Problem Set 11. 2: How Bonding affects Physical Properties



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  1. Problem Set 11.2: How Bonding affects Physical Properties

  1. Describe the type of bonding in sodium metal and explain why sodium is a good conductor of electricity

Metals are held together by the attraction of the positive ions for the sea of electrons

Because electrons are not localized (attached to a specific atom) they can move around relatively easily. This makes metals highly conductive

  1. Name and describe the bonding within and between the molecules in liquid fluorine

Flourine atom bond covalently to make the F2 molecule. The molecule is non-polar because the electrons are shared evenly between the two atoms. The intermolecular forces experienced by liquid fluorine are dispersion forces which are weak forces due to temporary induced dipole.

  1. Write the electron structures of both sodium and fluorine and describe how the atoms combine to form sodium fluoride

Na: 2.8.1

F: 2.8.7

Sodium’s outer electron transfers to fluorine to make oppositely charged ions with stable octets. The ions attract to make NaF

  1. Explain why sodium fluoride doesn’t conduct electricity until it is heated above its melting point.

Ions are not free to move in a solid but they can move in a molten (liquid) state

  1. List 3 kinds of intermolecular forces and rank them in order of increasing strength

Dispersion forces < dipole-dipole forces < hydrogen bonds

  1. What is a temporary induced dipole?

It is a dipole that forms because of random fluctuations of electrons in a molecule. If a bunch of electrons move to one side of a molecule, they induce a dipole on an atom nearby since it’s electrons will move away from the other molecule creating oppositely charged poles

  1. Identify the strongest type of intermolecular force in the following:

    1. CH3Cl Dipole-dipole

    2. CH4 Dispersion forces

    3. CH3OH Hydrogen bonding

  1. Rank the strength of the following compounds from strongest to weakest intermolecular forces. PCl3, NiCl2, I2, HF

I2 has dispersion forces since it is non-polar

NiCl2 will have ion-ion forces since it is most likely an ionic bond (presence of metal makes an ionic bond)

PCl3 (P-Cl bond has electronegativity difference of 0.9 and molecule is not symmetrical due to lone pair on P)

HF (H-F bond has electronegativity difference of 1.9) so it’s more polar than P-Cl and has the ability to hydrogen bond.


NiCl2 > HF > PCl3 > I2



  1. Which of the above 4 compounds are likely to dissolve in water?

NiCl2, HF and possibly PCl3 since ionic and polar compounds are able to dissolve in water

  1. Which of the above 4 compounds are likely to dissolve in a non-polar solvent?

I2

  1. In ethanol, C2H5OH(l),there are covalent bonds, hydrogen bonds and van der Waals. forces. Which bonds or forces are broken when ethanol is vaporized?

Van der Waals forces and hydrogen bonds are broken when ethanol evaporates. When a substance evaporates it changes from a liquid to a gas but it does not change chemical composition so only intermolecular forces are broken – not covalent bonds.

  1. True or false: Elements that covalently bond have very low electronegativity values

False. Just because the elements have low electronegativity values doesn’t mean they will covalently bond. The difference in electronegativity is what matters. For example, Fluorine has a very high electronegativity and makes covalent bonds when it bonds with itself. Sodium which has a very low electronegativity makes ionic or metallic bonds – not covalent bonds.

  1. State and explain the relative conductivity of the following substances

  1. Cu (s)

  2. Hg(l)

  3. H2 (g)

  4. LiOH(aq)

Relative conductivity:

Cu and Hg are both metals and therefore bond via metallic bonding which is a bunch of cations held together by an electron sea. Because the electrons are delocalized, the electrons can move easily and therefore, conduct. Hg’s conductivity is slightly lower than that of Cu because it doesn’t like to share electrons (which is why it’s a liquid at room temperature- doesn’t have the strong metallic bonds formed through delocalization). LiOH (aq) also has relatively high conductivity since ionic compounds separate into ions when they are dissolved in water. The presence of ions in water makes the water a good charge carrier. H2 is a terrible conductor because electrons are very localized between the hydrogen atoms.

  1. State and explain the relative solubility in water of the following substances

  1. CH3CH3 not soluble – non-polar

  2. CH3OCH3 not very soluble – has some ability to make dipole-dipole forces

  3. CH3CH2OH pretty soluble because of ability to hydrogen bond

  4. CH3CH2CH2CH2OH not soluble because of large non-polar carbon chain

  1. State and explain the relative boiling points for the following substances

Note that the following substances have about the same mass (about 60 amu)

  1. C4H10 non-polar so very low boiling point

  2. CH3CH2CH2OH capable of hydrogen bonding so has highest boiling point

  3. CH3OCH2CH3 primary intermolecular force is dipole dipole which is weaker than H-bonds for molecules of same mass. So b.p is between that of substances of a and b.




  1. Why is important that the molecules in the previous question have the same mass?

Because than you can compare intermolecular forces more easily. Molecules of lower mass may have lower boiling points because they are lighter and therefore need less energy to evaporate even though the forces between the molecules might be stronger than another molecule with a large mass.

  1. Which molecule has the most ionic character

  1. MgS

  2. HCl

  3. CO2

  4. CaO

CO2 isn’t even a contender since two non-metals make a covalent bond.

Hydrogen often forms covalent bonds with non-metals. It has a much higher electronegativity than other alkali metals and alkaline earth metals. So CaO or MgS must have the most ionic character. Remember that the most ionic character is determined by the difference in electronegativity so we need to identify the pair that have the greatest difference. The trend is that electroengativity increase across the periodic table and decreases down the periodic table. So Oxygen is more electronegative than Sulfur and Magnesium is more electronegative than Calcium so CaO will have a bigger difference and therefore the most ionic character.

