Revision Flashcards
1
Q
Alcohol series
A
Contains functional hydroxyl (OH) group
2
Q
Organic acids
A
Contains functional carboxylic group (COOH)
H-O-C=O
3
Q
In acidic medium S2O32- will convert into?
A
S4O62-
4
Q
Buffers
A
A buffer solution is an aqueous solution consisting of a mixture of a weak acid and its conjugate base, or vice versa. Its pH changes very little when a small amount of strong acid or base is added to it.
5
Q
2KMnO4(s) →
A
→ K2MnO4(s) + MnO2 (s) + O2(g)
6
Q
What setup is used to find the molar mass of Mg
A
Eudiometer
7
Q
named the fundamental unit carrying electricity as
“electrons”
A
Johnstone G. Stoney
8
Q
calculated a value of 1.76 x 10*8 coulombs per gram (C g-1) for the ratio of the electron’s electrical charge to its mass.
A
J. J. Thomson
9
Q
succeeded in measuring the charge of an electron as 1.602 x 10-19 C by performing the oil drop experiment
A
Robert Millikan
10
Q
experimentally proved the existence of positive charges in matter.
A
Eugen Goldstein
By using a perforated cathode
11
Q
showed that radioactive materials produce three types of emissions
A
Lord Ernest Rutherford
12
Q
Who were Rutherford’s associates
A
Geiger and Marsden
13
Q
discovered the ‘neutron
A
Sir James Chadwick
14
Q
found that the number of positive charges on the nucleus increases in atoms by single electron units
A
Jeffrey Moseley
15
Q
Discovered the first isotopes (of neon) using mass spectrometer
A
J. J. Thomson and Francis William Aston
16
Q
Wavelength range of visible range
A
400-750 nm
17
Q
Wave particle duality
A
Louis de Broglie
18
Q
ionization energy of Na
A
496 kJmol-1
19
Q
Hund’s rule
A
The Hund’s rule states that for degenerate orbitals, the lowest energy is attained when the number of electrons having the same spin is maximized.
20
Q
published nearly identical periodic tables
A
Dmitri Ivanovich Mendeleev in Russia and Lothar Meyer in Germany
21
Q
16th group elements are also known as
A
Chalcogens
22
Q
Who put forwarded VSEPR theory
A
Gillespie and Nyholm
23
Q
zincate
A
ZnO2 2-
24
Q
aluminate
A
AlO2ˉ
25
chlorite
ClO2ˉ
26
chlorate
ClO3ˉ
27
phosphate
PO4 3-
28
HClO
hypochlorous acid
29
HClO2
chlorous acid
30
HClO3
chloric acid
31
HClO4
perchloric acid
32
Standard enthalpy change of combustion
It is the enthalpy change when one mole of an element or a compound in the standard state undergoes complete combustion with excess oxygen (or one may say air) in the standard state to give the products in the standard state
33
Standard enthalpy change of bond dissociation
It is the enthalpy change when a gaseous compound in the standard state undergoes dissociation to gaseous atoms or components in the standard state by breaking a mole of bonds.
CH4(g) = CH3(g) + H(g)
H2(g) = 2H(g)
34
Standard enthalpy change of neutralization
It is the enthalpy change when a mole of an aqueous H+ ions and a mole of an aqueous OH- ions in the standard state react to form a mole of liquid water
35
Standard enthalpy change of solvation
It is the enthalpy change when a mole of gaseous ions in the standard state changes into a solution to form 1.0 mol dm-3
36
Standard enthalpy change of hydration
It is the enthalpy change when a mole of gaseous ions in the standard state changes into a solution with water with the concentration of 1.0 mol dm-3
Na+(g) + water Na+(aq)
37
Standard enthalpy change of dissolution
It is the enthalpy change when a mole of a substance in the standard state dissolved in a solvent to form a solution with the concentration of 1.0 mol dm-3
NaCl(s) + water NaCl(aq)
38
Standard enthalpy change of sublimation
It is the enthalpy change when a mole of a solid element or a mole of a solid compound in the standard state converts completely into a gas at its standard state
39
Standard enthalpy change of evaporization
It is the enthalpy change when a mole of a liquid compound or an element in the standard state converts into a mole of gaseous compound or element at its standard state
40
Standard enthalpy change of fusion
It is the enthalpy change when a mole of a solid compound or an element in the standard state converts into a mole of liquid compound or element at its standard state.
