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Chemistry > Chemical Bonding and Structure > Flashcards

Chemical Bonding and Structure Flashcards

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1
Q

structure of ionic compound

A

+ve and -ve ions attracted to each other by strong electrostatic forces, building up into strong crystal lattice

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2
Q

polyatomic ions

A

ions formed form >1 element where charge = often delocalised over whole ion. -ve ions = known as acid radicals bc formed when acid loses 1/+ H+ ions

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3
Q

hydroxide

A

(OH)-

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4
Q

nitrate

A

(NO3)-

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5
Q

hydrogencarbonate

A

(HCO3)-

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6
Q

phosphate

A

(PO4)3-

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7
Q

sulfate

A

(SO4)2-

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8
Q

hydrogensulfate

A

(HSO4)-

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9
Q

carbonate

A

(CO3)2-

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10
Q

ethanoate

A

(CH3COO)-

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11
Q

what EN diff. needed for ionic compounds to form?

A

> 1.8

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12
Q

what is ionic bond?

A

sum of all electrostatic attractions and repulsions w/in the lattice

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13
Q

describe the properties of ionic compounds:

A
  • high MP bc large amount of E to break lattice
  • soluble in H2O bc hydration E of ion provides E to overcome lattice enthalpy
  • solid ionic compounds don’t conduct bc ions = fixed positions
  • molten ionic compounds conduct, bc ions = free to move and carry charge
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14
Q

covalent bond

A

electrostatic attraction between shared pair of electrons and nuclei of atoms making bond

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15
Q

Lewis structure

A

shows all valence electrons of molecule

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16
Q

coordinate/ dative bond

A

when electrons in shared pair originate from same atom

e.g. sulfur dioxide, sulfur trioxide

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17
Q

how does bond length and strength vary?

A

single bonds = longest and weakest

triple bonds = shortest and strongest

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18
Q

bond polarity

A

when one end of molecule = more electron rich bc atom w/ > EN exerts > attraction for electron pair. small diff. in charge. > EN diff. = > polarity

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19
Q

describe bond polarity in diatomic molecules:

A

no polarity bc electron pair shared equally bc both atoms exert identical attraction

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20
Q

VSEPR theory

A

Valence Shell Electron Pair Repulsion theory states that pairs of electrons arrange themselves around central atom to be as far away as possible. > repulsion between non-bonded pairs than bonded electron pairs
mostly refers to domains bc a triple bond counts as one domain/ pair

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21
Q

octet rule

A

tendency of an atom (except H) in molecule/ ion to have 8 valence electrons
exceptions inc. phosphorus, sulfur + boron

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22
Q

shape + bond angle for 2 domains:

A

linear, 180

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23
Q

shapes + bond angle for 3 domains:

A
  • trigonal planar (3 bonding pairs), 120

- bent/ v-shaped (2 bonding pairs, 1 lone pair), <120

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24
Q

shapes + bond angles for 4 domains:

A
  • tetrahedral (4 bonding pairs), 109.5
  • trigonal pyramidal (3 bonding pairs, 1 lone), <109.5 e.g. 107
  • bent/ v-shaped (2 bonding, 2 lone), «109.5 e.g. 105
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25
Q

shapes + bond angles for 5 domains:

A
  • trigonal bipyramidal (5 bonding pairs), 120
  • seesaw/ distorted tetrahedral (4 bonding pairs), <180
  • T-shaped (3 bonding pairs), «180
  • linear (2 bonding pairs), 180
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26
Q

shapes + bond angles for 6 domains:

A
  • octahedral (6 bonding), 90, 180
  • square pyramidal (5 bonding), <90, 180
  • square planar (4 bonding), 90, 180
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27
Q

resonance hybrids

A
  • extreme forms of the true structure, which lies somewhere in between i.e. when >1 Lewis structure can be drawn for a molecule
  • e.g. ozone, bc bond length between O atoms = between double bond + single bond
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28
Q

allotropes

A

occur when an element can exist in diff. crystalline forms

29
Q

allotropes of carbon

A

diamond, graphite (graphene), buckminsterfullerene

30
Q

diamond

A
  • each C atom covalently bonded to 4 other C atoms = giant covalent structure
  • bonds = equally strong w/ no plane of weakness so v. hard
  • electrons = localised, can’t conduct electricity
31
Q

graphite

A
  • each C atom strong bonds to 3 others = layers of hexagonal rings w/ weak bonds between layers
  • excellent lubricant bc layers slide over each other
  • delocalized electrons = good conductor
32
Q

graphene

A
  • single layer of hexagonally arranged C i.e. 1 atom thick of graphite
  • v. light
  • semiconductor
  • 200x stronger than steel
  • magnetic form = graphone
33
Q

buckminsterfullerene

A
  • 60 C atoms in hexagons/ pentagons = geodesic spherical structure
  • led to nanotechnology
34
Q

what does molecular polarity depend on?

A
  • EN of atoms
  • shape of molecule
  • if individual bonds = polar doesn’t always mean molecule = polar bc resultant dipole may cancel out individual dipoles
35
Q

van der Waals’ forces

A

-general term for intermolecular forces inc. dipole-dipole, dipole-induced dipole, LDF

36
Q

London Dispersion Forces

A
  • weakest intermolecular forces
  • instantaneous dipole-induced dipole forces existing between any atom
  • this is bc electrons can be unevenly spread, producing temporary dipoles, leading to weak attraction between particles
  • increases w/ increasing mass
37
Q

Dipole-dipole forces

A

-when polar molecules = attracted to each other by electrostatic forces

38
Q

H bonding

A
  • strongest intermolecular force
  • occurs when H bonded to small v. EN element e.g. O, N, F
  • electron pair drawn away from H by EN element, proton remains
  • proton attracts lone pair electrons from F, N, O = v. strong dipole-dipole attraction
39
Q

how does solubility depend on bonding type?

A
  • polar solvents dissolve polar substances
  • organic molecules have polar head, non polar C tail
  • as length of chain increases, solubility in water decreases
  • ethanol = good solvent for many substances bc contains both polar and non-polar ends
40
Q

how does conductivity depend on bonding type?

A
  • needs electrons/ ions that are free to move
  • metals and graphite hav delocalised electrons = v. good conductors
  • molten ionic salts also conduct electricity but chemically decomposed in the process
41
Q

why is BP a good indicator of strength of intermolecular force but not MP?

A
  • when liquid -> gas, attractive forces between particles broken
  • when solid -> liquid, crystal structure broken down, but still attractive forces between particles - they show strength but also determined by way particles pack in crystal state and presence of impurities (impurities weaken structure = lower MP)
42
Q

describe the MP and BP for covalent structures:

A

-cov. bonds = so macromolecular covalent structures = v. high MP + BP
-also depends on type of attraction forces between molecules:
H bonding > dipole-dipole > LDF

43
Q

MP and BP of ionic compounds

A
  • relatively high MP and BP bc ionic attractions

- doesn’t necessarily mean that compounds w/ smaller highly charged ions = higher MP and BP

44
Q

describe metallic bonds:

A
  • close packed lattice of +ve ions in sea of delocalised electrons
  • metallic bond = attraction that 2 neighbouring particles have for the electrons between them
  • malleable + ductile bc close packed layers of +ve ions can slide over each other without breaking bonds
45
Q

alloy

A

–metallic solid solution- made of >1 metal but some have carbon e.g. steel = Fe + C

46
Q

properties of alloys in comparison to pure metals:

A
  • -lower MP
  • -added metals hav diff. radius/ even charge, distorting structure of original metal so bonding = < directional
    • < ductile, malleable bc impurities disturb lattice
  • -harder as well
47
Q

examples of alloys:

A
  • small amounts of C added to Fe = steel w/ high tensile strength
  • if Cr is added, produces stainless steel, w/ increased resistance to corrosion
48
Q

what does MP of metals depend on?