  1. Which molecule has the greatest polarity?

  1. Fluorine obviously non-polar since difference in electroneg. is zero.

  2. Hydrogen fluoride most polarity (largest difference in Electronegativity and is not symmetrical)

  3. Hydrogen chloride less polar than HF since Cl has lower electronegativity than F because of shielding of outer electrons

  4. Tetrafluoromethane (CF4) non-polar molecule because it’s symmetrical.

  1. When hydrogen bonds with sulfur it makes a gas, when it bonds with oxygen, it makes a liquid. If sulfur and oxygen are in the same family, why are the properties of H2S and H2O so different?

Oxygen is much more electronegative than sulfur and therefore makes more polar bonds with hydrogen than sulfur does. The polar bonds cause strong intermolecular forces – hydrogen bonding that make water a liquid at room temp rather than a gas. Sulfur is less electronegative because it has one more energy level filled. The eight inner electrons of Sulfur’s second energy level shield the outer electrons a lot
Problem Set 11.1 Answers

  1. Determine which ion will form

    1. Na+1

    2. Mg+2

    3. Li+1

    4. Al+3

    5. B+3

    6. O-2

    7. N-3

    8. F-1

    9. Cl-1

    10. S-2

    11. P-3

    12. Br-1

  2. Explain why the ion formation by a transition metal cannot be predicted

Transition metals have more than one possible oxidation state

  1. Give the kind of bonding in the following comounds

    1. Ionic

    2. Ionic (between Na and SO42-) and covalent (between S and O)

    3. Ionic because between metal and non-metal. Some may classify as polar covalent (metals with high oxidation states form polar bonds when bonding with non-metals)

    4. Ionic

    5. Non-polar covalent (two non-metals with similar electronegativities)

    6. polar covalent (bond is polar b/c two non-metals with different electronegativities but molecule is non-polar)

    7. polar covalent

    8. polar covalent (however, molecule is non-polar)

    9. polar covalent (however, molecule is non-polar)

  2. valence of 6 gives charge of -2 and valence of 2 gets a charge of +2 so formula is Y2+X2-

  3. 4 valence electrons (needs 4 bonds or two double bonds) and 6 valence electrons (needs two bonds or one double bond) so must be QR2.

    1. Can’t be QR because R would have to donate four electrons to Q which would give it too many electrons

  4. Covalent bonds involve share electrons, ionic bonds involve a transfer of electrons

  5. bond order is the number of bonds (single = order of 1, double = order of 2). The higher the order, the stronger and shorter the bond.

  6. C-O If you have a data booklet you can look up electronegativity values and find the difference. An easier way, though is to see how far apart the elements are on the periodic table

    1. B-C EN = 0.5 (next to each other)

    2. C-O EN = 1.0 (two away from each other)

    3. N-O EN = 0.5 (next to each other)

    4. O-F EN = 0.5 (next to each other)



Molecule

Lewis

s = single bond

d = double

t = triple

lp = lone pairs


Shape

Bond Angle

Polar?

Resonance

F2

1 s

3 lp on each



Linear

180

No

No

O2

1 d

2 lp on each



Linear

180

No

No

N2

1 t

1 lp on each



Linear

180

No

No

H2O

2 s

2 lp on O



Bent

104.5

Yes

No

NH3

3 s

1 lp on N



Trigonal pyramidal

107

Yes

No

PH3

3 s

1 lp on P



Trigonal pyramidal

Predict 107 (but actually 94)

Yes

No

SiH4

4 s

0 lp


Tetrahedral

109.5

No

No

CO2

2 d

2 lp on each O



Linear

180

No

No

SCl2

2 s

3 lp on each Cl

2 lp on S


Bent

104.5

No

No

C2Cl2

1 d C=C

4 s C-Cl


3 lp on each Cl

Trigonal planar

120

No

No

SiCl4

4 s

3 lp on each Cl



Tetrahedral

109

No

No

SO2

2 d

1 lp on S

2 lp on each O


Bent

<120

Yes

No

C2H2

1 t

2 s C-H


Linear

180

No

No

C2H4

1 d C=C

4 s C-H



Trigonal planar

120

No

No

CO32-

1 d w/ 2 lp on O

2 s w/ 3 lp on O’s



Trigonal planar

120

No, but it is an ion so it has an overall charge

Yes

NO2-1

1 d w/ 2 lp on O

1 s w/ 3 lp on O



1 lp on N

Bent

<120

No, but it is an ion so it has an overall charge

Yes




  1. answered in table

  2. Boron can break the octet rule – it can have only 6 electrons around it

    1. Elements in the third row of the periodic table can also break the octet rule




  1. SiH4 = tetrahedral and PH3 = trigonal pyramidal. The bond angles are different even though they are both based on the tetrahedral electron arrangement because the PH3 molecule has a lone pair that repels bonding pairs more.

  2. Principles of VSEPR

    1. Valence electron pairs repel each other – lone pairs repel more than bonding pairs

    2. To determine the geometry, the number of bonding and lone pairs around the central atom are considered

    3. The location of the nuclei of each atom determine the shape

  3. Because it’s the location of the nuclei that determine the shape and double and triple bonds do not increase the number of nuclei

  4. polar bonds (difference in electronegativity) and symmetry of the molecule

  5. 4 lone pairs (2 on each oxygen) 8 shared pairs

  6. 120 degrees

  7. shapes:

    1. trigonal planar

    2. tetrahedral

    3. Trigonal pyramidal

    4. Bent

    5. bent



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