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Chemistry IB HL > Topic 4: Chemical Bonding and Structure > Flashcards

Topic 4: Chemical Bonding and Structure Flashcards

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1
Q

why do chemicals bond?

A

atoms interact by transferring/sharing electrons to achieve stability

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2
Q

ionic bond

A
  • aka electrovalent bond

- established by the transfer of electrons from a metallic atom to a non-metallic atom

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3
Q

conditions for the formation of ionic bonds

A
  • no of valence electrons: 1st atom should have 1-3 valence e’s while 2nd atom should have 5-7 valence e’s
  • net lowering of energy: energy released as a result of electron transfer and formation of ionic compounds
  • higher lattice energy: the greater the energy, the greater the strength of an ionic bond
  • a big difference in electronegativity
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4
Q

lattice energy

A

the amount of energy released when one mole of an ionic compound crystal is formed from its cations and anions in gaseous state

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5
Q

characteristics of ionic compounds

A
  • solids at room temp: 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
  • high m.pt/b.pt: high temps required to overcome attractive forces
  • soluble in water: water’s polar properties detach ions from the crystal lattice by their electrostatic pull
  • good electricity 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
  • equal electronegativity so no transferring of electrons occurs
  • equal electron affinity to equally attract the electron pair
  • both atoms should have 5-7 valence electrons
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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
  • the average distance between centres of the nuclei of 2 bonded atoms
  • measured in picometer (pm)/angstrom (Å)
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10
Q

factors affecting bond length

A
  • bond multiplicity: single, double, or triple bonds. it’s inversely proportional to bond length.
  • size of atom: directly proportional to bond length
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11
Q

bond angle

A

the average angle between orbitals containing bonding electron pairs around the central atom in a molecule

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12
Q

bond energy

A
  • the amount of energy required to break one mole of a certain type between atoms in gaseous state
  • expressed in kJ/mol
  • measures bond strength
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13
Q

factors affecting bond energy

A
  • bond length: inversely proportional relation

- size of bonded atom: inversely proportional relation

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14
Q

polar covalent bond

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

Coordinate bond/dative bond

A

a covalent bond in which both electrons in the shared pair of electrons come from the same atom

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16
Q

VSEPR Theory

A

Valence Shell Electron Pair Repulsion Theory

  • the geometry of a molecule is dependent on the no. of valence e-s around the central atom
  • electron pairs tend to repel each other as e- clouds are negatively charged, so pairs try to stay as far apart to have min. energy/max. stability
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17
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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18
Q

types of intermolecular forces

A
  • Van der Waals/London Dispersion Forces
  • dipole-dipole forces
  • hydrogen bonding
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19
Q

order of bond strength

A

ionic bonds > hydrogen bonds > dipole-dipole > London forces

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20
Q

London dispersion forces

A
  • weak force existing in all covalent compounds
  • relatively stronger between easily-polarised molecules than between molecules that don’t polarise easily
  • due to temporary dipole caused by random electron movement
  • has an inductive effect on neighbouring molecules
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21
Q

factors affecting London forces

A
  • no. of e-s: higher no. of e-s = higher distance between valence e-s and nucleus = lower attraction to nucleus = electron cloud is more easily polarised
  • size/vol of electron cloud: the bigger the cloud, the higher the polarisability of electrons
  • shape of molecule: straight chain molecules can interact with each other across the molecule
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22
Q

polarisability

A

the ease of distortion of the electron cloud of a molecule by an electric field

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23
Q

dipole-dipole force

A
  • occurs due to electrostatic attraction between molecules with permanent dipoles
  • exists in all polar molecules with permanent dipole movement
  • b.pt increases significantly
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24
Q

hydrogen bond

A

occurs as a result of interaction between a non-bonding electron pair on one of the atoms with a H atom in a different molecule that carries a high partial positive charge

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25
Q

effect of H-bonding on b.pt of N/O/F hydrides

A
  • going across the periods, b.pt increases due to increasing strength of London forces and increasing molecular mass
  • however, from period 3 to period 2 in all groups but 4, there’s a sharp increase in b.pt due to H-bonding
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26
Q

allotrope

A

different forms of the same element having different physical properties but similar chemical properties

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27
Q

diamond

A
  • 3-D covalent network
  • each C atom is covalently bonded to 4 other C atoms in a tetrahedral arrangement
  • electrons can’t move freely
28
Q

uses of diamond

A
  • cutting glass
  • making bores for rock drilling
  • grinding/polishing hard materials
29
Q

graphite

A
  • consists of 2-D hexagonal rings

- layers are connected by weak Van der Waal forces, causing graphite’s soft nature

