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1
Q

emission spectrum of hydrogen

A

coloured lines of discrete frequencies/energy levels in a dark background, which correspond to electron transition from higher to lower energy levels.
Coloured lines converge at higher frequencies as energy levels are close together at higher energy levels

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2
Q

lyman

A

emission to n = 1

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3
Q

balmer

A

emission to n = 2 (visible light)

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4
Q

paschen

A

emission to n = 3

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5
Q

oxides and water

A

acidic oxide = acid
basic oxide = base
amphoteric oxide = insoluble, reacts with both acid and base BALZ barium aluminium lead zinc

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6
Q

whenever you see a find Kc qns

A

ICE table damnit (values are in concentration)

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7
Q

Gibbs free energy unit

A

kJmol-1 (around 3 digit, some formula gives u in Jmol-1)

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8
Q

how to find conc. of iodine with formula

A

titration with sodium thiosulfate
I2 + 2S2O3(2-) -> 2I(1-) + S4O6(2-)
aiya see onenote search ‘reaction of iodine with thiosulfate

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9
Q

standard electrode potential

A

standard electrode potential of a half-cell is the potential difference generated when it is connected to the standard hydrogen electrode at standard conditions

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10
Q

finding pH or H+

A

14 - pOH
[H+] = sqrt(Ka x (C-x))

Kw = [H+][OH-]

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11
Q

(polar) Protic solvent favoured by

A 
solvates nucleophile and carbocation intermediate
Favours sn1 (reduces Ea of rds cause carbocation is more stable)
Reduces effectiveness of sn2 (nucleophile gets solvated)
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12
Q

(polar) Aprotic solvent favoured by

A 
solvates cation counterion 
favours sn2 (increase reactivity of nucleophile to donate a pair of electrons)
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13
Q

transition metal r/s with periodic table group

A

group no. is no. of electrons in 3d + 4s orbitals

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14
Q

Electrophilic Addition: why is there major product

A

tertiary>secondary>primary (carbocation intermediate)
__ carbocation is more stable due to positive inductive effect caused by more electron-donating alkyl groups that disperse the positive charge on carbocation

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15
Q

reagents and conditions: nucleophilic substitution

A

NaOH, heat under reflux
KOH, heat under reflux
rmb to add lone pair and negative charge on transition state for mechanism

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16
Q

why atoms can exceed octet rule

A

period 3 and onwards (presence of 3d orbitals)

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17
Q

strong acid + weak base =

A 
acidic salt 
(strong conjugate acid, POE to left, will react with OH to return back into a weak base, reduce pH)

salt hydrolysis: salt + h2o will give H3O+ thats why acidic

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18
Q

weak acid + strong base =

A 
basic salt 
(strong conjugate base, POE to left, will react with H to return back into a weak acid, increase pH)

salt hydrolysis: salt + h2o will give original acid n OH- thats why basic

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19
Q

strong acid + strong base =

A

neutral salt

both have weak conjugate base/acid, so [H] : [OH] remains same

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20
Q

weak acid + weak base

A

neutral salt

both have strong conjudate base/acid, [H] : [OH] stay same

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21
Q

given Ka of weak acid, how to find Kb?

A

Kb is of the strong conjugate base
Ka = Kw/Kb
pKa = pKw - pKb

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22
Q

s. f of absolute uncertainty

* rmb to follow d.p of absolute uncertainty

A

1sf, round mathematically, dont always round up
(like in IA)
*rmb to follow d.p of absolute uncertainty

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23
Q

first order wrt a rxt

A

constant half-life

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24
Q

half-cell potential meaning (more +/- is…) and Ecell formula

A

more positive = reduction
more negative = oxidation
Ecell = reduction - oxidation
*rmb not to swap the sign, just plug into formula

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25
Q

when temperature increase, what happens to rate of forward reaction (assuming exothermic rxn)

A

increase in rate of reaction. even though POE to left by LCP, increase in temp will increase rate of reaction anyway.

