Covalent Bonding - Firelands Elementary School

Covalent BondingCHAPTER 9

I. The Covalent BondCovalent bond: chemicalbond resulting from thesharing of valence electronsA.1.Occurs when atoms havesimilar attractions for electrons –neither atom can gain or loseelectrons, so they shareCovalent bond:

B.Molecules are formed when two or more atoms bondcovalently; considered individual units (ex: H2O is onemolecule)1.a.Diatomic molecules: molecules composed of only two atomsMay be composed of atoms of 2 elements (ex: NaCl, CO, NO)May be composed of 2 atoms of the same elementb.i.H O N Cl Br I orineiodineO2Cl2I2

2.a.Polyatomic molecules: composed of three or more atomsEx: H2O, CCl4, CO2C2H4FeOHClpolyatomicdiatomic*(with ionicbonding)diatomicCl2diatomicP4polyatomic

C. Single Covalent Bonds1.Formed when one pair of electrons is shared2.Ex: H2H H H:Hcan be written as H:H or H-H3.H:H or H-H is an example of a Lewis structurea.b.Use electron-dot diagrams to represent bonding in moleculesOne line or one pair of dots represents one covalent bondSingle bonds are also called sigma (σ) bonds4.a.Form when orbitals overlap: can be two “s”, one “s” and one “p”, ortwo “p”

D. Multiple Covalent Bonds1.Formed when two or more pairs of electrons are shared2.Examples:double bondtriple bondtriple bondDouble & triple bonds contain pi bonds, when parallel orbitalsoverlap to share electrons3.a.b.Double bonds: one sigma, one pi (two shared pairs)Triple bonds: one sigma, two pi (three shared pairs)

E. Strength of Covalent Bonds1.Bond length: distance between bonded atoms2.Rule: shorter bond length stronger bond3.Bond length is longest in single bondsshorter in double bonds,shortest in triple bonds4.So .weakeststrongerstrongest

5.a.b.c.Bond dissociation energy: amount of energy required to break aspecific covalent bondEx: F2 (single)159 kJ/molO2 (double)494 kJ/molN2 (triple)945 kJ/molAmount of energy available in a compound is the sum of all bonddissociation energies of all bonds in the compoundIn chemical rxns, bonds are formed, which gives off energy, andbroken, which takes in energy – leads to overall loss or gain ofenergyi.ii.If more energy is taken in (gained): endothermicIf more energy is given off (lost): exothermic

Classwork p. 244 #1-5 (see example problem 9-1 on same pagefor guidance) p. 247 #6-12

II. Naming MoleculesA. Binary Molecular Compounds1.First element in the formula is first – use entire element name2.Second element is named using the root of the element name &adding suffix “–ide”3.Prefixes are used to indicate the number of atoms of each typepresent in the compound

Examples:4.a.b.c.P2O5COCO2Some molecular compounds were named before the modernnaming system & are still known by their common names5.a.Examples:i.Water: H2O, dihydrogen monoxideii.Ammonia: NH3, nitrogen trihydrideiii.laughing gas: N2O, dinitrogen monoxide, nitrous oxide

B. Acids1.Formed by dissolving a dry molecular compound in water,producing H ions2.Naming Binary Acidsa.b.c.d.Binary acids contain hydrogen & one other elementOne word, followed by “acid”Use the prefix “hydro-” to represent hydrogen, then the root of thesecond element with the suffix “-ic”Examples:i.HCl: hydrochloric acidii.HBr: hydrobromic acidiii.HF: hydrofluoric acid

Naming Oxyacids3.a.b.c.d.e.Oxyacids contain hydrogen and an oxyanion (a polyatomic ioncontaining oxygen)One word, followed by “acid”Look up the name of the oxyanion (on PT)Use the root of the oxyanion followed by a suffix:i.”-ic” if the suffix of the ion is “-ate”ii.“-ous” if the suffix of the ion is “-ite”Examples:i.HNO3: nitric acidii.HNO2: nitrous acid

