chapter 9 2012 · 6. formed between elements with very different electronegativity. 6. formed...

AnnouncementsChem 7 Final Exam Wednesday, Oct 10 1:30-3:30AMChapter 1-1270 or 75 multiple choice questions

Exam III (Chapter 7-10)Wednesday, October 3, 2012Time: 6:00PM - 7:30PMSEC A 214A and 215ASKIPPING THIS STUFF

1. Classical distinction between energy and matter (p. 217)2. Numerical problems involving the Rydberg equation (equations 7.3/4). 3. Spectral analysis in the laboratory (pp. 226-227)4. Numerical problems involving the Heisenberg uncertainty principle (p. 231)5. Trends among the transition elements (p. 261)6. Trends in electron affinity (pp. 265-266)7. Pseudo-noble gas configuration (p. 269)8. Lattice energy (pp. 283-285)9. IR spectroscopy (p. 292)10. Numerical problems involving electronegativity (p. 296)11. Electronegativity with oxidation number (p. 297)12. Section 11.3: MO theory and electron delocalization13. All sections in chapter 12 except 12.3 (types of intermolecular forces).

Cadmium

Bromine

Iodine

Metals

Non-Metals

Covalent compounds result from chemical reactions between non-metals.

An ionic compound (salt) results from the chemical reaction between a metal and non-metal.

Ionic Compounds Covalent Compounds

1. Solid: high melting and boiling points (>600˚C).

1. Gas, liquid or solid with low melting points (<300˚C)

2. Soluble in polar solvents (like water) 2. Insoluble in polar solvents

3. Insoluble in non-polar solvents (like organic solvents)

3. Soluble in organic solvents (not water soluble)

4. Molten compounds conduct electricity very well 4. Non-electrolytes as liquids

5. Aqueous solutions are electrolytes

5. No charged particles and (electrolytes) in water

6. Formed between elements with very different electronegativity.

6. Formed between elements of similar electronegativity.

Ionic and covalent compounds have different physical properties.

Valence electrons are the electrons in the outershell (highest value of n quantum number) of the atom that dictate chemistry & in chemical bonding.

1A 1ns1

2A 2ns2

3A 3ns2np1

4A 4ns2np2

5A 5ns2np3

6A 6ns2np4

7A 7ns2np5

Group # of valence e-e- configuration

Given by the Group Number for

Group A

• For B group elements, the valence electrons are in the highest value ns orbital and the (n-1)d orbitals.

Lewis dot structures are used to depict valence electrons and bonding between atoms.

• A chemical symbol represents the nucleus and all core e-.

• A single dot around the symbol represents one valence e-.

+1 +2 +3 -1-2-3

Lewis structures are used to depict how ionic bonds are formed by chemists.

••

••

Mg••

Cl• ••

••

Cl•••

••

••Cl••

••••

Mg2+ -2

Formation of sodium chloride

O•••

•••

••O••

••

••Ba

2+ 2-Ba• • +

Formation of magnesium chloride

Formation of barium oxide

noble gas electronic configuration

1. Using spdf electron configurations

2. Orbital box diagrams

3. Lewis electron-dot symbols

Li (1s22s1) + F (1s22s22p5) Li+ 1s2 + F- (1s22s22p6)

Li1s 2s 2p

F1s 2s 2p

+

Li+

1s 2s 2p

F-

1s 2s 2p+

.+ F: ::Li. Li+ +

.F: ::

.

Chemists use different ways to conceptualize and draw ionic bonding. Know both!

NON-METALSMetals: Charge of Cation = Group Number (1A, 2A, 3A)

Non-metals: Charge of Anion = Group Number - 8

Ni2+

The for main group or representative elements the group number in the periodic table can tell us the charge of the metal cation and the charge of the non-metal anion that will be formed in an ionic reaction.

Transition Metalsstarts with filling of d-orbitals

Use partial orbital diagrams and Lewis symbols to depict the formation of Na+ and O2- ions from the atoms, and determine the formula of the compound.

SOLUTION:

3s 3p

Na

3s 3p

2s 2p

O2s 2p

O2-

2 Na+

:Na

Na+ O

.

:

..

.

2Na+ + O 2-

::

::

Na

.

1. Covalent bonding results from sharing of one or more electrons between non-metals atom.

share electrons

Occurs generally between non-metal elements forming molecules or molecular compounds.

• Every covalent bond has a characteristic bond length that leads to maximum stability.

Energy Released When Bond Formed

The formation of a covalent bond releases energy and reduces the potential energy (i.e minimizes energy).

Atoms Far Apart

Atoms to Close Together

F F lone pairslone pairs

single covalent bond

F F

single covalent bond

F F+

7e- 7e-

F F

8e- 8e-

8 e- or an “octet” around each atom

Lewis Dot Structure

Lewis Skeletal Structure

We also use Lewis structures and the “octet rule” to show covalent bonds and bonding between atoms.

Note the outer electrons are not shown!

