) Molar mass: mass of one mole of a compounds (numerically equivalent to formula mass but with unit sof g/mol) e.g. aspirin tablet contains 325g of acetylsalicylic acid (C9H804). How many acetylsalicylic molecules does it contain? Molar mass 180.16 g/mol = 1.80 x 10^3 molx 6.022 x 10^23 = 1.09 x 10^21 molecules Chemical Composition Page 123 Molecule formula tells how many moles of an atoms are in one mole of a molecule compounds
Chem 112
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Percent composition Page 123 Chemical formula percent composition Mass % = mass of element ‘x’ in 1 mol compound 100% mass of 1 mol of compounds e.g. Calculate mass perfect of Cl in Freon112, C2Cl4F2 203.82g/mol 4x35.45)g x 100% = 34.7g% 203.82g 1/22/2014 1/24/2014 Writing and balancing Page 142 Reactions involve chemical change & matter Provides info about reactions Formulas of reactants and products Physical states of reactants and products Relative number of reactant and products that are required can be used to determine mass of reactant and products (g) – gas, (l) – liquid e.g. 2Co2(s) + 3C(s) 4Co(s) + 3Co2(g)
Chem 112
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Chapter 4 Stoichiometry Page 157158 Numerical relationship between chemical amounts in a reaction Limiting reactants For reactants with multiple reactant, one might be completely used up before the other – when this reactant is used up, reaction stops Reaction that limits amount of product formed is called the limiting reaction 1/27/2014 Solutions 163164 Many chemical reactions carried out in solution because mixing facilitates the close contact between atoms, ions or molecules required for a reaction to occur Solution homogenous mixture of two or more components Solute matter which is dissolved in the solvent Dilute solutions have small amount of solute Molarity Page 165166 Concentration of a solution is given by its molarity Molarity (M) = moles of solute Volume of solution (in Solution dilutions Page 170 solutions are often stored as concentrated stock solutions
C= n litres)
V
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to make solutions of lower concentrations from these stock solutions, more solvent is added – amount of solute doesn’t change, just the volume of solution moles of solute in solution 1 = moles solute in solution 2 concentrations and volumes of the stock and new solutions are inversely proportional mol1=mol2 M1V1 = M2V2 * moles of solute are constant, only the total volume of solution changes Solution Stoichiometry Page 173 Can be used to convert between amount of reactants and/or products in a chemical reaction Concentration can be labeled with square brackets around formula [formula] e.g. [HCl(aq)]
Solubility Page 175 There are attractive forces between the solute species holding them together also attractive forces between the solvent molecules When we mix solute with solvent, there are attractive forces between solute species and solvent molecules If attractions between solute and solvent are strong enough – solute with dissolve Result is a solution with freemoving, charged species able to conduct electricity Electrolytes and Nonelectrolytes Page 176177 Materials that dissolve in water to form a solution that will conduct electricity are called electrolytes Materials tat dissolve in water to form a solution that will not conduct electricity are called nonelectrolytes Salt vs. sugar dissolved in water Page 177
Chem 112
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Acids Page 178179 Acids molecular compounds that ionize when dissolved in water Molecules are pulled apart by their attration to H2O When acids ionize, they form cations and anions Percentage of molecules that ionize caries from one acid to another Acids that ionize virtually 100% are called strong acids HCl H+ + Cl Acids that only ionize a small percentage are called week acids HF > H+ + Cl