Common Acids (MUST KNOW!)4.a.b.c.d.e.f.g.h.HCl: hydrochloric acid (stomach acid)H2SO4: sulfuric acid (acid rain)HNO3: nitric acid (fertilizer production)H2CO3: carbonic acidH3PO4: phosphoric acid (pop)HClO4: perchloric acid (explosive!)CH3COOH: acetic acid (vinegar!)(COOH) 2: oxalic acid (organic – plants)

Homework Copy Figure 9-9 (p. 251) p. 249 #13-17 p. 250 #18-22 p. 251 #23-29 Worksheet!

III. Molecular StructuresA. Structural Formulas1.Definition: molecular model using letter symbols & bonds to showrelative positions of atoms2.Can be predicted by drawing the Lewis structureB. Steps to draw a Lewis structure:1.2.Find the total number of valence electrons – this is the total numberof electrons available for bondinga.Divide # above by 2 to find the total number of electron pairsDetermine atom locationsa.The central atom will be the one that can form multiple bondsb.The terminal atoms (at the end) will usually only form singlebonds. Hydrogen is ALWAYS a terminal atom

3.4.5.Draw the skeleton structure – the symbols of atoms in their predictedarrangementPlace single bonds between the central atom and each of the terminalatomsa.Subtract the # of pairs used from the total # of bonding pairs foundin step 1– this # represents all of the lone pairs and any double ortriple bondsb.Place lone pairs around the terminal atoms if necessary to satisfy theoctet rule – if there are odd numbers on any terminal atoms whenyou’ve done this, that means there will be a double bond betweenthat atom & the central atomCheck how many electron pairs are around the central atom. If thereare not 4 pairs, there will need to be double or triple bonds.a.Change one or two of the pairs on a terminal atom to a double ortriple bond between that terminal atom & the central atom to satisfythe central atom’s octet

6.Examples:a.Water: H2Ob.Carbon dioxide: CO2c.Phosphate ion: PO43-

Homework Copy Figure 9-10 (p. 252) p. 255 #30-34 p. 874 #1 a-d under Section 9-1 p. 875 #4 a-f under Section 9-3

C. Resonance Structures1.Some compounds can have more than one valid Lewis structure –most often when there are both double & single bonds2.In these cases, the bonds “resonate” or switch off betweenstructures3.Results in a more stable configuration with average bondstrengths & average bond lengths (bond length is not actuallysingle or double – it’s in between)4.Differ only in the position of electron pairs, not the atom location

D. Exceptions to Octet Rule (threepossible situations)Molecule has odd number of valenceelectrons & cannot pair each electron1.a.Example: NO2, ClO2, NOSome compounds form with fewer than 8electrons present around an atom, withelectrons still paired up (even # 8)2.a.b.Example: BH3Coordinate covalent bond: when one atom thathas lone pairs will contribute an electron pair toanother atom, which is short on electrons

Expanded octet: central atom has more than 4 electron pairs3.a.b.c.These molecules will often have 5 or 6 pairs around the centralatomCommon in highly reactive atomsExample: PCl5, SF6

Homework p. 256 #35-38 p. 258 #39-48 Worksheet!

Lab 8.1 Tips Groups 1 & 5: start building your models Groups 2 & 6: start with aspirin Groups 3 & 7: start with acetaminophen Groups 4 & 8: start with ibuprofen Rotate models to the group with the next number (4 should rotate to1, 8 should rotate to 5) Each group has their own modeling kit, so you do NOT need to rotate that –just the 3 models that I built! Long gray connectors are double bonds, short gray are single bonds When you are finished, take your models apart & put them back inthe bag – count how m any of each part you have, check off on theindex card in your bag, and each group member should initial at thebottom of the card confirming that all parts are back in place

IV. Molecular ShapeA. VSEPR Model1.Valence Shell Electron Pair Repulsion: model based on anarrangement that minimizes the repulsion of shared & unsharedpairs of electrons around the central atom2.Atoms exist at fixed angles to one another3.All electron pairs repel each othera.b.Shared pairs repel each other equallyLone pairs occupy a slightly larger orbital than shared pairs, pushingthe shared pairs toward one another

