chem Flashcards
(149 cards)
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
2
Q
lyman
A
emission to n = 1
3
Q
balmer
A
emission to n = 2 (visible light)
4
Q
paschen
A
emission to n = 3
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
6
Q
whenever you see a find Kc qns
A
ICE table damnit (values are in concentration)
7
Q
Gibbs free energy unit
A
kJmol-1 (around 3 digit, some formula gives u in Jmol-1)
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
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
10
Q
finding pH or H+
A
14 - pOH
[H+] = sqrt(Ka x (C-x))
Kw = [H+][OH-]
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)
12
Q
(polar) Aprotic solvent favoured by
A
solvates cation counterion favours sn2 (increase reactivity of nucleophile to donate a pair of electrons)
13
Q
transition metal r/s with periodic table group
A
group no. is no. of electrons in 3d + 4s orbitals
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
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
16
Q
why atoms can exceed octet rule
A
period 3 and onwards (presence of 3d orbitals)
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
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
19
Q
strong acid + strong base =
A
neutral salt
both have weak conjugate base/acid, so [H] : [OH] remains same
20
Q
weak acid + weak base
A
neutral salt
both have strong conjudate base/acid, [H] : [OH] stay same
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
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
23
Q
first order wrt a rxt
A
constant half-life
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
25
when temperature increase, what happens to rate of forward reaction (assuming exothermic rxn)
increase in rate of reaction. even though POE to left by LCP, increase in temp will increase rate of reaction anyway.
26
acidic buffer
weak acid, strong base/neutral salt of conjugate base
27
basic buffer
weak base, strong acid/neutral salt of conjugate acid
28
coordination number
no. of ligands around a central (metal) ion
29
Colours of iodine in different states
solid: black
aqueous: brown
in organic solvent: purple
gaseous: purple
30
Index of hydrogen (IHD), no. of H2 molecules to make molecule non-cyclic (at all) and saturated
1. no. of pi bonds + no. of rings
| 2. [(2n+2)-(no. of H)+(no. of grp 15)-(no. of grp 17)]/2
31
Reagent/conditions: reduction of carboxylic acid/aldehyde/ketone to alcohol
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)
32
describe a covalent bond
electrostatic forces of attraction between positively charged nuclei and one or more shared pair of electrons.
33
Hrxn
Hf(products)-Hf(reactants)
BE(bonds broken)-BE(bonds formed)
Hcmbst(rxt) - Hcmbst(pdt)
*just recall the flow cycle diagram thing
34
first order rxn
conc/time: constant half-life
| rate/conc. : linear increase
35
zero order rxn
conc/time: linear decrease, decreasing half-life
| rate/conc: constant
36
second order rxn
conc/time: increasing half life
| rate/conc: exponential increase
37
how does catalyst increases rate of reaction
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.
38
buffer eqns (finding ka
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
39
at maximum buffer capacity....
vol. of weak acid x2 of vol. of strong base
and vice versa
```
pH = pKa
pOH = pKb
```
40
amphiprotic
can act as bronsted-lowry acid and base
41
when to use reversible arrow
equilibrium qns, acid/base qns, salt hydrolysis
42
why is halogenoalkane more reactive than alkanes?