Strong and weak electrolytes Page 177179 Strong electrolytes are compounds that siddolve completely as ions Ionic compounds = strong acids – their solutions conduct electricity well Weak electrolytes are compounds that dissolve mostly as molecules, but partially as ions Weak acids – their solutions conduct electricity, but not well When compounds containing polyatomic ion dissolve, the polyatomic ion stays together When will salt dissolve? Page 181 Predicting is a compounds will dissolve in water is not easy Best way to do it is to do some experiments and develop some rules based on those experimental results empirical method * learn table 4.1 (solubility rules)
Chem 112
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compounds are generally soluble in water compounds that are generally insoluble in water
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Precipitation reactions Page 184185 Reactions between aqueous solutions and ionic compounds that produce an insoluble ionic compound Insoluble product is called a precipitate e.g. 2 KI (aq) + Pb(NO3)2 (aq) PbI2 (s) + 2 KNO3 (aq) Soluble Soluble Soluble Insoluble KI = a strong electrolyte – fully dissolves in water NaCl = strong electrolyte – fully dissolves in water KCl = a strong electrolyte – fully dissolves in water NaI = a strong electrolyte – fully dissolves in water K^+ + I^ + Na^+ +Cl^ KI (aq) + NaCl (aq) no reaction No precipitate formation = no reaction ** learn procedure on page 188 e.g. 4.10 page 189
Chem 112
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write equation for reaction of potassium carbonate and nickel (II) chloride. Ni 2+ (aq), Cl – (aq), K + (aq), Co3 2 (aq) NiCl2 (aq) + K2Co3 (aq) NiCo3 (s) + 2KCl (aq) ** soluble = aqueous and insoluble = solid usually all NO3’s dissolve and all chlorides dissolve (with the exception of lead)
Ionic equations Page 191192 e.g. Pb(NO3)2 (aq) + 2KCl (aq) 2KNO3 (aq) + PbCl2 (s) Molecular equations Complete ionic equations Aqueous strong electrolytes are written as ions – soluble salts, strong acids, strong bases Insoluble substances, weak electrolytes, and nonelectrolytes are written as empirical formulas and as molecules respectively – solids, liquids, and gases are not dissolved Spectator ions and net ionic equation – ions that are the exact same form in the reaction and product of a reaction e.g. 4.12 Page 193 3SrCl2 (aq) + 2Li3PO4 (aq) Sr3(PO4)2 (s) + 6LiCl (aq) 3Sr2+ (aq) + 6Cl (aq) + 6Li+ (aq) + 2PO4 3 (aq) Sr3(PO4)2 (s) + 6Li+ (aq) + 6Cl (aq) ** complete ionic equation
Chem 112
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3Sr2+ (aq) + 2PO4 3 (aq) Sr3(PO4)2 (s) ** net ionic equation
Acids & bases Page 194195 Arrhenious definitions acids – provide H+ ions (prontons) in water H+ is donated to water molecule to form hydronium ion, H3O+ bases – inize in water to form OH ions bases like NH3, that do not containOH ions, produce OH by pulling H+ off water molecules ** common acids and bases – see table 4.2 (Page 196) Diprotic acids – they can give up 2 moles of protons whereas others give equivalent amount of protons AcidBase reactions Page 197198 Also called neutralization reactions because acid and base neutralize each other’s properties 2 HNO3 (aq) + Ca(OH)2 (aq) Ca(NO3)2 (aq) + 2 H2O (l) Net ionic equation for an acid base reaction H+ (aq) + OH (aq) H2O (l) ** as long as salt that forms is soluble in water
Chem 112
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e.g. 4.13 Page 199 write molecular and net ionic equation for reaction between aqueous HI and aqueous Ba(OH)2.