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26
Q

acidic buffer

A

weak acid, strong base/neutral salt of conjugate base

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27
Q

basic buffer

A

weak base, strong acid/neutral salt of conjugate acid

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28
Q

coordination number

A

no. of ligands around a central (metal) ion

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29
Q

Colours of iodine in different states

A

solid: black
aqueous: brown
in organic solvent: purple
gaseous: purple

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30
Q

Index of hydrogen (IHD), no. of H2 molecules to make molecule non-cyclic (at all) and saturated

A
  1. no. of pi bonds + no. of rings

2. [(2n+2)-(no. of H)+(no. of grp 15)-(no. of grp 17)]/2

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31
Q

Reagent/conditions: reduction of carboxylic acid/aldehyde/ketone to alcohol

A
  1. LiAlH4, in dry ether then dilute acid (strongest - all 3)
  2. NaBrH4, in ethanol then dilute acid (ketone + aldehyde)
  3. H2, Ni catalyst, heat (alkene + ketone + aldehyde)
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32
Q

describe a covalent bond

A

electrostatic forces of attraction between positively charged nuclei and one or more shared pair of electrons.

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33
Q

Hrxn

A

Hf(products)-Hf(reactants)
BE(bonds broken)-BE(bonds formed)
Hcmbst(rxt) - Hcmbst(pdt)
*just recall the flow cycle diagram thing

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34
Q

first order rxn

A

conc/time: constant half-life

rate/conc. : linear increase

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35
Q

zero order rxn

A

conc/time: linear decrease, decreasing half-life

rate/conc: constant

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36
Q

second order rxn

A

conc/time: increasing half life

rate/conc: exponential increase

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37
Q

how does catalyst increases rate of reaction

A

catalyst provides an alternative pathway with a lower activation energy. This results in a greater proportion of particles having energy greater than the activation energy, hence rate increases.

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38
Q

buffer eqns (finding ka

A

Acidic buffer:
pH = pKa + lg([salt or c.base]/[acid])
pOH = pKb + lg(acid/salt or c.base)

Basic Buffer:
pH = pKa (of c.acid) + lg([base]/[salt or c.acid])
pOH = pKb + lg([c.acid/salt] or [base])

*all in terms of concentration
at MBC, concentration of base n acid is 1/1, so pH = pKa, pOH = pKb

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39
Q

at maximum buffer capacity….

A

vol. of weak acid x2 of vol. of strong base
and vice versa

pH = pKa
pOH = pKb
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40
Q

amphiprotic

A

can act as bronsted-lowry acid and base

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41
Q

when to use reversible arrow

A

equilibrium qns, acid/base qns, salt hydrolysis

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42
Q

why is halogenoalkane more reactive than alkanes?

A

C-Halogen bond is weaker than C-H bond (except for C-F)

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43
Q

formation of NO2+

A

HNO3 + H2SO4 -> NO2+ + HSO4- + H2O

by right should be reversible arrow

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44
Q

nitrobenzene R grp, intermediate R grp, phenylamine R grp

A

nitrobenzene: NO2
intermediate: NH3+
phenylamine: NH2

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45
Q

reference compound in NMR (name n purpose)

A

tetramethylsilane
all H atoms in same environment, one strong integration trace, no splitting
inert/stable

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46
Q

anode/cathode: oxidation/reduction

A

anode: oxidation
cathode: reduction

true for both voltaic and electrolytic cells

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47
Q

standard enthalpy of hydration

A

1 mole of gaseous ions form 1 mole of aqueous ions

exothermic

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48
Q

standard enthalpy of solution

A

1 mole of solute is dissolved (aqueous, ions)

EC(soln) = EC (lattice) + EC(hyd)
cam be endo n exo, depending on LE

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49
Q

volume of 1 mole of gas at STP (273K, 1atm)

