chem random Flashcards
(137 cards)
1
Q
exo/endo energy profile diagrams
A
exo: reactants energy > products
enthalpy change arrow points down
as exo releases heat
endo: products energy > reactants
enthalpy change arrow points up
as energy is absorbed
2
Q
buffer equilibrium explanation
A
H2CO3 <-> H+ + HCO3-
increase in H+ reacts with salt
H+ + HCO3- -> H2CO3
equilibrium shifts left
increase in OH- reacts with alkali
H2CO3 + OH- -> HCO3-
(H+ + OH- -> H20)
equilibrium shifts right
3
Q
why would measured pH be different to calculated pH
A
dissociation is not negligible
or
large Ka
idk what that means either
4
Q
why is MgCl2 lattice enthalpy more exothermic than CaCl2
A
magnesium ion smalle than calcium ions so greater charge density and attraction between magnesium ion and chloride ion greater
(stronger ionic bond = more exothemric)
5
Q
why is enthalpy change of hydration of F- more exothermic than Cl-
A
F- is a a smaller sized ion so has greater attraction to water
(generally as charge of ion increases and atom size decreases - enthalpy change is more exothermic)
6
Q
TMS
A
standard for chemical shift measurements
7
Q
benefits of single stereoisomers in the synthesis of drugs
A
fewer side effects
increased pharamacological effectiveness
reduced costs in seperating stereoisomers
idk
8
Q
qualitative analysis test equations
A
carbonates test:
CaCO3 + 2H+ -> Ca2+ + CO2 + H20
sulfates test
Ba2+ + SO42- -> BaSO4
ammonium test
NH4+ + OH- -> NH3 + H20
9
Q
integration and splitting
A
integration = number of H per environment
splitting = hydrogens of any adjacent carbons + 1
CH3-CH2
CH3 carbons splitting is 2+1 so triplet?
10
Q
transition elements
A
ions with incomplete d subshell
11
Q
why would a mechanism not happen in one step
A
rate equation doesnt match overall equation
collision unlikely with more than 2 ions
12
Q
chromatography
A
balance between solubility in moving phase and retention of stationary phase
13
Q
naming esters
A
Esters are named by combining the alkyl part derived from the alcohol and the carboxylate part derived from the carboxylic acid
Example:
Ethanol (alcohol) + Propanoic acid (carboxylic acid) = Ethyl propanoate (ester)
ethyl-3-bromopropanoate
alcohol has higher priority than halogen but halogen higher priority over enoate
number halogen based on carbon its on NOT stick its on
14
Q
why are esters biodegradeable
A
ester bond can be hydrolysed with water
c-c are non polar and harder to break
15
Q
why do polyamides have a higher bp than polyalkenes
A
polyamide = H bonding
polyalkene = london forces
H bonds stronger in polyamides
16
Q
amino acids have varying Rf values
A
amino acids have different polarities so have different affinities for mobile and stationery phase
detect with ninhydrin
17
Q
how is K in rate affected by T
A
as Temp increases, value of K increases
18
Q
acid hydrolysis rules
A
break amides + esters -> alcohol
protonate NH if its at the end to NH3+
middle to NH2 (unbroken)
19
Q
determining purity
A
recrystalise and determine melting point and compare values to known library
pure substances have sharp melting points
20
Q
recrystallisation techniques
A
Recrystallisation
Dissolve impure solid in minimum volume of hot solvent
filter solution while hot to remove insoluble impurities
allow to cool so compound crystallises out of solution
filter cool mixture under reduced pressure to remove most of the soluble impurities
Wash with cold solvent to remove any insoluble impurities
and dry
21
Q
what happens to pH of buffer solution if volume increases slightly
A
pH constant as the ratio of [HA]/[A-] is the same
22
Q
is dissociation of water exo or endo?
