Unit 5 bonding
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Unit 5: Bonding. A perfect topic for friendship week. Warm-up (get it done before the music ends). Write net the rxn sulfuric acid added to sol’n of barium chloride. solution of tin(II) chloride added to an acidified solution of potassium permanganate. combustion of methane

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Unit 5: Bonding

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Unit 5 bonding

Unit 5: Bonding

A perfect topic for friendship week


Warm up get it done before the music ends

Warm-up (get it done before the music ends)

  • Write net the rxn

    • sulfuric acid added to sol’n of barium chloride.

    • solution of tin(II) chloride added to an acidified solution of potassium permanganate.

    • combustion of methane

  • Write an electron affinity reaction for Chlorine

  • Write F.I.P. reaction for rubidium


A bond

A bond

  • What holds two atoms together? What forces are responsible?

  • A bond, at its simplest form, is an attraction between the electrons is one atom for the protons in another atom.

  • This attraction causes a lower energy state than the atoms w/o bonds thus E is ALWAYS released when bonds form

    • What does this mean for bond breaking?


Bond types

Bond types

  • Covalent (molecular & Network)

    • What can you say generally about the electroneg. of elements in these compounds?

  • Ionic

    • What can you say generally about the electroneg. of elements in these compounds?

  • Metallic

    • What can you say generally about the electroneg. of elements in these compounds?


Metallic bonds

Metallic Bonds

  • Electronegativity:

  • Properties:

    • Malleable

    • Conductive

    • Ductile


Ionic bonds

Ionic bonds

  • Electronegativity:

  • Properties:

    • Volatility:

    • Solubility:

    • Conductivity:

    • Element type (explain):

    • State:

    • Crystal shape:

Crystal Lattice


Covalent bond types

Covalent bond types

  • Molecular covalent (most common type in this class)

    • “Small” packets atoms sharing e-

      • Ex: water, carbon dioxide

  • Network Covalent (AKA Network solid)

    • Vast interconnected array of atoms sharing e-

      • Example: Silicon dioxide


Bonding in sio 2

Bonding in SiO2


Network solids

Network Solids

  • Electronegativity:

  • Properties:

    • Solubility:

    • Volatility:

    • Conductivity:

    • Element type:

    • State:

    • Melting point:

    • Crystal shape:


Network solids c

Network Solids & C

  • Allotrope: differently arranged forms of an element

  • Carbon: 3 forms

    • Graphite: hexagonal sheets

    • Diamond: tetrahedral lattice

    • Buckminsterfullerene: Think soccer ball (alternating hexagonal and pentagonal rings)


Graphite diamond

Graphite & Diamond


Buckminsterfullerene

Buckminsterfullerene

  • He was an architect:


Molecular covalent bonds

Molecular Covalent Bonds

  • Electronegativity:

  • Properties:

    • Volatility:

    • Solubility:

    • Conductivity:

    • Element type (explain):

    • State:

    • Crystal shape:


Considering strength of attractions

Considering Strength of Attractions

  • Ionic compounds

    • Coulomb’s Law

      • A qualitative consideration

    • Hess’ Law

      • Lattice energy

  • Molecular covalent compounds

    • Factors that affect bond strength

      • Qualitative considerations

    • Enthalpy rxn using bond E


Ionic compounds

Ionic compounds

  • The attraction between ions can be calculated using an equation called Coulomb’s Law

  • What are some factors that increase ion – ion attraction?

    • Ion charge and Ion size

  • E = 2.31 x 10-19J nm ([Q1 Q2]/r)

  • What charges lead to what values of E?


Answer this in your notes

Answer this in your notes:

  • Which has the highest m.p. (justify your answer)?

    • LiF or BeO

    • NaF or RbI


Ionic lattice energies

Ionic: Lattice Energies

  • Lattice energy: energy of formation of 1 mole of a solid crystalline ionic compound when ions in the gas phase combine.

