Principles of Chemistry I
Principles of Chemistry I CHEM 1307
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Atoms Molecules and Ions electron proton neutron nucleus Nothing exists except atoms and empty space everything else is opinionquot Democritus 460 370 BC Success in Chem 1307 Read chapters in advance Write down concepts you don t understand Take notes Don t write down every word Make notes if something is confusing Review your notes regularly Practice Do homework every week Practice with end of chapter questions Seek Help Make time for office hours Supplemental instruction is useful Time 69 hours per week Chapter 1 Self Study Quantitative information that is numerical ie temperature at which a substance melts Qualitative nonnumerical observations ie color of a substance macroscopic can be seen by the eye ie mountains rocks sand large range of sizes microscopic too small to be seen clearly by the eye ie tiny plants and animals crystals particulate tiny particles molecules that make up matter and cannot be seen even with powerful optical microscopes Classifying Matter Pure Substance Substance that cannot be separated into other kinds of matter by any physical process Mixture Two or more substances can be homogeneous or heterogeneous separated by physical means ie filtration Elements 5 Pure Substances Element Simplest type of matter consists of one type of atom 117 known 90 naturally occurring Metals nonmetals metalloids Compound A pure substance that can be broken down into two or more other pure substances by a chemical process Elements combine to form compounds C02 H20 NH3 etc Ionic Metal Nonmetal ie NaCl CaSO4 Molecular Two nonmetals or a nonmetal metalloid ie CH3OH SiCI4 Diatomic Molecule Two of the same atoms chemically bound together H2 N2 02 F2 Clz Brz I2 Physical amp Chemical Properties Physical Properties that can be observed and measured without changing the substance chemically Examples of physical properties odor taste color melting point boiling point density Chemical All the chemical changes possible for a substance upon a chemical change the substance becomes something new Examples of chemical properties change in color iron turns from gray to red upon rusting change in smelltaste milk spoiling feel heat or see light striking a match Law of Conservation of Mass Law of Conservation of Mass Matter is neither created nor destroyed during a chemical reaction The total mass does not change during a chemical reaction reactantl reactantZ V total mass total mass calcium oxide carbon dioxide gt calcium carbonate Ca0 CO2 gt CaCO3 5608g 44009 gt 10008g Law of Definite Composition The percentage by mass of the elements in a compound is always the same no matter how the compound is created CaCO3 calcium carbonate Molar mass 4008 1201 3 X 1600 10008 Mass Fraction Percent by Mass parts 100 part parts 100 parts 040 calcium 40 calcium 012 carbon 12 carbon 048 oxygen 48 oxygen 100 part by mass 100 by mass The law of definite composition applies to every compound Dalton s Atomic Theow 1 Each element is made up of tiny individual particles called atoms Atoms can be divided into subatomic particles 2 Atoms are indivisible they cannot be created or destroyed Nuclear reactions 3 Atoms of the same element are identical they have the same mass and properties Elements have different isotopes which are NOT identical 4 Atoms of one element are different from atoms of any other element Different elements have similar chemical properties 5 Atoms of one element can combine with atoms of another element to form compounds with whole number ratios Law of Multiple Proportions Subatomic Particles In the 18205 Dalton s theory was challenged by the suggestion that atoms contain smaller parts called subatomic particles The Electron Discovered by Michael Faraday and William Crookes J J Thompson described the properties in 1897 Symbol Charge Mass e39 1 9109 x 103928 9 How are atoms neutral if they contain these electrons Predicted Model of the Atom For an atom to remain neutral it must also contain a positive charge Plum Pudding Model Proposed in 1904 by J