A
  • size of ion
  • charge of ion
  • way in which atoms = arranged in solid metal
49
Q

describe MPs of metals:

A
  • trend in grp 1: smaller metal ion = stronger bond, higher MP, also they melt below 181
  • most metals have high MPs but mercury = liquid
50
Q

sigma bond

A
  • 1st bond between 2 atoms

- when 2 atomic orbitals on diff atoms overlap head on bet. s, p and both

51
Q

pi bond

A
  • 2nd, 3rd bond

- 2 p orbitals overlap sideways on, above and below plane of nuclei

52
Q

what happens instead of resonance?

A
  • instead of forming double bond, electrons can delocalise over all atoms = energetically more favourable
  • delocalisation can occur when alternate and single bods occur bet. C atoms
53
Q

Formal Charge

A

FC = (no. valence electrons) - (no. lone electrons) - 1/2(no. bonding electrons)

54
Q

what assumption is formal charge based on?

A

assumes that all atoms in a molecule have the same EN

55
Q

what is formal charge used for?

A
  • to determine which out of several potential Lewis structures = preferred when several = possible
  • even if total formal charge is equal, preferred structure is when individual atoms have lowest formal charges
56
Q

describe bonds in ozone:

A

-‘one and a half’ bond between oxygen atoms, weaker than double bond in O2

57
Q

why is the ozone layer important?

A

diff. in bond enthalpies helps protect us from sun’s harmful UV radiation. rate of ozone production = equal to rate of ozone destruction- this process absorbs a lot of UV light- known as steady state

58
Q

describe how ozone protects us:

A
  • stratospheric ozone in equilibrium to O2, continually formed and decomposed.
  • strong double bond in O2 = broken by high E UV to form O atoms called radicals- possess unpaired electron and v. reactive so radicals can react w/ O2 to form ozone
  • weaker O3 bonds need lower E UV to break them
  • broken into O2, and O radical, which can then react w/ O3 to form 2 O2 molecules
59
Q

how to find wavelength of UV light necessary for O2 and O3 dissociation:

A

ʎ < 242 nm (highest energy)
UV radiation of the right frequency can break bonds
use E = hv and c = ʎv to find wavelength

60
Q

what causes destruction of ozone:

A

-CFCs and nitrogen oxides create holes in ozone

61
Q

how do CFCs catalyse ozone destruction?

A
  • high E UV in stratosphere -> homolytic fission of C-Cl to produce chlorine radicals
  • radicals break ozone -> more radicals so process continues
  • estimated that 1 CFC can catalyse breakdown of up to 100K molecules of ozone
62
Q

how do nitrogen oxides catalyse ozone destruction?

A
  • catalytically decompose ozone by radical mechanism

- oxygen radicals generated by breakdown of NO2 in UV

63
Q

hybridisation

A

-mixing of atomic orbitals to form new orbitals for bonding

64
Q

describe hybridisation in methane (CH4)

A
  • bond angles = 109.5
  • orbitals = far apart = same E (degenerate)
  • however, on C, 2s = lower E than 2p
  • hybridisation combines orbitals to produce 4 degenerate ones
  • therefore 1 electron promoted from 2s orbital to 2p orbital to form 4 sp3 orbitals
65
Q

describe hybridisation in ethene (C2H4)

A
  • C bonds to 3 atoms, so only 2s and 2 of 2p orbitals used
  • leaves unaffected p orbital
  • formation of double bond
  • 3 sp2 orbitals = far apart so 120, trigonal planar
  • unaffected p orbital then overlaps to form pi bond
66
Q

describe sp hybridisation and give an example:

A
  • 2s orbital hybridises w/ 1 2p orbital to form linear sp hybrid w/ 180
  • remaining p orbitals overlap to form 2 pi bonds
  • e.g. ethyne, N2
67
Q

how to predict type of bonding and hybridisation in carbon compounds:

A
  • each single bond = sigma
  • each double bond = 1 sigma and 1 pi
  • each C either side of a double bond = sp2 hybridised
68
Q

what can hybridisation show?

A

shape and bond angles of molecules

Chemistry (10 decks)

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