Topic 4: Chemical Bonding and Structure Flashcards

1
Q

why do chemicals bond?

A

to achieve stability

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2
Q

ionic bond

A

AKA electrovalent bond

  • the sum of all electrostatic attractions and repulsions in the ionic compound
  • established by the transfer of electrons from one atom to another
  • to form ions with complete valence shells
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3
Q

conditions for the formation of ionic bonds

A
  • no of valence electrons
  • formation of ionic bonds is exothermic due to electron transfer and bond formation
  • big diff in electronegativity
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4
Q

nature of ionic bonds

A
  • solid ionic compounds held in crystal lattice

- ↑ lattice energy = ↑ ionic bond strength

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5
Q

characteristics of ionic compounds

A
  • high m.pt/b.pt: bonds require a lot of energy to break
  • solids at room temp: due to strong electrostatic forces of attraction between locked ions in crystal lattice
  • brittle: ions of the same charge are beside each other, so the repulsive forces cause it to split
  • soluble in water: as water is a polar solvent, it detaches ions from the crystal lattice due to its electrostatic pull
  • good conductors in molten state: as ions are free to move about
  • low volatility
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6
Q

covalent bond

A
  • bond formed by mutual sharing of electrons between the combining atoms
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7
Q

conditions for covalent bond

A
  • must have 5-7 valence e-s
  • equal electronegativity (so no transferring of electrons occurs)
  • equal electron affinity to equally attract the electron pair
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8
Q

covalent bonding parameters

A
  • bond length
  • bond angle
  • bond energy
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9
Q

bond length

A

avg. distance between centres of nuclei of 2 bonded atoms

unit: picometer (pm)/angstrom (Å)

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10
Q

factors affecting bond length

A
  • bond multiplicity: ↑ no. of bonds (i.e. single/double/etc), ↓ bond length
  • atom size: ↑ size, ↑ bond length
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11
Q

bond energy

A
  • energy required to break one mole of covalently bonded atoms in gaseous state
  • measures bond strength
    unit: kJ/mol
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12
Q

factors affecting bond energy

A
  • bond length: ↑ length, ↓ bond energy

- size of bonded atom: ↑ size, ↓ bond energy

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13
Q

coordinate (dative) bonds

A
  • type of covalent bond

- both electrons in the shared pair of e-s come from the same atom

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14
Q

polar covalent bond

A
  • covalent bond in which electrons are shared unequally
  • due to unequal electronegativity
  • thus bonded atoms acquire a partial positive/negative charge
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15
Q

VSEPR Theory

A

Valence Shell Electron Pair Repulsion Theory

  • electron pairs tend to repel each other
  • pairs try to stay as far apart to have min. energy and max. stability
  • greater repulsion between non-bonded e-s
  • the geometry of a molecule is dependent on the no. of valence e-s around the central atom
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16
Q

resonance structure

A

when the molecule’s characteristic properties can be described by 2+ structures, the real structure is a resonance hybrid of the possible structures

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17
Q

allotrope

A
  • different forms of the same element

- they have differing physical properties but similar chemical properties

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18
Q

allotropes of carbon

A
  • diamond
  • graphite
  • fullerene (C60)
  • graphene
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19
Q

structure of diamond

A
  • giant 3-D covalent tetrahedral arrangement
  • each C is strongly bonded to 4 other Cs (tetrahedral)
  • no plane of weakness
  • all bonds are equally strong
  • electrons can’t move freely
  • so diamond can’t conduct electricity
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20
Q

uses of diamond

A
  • cutting glass
  • making bores for rock drilling
  • grinding/polishing hard materials
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21
Q

structure of graphite

A
  • consists of 2-D hexagonal rings
  • each C atom is strongly bonded to 3 other Cs
  • weak Van der Waal forces between layers
  • layers can slide over each other so graphite is an excellent lubricant
  • graphite can also conduct electricity due to delocalized e-s, making it a good conductor
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22
Q

uses of graphite

A
  • making electrodes
  • lubricant
  • pencil lead
23
Q

structure of fullerene

A
  • perfect sphere
  • consisting of 60 C atoms arranged in hexagonal and pentagonal rings
  • soluble in suitable organic solvents
  • forms coloured solutions in organic solvents (varying from red to brown to magenta)
  • has delocalised e-s but doesn’t conduct electricity
24
Q

