CH 101 TEST #2 STUDY GUIDE
CH 101 TEST #2 STUDY GUIDE CH 101
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CH 101 Test 2 Study Guide Ch 4 5 Chapter 4 Finding Patterns The Periodic Law and the Periodic Table 0 In 1869 Mendeleev noticed that certain groups of elements had similar properties 0 He found that when elements were listed in order of increasing mass these similar properties recurred in a periodic pattern 0 To be periodic means to exhibit a repeating pattern The Periodic Law Mendeleev summarized these observations in the periodic law 0 When the elements are arranged in order of increasing mass certain sets of properties recur periodicaly The Periodic Law 18 Ar 1 Ne 11 Na l Elements with similar properties recur in a regular pattern J 12 Mg 13 Al M St 15 Pl re S 17 Cl 19 K 20 Ca 5 6 7 a 9 B C N O F Mendeleev s Periodic Table Mendeleev s periodic table 0 Organized known elements of the time in a table format He arranged the rows so that elements with similar properties would fall in the same vertical columns 0 Contained some gaps which allowed him to predict the existence and even the properties of yet undiscovered elements Mendeleev predicted the existence of an element he called ekasilicon Gallium lelltaauminuml Germanium lekasllltoul Mendeleev s Meuclleleev s Actual prop erties predlctecl properties predicted Actual properties properties Atomic mass About 68 amu 6972 arrlu Atomic mass About 72 amu 7264 amu Melting point lLow 298 C Density 55 gicm3 535 gicm3 Density 59 gicm3h 590 gicm3h Formula of oxidle X02 Geog Formula of oxide X203 33203 Formula of clrloricle XCM GeCl4 Formula of chloride XCl3 GaC13 The Modern Periodic Table Its Format 0 The elements are listed in order ofincreasing atomic number rather than increasing relative mass as they were in Mendeleev s periodic table 0 Rows of the table are referred to as periods 0 Columns in the table are referred to as groups or a family 0 Elements in a group or family have similar properties 0 NOTE Mendeleev s periodic law predicts pattern but does NOT explain Why the patterns or similarity in properties occurs Quantum theory explains the Why 0 Elements in the periodic table are classi ed as the following o Metals o Nonmetals o Metalloids o The periodic table can also be divided into 0 Maingroup elements whose properties tend to be largely predictable based on their position in the periodic table In the periodic table this area is labeled by a number and the letter A 0 Transition elements and inner transition metals or transition metals whose properties tend to be less predictable based simply on their position in the periodic table In the periodic table this area is labeled by a number and the letter B Electron Con guration How an Atom s Electrons Occupy Orbitals Quantummechanical theory describes the behavior of electrons in atoms 0 The electrons in atoms exist in orbitals A description of the orbitals occupied by electrons is called an electron con guration H 151 Number of electrons in orbital Orbital Electron Con guration and Quantum Theory Connection Schrodinger s equation showed that hydrogen s one electron occupies the lowest energy orbital in the atom For multieectron atoms the equation cannot be exactly solved because of the for electron electron interactions that happen between two electrons However approximate solutions showed the orbitals to be hydrogenlike 0 Two additional concepts affect multielectron atoms electron spin and energy splitting of sublevels To understand electron arrangement around an atom s nucleus we need to account for the effects of electron spin m5 quantum number Electron Spin and the Pauli Exclusion Principle 0 Spin is a fundamental property of all electrons All electrons have the same amount of spin 0 The orientation of the electron spin is quantized it can be only in one direction or its opposite 0 Spin up or spin down 0 The electron s spin adds a fourth quantum number to the description of electrons in an atom called the spin quantum number m5 0 The spin quantum number is not part of the Schrodinger equation ms can have values of 12 or 12 Orbital diagrams use a square to represent each orbital and a halfarrow to represent each electron in the orbital By convention a halfarrow pointing up is used to represent an electron in an orbital with spin Up a halfarrow pointing down is used to represent an electron in an orbital with spin down Spins must cancel in an orbital Paired meaning the two electrons in an orbital must have opposite spins ie one with the magnetic eld the other against the magnetic eld The Pauli Exclusion Principle No two electrons in an atom may have the same set of four quantum numbers Therefore no orbital may have more than two electrons and they must have opposite spins Example Helium He Electron con guration Orbital diagram He ills2 1 L 15 n I m mS 39l l 0 0 5 l l 0 0 T 2 Sublevel Energy Splitting in Multielectron Atoms The sublevels in each principal energy shell of hydrogen all have the same energy or other single electron systems Orbitals with the same energy E are said to be degenerate For multielectron atoms the energies of the sublevels are split 0 Caused by charge interaction shielding and penetration The lower the value of the quantum number orbital quantum number the less energy the sublevel has E s orbital I 0 lt E p orbital I 1 lt E d orbital I 2 lt E forbital I 3 Coulomb s Law 32 i QED 4W30 F Coulomb s law describes the attractions and repulsions between charged particles 0 For like charges the potential energy E is positive and decreases as the particles get farther apart as rincreases o For opposite charges the potential energy is negative and becomes more negative as the particles get closer together 0 The strength of the interaction increases as the size of the charges increases Electrons are more strongly attracted to a nucleus with a 2 charge than to a nucleus with a 1 charge Shielding and Effective Nuclear Charge 0 Each electron in a multielectron atom experiences both the attraction to the nucleus and repulsion by other electrons in the atom These repulsions cause the electron to have a net reduced attraction to the nucleus it is shielded from the nucleus 0 The total amount of attraction that an electron feels for the nucleus is called the effective nuclear charge of the electron Shielding and Penetration 0 The closer an electron is to the nucleus the more attraction it experiences 0 The better an outer electron is at penetrating through the electron cloud of inner electrons the more attraction it will have for the nucleus o The degree of penetration is related to the orbital s radial distribution function 0 In particular the distance the maxima of the function are from the nucleus Experiences 6 62 6 1 full 3 clharge 39 D Q e Nucleus Experiences Nucleus net charge 0 of about 1 a e e a b Electron Spatial Distribution and Sublevel Splitting The radial distribution function shows that the 25 orbital penetrates more deeply into the 15 orbital than does the 2p The weaker penetration of the 2p sublevel means that electrons in the 2p sublevel experience more repulsive force they are more shielded from the attractive force of the nucleus The deeper penetration of the 25 electrons means electrons in the 25 sublevel experience a greater attractive force to the nucleus and are not shielded as effectively 15 2P 25 Penetration of25 Total radial probability I l l 0 200 400 600 800 r pm Shielding and Penetration Penetration causes