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chemistry > bonding and structure > Flashcards

bonding and structure Flashcards

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1
Q

why do atoms form chemical bonds

A

to increase stability, some achieve full outer shell, attraction between opposite charges

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2
Q

what are key properties of ionic compounds

A

generally soluble in water, high melting and boiling points, only conduct electricity when molten or dissolved, hard but brittle

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3
Q

why are ionic compounds soluble in water

A
  • water is polar molecule
  • energy released when polar waters form attractions to ions in lattice
  • energy overcomes electrostatic forces and breaks ions away
  • if not enough energy released then insoluble
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4
Q

why do ionic compounds have high melting and boiling points

A
  • lots of strong electrostatic forces between ions
  • require lots of energy to break
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5
Q

why do ionic compounds only conduct when molten or dissolved

A
  • ions can freely move and carry charge
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6
Q

why are ionic compounds hard but brittle

A
  • like charges repel if structure is distorted causing it to break apart
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7
Q

trend of ionic radii across period

A

decreases as more protons so electrons have greater attraction to nucleus and shielding stays the same

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8
Q

why do metals have high melting points

A

strong electrostatic attractions between cations and electrons which require lots of energy to overcome

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9
Q

why do metals conduct electricity

A

delocalised electrons which can move through structure and carry charge

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10
Q

why do metals conduct heat

A

delocalised electrons can move through structure and transmit kinetic energy

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11
Q

why are metals malleable and ductile

A

regular structure so rows of metal ions can slide over each other

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12
Q

how does electrons in outer shell affect strength of metallic bonding

A

more electrons, greater charge on metal ion and greater delocalised electrons, stronger electrostatic attraction between electrons and positive metal ions

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13
Q

how does ionic radius affect strength of metallic bonding

A

smaller ion, shorter distance between positive nucleus and delocalised electrons so stronger attraction

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14
Q

what do lone pairs affect

A

the shape of molecules, form dative covalent bonds

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15
Q

what is a dative covalent bond

A

two atoms share a pair of electrons but both the electrons are donated by one atom

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16
Q

how are covalent bonds formed in terms of orbitals

A

the overlap of orbitals

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17
Q

how are sigma bonds formed

A
  1. overlap of 2 s-orbitals
  2. overlap of an s and a p-orbital
    3.3 overlap of 2 p-orbitals end on
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18
Q

what is a sigma bond and what can it do

A

a single covalent bond that can rotate

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19
Q

how are pi bonds formed

A

overlap of 2 p-orbitals, electron density is concentrated above and below molecule

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20
Q

what is a pi bond

A

double covalent bond, cannot rotate

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21
Q

what affects the strength of a covalent bond

A

length of bond - stronger attraction, higher bond enthalpy, shorter bond length

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22
Q

properties of giant covalent e.g. diamond

A

high melting and boiling points, low electrical conductivity, high strength, insoluble

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23
Q

properties of graphite

A

high melting and boiling points, good electrical conductivity, strong but brittle, insoluble

