Unit 9: CHEMICAL BONDING
Chemistry 11 Bonding Notes KeyUnit 9: CHEMICALBONDINGUnit 9: Bonding:1. Electronegativity2. Intramolecular Bonding3. Intermolecular Bonding4. Drawing Lewis Structures5. Lewis Structures for Polyatomic Ions6. Exceptions to the Octet Rule7. Resonance Structures8. Molecular Shapes9. Molecular PolarityValence Electrons1. Electronegativity The outer electrons involved in bonding Number of valence electrons will usually be between1 and 7The representative metals are on the left side of thestaircase, and most of them have 1, 2, or 3 valenceelectrons.For metal atoms with 1, 2, or 3 valence electrons, it iseasier to give them away in order to becomeisoelectronic with a noble gas, resulting in a charge(cation).In general, non‐metals have 4‐7 valence electrons, so itis easier in most cases to gain electrons (by transfer orsharing) to become isoelectronic with a noble gas.Electronegativity The ability of an atom to attract and/or gainelectrons elements on the left of the Periodic Table like to loseelectrons to become isoelectronic with a Noble Gas,thus their ability to attract electrons is poor elements on the right of the Periodic Table like to gainelectrons to become isoelectronic with a Noble Gas,thus their ability to attract electrons is quite strong1
Chemistry 11 Bonding Notes Key The most electronegative elements are found onthe right (non‐metals), but strongest on the topright (fluorine) Electropositive elements are very poor at gainingelectrons (because they like to give up electronsinstead (Groups I, II, III) These elements are found on the left (metals), butstrongest on the bottom left (Francium).2. INTRAMOLECULAR BONDING Atoms bond in order to achieve a lowerenergy state (try to get full orbitals) We will focus on the representative elements(s & p blocks) (bonding for transition metals isnot covered in Chem 11)Two Categories of BondingIntramolecular Bonding refers to attraction ions or atoms have forone another WITHIN a moleculeIntermolecular Bonding refers to the attraction BETWEEN moleculesNote: Intramolecular bonds are, in general,much stronger than intermolecular bondsThe Octet Rule Every representative element strives to achieve ans2p6 arrangement (full orbital) in order to beisoelectronic with a noble gas 2 6 8 which gives us the “octet” rule 8 valence electrons 0 valence electrons Elements achieve an octet by chemical reactions andform one of 3 types of bonds: an ionic bond a polar covalent bond a non‐polar covalent bondA tug of war between atomsA complete transferof electronsIonic bondPolarcovalentδ bondNon‐polarcovalentbondA) Ionic Bondse‐ e‐δ Unequalsharing ofelectronsEqual sharing ofelectronsThe type of bond formed depends on theelectronegativity difference between the two atoms Involves a metal and a non‐metal Electronegativity differenceE.g. What happens when Na metal comes intocontact with a Cl atom?Na will give up an electron to ClElectronegativity calculation:3.0 0.9 2.1 (greater than 1.7)Cl NaTherefore, an IONIC bond(a complete transfer of an electron)Na ClNa Cl Now both Na and Cl have an octet!2
Chemistry 11 Bonding Notes KeyHow does it happen? Since Na has a very low Ionization Energy (wants togive up an electron) and Cl has a very highelectronegativity (wants to gain an electron) acomplete transfer of electrons will take place Na and Cl‐ ions that result are attracted and stay besideeach other due to electrostatic forces ( ‐ attraction) Many metals and non‐metals have anelectronegativity difference greater than 1.7, resultingin ionic bonds (metal cations and non‐metal anions)How does Ca bond with Br2?2.8 1.0 1.8 so ionic bondCa has 2 valence electrons to give up to belike a noble gas.Each Br has 7 valence electrons and needsone more to be like a noble gas.Therefore, one Ca will give up two electrons tobecome Ca2 and this will cause two Br atomsto each gain one electron to become two Br ionsCa Br22 Ca Ionic solids are brittle andhave very high melting points(difficult for ions to move dueto strong electrostaticattraction) Each ionic solid (salt) has aunique crystal latticearrangementB) Non‐Polar Covalent Bonds equal sharing of electrons to obtain octet Form between