the energies of sublevels in the same principal level to not be degenerate In the fourth and fth principal levels the effects of penetration become so important that the 5 orbital lies lower in energy than the dorbitals of the previous principal level The energy separations between one set of orbitals and the next become smaller beyond the 45 orbital The ordering can therefore vary among elements causing variations in the electron con gurations of the transition metals and their ions Penetration Distance from nucleus General Energy Ordering of Orbitals for MultiElectron Atoms Total radial probability General Energy Ordering of Orbitals for Mlultiaelectron Atoms 3p 3p 3p E Energy 3 E E F Electron Con guration for MultiElectron Atoms o Aufbau Principle 0 Energy levels and sublevels ll from lowest energy to highest I 5 p gt 1 f Electron con guration Orbital diagram Li 152251 1L 1 15 25 o Orbitals that are in the same sublevel have the same energy 0 There can be no more than two electrons per orbital Pauli exclusion principle Electron configuration Orbital diagram C 1522522102 1L ll 1 l 15 25 2p Hund s Rule 0 When lling orbitals that have the same energy degenerate place one electron in each orbital before completing pairs Electron con guration Orbital diagram C 1522522192 ll ii i l is 25 2 p Summarizing the Filling of Electrons in Atomic Orbitals Electrons occupy orbitals so as to minimize the energy of the atom therefore lower energy orbitals ll before higher energy orbitals o Orbitals ll in the following order 15252p353p453d4p554d5p65 Orbitals can hold no more than two electrons each When two electrons occupy the same orbital their spins are opposite o This is another way of expressing the Pauli exclusion principle no two electrons in one atom can have the same four quantum numbers When orbitals of identical energy are available electrons rst occupy these orbitals singly with parallel spins rather than in pairs Hund s rule 0 Once the orbitals of equal energy are halffull the electrons start to pair Electron Con guration Valence Electrons and the Periodic Table The electrons in all the sublevels with the highest principal energy shell are called the valence electrons Electrons in lower energy shells are called core electrons One of the most important factors in the way an atom behaves both chemically and physically is the number of valence electrons Si 152252219635232 Core Valence electrons electrons Core Electrons Valence Electrons and the Periodic Table The group number corresponds to the number of valence electrons The length of each quotblockquot is the maximum number of electrons the sublevel can hold The period number corresponds to the principal energy level of the valence electrons Outer Electron Configurations of Elements 118 1A 8A 1 2 H He Is1 2A 3A 4A 5A 6A 7A 152 3 4 5 6 7 8 9 10 Li Be B C N 0 F Ne 25 252 2522191 2522122 2522123 2522194 2522p 2522136 11 12 13 14 15 16 17 18 Na Mg Al Si P 3 Cl Ar 3539 352 3523pJl 3523192 3523193 3523194 3523105 31523136 Orbital Blocks and Their Position in the Periodic Table Orbital Blocks of the Periodic Table Groups 1 18 1A 8A 1 5 block elements block elements 2 1 H 2 D 1 13 14 15 16 17 He 1 15 2A B dblockeleme11ts Elfcblock elements 3A 4A 5A 6A 7A 152 3 4 5 6 7 3 9 10 2 Li Be 13 c N o 15 Ne 241 232 2422p1 2922 2322123 2532pquot 252245 2522116 11 12 13 14 15 15 17 1s 3 Na Mg 3 4 5 6 7 8 9 10 11 42 All 51 P 5 c1 Ar 351 351 318 4B 5B 613 713 l7 8B I 1B 23 3523p39 383 3533113 353391 3423193 3523p quotE 19 211 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 E 4 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 3 451 452 452341 4523412 452343 4513715 452325 452345 4523237 4523113 463le 4523471 452441 452442 4534p3 45241 l 452455 4524156 37 3s 39 4o 41 42 43 44 45 46 47 43 49 50 51 52 53 54 5 Rb Sr Y Zr Nb Mo Tc R11 Rh lPd Ag Cd 11 Su SI Te I Xe 551 551 5524511 552422 551424 5514415 5524515 5514127 5514123 44110 551421 55245110 552591 551552 55354J 551544 5525pS 5525196 55 55 57 72 73 74 75 76 77 78 79 so 81 32 33 94 35 86 6 C5 Ba La Hf Ta W Re OS 1139 Pt Au Hg Tl Plb Bi Po At R11 1551 652 651591 633542 652543 2544 652545 552545 531547 551549 65151110 15525410 552641 632692 gs25133 651604 6526 652596 7 87 88 89 104 105 1015 107 108 109 110 111 112 113 114 115 116 117 11s 7 Fr Ra Ac Rf Db 5g Bh Hs Mt Ds Rg C11 9 Fl 393 Lv 7 7 73 752 751641 72612 752543 752644 53 59 60 61 62 63 154 65 66 67 63 69 711 71 lisarlthanid es Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 5524f 5d 15524r3 65241r4 6524f 651416 65247 ell yew 6524 5514 390 55147 4524712 6524f 6324f 6524f3946d39 911 91 92 93 94 95 96 97 98 99 100 101 1112 1113 Actjnides Th 31 U Np Pu An i Cm Bk Cf Es Fm Md N0 Lr 732642 7535f164l 75251455111 751519541 7425 7525f 752919541 7525 75251 7325f 795 785 7535f 735f3945439 Summarizing Periodic Table Organization The periodic table is divisible into four blocks corresponding to the lling of the four quantum sublevels 5 p d and f o The group number of a maingroup element is equal to the number of valence electrons for that element The row number of a maingroup element is equal to the highest principal quantum number of that element Transition and Inner Transition Metals 7739an5iti0n meta5 0 block and inner transition meta5 fblock exhibit trends differing from those of maingroup elements 5 block and 0 block 0 Because of sublevel splitting the 45 sublevel is lower in energy than the 3d sublevel therefore the 45 orbital lls before the 3d orbital o The difference in energy is not large 0 Some of the transition metals have irregular electron con gurations in which the n5 only partially lls before the n 1d or doesn t ll at all 0 Therefore their electron con guration must be found experimentally Examples of Transition and Inner Transition Metals Electron Con gurations Expected Found Experimentaly Cr Ar45230 4 Cr Ar45130l5 Cu Ar45230 9 Cu Ar45130 10 Mo Kr55240 4 Mo Kr55140 5 Ru Kr5524a6 Ru Kr55140 7 Pd Kr552408 Pd Kr55040 10 Electron Con guration and Elemental Properties 0 The properties of the elements follow a periodic pattern 0 Elements in the same column have similar properties 0 The elements in a period show a pattern that repeats The quantummechanical model explains this because the number of valence electrons and the types of orbitals they occupy are also periodic Electron Con guration and Elemental Properties Noble Gases The noble gases have eight valence electrons Except for He which has only two electrons They are especially nonreactive He and Ne are practically inert The reason the noble gases are so nonreactive is that the electron con guration of the noble gases is especially stable Electron Con guration and Elemental Properties The Metals Metallic elements make up the majority of the elements in the periodic table Alkali Metals They have one more electron than the previous noble gas and occupy the rst column In their reactions the alkali metals lose one electron and the resulting electron con guration is the same as that of a noble gas Forming a cation with a 1 charge Alkaline Earth Metals They have two more electrons than the previous noble gas and occupy the second column In their reactions the alkaline earth metals lose two electrons and the resulting electron con guration is the same as that of a noble gas Forming a cation with a 2 charge Transition and Inner Transition Metals They are located in the d block area of the periodic table In chemical reactions they will lose electrons from 5 and