• •

OH H H O H H O H

Dash FormulaLewis Dot

C••

••

O ••• • O•

•

• •

• •CO O

•

••

•••

••

••

••

•CO O ••

•••

••

••

••CO O••

••

••

••

What is the Lewis dot structure and the dash formula for carbon dioxide?

What is the Lewis dot structure and dash formula for water, H2O and for carbon dioxide?

Bonding Pairs (double bond)

Non-bonding Pairs

• Bond Order is the number of bonds between two atoms.– Single bond, order = 1– Double bond, order = 2– Triple bond, order = 3

• Higher bond orders give rise to:– Shorter bond lengths between bonded atoms– Stronger bond energies (more energy needed to break)

H O HCO O••

••

••

••

Bond order is the number of covalent bonds between two atoms.

Higher bond orders give shorter bond lengths and require more energy to break a bond.

Bond LengthsTriple bond < Double Bond < Single Bond

Comparing Bond Length and Bond StrengthThe trends in atomic radius can be extended to the bond length in molecules in a family of compounds.

(a) S - F, S - Br, S - Cl

(b) C = O, C - O, C O

(c) C-Br, C-I, C-Cl

Using the periodic table, but not Tables 9.2 and 9.3, rank the bonds in each set in order of increasing bond length and bond strength:

Comparing Bond Length and Bond StrengthUsing the periodic table, but not Tables 9.2 and 9.3, rank the bonds in each set in order of decreasing bond length and bond strength:

(a) S - F, S - Br, S - Cl

(b) C = O, C - O, C O PLAN: (a) The bond order is one for all and sulfur is bonded to halogens;

bond length should increase and bond strength should decrease with increasing atomic radius. (b) The same two atoms are bonded but the bond order changes; bond length decreases as bond order increases while bond strength increases as bond order increases.

(a) Atomic size increases going down a group.

Bond length: S - Br > S - Cl > S - F

Bond strength: S - F > S - Cl > S - Br

(b) Using bond orders we get

Bond length: C - O > C = O > C O

Bond strength: C O > C = O > C - O

Bond length depends on the size of the bonded atoms.

Internuclear distance(bond length)




Covalent radius

72 pm

Covalent radius

114 pm

Covalent radius

133 pm

Covalent radius

100 pm

F2

Cl2

Br2

I2

72pm

100pm

114pm

133pm

VSEPRT explains the geometry of molecules but NOT how covalent bonds are formed with that geometry.

Molecular formula

Lewis structure

VSEPRT

GeometryHybrid orbitals

VSEPRT Valence BondTheoryVSEPRTLewis Structure

Write the Lewis dot and skeletals structure of nitrogen trifluoride (NF3).

Write the Lewis dot and skeletal structures of the carbonate ion (CO3

2-).

Write the Lewis dot and skeletal structures structure of the carbonate ion (BrO3

-).

Write the Lewis dot and skeletal structures structure of the carbonate ion HCN?

Write the Lewis structure of nitrogen trifluoride (NF3).

Step 1 – N is less electronegative than F --> N is central atom!

F N F

F

A = 5 + 21 = 26 valence electrons

Step 3 - Write structure with N central and three bonds and rest non-bonding octet electrons around the central atom.

Step 2 - Count valence electrons = A; Nitrogen = 5, Fluorine = 3 X 7 = 21

octet

octet

octet

octet

Step 4 - Write structure with N central and three bonds and rest non-bonding octet electrons.

]BrO3[–

Valence e- = 7 + 3(6) + 1 = 26

O Br O

O

Look at the formula sometimes it gives clues to the central atom

H C N

HCNValence e- = 1 + 4 + 5 = 10

Carbon is central atom, watch for hydrogen--1 bond

–

Write the Lewis structure of the carbonate ion (BrO3-).

Bond Enthalpies and ReactionsHess’s Law allows us to analyze the energetics of a reaction. This provides insight into the relative stability of products and reactants.

Exothermic-energy releasedEndothermic

The bond energy, BE is the positive amount of energy (enthalpy) required to break or make one mole of bonds in a gaseous covalent compound to form products in the gaseous state at constant T and P.

!H0 = +436.4 kJ/mol

!H0 = +242.7 kJ/mol

H2 (g) H (g) + H (g)

Cl2 (g) Cl (g)+ Cl (g)

HCl (g) H (g) + Cl (g) !H0 = +431.9 kJ/mol

Bond Dissociation Energy

Example

All chemical reactions give off a net heat of reaction that can be either (!Hrxn) absorbed or gained.

The source of heat is the breaking and making of chemical bonds!

H2 (g) + F2 (g) 2HF(g) !Hrxn = ?

Type of bonds broken

Number of bonds broken

Bond energy (kJ/mol)

Energy change (kJ)

H H 1 436.4 436.4F F 1 156.9 156.9Type of

bonds formedNumber of

bonds formedBond energy

(kJ/mol)Energy

change (kJ)H F 2 568.2 1136.4

!Hrxn = (436.4 + 156.9( – (2 x 568.2) = - 543.1 kJ

∆Hrxn = ΣBEreactant bonds broken − ΣBEproduct bonds formed

Calculating Bond Enthalpies In A ReactionCalculate the !Hrxn for the chlorination of methane to form chloromethane gas using bond energies.