non‐metal atoms, usually of thesame element (diatomic molecules) eg.Cl2, H2, etc. Electronegativity difference of 0.0 ‐ 0.2 Each covalent bond is made up of two electrons,one from each atomHHH H 2BrCl2: Each Cl atom has 7 valence electrons.If they can each gain 1 electron, they canhave an octet like a noble gas (like Neon).Thus, they each share one of theirelectrons.The two shared electronsare shared equallybetween the chlorinesbecause they have equalelectronegativities. If this happens many timesin the same vicinity, ions forma crystal lattice with positionsof ions fixed due toelectrostatic attractionH HHydrogen has 1 valence electron andcan either give it up to have 0, or itcan gain another to becomeisoelectronic with He (1s2). So, if eachH shares its electron with the other,they both have two electrons (likeHelium). This is why H comesnaturally as the diatomic molecule H2 Electrons are equally shared between atomsbecause electronegativity (pulling power) isidentical (or almost!) Can have a single, double or triple bonddepending on each atom’s needs to fill its octet(we'll learn more about this later in the ons/chang 7e esp/bom1s2 11.swf3
Chemistry 11 Bonding Notes KeyC) Polar Covalent Bonds Usually form between non‐metal atoms, butsometimes between a metal and a non‐metal ifthe electronegativity difference is between 0.2and 1.7 Each atoms offers up an electron for sharing,but because one atom is significantly moreelectronegative that the other, the electrons areshared unequally (they are closer to the moreelectronegative atom) The atom that is pulling the shared electronscloser results in a 'partial' negative end and theother end is 'partially' positive The partial negative end of a polar bond isalways at the atom with the highestelectronegativity (and vice versa) Example: H and ClHElectronegativity Difference:3.0 2.1 0.9Thus, polar covalent bondClH ClUNEQUAL SHARINGH ClPolar MoleculePartial charge at each end called a dipole (δ)electron poorregionδ electron richregionClFHe poorδ e richH Clδ δ 0.21.7(NON‐POLAR)H ClNa Cl 4
Chemistry 11 Bonding Notes KeyExamples What kind of intramolecular bond formsin each of the following? CS2 H2 O Al2O3 AsH3HOMEWORK: Part 1 of the Intramolecular andIntermolecular Bonding WorksheetRecall:Intramolecular bonding within a molecule holds atoms together asmolecules holds ions togetherIntermolecular bonding between molecules holds molecules to eachother3. Intermolecular BondingWater striders use the strongintermolecular forces of water(creating surface tension) to stayon the surface.In order to change a liquid to a gas, energy isrequired to overcome the intermolecularattractions. So to make tea or coffee, youneed to decrease the intermolecularattractions of water.DNA uses intermolecular forces to stay twisted in asmall space. If it completely unwound, a humanDNA would be over 3m long.5
Chemistry 11 Bonding Notes KeyIntermolecular Attraction Much weaker than intramolecular bonds(ionic and covalent bonds) Intermolecular bonds become weaker asdistance between molecules increases Three types: London Dispersion Forces ‐ weakest type ofintermolecular bonding (also called Van der Waalsforces) Dipole‐to‐Dipole Forces in polar molecules Hydrogen Bond is a special type of very strongdipole‐to‐dipole forces (Water has very strong H‐bonds)Instantaneous unevendistribution of electronsin He atom(dispersion force)Creates atemporarydipoleCauses aneighbouring Heatom to create atemporary dipoleWhat is the result of dispersion forces? As atoms or molecules get larger, they have moreelectrons in their electron clouds, therefore can createstronger dispersion forces, creating a strongerattraction between the moleculesNonpolar MoleculesLondonKinetic EnergyForceshas overcomeexertingLondon ForcesLondon Dispersion Forces Temporary dipoles (partial positive or negativecharges) that result from the random distribution &movement of the electrons around the particle A temporary dipole in one molecule will induce atemporary dipole in a neighbouring one. The twodipoles then attract each other Dispersion forces are temporary and work over veryshort distances All molecules have them, but they areovershadowed by