Unit 6

Unit 6

Unit 6Covalent BondingLesson 1:Covalent BondingCovalent bonds: atoms held together by sharing electrons.Molecules: neutral group of atoms joined together by covalent bonds.Diatomic molecule: molecule consisting of 2 atoms. Remember them: F2, Cl2, I2, Br2, H2, N2, O2Molecules tend to have lower melting and boiling points than ionic compounds.YouTube - Making Molecules with Atoms

Molecular FormulaShows how many atoms of each element a molecule contains.Naming binary molecular compoundsComposed of two nonmetals; often combine in more than one way. Ex. CO and CO2Prefixes are used to name binary molecular compounds.

PrefixMono-Di-Tri-Tetra-Penta-Hexa-Hepta-Octa-Nona-Deca-Number12345678910Binary Compounds Containing Two NonmetalsTo name these compounds:give the name of the less electronegative element first with the Greek prefix indicating the number of atoms of that element present After give the name of the more electronegative non-metal with the Greek prefix indicating the number of atoms of that element present and with its ending replaced by the suffix ide. Do not use the prefix mono- if required for the first element.4Binary Molecular CompoundsN2Odinitrogen monoxideN2O3dinitrogen trioxideN2O5dinitrogen pentoxide

ICliodine monochlorideICl3iodine trichloride

SO2sulfur dioxideSO3sulfur trioxideYouTube - Naming molecular compounds5Binary Molecular CompoundsContaining Two Nonmetals________________ diarsenic trisulfide

________________sulfur dioxide

P2O5____________________

________________ carbon dioxide

N2O5____________________

H2O____________________As2S3SO2diphosphorus pentoxideCO2dinitrogen pentoxidedihydrogen monoxide6Naming Binary CompoundsBinary Compound?Metal Present?Does the metal formmore than one cation?MoleculeUse Greek PrefixesIonic compound (cation has one charge only)Use the elementname for the cation.Ionic compound (cation has more than one charge) Determine the Charge of the cation; use a Romannumeral after the cation name.YesYesYesNoNoZumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 987Classwork #1:Do handout Naming MoleculesLesson 2: The Nature of Covalent BondingIntroduction with balloon activityoctet rule: electron sharing occurs usually so that atoms attain the electron configurations of noble gases.Single covalent bond: two atoms held together by sharing a pair of electrons. Shown as two dots or as a long dash.A pair of valence electrons that is not shared between atoms is called an unshared pair.HHOHOHorOHHDouble bonds: covalent bond formed by sharing two pairs of electronsTriple bonds: covalent bond formed by sharing three pairs of electrons.

hydrogen

chlorine

iodine

nitrogenCovalent bonding with equal sharing of electrons occurs in diatomic molecules formed from one element.A dash may replace a pair of dots.

Classwork: introduction to lewis structures.Lesson 3:Molecular StructureStructural formula: uses symbols and bonds to show relative position of atoms. Steps to determine Lewis structures for moleculesPredict the location of certain atoms.Hydrogen is always an end atomThe least electronegative atom is the central atom (usually the one closer to the left on periodic table)Find the total number of electrons available for bonding. (# of valence electrons of atoms in molecule)Determine the number of bonding pairs by dividing the total number of electron by 2

Place one bonding pair (single bond) between central atom and terminal atoms. Subtract pairs used in step 4 from bonding pairs in step 3. Place lone pairs around each terminal atom bonded to the central atom to satisfy the octet rule. Any remaining pairs are assigned to the central atom.If the central atom does not have an octet, convert one or two of the lone pairs on the terminal atoms to a double or a triple bond between central and terminal atom. Some elements like Be, B, Al do not form a complete octet, S and P can have more than 8 valence electrons.

Ex. 1 Draw the lewis structure for ammonia, NH3Hydrogen is an end atom and nitrogen is the central atom.Total number of valence electrons: (1 nitrogen x 5 valence electrons)+ (3 hydrogens x 1 valence electron)= 8 valence electrons.Total number of bonding pairs= 8/2 = 4 Draw single bond from each H to NNHHHEx. 1 Draw the lewis structure for ammonia, NH3Subtract the number of pairs of electrons used from the total pairs of electrons: 4-3 =1 pair available

One lone pair remains, hydrogen can have only one bond, assign the lone pair to the central atom, N.

NHHHEx. 2 Draw the lewis structure for carbon dioxide, CO2Oxygen atoms are end atoms and carbon is the central atom.Total number of valence electrons: (1 carbon x 4 valence electrons)+ (2 oxygen x 6 valence electron)= 16 valence electrons.Total number of bonding pairs= 16/2 = 8 Draw single bond from each C to OCOOEx. 1 Draw the lewis structure for carbon dioxide, CO2Subtract the number of pairs of electrons used from the total pairs of electrons: 8-2 =6 pair available

Add three pairs of electrons to each oxygen.