CH4(g) + Cl2(g) " CH3Cl(g) + HCl(g) !Hrxn = ?1. Write the balanced chemical equation and set up table below.2. Write Lewis structures break all bonds in products and reform all bonds in products. 3. Use the simple equation below to calculate !Hrxn


Type of bonds formed

Number of bonds formed


Energy change (kJ)




Energy change (kJ)

+ +




Energy change (kJ)




Energy change (kJ)

41

11

339431 431

339H-ClC-ClC-H 3 414 1232

414243

C-HCl-Cl 243

1656

Calculate the !Hrxn for the chlorination of methane to form chloromethane gas using bond energies.

!Hrxn = [1656kJ/mol + 243 kJ/mol] - [1232kJ/mol + 339 kJ/mol + 431 kJ/mol]!Hrxn = -114 kJ/mol (exothermic)


CH4(g) + Cl2(g) " CH3Cl(g) + HCl(g) !Hrxn = ?

Calculate the standard enthalpy change ∆H in kJ for the hydrogenation of ethyne (acetylene) to ethane.

+ 2H2(g) H3C CH3(g)





Energy change (kJ)




Energy change (kJ)

Calculate (in kJ) the standard enthalpy change ∆H for the hydrogenation of ethyne (acetylene) to ethane.

+ 2H2(g) H3C CH3(g)

C-C 1 347 347C-H 6 413 2478

C-H

H-H

212

413839432

839826

864

!Hrxn = [826kJ + 839kJ + 864 kJ] - [347 kJ/mol + 2478 kJ/mol] = -296 kJ/mol




Energy change (kJ)




Energy change (kJ)

Calculating Enthalpy Changes from Bond Energies

Use Table 9.2 to calculate !H0rxn for the following reaction:

CH4(g) + 3Cl2(g) CHCl3(g) + 3HCl(g)

1. Write the balanced chemical equation2. Write Lewis structures to see bonds broken and bonds formed. 3. Use the given Bond energies in the problem and a table account for bonds broken/formed to calculate !Hrxn





Energy change (kJ)




Energy change (kJ)

Calculating Enthalpy Changes from Bond Energies

Use Table 9.2 to calculate !H0rxn for the following reaction:

CH4(g) + 3Cl2(g) CHCl3(g) + 3HCl(g)

Calculate the number of bonds broken and bond formed using the bond energies found in Table 9.2.

4C-H = 4 mol(413 kJ/mol) = 1652 kJ

3Cl-Cl = 3 mol(243 kJ/mol) = 729 kJ

#!Hobonds broken = 2381 kJ

3 C-Cl = 3 mol(-339 kJ/mol) = -1017 kJ1 C-H =1 mol(-413 kJ/mol) = -413 kJ

#!H0bonds formed = -2711 kJ

3H-Cl = 3 mol(-427 kJ/mol) = -1281 kJ

!Horeaction = #!Ho

bonds broken + #!Hobonds formed = 2381 kJ + (-2711 kJ)

Electronegativity is an element’s inherent ability to draw electrons to itself when chemically bonded to another atom in a molecule (relative to Li).

Notice F, O, N, Cl, Br, C are highly electronegative with F being the most electronegative.

Group I and II metals are the least electronegative.

Differences in electronegativity between bonding atoms blur the distinction between covalent, polar covalent and ionic bonding types.

Covalent Bonding Ionic Bonding

Polar Covalent Bonding

!+ !-

The delta’s are used toshow a partial charge on atoms.

F2 HF LiF

The arrow is usedto denote a polar covalent bond. Arrow points negative

F F H F Li F

For a polyatomic molecule we must consider the vector sum of polar bonds in the molecule to see if there is a net dipole moment.

DipoleMoment

No NetDipoleMoment

DipoleMoment

DipoleMoment

No Net Dipole

polar

3.0

2.0

0.0

χ = |ENX - ENM|

!

Electronegativity values can be used to judge the extent of ionic or covalent bond character in a chemical bond.

Mostly Ionic

Polar Covalent

Mostly Covalent

absolute value of electronegativity difference between2-bond atoms

χ = |ENX - ENM|

How ionic or covalent a bond is between two atoms can be judged by looking at the the absolute value of the difference in electronegativity.

Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H2S; and the NN bond in H2NNH2.

Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic

H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent

N – 3.0 N – 3.0 3.0 – 3.0 = 0 Covalent

Ionic compounds are typically solids of high melting and boiling points.

Across a given period compounds become more covalent with lower melting points than their ionic counterparts.

Melting point of compounds increase across a period as compounds become less ionic and more covalent.

Across Period 3 chloride compounds go from high melting solids to very low boiling point gases (as it gets more covalent)

chapter 9 2012 · 6. formed between elements with very different electronegativity. 6. formed...

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