2HI (aq) + Ba(OH)2 (aq) BaI2(aq) + 2H2O(l)
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Titration Page 200 Solution’s concentration can be determined by reacting it with another chemical (using stoich) Endpoint equivalence point Indicator e.g. 4.14 Page 202 H2SO4 (aq) + 2KOH (aq) K2SO4 (aq) + H2O (l)
C=n n=CV V
0.158 mol x 0.02287 L = 0.0036135 L
mols of H2SO4 = 0.0036135 / 2 mol C=n = 0.00180605 = 0.0903 V 0.02000 L
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Gas Evolution Reactions Page 203 Some reactions form a gas directly from the ion exchange: K2S(aq) + H2SO4(aq) K2SO4(aq) + H2S(g) Other reactions form a gas by the decomposition of one of the ion exchange products into a gas & water Not stable K2SO3(aq) + H2SO4(aq) K2SO4(aq) + H2SO3(aq H2SO3(aq) H2O(l) + SO2(g)
Compounds that undergo gas evolving reactions Page 205 table 4.3
e.g 4.14 Page 206 write molecular equation for reaction of aqueaous nitric acid and sodium carbonate. 2HNO3(aq) + Na2CO3(aq) 2NaNO3(aq) + CO2(g) + H2O (l) Redox reactions Page 207 To convert free element into an ion, atoms must fain or lose electrons (if one atom loses electron, another must accept them) e.g. Page 208 2Na(s) + Cl2(g) 2NaCl(s) Oxidation: Na _> Na+ + 1e reaction is written with two half reactions! Reduction: Cl2 + 2 e 2Cl
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L.E.O = lose electrons oxidize G.E.R = gain electrons reduce
Electron bookkeeping Page 209 Need a method for determining how electrons are transferred Can assign number to each element in reaction (oxidation state) that allows us to determine electron flow in reaction Oxidation stated not ion charged Oxidation states are imaginary charges assigned based on a set of rules Ion charges are real and measurable Rules for Oxidation States Page 210211 1. free elements have an oxidation state =0 e.g 2Na(s) + Cl2(g) 2NaCl(s) 2. monatomic ions have an oxidation state equal to their charge 3. a) sum of oxidation states of all atoms in neutral compound is 0 b) sum of oxidation state of all atoms in polyatomic ion equals charge on ion (e.g. NO3) written charge before number (+1 NOT 1+) 2/3/2014 Charles Law Boyles Law Avogadro’s Law
Chem 112 The Ideal Gas Law Page 244 – 245
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Chem 112
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Ideal gas is a hypothetical gas where the molecules (or atoms) have no interaction with each other ‘R’ ideal gas constant 0.08206 atm x L mol x K PV = nRT (P1)(V1) = (P2)(V2) Boyle’s Law constant n and T Charles’s Law constant n and P Avogadro’s Law constant P and T d= m v
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Molar mass of a gas Page 251 to determine molar mass of unknown substance, one can heat a weighed sample until it becomes a gas and measure the T, P and V – then use the ideal gas law
M = m N e.g 5.8 Page 251
d=MP=m M = RT V
RTm VP
Chem 112 R = 62.364 L
01/07/2014 mmHg (mol)(K)
Mixture of Gases Page 252 When gases are mixed together, the molecules or atoms behave dependent of each other All gases in mixture have the same volume All completely fill the container therefore, each gas’s V = Vcontainer All gases in mixture are at the same T Therefore they have the same average kinetic energy In certain applications, the mixture can be thought of as one gas Air mixture, we can measure P, V an T of air as if it were a pure substance When working with gas mixtures – n can still be used but now refers to the total of moles Partial pressure Page 252 – 253 Pressure of single gas in a mixture of gases Dalton’s law of partial pressures sum of the partial pressures of all the gases in the mixture equals the total pressure P total = Pa + Pb P total = n total RT Pa = na RT V V
Chem 112
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Pa Pa = na na n total
(Pa)
e.g 5.10 Page 256 n (he) =
29.2g = 6.034 mol He M (He)
n N (o2) = 4.32 g = 0.1350 mol O2 M (O2) Ptotal=n (6.1699 mol)=(0.082058 Latmmol1K1)(298K)=12.1 atm V 12.5 L x He
= n He = 0.9781 x O2 = n n total n total
total = 6.1699 mol
total RT
O2 = 0.02188
Collecting gases Page 257 – 258 Gases can be collected by displacing water from a container Since water evaporates, water vapor is also present in the gas Partial pressure of water vapor (vapor pressure) depends on the temperature (read from table 5.4)
Chem 112
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e.g. 5.11 Page 259 Collecting gases over water Need to know partial pressure of water look on table Change kelvin temperature to Celsius P(O2) = P total – P (H2O) = 737.68 mmHg