41
Standard enthalpy change of atomization
It is the enthalpy change when an element in the standard state converts into a one mole of gaseous atoms at the standard state
½ Cl2(g) = Cl(g)
42
Standard enthalpy change of first ionization,
It is the enthalpy change when a mole of a gaseous mono-positive ions at standard state are formed by removing an electron from each atom that is most weakly bonded to the nucleus from a mole of gaseous atoms of an element in standard state
43
Standard enthalpy change of electron gain
It is the enthalpy change when a mole of gaseous mono-negative ions are formed by gaining electrons to a mole of gaseous atom at the standard state
44
Standard lattice dissociation enthalpy
It is the enthalpy change when one mole of a solid ionic compound is converted to its gaseous positive and negative ions at the standard state
NaCl(s) = Na+(g) + Cl-(g)
45
Test to identify water
Anhydrous copper sulfate is used which turns blue in the presence of water.
46
Increasing order of radius of isoelectronic species
Cations < neutral < anion
47
Soluble s block carbonate
BeCO3
All others are insoluble
48
Intensive and extensive properties
An intensive property does not depend on the system size or the amount of material in the system. E.g: density, hardness, temperature
extensive properties such as the mass, volume and entropy of systems are additive for subsystems
49
Special facts to memorize during identification of gases in an compressibility vs pressure graph
He like nobel gases will be close to the ideal gas line
NH3 like compounds having strong secondary interactions will show an abnormal deflection towards -Z
CH4, CO like compounds show a negative deviation at 1st reaching a minimum value and starts increases continuously crossing the Z=1 line at 1 point
50
ΔH(reaction) =
ΔH(bond dissociation) + ΔH(bond formation)
51
Excess NH3 + Cl2
N2 (g) + NH4Cl (s)
52
Elements showing deviations from the aufbau principle
Cr, Cu, Rh
53
NaNO2 + NH4Cl =
NaCl + N2 + H2O
54
Experiment to show the presence N2 in air
Burn a piece of Mg(s) in air, a red solid compound Mg3N2 is formed. Then add water, Mg3N2 + H20 = Mg(OH)2 + NH3
To confirm the evolved gas is NH3, add Nessler's reagent and observe the brown colour.
55
Which chloride gives blue colour to the flame test
CsCl(s)
Cs+
56
Standard experiment to identify NH3
When a filter paper is dipped in Nessler's reagent and introduced to the gas, the yellow coloured Nessler's reagent turns brown.
57
Boiling point of inorganic compound X which exists as a colourless liquid at room temperature is lower than that of water. When exposed to sunlight, X easily decomposes liberating a gas
H2S
58
H2S vs H20 - Boiling point
H2S < H2O
59
Melting point of metals depend on the
Number of unpaired d orbitals
Higher the number of unpaired d orbitals, higher will be the delocalisation of these electrons resulting in high melting points.
60
K2MnO4 colour
Green
61
6NaOH + 4S =
Na2S2O3 + 2Na2S + 3H2O
62
What do you mean by disproportion reaction
Same compound undergoes both oxidation and reduction
E.g: 2H2O2 → 2H2O + O2
63
Elements forming the bond with the highest ionic character
Should contain the lowest covalent character
CsF
64
Across a period which elements will have highest melting point
Group 14
Due to the formation of covalent bond network structures
65
5 oxides of Nitrogen
N20, NO2, N2O3, N2O4, N2O5
N2O3, N2O4, N2O5 - dimer structures of nitrogen
66
Vinegar like smell
CH3COOH (Acetic acid)
67
H2S gas test
Smells like rotten eggs, turns
lead acetate paper black.
68
SO2 gas test
Smells like burning sulphur (burnt matches)
turns acidified potassium dichromate paper green
69
HCl gas test
Pungent smell, white fumes with ammonia, white ppt with silver nitrate solution.
70
NH3 gas test
Characteristic smell, turns Nessler’s solution brown.
71
NO2 gas test
Reddish brown, turns ferrous sulphate(blue green) solution black brown.
72
Cl2 gas test
Greenish yellow, turns starch iodide paper blue.
73
Br2 gas test
Reddish brown, turns starch paper orange yellow.
74
I2 gas test
Dark violet, turns starch paper blue.