uses of fullerene

A
  • ferromagnetism

- super conductivity

25
graphene
- very thin: 1 atom thick! - 200x stronger than steel - single planar sheet of hexagonally-arranged C-atoms (similar to graphite in this sense) - excellent thermal and electrical conductivity (300x efficiency of copper) - a 1mm thick piece consists of 3 million stacked sheets - forms a carbon nanotube when rolled up - when rolled up more, becomes fullerene
26
types of intermolecular forces
- London dispersion forces - dipole-dipole forces - hydrogen bonding - ionic bonding - covalent bonding - metallic bonding
27
order of bond strength
ionic bonds > hydrogen bonds > dipole-dipole > London forces
28
types of van der waals forces
- London dispersion forces | - dipole-dipole forces
29
London dispersion forces
- weakest intermolecular force - existing in all covalent compounds - ↑ ease of polarisation = ↑ strength of London dispersion forces - caused by temporary dipole due to random electron movement - has inductive effect on neighbouring molecules
30
factors affecting London forces
- no. of e-s: ↑ no. of e-s = ↑ distance between valence e-s and nucleus = ↓ attraction to nucleus = ↑ ease of polarisation - size of e- cloud: ↑ size = ↑ ease of polarisation - shape of molecule: ↓ branching = ↑ ease of polarisation
31
polarisability
the ease of distortion of the electron cloud of a molecule by an electric field
32
dipole-dipole force
- slightly stronger than London dispersion forces - occurs due to electrostatic attraction between molecules with permanent dipoles - exists in all polar molecules with permanent dipole movement - b.pt increases significantly
33
hydrogen bond
- result of interaction - between un-bonded e- pair on one atom and a H atom in a different molecule that carries a high partial positive charge - due to the H atom being bonded to a small, highly electronegative element (e.g. O, N, F) - thus it's literally just a proton and the e- pair is attracted to it
34
effect of H-bonding on b.pt
- hydrogen bonds are strong dipole-dipole interactions - in Group 5, 6, and 7: hydrogen bonding causes m.pt >>>> compared to dipole-dipole bonding in other hydrides of that group - can be seen with water, HF, and NH3
35
Silicon Dioxide (SiO2), Quartz
- 3-D tetrahedral covalent network (just like diamond) | - each Si atom is covalently bonded to 4 O atoms, and each O atom is covalently bonded to 2 Si atoms
36
properties of silicon dioxide
- insoluble in water | - doesn't conduct electricity in solid state
37
solubility rules
- polar substances dissolve more easily in polar solvents (same rule applies to non-polar substances) - in organic molecules: ↑ length of chain = ↓ solubility to water
38
factors determining m.pt/b.pt
- bond strength - impurities - structure
39
metallic bonding
the electrostatic forces of attraction between the lattice of cations and the delocalised sea of electrons
40
factors affecting metallic bond strength
- no. of delocalised e-s: ↑ no. of delocalized e-s = ↑ bond strength - charge on cation: ↑ charge = ↑ bond strength - cation radius: ↑ radius = ↓ bond strength
41
properties of metallic compounds
- good conductors - shiny/lustrous: due to delocalised electrons reflecting light - high m.pt/b.pt - malleable (i.e. can be reshaped under pressure) - ductile (i.e. can be drawn out in a wire)
42
why are metals malleable and ductile?
- the close-packed cation layers can slide over each other | - and they can do so without breaking more bonds than are made
43
chemical bond
an attractive force that acts between 2 or more atoms to hold them together as a stable molecule
44
exceptions to octet rule
- elements before 3rd period may not fulfil octet rule - elements on and after 3rd period may exceed octet rule - because elements on and after 3rd period have d-orbitals
45
what is the electron domain/negative charge centre in VSEPR?
- double or triple bonded e-s are oriented together - they behave as a single unit in terms of repulsion - this is known as the electron domain/negative charge centre
46
what is the order of repulsion in VSEPR theory?
Lp-Lp > Lp-Bp > Bp-Bp Lp = lone pair Bp = bond pair - when a molecule has lone pairs of e-s, the bonding e- pairs are pushed closer - because the lone pair is free moving - but bond pairs are not, they must remain in the region between 2 atoms - this decreases Bp's bond angle
47
formal charge
no of atom valence e-s (according to periodic table) - no of valence e-s in a specified molecule (only includes 1 e- from each bond) - ideally formal charge should be as close to 0 as possible!
48
common alloys
- brass - bronze - solder - pewter - amalgam
49
components of brass
``` principal metal: copper added metal(s): zinc ```
50
components of bronze
``` principal metal: copper added metal(s): tin ```
51
components of solder
``` principal metal: lead added metal(s): tin ```
52
components of pewter
``` principal metal: tin added metal(s): can be copper, antimony, bismuth, or lead ```
53
components of amalgam
``` principal metal: mercury added metal(s): can be tin, silver, gold, or sodium ```