J Thompson Explained how the positive and negative parts of an atom were arranged Electrons thought to be floating in a sphere of positive charge Sphere of positive charge The Nuclear Atom Rutherford Detector Lead In 1911 Rutherford and his students performed a series of scattering experiments to understand the atom A beam of positively charged ocparticles He atoms with no negatively charged electrons were directed at a piece of gold foil to probe the structure of the gold atom Experiment done to prove the Plum Pudding Model Gold Foil Experiment Particles passing deflected through Alphaparticles Plum pudding atom Atoms in gold foil aaaa 7 Thorium ngher Edumlmn Rutherford s Expected Result Rutherford s Actual Result ocparticles are 7000 times the size of an electron All particles should have passed through the gold foil without deflection However some particles were deflected randomly and some were deflected back at acute angles The Nuclear Model of the Atom 1 Every atom contains an extremely small very dense nucleus 2 All of the positive charge and nearly all of the mass of an atom is concentrated in the nucleus 3 The nucleus is surrounded by a much larger volume of nearly empty space that makes up the rest of the atom 4 The space outside the nucleus is very thinly populated by electrons the total charge of which balances exactly the positive charge of the nucleus Subatomic Particles Protons p 1673 x 103924 g 1836 times the mass of the electron Electrical charge 1 Electrons e39 9109 x 103928 g Electrical charge 1 Neutrons n0 1675 x 103924 g Electrical charge Neutral Number of Protons and Neutrons Atomic Number The number of protons in an atom of a certain element Represented by the symbol 2 Atoms are electrically neutral so there are the same number of protons and electrons in an atom Mass Number The number of protons plus neutrons in an atom of a certain element Represented by the symbol A Nuclear Symbols Mass number p n0A omic Atomic Z igmbol 6 C 7 N 8 0 number ltpgt carbon12 nitrogen14 oxygen16 Isotopes Number of Neutrons Same number of protons different number of neutrons in an atom An atom of uranium 235 An atom of uranium 238 Natural Abundance For a given element the percentage of each isotope found on earth U 992742 146 neutrons 07204 143 neutrons 00054 142 neutrons If you count out 10000 uranium atoms from an average sample 9927 would be U238 72 would be U235 1 would be U234 Isotopes Practice Problems Determine the element and number of protons neutrons and electrons in the following a 1w pZ5 n 11 56 e39p5 b 41 20x pZ20 n 41 2021 C 127 53y pZ53 n 127 5374 Atomic Mass Unit In the early 18005 John Dalton proposed a relative scale of atomic masses The standard today uses the mass of carbon12 which has 6 protons and 6 neutrons thereby having a mass of 12 amu 1 amu 112 the mass of one carbon12 atom By comparison oxygen has 8 protons and 8 neutrons thereby having an atomic mass of 16 amu or 13329 X mass C 12 1 amu 166054 X 103924 grams Calculating Atomic Mass Atomic mass is based on the natural abundance of the isotopes of a certain element I Table 52 Percent Abundance of Some Natural Isotopes Symbol Mass amu Percent Symbol Mass amu Percent 1H 1007825035 999885 13F 1899840322 100 2014101779 010115 3197207070 9493 3116 301602931 0000137 3297145843 076 QIIe 400260324 99999863 3396786665 429 E C 12 exactly 9893 28 3596708062 002 1C 13003354826 1107 cl 34968852721 75178 EN 14003074002 99632 EC 3696590262 2422 EH 1500010897 0368 389637074 932581 3 0 1599491463 99757 fij 399639992 00117 150 1619991312 01038 409618254 617302 10 179991603 01205 o Zu 7Thomsun many 21mm Calculating Atomic Mass o 39 0 39 t 2 Atomc Weght Wymass isotope 1 Wymass isotope 2 Bromine has two naturally occurring isotopes Calculate the atomic mass of Br from the mass and natural abundance of each isotope EBr 78918338 5069 Br 80916291 4931 Make sure to convert percentages to decimals by dividing by 100 05069 X 78918338 400037 amu 04931 X 80916291 398998 amu atomic mass of bromine 799035 amu Value is close to 80 naturally occurring Br contains nearly equal amounts of both isotopes