A

22.7dm^3/mol

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50
Q

volume of 1 mole of gas at SATP (298K, 1atm)

A

24.8dm^3/mol

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51
Q

when finding empirical formula:

A

if given % or mass of different elements, take % or mass divided by Mr to find mole ratio

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52
Q

for Gibbs free energy, enthalpy change, entropy change, rmb to:

A

put negative or positive sign

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53
Q

reagents n conditions: reduction of nitrobenzene to form phenylamine

A

step 1: Sn (tin) in concentrated HCl, heat under reflux
C6H5NO2 + 3Sn + 7H+ -> C6H5NH3+ + 3Sn2+ + 2H2O

step 2: NaOH (aq)
C6H5NH3+ + OH- -> C6H5NH2 + H2O

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54
Q

reagents and conditions: electrophilic substitution of nitrobenzene from benzene

A

concentrated HNO3, and concentrated H2SO4 catalyst

50 deg cel

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55
Q

rate mechanism of electrophilic substitution of nitrobenzene from benzene

A 
step 1 : HNO3 + H2SO4 -> NO2+ + HSO4- + H2O
step 2 (slow step): curly arrow from benzene electrons to N in NO2-, C is bonded to H and NO2, resonance structure is a U with + charge instead of O
step 3: curly arrow from C-H bond into U+ , form nitrobenzene and H+
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56
Q

identifying mass spectrum species from mass spectrum, rmb to:

A

rmb to put positive charge on molecule. add up Mr of atoms in molecule

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57
Q

conjugated system of electrons

A

electrons are delocalised within entire system of molecules instead of within 1 molecule (graphite, graphene)

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58
Q

delocalisation

A

electrons are shared by more than 2 atoms in a molecule as compared to being localised between a pair of atoms

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59
Q

what is a transition metal? and properties

A

transition metals form at least 1 stable ION with a PARTIALLY filled d sub-level orbital.

thats why zinc (full orbital) and scandium (empty orbital) are not considered transition metals

  1. forms complex ion
  2. coloured ions/compounds
  3. variable oxidation states
  4. paramagnetism
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60
Q

intermolecular forces

A
  1. london dispersion forces (instantaneous dipole-induced dipole moment
  2. permanent dipole (dipole-dipole)
  3. hydrogen bonding
    * 1 and 2 are known as van der waals forces
61
Q

OH grp name

A

hydroxyl

62
Q

resonance

A

Resonance occurs when there is more than one possible Lewis structure, due to the positioning of delocalised pi-electrons across more than two atoms

Resonance is possible whenever a Lewis structure has a multiple bond and an adjacent atom with at least one lone pair.

63
Q

bond order

A

total number of bond pairs divided by number of bonding locations

64
Q

formal charge

A

no. of valence electrons - no. of bond PAIRS - no. of non-bonding electrons
* sum of FC will equal the total charge of molecule

most reasonable lewis structure:

  • FC diff (FCmax - FCmin) closest to zero
  • negative charges located on most electronegative atoms
65
Q

how ligands form bonds

A

dative covalent bond from __ to central metal ion (and is hence also a Lewis base)

66
Q

structural isomers

A

chain
positional
functional grp

67
Q

stereoisomerism

*rmb to draw 3d structure

A 
Configurational (cis-trans, EZ)
Optical isomerism (enantiomerism, enantiomers, diastereomers, meso compounds)
68
Q

EZ isomers

A

E is opposite side
Z is same side
when all substituents are different

69
Q

cis-trans isomers

A

cis is same side
trans is opposite site
- need at least 1 substituent that is same on both sides of restrictee bond
- restricted rotation about a bond (either presence of double/triple bond, or a ring structure)

70
Q

racemic mixtures

A

equimolar amounts of each enantiomer, no net rotation as they cancel each other out

71
Q

diastereomers

A

different configuration at one or more, but not all, chiral centres from a pair of enantiomers, will still be optically active

