Name_____________________



AP Chemistry 3: Chemical Bonding Name __________________________

A. Bonding (8.1 to 8.4)

1. why bond?

a. attached (bonded) state is a lower energy state

b. complete valence shell = lower energy state

1. metals lose valence electrons (+ ion)

2. nonmetals with 5-7 valence electrons gain electrons to complete valence shell (– ion)

3 nonmetals shares valence e- with nonmetals

a. 1-3 valence electrons share 1 for 1 ∴ doubling valence number

b. 4-7 valence electrons share to fill s and p orbitals (8 electrons = octet rule)

c. Lewis symbols

1. chemical symbol + dots for valence electrons

2. Na•, •Mg•, etc.

d. three major types of bonds

1. ionic bond: electrostatic attraction between cations and anions

2. covalent bond: shared electrons between non-metal atoms

3. metallic bond: metal atoms collectively share valence electrons

2. ionic bonding

a. metal and nonmetal: Na(s) + ½ Cl2(g) → NaCl(s)

b. cations and anions: Na+(aq) + Cl-(aq) → NaCl(s)

c. cation listed first in formula and name (number of ions is not indicated in name)

d. formula represents simplest ratio of ions to make a neutral compound—not a molecule

e. bond strength—energy needed to break the bond: E ∝ Q1Q2/r, where Q1 and Q2 are ion charges and r is distance between ions

3. covalent bonding

a. bonding atoms' orbitals overlap, which maximizes attraction between nuclei and bonding electrons

b. atoms share 2, 4 or 6 electrons

1. 2 (single), 4 (double), 6 (triple) bond

2. multiple bonds reduce bond distance

a. bond distance < sum of atomic radii

b. shorter bond distance = stronger bond

c. polar bond when electrons are not shared equally

1. electronegativity

a. measures atom's attraction for bonding electron pair (higher # = stronger)

b. relative scale where period 2 elements are 1.0 (Li) to 4.0 (F), with 0.5 intervals

1. noble gases are excluded

2. trend:

a. increase across period

b. decrease down groups

2. bond polarity

a. electronegativity difference between atoms result in uneven sharing of electrons ∴ partially positive side, δ+, and a partial negative side, δ–

b. notation

c. measured as dipole moment

3. bond strength increases with polarity

d. naming binary molecules

1. two types of nonmetals covalently bonded

2. δ+ atom is written first with element name

3. second element is given –ide ending

4. prefix used to indicate number of atoms

a. 1—mono, 2—di, 3—tri, 4—tetra, etc.

b. mono never used for first element

c. example: CO2 is carbon dioxide

5. common names: NH3 (ammonia), H2O2 (hydrogen peroxide) and H2O (water)

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

B. Lewis Structures (8.5 to 8.7)

1. shows the atoms in a molecule with their bonding and non-bonding electron pairs

a. bonding electrons (– single, = double, ≡ triple)

b. lone (non-bonding or unshared) electron pair (••)

[pic]

2. drawing Lewis structures with one central atom

|count the total number of valence electrons (subtract charge for ions) |

|CO2: 4 + 2(6) = 16 |

|IF2–: 7 + 2(7) + 1 = 22 |

|draw a skeleton structure |

|first element in formula is central, except H |

|single bonds to other atoms (max. 4) |

| |

|O–C–O [F–I–F]– |

|(ions are bracketed) |

|place electrons around each atom |

|8 total electrons |

|except H, Be and B or when total number of electrons is an odd number |

| |

|.. .. .. .. .. .. |

|:O – C – O: [:F – I – F:]– |

|.. .. .. .. .. .. |

| |

|count Lewis structure electrons (including bonding electrons) |

|if equal to valence electrons, stop |

|if valence e- < Lewis e-, add additional bonds to reduces # of electrons|

|by 2's |

|if valence e- > Lewis e-, add 2 or 4 electrons to central atom (3rd |

|period or higher); called expanded octet |

| |

|.. .. .. .... .. |

|O = C = O [:F – I – F:]– |

|.. .. .. .. .. |

|added bonds expanded octet |

3. when more than one Lewis structure is possible use formal charge to decide which is more likely

a. each atom is assigned its lone electrons plus half the bonding electrons

b. formal charge = valence e- – assigned e-

c. preferred structure

1. atoms have formal charges closest to zero

2. negative formal charge reside on the more electronegative atom (upper right most on the periodic table)

d. example: NCS-

| |[:::N–C≡S:]- | [::N=C=S::]- | [:N≡C–S:::]- |

|valence e- |5 4 6 |5 4 6 |5 4 6 |

|assigned e- |7 4 5 |6 4 6 |5 4 7 |

|formal |-2 0 +1 |-1 0 0 |0 0 -1 |

∴ [::N=C=S::]- is preferred because formal charges are closest to zero and negative charge is on the nitrogen (higher electronegativity)

e. technique can produce erroneous structures (experiments are required to determine actual structure)

