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1
Q

Alcohol series

A

Contains functional hydroxyl (OH) group

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2
Q

Organic acids

A

Contains functional carboxylic group (COOH)
H-O-C=O

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3
Q

In acidic medium S2O32- will convert into?

A

S4O62-

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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.

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5
Q

2KMnO4(s) →

A

→ K2MnO4(s) + MnO2 (s) + O2(g)

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6
Q

What setup is used to find the molar mass of Mg

A

Eudiometer

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7
Q

named the fundamental unit carrying electricity as
“electrons”

A

Johnstone G. Stoney

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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

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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

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10
Q

experimentally proved the existence of positive charges in matter.

A

Eugen Goldstein
By using a perforated cathode

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11
Q

showed that radioactive materials produce three types of emissions

A

Lord Ernest Rutherford

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12
Q

Who were Rutherford’s associates

A

Geiger and Marsden

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13
Q

discovered the ‘neutron

A

Sir James Chadwick

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14
Q

found that the number of positive charges on the nucleus increases in atoms by single electron units

A

Jeffrey Moseley

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15
Q

Discovered the first isotopes (of neon) using mass spectrometer

A

J. J. Thomson and Francis William Aston

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16
Q

Wavelength range of visible range

A

400-750 nm

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17
Q

Wave particle duality

A

Louis de Broglie

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18
Q

ionization energy of Na

A

496 kJmol-1

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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.

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20
Q

published nearly identical periodic tables

A

Dmitri Ivanovich Mendeleev in Russia and Lothar Meyer in Germany

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21
Q

16th group elements are also known as

A

Chalcogens

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22
Q

Who put forwarded VSEPR theory

A

Gillespie and Nyholm

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23
Q

zincate

A

ZnO2 2-

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24
Q

aluminate

A

AlO2ˉ

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25
Q

chlorite

A

ClO2ˉ

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26
Q

chlorate

A

ClO3ˉ

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27
Q

phosphate

A

PO4 3-

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28
Q

HClO

A

hypochlorous acid

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29
Q

HClO2

A

chlorous acid

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30
Q

HClO3

A

chloric acid

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31
Q

HClO4

A

perchloric acid

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32
Q

 Standard enthalpy change of combustion

A

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

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33
Q

Standard enthalpy change of bond dissociation

A

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)

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34
Q

Standard enthalpy change of neutralization

A

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

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35
Q

Standard enthalpy change of solvation

A

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

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36
Q

Standard enthalpy change of hydration

A

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)

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37
Q

Standard enthalpy change of dissolution

A

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)

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38
Q

Standard enthalpy change of sublimation

A

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

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39
Q

Standard enthalpy change of evaporization

A

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

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40
Q

Standard enthalpy change of fusion

A

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.

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41
Q

Standard enthalpy change of atomization

A

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)

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42
Q

Standard enthalpy change of first ionization,

A

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

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43
Q

Standard enthalpy change of electron gain

A

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

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44
Q

Standard lattice dissociation enthalpy

A

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)

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45
Q

Test to identify water

A

Anhydrous copper sulfate is used which turns blue in the presence of water.

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46
Q

Increasing order of radius of isoelectronic species

A

Cations < neutral < anion

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47
Q

Soluble s block carbonate

A

BeCO3
All others are insoluble

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48
Q

Intensive and extensive properties

A

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

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49
Q

Special facts to memorize during identification of gases in an compressibility vs pressure graph

A

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

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50
Q

ΔH(reaction) =

A

ΔH(bond dissociation) + ΔH(bond formation)

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51
Q

Excess NH3 + Cl2

A

N2 (g) + NH4Cl (s)

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52
Q

Elements showing deviations from the aufbau principle

A

Cr, Cu, Rh

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53
Q

NaNO2 + NH4Cl =

A

NaCl + N2 + H2O

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54
Q

Experiment to show the presence N2 in air

A

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.

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55
Q

Which chloride gives blue colour to the flame test

A

CsCl(s)
Cs+

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56
Q

Standard experiment to identify NH3

A

When a filter paper is dipped in Nessler’s reagent and introduced to the gas, the yellow coloured Nessler’s reagent turns brown.

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57
Q

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

A

H2S

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58
Q

H2S vs H20 - Boiling point

A

H2S < H2O

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59
Q

Melting point of metals depend on the

A

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.