A
endothermic as Kw increases with temperature
23
Q
why can a reaction be exothermic
A
More energy is released by forming bonds
than energy required when breaking bonds
24
Q
calculating formation from combustion (vice versa)
A
reverse the sign of the formation data and then plug into cycle
change signs according to direction of the cycle
25
average bond enthalpy
enthalpy change require to break one mole of gaseous covalent bonds
26
activation energy of reverse from enthalpy change and Ea of forward
Ea (reverse) = Ea(forward) - enthalpy change
27
CH3COCl produces most acidic solution
forms HCL in water so strong acid
28
electrophile vs nucleophile
electrophiles are electron PAIR accepters and nucleophiles are electron pair donors
29
directing groups
NO2 - 3,5 directing
NH2 - 2,4 directing
COOH electron withdrawing group
30
drawing wedged lines
always solid (colour them in)
31
factors affecting the values of lattice enthalpies
Decrease in (ionic) size
more negative LE/more exothermic
Increase in (ionic) charge /charge density
more negative LE OR more exothermic
Greater attraction between ions gives more negative LE
eg:
Cl– is larger (than F–)
Mg2+ is smaller (than Na+) AND Mg2+ has a greater charge (than Na+)
F– has greater attraction for Na+ ions
AND
Mg2+ has greater attraction for F– ions
32
The change that produces lattice enthalpy is spontaneous with negative entropy change
Why is this change able to take place spontaneously?
∆H is more negative than T∆S
33
esters
always form alcohol to carboxylic acids (cant have 2 carboxylic acids)
in a condensation reaction
34
addition polymerisation
break the c=c bond and make it c-c with groups off it (straight) and directly connect to next one (No condensation)
35
solubility of esters
if one compound has 2 ester bonds then say
hydrogen bonds form with water
it forms MORE hydrogen bonds with water
as there are more OHs
36
reduction via NaBH4
if there are multiple carbonyls reduce all of them to alcohols and use 4[H] for 2 carbonyls etc
37
retention time
The time taken (from the injection of the sample) for the component to leave the column
38
dehydration or elimination of H20 from alcohol
alcohol group removed and alkene double bond forms
but doesnt affect carboxylic acid
form of addition polymerisation
39
ester catalyst
can only be produced when a carboxylic acid and alcohol are heated under presence of sulfuric acid - acid catalyst
40
stereoisomer
same structural formula, different spatial arrangement
41
phenyl methanol
benzene + CH3 + OH
42
standard electrode potential definition
The EMF/voltage/potential difference of a half-cell compared with a standard
hydrogen half-cell
Standard conditions
Temperature of 298 K
AND concentrations of 1 mol dm–3
AND pressure of 101 kPa
43
why would reactions predicted from E values not take place
high activation energy or non standard conditions
44
methanoic acid as a fuel cell over hydrogen
Methanoic acid is a liquid AND easier to store
OR hydrogen is a gas AND harder to store/transport
2. Hydrogen is explosive/more flammable
3. HCOOH gives a greater cell potential/voltage
4. HCOOH has more public/political acceptance than
hydrogen as a fuel
45
hydrogen fuel cells
2H2(g) + O2(g) → 2H2O.
A fuel cell converts energy from reaction of a fuel with oxygen into a voltage/electrical energy
can be stored under pressure as a liquid
46
nitrate to amine equation
NO3 + 6[H] -> NH2 + 2H2O
47
systematic name
use oxidation states eg
NaNO2
- sodium nitrate (III) as nitrogen has an oxidation state of 3
48
determining solvent used in recrystallisation
compound has to be soluble in hot solvent but insoluble in cold solvent
49
differences and similarities in bonding of kekule and benzene
similarities:
both have sideway overlap of p orbitals
both have pi bond system above and below plane
differences:
kekule has localised alternating pi bonds but benzene has delocalised ring of pi electrons
50
evidence supporting delocalised model
Bond length
(C–C) bond length is between single (C–C) and double bond (C=C)
ΔH hydrogenation less (exothermic) than expected
Benzene is less reactive than alkenes
OR bromination of benzene requires a catalyst/halogen carrier
51
reaction of sodium carbonate vs NaOH
Na2CO3 only reacts with carboxylic acids to form a salt (not alcohol)
NaOH reacts with alcohol and carboxylic acid to produce salt (neutralisation reaction)