then 0 orbitals to form cations pblock metals They are located in the p block area lefthand side of the metalloids of the periodic table o In chemical reactions they will lose electrons from the 5 and p orbitals to form cations Mlajor Divisions of the Periodic Table Silicon Carbon lj Metals j Metalloids 39 Nonmetals 1A 1 I Chromium j MOI31361 Sulfur 1 2A St1 011t1u11391 301d 2 r 1 H 39339 H3 l ls 2 152 V r 3 4 39 393 2 5 2 10 2 1 Be U N 0 e l 2s 252 52219 252211 222p 2522155 11 12 v 15 15 12 15 Bramine 3 Na Mg SB 6B 1quot S C1 A1 i 35quot 352 S 6 523173 3523114 3523175 3523116 i ll 19 211 23 24 133 34 35 35 4 K Ca l l V Cr l 311 21111 Ga Ge As Se Br 4539 451 452551 4533112 4523513 4513215 4513515 415255 45231 41533113 4151351 l 4513211 452451 4524p 4524113 452454 452455 452405 11 3239 53 39 111 41 12 15 14 1395 45 12 13 9 50 51 52 53 54 3 Rib 5139 Y Zr Ni M0 Tc lRu Rh 11 Ag Cd 1 Sn Sb Te ll Xe 5s 552 5132151 522152 5114214 5114215 5124215 5915 59153 421 553945390 5524 510 51355 55352 5125153 525p4 513555 ifsp 1 0de 55 55 5 2392 2395 2394 25 25 2392 23 251 51 51 52 53 34 35 55 l 6 C5 Ba La Hf Ta W Re 05 1139 Pl A1 Hg Tl Pb Bi Po At Rm 5539 533 55351139 5525211 5525213 5525214 5315115 5515116 551511 551511quot 5515111 552551 5325321 5525132 5535133 552551 551511 552556 i 7 8239 83 59 1114 1115 1115 10239 1011 109 1110 111 112 113 114 115 115 11 118 l 7 Fr Ra A1 Rf Db Sg Bill H5 MI 5 Kg C111 1 F H Lv 1 1 F539 152 T5265 1 T5313ch2 3 5 2amp151quotF 1739s 2554 53 59 50 51 52 53 54 55 55 5239 53 59 211 21 Lanthanides Ce Pr Nd Pm Sm lEILl Gdl Tb Dy Ho lEr Tm Yb Lu 6324f15d1 5524f1 5531f 5sla1f5 5521156 5334f 524rF551 5522959 5334 W 5322115 5524fl2 6524f 5534145324fl1551 911 91 52 515 514 95 95 92 95 951 11111 1111 1112 1113 Actinides Th Pa U Np P111 Am C111 Bk Cf Es F1111 Md N0 M 253553 252514551 25233551 2525f395539 252515 25251 25351 551 7515 2125 1351 2535f 2515f 2525f 2535111551 Behavior and Electron Con guration of Metalloids and Nonmetals Metalloids 0 They are located in the d block area of the periodic table between the metal and nonmetal elements Sitting on the steps of the zigzag diagonal line indicated on the periodic table 0 Metalloids in chemical reactions can exhibit metallic or nonmetallic behaviors o Metalloids can either lose electrons from 5 and then p orbitals to form cations or gain electrons into their 0 orbitals to form anions Nonmetals 0 They are located in the upper righthand side of the periodic table p block area 0 ln chemical reactions nonmetal elements will gain electrons into the p orbitals resulting in their ions having the same electron con guration as a noble gas at the end of their period row 0 Nonmetals form anions Electron Con guration and Elemental Properties The Halogens Halogens 0 They are nonmetals 0 They have one fewer electron than the next noble gas 0 In their reactions with metals the halogens tend to gain an electron and attain the electron con guration of the next noble gas forming an anion with charge 1 o In their reactions with nonmetals they tend to share electrons with the other nonmetal so that each attains the electron con guration of a noble gas Electron Con guration and Ion Formation Ion formation can be predicted by an element s location in the periodic table These atoms form ions that will result in an electron con guration that is the same as that of the nearest noble gas Metals form cations positively charged atoms Alkali metals group 1A form only 1 cations Alkaline earth metals group 2A form only 2 cations Transition inner transition and oblock metals form a variety of charged cations Nonmetals form anions negatively charged atoms Halogens group 7A usually gain one electron to form 1 anions Other nonmetals can form a variety of charged anions Elements That Form Ions with lPredictable Charges 1A 2A 3A 4A 5A 6A 7A 8A 1 11 N3 02 r 2 3 2 2 Na Mg 33 413 53 6B 7B gsra 1E 23 A S 31 3 K Ca2 8amp2 Bra 4 1119 59 Telquot IE 5 38 Ba Periodic Trends Atomic Radii and Effective Nuclear Charge 0 There are several methods for measuring the radius of an atom and they give slightly different numbers 0 Van der Waals radius nonbonding o Covalent radius bonding radius 0 Atomic radius is an average radius of an atom based on measuring large numbers of elements and compounds L van dler Waals radius E Covalent radius j 2 X Krypton radius 228 pm Krypton solid Br radlus 2 114 pm 0 Atomic rao ius decreases across a period left to right 0 Adding electrons to the same valence shell 0 Effective nuclear charge increases 0 Valence shell held closer 0 Atomic rao ius increases down a group o Valence shell farther from nucleus 0 Effective nuclear charge fairly close Atomic Radii 300 I l Alkalimetals Is l A 25 Rb I 250 I K Period 5 Period 4 transition transition elements 21 l U Tf d 39 elements a l l 2 1150 39 6393 d A 39 u Ke loo Kquot W h ALI V 39 I Ne 50E 7 7 H0 quot Noble gases 0 I 390 l0 2039 30 40 50 60 Atomic number Effective Nuclear Charge The effective nuclear charge is a net positive charge that is attracting a particular electron Core electrons ef ciently shield electrons in the outermost principal energy level from nuclear charge Outermost electrons in the valence shell do not ef ciently shield one another from nuclear charge Zis the nuclear charge and Sis the number of electrons in lower energy levels Electrons in the same energy level contribute to screening but since their contribution is so small they are not part of the calculation Trend issgtpgtdgtf Zeffective Z 5 Screening and Effectilve lNucllear Charge e Valence 251 electron 6 Core 152 electron Effective nuclear charge gt a 3 2 is 1 NuclelN Nucleus 3 J Lithium Summarizing Atomic Radii Trend for MainGroup Elements The size of an atom is related to the distance the valence electrons are from the nucleus o The larger the orbital an electron is in the farther its most probable distance will be from the nucleus and the less attraction it will have for the nucleus 0 Traversing down a group adds a principal energy level and the larger the principal energy level an orbital is in the larger its volume 0 Quantummechanics predicts that the atoms should get larger down a column 0 The larger the effective nuclear charge an electron experiences the stronger the attraction it Will have for the nucleus 0 The stronger the attraction the valence electrons have for the nucleus the closer their average distance will be to the nucleus 0 Traversing across a period increases the effective nuclear charge on the valence electrons o Quantummechanics predicts that the atoms should get smaller across a pedod Summarizing Atomic Radii Trend for Transition Elements Atoms in the same group increase in size down the column Atomic radii of transition metals are roughly the same size across the 0 block Much less difference than across main group elements Valence shell n52 not the n 10l electrons Effective nuclear charge on the n52 electrons approximately the same Ions Magnetic Properties Electron con gurations that result in unpaired electrons mean that the atom or ion will have a net magnetic eld this is called paramagnetism Will be attracted to a magnetic eld Ag Kr 5514031 i ll ll ll ll ll 55 4 Electron con gurations that result in all paired electrons mean that the atom or ion will have no magnetic eld this is called diamagnetism Slightly repele0l by a magnetic eld Zn Ar 4323611 ll ll ll ll ll ll 45 3d Radii of Atoms and Their Ions l Cations Cation radius is smaller than its corresponding atom