COOEx. 1 Draw the lewis structure for carbon dioxide, CO2No lone pairs remain for carbon. Carbon does not have an octet, use a lone pair from each oxygen to form a double bond with the carbon atom.

COOCOOCW: lewis structures handout part 1Lesson :4 Exception to octet ruleSome molecules have an odd number of valence electrons and cannot form an octet around each atom.Some molecules form with fewer than eight electrons present around an atom. Ex. BoronSome compounds have central atoms with more than 8 electrons. This is called an expanded octet. Ex. S, Xe and PEx. 3 Draw the lewis structure for XeF4 (exception octet rule)F is an end atom and nitrogen is the central atom.Total number of valence electrons: (1 xenon x 8 valence electrons)+ (4 fluorines x 7 valence electron)= 36 valence electrons.Total number of bonding pairs= 36/2 = 18Draw single bond from each F to XeXeFFFFEx. 1 Draw the lewis structure for XeF4 (exception octet rule)Subtract the number of pairs of electrons used from the total pairs of electrons: 18-4 =14 pairs available14 lone pairs remain, place them around each fluorine so that each fluorine has 8 valence electrons

XeFFFFEx. 1 Draw the lewis structure for XeF4 (exception octet rule)There are 2 pairs of electrons still available, place around Xe which is capable of having more than 8 valence electron.

XeFFFFMolecular ShapeVSEPR (Valence shell electron pair repulsion) ModelThe repulsion between electron pairs in a molecule result in atoms existing at fixed angles from each other. (Remember balloon activity)Shared electron pairs repel each otherA greater repulsion occurs between unshared electron pairs and shared electron pairs. Bonding and Shape of Molecules:Count number of bonds and unshared pairs of electrons AROUND CENTRAL ATOM and then use table below to determine shape of molecule.

Number of BondsNumber of Unshared PairsShapeExamples2

3

4

3

20

0

0

1

2Linear

Trigonal planar

Tetrahedral

Pyramidal

BentBeCl2

BF3

CH4, SiCl4

NH3, PCl3

H2O, H2S, SCl2-Be- B C N: O::CovalentStructure27Use table on last slide to determine shape of molecule.HHOShape: bent2 bonds and 2 unshared pairsSO228Carbon tetrachlorideCClClClClCCl4C109.5oClClClClCarbon tetrachloride carbon tet had been used as dry cleaning solventbecause of its extreme non-polarity.Shape: Tetrahedral4 bonds and 0 unshared pairs.29Classwork: do in your notebookDetermine the shape for the following molecules (first draw the lewis structure for the molecule and then use the table on slide #7 to determine the shape taking in consideration the number of bonds and unshared pairs of electrons around the CENTRAL ATOM.)BF3 2. OCl2 3. CF4

4. NH3 5. BeI2Strength of covalent bonds:Depends on distance between bonded nuclei. The shorter the bond length the stronger the bond.Triple bond strength> double bond strength> single bond strength

Electronegativity and polarityThe type of bond can be predicted by using the electronegativity difference of the elements that are bonded.

For identical atoms, the bond they form is a nonpolar bond because the pair of electrons is shared equally.Bonds between different atoms can be ionic or covalent.If electronegativity difference is greater than 1.70 it is considered an ionic bond.If electronegativity diffence is less than 1.70 it is considered a covalent bond.If electronegativity difference is 0, the bond is nonpolar covalent.Polar covalent bondsForm when pair of electrons is not shared equally by bonding atoms (like a tug-of-war)Partial charges occur at the ends of the bond. Using the symbols - , partially negative, and +, partially positive, next to the model of a molecule indicates the polarity of the polar covalent bond.

Molecular PolarityMolecules are either polar or nonpolarThe nature of the covalent bond and the shape of the molecule result in a polar or nonpolar molecule.Symmetric molecules tend to be nonpolarH-H (H2) has a nonpolar bond thus is a nonpolar molecule.H2O has polar bonds and is a polar molecule.

CO2 has polar bonds but do to the molecules shape is a nonpolar molecule.

Polar bonds in this molecule are opposite to each other and cancel each other, so molecule is nonpolar.CH4 has polar bonds but do to the molecules shape is a nonpolar molecule. Polar bonds in this molecule are opposite to each other and cancel each other, so molecule is nonpolar.

CHCl3 has polar bonds but do to the molecules shape is a nonpolar molecule. Polar bonds in this molecule do not cancel each other, so molecule is polar.

Classwork