52
acid anhydride reaction
acid anhydride more likely to react with alcohol and form molecule + carboxylic acid by splitting the ethanoic anhydride
eg:
salicylic acid (benzene + COOH + OH)
salicylic acid + ethanoic anhydride
-> aspirin (carboxylic acid group intact and half the enthanoic anhydride reacts with OH) and ethanoic acid
53
distillation vs reflux
distillation allows partial oxidation as volatile components can escape
reflux allows complete oxidation as volatile components cannot escape
54
polarity
shorter carbon chains are less polar than longer ones
OH = more polar
55
acid + heat to an alcohol
removes OH and forms double bond
56
PV=NRT units
p = pa
v = m3
n = mol
R = 8.31 jkmol
T = k
57
alkane + amine naming
amino-alkane
eg:aminopropane
or 1,3diaminopropane
58
how does fe2+ transport o2 around the body
iron 2+ found in haemoglobin
oxygen forms a coordinate bond with iron(II)
when required O2 is replaced with CO2 via ligand substitution
CO forms stronger bonds than O2 so toxic
59
oxidation num of cis platin platinum
nh3 charge = 0
oxidation num of cis platin platinum is +2
60
role of cis platin for treating cancer
platin binds to DNA of cancer cell stopping replciation
61
pressures influence on Kp
Equilibrium shift to right
Effect of increased pressure on Kp expression:
Ratio (in Kp expression) decreases
as Denominator/bottom of Kp expression increases more (than numerator/top)
Equilibrium shift (Kp expression):
Ratio (in Kp expression) increases to restore Kp and Numerator/top of Kp expression increases to restore Kp
62
thiosulfate redox indicator
add starch
solution goes from blue black to colourless as iodine is used up
63
electrode potentials include:
include any ions that are aqueous in cells
+ always mention which E value is smaller and why it would shift to the left then
64
effect on cell potential if concentration of Ag+ increases
Ag+ + e- -> Ag(S)
Ag+ becomes more positive so e cell becomes smaller
as products are favoured so more electrons are being transported increasing voltage
65
sign of electrodes when oxidation or reduction occurs
Ni electrode is reduced = +ve
reduction - gain of e-
Ni electrode is oxidised = -ve
oxidation - loss of e-
66
why would the measured cell potential slowly change
concentrations of ions change
67
pH of hydrogen half cell
0
pH would decrease as current is delivered as H redox system is more negative so releases electrons
equilibrium shifts to increase H+ conc decreasing pH
68
hydrogen half cell (SHE)
2H+(aq) + 2e⁻ → H₂(g)
draw with platinum electrode surrounded by a system of hydrogen gas in a solution of H+
69
dichromate under reflux
Green solution Cr3+
Orange solution Cr2O7
70
always include charge on transition metal complexes
and on formulae of complexes!!!!!
71
. Explain the use of two deuterated compounds in NMR spectroscopy
CDCl3 used as a solvent (no hydrogens)
D2O used to identify OH OR NH
protons ✓
72
why do primary amines have higher BP than tertiary amines
tertiary has more branching so less surface area of contact, so weaker london forces
primary and secondary amines have hydrogen bonds which are stronger than london forces
so less energy required to overcome intermolecular forces
73
if an NMR question says "singlet at 9.5ppm) say
Peak at (δ) 9.6 shows H–C=O
AND
No H on adjacent C atom as peak
is singlet
ALWAYS MENTION SINGLETS EFFECT ON ADJACENT HYDROGENS
74
reduction of an aldehyde equation
CH3(CH2)3CHO + 2[H] →
CH3(CH2)3CH2OH
75
State the region of the electromagnetic spectrum used in 1H NMR spectroscopy
radio waves
76
which compound is used as a standard in NMR
TMS - tetramethylsilane
Si(CH3)4
77
gdm3 -> moldm3
= divide by mr
78
incomplete combustion products
CO + H2O (no CO2)
79
acyl chloride reactions
acyl chloride + alcohol = ester
acyl chlorides dont react with carboxyl group - only alcohol
acyl chloride + water = carboxylic acid
80
amino acids in polar solvents affect on retention time
more polar solvents mean amino acids interact more and moves up stationary phase so shorter retention time and larger rf
idk why
81
if gibbs question involves standard free energy change without temp
use a temperature of 298k (25 degrees) NOT 273
82
dynamic equilibrium
the rate of the forward reaction is the same as the rate of the backward reaction in a closed system, and the concentrations of the reactants and products are constant.