C-Halogen bond is weaker than C-H bond (except for C-F)
43
formation of NO2+
HNO3 + H2SO4 -> NO2+ + HSO4- + H2O
| by right should be reversible arrow
44
nitrobenzene R grp, intermediate R grp, phenylamine R grp
nitrobenzene: NO2
intermediate: NH3+
phenylamine: NH2
45
reference compound in NMR (name n purpose)
tetramethylsilane
all H atoms in same environment, one strong integration trace, no splitting
inert/stable
46
anode/cathode: oxidation/reduction
anode: oxidation
cathode: reduction
true for both voltaic and electrolytic cells
47
standard enthalpy of hydration
1 mole of gaseous ions form 1 mole of aqueous ions
| exothermic
48
standard enthalpy of solution
1 mole of solute is dissolved (aqueous, ions)
EC(soln) = EC (lattice) + EC(hyd)
cam be endo n exo, depending on LE
49
volume of 1 mole of gas at STP (273K, 1atm)
22.7dm^3/mol
50
volume of 1 mole of gas at SATP (298K, 1atm)
24.8dm^3/mol
51
when finding empirical formula:
if given % or mass of different elements, take % or mass divided by Mr to find mole ratio
52
for Gibbs free energy, enthalpy change, entropy change, rmb to:
put negative or positive sign
53
reagents n conditions: reduction of nitrobenzene to form phenylamine
step 1: Sn (tin) in concentrated HCl, heat under reflux
C6H5NO2 + 3Sn + 7H+ -> C6H5NH3+ + 3Sn2+ + 2H2O
step 2: NaOH (aq)
C6H5NH3+ + OH- -> C6H5NH2 + H2O
54
reagents and conditions: electrophilic substitution of nitrobenzene from benzene
concentrated HNO3, and concentrated H2SO4 catalyst
| 50 deg cel
55
rate mechanism of electrophilic substitution of nitrobenzene from benzene
```
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+
```
56
identifying mass spectrum species from mass spectrum, rmb to:
rmb to put positive charge on molecule. add up Mr of atoms in molecule
57
conjugated system of electrons
electrons are delocalised within entire system of molecules instead of within 1 molecule (graphite, graphene)
58
delocalisation
electrons are shared by more than 2 atoms in a molecule as compared to being localised between a pair of atoms
59
what is a transition metal? and properties
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
60
intermolecular forces
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
OH grp name
hydroxyl
62
resonance
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
bond order
total number of bond pairs divided by number of bonding locations
64
formal charge
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
how ligands form bonds
dative covalent bond from __ to central metal ion (and is hence also a Lewis base)
66
structural isomers
chain
positional
functional grp
67
stereoisomerism
| *rmb to draw 3d structure
```
Configurational (cis-trans, EZ)
Optical isomerism (enantiomerism, enantiomers, diastereomers, meso compounds)
```
68
EZ isomers
E is opposite side
Z is same side
when all substituents are different
69
cis-trans isomers
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
racemic mixtures
equimolar amounts of each enantiomer, no net rotation as they cancel each other out
71
diastereomers
different configuration at one or more, but not all, chiral centres from a pair of enantiomers, will still be optically active
72
meso compound
contain chiral centres but is optically inactive due to plane of symmetry
73
reagents and conditions: hydrogenation of alkenes
1. Ni catalyst at high temp (150degcel) and pressure
2. Pd (palladium) catalyst, r.t.p
3. Pt (platinium) catalyst, r.t.p
74
reagents and conditions: (alkenes) electrophilic addition of halogens, hydrogen halides, inter-halogen compounds
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
reagents and conditions: free radical substitution
```
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
reagents and conditions: electrophilic addition of steam in alkenes
industrial: H2O(g), concentrated H3PO4 catalyst, 300 degcel, 65 atm
laboratory: concentrated H2SO4, followed by boiling with H2O (l)
77
reagents and conditions: elimination of alcohols to form alkenes
1. excess concentrated H2SO4, 170-180 degcel
2. Al2O3 catalyst, 350 degcel
3. Excess concentrated H3PO4, 250 degcel
78
reagents and conditions: elimination of halogenoalkanes
1. Alcoholic KOH, heat
| 2. Alcoholic NaOH, heat
79
preparation of halogenoalkanes
1. free radical sub
2. electrophilic substitution of hydrogen halide in non-polar solvent
3. Halogenation in non-polar solvent
80
primary halogenoalkanes prefer nucleophilic sub by sn1/sn2?
sn2
81
tertiary halogenoalkanes prefer nucleophilic sub by sn1/sn2?
sn1
82
reagents and conditions: oxidation of alcohols n colour change observed
```
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
what does A in arrhenius equation account for?
frequencies of collisions between reactant molecules. it is a constant
84
enthalpy change of atomisation
energy required to form one mole of gaseous atoms from element in standard state (always endo)
85
limitation of using Bond Enthalpy to find enthalpy change of rxn
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
lattice energy enthalpy change
one mole of solid ionic compound is separated into gaseous IONS. always endothermic
87
standard hydrogen electrode
Pt electrode
1.00moldm-3 H+ (aq)
H2 gas at 100kPa (1 bar)
88
halogen standard state
fluorine - gas
chlorine - gas
bromine - liquid
iodine - solid
89
ether functional grp
R-O-R'
90
alkyne functional grp
C triple bond C
91
rate constant unit (k)
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
lewis base donate or accept e- pair?
donate
93
esterification reagents n conditions
alcohol+carboxylic acid, concentrated H2SO4 catalyst, heat under reflux
94
carbonyl functional grp
C=O
95
naming esters
Name the alcohol part before the acid part, replacing "oic acid" with "oate"
96
how to determine sp1, sp2, or sp3 (hybridisation)
look at number of electron domains that the atoms has
97
calculate change in entropy
S(pdt) - S(rxt)
98
why does Kc change with temperature change?