Calculating Isotope Abundances Antimony Sb has two stable isotopes 121Sb 120904 amu and 1233b 122904 amu What are the relative abundances of these isotopes From the periodic table the atomic mass of Sb is 121760 amu Because the atomic mass is closer to 121 than 123 we can determine that isotope 121Sb is more abundant Atomic Weight Wymass isotope 1 ltWmwezgtmass isotope 2 Atomic Weight 121760 x120904 y122904 The sum of the abundance must equal 1 xy1 121760 x120904 1x122904 121760 120904x 122904 122904x 121760 122904 120904x 122904x Solving for x x 0572 therefore 121Sb has 5720 abundance 123Sb has 4280 abundance Looking For Trends in the Elements As elements were discovered scientists were trying to organize them in one place In 1869 Dmitri Mendeleev discovered that if elements were arranged by atomic mass there was a recurring pattern 65 elements had been discovered when Mendeleev began his periodic table Elements with similar properties were placed in rows and columns The increase in atomic mass was not the perfect trend some elements were placed out of order to match their properties In 1913 H G J Moseley determined that elements should be ordered according to atomic number Periodic Table The Key to Chemistry columns groups elements with similar properties 1A 7A SA 1 17 18 2Alt CurreniUSusage gt3A AA 5A 6A 2lt lUPACnotalicn gt13 14 15 16 rows Be periods gt increasing Z atomic number Grouping the Elements 1A 8A 1 Metals 18 1 2A Meta oids 3A 4A 5A 6A 7A 2 H 2 13 14 15 16 17 He Nomnetals 3 4 5 6 7 8 9 10 Li Be B C N 0 F Ne 11 12 313 4B 58 6B 7B 88 1 1B 2B 13 14 15 16 17 18 Na M 3 4 5 6 7 8 9 10 11 12 A1 Si P 5 C1 Ar 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 K Ca Sc Ti V Cr M11 Fe Co Ni Cu Zn Ga Ge As Se Br Kr 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 Cs Ba La Hf Ta W Re 05 Ir Pt Au Hg T1 Pb Bi Po At Rn 87 88 89 104 105 106 107 108 109 110 111 112 114 116 Fr Ra Ac Rf Db Sg Bh Hs Mt Lamhamde 58 59 60 61 62 63 64 65 66 67 68 69 70 71 5 Ce Pr Nd Fm 8111 Eu Gd Tb Dy Ho Er Tm Yb Lu A mid 90 91 92 93 94 95 96 97 98 99 100 101 102 103 C es Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr Properties of Metals amp Nonmetals Metals good conductors of heat and electricity can be hammered into sheets malleable form compounds with nonmetals but NOT other metals lose electrons to become positively charged usually shiny Nonmetals do not conduct heat or electricity brittle form compounds with metals AND nonmetals usually gain electrons to become negatively charged Metalloids have the physical appearance and properties of a metal behave chemically like a nonmetal Further Classifications Alkali Metals 1 Alkaline Earths 2 13 14 15 16 7 3A 4A 5A DA 5 Transition Metals o 3456789101112 Lanthanides and Actinides Quick Introduction to the Main Group Group 1 Alkali Metals Silver colored solid at room temperature Very reactive especially with water Group 2 Alkaline Earths React with water to produce alkaline solutions Ca Mg very abundant in the earth s crust Radium Ra is radioactive and used to treat cancer Group 3 Boron is a metalloid while the rest of the group are metals Aluminum is the most abundant metal in the earth s crust Group 4 Elements in this group are metals nonmetals and metalloids Carbon is the basis of living things exists as multiple allotropes most common graphite and diamond Allotropes Different forms of the same element that exist in the same physical state under the same temperature and pressure Each allotrope has its own properties a Graphite 2 Diamond 3 Burkyballs re a mksCole Sewage Lear an Graphite consists of sheets of carbon that are weakly connected while diamond forms a hard dense structure Quick Introduction to the Main Group Group 5 Nitrogen makes up 79 of the earth s atmosphere Phosphorus is used in bones teeth and DNA multiple allotropes exist including red and white phosphorus Group 6 Oxygen makes up 20 of the earth s atmosphere and has an allotrope called ozone O3 Most energy that powers life on earth is derived from reactions in which oxygen combines with other substances S Se Te called chalcogens stinky