radius The loss of electrons results in the remaining electrons in the atom to experience a larger effective nuclear charge than the neutral atom Traversing down a group increases the n 1 level causing the cations to get larger Traversing to the right across a period increases the effective nuclear charge for isoelectronic cations causing the cations to get smaller Radii of Atoms and Their Cations pm Group 1A Group 2A Group SA Be 1332 0 152 61 112 31 o a 5 m 85 23 41 Na Na Mg Mg A1 Art 136 95 160 65 143 50 K K 227 133 Rb Rb 2 18 148 0 Sr 5 n Radii of Atoms and Their Ions l Anions When atoms form anions electrons are added to the valence shell These new valence electronsquot experience a smaller effective nuclear charge than the old valence electronsquot The result is that the anion is larger than the atom Traversing down a group increases the 17 level causing the anions to get larger Traversing to the right across a period decreases the effective nuclear charge for isoelectronic anions causing the anions to get larger Radii of Atoms and Their Anions pm Group 6A Group 7A 0 03 if 73 140 72 136 C1 C1 99 r 14 1 I 33 B 1 1 6 CD 181 litquot 95 Jr 216 143 221 Ions Ionic Radii Summary Ions in the same group have the same charge Ion size increases down the column Higher valence shell larger Cations are smaller than neutral atoms anions are larger than neutral atoms Cations are smaller than anions Except Rb and Cs which are bigger than or the same size as F and 02 Larger positive charge smaller cation For isoelectronic species lsoeectronic same electron con guration Larger negative charge larger anion For isoelectronic species Periodic Trend Ionization Energy Potential Ionization Energy IE o It is the minimum energy needed to remove an electron from an atom or ion in the gas phase It is an endothermic process requires the input of energy to remove the electron Valence electron easiest to remove lowest IE First ionization energy energy to remove electron from neutral atom All atoms have rst ionization energy Mg IE1 a M1g 1 e39 0 Second lE energy to remove from 1 ion etc M1g IE2 a M2g 1 e 0 First Ionization Energies 2500 He 1 39P a I Ne Noble gases 2000 1 z 1 To l E l Ar 3 l I 0 1500 i 1 g l l a I Periool 4 K1 Period 5 M l 1 transition 39i transition X i i i 1 elements 1 elements 2 1000 l i 7 i i 7 i 7 7 Vi 1 ll l I 7 ll r i li l l I 39 E l l W l M I ll I l l v i ML l 1 Na 1 Rb Alkali metals l 0 l O 10 20 30 40 50 Atomic number The larger the effective nuclear charge on the electron to be removed the more energy it takes to remove it The farther the most probable distance the electron is from the nucleus the less energy it takes to remove it 0 Trend 0 First lE decreases down the group Valence electron is farther from nucleus 0 First lE generally increases across the period Effective nuclear charge increases Summary on First Ionization Energy for MainGroup Elements 0 First ionization energy generally 0 decreases as we move down a column or family in the periodic table because electrons in the outermost principal level are increasingly farther away from the positively charged nucleus and are therefore held less tightly 0 increases as we move across left to right a row or period in the periodic table because electrons in the outermost principal energy level experience a greater effective nuclear charge Zeff The strength of attraction is related to the most probable distance the valence electrons are from the nucleus and the effective nuclear charge the valence electrons experience 0 The larger the orbital an electron is in the farther its most probable distance will be from the nucleus and the less attraction it will have for the nucleus Quantummechanics predicts that the atom s rst ionization energy should get 0 lower down a column Traversing across a period increases the effective nuclear charge on the valence electrons 0 larger across a period First Ionization Energy Exceptions to the Trend GENERAL trend for rst ionization energy of maingroup elements is that as you go across a period ionization energy increases 0 Exceptions 2A to 3A and 5A to 6A Exceptions are a usually a result of the type of o Orbital 5 p d or f and its shielding ability 0 Repulsion factors associated with electrons occupying degenerate orbitals ie p orbitals Be ls 25 2p B ls 25 2p 0 B has smaller rst ionization energy than Be due to electron position 20 for B and 25 for Be 0 The electron in 20 orbitals has less shielding i e less effective nuclear charge factor and therefore requires less energy for its removal than an electron in a 25 orbital First ionization energy s GENERAL trend is that as you go across a period ionization energy increases N ls Zs 2p N ls Zs 2p 0 To ionize N you must break up a halffull sublevel which costs extra energy C ls Zs 2p 3 ls Zs 2p 0 When you ionize 0 you get a halffull sublevel which costs less energy Trends of Second and Successive Ionization Energies They depend on the number of valence electrons an element has 0 Group lA elements CANNOT have successive ionization potentials eg have only one valence electron 0 Group 2A can have second ionization potential but not third 0 Removal of each successive electron costs more energy 0 Shrinkage in size due to having more protons than electrons 0 Outer electrons closer to the nucleus therefore harder to remove There s a regular increase in energy for each successive valence electron Periodic Trend Electron Af nities EA for MainGroup Elements 0 Electron Af nity o It is the energy associated with the addition of an electron to the valence shell of an atom that is in the gas phase Mg l e a M1 g EA o It is de ned as exothermic the release of energy but may actually be endothermic in take of energy Some alkali earth metals and all noble gases electron af nities are endothermic oThe more energy that is released the larger the electron af nity The more negative the number the larger the EA 0 General Trend for MainGroup Elements 0 EA increases across a period EA becomes more negative from left to right Haogens have the highest EA for any period Periodic Trend Electron Af nities EA for MainGroup Elements Summarizing Electron Af nity for MainGroup Elements O 0 Most groups columns of the periodic table do not exhibit any de nite trend in electron af nity Among the group 1A metals however electron af nity becomes more positive as we move down the column adding an electron becomes less exothermic Alkali metals group 1A decrease electron af nity down the column Generally irregular increase in EA from second period to third period Group SA generally has lower EA than expected because extra electron must pair Groups 2A and 8A generally have very low EA because added electron goes into higher energy level or sublevel Electron af nity generally becomes more negative adding an electron becomes more exothermic as we move to the right across a period row in the periodic table Characteristics of Metals versus Nonmetals Metals Nonmetals Malleable and ductile Brittle in solid state Shiny lustrous Dull nonre ective solid re ect light surface Conduct heat and electricity Electrical and thermal Most oxides basic insulators and ionic Most oxides acidic and Form cations in solution molecular Lose electrons in reactions Form anions and polyatomic oxidized a n io n s Gain electrons in reactions reduced Periodic Trend Metallic Character properties of a metal 0 More malleable and ductile better conductor and easier to ionize Metallic character decreases