other, stronger intermolecularforcesin polar molecules or ionic bonding in ioniccompounds Important for noble gases and non‐polar molecules,as it is their only form of intermolecular bonding Size of the London Force depends on the number ofelectrons in the atoms or molecules Large atoms and molecules have a larger electroncloud. Because of this they distort and polarize moreeasily, thus stronger dispersion forcesLondon Forces and Boiling Points Boiling points (C) of halogens (non‐polar molecules)Cl 2Br2I2 many electronsfew electrons F2‐188‐3559184 At room temperature, F 2 and Cl2 are gases, Br2 is aliquid, and I 2 is a solid Boiling points (C) of Noble gasesNeArKrfew electrons He‐269 ‐246 ‐186‐152Xe‐108many e ‐attractionThe weaker the dispersion forces, the easier the particles can8become separated (the lower the temperature at which they willchange state).6
Chemistry 11 Bonding Notes KeyDipole‐to‐Dipole AttractionsIn polar covalent molecules, differences inelectronegativity lead to unequal sharing ofelectrons which results in permanent dipoles (partialcharges). Polar covalent molecules also have LondonDispersion Forces, but they are overshadowed byDipole to Dipole Attractions the partial negative end of one molecule isattracted to partial positive end of another dipole‐dipole forces are stronger than dispersionforces, and thus polar molecules have, in general,higher melting points & boiling points than non‐polar molecules and atomsHCl moleculesdipole to dipoleintermolecular forcea) The interaction of twopolar moleculesb) The interaction of manydipoles in a liquidIn general, the larger the dipole dipoleinteraction, the higher the boiling point.Hydrogen Bonding a special type of dipole‐dipole intermolecular bond Strongest intermolecular bond Occurs in molecules that have N‐H, O‐H or H‐Fpolar covalent intramolecular bonds (Hydrogenbonds are FON to learn!) permanent dipoles created are especially powerfulas the hydrogen nucleus is essentially naked (due tothe more electronegative F, O, or N pullinghydrogen's only electron away ‐ unequal sharing) the naked hydrogen (partial positive charge)creates an intermolecular hydrogen bond with aneighbouring partial negative charge(a) The polar water molecule(b) Hydrogen bonding among watermoleculeshydrogen bond7
Chemistry 11 Bonding Notes KeyMolecules with hydrogen bonding haveunusually highmelting and boiling points due to the hydrogen bondsholding particles close togetherhydrogen bondingdipole dipolebondingWhat if the intramolecular bonds are ionic?Then a crystal lattice is formed, and all thebonds are ionic (intra and inter)as you cannot tell them apart!However, we class ionic bondsonly as intramolecular. any substance that has hydrogen bonding hasan unusually high solubility in water since water has hydrogen bonds, it likes tointeract with other substances that havehydrogen bonds water has strong surface tension due tohydrogen bonds (water striders)Deciding Which Intermolecular Bonds Exist inCertain Compounds First, using the electronegativity table, decide whichintramolecular bond exists within your molecule. If the intramolecular bond is:a. Ionic, then the intermolecular bonds are also ionic(crystal lattice), however ionic bonds are all consideredintramolecular bondsb. Non‐polar covalent: then the intermolecular bonds areLondon Dispersion Forces.c. Polar covalent: then the intermolecular bonds areDipole‐Dipole Forces, unless the intramolecular bonds arebetween H‐O, H‐F, or H‐N, in which case they are HydrogenBondsTry these: HClIntra: Polar CovalentInter: Dipole‐Dipole I2Intra: Non‐polar CovalentInter: London Forces NH3Intra: Polar CovalentInter: Hydrogen NaBrIntra: IonicInter: Ionic8
Chemistry 11 Bonding Notes KeyHOMEWORK: Complete Part 2 of the Intramolecular andIntermolecular Bonding Worksheet Hebden p. 180‐182 #73, 76, 80, 81, 83Lewis Structures Diagrams that show how valence electrons aredistributed in an atom, ion or molecule Also called electron dot diagrams Each valence electron is represented as a dot Electron pairs shared between atoms form bonds(represented with a line) Unshared pairs are called “lone pairs”4.Drawing LewisStructuresLewis Bonding Valence e‐ are the players inbonding Valence e‐ transfer leads toionic bonds. Sharing of valence e‐ leads tocovalent bonds. e‐ are transferred or shared tomake each atom isoelectronic witha noble gas the octet. (s2p6)The Octet RuleAtoms tend to gain, lose, or share electronsuntil they have eight valence electrons (ands2p6 arrangement).Hydrogen is anexception. It likesto gain oneelectron to have 2electrons (likeHelium) s29
