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Chapter 9: Chemical Bonds Types of Bonds Ionic –Metal and nonmetal –Electron transfer –Infinite lattice Covalent –Nonmetal and nonmetal –Shared electrons.

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Presentation on theme: "Chapter 9: Chemical Bonds Types of Bonds Ionic –Metal and nonmetal –Electron transfer –Infinite lattice Covalent –Nonmetal and nonmetal –Shared electrons."— Presentation transcript:

1 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

2 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 (Mg 2+, Ca 2+, etc) Al (group III) --> Al 3+ Group VII (17) --> -1 anions (F –, Cl –, Br –, etc) Group VI (16) --> -2 anions (O 2–, S 2–, etc) Group V (15) --> -3 anions (N 3–, P 3– ) e.g. Na 2s 2 2p 6 3s 1 --> Na + 2s 2 2p 6 {~Ne} Cl 3s 2 3p 5 --> Cl – 3s 2 3p 6 {~Ar} Remember Ch. 2 Figure 2.14 (p62)

3 Ionic Bonds: Lewis Symbols simple notation for showing number of valence electrons ClO group VII (7 valence e – )group VI (6 valence e – ) 3s 2 3p 5 2s 2 2p 4 O Cl e.g. use Lewis symbols to illustrate the formation of a compound of sodium and sulfur Na 2 S Na + combined with S 2–

4 Lattice Energy Ionic Bonding; electrostatic attraction of positive (cation) and negative (anion) ions Neutral atoms cation + anion Ionic compound Lattice energyenergy released when gaseous ions combine 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 (IE + EA) e – transfer (lattice energy)

5 Born-Haber Cycles Lattice energy can be calculated using a Born-Haber cycle; a hypothetical series of steps describing the formation of an ionic compound from the elements. Here,  H° lattice =  H° f – (  H° step1 +  H° step2 +  H° step3 +  H° step4 )

6 Covalent Bonding 1. Covalent Bond Formation –results from sharing of one or more pairs of electrons between two atoms –Examples: 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.

7 Examples

8 Multiple Bonds “double” and “triple” bonds double bondsharing of 2 pairs of electrons between two atoms triple bondsharing of 3 pairs of electrons between two atoms Type of Bondsingle doubletriple Bond order Examples: O 2 {O=O double bond} N 2 {N ≡ N triple bond} CO 2 {two C=O double bonds} bond energy/bond strength bond distance

9 Electronegativity and Bond Polarity electronegativitytendency of an atom in a molecule to attract electrons to itself e.g. Cl is more electronegative than H, so there is partial charge separation in the H-Cl bond: 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 3-D shape (Later) Electronegativity Increases Periodic Table

10 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 but only after the octets of any 2nd row elements are completed

11 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!) NH 4 + CO All nonzero formal charges must be shown in a Lewis structure. Put a circle around the formal charge.

12 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 but only after the octets of any 2nd row elements are completed

13 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 HCO 2 – 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) All reasonable resonance structures must be shown in a Lewis structure.

14 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)

15 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 3 rd row or higher

16 Sample Problems Write Lewis Electron Dot Structures (including formal charges and/or resonance as needed) for the following compounds and ions. PF 3 HCNSF 5 – NO 2 – SOCl 2 O 3 HNO 3 H 2 CON 3 –

17 Sample Problems Write Lewis Electron Dot Structures (including formal charges and/or resonance as needed) for the following compounds and ions. PF 3 HCNSF 5 – NO 2 – SOCl 2 O 3 HNO 3 H 2 CON 3 –

18 Bond Energies and Estimating  H Bond Energy—energy required to break 1 mole of the bond in the gas phase, e.g. HCl (g)  H (g) + Cl (g)  H = 431 kJ/mol CH 4(g)  CH 3(g) + H (g)  H = 438 kJ/mol Endothermic! Estimating  H rxn using bond energies;  H rxn = ∑(  H bonds broken) – ∑(  H bonds formed)

19 Sample Problem Ethanol is a possible fuel. Use average bond energies to calculate  H rxn for the combustion of gaseous ethanol.

20 Sample Problem Ethanol is a possible fuel. Use average bond energies to calculate  H rxn for the combustion of gaseous ethanol. CH 3 CH 2 OH (g) + 3 O 2(g)  2 CO 2(g) + 3 H 2 O (g)  H rxn = kJ/mol


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