75
Define critical temperature
Temperature at and above which the vapour of the substance cannot be liquefied, no matter how much pressure is applied.
76
Critical pressure
Is the pressure required to liquefy a vapour at critical temperature
77
Compounds formed when group 1 and 2 metals react with aqueous halides
Forms aqueous colourless solutions with halides
78
TiO2
white
79
CuCl(s)
White in colour
Used as a catalyst
80
Ionic lattices when fused conducts electricity by
Existence of positive and negative ions
Metals conduct electricity due to the presence of mobile electrons
81
In basic medium orange colour K2Cr2O7 reduces into
Yellow colour K2Cr2O4
82
Ability of hydrolysis depends on
Covalent character of compounds
Higher the covalent character higher the ability to be hydrolysed.
83
Although N2 is a inert gas when Al is combusted in air, N2 reacts with Al
Energy released during the combustion of Al with O2 is used to break the bonds between O2
84
Which metal sulfate can be used in purification of water
Al2(SO4)3
85
Na2S203 + H2SO4
Na2SO4 + SO2 + S + H2O
86
Relationship between dissolution, hydration and lattice enthalpy
Dissolution = hydration - lattice enthalpy
87
S + NaOH
Na2S + Na2S2O3 + H20
88
H2O2 + PbS
PbSO4 + H2O
89
SO2 + H2S
S + H20
90
(NH4)2S ---->
NH3 + H2S
91
Al2O3 + NaOH
NaAlO2 + H20
92
For a alkyl group connected to a benzene ring to be oxidised it should have
at least 1 H atom connected to it
Tertiary alkyl groups do not get oxidised.
93
Which group 2 metal gives its peroxide with excess oxygen
Ba
94
BaSO4
White ppt
BaSO4is extremely insoluble in water, alkalies, or acids, but is slightly soluble in hot, concentrated sulfuric acid.
95
BaCO3
White ppt
Barium carbonate is soluble in acid, including dilute acetic acid, in strong bases, and in aqueous ammonia.
96
BaNO3
White ppt
Soluble in hot water
97
which hydroxide is used to identify CO2
Ca(OH)2
98
what is used in the manufacture of baby soap
NaOH and *KOH
99
Ag -White ppt that turns black on standing
AgS203 (white solid)--------> Ag2S (black solid)
100
Yellow ppt of Ag dissolves in dil.nitric acid as well as in ammonium hydroxide
AgCrO4
101
BaSO3
White ppt
Dissolves in di.HCl forming colourless, pungent smelling gas SO2
102
PbSO4
PbSO4 is white and insoluble in HNO3 and hot water.
103
Pb - Yellow ppt that dissolves in hot water
PbI2
104
Which radicals are not formed
Hydrogen radicals
105
Most abundant s block metals on earth's crust
Na and K (In order)
106
Most abundant element occurring in the oceans is
Na, Mg, Ca, K
107
What would you expect to observe when a dil.NH3 solution is added to an aqueous solution of Mn2+ and then left exposed to air
Pale pink solution makes white precipitate
This precipitate turns brown or blackish brown on exposure to air.
108
Identification of Pb2+
Add K2CrO4 to a solution, yellow ppt is formed which is soluble in NaOH
Add KI to a solution, yellow ppt which dissolves on boiling and precipitate again as golden spangles on cooling
109
yellow ppt which dissolves on boiling and precipitate again as golden spangles on cooling
PbI2
110
Identification of NO 3-
Add conc.H2SO4 and heat, reddish-brown gas is formed
To a solution add freshly prepared solution of FeSO4 followed by conc.H2SO4 down the sides of the test tube or brown test is formed.
111
Why Mg,Be do not produce colours under flame test
When excited electrons falls back to ground state, energy is released as photons but this energy in not in the visible region (UV region)
112
Variation of reactivity of 1st group elements down the group (reason)
Increases down the group
When going down the group no. of energy levels increases causing the shielding effect to increase. This results in the decrease in effective nuclear charge decreasing 1st ionisation energy)
113
Points to identify K
Lilac colour to the fame test.
Formation of superoxide.
Used in the production of baby soap.
114
Na + H2O(g)
Na2O + H2
115
Na2O2 + H2O
NaOH + H2O2
116
KO2 + H2O
KOH + H2O2 + O2
117
Only group IIA element forming peroxide with O2
Barium
118
Mg + H2O(g)
MgO + H2
119
Be + H2O(g)
BeO + H2
120
Na is used in the production of
Yellow bulb / Sodium lamp
121
Performing flame test
1.Dip the clean wire loop (made up of Pt, Nichrome) in dil.HCl until it does not give a colour when held against the naked flame.