Chemistry 11 Bonding Notes Key Recall how to determine the valence electron for the elements based on position in theperiodic table.012 Lewis Dot Diagrams represent the valence electrons, as they are theelectrons that are available for reaction.3 4 5 6 7Ionic vs Covalent Compounds Formation of sodium chloride:Na ClNa Cl A valid Lewis structure should have an octet for eachatom except hydrogen (which should have a duet)H2 Formation of hydrogen chloride:H ClH ClA metal and a nonmetal transfer electrons to forman ionic compound. Two nonmetals shareelectrons to form a covalent compound.Double and Triple Bonds Atoms can share four electrons to form a doublebond or six electrons to form a triple bond.O2OOO OO ON2NNNNNLewis Structures For Covalent CompoundsN Triple bonds are the shortest and strongestwhereas single bonds are the longest andweakestHHH HH HCl2How to draw a Lewis Structureusing the 'hnsb' SystemWe'll use NCl3 as an example: 1. Find the number of valence electrons you Have. To do this, determine the sum of the number ofvalence electrons for each element in the structure. Thisnumber will always be even. N has 5 valence, and each Cl has 7 , so 26 e‐ 2. Find the total number of electrons Neededby all elements in the structure to satisfy theOCTET rule. Every element needs 8 (excepthydrogen which needs 2).N needs 8 , each Cl needs 8 , so 32 e‐10
Chemistry 11 Bonding Notes Key How to draw a Lewis Structure3. Find the total number of electrons Sharedbetween all elements. This is found byNeed – Have. For NCl 3 , 32 ‐ 26 6 e ‐ 4. Find the number of Bonds by dividingShared by 2.For NCl 3 , 6 / 2 3 bondsOther helpful hints:A. Connectivity ‐ From the ChemicalFormula, determine the atomconnectivity for the structure.i. Given a chemical formula, ABn , A is thecentral atom and B surrounds the A atom.i.e., In NH3 , NCl3 & NO2 the N is thecentral atom in the structureii. H and F are never central atoms.So, in NCl3 , N is central and Cl are terminalB. You can calculate the number of Non ‐Bonding Electrons (nbe) byHave – SharedIn NCl3 , 26 ‐ 6 20 e‐C. The number of bonds each particularatom tends to make depends on the numberof its valence electrons. See next slide.So, N likes to make 3 bonds, and each Cl 1bond. Thus the central N will have a singlebond to each of the Cl atoms.Example: NCl3h: 26n: 32s: 6b: 3nbe: 20central: N (3 bonds)terminal: Cl (1 bond)ClNUsing the hnsb SystemClClCH4CH H H Hh: 8Cl N ClClCl N ClCln: 16s: 8b: 4nbe: 0bonds: C makes 4 bonds and each H makes 1 bond11
Chemistry 11 Bonding Notes KeyF2h: 14n: 16s: 2b: 1nbe: 12bonds: F 1N2h: 10n: 16s: 6b: 3nbe: 4bonds: N 3HFh: 8n: 10s: 2b: 1nbe: 6bonds: H 1, F 1H 2Oh: 8n: 12s: 4b: 2nbe: 4bonds: H 1, O 2Sometimes, the way a formula iswritten, gives some information as tohow the atoms are connected.FFNNHFHHOHClh: 8n: 10s: 2b: 1nbe: 6bonds: H 1, Cl 1CO2h: 16n: 24s: 8b: 4nbe: 8bonds: C 4, O 2NH3 h: 8n:14s: 6b: 3nbe: 2bonds: N 3, H 1COCl2h: 24n: 32s: 8b: 4nbe: 16bonds: C 4, O 2, Cl 1HOClh: 14n: 18s: 4b: 2nbe: 10bonds: H 1, O 2, Cl 1CH3OHh: 14n: 24s: 10b: 5nbe: 4bonds: C 4, H 1, O 2HClCOONHHHCOClClHOClCHHHOH12
Chemistry 11 Bonding Notes KeyHOMEWORK:Lewis Structures Worksheet Part 1, Set A onlyCovalent Bonding Pictoral Summary5. Lewis Structuresfor Polyatomic IonsLewis Structures for Polyatomic Ions For negatively charged ions, add the correctamount of electrons to your Have group For positively charged ions, subtract thecorrect amount of electrons from your Havegroup Follow the same rules as before (note: you willalways Have an even number of electrons)For most neutral structures, C will make 4bonds, N will make 3, O will make 2, andthe halogens & hydrogen will make 1.For polyatomic ions, the atom's first choicewould be to keep to the # of bonds listedabove, but this is not always the case: C may sometimes only make 3 bonds N can make 2, 3, or 4 bonds O can make 1, 2, or 3 bonds Cl, Br, I can make 1 or 3 bondsCN‐h: 4 5 1 10n: 16s: 6b: 3nbe: 413
