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Unit 2: Bonding and Intermolecular Forces
Metals: A type of naturally found element that (properties):
● Mostly silver ○ NEVER WHITE ● has a high melting and boiling point ○ Strong attractive forces ● is lustrous, ductile, and malleable ○ b/c of flexible sea ● Can Conduct* electricity ● With a crystal lattice structure ○ Cation metal atoms “swimming” in a sea of delocalized electrons that were before valence electrons ● Particle Diagram = Gray (Grid because SOLID at room temp.) ○ Write in charges (in bubbles)
What is happening at the microscopic level when something melts? ● Particles gain enough energy to leave the rigid crystal structure and Overcome attractive forces of the electron sea. They can move around now!
What is happening at the microscopic level when something boils? ● Particles gain enough energy to move freely
Why can a metal conduct electricity? ● In order for a substance to conduct electricity, it must have charged particles that are able to move freely. In a metal, these charged particles are the electrons in the “sea.” charged particles are the electrons in the “sea.”
An ALLOY is a mixture of 2 or more metals:
● Interstitial Alloy ○ particles/atoms of the other element insert themselves in between the existing metal. ○ Occurs when the atom sizes are quite different ■ Different number of energy levels ● Substitutional Alloy ○ kicks out some of the orig. atoms and replaces them with atoms of the secondary metal ○ Occurs when the atom sizes are almost the same ■ The same number of energy levels ● In a Particle Diagram (probably won’t have to show charges and e-) ○ Nonmetals DON’T give electrons away
Ionic compounds:
● Ionic bonds form between nonmetals and metals by the attraction of oppositely charged ions ○ COMPLETE transfer of one or more electrons to the other atom ■ The atom with a higher electronegativity gets the e- ● Particles arrange themselves in a regular repeating pattern called a crystal ○ Like particles repel and opposites attract ○ Each particle is surrounded by oppositely charged particles ○ No Individual Molecules, one (formula) unit ● Particle Diagrams a Rectangle with connecting Dashed lines (show attraction) ○ Grid B/C solid ○ Pay attention to the size of the atoms! ■ Electron lost/gained ~ energy level count ○ Attractive forces with a dotted line ● The Stronger Ion-Ion attraction comparing 2 compounds is when the ions in the compound ○ have a greater magnitude in charge ○ Smaller distance from nuclei to nuclei ○ Coulomb’s law
Ion Properties
○ Brittle ■ Trying to shape it causes layers of the lattice line up with like charges and they repel ○ Usually white (except w/ transition metals)
Lewis Structures (contd.):
● RESONANCE ○ When there are 2 identical atoms but seem to have a different bond order, they indeed have the same bond order ■ That 2nd or 3rd bond is delocalized (kinda swarms the identical atom branches) ■ So the Bond order is for ex/ 1.5, 1.33, etc ○ The most stable Resonance structure is one with the lowest formal charges ■ Split the bonds into 1 electron per atom ■ Count the number of electrons around an atom and calculate the difference between the regular amount of valence electrons ■ If the charges are similar, the extra electrons should be going to the most electronegative atoms
Electron Domains:
● How many links does it have around (2-6) ○ Lone pair ○ Connection (no matter # of bonds) between two atoms
Hybridization:
● Orbitals blend to form identical bonds ○ Based on observation ● Up to 4 domains ● “Exponents” add to # of domains - 1 ○ 2 domains - sp ■ 2 p orbitals left > form double or triple bonds ○ 3 domains - sp^2 ■ 1 p orbital left > form double bonds ○ 4 domains - sp^3
Sigma and Pi bonds:
● Single bond: sigma ● Additional bonds: pi
Polarity: ⇸
● DOT STRUCTURES ● Do Dipoles Cancel? ○ Is there another EQUAL dipole 180° away? ■ If there is another dipole 180° away, but of a different electronegativity. It is still slightly polar ○ If it does: nonpolar ■ Or if the bond is between two identical atoms
Polarity (contd.):
● Depends on Bond Angles ● Draw Partial Charges ○ The more uneven electron sharing (difference in electronegativity) the greater the partial charge ● C-H ALWAYS nonpolar
*Ionic (overall charged) Covalent Compounds DON’T have dipoles/partial charges! Only for Neutral
Atoms*
IN IC Compounds: the extra electrons ALWAYS go to the more electronegative atom(s)!
