ElectronSpinEach orbital can hold no more than two electrons. The two electrons in a particular orbital differ inone way, namely, they have different spins. Electrons can "spin" in one of two direction, one pointingupward and one pointingdownward.Forthe Is orbital containing2electrons, it canbe illustrated intwo ways,i.e.IsNor1s?Howtoillustrate the 2porbitals that contain6electrons?3.12WritingElectronicConfigurationsforAtonsThe electronic configurations for an atom is written by listing the orbitals occupied by electrons inthe atom along with the number of electrons in each orbitals. Three Rules which must be followed inwritingelectronic configurationsarePauli principle,Aufbau principle, and Hund's rulePauliPrinciple:Eachorbitalmaycontaintwoelectrons.Itispossibleforanorbital tocontainnoelectronsor justoneelectron,butnomorethantwoelectrons.Aufbau Principle: Orbitals are filled by starting with the lowest-energy orbitals first For example,Is orbitalsarefilledbefore2s orbitalswhichinturnarefilledbefore2porbitalsHund's Rule : When orbitals of equal energy, such as the three p orbitals, are beingfilled, electronstend to have the same spin. The electrons occupy different orbitals so as to remain as far apart as possible.This is reasonable, since electrons have like charges and tend to repel each other. The electrons do notpair up until there is at least one electron in eachofthe equal-energy orbitals·e.g.,2p,ppyL pExamplesofElectronicConfigurationsofAtoms.H)Is! or Is .·He)1s? or ls.B)1s2252plor2p12py2p:.25111s ↑±2p_2p,2p::C)1s?2s?2p2plor2s↓1s T±10
10 Electron Spin Each orbital can hold no more than two electrons. The two electrons in a particular orbital differ in one way, namely, they have different spins. Electrons can “spin” in one of two direction, one pointing upward and one pointing downward. For the 1s orbital containing 2 electrons, it can be illustrated in two ways, i.e., 1s or 1s 2 How to illustrate the 2p orbitals that contain 6 electrons? 3.12 Writing Electronic Configurations for Atoms The electronic configurations for an atom is written by listing the orbitals occupied by electrons in the atom along with the number of electrons in each orbitals. Three Rules which must be followed in writing electronic configurations are Pauli principle, Aufbau principle, and Hund’s rule. Pauli Principle: Each orbital may contain two electrons. It is possible for an orbital to contain no electrons or just one electron, but no more than two electrons. Aufbau Principle: Orbitals are filled by starting with the lowest-energy orbitals first. For example, 1s orbitals are filled before 2s orbitals which in turn are filled before 2p orbitals. Hund’s Rule : When orbitals of equal energy, such as the three p orbitals, are being filled, electrons tend to have the same spin. The electrons occupy different orbitals so as to remain as far apart as possible. This is reasonable, since electrons have like charges and tend to repel each other. The electrons do not pair up until there is at least one electron in each of the equal-energy orbitals. • e.g., 2p4 , px py pz Examples of Electronic Configurations of Atoms • H) 1s 1 or 1s . • He) 1s 2 or 1s • B) 1s 2 2s 2 2px 1 or 2px 2py 2pz 2s 1s • C) 1s 2 2s 2 2px 12py 1 or 2px 2py 2pz 2s 1s
Chapter 4 ChemicalBonding4.1IntroductionChemical bondsaretheattractiveforces which join atoms.Thedistancebetweenthecenters oftwoatoms joined by a chemical bond is between 70 pm and 300pm. The energy needed to break a chemicalbondbetweentwoatoms is calledthebondenergyChemical compounds areconveniently divided into two broad classes,called ionic compounds andcovalentcompounds.4.2Types ofCompoundsCompounds can be classified as ionic or covalent by examining two physical properties, meltingpoint and the ability to conductelectricity.lonic Compounds have very high melting point and are goodconductors of electricity when they are either melted or dissolved in water. Covalent compounds