Topic 4: Chemical Bonding and Structure Flashcards
why do chemicals bond?
to achieve stability
ionic bond
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
conditions for the formation of ionic bonds
- no of valence electrons
- formation of ionic bonds is exothermic due to electron transfer and bond formation
- big diff in electronegativity
nature of ionic bonds
- solid ionic compounds held in crystal lattice
- ↑ lattice energy = ↑ ionic bond strength
characteristics of ionic compounds
- 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
covalent bond
- bond formed by mutual sharing of electrons between the combining atoms
conditions for covalent bond
- must have 5-7 valence e-s
- equal electronegativity (so no transferring of electrons occurs)
- equal electron affinity to equally attract the electron pair
covalent bonding parameters
- bond length
- bond angle
- bond energy
bond length
avg. distance between centres of nuclei of 2 bonded atoms
unit: picometer (pm)/angstrom (Å)
factors affecting bond length
- bond multiplicity: ↑ no. of bonds (i.e. single/double/etc), ↓ bond length
- atom size: ↑ size, ↑ bond length
bond energy
- energy required to break one mole of covalently bonded atoms in gaseous state
- measures bond strength
unit: kJ/mol
factors affecting bond energy
- bond length: ↑ length, ↓ bond energy
- size of bonded atom: ↑ size, ↓ bond energy
coordinate (dative) bonds
- type of covalent bond
- both electrons in the shared pair of e-s come from the same atom
polar covalent bond
- covalent bond in which electrons are shared unequally
- due to unequal electronegativity
- thus bonded atoms acquire a partial positive/negative charge
VSEPR Theory
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
resonance structure
when the molecule’s characteristic properties can be described by 2+ structures, the real structure is a resonance hybrid of the possible structures
allotrope
- different forms of the same element
- they have differing physical properties but similar chemical properties
allotropes of carbon
- diamond
- graphite
- fullerene (C60)
- graphene
structure of diamond
- 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
uses of diamond
- cutting glass
- making bores for rock drilling
- grinding/polishing hard materials
structure of graphite
- 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
uses of graphite
- making electrodes
- lubricant
- pencil lead
structure of fullerene
- 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
uses of fullerene
- ferromagnetism
- super conductivity
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
types of intermolecular forces
- London dispersion forces
- dipole-dipole forces
- hydrogen bonding
- ionic bonding
- covalent bonding
- metallic bonding
order of bond strength
ionic bonds > hydrogen bonds > dipole-dipole > London forces
types of van der waals forces
- London dispersion forces
- dipole-dipole forces
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
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
polarisability
the ease of distortion of the electron cloud of a molecule by an electric field
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
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
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»_space;» compared to dipole-dipole bonding in other hydrides of that group
- can be seen with water, HF, and NH3
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
properties of silicon dioxide
- insoluble in water
- doesn’t conduct electricity in solid state
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
factors determining m.pt/b.pt
- bond strength
- impurities
- structure
metallic bonding
the electrostatic forces of attraction between the lattice of cations and the delocalised sea of electrons
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
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)
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
chemical bond
an attractive force that acts between 2 or more atoms to hold them together as a stable molecule
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
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
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
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!
common alloys
- brass
- bronze
- solder
- pewter
- amalgam
components of brass
principal metal: copper added metal(s): zinc
components of bronze
principal metal: copper added metal(s): tin
components of solder
principal metal: lead added metal(s): tin
components of pewter
principal metal: tin added metal(s): can be copper, antimony, bismuth, or lead
components of amalgam
principal metal: mercury added metal(s): can be tin, silver, gold, or sodium
