Chapter 1: Fundamental Concepts - personal.tcu.edu

Chapter 1: Fundamental Concepts - personal.tcu.edu

Chapter 9: Chemical Bonds Types of Bonds Ionic Metal and nonmetal Electron transfer Infinite lattice Covalent Nonmetal and nonmetal Shared electrons Individual molecules Metallic Metal and metal electron sea Infinite lattice Ionic Bonds: Octet Rule In forming ionic compounds, atoms tend to gain or lose electrons in order to achieve a stable valence shell electron configuration of 8 electrons. Group I metals --> +1 cations (Li+, Na+, etc) Group II metals --> +2 cations (Mg2+, Ca2+, etc) Al (group III) --> Al3+ Group VII (17) --> -1 anions (F, Cl, Br, etc) Group VI (16) --> -2 anions (O2, S2, etc)

Group V (15) --> -3 anions (N3, P3) e.g. Na 2s22p63s1 --> Na+ 2s22p6 {~Ne} Cl 3s23p5 --> Cl 3s23p6 {~Ar} Remember Ch. 2 Figure 2.14 (p62) Ionic Bonds: Lewis Symbols simple notation for showing number of valence electrons Cl O group VII (7 valence e) group VI (6 valence e) 3s23p5 2s22p4 Cl O

e.g. use Lewis symbols to illustrate the formation of a compound of sodium and sulfur Na2S Na+ combined with S2 Na + S + Na 2 Na + S 2- Lattice Energy Ionic Bonding; electrostatic attraction of positive (cation) and negative (anion) ions e transfer

Neutral atoms cation + anion (IE + EA) Ionic compound Lattice energy combine (lattice energy) energy released when gaseous ions to form crystalline solid (an ionic compound) e.g. lattice energy of NaCl is 787 kJ/mol: Na+(g) + Cl(g) --> NaCl(s) H = 787 kJ Bigger ions smaller lattice energy Higher charge larger lattice energy Born-Haber Cycles Lattice energy can be calculated using a BornHaber cycle; a hypothetical series of steps describing the formation of an ionic compound

from the elements. Here, Hlattice = Hf (Hstep1 + Hstep2 + Hstep3 + Hstep4) Covalent Bonding 1. Covalent Bond Formation results from sharing of one or more pairs of electrons between two atoms Examples: H + H H H H + F H F

or or H H H F 2. Octet Rule -- for covalent bonding In forming covalent bonds, atoms tend to share sufficient electrons so as to achieve a stable outer shell of 8 electrons around both atoms in the bond. Examples H C + 4 H H C H or H

N + 3 H H + 2 H H H C H H N H H O H O H unshared e- pairs or "lone pairs"

Multiple Bonds double and triple bonds double bond atoms triple bond atoms sharing of 2 pairs of electrons between two sharing of 3 pairs of electrons between two bond energy/bond strength Type of Bond Bond order single double bond distance 1 2 Examples: O2 {O=O double bond} N2 {NN triple bond}

CO2 {two C=O double bonds} triple 3 Electronegativity and Bond Polarity electronegativity attract tendency of an atom in a molecule to electrons to itself Electronegativity Increases Periodic Table e.g. Cl is more electronegative than H, so there is partial charge separation in the H-Cl bond: + H Cl or

H Cl The H-Cl bond is described as polar and is said to have a dipole The entire HCl molecule is also polar as a result More complex molecules can be polar or nopolar, depending on their 3D shape (Later) Lewis Electron Dot Structures General Procedure -- stepwise process --Write the skeletal structure (which atoms are bonded?) --H atoms terminal --More electronegative atoms terminal (e.g. halogens) --Count all valence electrons (in pairs) and charges --Place 2 electrons (1 pair) in each bond --Complete the octets of the terminal atoms --Use multiple bonds if needed to complete the octet of the central atom, or --Put any remaining electron pairs on the central atom --Show formal charges Apply the OCTET RULE as follows: --H never has more than 2 electrons (I.e. one bond) --2nd row elements (e.g. C, N, O) almost always have an octet and never have more than 8 electrons (sometimes boron has only 6) --3rd row and higher elements can have more than 8 electrons

Formal Charge The apparent charge on an atom in a covalent bond = (# of valence e in the isolated atom) - (# of bonds to the atom) - (# of unshared electrons on the atom) minimize formal charges whenever possible (but the octet rule takes priority!) NH4+ CO + H H N C O H H Put a circle around the formal charge. All nonzero formal charges must be shown in a Lewis structure.

Lewis Electron Dot Structures General Procedure -- stepwise process --Write the skeletal structure (which atoms are bonded?) --H atoms terminal --More electronegative atoms terminal (e.g. halogens) --Count all valence electrons (in pairs) and charges --Place 2 electrons (1 pair) in each bond --Complete the octets of the terminal atoms --Use multiple bonds if needed to complete the octet of the central atom, or --Put any remaining electron pairs on the central atom --Show formal chargesand resonance forms as needed Apply the OCTET RULE as follows: --H never has more than 2 electrons (I.e. one bond) --2nd row elements (e.g. C, N, O) almost always have an octet and never have more than 8 electrons (sometimes boron has only 6) --3rd row and higher elements can have more than 8 electrons Resonance When multiple bonds are present, a single Lewis structure may not adequately describe the compound or ion -- occurs whenever there is a choice of where to put a multiple bond. e.g. the HCO2 ion is a resonance hybrid of two contributing resonance structures

The C-O bond order is about 1.5 (average of single and double bonds) O H C O H O C O All reasonable resonance structures must be shown in a Lewis structure. Lewis Dot Structure Checklist Correct total # of valence electrons Correct connectivity

NO atoms with > octet (or duet) in group 2 or below All atoms have octets (or duets), if possible Lone electrons clearly shown Nonzero formal charges included (and circled) Note it should be 2+ not +2 If the charge is 1+ or 1- do not write the 1! Important resonance structures included Ions have brackets and overall charge (not circled) Exceptions to the Octet Rule Odd-Electron Species radicals or free radicals Incomplete Octets Group 3 elements often have 6 electrons total, e.g. AlCl 3 Often form bonds to complete the octets Expanded octets Only 3rd row or higher Sample Problems Write Lewis Electron Dot Structures (including formal charges and/or resonance as needed) for the following compounds and ions. PF3 HCN

SF5 NO2 SOCl2 O3 HNO3 H2CO N 3 Sample Problems Write Lewis Electron Dot Structures (including formal charges and/or resonance as needed) for the following compounds and ions. H PF3 C N HCN

SF5 N NO2 N F F P F F F SOCl2 N N F

F O O F S O O O3 HNO3 H2CO N 3 O O O

O O O O Cl S Cl N O H C O O O N N

N H N N N N N N N N 2- O O O H

2- H Bond Energies and Estimating H Bond Energyenergy required to break 1 mole of the bond in the gas phase, e.g. HCl(g) H(g) + Cl(g) H = 431 kJ/mol CH4(g) CH3(g) + H(g) H = 438 kJ/mol Endothermic! Estimating Hrxn using bond energies; Hrxn = (H bonds broken) (H bonds formed) Sample Problem Ethanol is a possible fuel. Use average bond energies to calculate Hrxn for the combustion of gaseous ethanol. Sample Problem Ethanol is a possible fuel. Use average bond energies to calculate Hrxn for the combustion of gaseous ethanol. CH3CH2OH(g) + 3 O2(g) 2 CO2(g) + 3 H2O(g) Hrxn = -1245 kJ/mol

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