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

C. VSEPR Model (9.1 to 9.3)

1. rules

a. maximum separation between electron pairs

b. atom positions define molecular geometry

c. lone electron pairs squeeze bond angle

(actual angle < ideal angle)

|Electron |Domain Geometry |– |: |Molecular Geometry |Bond |

|Domains | | | | |Angle |

|2 | |2 |0 | |180o |

| | | | | | |

|3 | | | | |120o |

| | | | | | |

| | |3 |0 | | |

| | | | | | |

| | | | | | |

| | | | | | |

| | |2 |1 | | |

|4 | | | | |109.5o |

| | | | | | |

| | |4 |0 | | |

| | | | | | |

| | | | | | |

| | | | | | |

| | |3 |1 | | |

| | | | | | |

| | | | | | |

| | | | | | |

| | | | | | |

| | |2 |2 | | |

|5 | | | | |90o |

| | | | | |120o |

| | |5 |0 | | |

| | | | | | |

| | | | | | |

| | | | | | |

| | |4 |1 | | |

| | | | | | |

| | | | | | |

| | | | | | |

| | | | | | |

| | |3 |2 | | |

| | | | | | |

| | | | | | |

| | | | | | |

| | |2 |3 | | |

|6 | | | | |90o |

| | | | | | |

| | |6 |0 | | |

| | | | | | |

| | | | | | |

| | | | | | |

| | |5 |1 | | |

| | | | | | |

| | | | | | |

| | | | | | |

| | |4 |2 | | |

| | | | | | |

2. polar molecules

a. lone electron pairs distort symmetry except for sp3d-linear and sp3d2-square planar

b. different perimeter atoms

c. polar interactions increase water solubility, increase melting and boiling temperatures, decrease evaporation (volatility)

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

D. Valence-Bond Theory (9.4 to 9.5)

1. explains electron domain geometries in terms of electron orbitals

2. atomic orbitals from bonding atoms merge, which allows single electrons from each atomic orbital to occupy overlapping area and simultaneously attract both nuclei (i.e. H–H: overlap of 1s orbitals)

3. more complex molecules require a fusion of s and p orbitals into equivalent (hybrid) orbitals

a. explains why covalent bonds around an atom are all the same even if electrons were originally in different shaped (s, p and d) atomic orbitals

b. 1 s + 1 p form 2 sp hybrids (2 electron domains)

ground state excited state hybridized state

|↑↓ | | | | |→ |↑ | |↑ | |

e. expanded octet hybridization

1. 1 s + 3 p + 1 d = 5 sp3d hybrids (5 domains)

2. 1 s + 3 p + 2 d = 6 sp3d2 hybrids (6 domains)

4. not all valence electrons enter hybrid orbitals

a. one electron pair per bond enters a hybrid orbital

1. sigma bond (σ)

2. electrons located between bonding atoms

b. lone pairs of electrons enter hybrid orbital

c. remaining bonding pairs of electrons from multiple bonds remain in pure p orbitals

1. pi bond (π)

2. electrons located above/below bonding atoms

d. example: ::O=C=O::

p p

sp2 p p sp2

sp2 sp sp sp2

p p sp2

sp2 p p

e. π bond electrons can spread out across entire molecule (delocalized)

1. NO3- has one π bond, which is shared evenly and simultaneously between 3 O's

2. multiple Lewis structures show all possible locations for π bonds = resonance forms

[pic]

3. bond order = sigma bond + share of π bonds

(each N–O bond has bond order = 1 1/3)

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

| |

E. Simple Organic Molecules—Hydrocarbons (25.1 to 25.6)

1. general properties

a. contain C and H

b. nonpolar, flammable (fuels)

2. formulas and names

a. number of carbons in parent chain

|1 |2 |3 |4 |5 |

|meth |eth |prop |but |pent |

|6 |7 |8 |9 |10 |

|hex |hept |oct |non |dec |

b. bond between carbons

1. alkanes (all single bonds) end in “ane”

2. alkenes (1 or more double bonds)

1 double end in “ene”, 2 double end in "diene"

3. alkynes (1 or more triple bonds) end in “yne”

4. cyclical

a. 3 to 6 carbon ring with single bonds between carbons: prefix "cyclo"

b. 6 carbon ring with 3 shared π bonds: benzene (called aromatic hydrocarbon)

c. branches

1. C-branches—“yl”

2. benzene branch—"phenyl"

3. location of branches

a. number of the parent carbon

b. lowest number possible

c. dash: # – word, comma: #, #

d. number of branches (2—di, 3—tri, etc.)

3. condensed structural formula

a. hydrogens are written after the carbon

b. branches are in parentheses after hydrogens

c. example: 4-ethyl-2-methyl-1-hexene

CH2C(CH3)CH2CH(C2H5)CH2CH3

d. semi-condensed (shows branches and bonds)

CH3 C2H5

| |

CH2=C–CH2–CH–CH2–CH3

4. functional groups

a. dramatically modify properties of hydrocarbon

b. haloalkanes: halogen replaces one or more H

1. reduces reactivity (flammability)

2. named as a branch with an “o” ending

c. oxygen containing groups

1. hydroxyl group (C–OH)

a. water soluble

b. alcohols (-ol ending)

c. acids (-anoic acid ending)

2. carbonyl group (C=O)

a. aldehydes (-al ending)

b. ketones (-one ending)

c. esters (-oate ending )

3. ethers have C–O–C (-yl -yl ether ending)

4. increases polarity: C–OH > C=O > C–O–C

d. amines

1. replace H in ammonia with hydrocarbon group = amine (CH3NH2 = methylamine)

2. when NH2 branches off hydrocarbon = amino

CH3CH(NH2)CH2CH3 (2-aminobutane)

3. weak bases (neutralize acids—absorb H+)

5. isomerism

a. structural isomers: same molecular formula, different structure and name

1. move double/triple bond position

2. move branch

3. form cycloalkane from alkene

b. geometric isomers: same molecular formula and structure, but different spatial arrangement around the >C=C ................
................

In order to avoid copyright disputes, this page is only a partial summary.

Google Online Preview   Download