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60
Q

K2MnO4 colour

A

Green

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61
Q

6NaOH + 4S =

A

Na2S2O3 + 2Na2S + 3H2O

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62
Q

What do you mean by disproportion reaction

A

Same compound undergoes both oxidation and reduction
E.g: 2H2O2 → 2H2O + O2

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63
Q

Elements forming the bond with the highest ionic character

A

Should contain the lowest covalent character

CsF

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64
Q

Across a period which elements will have highest melting point

A

Group 14
Due to the formation of covalent bond network structures

65
Q

5 oxides of Nitrogen

A

N20, NO2, N2O3, N2O4, N2O5

N2O3, N2O4, N2O5 - dimer structures of nitrogen

66
Q

Vinegar like smell

A

CH3COOH (Acetic acid)

67
Q

H2S gas test

A

Smells like rotten eggs, turns

lead acetate paper black.

68
Q

SO2 gas test

A

Smells like burning sulphur (burnt matches)

turns acidified potassium dichromate paper green

69
Q

HCl gas test

A

Pungent smell, white fumes with ammonia, white ppt with silver nitrate solution.

70
Q

NH3 gas test

A

Characteristic smell, turns Nessler’s solution brown.

71
Q

NO2 gas test

A

Reddish brown, turns ferrous sulphate(blue green) solution black brown.

72
Q

Cl2 gas test

A

Greenish yellow, turns starch iodide paper blue.

73
Q

Br2 gas test

A

Reddish brown, turns starch paper orange yellow.

74
Q

I2 gas test

A

Dark violet, turns starch paper blue.

75
Q

Define critical temperature

A

Temperature at and above which the vapour of the substance cannot be liquefied, no matter how much pressure is applied.

76
Q

Critical pressure

A

Is the pressure required to liquefy a vapour at critical temperature

77
Q

Compounds formed when group 1 and 2 metals react with aqueous halides

A

Forms aqueous colourless solutions with halides

78
Q

TiO2

A

white

79
Q

CuCl(s)

A

White in colour
Used as a catalyst

80
Q

Ionic lattices when fused conducts electricity by

A

Existence of positive and negative ions
Metals conduct electricity due to the presence of mobile electrons

81
Q

In basic medium orange colour K2Cr2O7 reduces into

A

Yellow colour K2Cr2O4

82
Q

Ability of hydrolysis depends on

A

Covalent character of compounds
Higher the covalent character higher the ability to be hydrolysed.

83
Q

Although N2 is a inert gas when Al is combusted in air, N2 reacts with Al

A

Energy released during the combustion of Al with O2 is used to break the bonds between O2

84
Q

Which metal sulfate can be used in purification of water

A

Al2(SO4)3

85
Q

Na2S203 + H2SO4

A

Na2SO4 + SO2 + S + H2O

86
Q

Relationship between dissolution, hydration and lattice enthalpy

A

Dissolution = hydration - lattice enthalpy

87
Q

S + NaOH

A

Na2S + Na2S2O3 + H20

88
Q

H2O2 + PbS

A

PbSO4 + H2O

89
Q

SO2 + H2S

A

S + H20

90
Q

(NH4)2S —->

A

NH3 + H2S

91
Q

Al2O3 + NaOH

A

NaAlO2 + H20

92
Q

For a alkyl group connected to a benzene ring to be oxidised it should have

A

at least 1 H atom connected to it

Tertiary alkyl groups do not get oxidised.

93
Q

Which group 2 metal gives its peroxide with excess oxygen

A

Ba

94
Q

BaSO4

A

White ppt

BaSO4is extremely insoluble in water, alkalies, or acids, but is slightly soluble in hot, concentrated sulfuric acid.

95
Q

BaCO3

A

White ppt

Barium carbonate is soluble in acid, including dilute acetic acid, in strong bases, and in aqueous ammonia.

96
Q

BaNO3

A

White ppt

Soluble in hot water

97
Q

which hydroxide is used to identify CO2

A

Ca(OH)2

98
Q

what is used in the manufacture of baby soap

A

NaOH and *KOH

99
Q

Ag -White ppt that turns black on standing

A

AgS203 (white solid)——–> Ag2S (black solid)

100
Q

Yellow ppt of Ag dissolves in dil.nitric acid as well as in ammonium hydroxide

A

AgCrO4

101
Q

BaSO3

A

White ppt

Dissolves in di.HCl forming colourless, pungent smelling gas SO2

102
Q

PbSO4

A

PbSO4 is white and insoluble in HNO3 and hot water.

103
Q

Pb - Yellow ppt that dissolves in hot water

A

PbI2

104
Q

Which radicals are not formed

A

Hydrogen radicals

105
Q

Most abundant s block metals on earth’s crust

A

Na and K (In order)

106
Q

Most abundant element occurring in the oceans is

A

Na, Mg, Ca, K

107
Q

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

A

Pale pink solution makes white precipitate
This precipitate turns brown or blackish brown on exposure to air.