left to right across a period Metallic character is how closely an element s properties match the ideal o Metals found at the left of the period and nonmetals to the right found at the bottom Metallic character increases down the column 0 Nonmetals found at the top of the middle main group elements and metals Trends in Metallic Character Metallic character decreases Metals Metalloids Du Nonmetals 4A 5A 6A 7A 8A 18 1 1 2A 3A H 2 1 31 14 15 16 17 He 3 2 3 4 5 6 7 8 9 10 g 11 Be 3 C N O F Ne EL 3 11 12 38 4B 3B 6B 7B IESB 113 23 13 14 15 16 17 18 3 Wm Na Mg 3 4 5 6 7 8 9 10 11 12 Al Si P 5 c1 Ar E 4 19 20 21 22 231 24 25 26 27 28 29 30 31 32 33 3 41 35 36 55 31 K Ca Sc T1 V 1 Mn Fe Co Ni Cu 211 Ga Ge As Se Br Kr S 5 37 38 3 9 10 411 42 43 44 45 46 47 48 49 50 51 52 53 54 g Rb Sr Y Zr Nb M0 Tc Ru R11 ch1 Ag Cd In Sn Sb 116 I Xe 3 3 6 55 56 57 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 5 CS Ba La Hf Ta W Re Os Ir Pt Au llllg Tl lPlb Bi Po At Rn 1 3 7 87 88 89 104 105 106 107 108 109 110 111 112 113 114 115 116 117 118 g 1E Ra Ac JRE Db Sg 311 ES Mt D s R3 C111 1131 Lv Lamhanide 7 58 59 60 61 62 63 64 65 66 67 68 69 70 71 39 Ce Pr Ndl Pm Sm Eu Gd Tb Dy H0 Er Tm Yb Lu Actmides 90 91 92 93 94 95 96 97 98 99 100 101 1 02 103 Th Pa U Nlp P11 Am Cm Bk Cf Es Fm Md No Lr Metallic Elements 0 Very low ionization energies found in seawater O Ionization energy decreases down the column Electron af nity decreases down the column Except for the noble gases metals generally have smaller rst ionization Good reducing agents easy to oxidize Very reactive not found uncombined in nature React with nonmetals to form salts Compounds generally soluble in water therefore metal ions are energies and nonmetals generally have larger electron af nities strongly decrease across a period because the valence electrons are held more strongly and the electron af nity increases 0 All very low MP for metals Except K Density increases down the column Atomic radius increases down the column Melting point decreases down the column Quantummechanics predicts that the atom s metallic character should increase down a column because the valence electrons are not held as Chapter 5 Elements to Molecules When two or more elements combine a molecule is formed Elements can be the same The molecular compound of oxygen 02 Elements can be different The molecular compound of water H20 Molecules are compounds The great diversity of substances that we nd in nature is a direct result of the ability of elements to form compounds Hydrogen Oxygen and Water The dramatic difference between the elements hydrogen and oxygen and the compound water is typical of the differences between elements and the compounds that they form When two or more elements combine to form a compound an entirely new substance results T i t39 selected Hydrogen Oxygen a er Properties Boiling Point 253 C 83 C 100 cquotC State at Room Gas Gas Liquid Temperature Flammability Expilosive Necessary Used to for extinguish combustion flame Law of De nite Proportion Formation of Molecules A hydrogen oxygen mixture can have any proportions of hydrogen and oxygen gas Water H20 has a de nite proportion of hydrogen to oxygen A water molecule always is composed of two hydrogen atoms to every one oxygen atom A ratio of 2 hydrogen atoms 1 oxygen atom Minutes and Camppundits i3939lll139lgill39l and hygmn i39ln39ilixtun r39lr39ir39 l39nr 54 Cum anumI 39I his can have any ratin niquot Mu Inuitmica 111m 1 air mail atquot l lju dwgun 5L trifrgt39u l lii i ugun l Elf511 i1I39II I uiI lu tIMquotgcn HI LL39 iU INII Types of Chemical Bonds Compounds are composed of atoms held together by chemical bonds Chemical bonds result from the attractions between the charged particles the electrons and protons that compose atoms Chemical bonds are classi ed into three types Ionic Covalent Metallic lonic versus Covalent Covalent bonding K if i I tquot d tt l H20 molecules Table salt Na lfs 0 Metal and Nonmetal Nonmetal and Nonmetal O Ionic Bonds Between a Metal Atom and Nonmetal Atom Ionic bonds occur between metals and nonmetals They involve the transfer of electrons from one atom to another When a metal interacts with a nonmetal it can transfer one or more of its electrons to the nonmetal The metal atom then becomes a cation The nonmetal atom becomes an anion O Ionic Bonds Ionic Compounds These oppositely charged ions attract one another by electrostatic forces and form an ionic bono The result is an ionic compound which in the solid phase is composed of a lattice Le a regular threedimensional array of alternating cations and anions O The Formation of an Ionic Compound quotSodium a metal Chlorine Ila nommelal loses an Electron 3 gain an EIECIIDFL v Nut militia Hulitrtillfl 6 1anl1e minim L7 li Nithlifhf1i lUi i39IIE39IIII1 V mlitm Radium Hull Chlorine gas 1 UPEGEETEI charged h 7 quot ions arel ueld together 7 lbw ionic tmritli forming v a crystalline la ttite Sodium chloridetfllllt O Covalent Bonds Bonds between Nonmetal Atoms Covalent bonds occur between two or more nonmetals They involve the sharing of electrons between two atoms When a nonmetal bonds with another nonmetal neither atom transfers its electron to the other Instead the bonding atoms share some of their electrons The covalently bound atoms compose a molecule Hence they are referred to as molecular compounds Molecular compounds are composed of atoms covalently bonded to each other Representing Compounds Chemical Formulas and Molecular Models A compound is represented with its chemical formula Chemical formula indicates the type and number of each element present in the compound Water is represented as H20 Carbon dioxide is represented as C02 Sodium chloride is represented as NaCI Carbon tetrachloride is represented as CC4 Types of Chemical Formulas Chemical formulas can generally be categorized into three different types Empirical formula Molecular formula Structural formula An empirical formula gives the relative number of atoms of each element in a compound 0 It is the simplest wholenumber ratio representation of the type and number of elements present in a molecule A molecular formula gives the actual number of atoms of each element in a molecule of a compound 0 For C4H8 the greatest common factor is 4 The empirical formula is therefore CH2 0 For BzH6 the greatest common factor is 2 The empirical formula is therefore BH3 o For CCI4 the only common factor is 1 so the empirical formula and the molecular formula are identical Types of Chemical Formulas Structural A structural formula is a sketch or diagram of how the atoms in the molecule are bonded to each other It uses lines to represent covalent bonds and shows how atoms in a molecule are connected or bonded to each other Example The structural formula for H202 is shown below H o o H Types of Chemical Formulas Summary The type of formula used depends on how much is known about the compound and how much information is to be communicated A structural formula communicates the most information o It conveys the type and actual number as well as the arrangement of the atoms in the molecule 0 It s a visual picture of the compound 0 A molecular formula conveys the actual type and number of elemental atoms in the compound o It does not tell you how each of the atoms are bonded to each other 0 An empirical formula communicates the least o It conveys the simplest wholenumber relative relationship of atom to atom in the moecue Molecular Models 3D Representation