poisonous Group 7 Halogens Very reactive highly toxic Most exist in diatomic form Group 8 Noble Gases Least reactive elements all exist as gases Formulas Molecular Formula Describes the number of atoms of each element in a molecule but gives no structural information about how the atoms are connected Ethanol C2H6O Condensed Formula Indicates how atoms are grouped together which elements are bonded to each other Ethanol CH3CHZOH Structural Formula Shows how all the atoms are attached within a molecule The lines between atoms represent chemical bonds that hold the molecule together H C C O H H H Ion Change in the Number of e39 Ion charged atom Cation Positively charged ion formed by losing electrons of protons gt of electrons Anions Negatively charged ion formed by gaining electrons of protons lt of electrons L Gain a one electron one electron 4 x V x Silicon cation Neutral silicon atom Silicon anion How to Predict Ionization Metals tend to lose electrons to become positively charged cations Nonmetals tend to gain electrons to become negatively charged anions Li atom 3 protons and 3 electrons 3e 3913P 53n Lithium Li gt e Li39 cation 3 protons e 4 and 2 electrons 39l3p 8311 Lithium ion L1 L1 Li quot39 3P 31 3n 3n Tl 7 F F 9P 9P 10 10h El Monatomic Ions A monatomic ion is formed when a single atom gains or loses e39 Main group metals form positive ions having a charge equal to the group number of the metal A magnesium ion Group 2 is formed when Mg loses two electrons Mg gt Mg2 2 e39 Magnesium ion Nonmetals form negative ions having a charge equal to the group number of the nonmetal minus 8 Nitrogen Group 5 has a 5 8 3 charge when ionized N3e39 gtN339 nitrogen nitride ion Naming Monatomic Ions Naming Cations The name of a monatomic cation is the element name followed by the word ion ie Na is the sodium ion Naming Transition Metals The charge is included in the name of an ion only when the ions of an element exhibit more than one common charge ie Fe2 is the ironII ion while Fe3 is the ironIII ion Naming Anions The name of a monatomic anion is the element name changed to end in ide followed by the word ion ie P339 is the phosphide ion Monatomic Anion Review Table Name Symbol Base name Anion name Symbol Hydrogen H Hydr Hydride H39 Fluorine F Fluor Fluoride F39 Chlorine C Chlor Chloride CI39 Bromine Br Brom Bromide Br39 Iodine I Iod Iodide 139 Oxygen 0 OX Oxide OZ39 Sulfur S Sulf Sulfide SZ39 Nitrogen N Nitr Nitride N339 Phosphorous P Phosph Phosphide P339 Carbon C Carb Carbide C439 Exceptions Transition metals form cations but their charges are not easily predicted Many TMs form several different ions Co Co2 2 e39 cobaltII ion Co C03 3 e39 cobaltIII ion i No space between last letter of element name and parenthesis Hydrogen can lose an electron to form a hydrogen ion H Hydrogen can gain an electron to form a hydride ion H Hydrogen most commonly forms the hydrogen ion H Metals with Multiple Charges Element Ion Formula Systematic Name Common Name Copper Cu1 copperI cuprous Cu2 copperII cupric Cobalt Co2 cobaltII Co3 cobalt III Iron Fe2 ironII ferrous Fezquot ironIII ferric Manganese Mn2 manganeseII Mn3 manganeseIII Tin Sn2 tinII stannous Sn4 tinIV stannic Lead Pb2 leadII Pb4 leadIV Mercury H922 mercuryI mercurous Hg2 mercuryII mercuric Common Polyatomic Ions Composed of a two or more different atoms with an overall charge Cations Ammonium NH4 Hydronium H3O Perchlorate CIO439 An39ons39 Chlorate CIO339 Chlorite 002 Carbonate C032 Hypochlorite 00 Hydrogen Carbonate HCO339 These names can be applied to other Chromate CI O42 halogens ie 10439 periodate Acetate C2H3OZ39 Cyanide CN39 Dichromate Cr207239 Peroxide 02239 Hydride H Phosphate PO4339 Hydroxide OH Hydrogen Phosphate HPO4239 Nitrate NO339 Sulfate 504239 Nitrite N05 Sulfite 503239 nAI M a n a gal A kilnN Ionic Compounds Ionic compounds are formed by one metal and one nonmetal Chemical compounds are electrically neutral so for ionic compounds the formula must have an equal number of positive and negative charges Write the formula for the cation followed by the formula for the anion omitting the charges