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1

Orbitals

mathematical predictions of where electrons might be

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2

What orbitals will first-row orbitals have? (H, He)

S orbitals only

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3

What orbitals will second-row orbitals have? (B,C,N,O,F)

1s,2s,2p orbitals

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4

What is the octet rule?

Max of 8 electrons

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5

ionic bonding

electrostatic interaction of 2 ions (one positive and one negative)

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6

Cation

positive ion

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7

Anion

negative ion

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8

The most common ionic bonding is between

metal and nonmetal

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9

covalent bonding

Not all bonds are the same, and NOT all are equal

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10

formal charge equation

valence electrons minus [# of lone pairs + (1/2 # of shared electron bonds)]

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11

Molecular shape

angles will change with lone pairs and bondage

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12

Bond length

average length between 2 bonded atoms

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13

What does bond length do when atomic radius increases?

Bond length decreases with atomic radius

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14

Atomic radius decreases as you move

left to right across periodic table

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15

The atomic radius increases as you move

down the periodic table

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16

Bond angle

# of groups (lone pairs or bonded atoms) will determine geometry

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17

2 groups is what geometry

linear, 180

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18

3 groups is what geometry

trigonal planar, 120

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19

4 groups is what geometry

tetrahedral, 109.5

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20

Heteroatom

any atom other then C or H

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21

Hybridization

rearrangement of atomic orbitals into molecular orbitals

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22

2 groups and 2 hybrid orbitals means a hybridization of

sp

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23

3 groups and 3 hybrid orbitals means a hybridization of

sp^2

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24

4 groups and 4 hybrid orbitals means a hybridization of

sp^3

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25

TRUE or FALSE. Lone pairs count as a group when determining hybridzation.

True

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26

As bond length increases, bond strength…

decreases

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27

Shorter bond is a

stronger bond

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28

TRUE or FALSE. Triple bonds are the strongest.

True

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29

More s-character means a

stronger and shorter bond

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30

TRUE or FALSE. sp is the shortest and strongest bond.

true

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31

Polarity

unequal sharing of electron density

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32

High electron density likely to have

electron between 2 atoms

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33

An electronegative atom connected to an electropositive atom leads to

a polar molecule

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34

Nonpolar molecule

does NOT contain electronegative atoms

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35

Polar molecules have a net

dipole

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36

Resonance structure

2 or more Lewis Structures with some placement of atoms, but different arrangement of electrons (bonds and lone pairs)

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Curved arrow notation

move electrons in pairs in resonance structures

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38

sigma bond

one bond

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pi bond

double bond

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one sigma bond and 2 pi bonds

triple bond

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41

conjugation

occurs where p-orbitals (pi bonds) overlap on 3 or more adjacent carbons

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42

Special resonance allows for

delocalization (spreading out) of electron density

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43

aromaticity

special type of conjugation, whatever type of charge you start with is the charge you must end with

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44

The better resonance form has

lower energy

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45

Better resonance forms have more bonds and less charges which leads to more

stablity

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46

Hydrocarbons

made up of only hydrogens and carbons

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47

Aliphatic hydrocarbons

alkanes, sigma bond

alkenes, 1 sigma, and 1 pi bond

alkynes, 1 sigma and 2 pi bonds

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R is a substituent that means

any hydrocarbon

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1 degree

primary carbon, bound to 1 other carbon

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2 degree

secondary carbon, bound to 2 carbons

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3 degree

tertiary carbon, bound to 3 carbons

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52

4 degree

quaternary carbon, bound to 4 carbons

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53

Alkyl halides

R-X

X is F, Cl, Br, I

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54

Alcohols (hydroxyl groups or -OH)

R-OH

Ex: methanol

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55

Ethers (alkoxy groups)

R1-O-R2 (bent structure)

Ex: Dimethyl ether

H3C-O-CH3 (bent structure)

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Amines (amino group) 1, 2, or 3 amines

Ex: methyl amine

H3C-NH2

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Thiols (marcato groups)

R-SH

Ex: H3C-SH

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58

Sulfide (alkylthio group)

R1-S-R2

Ex: Dimethyl sulfide

H3C-S-CH3

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Aldehyde (end in “al”)

Ex: ethanal

Must have hydrogen

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60

Ketone (carbonyl groups)

Ex: acetone

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Ketone structure

can also be R1 and R2

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carboxylic acid (carboxy groups)

Ex: acetic acid

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carboxylic acid structure

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Esters

Ex: ethyl acetate

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Ester structure

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amide

Ex: methyl amide

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67

amide structure

The R2 and R3 can also be H’s

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68

acid chloride

Ex: acetyl chloride

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69

acid chloride structure

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70

resonance structures are better when they have

fewer charges and more bonds

full octets

negative charges on electronegative atoms

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71

Bronsted Lowry acid

donates H+ (proton) or H3O+

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72

Bronsted Lowry base

accepts H+ (lone pair electron or pi bond)

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73

Example of Bronsted Lowry acid

H-Cl

H-Br

H2SO4

H3O+

H2O

Acetic Acid

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74

Example of Bronsted Lowry Bases

H2O (4 lone pairs on Oxygen)

H3N (2 lone pairs on Nitrogen)

Hydroxide, negative charge

Methoxide, negative charge

Acetone

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75

Reaction of Bronsted Lowry acids and bases

H-A + B ⇆ A^- + H-B^+

acid base C.B C.A

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76

Another example of Bronsted Lowry acids and bases

H-Br + Methanol ⇆ Br^- + (Methanol + extra H)

Acid Base C.B C.A

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77

Lewis acid

accepts an electron pair (lone pair or pi bond)

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78

Lewis base

donates an electron pair

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79

TRUE or FALSE. Transition metals can accept a lone pair of electrons.

true. B, Al, Ga, etc.

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80

Lewis acid and base reaction example

H-A + water ⇆ A +H3O

Acid Base C.B C.A

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81

What is the K equilibrium formula

products divided by starting material

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82

Example of K equilibrium formula

[H30+][A-] divided by [H-A][H2O], water is liquid therefore not included so equation is

[H3O+][A-] divided by [HA]

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83

Ka is the

acidity constant

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84

A bigger Ka means

equilibrium lies towards the right (product-favored)

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85

A higher Ka means

a stronger HA (acid)

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86

pKa =

-log(Ka)

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87

smaller pKa means a

stronger acid

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88

Larger pKa means a

weaker acid

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89

Strong acid wants to

donate a proton

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90

A strong acid makes a

weak conjugate base

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91

A strong base wants to

accept a H+ (Become B-H)

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92

A strong base makes a

weak conjugate acid

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93

TRUE or FALSE. Equilibrium favors the side with the weaker acids/bases because they don’t react well.

true

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94

Equilibrium favors the

higher pKa value (weaker acid)

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95

Acidity increases from

left to right across periodic table

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96

The strongest conjugate base is the least

stable anion

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97

A larger ionic radius allows us to better distribute the

negative charge over the surface area

distribute charge = more stable anion

increase acidity

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98

The most negative pKa value has the

biggest ionic radius

H-F < H-Cl < H-Br < H-I

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99

Inductive effect

presence of a more electronegative atom has an electron-withdrawing effect

electron density being towards pulled the more electronegative atom

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100

How do you determine acid strength?

element effects

inductive effects

resonance effects

hybridization effects

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