of a Molecule A molecular model is a more accurate and complete way to specify a compound A balland stick molecular model represents atoms as balls and chemical bonds as sticks how the two connect re ects a moecue s shape The balls are typically colorcoded to speci c elements d Hydrogen 0 Phosphorus Sulfur Chlorine O o In a space lling molecular model atoms ll the space between each other to more closely represent our best estimates for how a molecule might appear if scaled to visible size l1 CH4 H C H 7 Molecular formula Structural formula Ball andstick model Space lling model Ways of Representing a Compound TABLE 5 Benzene Acetylene quotjlurmsei and Ammonia Structural Formula Ballisanda Stick Model Molecular Formula SpacesFililing Model Name of Empirical Formula Compound H r I Benzene CH CEHIE I Hi 1 C c X H C H a g i 1 H c Acetylene CH CgHz lH CEC H Q U Oa i H OH quotJ I J7 HDCH Q e Glucose CH20 CEH1205 J V H t 0H oaro c t H2C 1 OH V Ammonia NH NH H H Q i so Lewis Structure Model Representing a Substance s Valence Electrons The Lewis Model 0 Valence electrons are represented as dots 6 dots representing 0 valence electrons Lewis electrondot structures Lewis structures depict the structural formula with its valence electrons Lewis structures focus on valence electrons because chemical bonding involves the transfer or sharing of valence electrons between two or more atoms Lewis Theory and Bonding The Lewis Model 0 It is used mainly to illustrate covalentbonded molecular compounds but can be used to illustrate simple ionic compounds Lewis Symbols 0 They can be used to represent the transfer of electrons from a metal atom group 1A group 3A metals to a nonmetal atom resulting in ions that are attracted to each other and therefore bond K 1522522p63523p6451 K 1522522p63523p6450 l O Octet in previous level Octet Rule A Guideline for Molecule Formation 0 When atoms bond they tend to gain lose or share electrons to result in a noble gas like con guration 0 nsznp6 Nonmetals period 2 elements must obey the octet rule ie eight valence electrons around each atom in the molecule 1 Exceptions to the octet rule expanded octets 0 They involve the nonmetals elements located in period 3 and below Nonmetals period 3 on down in the periodic table follow the octet rule when they are not the center atom The center atom is the atom in the molecule that the other elements individually bond attach to 0 When they are the center atom they can accommodate more than eight electrons 0 Using empty valence dorbitals that are predicted by quantum theory 0 Exceptions to the octet rule 0 H Li Be and B attain an electron con guration like that of He He can have ONLY two valence electrons a m Li loses its one valence electron H shares or gains one electron Though it commonly loses its one electron to become H Be loses two electrons to become Be2 Though it commonly shares its two electrons in covalent bonds resulting in four valence electrons B loses three electrons to become B3 Though it commonly shares its three electrons in covalent bonds resulting in six valence electrons Lewis Theory Predictions for Ionic Bonding Lewis theory predicts the number of electrons a metal atom should lose or a nonmetal atom should gain in order to attain a stable electron arrangement 0 The octet rule 0 Octet rule guideline allows us to predict the o formulas of ionic compounds that result 0 relative strengths of the resulting ionic bonds based on Coulomb s law Ionic Bonding Model versus Reality 0 Lewis theory 0 lmplies that the positions of the ions in the crystal lattice are critical to the stability of the structure 0 Predicts that moving ions out of position should therefore be dif cult and ionic solids should be hard Ionic solids are relatively hard compared to most molecular solids o lmplies that if the ions are displaced from their position in the crystal lattice repulsive forces should occur 0 Predicts that the crystal will become unstable and break apart Lewis theory predicts that ionic solids will be brittle lonic solids are brittle when struck they shatter o lmplies that in the ionic solid the ions are locked in position and cannot move around 0 Predicts that ionic solids should not conduct electricity To conduct electricity a material must have charged particles that are able to ow through the material Ionic solids do not conduct electricity 0 lmplies that in the liquid state or when dissolved in water the ions will have the ability to move around 0 Predicts that both a liquid ionic compound and an ionic compound dissolved in water should conduct electricity Ionic compounds conduct electricity in the liquid state or When dissolved in water Ionic Bonding and the Crystal Lattice The extra energy that is released comes from the formation of a structure in which every cation is surrounded by anions and vice versa 0 This structure is called a crystal lattice The crystal lattice is held together by the electrostatic attraction of the cations for all the surrounding anions o Electrostatic attraction is nondirectional force Therefore there is no ionic molecule 0 The chemical formula is an empirical formula simply giving the ratio of ions based on charge balance The crystal lattice maximizes the attractions between cations and anions leading to the most stable arrangement Energetics of Ionic Bond Formation Using NaCl as an Example The ionization energy of the metal is endothermic Nas gt Nag 1 e AH 496 kJmol The electron af nity of the nonmetal is exothermic 12C2g 1 e gt Cl g AH 244 kJmol Generally the ionization energy of the metal is larger than the electron af nity of the nonmetal therefore the formation of the ionic compound should be endothermic But the heat of formation of most ionic compounds is exothermic and generally large Why is this Nas 12C2g gt NaCls AH f 411 kJmol Crystal Lattice and Lattice Energy of NaCl Lattice Energy 0 The extra stability that accompanies the formation of the crystal lattice is measured as the lattice energy o It is the energy released when the solid crystal forms from separate ions in the gas state Always exothermic o Lattice energy is NOTmeasured directly but is calculated from knowledge of other processes o It depends directly on the size of charges and inversely on distance between ions Lattice Energy of an ionic Compound r Gaseous iliil lisl coalesce Heat is emitted Mia g erg gt maple iiiiquot iatticccncrg Conductivity of NaCl In NaCI5 the ions are stuck in position and not allowed to move to the charged rods In NaCIaq the ions are separated and allowed to move to the charged rods O NaCIs Naleaq Ionic Compounds IonC compounds are composed of cations metals and anions nonmetals bound together by ionic bonds 0 Examples of ionic compounds NaBr A2C033 CaHPO4 and MgSO4 The basic unit of an ionic compound is the formula unit the smallest electrically neutral collection of ions 0 Example The ionic compound table salt with the formula unit NaCl is composed of Na and Cl ions in a onetoone ratio 0 Summarizing IonC Compound Formulas o Ionic compounds always contain positive and negative ions 0 In a chemical formula the sum of the charges of the positive ions cations must equal the sum of the charges of the negative ions anions o The formula of an ionic compound re ects the smallest wholenumber ratio of Ions Naming Ionic Compounds Ionic compounds can be categorized into two types depending on the metal in the compound o The rst type contains a metal whose charge is