And don’t forget to write the bracket charge!!
Intermolecular Forces:
● Determines if a substance is solid, liquid, or gas ● Determines higher boiling or melting points ● Dipole-Dipole Attraction ○ Super Strong: Hydrogen Bonding ■ O-H, F-H, N-H ○ *One molecule can have more than one dipole-dipole attraction ● London Dispersion Forces (LDF) ○ All molecules have this ○ a temporary dipole-dipole attraction when electrons in a molecule unevenly travel to one side ■ SUPER RANDOM ■ SUPER WEAK ○ LDF can be strong(er) when: ■ A large molecule has MANY electrons/electron cloud ■ “More polarizable” ● Ion-dipole attraction ○ Positive ion attracted to negative dipole ○ Usually stronger than dipole-dipole
Network Solid:
○ All molecules are covalently bonded to the other ○ The highest melting and boiling points ○ Si, C ○ Some conduct electricity in solid state
180° Linear
120° Trigonal Planar
120° Bent
N/A Linear
109.5° Tetrahedral
109.5° Trigonal Pyramid
109.5° Bent
(Rare) -
90° & 120° Trigonal Bipyramidal
180°, 90°, &120° Seesaw
90° T-Shaped
180° (Rare) Linear
(Rare) -
90° Octahedral
90° Square Pyramidal
90° Square Planar
90° T-Shaped
180° (Rare) Linear
N/A (Rare)
Example Problems:
Covalent~
● Point 1 takes out metal and network solid
○ Point 2 also takes out metal ○ “brittle”
● Point 3 and Point 4 takes out ionic substance
○ Point 4 b/c Ionic have REALLY high melting points
● Network solid
○ Point 1 rules out metal ~ Metals not white ○ Point 2 (and 3) rule out covalent and ionic
- Account for each of the
following observations about pairs of substances. In your answers, use appropriate principles of chemical bonding and/or intermolecular forces. In each part, your answer must include references to both substances. a. Even though NH 3 and CH 4 have similar molecular masses, NH 3 has a much higher normal boiling point (−33°C) than CH 4 (−164°C). b. At 25°C and 1.0 atm, ethane (C 2 H 6 ) is a gas and hexane (C 6 H 14 ) is a liquid. c. Si melts at a much higher temperature (1,410°C) than Cl 2 (−101°C). d. MgO melts at a much higher temperature (2,852°C) than NaF (993°C).
A) NH 3 has a much higher boiling point than CH 4 because it exhibits very strong dipole-dipole attractions, so strong that this type of dipole-dipole attraction is known as hydrogen bonding. CH 4 has a much weaker attraction between its molecules known as London dispersion forces. The difference in attractive forces accounts for the difference in boiling points.
B) Both C 2 H 6 and C 6 H 14 are nonpolar molecules that don’t exhibit any attractions, except for London dispersion forces. C 6 H 14 is a liquid because it has so many electrons that at times the LDFs in the substance are strong enough to hold the molecules closer together than in a gas. In C 2 H 6 there are not many electrons, at least comparatively, so its LDFs are quite weak and it can be a gas in normal conditions of 25°C and 1.0 atm.
