Chemical Bonding - WordPress
Chemical Bonding
Introduction to Chemical Bonding
Chemical bond ! is a mutual electrical attractionbetween the nuclei and valence electrons of differentatoms that binds the atoms togetherWhy are most atoms chemically bonded to each other?
As independent particles, they are at relativelyhigh potential energyNature, however, favors arrangements in whichpotential energy is minimizedThis means that most atoms are less stableexisting by themselves than when they arecombinedBy bonding with each other, atoms decrease inpotential energy, thereby creating more stablearrangements of matter
Types of Chemical BondingBond – valence electrons rearranged tomake atom more stableWay they are rearranged depends on typeof bondIonic bonding ! chemical bonding thatresults from the electrical attraction between
Covalent bonding ! results from thesharing of electron pairs between twoatomsIn purely covalent bond, electrons sharedequally between two atoms
Ionic or Covalent?Bonding is rarely purely one or the otherDepending on how strongly the atoms attractelectrons, falls somewhere betweenElectronegativity (EN)! atom’s ability to attractelectronsDegree of bonding between atoms of 2 elementsbeing ionic or covalent estimated by calculatingdifference in elements’ ENs
ExampleFluorine’s EN 4.0,Cesium’s EN 0.74.0-0.7 3.3According to table, F-Cs isionicThe greater the difference,the more ionic the bond
Bonding between atoms with ENdifference of less than or equal to ( ) 1.7has ionic character less than or equal to( ) 50%Classified as covalentBonding between atoms of same elementis completely covalent
Nonpolar-covalentH-H bond has 0% ionic characterNonpolar-covalent bond ! a covalent bond in which thebonding electrons are shared equally by the bonded atoms,resulting in a balanced distribution of electrical charge0-5% ionic character (0-0.3 EN difference) is nonpolar-covalentbond
Polar-covalent BondsBonds that have significantly differentEns, electrons more strongly attracted bymore-EN atomThese bonds are polar ! they have anuneven distribution of charge
Covalent bonds with 5-50% ionic character (0.3-1.7 ENdifference) are polarδ δ Polar-covalent bond ! covalent bond in which thebonded atoms have an unequal attraction for the sharedelectrons
Sample ProblemUse electronegativity differences to classifybonding between sulfur, S, and the followingelements: hydrogen, H; cesium, Cs; andchlorine, Cl. In each pair, which atom will bemore negative?
Bonding betweensulfur andEN differenceBond typeMore-negativeatomHydrogen2.5-2.1 0.4Polar-covalentSulfurCesium2.5-0.7 1.8IonicSulfurChlorine3.0-2.5 0.5Polar-covalentchlorine
Practice ProblemUse electronegativity differences to classifybonding between chlorine, Cl, and thefollowing elements: calcium, Ca; oxygen, O;and bromine, Br. Indicate the more-negativeatom in each pair.