72
Q

meso compound

A

contain chiral centres but is optically inactive due to plane of symmetry

73
Q

reagents and conditions: hydrogenation of alkenes

A
  1. Ni catalyst at high temp (150degcel) and pressure
  2. Pd (palladium) catalyst, r.t.p
  3. Pt (platinium) catalyst, r.t.p
74
Q

reagents and conditions: (alkenes) electrophilic addition of halogens, hydrogen halides, inter-halogen compounds

A

Halogen, r.t.p in the dark

organic/non-polar solvent like CCl4 and polar solvent like H2O have different products, check onenote if not sure

75
Q

reagents and conditions: free radical substitution

A 
X2, UV light. will decolourise Cl2 or Br2 gas
after initiation (homolytic fission of Cl-Cl bond), dont need UV light anymore as it will propagate
76
Q

reagents and conditions: electrophilic addition of steam in alkenes

A

industrial: H2O(g), concentrated H3PO4 catalyst, 300 degcel, 65 atm
laboratory: concentrated H2SO4, followed by boiling with H2O (l)

77
Q

reagents and conditions: elimination of alcohols to form alkenes

A
  1. excess concentrated H2SO4, 170-180 degcel
  2. Al2O3 catalyst, 350 degcel
  3. Excess concentrated H3PO4, 250 degcel
78
Q

reagents and conditions: elimination of halogenoalkanes

A
  1. Alcoholic KOH, heat

2. Alcoholic NaOH, heat

79
Q

preparation of halogenoalkanes

A
  1. free radical sub
  2. electrophilic substitution of hydrogen halide in non-polar solvent
  3. Halogenation in non-polar solvent
80
Q

primary halogenoalkanes prefer nucleophilic sub by sn1/sn2?

A

sn2

81
Q

tertiary halogenoalkanes prefer nucleophilic sub by sn1/sn2?

A

sn1

82
Q

reagents and conditions: oxidation of alcohols n colour change observed

A 
acidified KMnO4 (stronger) (purple to colourless)
acidified K2Cr2O7 (weaker) (orange to green)

primary alcohols
to aldehyde:
acidified K2Cr2O7, heat with immediate distillation

to carboxylic acid:
acidified KMnO4, heat under reflux
acidified K2Cr2O7, heat with reflux

secondary alcohols:
to ketone:
acidified K2Cr2O7/KMnO4, heat under reflux

tertiary alcohols do not undergo oxidation

83
Q

what does A in arrhenius equation account for?

A

frequencies of collisions between reactant molecules. it is a constant

84
Q

enthalpy change of atomisation

A

energy required to form one mole of gaseous atoms from element in standard state (always endo)

85
Q

limitation of using Bond Enthalpy to find enthalpy change of rxn

A
  1. BE is an average value for polyatomic molecules, diatomic is exact
  2. BE is only for gas molecules, if not gas then need to add enthalpy change of vaporisation (phys)
86
Q

lattice energy enthalpy change

A

one mole of solid ionic compound is separated into gaseous IONS. always endothermic

87
Q

standard hydrogen electrode

A

Pt electrode
1.00moldm-3 H+ (aq)
H2 gas at 100kPa (1 bar)

88
Q

halogen standard state

A

fluorine - gas
chlorine - gas
bromine - liquid
iodine - solid

89
Q

ether functional grp

A

R-O-R’

90
Q

alkyne functional grp

A

C triple bond C

91
Q

rate constant unit (k)

A

zero: mol dm^-3 s^-1
first: s-1
second: mol^-1 dm^3 s^-1
basically rmb that “rate” itself is moldm^-3 s^-1

92
Q

lewis base donate or accept e- pair?