Chemistry 11 Bonding Notes KeyFor polyatomic ions, when finished drawingthe Lewis structure, put square bracketsaround it, and the ion charge in the top rightcorner outside of the brackets.NO2 h: 5 6 6 1 18n: 24s: 6b: 3nbe: 12 H3O h: 6 3 1 8n: 14s: 6b: 3nbe: 2 HOMEWORK:Lewis Structures Worksheet Part 1 Set B only6. Exceptions to the Octet RuleThere are numerous exceptions to the octet rule.Electron Deficiencies:Be can be stable with only 4 valence electronsB can be stable with only 6 valence electronsBeF2 Valence Shell Expansion (more than octet): P and Cl can be stable with 10 valence electrons S can be stable with 12 valence electronsThe hnsb system does not work for these exceptions!BF314
Chemistry 11 Bonding Notes KeyValence Shell ExpansionClF3How can these elements have more than an octet?For third‐row elements P, Cl and S, the energetic proximityof the 3d orbitals allows for the participation of theseorbitals in bonding. When this occurs, more than 8electrons can surround a third‐row element. Cl and P can have 10e‐ around it S can have 12e‐SF67. Resonance StructuresWe have assumed up to this point that there is only onevalid Lewis structure for each formula.There are formulas for which more than one Lewisstructure is possible: Different atomic linkages: Structural Isomersex: C4 H10Resonance StructuresO3h: 18n: 24s: 6b: 3nbe: 12 Same atomic linkages, different bonding:Resonance In this example, O3 has two resonance structures:Each structure is perfectly valid, soconceptually, we think of the bondingbeing an average of these twostructures, as the bond lengths areactually the same (see picture below).Electrons are delocalized betweenthe oxygen atoms such that onaverage, the bond strength isequivalent to 1.5 O O bonds.CO32 h: 24n: 32s: 8b: 4nbe: 162 15
Chemistry 11 Bonding Notes KeyHOMEWORK:Lewis Structures Worksheet Part 1 Set C onlyHow do you determine shape? Both bonding pairs and lone pairs need theirspace around the atom ‐ VSEPR Valence Shell Electron Pair Repulsion Electron pairs repel each other ‐‐ thus,maximize distance from other pairs Shape depends on arrangement of bondsand lone pairs around central atom Draw Lewis Structure and then use it todetermine the 3D shape8. Molecular ShapesLinear Bond angle of 180o All 2‐atom moleculesare linear 3‐atom molecules withno lone pairs on centralatom E.g. CO, H2, CO2, BeCl2 (doesn’t obeyoctet rule)Trigonal PlanarTetrahedral Bond angle 120 3 atoms surroundingcentral atom with nolone pairs BF3 (does not obeyoctet rule) CO32‐ Bond angle 109.5o When 4 atomssurround a centralatoms which has nolone pairs E.g. CH4, CCl4o16
Chemistry 11 Bonding Notes KeyTrigonal Pyramidal Bond angle 107.3 lone pairs demand alittle more space than abond 3 atoms surroundcentral atom with onelone pair E.g. NH3, AsH3, H3O oOther ShapesTrigonal bipyramidal: E.g. PCl5Octahedral: E.g. SF6Angular (Bent) Bond angle 104.5o 3 atoms present, witheither 1 or 2 lone pairson central atom E.g. H2O, ClO2‐, O3104.5Predict the shape of the following: H2S AngularHCl LinearSiH4 TetrahedralSO2 AngularHCN LinearPCl3 Trigonal PyramidalCO32‐ Trigonal PlanarNH4 Tetrahedral17
Chemistry 11 Bonding Notes KeyHybrid Shapes:Some structures are a hybrid of 2 or moreshapesCH3COOH:HOMEWORK:Lewis Structures Worksheet Part 2 Sets A, B, & C tetrahedral trigonal planar bent9. Molecular Polarity Earlier this chapter we predicted the polarityof a bond (ionic, polar, or non polar) using theelectronegativity difference between the twoelements Ionic bonds result in full charge (transfer) Polar Covalent bonds result in partial charges– dipoles (unequal sharing) Non polar Covalent bonds result in nocharges (equal sharing) The method uses arrows to show bonddipolesδ δ To determine the polarity of the entiremolecule, we must know the shape of themolecule and the net result of all of the bondpolarities. Two results possible: If there is no charge concentration in any onedirection in the molecule, then NO net dipoleexists, resulting in a non polar molecule If there is a concentration of / charge onany side of the molecule, then a net dipoleexists, resulting in a polar moleculeCO2 , BCl3 , and CCl4 are allnon polar molecules Are the following molecules polar or non polar?CO2, BCl3, CCl4In all cases, the dipoles cancel out,due to the symmetry of the molecule.Thus, there is no net dipole.18