Bonding betweenchlorine andEN differenceBond typeMore-negativeatomCalcium3.0-1.0 2.0IonicChlorineOxygen3.5-3.0 0.5Polar-covalentOxygenBromine3.0-2.8 0.2NonpolarcovalentChlorine
Section 2 – Covalent Bonding and MolecularCompounds
Many chemical compounds are moleculesMolecule ! neutral group of atoms that areheld together by covalent bondsSingle molecule of compound is individual unitCapable of existing on its ownMay consist of 2 or more atoms of same elementor two or more different atomsMolecular compound ! chemicalcompound whose simplest units are molecules
Formation of Covalent BondBonded atoms have lower potentialenergy than unbonded atomsAt large distance atoms don’t influenceeach otherPotential energy set at 0
Each H has ( ) protonNucleus surrounded by (-) electronAs atoms near each other, chargedparticles start to interact
Approaching nuclei andelectrons are attracted toeach otherDecrease in total potentialenergyAt the same time, twonuclei and two electronsrepel each otherIncrease in potentialenergy
The amount of attraction/repulsion depends on howclose the atoms are to each otherWhen atoms first “see” each other, electron-protonattraction stronger than e-e or p-p repulsionsSo atoms drawn to each other and potential energylowered
Attractive force dominates until adistance is reached where repulsionequals attractionValley of the curve
Closer the atoms get, potential energyrises sharplyRepulsion becomes greater thanattraction
Bonded atoms vibrate a bitAs long as energy stays close to minimum theystay covalently bondedBond length ! the distance between twobonded atoms at their minimum potential energy(average distance between two bonded atoms)
To form covalent bond, hydrogen atomsneed to release energyAmount of energy equals differencebetween potential energy at zero level(separated atoms) and at bottom of valley(bonded atoms)
Same amount of energy must be added toseparate bonded atomsBond energy ! energy required tobreak a chemical bond and form neutralisolated atoms
Units of bond energy usually kJ/molIndicates energy required to break one moleof bonds in isolated moleculesEx. 436 kJ/mol is energy needed to break HH bonds in 1 mol hydrogen molecules andform 2 mol of separated H atomsBond lengths and bond energies vary with thetypes of atoms that have combined
All individual H atoms contain single,unpaired e in 1s orbitalSharing allows electrons to experienceeffect of stable electron configuration of2helium, 1s
The Octet RuleNoble-gas atoms have minimum energy existingon their own b/c of electron configurationsOuter orbitals completely fullOther atoms fill orbitals by sharing electronsBond formation follows octet rule ! chemicalcompounds tend to form so that each atom, bygaining, losing, or sharing electrons, has an octetof electrons in its highest occupied energy level
Example: Bonding of Fluorine2 F atoms bond to form F27 e in highest energy level
Example: HCl
Exceptions to Octet RuleMost main-group elements form covalent bondsaccording to octet ruleEx. H-H only 2 electrons21Boron, B, has 3 valence electrons ([He]2s 2p )Boron tends to form bonds where it is surroundedby 6 e- (e- pairs)Others can be surrounded by more than 8 whenbonding to highly electronegative elements
Electron Dot NotationTo keep track of valence electrons, it is helpfulto use electron dot notation ! electronconfiguration notation in which only the valenceelectrons of an atom of a particular element areshown, indicated by dots placed around theelement’s symbolInner-shell electrons NOT shown
Sample Problem 1Write the electron-dot notation forhydrogen.A hydrogen atom has only one occupiedenergy level, the n 1 level, which contains asingle electron. So, e-dot notation is writtenasH
Sample Problem 2Write the e-dot notation for nitrogen.23Group notation for nitrogen’s family is ns npwhich means nitrogen has 5 valenceelectrons. E-dot notation is written as
Lewis StructuresE-dot notation can also be used to represent moleculesEx. H2 represented by combining notations of 2 individualH atomsPair of dots representsebeing shared
F2Each F atom surrounded by 3 pairs ethat are not shared in bondsUnshared (lone) pair ! pair of e- thatis not involved in bonding and thatbelongs completely to one atom
Lewis StructuresPair of dots representing shared pair incovalent bond often replaced by long dashH HH HLewis structures ! formulas in whichatomic symbols represent nuclei and innershell electrons, dot-pairs or dashes betweentwo atomic symbols represent electron pairs incovalent bonds, and dots next to only oneatomic symbol represent unshared electrons
Common to write Lewis structures that showonly shared e using dashesStructural formula ! indicates the kind,number, arrangement, and bonds but not theunshared pairs of ato in a moleculeF-FH - Cl
Lewis structures and structural formulas formany molecules can be drawn if you know thecomposition of the molecule and which atomsare bonded to each otherSingle bond ! covalent bond made bysharing of one pair of e between 2 atoms
Sample ProblemDraw the Lewis structure of iodomethane,CH3I.1. Determine type and number of atoms inmolecule.1 C, 1 I, 3 H
2. Determine the total number of valencethe atoms to be combined.C 1 x 4e- 4eI 1 x 7e- 7e-ein
3. Arrange the atoms to form a skeletonstructure for the moleculeIf carbon is present, it is the central atomOtherwise, the least-electronegative atom iscentral (except for hydrogen which isNEVER central)Then connect the atoms by electron-pairbonds.