2.Dip a clean wire loop into a solid sample of the compound being tested
3.put the loop into the edge of the blue flame from a Bunsen burner
observe and record the flame colour produced
122
Cu flame test colour
Blue green (apple green)
123
Identification tests for H2,O2 and CO2
Place a burning splint near the opening of a test tube. If a popping noise occurs, it's probably hydrogen. Place a glowing splint in the test tube, and if it reignites, it could be oxygen. Place a burning splint into a test tube, and if it goes out, it could be carbon dioxide.
124
Only group IA element reacting with N2
Li
125
LiHCO3 decomposition
Li2O + CO2 + H2O
126
Only group IIA covalent carbonate
BeCO3
Slightly soluble in water
127
Pb2+ flame test color
Blue/white
128
2 white ppts that turns black on heating
AgS2O3
PbS2O3
129
Elements existing as gases at room temperature
11 elements
H, He, N, O, F, Ne, Cl, Ar, Kr, Xe, Rn
130
Elements existing as liquids at room temperature
Only 2 elements
mercury (a metal) and bromine (a halogen)
131
Solubility of an ionic compound depends on
Hydration energy and dissolution energy
132
Natural source and disinfectant produced by CaCO3
Lime stone
Bleaching powder
133
Types of reactions under organic chemistry
Addition
Substitution
Redox
Elimination
Acid/base reaction
Radical
134
Benzene ring with COR
acyl benzene
135
Alcohols + carboxylic acid with conc.H2SO4 and heat
Produces esters
136
Reaction of phenol with carbonates and bicarbonates
No reaction
137
Brady reagent
2,4-DNP or 2,4-dinitrophenylhydrazine
138
Clemmenson reduction
Zn(Hg) / conc.HCl
139
Reaction of amines with aldehydes and ketones
Amines show nucleophilic addition followed by elimination with aldehydes and ketones. The products are called imines.
140
Formation of diazonium salt from aniline occurs at which temperature
0-5
141
Al + NaOH + H2O =
NaAlO2 + H2
142
How to convert an aldehyde or a ketone into an alkane
By Clemmenson's reduction
Zn (Hg) / conc.HCl
143
What does the valency of an element of a compound means
Is the number of electrons participating in bonding of that particular element
144
What is the special feature that a compound should contain in order to colour the CCl4 layer violet when shaken with CCl4 and acidified solution of KI
It should undergo reduction
Compounds undergoing oxidation will not colour the CCl4 layer violet
145
Empirical formula and molecular formula
Is the simplest type of formula derived from the composition of the atomic masses of the elements in a compound.
Molecular formula is the formula that shows the actual number of atoms of each element in a molecule of a compound
146
Which quantum numbers are associated with the shape of an orbital
mainly the azimuthal or the angular momentum quantum number
but in d and f subshells ml (magnetic quantum number is also involved)
147
Strength of a metallic bond depends on
No. of electrons donated by atoms to create the metallic bond
Ionic radius
Ionic nature
148
Hydrogen bonds are formed between
H at one out of O,N,F
149
Bond angle of H2O vs H2S
H2O has a higher bond angle
O is more electronegative than S, because of this the repulsion between lone pairs and bonding electrons will be higher
150
Approximate bond angles of H2S, H2O and NH3
H2S - 90
H2O - 104.5
NH3 - 107
151
Which element has the highest melting point among S block metals
Be
152
State the condition under which the heat change of a reaction is equal to its enthalpy change
Under constant pressure
153
Upon addition of HNO3 and H, Fe3+ forms a complex which is dark blue
K4[Fe(CN)6]
Formed compound is Fe4[Fe(CN)6]3+
154
At the equivalence point of a equilibrium mixture
No. of acids = No.of base
Only the salt will be remained in the equilibrium mixture
155
Covalent character increases when the electronegativity difference between elements of a compound
Decreases
156
All metals
Conducts electricity
157
All d block elements are
metals
158
If the atoms of a molecule are in the same plane the possible hybridizations are
Sp and sp2
159
If the atoms of a molecule are not in the same plane the possible hybridizations are
Sp3