Chemistry 11 Bonding Notes KeyAre the following molecules polar ornon polar? HCl, NH3, H2OH ClHCl:net dipoleNH3:H2O:Polar or non polar?COCl2net dipolenet dipolenet dipolepolarpolarpolarAll of the particles we've investigated sofar are molecules (neutral structures).POLAR!Polar particles dissolve in polar liquids.For example, the polar molecule CH3COOH(acetic acid) dissolves in water (which is polar)to make vinegar. But non polar canola oil doesnot dissolve in water.Non polar molecules dissolve in non polarliquids.Polar liquids are miscible (mix with) otherpolar liquids, but are immiscible (don't mixwith) non polar liquids.Ions have a charge, and are thereforepolar molecules with no furtherinvestigating needed.H2 OCuCl2polarI2non polarpolarCuCl2 dissolvesin H2Opolar polarI2 does notdissolve in H2Ononpolar polarC6H12non polarCuCl2 does notdissolve in C6H12polar nonpolarI2 does dissolvein C6H12nonpolar nonpolarSolids: CuCl2 Liquids: H2OC6H12I2http://www.dlt.ncssm.edu/core/Chapter10 Intermolecular Forces/Chapter10 Animations/Polar vs Nonpolar.html19
Chemistry 11 Bonding Notes KeyHOMEWORK:Lewis Structures Worksheet Part 3 Sets A, B, & C20
8 Molecules with hydrogen bonding have unusually high melting and boiling points due to the hydrogen bonds holding particles close together hydrogen bonding dipole dipole bonding any substance that has hydrogen bonding has an unusually high solubility in water sin
Modern Chemistry 1 Chemical Bonding CHAPTER 6 Chemical Bonding SECTION 1 Introduction to Chemical Bonding OBJECTIVES 1. Define Chemical bond. 2. Explain why most atoms form chemical bonds. 3. Describe ionic and covalent bonding. 4. Explain why most chemical bonding is neither purely ionic or purley 5. Classify bonding type according to .
In Grade 9, you have learned about chemical bonding and its types such as ionic, covalent and metallic bonding and their characteristics. In this unit, we will discuss some new concepts about chemical bonding, like molecular geometry, theories of chemical bonding and much more. Activity
comparing with Au wire bonding. Bonding force for 1st bond is the same range, but approx. 30% higher at 2nd bonding for both Bare Cu and Cu/Pd wire bonding but slightly lower force for Bare Cu wire. Bonding capillary is PECO granular type and it has changed every time when new cell is used for bonding
from electric shock. Bonding and earthing are often confused as the same thing. Sometimes the term Zearth bonding is used and this complicates things further as the earthing and bonding are two separate connections. Bonding is a connection of metallic parts with a Zprotective bonding conductor. Heres an example shown below.
bonding in several substances. Explain, in terms of valence electrons, why the bonding in methane (CH4) is similar to the bonding in water (H2O). In both CH4 & H2O the valence electrons are shared to form covalent bonds. Explain, in terms of valence electrons, why the bonding in HCl is different than that bonding in NaCl.
MERRYLAND HIGH SCHOOL ENTEBBE S.2 CHEMISTRY NOTES BONDING AND STRUCTURE NOTES BONDING Bonding is the chemical combination of atoms or elements to form compounds. The force of attraction holding atoms or elements together in a molecule/crystal is referred to as a chemical bond. Chemical bonding /combination occurs mainly in four forms
Pure covalent bonding only occurs when two nonmetal atoms of the same kind bind to each other. When two different nonmetal atoms are bonded or a nonmetal and a metal are bonded, then the bond is a mixture of cova-lent and ionic bonding called polar covalent bonding. Covalent Bonding In METALLIC BONDING the valence electrons are
non-bonding e 0 1/2 bonding e 1 formal charge 0 O: orig. valence e 6 non-bonding e 4 1/2 bonding e 2 formal charge 0 Example: H 2 O H:O:: Total valence electrons Formal Charge Total non-bonding