4. Add unshared pairs of electrons so thateach hydrogen atom shares a pair ofelectrons and each other nonmetal issurrounded by 8 electrons.
5. Count the electrons in the structure to be surethat the number of valence e- used equals thenumber available. Be sure the central atomand other atoms besides H have an octect.There are eight e- in the four covalent bondsand six e- in the three unshared pairs, givingthe correct total of 14 valence electrons
Practice ProblemDraw the Lewis structure of ammonia,NH3.
Practice ProblemDraw the Lewis structure for hydrogensulfide, H2S.
Multiple bonds !double and triple bondsDouble bonds have higher bond energies andare shorter than single bondsTriple bonds have higher bond energies and areshorter than double bonds
Multiple Covalent BondsAtoms of same elements (especially C, Nand O) can share more than one e- pairDouble bond ! covalent bond made bythe sharing of two pairs of e- between twoatoms
Triple bond ! covalent bond made bysharing of 3 pairs of e- between 2 atomsEx. N2Each N has 5 valenceEach N shares 3 e- with other
Practice ProblemDraw the Lewis structure for methanal,CH2O, which is also known asformaldehyde.
Practice ProblemDraw the Lewis structure for carbondioxide.
Practice ProblemDraw the Lewis structure for hydrogencyanide, which contains one hydrogenatom, one carbon atom, and one nitrogenatom.
Resonance StructuresSome molecules/ions cannot berepresented correctly by single LewisstructureEx. Ozone (O3)Each structure has one single and onedouble bond
Chemists used to think ozone spends timealternating or “resonating” between twostructuresNow know that actual structure is something likean average between the twoResonance ! bonding in molecules or ionsthat cannot be correctly represented by a singleLewis structure
To indicate resonance, double-headed arrowplaced between resonance structures
Section 3Ionic Bonding and Ionic Compounds
Ionic BondingIonic compound ! composed of positive andnegative ions that are combined so that thenumbers of positive and negative charges areequalMost exist as crystalline solids, a 3-D network of( ) and (-) ions mutually attracted to one another
Different from molecular compound b/c ioniccompound not made of independent, neutralunitsChemical formula represents simplest ratio ofcompound’s combined ions that give electricalneutrality
Chemical formula of ionic compound shows ratioof ions present in ANY sample of ANY sizeFormula unit ! simplest collection of atomsfrom which an ionic compound’s formula can berecognizedEx. NaCl is formula unit for sodium chlorideOne sodium cation and one chlorine anion
Ratio of ions in formula depends on chargesof ions combinedEx. Calcium and fluorine2 Ca1F total 11F toSo need 2equal 2 (-2) 0Formula unit is CaF2
Formation of Ionic CompoundsE-dot notation can be used to demonstratechanges that take place in ionic bondingDo not usually form by combination of isolatedions
Sodium readily gives up 1 eChlorine readily accepts 1e-
Characteristics of IonicIn ionic crystals, ions minimize potential energyby combining in orderly arrangement called acrystal lattice
Attractive forces: between oppositely chargedions (cations and anions) and between nucleiand electronsRepulsive forces: between like-charged ionsand between electronsCrystal lattice structure represents balancebetween these two forces
Within arrangement, each Na is surroundedby 6 ClAt the same time, each Cl- is surrounded by6 Na
3-D arrangements of ions and strengths ofattraction are different with sizes and charges ofions and number of ions of different chargesEx. CaF2, there are 2 anions for each cation2 Each Ca is surrounded by 8 F2 Each F is surrounded by 4 Ca
Lattice EnergyTo compare bond strengths in ionic compounds,chemists compare amounts of energy releasedwhen separated ions in gas form crystalline solidLattice energy ! energy released when onemole of an ionic compound is formed fromgaseous ions
Comparison of Ionic and Molecular CompoundsForce that holds ions together in ioniccompounds is very strong overall betweenopposite chargesMolecular compound – bonds making up eachmolecule also strong, but forces betweenmolecules not strong