B. Possible Geometric Configurations1.Linear: atoms in a straight line or form 180 anglea.b.2 pairs of shared electrons; no lone pairsExamples: Cl2, BeF2Trigonal planar: atoms form a triangle at 120 from each other2.a.b.3 pairs of shared electrons; no lone pairsExamples: BF3, AlCl3Tetrahedral: atoms form a three-sided pyramid at 109.5 fromeach other3.a.b.4 pairs of shared electrons; no lone pairsExample: CH4

Trigonal pyramidal: atoms form an uneven pyramid with bondangles of 107.3 4.a.b.3 pairs of shared electrons; 1 lone pairExamples: PH3, NH3Bent: atoms form a bent shape with bond angles of 104.5 5.a.b.2 pairs of shared electrons; 2 lone pairsExample: H2OTrigonal bipyramidal: three atoms around central in the samehorizontal plane, and one atom directly above & one atom directlybelow the central in the vertical plane6.a.b.c.Forms 2 pyramids (top & bottom); bond angles are 90 between thevertical and horizontal atoms, and 180 between atoms in the sameplane5 pairs of shared electrons; no lone pairsExamples: NbBr5, PCl5

Octahedral: four atoms around central in the same horizontalplane, and one atom directly above & one atom directly below thecentral in the vertical plane7.a.b.c.Bond angles are 90 between the central atom and each terminalatom6 pairs of shared electrons; no lone pairsExample: SF6HOMEWORK: copy Table 9-3 on p. 260 – add a columnbetween “Example” and “Total Pairs” for Lewis structures &draw those!

C. Hybridization1.Atomic orbitals merge to form hybrid/combination orbitals atbonding sites2.Explains why all bonding orbitals around certain atoms are thesame instead of having bonding “s” orbitals, bonding “p” orbitals,etc.3.This involves moving electrons to higher orbitals in order to freeup bonding sites

Variations of hybrid orbitals4.a.b.c.d.e.sp orbitals: 1 “s” 1 “p” orbital merge; forms triple bonds; linearstructuresp2 orbitals: 1 “s” 2 “p” orbitals merge; forms double bonds;planar structuresp3 orbitals: 1 “s” 3 “p” orbitals merge; forms single bonds;tetrahedral structuresp3d orbitals: 1 “s” 3 “p” 1”d” orbital merge; forms single bonds;trigonal bipyramidal structuresp3d2 orbitals: 1 “s” 3 “p” 2”d” orbital merge; forms singlebonds; octahedral structure

Homework p. 262 #49-59

V. Electronegativity & PolarityA. Electronegativity1.Electronegativity (EN): relative ability of an atom to attractelectrons in a chemical bonda.b.c.EN is an average value; slightly different for different bondsMost active metals have lowest EN, most active nonmetals havehighest ENEN is a dimensionless quantity (not positive or negative)

B. Bond Character1.Atomic bonds are not just purely ionic or covalent; it is acontinuum of bond types2.Bond strength depends on the EN difference between the twoatomsa.b.c.Greater difference stronger bondWhen EN difference is large, one atom has a higher attraction & pullselectrons from the other atom to itself – forms ionic bondWhen EN difference is small, atoms have nearly the same attractionfor electrons, so they share – forms covalent bond

When EN difference is 1.70, a bond is 50% ionic & 50% covalent3.a.b.c.EN difference 1.70 shows ionic characterEN difference 1.70 shows covalent characterExamples:i.NaCl:Na 0.93Cl 3.16difference 3.16 – 0.93 2.232.23 1.70ionicii.H2O:H 2.20O 3.44difference 3.44 – 2.20 2.231.24 1.70covalentiii.O2:O 3.44O 3.44difference 3.44 – 3.44 zerozero 1.70100% covalent / nonpolarcovalent