Use subscripts to show the number of each ion in the formula so that the overall charge is equal to zero a Use the fewest number of ions possible ie CaCl2 not CazCl4 b If a polyatomic ion is needed more than once enclose the ion in parentheses and place the subscript after the closed parenthesis ie MgOH2 not MgOH2 Parentheses are ONLY used with polyatomic ions Names of Ionic Compounds Rules 1 Write the name of the cation 2 Write the name of the anion WARNING Prefixes are not included in the names of ionic compounds CaBr2 calcium bromide M9CN2 magnesium cyanide A253 aluminum sulfide Sr3PO42 strontium phosphate The Catch When naming an ionic compound containing a metal that can have more than one ionic charge the compound name includes the charge of the metal FeI3 ironIII iodide Formulas for Ionic Compounds A net zero charge is achieved by combining cations and anions so that the overall positive and negative charges balance Write the formula for the combination of K and 0239 1 K 1 charge 1 0239 2 charge To get charge balance we need 2 charge so 2 K Formula is K20 Write the formula for the combination of Ba2 and PO4339 1 Ba2 2 charge 1 PO4339 3 charge To get charge balance we need 3 x Ba2 and 2 x PO4339 I I 1 Ionic Compound Names and Formulas What is the charge on each of the ions 2 X Fe3 6 Fe2so43 3 x 5042 6 Li2C03 2 x Li 2 1 x c032 2 What is the formula for each ionic compound aluminum sulfide AIZS3 titaniumIV hydroxide TiOH4 ammonium phosphate NH43PO4 Acids and Ionization HX H x The acid appears to ionize or separate into ions To form H the neutral hydrogen atom must lose an electron Now the hydrogen ion consists of only one proton Acids may be classified by the number of H ions a single molecule can lose HX gt H X39 monoprotic HZY gt2 H YZ39 diprotic H32 gt3 H Z339 triprotic Polyprotic Acids Polyprotic acids do not lose all their hydrogens at once Upon dissolution in water the ionization occurs stepwise Each intermediate anion formed is a stable chemical species Example only 1 hydrogen ionized now the second hydrogen ionized H3PO4 H H2P0439 gtH HPO4239 phosphoric dihydrogen hydrogen acid phosphate phosphate 1 charge balances H charge H P043 3 charge balances loss of 3 total H phOSphate i0 Naming Binary Acids Hydrogen and a nonmetal NO OXYGEN Rules for Naming Binary Acids 1 Name hydrogen first using hydro 2 Then name nonmetal using ic ending Examples HBr hydrobromic acid HF hydrofluoric acid HI hydroiodic acid These are the names when the acids are dissolved in water Naming Binary Acids Periodic Trends H25 hydrosulfuric acid HZSe hydroselenic acid HzTe hydrotelluric acid Resulting Anions from First Ionization HS39 hydrogen sulfide ion HSe39 hydrogen selenide ion HTe39 hydrogen telluride ion Oxyacids and Oxyanions An acid that contains oxygen in addition to other elements is an oxyacid When H is removed from an oxyacid the oxygen stays with the nonmetal as part of an oxyanion HNO3 gt H NO339 nitric acid nitrate ion Rules for Naming Oxyacids Ending in ic The formula of the anion is the formula of the acid without the hydrogens There is a negative charge equal to the number of ionizable hydrogens in the acid The name of the anion is the name of the central element of the acid changed to end in ate Oxyacids to Know and Love Acid Ionization Equation Ion Name Chloric acid Nitric acid Sulfuric acid Carbonic acidtilt Phosphoric acid i HCio3 a H C10 HNO3 gt w No 142504 gt 2 in 3043 HZCO3 a 2 H co H3PO4 gt 3 H P043 Chlorate ion Nitrate ion Sulfate ion Carbonate ion Phosphate ion The charge on the anion relates to how many hydrogen atoms are in the acid formula The ion names for these acids all end with ate Acid Names Relate to Periodic Groups Sulfuric acid is composed of hydrogen oxygen and sulfur a nonmetal from Group 6 What other elements are found in Group 6 selenium Se tellurium Te Based on our knowledge of acids what are the name and formula for the oxyacids formed with Se and Te selenic acid HZSeO4 selenate ion SeO4239 telluric acid H2TeO4 tellurate ion TeO4239 What can we say about oxyacids formed from halogens chloric acid HCIO3 chlorate ion