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Chemistry 20 IB > Unit 1: Bonding (Textbook) > Flashcards

Unit 1: Bonding (Textbook) Flashcards

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0
Q

D: Octet rule

A

Eight electrons filling the valence level = stable

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1
Q

D: Valence orbital

A

The volumes of space that can be occupied by electrons in an atom’s highest energy level.

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2
Q

D: Electronegativity

A

Relative ability of an atom to attract a pair of bonding electrons in its valence level.

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3
Q

D: Covalent bond

A
  • Forms between two nonmetals

- The attraction between two nuclei for a shared pair of bonding electrons.

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4
Q

D: Ionic bond

A

The attraction between a specific cation and anion.

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5
Q

Why do ionic bonds form a crystal lattice?

A

When enough ions assemble to form a visible amount of compound, we are able to see a repeating crystal shape.

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6
Q

D: Metallic bonding

A

When both types of colliding atoms have relatively low electronegativities, the atoms share electrons but no chemical reaction takes place.

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7
Q

5 Properties of a metallic bond

A 
Shiny
Malleable 
Conductor 
Crystal lattice 
Low e/n
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8
Q

In metallic bonds, why are valence electrons free to move between the atoms?

A

The valence electrons are not held strongly by their atoms

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9
Q

Describe the phrase “sea of mobile electrons”

A

Wherever the electrons move in a metallic bond, they hold the atoms together. The “glue” that holds the structure together

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10
Q

Why are metallic bonds malleable?

A

The atoms can be moved around each other without moving them further apart from each other so the bonds are not weakened or broken by the changing shape of the solid.

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11
Q

VSEPR Theory (5)

A
  1. Only the valence electrons on the central atom are important for molecular shape
  2. Valence electrons are paired
  3. Bonded/lone pairs of electrons are treated equally
  4. Valence electrons pairs repel
  5. Molecular shape is determined by the positions of electron pairs when they’re a max distance apart
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12
Q

D: Polar molecule

A

Negative charge is not distributed equally among the atoms = partial positive and negative charges on opposite sides of the molecule

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13
Q

D: Nonpolar molecule

A

A molecule with symmetric electron distribution

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14
Q

D: Nonpolar covalent bond

A

If the bonded atoms have the same electronegativity they will attract any shared electrons equally

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15
Q

D: Polar covalent bond

A

If atoms have different e/n

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16
Q

If x > 1.7

A

Ionic

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17
Q

If x < 1.7

A

Polar covalent

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18
Q

If x =0

A

Nonpolar covalent

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19
Q

D: Bond dipole

A

The charge separation that occurs when the e/n difference of two bonded atoms shifts the shared electrons making one end negative and the other positive

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20
Q

Explain “like dissolves like”

A

Nonpolar solute dissolves in nonpolar solvent. Vice versa.

21
Q

D: Intermolecular forces

A

The forces of attraction and repulsion between molecules

22
Q

What are “van der Waals” forces?

A

DD

LDF

23
Q

Properties of DD forces

A
  • Attraction between dipoles
  • Weakest IMF
  • Only for polar molecules
24
Q

LDF

A

Electron movement within a molecule produces a momentary dipole

25
Q

D: Isoelectronic molecules

A

Molecules with the same number of electrons

26
Q

D: Hydrogen bond

A
  • Hydrogen with fluorine, nitrogen or oxygen

- Strongest IMF

27
Q

Why is water more dense than ice?

A

Water contracts when freezing. HB hold water molecules in a hexagonal lattice = low density

28
Q

Where are ionic compounds abundant in nature?

A

Fresh and salt water

29
Q

5 Properties of ionic compounds

A
  • Rigid structure
  • Solid at room temp.
  • High bp and mp
  • Crystal lattice
  • Brittle
30
Q

Why are ionic compounds brittle?

A

Rigid structure = ions cannot be rearranged without breaking the ordered structure of the crystal lattice

31
Q

5 Properties of metals

A 
Shiny 
Conductor 
Malleable
Low to high bp 
Crystal lattice
32
Q

Bonding Theory in metals (3)

A
  1. Low e/n = loosely held electrons
  2. Empty valence orbitals = electron mobility
  3. Electrostatic attraction between positive ions and negative electrons = strong bonds
33
Q

Why are metals lustrous?

A

Valence electrons absorb and re-emit light

34
Q

Why are metals malleable?

A

Nondirectional bonds mean that planes of ions can slide over each other while remaining bonded

35
Q

Why are metals able to conduct electricity?

A

Valence electrons can freely move throughout the metal

36
Q

Why do metals form a crystalline structure?

A

Electrons provide the “electrostatic glue” holding the metal ions together producing structures that are continuous and closely packed

37
Q

Properties of molecular solids

A
  1. Not conductors (in pure form and solution)
  2. Relatively low mp
  3. Not very hard
38
Q

Why can’t molecular compounds conduct electricity?

A

Individual entities are neutral

39
Q

2 Examples of covalent network crystals

A

Diamond

Quartz

40
Q

Properties of covalent network crystals

A
  1. Very hard
  2. Brittle
  3. Very high mp
  4. Insoluble
  5. Not conductors
41
Q

Why are covalent network crystals considered brittle?

A

Do not bend under pressure

42
Q

Describe covalent network crystal shape of a diamond

A

Carbon atoms in a large tetrahedral network with each carbon covalently bonded to four other carbon atoms

43
Q

Why are covalent network crystals so hard?

A

Interlocking structure meaning atoms are not easily displaced

44
Q

Why do covalent network crystals have such high melting points?

A

Many covalent bonds need to be broken which requires a lot of energy

45
Q

3 Properties of molecular crystals

A
  1. Soft
  2. Low mp
  3. Nonconducting (solid, liquid and solution)
46
Q

Forces/bond in molecular crystals

A

LDF
DD
HB

47
Q

5 Properties of metallic crystal

A
  1. Soft to very hard
  2. Conductor (solid and liquid)
  3. Ductile
  4. Malleable
  5. Lustrous
48
Q

3 Properties of ionic crystal

A
  1. Hard, brittle
  2. Medium to high mp
  3. Conductor (solid and solution)
49
Q

3 Examples of covalent network crystal

A

C(s)
SiC(s)
SiO2(s)

Chemistry 20 IB (4 decks)

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