(assuming temeprature rise)
| favour endo, POE shifts to the left, backward rate more than forward rate, [rxt] more than [pdt], so Kc reduces
99
when doing electrolysis
always just write out overall eqn, where all numbers are integers, so can find n (no. of electrons) for change in G
100
sn1 stands for
unimolecular nucleophilic substiution
101
sn2 stands for
bimolecular nucleophilic substitution
102
when drawing nucleophilic substiution,
always freaking use 3d diagram pls
103
nmr use what em wave
radio waves
IR - infrared
Uv vis - UV bruh
104
forming different lewis structures to test FC
- electron pushing method
- can be enantiomers (just mirror)
- just shift the double bond around
105
Difference Between Formal Charge and Oxidation State
- 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
why are transition metals coloured
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
disproportionation
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
Biochemical Oxygen Demand (BOD)
- 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
amphoteric oxide w/ acid and base
with acid: salt (neutralisation)
| with base: sum bs (not hydroxide)
110
when showing Kc expression
[pdt]/[rxt]
| rmb: concentration!!!! not moles
111
balancing chemical equations
```
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
is neutralisation a redox reaction
no
113
hydration energy
charge/radius (similar to lattice energy and charge density)
114
why is reactants before slow step included in rate expression
cause that means slow step probably involves intermediate in slow step. include intermediate by using Kc of first fast step
115
Paramagnetism
Paramagnetism is caused by unpaired electrons, will be attracted by magnetic field
116
Diamagnetism
Diamagnetism is caused by paired electrons, will be repelled by magnetic field
All substances have diamagnetism as all hv paired electrons
117
sample qns: what ion is likely to have a more exothermic enthalpy of hydration
the ion with greater charge density, stronger bond with water molecules
118
standard enthalpy change unit
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
why is enthalpy change calculated using Hf or Hc accurate/inaccurate
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
radical dot before or after atom
before atom, after a molecule. idk why
121
physical n chemical evidence for benzene structure
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
how to know what rxts at cathode given an (aqueous) electrolyte
for metals, can use reactivity series
PSCMAZITLHCSG
basically only copper preferred over hydrogen
less reactive will reduce
123
factors that affect change in E for d-d orbital splitting
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
why got colourless metal complexes
1. no d-d electron transition (empty/full d orbitals)
| 2. change in E is out of visible light region
125
neutralisation with at least 1 strong component, reaction is...
to completion
126
drastic change at start of ph graph is when?
when a weak component is used at the start
127
when labelling voltaic cell:
label electrolye, salt bridge (flow of cation/anion), flow of electrons
128
order of drawing 3d diagram
line wedge dash
129
chemicals that accept electron pair (lewis acid examples)
BF3, AlF3, SiBr4, SiF4
130
When to include state symbol in chem eqn
unless stated in qns, only for energetics qns
131
reactivity of halogenoalkanes
tertiary > secondary > primary (cause sn1 is more reactive than sn2)
132
why is 3 and 4 member rings so unstable?
bond angle less than VSEPR angle, creates angle strain
133
resonance hybrid
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
ozone radicals formation steps
radicals are always neutral, instead of counting electrons just count oxygen atoms
O2 -> 2O.
O3 -> O2 + O.
135
unit for pressure in pv=nrt
Pa
| rmb volume is in m3
136
oxidation of halides for electrolysis
bromide n iodide are always preferentially discharged for dilute/concentrated
chloride only for concentrated discharge
137
why does iron form many different coloured complex ions
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
value of delta G at equilibrium
delta G = 0 as rate of forward reaction is equal to rate of backward reaction
139
bonding in ozone
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
bonding in ozone
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
colour of NO2 (nitrogen dioxide gas)
reddish brown
142
the plus sign on oxidation state/charge
O.S +2, charge 2+
143
order of colour wheel
ROYGBV
| roy give bj (in) vietnam
144
conformational isomers
different relative positions of substituents by rotation about a sigma bond
145
nitrogen (IV) oxide and water
2NO2 + H2O -> HNO3 + HNO2
to form NO2, N2/NO will react with O2
146
sulfur dioxide or sulfur trioxide with water
SO2 (g) + H2O (l) ⇌ H2SO3 (aq)
| SO3 (g) + H2O (l) → H2SO4
147
why metal gives different colour flames
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
average bond enthalpy
average bond enthalpy is the energy needed to break one mole of a bond in a gaseous molecule averaged over similar compounds
149
Frequency/pre-exponential factor A
Indicates frequency of collisions and probability that collisions have proper orientations