CIO339 iodic acid HIO3 iodate ion IO339 Change in the Number of Oxygens The number of oxygen atoms in oxyacids can change The ion suffix changes as a result The ion charge does NOT change 504239 sulfate ion 503239 sulfite ion NO339 nitrate ion NOZ39 nitrite ion Oxyacids formed from halogens can gain or lose four oxygens The ion prefix and suffix change as a result BrO439 perbromate ion BrO339 bromate ion BrOZ39 bromite ion BrO39 hypobromite ion Change in the Number of Oxygens Acid Ionization Equation Anion Name Perchloric acid HCIO4 gt w ClO4 Perchlorate ion Chloric acid l lClO3 a H C10 Chlorate ion Chlorous acid HClO2 gt H C102 Chlorite ion Hypochlorous acid HClO gt H ClO Hypochlorite ion Hydrochloric acid HCl a H Cl Chloride ion prefix suffix oxyanions per ate XO439 ate XO339 NO339 CO3239 5042 P043 ite XOZ39 N0239 5032 hypo ite XO39 X F Cl Br I Ionic Compounds Acid Anions What is the name of each ionic compound cr3PO42 NaHCO3 AIBrO3 CaH2PO42 K104 chromiumII phosphate sodium hydrogen carbonate aluminum hypobromite calcium dihydrogen phosphate potassium periodate Properties of Ionic Compounds When a substance having a negative charge is brought near a substance having a positive charge there is a force of attraction between them Electrostatic Force Force of attraction or repulsion caused by electric charges 1 n 1 n 1 5 gt gt 39 Li F LiF Coulomb s Law charge on and ions charge on electron k nen39e d2 I distance between ions Force of Attraction proportionality constant Electrostatic Force As the ion charges increase the attraction between ions increases 1 41 f2 Is inn charge increases farm of attraction mcreasos As the distance between ions becomes smaller the attraction between ions increases As distance increases force of attraction decreases Electrostatic Force Ionic compounds do not consist of simple pairs of ions An ionic solid is made up of millions of ions arranged in an extended 3D network called a crystal lattice Each ion is surrounded by oppositely charged ions and is held tightly in its allocated position Molecular Compounds Molecular or covalent compounds are typically formed between two nonmetals or a nonmetal and a metalloid While ionic compounds are generally solids at RT molecular compounds range from solids to liquids to gases The bonding electron pairs are shared between the atoms no ionization occurs Rules for Naming The first word is the name of the element appearing first in the chemical formula including a prefix to indicate the number of atoms of that element in the molecule The second word is the name of the element appearing second in the chemical formula changed to end in ide including a prefix to indicate the number of atoms of that element in the molecule Numerical Prefixes Used in Naming Number Prefix Number Prefix 1 mono 6 hexa 2 di 7 hepta 3 tri 8 octa 4 tetra 9 nona 5 penta 10 deca The letter 0 in mono and the letter a in prefixes four through ten are omitted if the resulting word sounds better 05 is pentoxide NOT pentaoxide NBr3 nitrogen tribromide PCI5 phosphorus pentachloride SZF10 disulfur decafluoride Common Molecular Compounds Compound Common Name Compound Common Name CH4 methane NZH4 hydrazine CZH6 ethane PH3 phosphine C3H8 propane NO nitric oxide C4H10 butane N20 nitrous oxide NH3 ammonia H20 water The Mole Avogadro s Number When talking about doughnuts we can use a dozen as our unit However atoms and molecules are tiny so we need a MUCH large number Chemists use the mole mol as a way to count atoms or molecules 1 mol 602 x 1023 atoms elements 1 mol 902 x 1023I molecules compounds Avogadro s number Molar Mass Express chemical quantities at the macroscopic level Molar mass mass in grams of one mol of a substance Copper Cu The atomic mass of one Cu atom is 6355 amu 1 mol of copper atoms weighs 6355 g The molar mass of copper is 6355 gmol What is the molar mass of methane Formula CH4 CH4 1201 gmol 41008 gmol 1604 gmol Using the Mole What mass of lead is equivalent to 250 moles of lead 250M x 2072 g Pb 518 g Pb 1W How many moles of tin are in 366 g of tin How many atoms of tin are in the sample 366g Trx 1 mol Sn 