Lecture 17 Flashcards

1
Q

Electronegativity

A

A blue describing how well an atom can attract electrons in a molecule
- derived directly from bond energies

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2
Q

Electron affinity

A

Amount of energy released when an electron is added to an atom

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3
Q

Pauling electronegativity scale

A

A numerical scale of electronegativies based on bond energy calculations for different elements joined by covalent bonds

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4
Q

Why are smaller elements more electronegative?

A

Positively charged nuclei are closer to electron density

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5
Q

Ionic bond

A

No sharing e-
Large electronegativity difference
ex. NaCl, KBr

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6
Q

Polar covalent

A

Unequally sharing e-
Electronegativity difference in the middle
ex. HCl, HI

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7
Q

No polar covalent

A

Equal sharing e-
Small electronegativity difference
ex. H2, Cl2

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8
Q

ΔEN values for mostly ionic compounds

A

1.7-3.3

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9
Q

ΔEN values for polar covalent compounds

A

0.4-1.7

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10
Q

ΔEN for mostly covalent (no polar) compounds

A

0-0.4

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11
Q

Bo bonds need to be only ionic or only covalent?

A

No there is a continuum

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12
Q

Electronegativity

A

The ability of an atom in a covalent bond to attach the shared electron pair

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13
Q

Metallic bonding

A

Metal is held together by the attraction between metal cations and the sea of valence electrons

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14
Q

Group 2 had higher ________ than group 1

A

Melting points

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15
Q

Why do metals bend instead of crack?

A

Metal ions slide past each other through electron sea and end up in a new position

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16
Q

Properties of metals

A
  • moderate to high melting points
  • much higher boiling points
  • good conductors of electricity (in solid or liquid state)
  • good conductors if heat
17
Q

The melting points of metals are only moderately high because…

A

The melting process does not break metallic bonds