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AP Chem > Unit 2 > Flashcards

Unit 2 Flashcards

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1
Q

Electronegativity

A

A measure of the ability of an atom to attract shared electrons

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2
Q

When does electronegativity increase

A

Electronegativity increases as you move across a period on the periodic table, from left to right. This is because
the atomic radius is decreasing while the number of protons (and effective nuclear charge) is increasing

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3
Q

Nonpolar covalent bond

A
  • equal sharing of electrons
  • electronegativity difference <0.5
  • no bond dipole moment
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4
Q

Polar covalent bond

A
  • unequal sharing of electrons
  • electronegativity difference 0.5-1.7
  • partial bond dipole moment
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5
Q

Ionic bond

A
  • transferred electrons
  • electronegativity difference >1.7
  • positive and negative bond dipole moment
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6
Q

Ionic bonds are stronger when. . .

A

Charges are larger and the ions are smaller ( this can be explained by electrostatic force/Coulomb’s law equation

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7
Q

Properties of ionic substances

A
  • Form crystals (lattice of positive and negative ions)
  • High melting and boiling points
  • Hard
  • Brittle
  • Conduct electricity when dissolved and when molten (melted)
  • Good insulators as a solid
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8
Q

How do metallic bonds attract

A

Due to multiple metallic cations being attracted to a delocalized sea of valence electrons

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9
Q

When is the IMF stronger for metallic bonds

A

The IMF is stronger when there are smaller metallic cations and when there are more valence electrons

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10
Q

Properties of metallic substances

A
  • Shiny (Luster)
  • Malleable and ductile
  • Conduct heat and electricity
  • Metallic oxides are basic and ionic
  • Lose electrons to form cations
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11
Q

Properties of Nonpolar and Polar Covalent Molecules

A

x Non-lustrous, various colors
x Brittle, hard or soft
x Poor conductors
x Nonmetallic oxides are acidic and covalent
x Form anions by gaining electrons

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12
Q

Polar covalent bond rules

A

a) The atom with the higher electronegativity will develop a partial negative charge relative to the other
atom in the bond.
b) In single bonds, greater differences in electronegativity lead to greater bond dipoles.
c) All polar bonds have some ionic character, and the difference between ionic and covalent bonding is
not distinct but rather a continuum.

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13
Q

Bond energy

A

Bond energy is the energy required when breaking a bond, or the energy released when a bond is formed

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14
Q

How is bond energy different

A
  • Larger atomic radii increase the bond length. Longer bond length decreases the bond energy
  • Increasing the bond order increases the bond energy, as there are more electrons involved and therefore greater
    coulombic attraction and the bond length has decreased
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15
Q

Lattice energy

A

Energy to separate ions in ionic compounds

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16
Q

How to calculate lattice energy

A

Larger charges = more attraction = more energy required to separate the ions.
Smaller radii = more attraction = more energy required to separate the ions.

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17
Q

Why does KBr (672 kJ/mol) have a higher lattice energy than KI (632 kJ/mol)?

A

Bromide ions (Br-) have a smaller ionic radius than iodide ions (I-). This results in a smaller distance
between the potassium ion and the bromide ion. The smaller distance increases the coulombic attraction
resulting in higher lattice energy.

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18
Q

How are metals composed

A

Metals are composed of cations that are embedded in delocalized sea of valence electrons. The number of valence electrons determines the amount of electrons in the delocalized sea of electrons.

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19
Q

Alloys

A

Mixtures of metals. There are two types interstitial and substitution. The difference between them is the size of the atoms that are being added to the metal

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20
Q

Attraction of alloys

A

When the charge on the cations and the
number of electrons increases the attractions are greater. Additionally, when the ionic radius decreases the attraction increases.