A

donate

93
Q

esterification reagents n conditions

A

alcohol+carboxylic acid, concentrated H2SO4 catalyst, heat under reflux

94
Q

carbonyl functional grp

A

C=O

95
Q

naming esters

A

Name the alcohol part before the acid part, replacing “oic acid” with “oate”

96
Q

how to determine sp1, sp2, or sp3 (hybridisation)

A

look at number of electron domains that the atoms has

97
Q

calculate change in entropy

A

S(pdt) - S(rxt)

98
Q

why does Kc change with temperature change?

A

(assuming temeprature rise)

favour endo, POE shifts to the left, backward rate more than forward rate, [rxt] more than [pdt], so Kc reduces

99
Q

when doing electrolysis

A

always just write out overall eqn, where all numbers are integers, so can find n (no. of electrons) for change in G

100
Q

sn1 stands for

A

unimolecular nucleophilic substiution

101
Q

sn2 stands for

A

bimolecular nucleophilic substitution

102
Q

when drawing nucleophilic substiution,

A

always freaking use 3d diagram pls

103
Q

nmr use what em wave

A

radio waves
IR - infrared
Uv vis - UV bruh

104
Q

forming different lewis structures to test FC

A
  • electron pushing method
  • can be enantiomers (just mirror)
  • just shift the double bond around
105
Q

Difference Between Formal Charge and Oxidation State

A
  • formal charge is the charge of an atom in a molecule (assuming electrons are shared equally between 2 atoms)
  • oxidation state is the number of electrons an atom loses or gains or shares with another atom.
106
Q

why are transition metals coloured

A

ligands cause d-orbitals to undergo d-d orbital splitting, d-electrons transit to higher energy level by absorbing a wavelength of light within the visible region of spectrum. The colour of light absorbed is complementary to the colour observed.

107
Q

disproportionation

A

when a species is simultaneously oxidised and reduced, hence Ecell of each reaction should be calculated (red - oxid) to determine if feasible (>0) or not (<0)

108
Q

Biochemical Oxygen Demand (BOD)

A
  • The amount of oxygen that would be consumed if all the organic matters were oxidised (given oxygen) in a sample of water at a particular temperature over a period of 5 days
    ○ Or basically, how much living matter is inside
109
Q

amphoteric oxide w/ acid and base

A

with acid: salt (neutralisation)

with base: sum bs (not hydroxide)

110
Q

when showing Kc expression

A

[pdt]/[rxt]

rmb: concentration!!!! not moles

111
Q

balancing chemical equations

A 
balance all non H/O atoms
balance O by adding h2o on other side
balance H by adding H+ 
balance charges by adding e-
balance e- for both half eqns
112
Q

is neutralisation a redox reaction

A

no

113
Q

hydration energy

A

charge/radius (similar to lattice energy and charge density)

114
Q

why is reactants before slow step included in rate expression

A

cause that means slow step probably involves intermediate in slow step. include intermediate by using Kc of first fast step

115
Q

Paramagnetism

A

Paramagnetism is caused by unpaired electrons, will be attracted by magnetic field

116
Q

Diamagnetism

A

Diamagnetism is caused by paired electrons, will be repelled by magnetic field
All substances have diamagnetism as all hv paired electrons

117
Q

sample qns: what ion is likely to have a more exothermic enthalpy of hydration

A

the ion with greater charge density, stronger bond with water molecules

118
Q

standard enthalpy change unit

A

kJ, no need mol-1 cause standard is 1 mole (same for entropy)
kj = standard!!! cause its 1 mole MEANS must divide by number of moles

119
Q

why is enthalpy change calculated using Hf or Hc accurate/inaccurate

A

accurate:
no approximations were made in the cycle
data in table has small uncertainties
values are specific to the compounds

inaccurate:
values were experimentally determined, will have uncertainties
different sources will have different values for enthalpy change

120
Q

radical dot before or after atom

A

before atom, after a molecule. idk why

121
Q

physical n chemical evidence for benzene structure

A

all bond lengths are same
all bond angles are same
planar geometry

chemical
can only do nucleophilic substitution reactions not addition
benzene less exothermic when undergoing hydrogenation