A) F 2 and I 2 are both nonpolar and experience LDFs. Since I 2 is solid at normal conditions, its LDFs are stronger than F 2 ’s. I 2 ’s LDFs are stronger because it has a larger electron cloud making it more polarizable. B) Both NaF and CsCl are ions that experience ion-ion attractions. The ions in each compound also have the same magnitude of charge: Na has +1, F has -1 and Cs has +1, Cl has -1. The difference is that the distance from the nucleus of Na to the nucleus of F is shorter than the distance from the nucleus of Cs to the nucleus of Cl. Since NaF has a shorter distance, its attractive force between the ions is larger. The stronger attractive force makes it harder to make particles move farther away from each other, so NaF’s melting point is higher. C) What determines the shape of a molecule is the number of domains and the bond angle. ICl 4 -^ has 6 domains, 4 bonds and 2 lone pairs. If a molecule has 6 domains, each domain is 90° from the other. In this case the shape would look like a flat square, if you connect each atom on the outside. That’s why it’s square planar. BF 4 - has only 4 domains, all 4 bonds. When a molecule has 4 domains, each domain is 109.5° apart. If imagined in a 3-D way, it looks like a tripod. The official name for this “tripod” is tetrahedral.
A) K, potassium, is a metal and metals conduct electricity in solid forms because of their structure. The atoms “give” their valence electrons to a sea of electrons in which the atoms, now cations, are suspended and that is their means of bonding. In order to conduct electricity, a substance must have charged particles that can move around freely, which is the role the electrons play in this scenario. KNO 3 , however, is an ionic compound. Each ion in the compound, K+^ and NO 3 - , are strongly held together by ion-ion attractions. While there are charged particles, these particles can not move freely, thus as a solid KNO 3 can’t conduct electricity. C) CCl 4 and CBr 4 are both nonpolar and experience only LDFs. Since CBr 4 has a higher boiling point, it
must experience stronger LDFs. It experiences higher LDFs because it has a larger electron cloud making it
more polarizable. That’s why it would be harder to pull away the particles in CBr 4.
D) I 2 and Br 2 are both nonpolar and experience only LDFs. Since I 2 has a higher boiling point, it must
experience stronger LDFs. It experiences higher LDFs because it has a larger electron cloud making it more
polarizable. That’s why it would be harder to pull away the particles in I 2. The bond energies do not affect the
boiling point because boiling does not involve breaking bonds.
A) MgO has a higher melting point than MgCl 2. The charges of the ions in MgO are larger than the ones in MgCl 2 : Mg has -2 and O has -2 while Mg has +1 and each Cl has -1. The greater magnitude in charges in MgO creates a stronger ion-ion attraction between its ions. Additionally, chlorine has 3 energy levels while oxygen only has 2. The additional energy level creates a larger distance between the nucleus of chlorine and magnesium than the distance between the nucleus of oxygen and magnesium. Consequently, the additional distance creates an even weaker attractive force between the ions in MgCl 2. Therefore, MgO has a higher melting point than MgCl 2. B) MgCl 2 has a lower melting point than MgF 2. Even though the ions in both the compounds have the same charges, +2 & -2, the distance between the nuclei of the ions differs. Chlorine has 3 energy levels while fluorine has 2. The additional energy level creates a larger distance between the nucleus of chlorine and magnesium than the distance between the nucleus of fluorine and magnesium. Consequently, the additional distance creates a weaker attractive force between the ions in MgCl 2. Due to the weaker attractive force, the melting point of MgCl 2 is lower than the melting point of MgF 2. C) The bond length of Br 2 is higher than the bond length of F 2 because the distance from the nucleus of one Br atom to the other is larger than the distance from the nucleus of one F atom to the other. The distance is larger because Br has 2 more energy levels, which make it have a larger atomic radius. D) The bond length of F 2 is higher than the bond length of N 2 because the distance from the nucleus of one F atom to the other is larger than the distance from the nucleus of one N atom to the other. The distance is larger because the F-F bond is single, a bond order of 1, while the N-N bond is triple, a bond order of 3. The higher bond order means the shorter the bond length, so that’s why N 2 ’s bond length is larger than F 2 ’s.
