Brainscape's Knowledge GenomeTM

Browse over 1 million classes created by top students, professors, publishers, and experts.


See full index

AP Chem > Chapter 8 and 9.1- Bonding > Flashcards

Chapter 8 and 9.1- Bonding Flashcards

1
Q

3 different terms effectively meaning the same thing…
Define:
1. Dipole
2. Dipole Moment
3. Polar Bond

A
  1. Something with 2 poles
  2. A polar molecule (has a positive end and a negative end)
  3. Polar bond (bond with unequal electron sharing)
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
2
Q

Formal charge and which formal charges are best

A

Group number - (dots + lines)
Best structure = closest to formal charge of 0 on ALL ATOMS, and if tied best = most negative charge on the most electronegative element

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
3
Q

Lewis structures:
1. Which element makes up the center?
2. Where do hydrogens go in a compound?

A
  1. Least electronegative ones, carbon if it is present
  2. If the center has been filled, the most electronegative ions
How well did you know this?
1
Not at all
2
3
4
5
Perfectly
4
Q

Free radical

A

element with an unpaired lone electron

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
5
Q

When determining shapes using VSEPR Theory single, double, and triple bonds all count as…

A

single bonds

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
6
Q

How are dipole moments drawn? (individual AND overall)

A

Overall: Arrows with perpendicular lines at the bottom pointing towards the negative end
INdividual: For EACH polar bond use that weird little squiggle guy for positive and negative parts

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
7
Q

Heat of reaction of bonds

A

Sum of bonds broken - sum of bonds formed (in kj/mol)

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
8
Q

Breaking bonds is always…
Forming bonds is always…
(exothermic/endothermic)

A

Breaking bonds is endothermic
Forming is exothermic

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
9
Q

Bond order

A

Number of lines/number of locations

Tells us if lines are single double or triple, and therefore the energy length and energy. Higher values mean more attraction and shorter lengths, whereas the opposite is true for lower values.

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
10
Q

When is a bond ionic?

A

metal-nonmetal OR polyatomic ions

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
11
Q

Lone electron pairs repel other domains (more/less) than atoms (VSEPR)

A

more

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
12
Q

What are the electronegativity differences for nonpolar, polar, and ionic bonds respectively? (ex. Fluorine is 1.9 units more electronegative than Oxygen)

A

Nonpolar: 0.3 or below
Polar: 0.3-1.7
Ionic: 1.7-3.3

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
13
Q

What are the most electronegative elements?

A

Cl, N, O, F (in increasing order)

Charge (almost) always creates a bigger difference than energy level

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
14
Q

Difference between ionic, covalent, and metallic bonds?

A

Ionic: metal-nonmetal, moreso “stealing” of electrons rather than sharing, held together via opposite charge creating electrostatic attraction

Covalent: nonmetal-nonmetal, sharing of electrons to create stability

Metallic: like an electron “soup,” electrons don’t stay with any particular atom, using the force of attraction between the delocalized electrons and cations to hold the bond together

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
15
Q

As covalent bond interatomic distance decreases energy (increases/decreases)

A

increases until a point of maximum stability and lowest energy

Pulling the atoms too close will result in the force of repulsion between nuclei to overcome interatomic electron-nucleus attraction

How well did you know this?
1
Not at all
2
3
4
5
Perfectly
16
Q

What is the purpose of resonance structures?

A
  • To provide stability by sharing the uneven formal charge between atoms
  • The stability of a structure (closeness to 0 on everything) determines the time spent in the structure
17
Q

What are the exceptions to the octet rules?

A
  • 3rd period and below NONMETALS can use the d orbitals and have up to 6 bonds
  • Boron can have 3 bonds
  • Beryllium can have 2 bonds
  • Hydrogen can have 1 bond
18
Q

Define domain

A

area of an electron (lone pairs, single bonds, double bonds, and tripe bonds all count as one)

19
Q

Name the bonds that make up single, double, and triple bonds

A

single: sigma bond
Double: 1 sigma bond, 1 pi bond
triple: 1 sigma bond, 2 pi bonds

20
Q

Polar vs. nonpolar covalent bonds

A

Polar - Unequal sharing due to diff. electronegativity, creating separate geometric centers of average positive and negative charge
Nonpolar - Equal pull/electronegativity of molecules

21
Q

When asked about lattice energy, the related equation is…

A

Coulomb’s law (relating charge first, and if the same then distance to force of attraction/energy holding the lattice together) (multiply the absolute value of charges of the atoms (do not take into account the # of atoms))

22
Q

How might orbitals change as an atom goes from a free state to a component of a molecule?

A

Hybridization of orbitals due to shifting electrons to arrange them in the manner with the least energy when existing free vs. in a molecule

23
Q

What energy level is 2sp3 closest to? 2sp2? 2sp1?

A

2sp3 is closest to 2p, being 75% p character, with 2sp2 being 67% p character and 2sp being in the middle with 50% p and 50% s character

24
Q

Which orbitals are atomic? Which are molecular?

A

Standard orbitals are atomic whereas hybrid orbitals (ex. sp3) are molecular

25
Q

When do Lewis diagrams/atoms have a lone pair before a single electron in each valence orbital?

A

In the case of the carbon family hybridizing its electrons in sp3, promoting all electrons into equal degenerate orbitals

26
Q

Breaking bonds are (exothermic/endothermic)

A

endothermic

27
Q

In determining bond length for covalent vs. ionic bonds…

A

Covalent: Prioritize whether bonds are single, double, or triple, then radius
Ionic: Ionic “bond lengths” should be determined by charge then radius (charge overcomes distance)

28
Q

Define covalent networks

A

Covalent molecules in a network (think crystal lattice w/o the whole + or - charges)

Most often made by Carbon or Silicon due to ability to bond to 4 molecules

29
Q

Is C-H polar?

A

nO

30
Q

Can halogens form double bonds?

A

non

31
Q

What is a sigma bond? a pi bond?

A
  • A sigma bond is a direct sharing of electrons via overlapping orbitals
  • A pi bond is a pair of parallel p orbitals that do not overlap and are perpendicular to the sigma bonds (attraction above AND below a sigma is 1 pi bond)
32
Q

Formula for determining net bond energy

A

Sum of REACTANTS - sum of products

33
Q

Does electronegativity have an impact on bond length?

A

No

AP Chem (17 decks)

Brainscape helps you reach your goals faster, through stronger study habits.
© 2024 Bold Learning Solutions. Terms and Conditions