2 Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons thatparticipate in chemical bonding.Group# of valence e-e- configuration1A1ns12A2ns23A3ns2np14A4ns2np25A5ns2np36A6ns2np47A7ns2np5
3 Lewis Dot Symbols for the Representative Elements & Noble Gases
4 The Ionic BondIonic bond: the electrostatic force that holds ions together in an ionic compound.Li+F-Li+F1s21s22s22p61s22s11s22s22p5[He][Ne]LiFLiLi+ + e-e- +F-F-Li+ +Li+
5 The mineral corundum (Al2O3). 9.1Use Lewis dot symbols to show the formation of aluminum oxide (Al2O3).The mineral corundum (Al2O3).
6 Electrostatic (Lattice) Energy Lattice energy (U) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions.E is the potential energyQ+ is the charge on the cationE = kQ+Q-rQ- is the charge on the anionr is the distance between the ionsCompoundLattice Energy(kJ/mol)Lattice energy increases as Q increases and/oras r decreases.Q: +2,-1Q: +2,-2MgF2MgO29573938LiFLiCl1036853r F- < r Cl-
7 Born-Haber Cycle for Determining Lattice Energy DHoverall = DH1 + DH2 + DH3 + DH4 + DH5o
9 Why should two atoms share electrons? A covalent bond is a chemical bond in which two or more electrons are shared by two atoms.Why should two atoms share electrons?7e-7e-8e-8e-FF+FLewis structure of F2lone pairsFsingle covalent bondsingle covalent bondF
10 Lewis structure of water single covalent bonds2e-8e-2e-H+O+HOHorDouble bond – two atoms share two pairs of electrons8e-8e-8e-OCorOCdouble bondsTriple bond – two atoms share three pairs of electrons8e-N8e-orNtriple bond
11 Lengths of Covalent Bonds Bond LengthsTriple bond < Double Bond < Single Bond
13 Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atomsHFelectron richregionelectron poorregione- poore- richFHd+d-
14 Electron Affinity - measurable, Cl is highest Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.Electron Affinity - measurable, Cl is highestX (g) + e X-(g)Electronegativity - relative, F is highest
15 The Electronegativities of Common Elements Most electronegative element!
16 Variation of Electronegativity with Atomic Number
17 Classification of bonds by difference in electronegativity Bond TypeCovalent0 < and <2Polar Covalent 2IonicIncreasing difference in electronegativityCovalentshare e-Polar Covalentpartial transfer of e-Ionictransfer e-
18 9.2Classify the following bonds as ionic, polar covalent, or covalent:the bond in HClThe bond in KFthe CC bond in H3CCH3
19 Writing Lewis Structures Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center.Complete an octet for all atoms except hydrogen.If structure contains too many electrons, form double and triple bonds on central atom as needed.
20 NF3 is a colorless, odorless, unreactive gas. 9.3Write the Lewis structure for nitrogen trifluoride (NF3) in which all three F atoms are bonded to the N atom.NF3 is a colorless, odorless, unreactive gas.
21 HNO3 is a strong electrolyte. 9.4Write the Lewis structure for nitric acid (HNO3) in which the three O atoms are bonded to the central N atom and the ionizable H atom is bonded to one of the O atoms.HNO3 is a strong electrolyte.
22 9.5Write the Lewis structure for the carbonate ion ( ).
23 An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.formal charge on an atom in a Lewis structure=total number of valence electrons in the free atom-total number of nonbonding electrons-12total number of bonding electrons()The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.
25 Formal Charge and Lewis Structures For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present.Lewis structures with large formal charges are less plausible than those with small formal charges.Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms.
26 9.7Formaldehyde (CH2O), a liquid with a disagreeable odor, traditionally has been used to preserve laboratory specimens. Draw the most likely Lewis structure for the compound.
27 A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure.O+-O+-
28 9.8Draw three resonance structures for the molecule nitrous oxide, N2O (the atomic arrangement is NNO).Indicate formal charges.Rank the structures in their relative importance to the overall properties of the molecule.
29 Exceptions to the Octet Rule The Incomplete OctetBe – 2e-2H – 2x1e-4e-BeH2HBeB – 3e-3F – 3x7e-24e-3 single bonds (3x2) = 69 lone pairs (9x2) = 18Total = 24FBBF3
30 Exceptions to the Octet Rule Odd-Electron MoleculesN – 5e-O – 6e-11e-NONOThe Expanded Octet (central atom with principal quantum number n > 2)SFS – 6e-6F – 42e-48e-6 single bonds (6x2) = 1218 lone pairs (18x2) = 36Total = 48SF6
31 AlI3 has a tendency to dimerize or form two units as Al2I6. 9.9Draw the Lewis structure for aluminum triiodide (AlI3).AlI3 has a tendency to dimerize or form two units as Al2I6.
32 PF5 is a reactive gaseous compound. 9.10Draw the Lewis structure for phosphorus pentafluoride (PF5), in which all five F atoms are bonded to the central P atom.PF5 is a reactive gaseous compound.
33 9.11Draw a Lewis structure for the sulfate ion in which all four O atoms are bonded to the central S atom.
34 9.12Draw a Lewis structure of the noble gas compound xenon tetrafl uoride (XeF4) in which all F atoms are bonded to the central Xe atom.
35 Single bond < Double bond < Triple bond The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond enthalpy.Bond EnthalpyH2 (g)H (g)+DH0 = kJCl2 (g)Cl (g)+DH0 = kJHCl (g)H (g)+Cl (g)DH0 = kJO2 (g)O (g)+DH0 = kJON2 (g)N (g)+DH0 = kJNBond EnthalpiesSingle bond < Double bond < Triple bond
36 Average bond enthalpy in polyatomic molecules H2O (g)H (g)+OH (g)DH0 = 502 kJOH (g)H (g)+O (g)DH0 = 427 kJAverage OH bond enthalpy =2= 464 kJ
37 Bond Enthalpies (BE) and Enthalpy changes in reactions Imagine reaction proceeding by breaking all bonds in the reactants and then using the gaseous atoms to form all the bonds in the products.DH0 = total energy input – total energy released= SBE(reactants) – SBE(products)exothermicendothermic