havemuchlowermeltingpointandarepoorconductorsofelectricity.4.3Formation of IonsIons are electrically charged species fomed when a neutralatom eithergains or loses one ore moreelectrons.Cations,orpositive ions,fom when atoms lose oneormore electrons.Anions,or negativeions,formwhenatomsgainoneoremoreelectrons.An ionic compound is an electrically neutral compound which consists of cations and anions heldtogether byforces ofelectric attraction.Stable Noble Gas ConfigurationsThe atoms of representative elements tend to lose or gain electrons so that their electronicconfigurations become identical to those ofthe noble gas nearest to them in the periodic table.Cation Formation:The metallic element of group IA have the general electronic configuration nsTo obtain a stable noble gas configuration they lose this highest-energy electron. e.g.→ Lit) 1s?Li) 1s°2s'-(He) 1s?SimilarlyforthegrouplIA elements with thegeneral electronic configuration ns,we have,forexample,-2eMg2+)1s2s?2pMg)1s222p*3s2Anion Formation: The nonmetallic element of groups VIA and VIIA gain electrons to form negativeionswith stable,noblegaselectronicconfigurations.+e>F)1s22s°2p° (Ne)F) 1s°2s2pe.g.2-0-) 15*252p (Ne)O) 1s*2s2p4.4Polyatomic ionsIt is possible for ions to include two or more atoms. Such polyatomic ions behave as if they weremonatomicionsandinfactareoftencomponentsof ioniccompounds.Themostfrequentlyencounteredpolyatomic cation is the ammonium ion,NH,.Several anions havenames thatend in-ide,including thesethree:OH-(hydroxide),CN-(cyanide)andO,2-(peroxide).11
11 Chapter 4 Chemical Bonding 4.1 Introduction Chemical bonds are the attractive forces which join atoms. The distance between the centers of two atoms joined by a chemical bond is between 70 pm and 300pm. The energy needed to break a chemical bond between two atoms is called the bond energy. Chemical compounds are conveniently divided into two broad classes, called ionic compounds and covalent compounds. 4.2 Types of Compounds Compounds can be classified as ionic or covalent by examining two physical properties, melting point and the ability to conduct electricity. Ionic Compounds have very high melting point and are good conductors of electricity when they are either melted or dissolved in water. Covalent compounds have much lower melting point and are poor conductors of electricity. 4.3 Formation of Ions Ions are electrically charged species formed when a neutral atom either gains or loses one ore more electrons. Cations, or positive ions, form when atoms lose one or more electrons. Anions, or negative ions, form when atoms gain one ore more electrons. An ionic compound is an electrically neutral compound which consists of cations and anions held together by forces of electric attraction. Stable Noble Gas Configurations The atoms of representative elements tend to lose or gain electrons so that their electronic configurations become identical to those of the noble gas nearest to them in the periodic table. Cation Formation: The metallic element of group IA have the general electronic configuration ns1 . To obtain a stable noble gas configuration they lose this highest-energy electron. e.g. Li) 1s2 2s1 Li+ ) 1s2 (He) 1s2 Similarly for the group IIA elements with the general electronic configuration ns2 , we have, for example, Mg) 1s2 2s2 2p6 3s2 Mg2+ ) 1s2 2s2 2p6 Anion Formation: The nonmetallic element of groups VIA and VIIA gain electrons to form negative ions with stable, noble gas electronic configurations. e.g. F) 1s2 2s2 2p5 F - ) 1s2 2s2 2p6 (Ne) O) 1s2 2s2 2p4 2e O2- ) 1s2 2s2 2p6 (Ne) 4.4 Polyatomic ions It is possible for ions to include two or more atoms. Such polyatomic ions behave as if they were monatomic ions and in fact are often components of ionic compounds. The most frequently encountered polyatomic cation is the ammonium ion, NH4 + . Several anions have names that end in -ide, including these three: OH- (hydroxide), CN- (cyanide) and O2 2- (peroxide)