Because of bond strength difference, molecularcompounds melt at lower temperaturesIonic compounds have higher melting andboiling points
Ionic compounds are hard but brittleSlight shift of one row of ions causes largebuildup of repulsive forcesRepulsive forces make layers split completely
In solid state ions cannot move – compoundsare not electrical conductorsMolten state – ions can move freely and cancarry electric currentMany ionic compounds dissolve in waterAttraction to water molecules overcomesattraction to each other
Polyatomic IonsCertain atoms bond covalently to each other toform group of atoms that has molecular ANDionic characteristicsPolyatomic ion ! a charged group ofcovalently bonded atoms
Lewis Structures of Polyatomic IonsPolyatomic ions combine with ions of opposite chargeto form ionic compoundsTo find Lewis structure, follow previous instructionsexceptIf ion is negative, add to the total number of valenceelectrons a number of e same as the ions negativechargeIf ion positive, subtract same number of e- as thepositive charge
Section 4Metallic Bonding
Metallic Bonding is DifferentMetals have unique property of highlymovable electrons (why they conductelectricity so well)In molecular compounds e cannot move, heldin shared bondIn ionic compounds, e cannot move, held toindividual ions
Metallic-Bond ModelHighest energy levels of most metal atoms onlyoccupied by few eEx. s-block metals have one or two valence ewhere all 3 p orbitals are emptyd-block metals have many empty d orbitals justbelow highest energy level
Overlapping OrbitalsWithin metal, empty orbitals in outer energylevels overlapAllows outer e to move freelye are delocalized ! do not belong to any oneatomMetallic bonding ! chemical bonding thatresults from attraction between metal atoms andsurrounding sea of electrons
Metallic PropertieseFreedom of to move around causes highelectrical and thermal conductivity
b/c many orbitals separated by very smallenergy differences, metals can absorb widerange of light frequenciesAbsorption of light excites e to higher energylevelse immediately fall back down to lower levels,giving off light (why metals are shiny)
Malleability ! ability of a substance to behammered or beaten into thin sheetsDuctility ! ability of a substance to bepulled into wiresBoth possible because of structure, one lineof metal atoms can slide without breakingbondsNot possible with ionic crystal structures
Metallic Bond StrengthBond strength varies with nuclear charge of metalatoms and number of e- in metal’s e- seaBoth factors reflected as heat of vaporizationWhen metal vaporized, bonded atoms in solidstate converted to individual atoms in gas stateHigher heat of vaporization, higher bond strength
Section 5Molecular Geometry
Molecular GeometryProperties of molecules depend on bonding ofatoms and the 3-Dimensional arrangement ofmolecule’s atoms in spacePolarity of each bond, along with geometry ofmolecule, determines molecular polarity! uneven distribution of molecular chargeStrongly influences forces that act BETWEENmolecules
VSEPR TheoryDiatomic molecules must be linear (only twoatoms)To predict geometries of more complexmolecules, consider locations of all e- pairssurrounding bonded atomsThis is basis of VSEPR
“Valence-shell, electron-pair repulsion”VSEPR theory ! repulsion between the setsof valence-level e- surrounding an atom causesthese set to be oriented as far apart as possibleHow does this account for molecular shape?Let’s consider only molecules with no unsharedvalence e- on central atom
Ex. BeF2Be doesn’t follow octect ruleBe forms covalent bond with each F atomSurrounded by only two electron pairs it shareswith F atomsAccording to VSEPR, shared pairs oriented asfar away from each other as possible
Distance between e- pairs maximized if bonds toF are on opposite sides of Be, 180 apartSo, all 3 atoms lie in straight line – molecule islinear
If we represent central atom in molecule by “A”and atoms bonded to “A” are represented by “B”then BeF2 is an example of an AB2 moleculeAB2 is linearWhat would AB3 look like?
