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wow woah insane its the best science flash cards in the observable universe confirmed > Chemical structure and bonding > Flashcards

Chemical structure and bonding Flashcards

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1
Q

Properties of metal

A
  • high melting point
  • good conductors of electricity and heat
  • malleable and ductile
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2
Q

Metallic bonding

and what increases its strength

A

Strong electrostatic attraction between metal cations and delocalised electrons

Strength increased by:

  • more delocalised electrons
  • smaller ion size
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3
Q

Metallic structure

A

Giant lattice of metal cations with delocalised electrons

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4
Q

Why metal conducts electricity

and what makes it better at conducting

A

Delocalised electrons can carry negative charge through the metal

Conductivity increased by:

  • more delocalised electrons
  • smaller ion size (for higher electron density)
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5
Q

Why metal is a good thermal conductor

A

Delocalised electrons quickly transfer kinetic energy around the metal

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6
Q

Why metal is malleable and ductile

A
  • metallic bonding relies on delocalised electrons so it is non-directional
  • if layers of the lattice slide over each other, bonding stays intact
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7
Q

Ionic bonding

and what increases its strength

A

Strong electrostatic attraction between oppositely charged ions

Strength increased by:

  • larger ion charges
  • smaller ion size
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8
Q

Properties of ionic compounds

A
  • high melting point
  • conducts electricity when molten or dissolved
  • brittle
  • soluble in water
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9
Q

Why ionic compounds conduct electricity only when molten or dissolved

A
  • solid ionic compounds have charged particles, but they are not free to move
  • as part of a liquid, the ions can move freely and carry charge
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10
Q

Why ionic compounds are brittle

A
  • ionic bonding is between oppositely charged ions

- if layers of the lattice slide over each other, ions with the same charge repel, breaking the bonding

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11
Q

Why ionic compounds are soluble in water

A
  • water is a polar molecule

- poles of water molecules surround oppositely charged ions, pulling them out of the lattice

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12
Q

Demonstration of the existence of ions

A

Electrolysis of green copper chromate solution separates blue copper ions from yellow chromate ions

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13
Q

Covalent bonding

and what increases its strength

A

Strong electrostatic attraction between two nuclei and a bonding pair of electrons

Strength increased by:
- shorter bond length (smaller atomic radii)

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14
Q

Sigma bond

A
  • covalent bond formed by the overlap of two orbitals directly between the nuclei
  • the first type of covalent bond to form
  • stronger than a pi bond
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15
Q

Pi bond

A
  • covalent bond formed by a sideways overlap of two p orbitals
  • only forms after a sigma bond is present
  • present in double or triple covalent bonds
  • weaker than a sigma bond
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16
Q

Electronegativity

and what increases it

A

An element’s ability to attract a bonding pair of electrons

Increased by:
- higher nuclear charge
- smaller atomic radius
- less electron shielding
(highest in top right of table)
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17
Q

Polar covalent bond

A
  • covalent bond between atoms of different electronegativity (different elements)
  • EN diff causes a dipole to form
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18
Q

Dative covalent bond

A
  • covalent bond where both bonding electrons are from the same atom
  • functionally the same as any other covalent bond
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19
Q

Octet rule

A
  • atoms usually want 8 electrons in their outer shell
  • C, N, O, and F always follow this
  • H, Be, B, and elements past period 3 do not follow
20
Q

Polarising power of a cation

A

Ability of a cation to distort an anion’s electrons (increases as EN increases)

21
Q

Polarisability of an anion

A

Tendency of an anion to be distorted by a cation (increases as EN decreases)

22
Q

London forces

A
  • occurs in all molecules
  • electron density constantly moves around in molecules
  • when it is higher on one side, a dipole is formed
  • this induces more dipoles on neighbouring molecules, which are then attracted
23
Q

What makes london forces stronger

A
  • more electrons (larger molecule)

- more points of contact (flatter molecule / less branches)

24
Q

Permanent dipole forces

A
  • occurs only in polar molecules

- attraction between the dipoles of polar molecules

25
Q

Hydrogen bonding

A
  • H is bonded to something very electronegative (N, O, or F)
  • a lone pair of electrons is present
  • strong attraction between the very positive H pole and the negative electrons
  • the strongest intermolecular force
26
Q

Hydrogen bonding in water

A
  • HF has 3 lone pairs per molecule but only 1 H
  • ammonia has 3 Hs but only 1 lone pair
  • water has 2 lone pairs and 2 Hs so can form 2 hydrogen bonds per molecule
27
Q

What is required for two liquids to mix?

A

They must have the same intermolecular forces

28
Q

Intermolecular forces in alcohols

A
  • all alcohols can hydrogen bond due to the OH group
  • as the carbon chain gets longer, london forces increase and hydrogen bonding decreases
  • larger alcohols are less soluble in water
29
Q

Why is ice less dense than water?

A

Hydrogen bonds cause the solid molecules to form a crystal structure that is less dense than water

30
Q

Structure of diamond

A
  • giant covalent lattice
  • 4 sigma covalent bonds per carbon
  • very hard with high melting point
31
Q

Structure of graphite

A
  • giant covalent lattice
  • 3 sigma covalent bonds per carbon with delocalised 4th electrons
  • conducts electricity
  • layers easily slide over each other
  • high melting point
32
Q

Structure of fullerene

A
  • large ball shaped molecule
  • 3 bonds per carbon with delocalised 4th electrons
  • cannot conduct electricity because molecules
  • low melting point
33
Q

Repulsion between electron pairs

A
  • bonding regions cause less repulsion

- lone pairs cause more repulsion

34
Q

Shape of a carbon dioxide molecule

A
  • linear
  • 2 bond regions, 0 lone pairs
  • bond angle of 180
35
Q

Shape of a boron trifluoride molecule

A
  • trigonal planar
  • 3 bond regions, 0 lone pairs
  • bond angles of 120
36
Q

Shape of a sulfur dioxide molecule

A
  • v shaped
  • 2 bonding regions, 1 lone pair
  • bond angle of 119
37
Q

Shape of a water molecule

A
  • v shaped
  • 2 bonding regions, 2 lone pairs
  • bond angle of 104.5
38
Q

Shape of a methane molecule

A
  • tetrahedral
  • 4 bonding regions, 0 lone pairs
  • bond angles of 109.5
39
Q

Shape of an ammonia molecule

A
  • trigonal pyramidal
  • 3 bonding regions, 1 lone pair
  • bond angles of 107.5
40
Q

Shape of a phosphorus pentachloride molecule

A
  • trigonal bipyramidal
  • 5 bonding regions, 0 lone pairs
  • bond angles of 120, 180, 90
41
Q

Shape of a sulfur hexafluoride molecule

A
  • octahedral
  • 6 bonding regions, 0 lone pairs
  • bond angles of 90, 180
42
Q

What makes a molecule polar?

A
  • polar covalent bonds

- shape of the molecule (so that dipoles don’t cancel each other out)

43
Q

Example of a non-polar molecule with polar bonds

A

Methane, tetrachloromethane, carbon dioxide

44
Q

Example of a polar molecule

A

Water, chloromethane, ammonia

45
Q

Types of bonds present in single, double, and triple covalent bonding

A

Single - 1 sigma
Double - 1 sigma, 1 pi
Triple - 1 sigma, 2 pi

46
Q

Formation of ions

A

Electrons lost&raquo_space; cation formed

Electrons gained&raquo_space; anion formed

47
Q

Why ionic radius decreases across a period (for isoelectronic ions)

A
  • shielding stays the same
  • nuclear charge increases
  • attraction between nucleus and outer shell increases

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