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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 Individual molecules Metallic Metal and metal “electron sea” Infinite lattice. Ionic Bonds: Octet Rule.

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chapter 9 chemical bonds
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
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

O

Cl

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

e.g. use Lewis symbols to illustrate the formation of a compound of

sodium and sulfur

Na2S Na+ combined with S2–

lattice energy
Lattice Energy

Ionic Bonding;

electrostatic attraction of positive (cation) and negative (anion) ions

Neutral atoms cation + anion

Ionic compound

Lattice energy energy 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)DH = –787 kJ

Bigger ions  smaller lattice energy

Higher charge  larger lattice energy

e– transfer

(IE + EA)

(lattice energy)

born haber cycles
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, DH°lattice = DH°f – (DH°step1 + DH°step2 + DH°step3 + DH°step4)

covalent bonding
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.
multiple bonds
Multiple Bonds
  • “double” and “triple” bonds

double bond sharing of 2 pairs of electrons between two atoms

triple bond sharing of 3 pairs of electrons between two atoms

Type of Bond single double triple

Bond order 1 2 3

Examples: O2 {O=O double bond}

N2 {N≡N triple bond}

CO2 {two C=O double bonds}

bond energy/bond strength

bond distance

electronegativity and bond polarity

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Electronegativity and Bond Polarity

electronegativity tendency 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)
lewis electron dot structures
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
formal charge
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

Put a circle around the formal charge.

All nonzero formal charges must be shown in a Lewis structure.

lewis electron dot structures1
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

and 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
resonance
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)

All reasonable resonance structures must be shown in a Lewis structure.

lewis dot structure checklist
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
Exceptions to the Octet Rule
  • Odd-Electron Species
    • “radicals” or “free radicals”
  • Incomplete Octets
    • Group 3 elements often have 6 electrons total, e.g. AlCl3
    • Often form bonds to complete the octets
  • Expanded octets
    • Only 3rd row or higher
sample problems
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 N3–

sample problems1
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 N3–

bond energies and estimating d h
Bond Energies and Estimating DH
  • Bond Energy—energy required to break 1 mole of the bond in the gas phase, e.g.
  • HCl(g) H(g) + Cl(g)DH = 431 kJ/mol
  • CH4(g)  CH3(g) + H(g)DH = 438 kJ/mol
  • Endothermic!
  • Estimating DHrxn using bond energies;

DHrxn = ∑(DH bonds broken) – ∑(DH bonds formed)

sample problem
Sample Problem

Ethanol is a possible fuel. Use average bond energies to calculate DHrxn for the combustion of gaseous ethanol.

sample problem1
Sample Problem

Ethanol is a possible fuel. Use average bond energies to calculate DHrxn for the combustion of gaseous ethanol.

CH3CH2OH(g) + 3 O2(g) 2 CO2(g) + 3 H2O(g)

DHrxn = -1245 kJ/mol

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