CommonanionsandtheirnamesFormulaNameFormulaNameWith-1 chargeWith-2 chargeS2-F-FluoridesulfideCrCO;2-ChloridecarbonateBrSO,2-bromidesulfite1SO,2-IodidesulfateCrO,2-NO2NitritechromateCr20,2-NO3NitratedichromateHCO3SiOs2-bicarbonatesilicateCIO-HypochloriteCIO2ChloriteWith -3 chargePO43-CIO4perchloratephosphateMnO4permanganate4.5lonic CompoundsThe name of an ionic compound is the name ofthe cation followed bythe nameofthe anionSum of charges on cations - Sum of charges on anionsSodium chloride:NaCI = Nat+ CIMagnesium Chloride:MgClz= Mg2++2CIBarium phosphate:Ba:(PO4)24.6CovalentBondingTheoriesCovalent compounds are sometimes called molecular compounds. A covalent bond between twoatoms is formed by the sharing of one or more pairs of electrons.This is unlike an ionic bond, formationof which involves a transfer of electrons.Using themodern orbital pictureof theatom, one can explainhowacovalentbondforms.Covalent bonding in H2A H atom has a ls orbital containing one electron. When two H atoms get closer and closer, their Isorbital begin to overlap. The two ls orbitals merge to form a molecular orbital of increased electrondensity. The two electrons in the molecular orbital are shared by two H atoms.Types of covalent bondsSigma (o) MOs fom from the overlap of s with s and p with s and from the head-to-head overlap oftwop orbitals.Thepi(元)MOsformfromtheside-to-sideoverlapoftwoporbitals00s+pp+ps+sp+p元bondbond(s)12
12 Common anions and their names Formula Name Formula Name With –1 charge With –2 charge F - Fluoride S 2- sulfide Cl- Chloride CO3 2- carbonate Br- bromide SO3 2- sulfite I - Iodide SO4 2- sulfate NO2 - Nitrite CrO4 2- chromate NO3 - Nitrate Cr2O7 2- dichromate HCO3 - bicarbonate SiO3 2- silicate ClO- Hypochlorite ClO2 - Chlorite With –3 charge ClO4 - perchlorate PO4 3- phosphate MnO4 - permanganate 4.5 Ionic Compounds The name of an ionic compound is the name of the cation followed by the name of the anion. Sum of charges on cations = Sum of charges on anions Sodium chloride: NaCl = Na + + Cl− Magnesium Chloride: MgCl2 = Mg2+ + 2Cl− Barium phosphate: Ba3(PO4)2 4.6 Covalent Bonding Theories Covalent compounds are sometimes called molecular compounds. A covalent bond between two atoms is formed by the sharing of one or more pairs of electrons. This is unlike an ionic bond, formation of which involves a transfer of electrons. Using the modern orbital picture of the atom, one can explain how a covalent bond forms. Covalent bonding in H2 A H atom has a 1s orbital containing one electron. When two H atoms get closer and closer, their 1s orbital begin to overlap. The two 1s orbitals merge to form a molecular orbital of increased electron density. The two electrons in the molecular orbital are shared by two H atoms. Types of covalent bonds Sigma () MOs form from the overlap of s with s and p with s and from the head-to-head overlap of two p orbitals. The pi() MOs form from the side-to-side overlap of two p orbitals. s+s s+p p+p p+p bond(s) bond
4.7LewisElectronDotStructuresLike molecular orbital theory, the electron dot theory, proposed by the American chemist G.N. Lewis.describes a covalent bond as a shared pair of electrons. The Lewis theory predicts the likelihood offormation of covalentmolecules by establishing a criterion for their stabilityThecriterionisthat anelectronic configuration ofeach atombe the same as thatof one of the noblegases.Octet Rule: Each atom in the bond must be surrounded by eight electrons or, if the atom is H, by twoelectrons.This so-called octet rule is followed by most covalent compounds.The electrons included inthe Lewis structures are those which areinthehighest-energylevel of each atom;these aretheelectronsavailableforbondingand are called valenceelectronsExamplesoftheLewisStructuresH, molecule: H:HThe electron pair which joins the two atoms is single covalentbond.HH:C:HH:0:H:F.