30
Q

uses of graphite

A
  • making electrodes
  • lubricant
  • pencil lead
31
Q

fullerene

A
  • perfect sphere consisting of 5 and 6 membered C rings
  • soluble in suitable organic solvents
  • forms coloured solutions in organic solvents (varying from red to brown to magenta)
  • has delocalised electrons but doesn’t conduct electricity
32
Q

uses of fullerene

A
  • ferromagnetism

- super conductivity

33
Q

graphene

A
  • one of the thinnest and strongest 2D materials
  • consists of a single planar sheet of C-atoms hexagonally arranged, and is only one atom thick
  • excellent thermal and electrical conductivity (300x the efficiency of copper)
  • a 1mm thick piece consists of 3 million stacked sheets
  • if a sheet is rolled up. it forms a carbon nanotube. if that is rolled up more, it becomes fullerene.
34
Q

uses of graphene

A

currently research is being done to optimise it. future applications include:
- graphene-plastic composite materials to replace metals in the aerospace industry

  • LCDs and flexible touch screens for mobile devices
35
Q

Silicon Dioxide (SiO2), Quartz

A
  • 3-D covalent network

- each Si atom is covalently bonded to 4 O atoms, and each O atom is covalently bonded to 2 Si atoms

36
Q

properties of silicon dioxide

A
  • insoluble in water

- doesn’t conduct electricity in solid state

37
Q

metallic bonding

A

the electrostatic forces of attraction between the lattice of cations and the delocalised sea of electrons

38
Q

factors affecting metallic bond strength

A
  • no. of delocalised e-s: directly proportional
  • charge on the cation: directly proportional
  • radius of the cation: inversely proportional
39
Q

properties of metallic bonds

A
  • good electrical conductivity
  • shiny/lustrous: due to delocalised electrons reflecting light
  • high m.pt
  • malleable and ductile
40
Q

chemical bond

A

an attractive force that acts between 2 or more atoms to hold them together as a stable molecule

41
Q

exceptions to octet rule

A
  • 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
42
Q

procedure for drawing Lewis dot structures

A
  1. Find the total no. of valence electrons; add an e- for every negative charge, subtract an e- for every positive charge
  2. Decide the central atom (least electronegative atom, except for H) and draw bonds; subtract the electrons used in the bonds from the calculated total in step 1
  3. Assign leftover electrons to terminal atoms and subtract those from the results of step 2
  4. If necessary, assign leftovers to the central atom. If central atom doesn’t have an octet, create multiple bonds. If it follows or exceeds the octet rule, you’re done! (remember to check formal charge
43
Q

How to predict VSEPR:

A
  1. Draw the dot structure, showing placement of valence e-s
  2. Count the no of “e- clouds” surrounding the central atom
  3. Predict the geometry of the electron clouds around the central atom
  4. Ignoring lone pairs, predict the geometry of the molecule/ion
44
Q

what is the electron domain/negative charge centre in VSEPR?

A
  • 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
45
Q

what is the order of repulsion in VSEPR theory?

A

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
46
Q

formal charge

A

no of atom valence electrons - no of valence electrons in the Lewis dot structure

  • ideally formal charge should be as close to 0 as possible!
47
Q

electron cloud

A

region of electron density

48
Q

Hybridisation

A
  • the process of intermixing orbitals of slightly different energies
  • to redistribute those energies and give a new set of orbitals
  • of equivalent energies and shape
49
Q

hybridised orbital

A

the new orbital formed from hybridisation

50
Q

why does Ca have a greater m.pt value than K?

A
  • Ca ionic charge is greater than K
  • Ca has more delocalized e-s
  • greater attraction of delocalized e-s and Ca2+
51
Q

why does Na2O have a higher m.pt than SO3?

A
  • Na2O is ionic, so it has strong bonds of electrostatic attraction between Na & O
  • SO3 has weak intermolecular/Van der Waals/London/dispersion/dipole-dipole attractions
  • intermolecular/Van der Waals/London/dispersion/dipole-dipole attractions more easily broken than ionic bonds
52
Q

bonding in carbon allotropes

A
  • covalent bonds

- Van der Waals/London/dispersion forces

53
Q

carbon allotropes containing delocalized e-s

A
  • graphite

- fullerene

54
Q

carbon allotropes containing no delocalized e-s

A

diamond

55
Q

structure of diamond

A

network/giant structure/macromolecular/3-D

56
Q

structure of graphite

A

layered/2-D/planar

57
Q

structure of fullerene

A
  • made up of molecules/spheres of atoms

- arranged in hexagons/pentagons

58
Q

bond angles of diamond

A

109 deg

59
Q

bond angles of graphite

A

120 deg

60
Q

bond angles of fullerene

A

109-120 deg

61
Q

no of atoms each C atom in diamond is bonded to

A

4

62
Q

no of atoms each C atom in graphite is bonded to

A

3

63
Q

no of atoms each C atom in fullerene is bonded to

A

3

64
Q

structure of SiO2

A
  • quartz has a giant/macromolecular/network structure
65
Q

bonding in SiO2

A
  • each Si atom is covalently bonded to 4 O atoms

- each O atom is covalently bonded to 2 Si atoms

Chemistry IB HL (11 decks)

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