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21
Q

Interstitial

A

The atoms added to the metal are small and fit in
between the metal atoms in the existing holes
(interstices). (H, B, C and N are commonly added)

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22
Q

Substitutional

A

The atoms added to the metal have similar radii so
they replace the atoms in the lattice

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23
Q

Formal charge

A

Formal Charge = (# Valence Electrons) - (# lone electrons +# bonds connected)

24
Q

The most valid structure

A
  • Formal Charges as close to zero for all atoms. (Most preferable is “no formal charge.”)
  • Negative Formal Charges will be placed on the most electronegative atom
25
Q

How to represent resonance

A

Draw all the structures of the molecule, keeping the arrangement of the atoms the same
but changing the location of the electrons. Double arrows are drawn between the structures indicating that the
structures are superimposed to make a hybrid.

26
Q

Resonance structures

A

when drawing a Lewis Structure you might find that there is more than one arrangement of
bond/electrons that are equally valid; you might find that a double (or triple) bond could be placed in one of two or more places in a molecule or ion. Look at the NO21- ion shown; both images are valid so both must be included when drawing a Lewis structure

27
Q

VSPER model

A

The valence shell electron-pair repulsion (VSPER) theory predicts the geometries of molecules and polyatomic
ions

28
Q

Linear

A
  • 0 lone pairs
  • 2 lines
  • 180
29
Q

Trigonal planar

A
  • 0 lone pairs
  • 3 lines
  • 120
30
Q

Bent one lone pair

A
  • 2 lines
  • 1 lone pair
    -<120
31
Q

Tetrahedral

A
  • 4 lines
  • 1 front one back
  • 109.5
32
Q

Trigonal pyramidal

A
  • one lone pair
  • 3 lines
  • one front one back
  • 109.5
33
Q

Bent 2 lone pairs

A
  • 2 lone pairs one front one back
  • 2 lines
  • <109.5
34
Q

Trigonal bipyramidal

A
  • no lone pairs
  • 5 lines 1 front 1 back
35
Q

Seesaw

A
  • 1 lone pair
  • 4 lines 1 front 1 back
  • <90, <120
36
Q

T shape

A
  • 2 lone pairs 1 front 1 back
  • 3 lines
  • <90 between electrons
37
Q

Linear 3 lone pairs

A
  • 3 lone pairs 1 front 1 back
  • 2 lines
  • 180
38
Q

Octahedral

A
  • no lone pairs
  • 6 lines 2 front 2 back
  • 90
39
Q

Square pyramidal

A
  • 1 lone pair
  • 5 lines 2 front 2 back
  • <90
40
Q

Square planar

A
  • 2 lone pairs
  • 4 lines 2 front 2 back
  • 90
41
Q

T shape .2

A
  • 3 lone pairs 1 front 2 back
  • 3 lines one front
  • <90 between 2 electrons
42
Q

Linear 4 lone pairs

A
  • 4 lone pairs 2 front 2 back
  • 2 lines
  • 180
43
Q

How do lone pairs effect angles

A

Lone pairs of electrons repel more than bonds and tend to compress the angle between bonding atoms.

44
Q

How to find hybridization

A

the sum of the superscripts on the hybrid orbitals equals the number of electron clouds around the central atom

45
Q

Order of hybridization

A

s, sp, sp2, sp3, sp3d

46
Q

Single bond contains. . .

A

Contains one sigma bond

47
Q

Double bond contains. . .

A

One sigma one pi bond

48
Q

How are double and triple bonds formed

A

Formed from unhybridized p orbitals and are called pi bonds

49
Q

Triple bond contains. . .

A

One sigma two pi

50
Q

Rank bonds from highest energy

A

Triple bonds > double bonds > single bonds

51
Q

Bond order

A

Number of bonds between 2 atoms

52
Q

Bond order of resonance bonds

A

bonds that have resonance are experimentally determined to be the same length (and order of 2 and 1 would be 1.5 for both)

53
Q

Bond polarity described

A

When electrons are shared in a covalent bond, the shared electrons spend more time around the more electronegative element in the bond resulting in a polar molecule

54
Q

Partial charges

A

Polar covalent bonds gives the more electronegative element a partial negative charge and the less electronegative element a slightly positive charge

55
Q

Symmetrical shapes

A
  • linear, Trigonal planar, tetrahedral, Trigonal biqyramidal, square planar, and octahedral

AP Chem (1 decks)

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