+F:-:F:F:iiH,O:CH4:F,:..:Ci:P:CI:H:N:H:ci:ii.PCl3:NH,:?Theansweris4.8MultipleCovalentBondSometimes more than one electron pair must be placed between two atoms to satisfy the octet ruleBonds that include more than one electron pair are called multiple covalent bonds. In double bonds, thereare two electron pairs and in triple bonds there are threeH:C.:C:HH:C:::C:Hiie.g.,CHa:C,H2:4.9Exceptions ofLewis TheorySome compoundsdoexist even though Lewis Structures whichfollowtheoctet rule cannot bedrawnforthem.Theonly way to drawLewis structuresfor thesemolecules is to violatethe ruleofeightaroundtheircentralatom.ci.a::F:B:F::c.pa::F::CI:BF3:.E.g.,PCls:4.10ElectronegativityandPolarBondsWhen an electron pair (or pairs)involved in a covalent bond is shared by two identical atoms, thesharing is equal. When an electron pair is shared by two different atoms, one atom may have a greaterattractionfortheelectronpairthantheotheratom.Theatomwiththegreaterattractionfortheelectron8+...8H:C:pair with assumea partial negative charge relative to the other atom.E.g,HCl:13
13 4.7 Lewis Electron Dot Structures Like molecular orbital theory, the electron dot theory, proposed by the American chemist G.N. Lewis, describes a covalent bond as a shared pair of electrons. The Lewis theory predicts the likelihood of formation of covalent molecules by establishing a criterion for their stability. The criterion is that an electronic configuration of each atom be the same as that of one of the noble gases. Octet Rule: Each atom in the bond must be surrounded by eight electrons or, if the atom is H, by two electrons. This so-called octet rule is followed by most covalent compounds. The electrons included in the Lewis structures are those which are in the highest-energy level of each atom; these are the electrons available for bonding and are called valence electrons. Examples of the Lewis Structures H2 molecule: H:H The electron pair which joins the two atoms is single covalent bond. F2: F + F F F H2O: H O H CH4: H C H H H • PCl3: Cl P Cl Cl NH3: ? The answer is H N H H 4.8 Multiple Covalent Bond Sometimes more than one electron pair must be placed between two atoms to satisfy the octet rule. Bonds that include more than one electron pair are called multiple covalent bonds. In double bonds, there are two electron pairs and in triple bonds there are three. e.g., C2H4: H H C C H H C2H2: H C C H 4.9 Exceptions of Lewis Theory Some compounds do exist even though Lewis Structures which follow the octet rule cannot be drawn for them. The only way to draw Lewis structures for these molecules is to violate the rule of eight around their central atom. • E.g., PCl5: P Cl Cl Cl Cl Cl BF3: B F F F 4.10 Electronegativity and Polar Bonds When an electron pair (or pairs) involved in a covalent bond is shared by two identical atoms, the sharing is equal. When an electron pair is shared by two different atoms, one atom may have a greater attraction for the electron pair than the other atom. The atom with the greater attraction for the electron pair with assume a partial negative charge relative to the other atom. E.g., HCl: H Cl + −
Bonds suchas the one in HClin which the sharing between atoms isnotequal are polarcovalentbonds. An extreme case ofthe polar covalent bond is the ionic bond, in whichelectron transfer hasoccured, producing ions with full charge.The other extreme case is the nonpolar covalent bond (as in H2,F2, and N2).ElectronegativityThe degree of attraction an atom asfor abonding electron pair is the electronegativityof theatomLinus Pauling, whose contributions to chemical bonding theory eamed him a Nobel Prize in 1954,assigned numbers to represent the electronegativity of atoms; the higher the number , the greater theelectronegativity. The atom with the highest electronegativity, 4.o, is Fluorine. The greater theelectronegativity difference between two atoms, the more polar the bond that foms between them. Whentheelectronegativitydifferenceisgreaterthan.7,thebondbetweentheatomsisconsideredtobeionic.Predictionofpolarity ofbondsH-HC-SNa--CICa--0ElectronegativityCa (1.0) 0(3.5)H(2.1)C(2.5) S(2.5)Na(0.9) CI(3.0)2.10.02.40.0DifferencePolarityionicionicnonpolarnopolar4.11PolarityofMoleculesSome important properties of compounds depend on whether or not their molecules are polar. To findout if a molecule is polar we check to see if it contains any polar bonds and then find out howthe polarbonds are arranged in themolecule. In very symmetrical molecules polar bonds may cancel one anotherso that the molecule as a whole is nonpolar.Nonpolar MoleculesMolecules that contain only nonpolar bonds must benonpolar.Some nonpolar moleculesdo containpolarbonds,but theyare so symmetrical thatthepolarities cancel,e.g.CF,and CO,4.12Naming binary covalentcompoundsCovalent compounds which contain two nonmetals are called binary covalent compounds.Theirnamesconformtoaspecial systemsimilartothatfornamingioniccompounds.Thenameoftheelementwrittenontheleftoftheformula(usuallytheleastelectronegativeelement)issimplythenameoftheelement itself.Thenameoftheotherelementwrittenontheright(usuallythemostelectronegativeelement)ismodifiedwiththesuffix-ideNames of some covalent binary compoundsFormulaFormulaTrivialPropernameTrivial namePropernamenameCONH3Carbon monoxideNitrogen trihydrideammoniaCO2H20Carbon dioxideDihydrogen oxidewaterNOSO3Nitrogen oxideNitric oxideSulfur trioxideNO2Nitrogen dioxideCH4Carbon tetrahydridemethaneN20Nitrous oxideSO2Dinitrogen oxideSulfur dioxide14
14 Bonds such as the one in HCl in which the sharing between atoms is not equal are polar covalent bonds. An extreme case of the polar covalent bond is the ionic bond, in which electron transfer has occurred, producing ions with full charge. The other extreme case is the nonpolar covalent bond (as in H2, F2, and N2). Electronegativity The degree of attraction an atom as for a bonding electron pair is the electronegativity of the atom. Linus Pauling, whose contributions to chemical bonding theory earned him a Nobel Prize in 1954, assigned numbers to represent the electronegativity of atoms; the higher the number , the greater the electronegativity. The atom with the highest electronegativity, 4.0, is Fluorine. The greater the electronegativity difference between two atoms, the more polar the bond that forms between them. When the electronegativity difference is greater than 1.7, the bond between the atoms is considered to be ionic. Prediction of polarity of bonds Na-Cl Ca-O H-H C-S Electronegativity Na(0.9) Cl(3.0) Ca (1.0) O(3.5) H(2.1) C(2.5) S(2.5) Difference 2.1 2.4 0.0 0.0 Polarity ionic ionic nonpolar nopolar 4.11 Polarity of Molecules Some important properties of compounds depend on whether or not their molecules are polar. To find out if a molecule is polar we check to see if it contains any polar bonds and then find out how the polar bonds are arranged in the molecule. In very symmetrical molecules polar bonds may cancel one another so that the molecule as a whole is nonpolar. Nonpolar Molecules Molecules that contain only nonpolar bonds must be nonpolar. Some nonpolar molecules do contain polar bonds, but they are so symmetrical that the polarities cancel, e.g. CF4 and CO2. 4.12 Naming binary covalent compounds Covalent compounds which contain two nonmetals are called binary covalent compounds. Their names conform to a special system similar to that for naming ionic compounds. The name of the element written on the left of the formula (usually the least electronegative element) is simply the name of the element itself. The name of the other element written on the right (usually the most electronegative element) is modified with the suffix -ide. Names of some covalent binary compounds Formula Proper name Trivial name Formula Proper name Trivial name CO Carbon monoxide NH3 Nitrogen trihydride ammonia CO2 Carbon dioxide H2O Dihydrogen oxide water NO Nitrogen oxide Nitric oxide SO3 Sulfur trioxide NO2 Nitrogen dioxide CH4 Carbon tetrahydride methane N2O Dinitrogen oxide Nitrous oxide SO2 Sulfur dioxide