0308 moles Sn 11871g 8n 0308310 x 602 x 1023 atoms Sn 185 x 1023 atoms Sn 1W Using the Mole A tablet of aspirin contains 325 mg of the drug How many moles of aspirin have you ingested What does one molecule of aspirin weigh C9H8O4 91201 gmol 81008 gmol 41600 gmol 1802 gmol OTOH 0325 W x 1 mol C9H8O 000180 mol chgo4 CCxcKCCHB 180 II I I HCCCH O L 1 602 x 1023 molecules Using the Mole How many atoms of carbon are in 165 g oxalic acid oxalic acid H2C204 165 2 4x 1 x 602 x 1023 mole ules 110 x 1023 molecules 903W 1 m 2 4 110 x 1023mmLQeules x 2 C atoms 220 x1023 C atoms 1 How many atoms of oxygen 110 x 1023leQ UIESX 4 O atoms 440 x 1023 O atoms lizizwaieccrrer Percent Composition The percentage by mass of each element in the compound Calculate the percentage of each element in NaZSO4 Calculate the molar mass NaZSO4 22299 gmol 3207 gmol 41600 gmol 14205 gmol Determine the percent composition 00 Na g 3820 100 W 100 3237 00 00 s g Nizsso4x 100 1 100 2258 00 g 0 4 X 53900 g 0 x 100 4505 00 O g NaZSO4X 100 14205 9 NaZSO4 Percent Composition What is the percent mass of each element in propane What mass of carbon is contained in 454 g of propane Calculate the molar mass of C3H8 C3H8 31201 gmol 81008 gmol 4409 gmol Determine the percent mass of each element g C 3 X 1201 g C O o g H 8 X 1008 g H o 0 Mass of carbon in 454 g propane 4549 63H839 x 817ng 371 9C 1 8 The Empirical Formula The empirical formula shows the simplest whole number ratio of atoms in a formula Eugenol is a major component in oil It is composed of 7314 0o C 737 0o H and the rest 0 What is the EF Convert the percentages into moles of each element 7314ng X M 6089 mol C 1201 ng 737944 X 1 mol H 731 mol H 100894 194996 X 1 mol 0 1218 mol 0 160090 Express the moles as the smallest possible ratio 6089 500 731 600 1218 100 1218 1218 1218 Write the EF using the values above as subscripts c H 0 Must have a whole 5 5 number ratio The Molecular Formula The molecular formula is the actual ratio of atoms in a compound Eugenol has an EF of C5H60 and a molar mass of 1642 gmol Calculate the mass of the EF C5H6O 51201 gmol 61008 gmol 1600 gmol 8210 gmol Determine the number of EF units in the molecule Molar mass compound 1642 mol 2 Molar mass EF 8210 gmol Write the molecular formula C5H6O2 C10quotI1202 The molar mass must be known to find the MF Formula Determination From Mass Data Gallium oxide GaxOy forms when gallium is combined with oxygen If 125 g Ga reacts with oxygen to form 168 g GaxOy what is the formula of the product Determine the amount of oxygen in the product 168 g GaxOy 125 g Ga 043 g oxygen Convert the masses into moles of each element 125 9825 X 1 mol Ga 00179 mol Ga 6972 nga 04396 X 1 mol 0 0027 mol 0 1600 96 Express the moles as the smallest possible ratio 00179 100 0027 150 00179 00179 Write the EF using the values above as subscripts Must have a whole G3015 Ga203 number ratio Hyd rates Some molecules form solids that include water molecules as part of the crystal structure The compounds are called hydrates Water can be removed from the compounds by heating forming anhydrous compounds Example CoCl2 anhydrous is a blue compound CoCl2 can combine with 6 waters to form a hydrate CoCl2 39 H20 is a deep red solid number of water molecules in a hydrate is shown after the anhydrous formula separated by a dot Equation for Dehydration 396 H20 a 6 H20 Naming Hydrates Rules 1 Name the anhydrous compound 2 Use a prefix to indicate the number of water molecules followed by the word hydrate Examples CuSO4 395 H20 copperII sulfate pentahydrate NaZCO3 10 H20 sodium carbonate decahydrate nickel chloride hexahydrate NiCl2 6 H20 Formulas for Hydrated Compounds To determine the number of water molecules bound to CuSO4 you weigh out 1023 g of the hydrate After heating thoroughly only 0654 g of anhydrous CuSO4 remain How many water molecules were bound Mass of hydrate 1023 g Mass of CuSO4 0654 g Mass of water 0369 g Convert the masses into moles of each compound 0654532550 x 1 mol Cuso4 000410 mol Cuso4 1 4 03699436 x 1 mol H20 00205 mol H20 1 I 2 Make a ratio of the values and determine the EF Cuso4 5 H20