108
Q

Identification of Pb2+

A

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
Q

yellow ppt which dissolves on boiling and precipitate again as golden spangles on cooling

A

PbI2

110
Q

Identification of NO 3-

A

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
Q

Why Mg,Be do not produce colours under flame test

A

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
Q

Variation of reactivity of 1st group elements down the group (reason)

A

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
Q

Points to identify K

A

Lilac colour to the fame test.
Formation of superoxide.
Used in the production of baby soap.

114
Q

Na + H2O(g)

A

Na2O + H2

115
Q

Na2O2 + H2O

A

NaOH + H2O2

116
Q

KO2 + H2O

A

KOH + H2O2 + O2

117
Q

Only group IIA element forming peroxide with O2

A

Barium

118
Q

Mg + H2O(g)

A

MgO + H2

119
Q

Be + H2O(g)

A

BeO + H2

120
Q

Na is used in the production of

A

Yellow bulb / Sodium lamp

121
Q

Performing flame test

A

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
Q

Cu flame test colour

A

Blue green (apple green)

123
Q

Identification tests for H2,O2 and CO2

A

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
Q

Only group IA element reacting with N2

A

Li

125
Q

LiHCO3 decomposition

A

Li2O + CO2 + H2O

126
Q

Only group IIA covalent carbonate

A

BeCO3

Slightly soluble in water

127
Q

Pb2+ flame test color

A

Blue/white

128
Q

2 white ppts that turns black on heating

A

AgS2O3
PbS2O3

129
Q

Elements existing as gases at room temperature

A

11 elements
H, He, N, O, F, Ne, Cl, Ar, Kr, Xe, Rn

130
Q

Elements existing as liquids at room temperature

A

Only 2 elements
mercury (a metal) and bromine (a halogen)

131
Q

Solubility of an ionic compound depends on

A

Hydration energy and dissolution energy

132
Q

Natural source and disinfectant produced by CaCO3

A

Lime stone
Bleaching powder

133
Q

Types of reactions under organic chemistry

A

Addition
Substitution
Redox
Elimination
Acid/base reaction
Radical

134
Q

Benzene ring with COR

A

acyl benzene

135
Q

Alcohols + carboxylic acid with conc.H2SO4 and heat

A

Produces esters

136
Q

Reaction of phenol with carbonates and bicarbonates

A

No reaction

137
Q

Brady reagent

A

2,4-DNP or 2,4-dinitrophenylhydrazine

138
Q

Clemmenson reduction

A

Zn(Hg) / conc.HCl

139
Q

Reaction of amines with aldehydes and ketones

A

Amines show nucleophilic addition followed by elimination with aldehydes and ketones. The products are called imines.

140
Q

Formation of diazonium salt from aniline occurs at which temperature

A

0-5

141
Q

Al + NaOH + H2O =

A

NaAlO2 + H2

142
Q

How to convert an aldehyde or a ketone into an alkane

A

By Clemmenson’s reduction
Zn (Hg) / conc.HCl

143
Q

What does the valency of an element of a compound means

A

Is the number of electrons participating in bonding of that particular element

144
Q

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

A

It should undergo reduction

Compounds undergoing oxidation will not colour the CCl4 layer violet

145
Q

Empirical formula and molecular formula

A

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
Q

Which quantum numbers are associated with the shape of an orbital

A

mainly the azimuthal or the angular momentum quantum number

but in d and f subshells ml (magnetic quantum number is also involved)

147
Q

Strength of a metallic bond depends on

A

No. of electrons donated by atoms to create the metallic bond
Ionic radius
Ionic nature

148
Q

Hydrogen bonds are formed between

A

H at one out of O,N,F

149
Q

Bond angle of H2O vs H2S

A

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
Q

Approximate bond angles of H2S, H2O and NH3

A

H2S - 90
H2O - 104.5
NH3 - 107

151
Q

Which element has the highest melting point among S block metals

A

Be

152
Q

State the condition under which the heat change of a reaction is equal to its enthalpy change

A

Under constant pressure

153
Q

Upon addition of HNO3 and H, Fe3+ forms a complex which is dark blue

A

K4[Fe(CN)6]

Formed compound is Fe4[Fe(CN)6]3+

154
Q

At the equivalence point of a equilibrium mixture

A

No. of acids = No.of base

Only the salt will be remained in the equilibrium mixture

155
Q

Covalent character increases when the electronegativity difference between elements of a compound

A

Decreases

156
Q

All metals

A

Conducts electricity

157
Q

All d block elements are

A

metals

158
Q

If the atoms of a molecule are in the same plane the possible hybridizations are

A

Sp and sp2

159
Q

If the atoms of a molecule are not in the same plane the possible hybridizations are

A

Sp3

Chemistry (15 decks)

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