1
Q
4 allotropes of carbon
A
graphene, graphite, diamond, C60
2
Q
bonding of graphite
A
- each carbon is bonded to 3 other carbons
- all carbons are sp2 hybridized
- layers are held together by London Dispersion Forces
- weak intermolecular forces
3
Q
bonding of diamond
A
- each carbon is bonded to 4 other carbons
- all carbons are sp3 hybridized
4
Q
bonding of C60
A
- each carbon is bonded to 3 other carbons
- all carbons are sp2 hybridized
- carbons form 60 member sphere
5
Q
bonding of graphene
A
- each carbon is bonded to 3 other carbons
- all carbons are sp2 hybridized
6
Q
physical properties of graphene (that make it useful)
A
- good electrical conductor
- good thermal conductor (best)
- strong
- high melting point
- flexible
7
Q
physical properties of graphite
A
- good electrical conductor
- low thermal conductivity
- high melting point
8
Q
physical properties of diamond
A
- not electrical conductive
- very high thermal conductivity
- high melting point
9
Q
physical properties of C60
A
- semi electrical conductor
- very low thermal conductivity
- low melting point
10
Q
elemental silicon
A
- lot like diamond
- each silicon atom is covalently bonded to 4 other silicon atoms
- all silicon is sp3 hybridized
- results in giant lattice structure
11
Q
silicon dioxide (SiO2)
A
- silica / quartz
- giant covalent structure based on tetrahedral arrangement
- each silicon atom is bonded to 4 oxygen atoms
- each oxygen atom is bonded to 2 silicon atoms
- strong
- insoluble
- high melting point
- non-conductor of electricity
12
Q
carbonate
A
C03 2-
13
Q
sulfate
A
SO4 2-
14
Q
nitrate
A
NO3 -
15
Q
hydroxide
A
OH -
16
Q
hydrogencarbonate
A
HCO3 -
17
Q
ammonium
A
NH4 +
18
Q
phosphate
A
PO4 3-
19
Q
transitional metals
A
- have electron configuration that allows them to lose different # of electrons from their d sub-shell
- form stable ions with different charges -> these ions would have distinct properties such as colors
20
Q
Silver ion charge
A
- Ag+ always
- never write oxidation number
21
Q
oxidation number
A
symbols come before numbers
e.g. +2 instead of 2+
22
Q
Tin
A
Sn
23
Q
Mercury
A
Hg
24
Q
formation of ionic bonds
A
- electrostatic attraction between oppositely charged ions
- ions form through the process of ionization: electrons are transferred between atoms
25
ionic lattice
- 3D crystalline structure formed by ionic compounds
- always involves fixed arrangement of ions based on repeating units
26
coordination number
number of ions that surround a given ion in a lattice
27
lattice energy
- measure of strength of attraction between ions in a lattice
- great for ions that are small & highly charged (up, right on periodic table) bc they have larger charge density
28
physical properties of ionic compounds
- higher MP & BP: electrostatic attractions btwn ions in lattice are strong -> require large amount of energy to break
- solid at room temp
- conduct electricity when molten or in aqueous solution since ions are free to move
- cannot conduct electricity when solid as ions are held in fixed positions & cannot move
29
melting point & boiling point
- higher when charge on ion is greater <- due to increased attraction btwn ions
e.g. MP of MgO > Na2O (Mg: 2+, Na: 1+)
30
solubility
- determined by the degree to which separated particles of solute are able to form bonds/attractive forces with solvent
- ionic compounds: polar solvents
- covalent compounds: polar compounds -> polar solvents, nonpolar compounds -> nonpolar solvents
- metallic compounds: insoluble
31
hydration
- told by (aq)
- substance now exists as individual ions surrounded by water molecules
32
brittleness
- brittle: shatter when large force is applied
- ionic compounds: brittle bc movement of ions within lattice places ions of same charge next to each other -> repulsive forces cause the crystal to split
33
electronegativity
≤ 0.4: non-polar
<1.8: covalent
≥ 1.8: ionic
34
aluminum chloride (AlCl3)
- solid: ionic, AlCl3
- molten (liquified by heat): separates into Al2Cl6 molecules instead of its ions
- gas: ionic, AlCl3
35
allotropes
different forms of same chemical element in same physical state
36
delocalized
when electrons are not restricted to one location
37
giant molecular structure / network covalent structure
- crystalline lattice in which atoms are linked by covalent bonds
38
relative strength of intermolecular forces
hydrogen bonds > dipole-dipole forces > London dispersion forces
39
London dispersion forces
- temporary/instantaneous dipole: weak dipole formed when density of electrons at one point is greater over one region of molecule/atom
- induced dipole: caused when temporary dipole affects the electron distribution of a neighbor atom/molecule
- LDF: result of induced dipole
- strength increases with molecular size bc greater # of electrons increase chance of temporary dipoles developing
- only force that exists btwn non-polar molecules
40
physical properties of non-polar molecules
- gas at room temp
- low MP & BP as relatively little E is required to break weak intermolecular forces
41
dipole-dipole attraction
- permanent dipole: permanent separation of charge within bonds due to electronegativity difference btwn bonded atoms (in polar molecules), one end is δ+ while the other end has δ-
- dipole-dipole attraction: force of attraction caused when permanent dipoles result in opposite charges on neighboring molecules attracting each other
- vary depending on distance & relative orientation of dipoles
- cause MP & BP of polar substances to be higher than those of non-polar substances of comparable molar mass
- lead to solubility of polar solutes in polar solvents
42
Van der Waals’ forces
- umbrella term to include LDF, dipole-dipole attractions & dipole-induced dipole forces
43
hydrogen bond
- molecule containing hydrogen bonds to very electronegative atom (F, O or N) -> molecules attracted to each other by a particularly strong type of intermolecular force
- large EN difference results in electron pair being pulled away from H -> H has exposed proton & exerts strong attractive force on lone pair in electronegative atom of a neighboring molecule
44
conductivity of covalent compounds
- not able to conduct electricity in solid & liquid state since they don't contain ions
45
metalic bond
- electrostatic attraction btwn lattice of positive ions & delocalized electrons
- strength depends on # of delocalized ions, charge of cation & radius of cation
46
metal alloy
- substance that combines more than one metal or mixes a metal with other non-metallic elements
e.g. brass: copper & zinc
47
physical properties of metallic compounds
- good electrical conductivity: delocalized electrons are highly mobile so they can move through metal structure under applied voltage
- good thermal conductivity: delocalized electrons & close packed cations enable efficient transfer of heat energy
- high MP & BP: lots of energy is required to break strong metallic bonds & separate atoms
- ductile (can be drawn out into threads): movement of delocalized electrons are non-directional and random so metallic bonds remain intact while conformation changes under applied pressure
- malleable (can be shaped under pressure)
48
sigma bond
- formed by direct head-on/end-to-end overlap of atomic orbitals resulting in electron density concentrated between the nuclei of the bonding atoms
- hybridized
49
pi bond
- sideways overlap of atomic orbitals resulting in electron density above and below the plane of the nuclei of the bonding atoms
- unhybridized, p bond
- weaker than sigma bonds, can only form after one sigma bond has been formed
IB CHEM HL
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