83
enthalpy change of atomisation definition + iodine equation
½ I2(s) → I(g)
the enthalpy change when one mole of gaseous atoms are formed from an element in its standard state
84
effect of pressure or concentration on Boltzmann distribution
no change
85
enthalpy change of hydration
) 1 mole of gaseous ions OR 1
mole of hydrated ions / aqueous ions
gaseous ions forming aqueous / hydrated ions
86
why is entropy positive in solution enthalpies
aqueous particles are more disordered than solid
solid -> aqueous - increasing disorder
87
The enthalpy and entropy changes of a reaction both have a negative sign.
Discuss how the feasibility of this reaction will change as the temperature increases.
reaction less feasible and G increases
so NEGATIVE TS becomes more positive
ALWAYS MENTION THE NEGATIVE
88
why can an aromatic ring not form more than one E/Z isomer
ring would be strained and break down as you cannot form a ring if high priority groups are on opposite sides (only forms Z)
89
naming aldehydes
butanAL
90
4 reaction mechanism and examples
nucleophilic substitution =alcohol -> haloalkane
nucleophilic addition is carbonyl mechanism
electrophilic substitution = friedal crafts
electrophillic addition = double bond + HBr
91
when to purify using dissolving in solvent and when to separate organic and aqueous layers of a separating funnel
sperating funnel for when product is insoluble in water and you just need to separate
dissolving in minimum volume of solvent to remove impurities (always say for purification )
ask Miss
92
is enthalpy change of solution and hydration exo or Endo
solution = 1 mol of an ionic compound dissolves entirely in water= exo or Endo so check before doing a cycle
hydration = 1 mol of gaseous ions dissolve in water = always exo
93
enthalpy change of solution equations if temperature of the water increase is reaction Endo or exo
exo as the water temp has increased as the solutes have lost energy and given it to the water
so negative value
94
colours of halogens in solid
chlorine = green yellow
bromine = orange brown
iodine = shiny grey
95
colours of halogens in aqueous solution
chlorine = pale green / colourless
bromine = orange
iodine = brown
96
colours of halogens in cyclohexane
chlorine = pale green
bromine = orange
iodine = violet
97
hydration enthalpy cycles remember
Endo -L (on the left
exo -L + S (exoLs)
98
why do hydration enthalpies become less exothermic down group 1
size of ion increases/charge density decreases
electrostatic force of attraction between metal ion and delta negative oxygen of water decreases
99
whys it difficult to predict whether the enthalpy change of solution becomes more or less exo from MgF2 to MgI2
halide ion gets larger down the group
lattice enthalpy gets less exothermic down group (less attraction for Mg2+)
hydration enthalpy less exothermic down group (less attraction of halide to water)
difficult to predict whether lattice enthalpy or hydration enthalpy has a bigger effect
100
bond enthalpy equation
reactants - products
101
predict fastest rate of hydrolysis of bromo/iodo/chloro alkanes
I>Br>Cl
C-I is the weakest bond
C-Cl is a stronger bond
102
alkaline hydrolysis of a haloalkane
halo alkane + OH- -> alcohol instead of halogen and halogen ion
103
features of a dynamic equilibrium
occurs in a closed system
rates of forward and backwards reaction are equal
conc of products and reactants are equal
104
exceptions to electron config filling rule
Chromium (Cr):
Instead of [Ar] 4s² 3d⁴, it is [Ar] 4s¹ 3d⁵
The 4s² 3d⁴ configuration would be less stable, as it leaves the d-subshell with an incomplete half-filled configuration
Copper:
Instead of [Ar] 4s² 3d⁹, it is [Ar] 4s¹ 3d¹⁰
(Still a TM as Cu2+ has incomplete)
The 4s² 3d⁹ configuration would be less stable, as it would leave the d-subshell with an incomplete filling.