The 3 A-B bonds stay farthest apartby pointing to corners of equilateraltriangle, giving 120 between bonds trigonal-planar geometry
AB4 molecules following octect rule bysharing 4 e- pairs with B atomsDistance between e- pairs maximized ifeach A-B bond points to one of 4 cornersof tetrahedron (tetrahedral geometry)Angle is 109.5
Sample ProblemUse VSEPR theory to predict the moleculargeometry of aluminum chloride, AlCl3.
This molecule is an exception to the octet rulebecause in this case Al forms only three bondsAluminum trichloride is an AB3 type of moleculeTherefore, according to VSEPR theory, it shouldhave trigonal-planar geometry
Practice ProblemUse VSEPR theory topredict the moleculargeometry of thefollowing molecules:a. HIlinearb. CBr4tetrahedralc. AlBr3Trigonal-planard. CH2Cl2tetrahedral
VSEPR and Unshared e- PairsAmmonia, NH3, and water, H2O, are examplesof molecules where central atom has bothshared and unshared e- pairsHow does VSEPR account for the geometries?
Lewis structure of ammonia shows in addition to3 e- pairs it shares with 3 H atoms, the central Nhas one unshared pair of eVSEPR theory says that lone pair occupiesspace around N atom just as bonding pairs doSo, as an AB4 molecule, e- pairs maximizeseparation by assuming 4 corners of tetrahedron
Lone pairs occupy space but description ofshape of molecule refers to positions of atomsonlySo, molecular geometry of ammonia moleculeis pyramid with triangular baseGeneral formula is AB3EE is unshared e- pair
Water molecule has 2 unshared e- pairsIt is AB2E2 moleculeA (O) is at center of tetrahedron2 corners occupied by B (H)Other 2 corners occupied by E (unsharede-)
Molecular ShapeAtoms bonded tocentral atomLone pairs of electronsBond angleLinear20180 Bent or Angular21Less than 120 Trigonal-planar30120 Tetrahedral40109.5
Molecular ShapeAtoms bonded tocentral atomLone pairs of electronsBond angleTrigonal-pyramidal31Less than 109.5 Bent or Angular22Less than 109.5 Trigonal-bipyramidal5090, 120, and 80 Octahedral6090 and 180
Sample ProblemUse VSEPR theory to predict the shapeof a molecule of carbon dioxide, CO2.
HybridizationVSEPR theory useful for explaining shapes ofmoleculesDoesn’t tell the relationship between geometryand orbitals of bonding electronsModel used to explain this is hybridizationMixing of two or more atomic orbitals of similarenergies on the same atom to make new orbitalsof equal energies.
Methane (CH4) provides good exampleOrbital notation of C shows it has 4valence e2 in 2s and 2 in 2pWe know from experiments that methanehas tetrahedral geometryHow does C form 4 equal covalent bonds?
2s and 2p orbitals have different shapesThese orbitals hybridize to form four new, identical3orbitals called sp orbitalsSuperscript 3 shows that 3 p orbitals were includedin hybridization3Sp all have same energyMore than 2sLess than 2pHybrid orbitals ! orbitals of equal energy madeby the combination of two
Ionic bonding! chemical bonding that results from the electrical attraction between . Covalent bonding! results from the sharing of electron pairs between two atoms In purely covalent bond, electrons shared equally between two atoms. Ionic or Covalent?
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.
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
A)Metallic bonding B)hydrogen bonding C)covalent bonding D)ionic bonding 26.The particle diagram below represents a solid sample of silver. Which type of bonding is present when valence electrons move within the sample? A)ionic B)metallic C)nonpolar covalent D)polar covalent 27.Which type of bonding is present in a sample of an element that is .