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Reactions involving gases PV = nRT Page 260 When gases are at STP, use 1 mol = 22.4 L e.g. 5.12 Page 261 if 4.5L of O2 was formed at P = 745 mmHg and T = 308 K, than how many grams of Ag2O must have been decomposed? (R=62.364 LmmHgmol1K 1) n (O2) = n(Ag2O) n (O2) = 2 (0.1776 mol) = 2.01776 mol m = nM 231.74 g/mol = (82.3 g) M e.g. 5.13 Page 254
PV 245mmHg – 4.58L = 0.1776 mol RT R x 308 K
Chem 112
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how many grams of water form when 1.24 L of H2 at STP completely react with oxygen? 2H2 + O2 2H2O 1.24 L = 0.05533 mol 22.4 L mass H2O = n(H2O) M(H2O) n(H2O) 18.02 g/mol = 0.997 g of H2O
Kinetic molecular theory Page 264 – 266 Basic postulates are 1) size of a particle is very small 2) average kinetic energy of a particle is proportional to temperature (in K) 3) all collisions are elastic Can use postulates to understand the ideal gas equation (and Boyles, Charles’ & Avogadro’s Laws), as well as gas mixtures According to kinetic molecular theory – particles of different masses must have the same average KE at a given T KE = 1/2 MV2 Charles’ Law with KE Page 266 – 268 Increasing T increases their average speed, causing gas atoms/molecules to hit the wall harder and more frequently To keep P constant – the V increases Ideal vs. Real gases Page 296
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Real gases do not behave like ideal gases at high P or low T Assumptions in ideal gas law No attractions between gas moles (or atoms) Gas molecules (or atoms) do not take up space Assumptions are not valid at ow T and/or high P
Effect of P Molecular Volume Page 270 At high P amount of space occupied by molecules or atoms is significant amount of total volume Volume occupied by atoms or molecules themselves make real volume larger than idea gas law would predict Effect of Intermolecular Attractions Page 271 At low T attractions between molecules or atoms are significant Intermolecular attractions makes the real pressure less than the ideal gas law would predict Intermolecular attractions makes molar volume of a real gas smaller than predicted by ideal gas law
Chem 112
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Chapter 6 Thermochemistry 1. Energy & work Page 279 Energy – capacity to do work (SI units: Joules) Work – Page 280 Classification of Energy – think of energy as something an object possesses Heat (q) – one way energy can be transferred. Heat flows when two objects of different temperature are in contacts. Other way s energy transferred by work (w) Kinetic energy Energy of motions Energy that is transferred between two or more things Thermal energy Energy associated with temperature change Also a type of kinetic energy Potential energy – e.g holding a billiard ball above the ground, gravitational potential energy (obvious) or the chemical energy stored in a propane tank (not so obvious) Some forms of energy Page 280
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Electrical – Kinetic energy associated wit flow of electrical charge Heat or thermal energy – Kinetic energy associated with molecular motion Light or radiant energy – Kinetic energy associated with energy transitions in an atom or molecule Nuclear – potential energy in nucleus of atoms (stored energy) Chemical – potential energy in the attachment of atoms or because of their position
Units of energy Page 282 Amount of kinetic energy an object has is directly proportional to its mass and velocity KE = ½ mv^2 Joule (J) – amount of energy needed to move 1 kg mass a distance of 1 meter J = N x m = kg x m2/s2 Calorie (cal) – amount of energy needed to increase the temperature of one gam of water by 1 degree Celsius Kcal = energy needed to increase temperature of 1000g or water 1 degree Celsius Food calories (cal) = kcal Energy conversion factors no need to memorize 1 cal = 4.184 J 1 calorie (cal) = 1 kcal 1 kilowatthour (kWh) = 3.60 x 10^6 J Laws of Conservation of energy Page 283 Energy cannot be created or destroyed – 1st law of thermodynamics “the total energy of universe is constant”
Chem 112 Energy Flow and Conservation of Energy Page 284 Important for chemical changes System – material or process we are studying Surrounding – everything else in universe Conservation of energy Total energy change in system and surrounding must be zero (delta) energy univerise = 0 = (delta) energy system + (delta) energy surroundings
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