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24
Q

properties of graphene

A

high melting and boiling points, good electrical conductivity, strong, insoluble

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25
properties of simple molecular
low melting and boiling points, bad electrical conductivity, brittle, slightly soluble
26
what does VSEPR state (valence shell electron pair repulsion theory)
- the shapes of molecules are caused by the repulsion between pairs of electrons, both bonding and lone pairs, that surround the central atom
27
factors affecting electronegativity
nuclear charge, distance between bonding pair of electrons and nucleus, electron shielding
28
London forces
- weakest IMF - exists between all molecules - caused by fluctuations of electron cloud - instantaneous dipole - induced dipole interactions - constantly breaking and forming due to random electron cloud fluctuations
29
factors affecting strength of London forces
- number of electrons within molecule (greater number, greater dipoles so stronger forces) - shape of the molecule (number of points of contact between molecules)
30
what is hydrogen bonding
strongest IMF, type of permanent dipole interaction, only between H on one molecule bonded to NOF (electronegative) which must have lone pair
31
ionic lattice structure
regular arrangement of ions in fixed ratio, ions are touching and held by strong electrostatic forces, attractions not directional
32
how to test the migration of ions
draw pencil line on paper soaked in solution attached to wires on a battery, place sample on pencil line and see which way the colours migrate (positive side they are negative ions and vice versa)
33
what is a covalent bond
electrostatic attraction between 2 nuclei and the shared pair of electrons between them - electrons that form bonds are called bonding pairs, otherwise lone pairs
34
another name for dative covalent bond
co-ordinate bond
35
what does valence shell electron repulsion theory state
- shapes of molecules are caused by repulsion between the pairs of electrons, both bonding and lone pairs, that surround the central atom - electron pairs arrange themselves around central atom to minimise repulsion and maximise separation
36
strength of repulsion of electron pairs in order
lone pair/lone pair > lone pair/bonding pair > bonding pair/bonding pair
37
how to use VSEPR to determine shape
1. draw dot and cross diagram 2. identify central atom, number of bonding pairs and number of lone pairs (determines total number of groups of electrons around central atom) 3. assign name and bond angle (double and triple bonds are treated as one 'group' of electrons)
38
linear
bond angle 180 2 groups of electrons around central atom and both are bonding pairs
39
trigonal planar
bond angle 120 3 bonding pairs of electrons
40
tetrahedral
bond angle 109.5 4 bonding pairs of electrons
41
trigonal bipyramid
bond angles 90 and 120 5 bonding pairs of electrons
42
octahedral
bond angle 90 6 bonding pairs of electrons
43
bent/v shaped
bond angle 104.5 4 groups of electrons (2 bonding 2 lone)
44
trigonal pyramid
bond angle 107 4 groups of electrons (3 bonding 1 lone)
45
how many degrees do two lone pairs repel by
2.5
46
square pyramid
bond angle 89 5 bonding 1 lone
47
square planar
bond angle 90 4 bonding 2 lone
48
electronegativity meaning
the measure of the ability of an atom to attract the bonding pair of electrons in a covalent bond
49
why does electronegativity increase across period
more protons across period some attracts electrons more (delta negative further across)
50
why does electronegativity decrease down gorup
more shielding
51
permanent dipole meaning
charge difference between atoms
52
Pauling scale
BOND TYPE - ELECTRONEGATIVITY DIFFERENCE covalent = 0 polar covalent = 0-1.8 ionic = >1.8
53
electron density distribution in non-polar covalent
evenly distributed
54
electron density distribution in polar covalent
not evenly distributed
55
electron density distribution in ionic
2 distinct regions of electron density
56
what is an induced dipole
dipoles caused by neighbouring molecules
56
what is an instantaneous dipole
a temporary dipole
57
why do melting and boiling points increase down group
more electrons in molecules, greater instantaneous and induced dipoles, stronger London forces, more energy needed to overcome
58
permanent dipole interactions
stronger than London forces, occur between polar molecules
59
how to draw hydrogen bonding
- consider shape of molecule - draw lone pairs of electrons - label partial charges - use dotted line to represent hydrogen bond (start at lone pair of electrons, drawn at 180 degrees)
60
non-organic solvents and solubility (H2O)
London forces and H bonding, ionic compounds generally soluble in it, non-polar substances insoluble in it, compounds with H bonds and organic compounds with N or O generally soluble in it (decreases as C chain increases)
61
organic substances and solubility (hexane)
London forces, ionic compounds insoluble in it, compounds with H bonding and organic compounds with N or O generally insoluble, non-polar substances soluble in it
62
how do solutes dissolve in a solvent
IMFs between solute particles break, IMFs between solvent particles need to break (requires energy), IMFs form between solute and solvent particles to form solution (releases energy) - can only dissolve in enough energy is released when IMFs form in solution to overcome IMFs in solute and solvent
63
how do ionic compounds dissolve in water
polar water molecules attracted to ions on outside of ionic lattice (ion-dipole attraction), energy released to overcome electrostatic forces between ions and H bonds in water (enthalpy of hydration)
64

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