invariant from one compound to another Whenever the metal in this rst type of compound forms an ion the ion always has the same charge lOiNIC COMPOUN Metals with Invariant Cation Charges Common Nonmetal Anions TABLE 52 Metals Whose Charge Is Invariant from One Compound to Another Ll LT Lithlu m l A Na Na Sodiu m l A K K39 Potassium 39lA Rb Rb Fiulaidi wm ilA Cs 33 Cesi mm l A Be Be Beryllium 2A Wig I Vlg2 Magnesium 2A Ca Ca2 Calcium 2A Sr Srz Strontium 2A Ba Ba2 Barium 2A Al AI3 Aluminum 3A Zn Zn2 Zinc Sc 503 Scandium Ag Agquot Silver The charge of these metals cannot be inferred from their group number HSilver semetimes forms compounds with other charges but these are rare TABLE 53 Some Common Monoatomio Anions Symbol Base N onmetal 7 for Ion Name Fluorine F fluor Fluonole Chlorine CI chlor Chloride Bromine Br brom Bromide Iodine iool Iodide Oxygen 02 ox Oxide Sulfur 82 sullf Sulfide Nitrogen N3 nitr Nitride Phosphorus P3 phosph Phosphide Naming Binary Ionic Compounds of Type I Cations Binary compounds contain only two different elements The names of binary ionic compounds take the following form base name of i anion nonmetal l ide o For example the name for KCI consists of the name of the cation potassium followed by the base name of the anion Chlor with the ending io e o KCI is potassium chloride The name for CaO consists of the name of the cation calcium followed by the base name of the anion 0X with the ending io e o CaO is calcium oxide Multivalent Metals Naming Type II Ionic Compounds 0 The metals in this category tend to have multiple charges ie multivalent cations 0 Their charge cannot be predicted as in the case of most representative elements and must be noted in their name Multivalent metals 0 Transition and inner transition metals 0 Iron Fe forms a 2 cation in some of its compounds and a 3 cation in others FeSO4 Here iron is a 2 cation Fe2 Fe2SO43 Here iron is a 3 cation Fe3 Many of the p block metals 0 Not all p block metals are multivalent 0 Some maingroup metals such as Pb TI and Sn form more than one type of cation Type II Cation TEBLE Some Metals That Form Cations with Different Charges Older Namequot Chromium Crquot Chromiumlll Chromous Crquot Chromiumllll Chromlc Iron Fe Ironllll Ferrous Fe3 Ironlllll Ferric Cobalt 302 Cobaltlll Cobaltous 003 Cobaltllll Cobaltlc Copper Cu Copperll Cuprous Cu Copperlll Cupric Tin Sn2 Tlnlll Stannous Sn4 TinllVl Stannic Mercury H922 Mercury Mercurous Hg2 Mercuryill Merourio Lead Fe2 Leadllll Plumbous Pb4 LeadilV Plumbic An older naming system substitutes the names found in this column for the name of the metal and its charge Under this system chromiumtlll oxide is named chromous oxide In this system the suffix ous indicates the ion with the lesser charge and ic indicates the ion with tlhe greater charge We will not use the older system in this text Naming Type II Binary Ionic Compounds For these types of metals the name of the cation is followed by a roman numeral in parentheses that indicates the charge of the metal in that particular compound For example we distinguish between Cu2 and Cu as follows Cu2 Copperll Cu Copperl The full names for compounds containing metals that form more than one kind of cation have the following form Cu20 Copperl oxide CuO Copperll oxide quot base name of anion nonmetal l 1 tag Naming Type II Binary Ionic Compounds Example CrBr3 To name CrBr3determine the charge on the chromium Total charge on cation total anion charge O Cr charge 3Br charge O Since each Br has a 1 charge then Cr charge 3 1 O Cr charge 3 O Cr 3 Hence the cation Crs is called chromiumlll and Br is called bromide The name for CrBr3 is chromiumlll bromide Naming Type II Binary Ionic Compounds Example ShClz To name SnCI2 determine the charge on the tin Total charge on cation total anion charge O Sn charge 2CI charge O Since each CI has a 1 charge then Sn charge 2 1 O Sn charge 2 O Sn 2 Hence the cation Sn2 is called tinll and CI is called chloride The name for SnC2 is tinll chloride Oxyanions Most polyatomic ions are oxyanions anions containing oxygen and another element Notice that when a series of oxyanions contains different numbers of oxygen atoms the oxyanions are named according to the number of oxygen atoms in the ion If there are two ions in the series then the one with more oxygen atoms has the ending ate and the one with fewer has the ending ite For example N03 is nitrate SO42 is sulfate N02 is nitrite SO32 is sul te If there are more than two ions in the series then the pre xes hypo meaning less than and per meaning more than are used CIO hypochlorite BrO hypobromite ClOz chlorite BrOz bromite C03 chlorate Br03 bromate CIO4 perchlorate BrO4 per o romate Naming Ionic Compounds Containing Polyatomic Ions Ionic compounds that contain a polyatomic ion rather than a simple anion eg CI are named in the same manner as binary ionic compounds except that the name of the polyatomic ion used For example NaN02 is named according to its cation Na sodium and its polyatomic anion NOZ nitrite Hence NaNOz is sodium nitrite Common Polyatomic Ions TABLE 55 Some Comtmon Polyatomic Ions Formula Formula Acetate C2H3 02 Hypochlorite CIO Carbonate 3032 Chlorite Clog Hydrogen carbonate HCO Chlorate Clog lor bicarbonate Hydroxide I OH Perchlorate CO4 Nitrite NOf Permanganate Mn04 Nitrate N0 Sulfite 8032 Chromate Cr042 Hydrogen sulfite or HSO3 bisulfite Dichromate 02072 Sulfate 5042 Phosphate P04 Hydrogen sulfate or HSO4 bisulfate Hydrogen phosphate I HPO42 Cyanide CN Dihydrogen H2P04 Peroxide of phosphate Ammonium NH4 Hydrated Ionic Compounds Hydrates are ionic compounds containing a speci c number of water molecules associated with each formula unit For example the formula for epsom salts is MgSO4 7H20 Its systematic name is magnesium sulfate heptahydrate Another exampleCoC2 6H20 is cobaltll chloride hexahydrate Hydrates Common hydrate pre xes o hemi 12 o mono 1 0 di 2 o tri 3 o tetra 4 o penta 5 o hexa 6 o hepta 7 o octa 8 Other common hydrated ionic compounds and their names are as follows 0 CaSO4 39 12H20 is called calcium sulfate hemihydrate o BaC2 6H20 is called barium chloride hexahydrate o CuSO4 6H20 is called copper sulfate hexahydrate Covalent Bonding Bonding and Lone Pair Electrons Electrons that are shared by atoms are called bonding pairs Electrons that are not shared by atoms but belong to a particular atom are called lone pairs 0 Also known as nonbonding pairs Bonding pair H H k Lone pair Duet Single Covalent Bonds 0 When two atoms share one pair of electrons the result is called a single covalent bond 0 Two electrons One atom may use more than one single bond to ful ll its octet 0 To different atoms 0 H only duet Double Covalent Bond When two atoms share two pairs of electrons the result is called a double covalent bond Four electrons between the two atoms Example 02 Octet Octet Elements that can doublebond with each other and themselves are C N O S and P Triple Covalent Bond When two atoms share three pairs of electrons the result is called a triple covalent bond Six electrons between the two atoms Example N2 N iN 2NENZ Elements that can triplebond with each other and themselves are C N O and S Covalent Bonding Model versus Reality Lewis Theory implies that some combinations should be stable Whereas others should not Stable combinations result in octets allows us to predict the formulas of molecules of covalenty bonded substances Hydrogen and the halogens are all diatomic molecular elements as predicted