  • Can be found two ways

    • ΔH lattice = k (Q1 x Q2) / d

      k = a constant for the compound

      Q1 = charge on cation

      Q2 = charge on anion

      d = distance between the centers of the atoms


Unit 5 bonding

  • Hess’s law is more commonly used

    Calculate the lattice energy, ΔELattice, of NaBr(s),

    Na+(g) + Br-(g) NaBr(s)ΔELattice = ?


Unit 5 bonding

  • Values needed: (write equations)

    • Energy to turn metal into a gas

    • Energy to turn nonmetal into a gas (if not already a gas) [it is assumed to be g here]

    • Energy to split the nonmetal into single atom if not already

    • Ionization energy for the metal

    • Electron affinity for the nonmetal (this is like an atoms electronegativity except measured differently so the unit is in kJ)

    • ΔHf of the ionic compound

      Let’s go back and see if you can solve the last problem (-751kJ????)


Unit 5 bonding

  • Hess’s law is more commonly used

    Calculate the lattice energy, ΔELattice, of NaBr(s),

    Na+(g) + Br-(g) NaBr(s)ΔELattice = ?

    given the following thermochemical equations.

    Na(s) Na(g)ΔHf = 107 kJ

    Na(g) Na+(g) + e-ΔIE = 496 kJ

    ½ Br2(g) Br(g) ΔHf = 112 kJ

    Br(g) + e-  Br-(g) ΔEA = -325 kJ

    Na(s) + ½ Br2(g) NaBr(s)ΔH = -361 kJ


Covalent bond energies and some other info on covalent compounds

Covalent Bond Energies (and some other info on covalent compounds)

  • Bond order

    • Single bonds are 1st order

    • Double bonds are 2nd order

    • Triple bonds are 3rd order

      In molecules with resonance the bond order is not as easy to determine consider O3

      Here bond order is determined by the number of shared pairs between the two atoms divided by the number of links (not bonds) between the atoms. So O3 actually has bonds that are 1.5 order.


Unit 5 bonding

  • Bond length

    • Always average values because they can vary

    • Follow predictable trends H bonds to halides for example H-F< H-Cl < H-Br < H-I Do you see a trend? Or C-C > C-N > C-O > C-F. Do you see another trend?

    • As bond order between two elements increases bond length decreases.

      C-O > C=O > C≡O


The connection

The Connection

  • Bond order (and bond length) effect bond strength (AKA bond energy)

  • Bond Energy (D)-enthalpy change for breaking a bond in a molecule with the reactants and products in the gas phase under standard conditions

  • D is always a positive value or always endothermic (forming bonds is always exothermic)

  • If the amount of energy evolved in forming bonds is greater than the total amount of energy required to break the bonds the overall reaction is exothermic

  • The enthalpy for a reaction can be estimated by the equation

    ΔHrxn = ∑ D (bonds broken) - ∑ D (bonds formed)


Average bond energies

Average Bond Energies


Enthalpy of reaction

Enthalpy of Reaction

  • Δ H = Σ bond energies of reactants – Σ bond energies of products

    Practice: using bond energies from page 365, determine the enthalpy of reaction for the combustion of hydrogen.

  • What bonds were broken (look at reactants)

  • What bonds were formed (look at products)


Enthalpy of reaction1

Enthalpy of Reaction

  • Δ H = Σ bond energies of reactants – Σ bond energies of products

    Practice: using bond energies from page 365, determine the enthalpy of reaction for the combustion of methane.

  • What bonds were broken (look at reactants)

  • What bonds were formed (look at products)


Unit 5 bonding

4 C – H bonds broken

2 O=O bonds broken

2 C=O bonds formed

4 H-O bonds formed

413 x 4 = 1652

498 x 2 = 996

Sum = 2648

Subtract

2 x 745 = 1490

4 x 463 = 1852

Sum = 3342

Total = -694 kJ/mol


In class problems

In class problems


Unit 5 bonding

HW

  • Web assign


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