122
Q

how to know what rxts at cathode given an (aqueous) electrolyte

A

for metals, can use reactivity series
PSCMAZITLHCSG
basically only copper preferred over hydrogen
less reactive will reduce

123
Q

factors that affect change in E for d-d orbital splitting

A
  1. type of metal ion
  2. oxidation state (ig its kinda like no. of electrons)
  3. type of ligand
  4. geometry (or coordination number)
124
Q

why got colourless metal complexes

A
  1. no d-d electron transition (empty/full d orbitals)

2. change in E is out of visible light region

125
Q

neutralisation with at least 1 strong component, reaction is…

A

to completion

126
Q

drastic change at start of ph graph is when?

A

when a weak component is used at the start

127
Q

when labelling voltaic cell:

A

label electrolye, salt bridge (flow of cation/anion), flow of electrons

128
Q

order of drawing 3d diagram

A

line wedge dash

129
Q

chemicals that accept electron pair (lewis acid examples)

A

BF3, AlF3, SiBr4, SiF4

130
Q

When to include state symbol in chem eqn

A

unless stated in qns, only for energetics qns

131
Q

reactivity of halogenoalkanes

A

tertiary > secondary > primary (cause sn1 is more reactive than sn2)

132
Q

why is 3 and 4 member rings so unstable?

A

bond angle less than VSEPR angle, creates angle strain

133
Q

resonance hybrid

A

charge on delocalised electrons only if molecule is charged (like ozone is neutral so only have dotted line no charge). can put charge in terms of [_]- like that

134
Q

ozone radicals formation steps

A

radicals are always neutral, instead of counting electrons just count oxygen atoms
O2 -> 2O.
O3 -> O2 + O.

135
Q

unit for pressure in pv=nrt

A

Pa

rmb volume is in m3

136
Q

oxidation of halides for electrolysis

A

bromide n iodide are always preferentially discharged for dilute/concentrated
chloride only for concentrated discharge

137
Q

why does iron form many different coloured complex ions

A

iron can form several stable ions with incomplete d-orbitals that each has a different oxidation state. The different oxidation state results in energy difference between split d-d orbitals to change, leading a different colour of light to be absorbed and observed.

138
Q

value of delta G at equilibrium

A

delta G = 0 as rate of forward reaction is equal to rate of backward reaction

139
Q

bonding in ozone

A

delocalisation of pi-electrons, both O-O bonds have bond order of 1.5 n are of equal length. stronger than single bond but weaker than double bond, both bonds are of equal strength

140
Q

bonding in ozone

A

delocalisation of pi-electrons, both O-O bonds have bond order of 1.5 n are of equal length. stronger than single bond but weaker than double bond, both bonds are of equal strength

141
Q

colour of NO2 (nitrogen dioxide gas)

A

reddish brown

142
Q

the plus sign on oxidation state/charge

A

O.S +2, charge 2+

143
Q

order of colour wheel

A

ROYGBV

roy give bj (in) vietnam

144
Q

conformational isomers

A

different relative positions of substituents by rotation about a sigma bond

145
Q

nitrogen (IV) oxide and water

A

2NO2 + H2O -> HNO3 + HNO2

to form NO2, N2/NO will react with O2

146
Q

sulfur dioxide or sulfur trioxide with water

A

SO2 (g) + H2O (l) ⇌ H2SO3 (aq)

SO3 (g) + H2O (l) → H2SO4

147
Q

why metal gives different colour flames

A

frequency of light is emitted when electrons transmit from higher to lower atomic energy level, energy difference is difference for each metal.
NOT energy difference between d-orbitals, is between subshells

148
Q

average bond enthalpy

A

average bond enthalpy is the energy needed to break one mole of a bond in a gaseous molecule averaged over similar compounds

149
Q

Frequency/pre-exponential factor A

A

Indicates frequency of collisions and probability that collisions have proper orientations

hl chem (2 decks)

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