by Lewis theory Oxygen generally forms either two single bonds or a double bond in its molecular compounds There are some stable compounds in which oxygen has one single bond and another in which it has a triple bond but it still has an octet Lewis Theory of Covalent Bonding implies that the attractions between atoms are directional The shared electrons are most stable between the bonding atoms predicts that covalenty bonded compounds Will be found as individual molecules Rather than an array like ionic compounds Compounds of nonmetals are made of individual molecule units Molecular Compounds Formulas and Names The formula for a molecular compound cannot readily be determined from its constituent elements because the same combination of elements may form many different molecular compounds each with a different formula Nitrogen and oxygen form all of the following unique molecular compounds 39 NO N02 N20 N203 N204 and N205 Molecular Compounds Molecular compounds are composed of two or more nonmetals Names of Molecular Compounds Write the name of the element with the smallest group number rst If the two elements lie in the same group then write the element with the greatest row number rst The pre xes given to each element indicate the number of atoms present Binary Molecular Compounds name of base name of 1st a 2nd element element ide These pre xes are the same as those used in naming hydrates mono 1 hexa 6 di 2 hepta 7 tri 3 octa 8 tetra 4 nona 9 penta 5 deca 10 If there is only one atom of the rst elementin the formula the pre x mono is normally omitted Formula MassMolecular Mass of a Compound Molecular Mass The mass of an individual molecule or formula unit is known as molecular mass or molecular weight of the compound It is the mass of ONE MOLE of that compound Determining a Compound s Molecular Mass Sum of the masses of the atoms in a single molecule or formula unit Example What is the molecular mass of water H20 2101 amu H 1600 amu O 1802 amu One mole of water has a molecular mass of 1802 grams Problem Solving Calculating Formula Mass Number of atoms Atomic mass Number of atoms Atomic mass Formula mass 1 J of Ist element in X of of 2nd element in X of chemical formula list element chemical formula 2nd element Using Molar Mass to Count Molecules by Weighing Molar mass in combination with Avogadro s number can be used to determine the number of atoms in a given mass of the element Use molar mass to convert to the amount in moles Then use Avogadro s number to convert to number of molecules gC02 11101 C02 I C02 molecules 1 11101C02 60g X 1023 CO2 molecules 4401 g C02 1 11101C02 Composition of Compounds 0 A chemical formula in combination with the molar masses of its constituent elements indicates the relative quantities of each element in a compound 0 The percentage of each element in a compound can be determined from o 1 the formula of the compound and o 2 the experimental mass analysis of the compound molecular mass of element Z 0 mass of elementZ x 100 0 mass of 1 mol of compound 0 The percentages may not always total to 100 due to rounding Conversion Factors from Chemical Formula Chemical formulas contain within them inherent relationships between numbers of atoms and molecules Or moles of atoms and molecules 1 11101 CC12F2 2 11101 C1 These relationships can be used to determine the amounts of constituent elements and molecules Such as percent composition Determining a Chemical Formula from Experimental Data 0 Empirical Formula 0 Simplest wholenumber ratio of the atoms of elements in a compound 0 Can be determined from elemental analysis Masses of elements formed when a compound is decomposed or that react together to form a compound Combustion analysis Percent composition 0 Note An empirical formula represents a ratio of atoms or a ratio of moles of atoms not a ratio of masses Finding an Empirical Formula 1 Convert the percentages to grams a If not given assume you start with 100 g of the compound b Example 245 C means 245 g C 2 Convert mass grams to moles a Use molar mass of each element b Example 245 g C x 1 mol C1201 grams 200 mol C 3 Divide all by the smallest number of moles to obtain the atomtoatom ratio for each of the elements in the compound a If the result is within 01 of a whole number round to the whole number 4 Multiply all mole ratios by a number to make all whole numbers a If ratio is 5 multiply all by 2 if the ratio is 33 or 67 multiply all by 3 and so on b Skip if already whole numbers From Empirical to Molecular Formulas for Compounds The molecular formula is a multiple of the empirical formula It is the actual formula of the compound Knowing the molecular formula you can determine the molecular mass of the compound To determine the molecular formula you need to know the empirical formula and the molar mass of the compound Molecular formula empirical formula n where n is a positive integer Molecular Formulas for Compounds The molar mass is a wholenumber multiple of the empirical formula molar mass the sum of the masses of all the atoms in the empirical formula l7 molar mass empirical formula molar mass Combustion Analysis A common technique for analyzing compounds is to burn a known mass of compound and weigh the amounts of product made 0 This is generally used for organic compounds containing C H or O By knowing the mass of the product and composition of constituent element in the product the original amount of constituent element can be determined 0 All the original C forms C02 the original H forms H20 and the original mass of O is found by subtraction Once the masses of all the constituent elements in the original compound have been determined the empirical formula can be found Organic Compounds Early chemists divided compounds into two types organic and inorganic Compounds from living things were called organic compounds from the nonliving environment were called inorganic Organic compounds are easily decomposed and could not be made in the lab Inorganic compounds are very dif cult to decompose but are able to be synthesized Modern Organic Compounds Today organic compounds are commonly made in the lab and we nd them all around us Organic compounds are mainly made of C and H sometimes with O N P S and trace amounts of other elements The main element that is the focus of organic chemistry is carbon Structural formula Spacefilling model i H lj H H 7 o Methane CH4 Carbon Bonding Carbon atoms bond almost exclusively covalently 0 Compounds with ionic bonding C are generally inorganic When C bonds it forms four covalent bonds 0 four single bonds two double bonds one triple bond and one single bond etc Carbon is unique in that it can form limitless chains of C atoms both straight and branched and rings of C atoms Carbon Bonding of Organic Molecules H H H H H H l l l l chgCECH chgcgch l l H H H H H H H H Propane C3H3 Isolbllt me C4Hlo H H H H Ethene C2H4 Ethyne 32le Common Hydrocarbons Cyclohexalne C lng O Acetic acid CH 3COOH ff 56 Colman Hydrocarbons Molestiller Structural Spaoea lling Formula Formula Mode Eamon Uses Methane CH4 Primary component of natural gas Propane GgHg LP gas for grills and outdoor stoves nB utane C4Hm Common fuel for lighters nPentane Ci5H12 Component of gasoline H H a f Ethene 02H4 C C Ripening agent in fruit H H Ethvne 02H H C E C 1H Fuel for welding torches k The quotnquot in the names of these hydrocarbons stands for normal which means straight chain
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