105
purpose of HNO3 in qualitative analysis
reacts with carbonate / CO3 2−
AND
carbonate forms a white ppt
106
copper + sulfuric acid equation
Cu + 2H2SO4 → CuSO4 + SO2 (g) + 2H2O
107
why are sc and zn not TM
Scandium / Sc and zinc / Zn are not transition elements
Electron configurations of ions:
Sc3+ AND 1s2 2s2 2p6 3s2 3p6 \
Zn2+ AND 1s2 2s2 2p6 3s2 3p6 3d10
Sc3+ AND d sub-shell empty
Zn2+ AND d sub-shell full
108
oxidation num of monoatomic ions
same as charge
eg: Na+ = 1+
109
copper 2+ + HCl equation
[Cu(H2O)6] 2+ + 4Cl− → [CuCl4] 2− + 6H2O
110
in calorimetry, how is temp and enthalpy change of solution affected by doubling mass
temperature increase doubles (larger temp increase to heat on a larger scale)
enthalpy change of solution value remains constant
SAME ENERGY RELEASED PER MOL
(its expressed per mol so solute mass doesn't affect it)
111
enthalpy change of hydration
enthalpy change required for 1 mol of gaseous ions to form 1 mol of aqueous hydrated ions
112
appropriate num of significant figures
you should use the smallest number of significant figures among the numbers used in the calculation
113
in calorimetry, how is temp and enthalpy change of solution affected by doubling VOLUME
ΔT is less
AND
ΔsolH is the same
Reason for ΔT less
(same) energy/heat spread over larger volume (of
water)
Reason for ΔsolH same
Same energy released per mole of H2SO4
114
carbonate test - (first)
add HNO3 - nitric acid forming effervescence from CO2
2H+ + CO3 2– → CO2 + H2O
115
sulfate test (second)
Add Ba(NO3)2 - barium nitrate
white ppt = sulfate
Ba2+ + SO4 2– → BaSO4
116
solubility of halide ppt
AgCl = soluble in dilute ammonia
AgBr = soluble in concentrated ammonia
AgI = insoluble in concentrated ammonia
117
ammonium test
heat with NaOH
gas turns damp red litmus blue
NH4+ + OH− → NH3 + H2O
118
CO2 bonding and structure
simple molecular
weak London forces
less energy require to overcome intermolecular forces than covalent
119
why is magnesiums melting point higher than sodium's
Magnesium has more outer electrons
Magnesium ions have a greater (positive) charge density
Magnesium has a greater attraction between ions
and delocalised electrons
120
why is phosphorous mp greater than chlorine (P4 vs Cl2)
phosphorous has more electrons
so stronger London forces
so more energy required to overcome intermolecular/London forces
121
silicon
giant covalent with covalent bonds BETWEEN ATOMS
122
plotting successive ionisation Energies
aluminium eg:
1s2
2s2 2p6
3s2 3p1
(lines are jumps)
aluminium has large increase in IE at 3 to 4 (4th in 2 not 3)
and 11 to 12
123
ionic vs covalent bond
Ionic bonds typically form between metals and nonmetals, while covalent bonds form between two or more nonmetals
124
Potassium is placed immediately after argon in the periodic table - why?
Potassium (atoms) have one more proton (than
argon)
as periodic table arranged by atomic number (proton num)
125
halogen vs halide layer
halogen in organic layer and halide ion in aqueous layer
so brown to violet means
Br2 + 2I– → 2Br– + I2
126
bleach disproportionation equation
Cl2 + 2NaOH -> NaClO + NaCl + H20
conditions = COLD DILUTE NaOH
127
two routes to make Ba(OH)2
add water:
Ba + 2H2O → Ba(OH)2 + H2
add oxygen then water
2Ba + O2 → 2BaO
BaO + H2O → Ba(OH)2
128
barium nitride + water equation
Ba3N2 + 6H2O → 3Ba(OH)2 + 2NH3
129
electronegativity
ability of an atom to attract a pair of electrons in a covalent bond
130
anomalous properties of ice
ice less dense than water
as molecules are held in an open lattice structure held apart by hydrogen bonds
ice has a relatively high melting point
hydrogen bonds are stronger than other intermolecular attractions so more energy needed to overcome hydrogen bonding
131
why are molecules polar
difference In electronegativities between atoms
molecule is not symmetrical and dipoles don't cancel out
132
activation energy form lnk gradient
Ea = -gradient * R
133
factors affecting KP
ONLY TEMPERATURE not pressure
134
why can a molecule such as SO3 have polar bonds but not molecules
differences in electronegativities between sulfur and oxygen
SO2 has lone pair so non linear and asymmetrical so dipoles cant cancel
SO3 has dipoles that can cancel
135
bond enthalpy equation
reactants - products
136
why can a reaction be endothermic
more energy is required for bond breaking than released by bond making
so products have more energy (less energy released)
137
why can a reaction be exothermic
more energy is released in bond formation than required in bond breaking
