NEB Class 11 · Past paper
The complete NEB Class 11 2074 exam paper for Chemistry, all 33 questions with solved model answers.
Tap a question to open its answer.
How many oxygen molecules are required to produce 220 mg of CO2 according to the reaction C + O2 -> CO2?
$$ \begin{aligned} \ce{C + O2 -> CO2} \ n_{\ce{CO2}} &= \dfrac{0.220}{44} \ &= 0.005\ \text{mol} \end{aligned} $$
From the equation, $1$ mol $\ce{O2}$ gives $1$ mol $\ce{CO2}$, so $n_{\ce{O2}}=0.005\ \text{mol}$.
$$ \begin{aligned} \text{molecules of }\ce{O2} &= 0.005\times6.022\times10^{23} \ &= \mathbf{3.011\times10^{21}} \end{aligned} $$
State Gay-Lussac's law of gaseous volumes.
Gay-Lussac's law of gaseous volumes: when gases react, they do so in volumes that bear a simple whole-number ratio to one another and to the volumes of the gaseous products, provided all volumes are measured at the same temperature and pressure.
Example:
$$ \ce{H2(g) + Cl2(g) -> 2HCl(g)} $$
; one volume of hydrogen and one volume of chlorine give two volumes of hydrogen chloride ($1:1:2$).
Give proper reasons: i) Evaporation takes place from the surface of a liquid. ii) Liquid drops are spherical in shape.
i) Only the molecules at the surface of a liquid have a free face and can escape into the vapour; the more energetic surface molecules overcome the surface attractions and leave. Molecules inside are surrounded on all sides, so evaporation is a surface phenomenon.
ii) Because of surface tension, a liquid tends to have the minimum possible surface area. For a given volume a sphere has the least surface area, so free liquid drops take up a spherical shape.
Differentiate between hygroscopic and deliquescent substances, giving an example of each.
Hygroscopic substance: absorbs moisture from the air but does not dissolve in it (stays solid), used as a drying agent. Example: conc. $\ce{H2SO4}$, $\ce{CaO}$, silica gel.
Deliquescent substance: absorbs so much moisture from the air that it dissolves in the absorbed water, forming a solution. Example: $\ce{NaOH}$, $\ce{CaCl2}$, $\ce{MgCl2}$.
What information is provided by the magnetic quantum number? Give the values of l and m for a 3px orbital.
Magnetic quantum number ($m$): it gives the orientation (spatial direction) of an orbital in space relative to the applied magnetic field, and hence the number of orbitals in a subshell ($m$ ranges from $-l$ to $+l$).
For a 3p orbital: $l=1$, and $m$ can be $-1, 0, +1$. The $3p_x$ orbital corresponds to $l=1,\ m=+1$ (one of the $m=\pm1$ set).
Calculate the atomic number and the number of p-electrons of an atom whose valence shell electronic configuration is 4s2.
Valence shell $4s^2$ means the full configuration is $\ce{1s^2 2s^2 2p^6 3s^2 3p^6 4s^2}$.
$$ \begin{aligned} \text{Total electrons} &= 2+2+6+2+6+2 \ &= 20 \end{aligned} $$
so atomic number $Z=20$ (calcium).
$$ \begin{aligned} p\text{-electrons} &= \underbrace{2p^6}{6}+\underbrace{3p^6}{6} \ &= \mathbf{12} \end{aligned} $$
What is meant by controlled nuclear fission? Mention one important application.
Controlled nuclear fission: a fission chain reaction in which the number of neutrons causing further fission is regulated (e.g. by absorbing surplus neutrons with control rods) so that energy is released steadily rather than explosively.
Application: it is used in nuclear reactors to generate electricity (and to produce radioactive isotopes).
Write down the Lewis structure of i) N2O ii) HClO4.
(i) $\ce{N2O}$: a linear molecule; the two nitrogens are triple/double bonded and the terminal N bonds to O, with resonance $\ce{:N#N-O:} \leftrightarrow \ce{:N=N=O:}$ (total 16 valence electrons).
(ii) $\ce{HClO4}$ (perchloric acid): the central Cl is bonded to four O atoms - three by double bonds and one $\ce{-O-H}$ single bond: skeleton $\ce{H-O-Cl(=O)(=O)=O}$ (Cl in the $+7$ state).
How would you explain the polarity of a molecule on the basis of dipole moment?
The dipole moment ($\mu = q\times d$) measures the separation of positive and negative charge in a molecule; it is a vector.
Thus the magnitude of the net dipole moment tells us whether, and how strongly, a molecule is polar.
Which ion would you expect to have the larger size and why? Mg^2+ or Na^+.
$\ce{Na+}$ is larger than $\ce{Mg^{2+}}$.
Both ions are isoelectronic (each has 10 electrons, the neon configuration). But $\ce{Na+}$ has 11 protons while $\ce{Mg^{2+}}$ has 12 protons. The greater nuclear charge of $\ce{Mg^{2+}}$ pulls the same 10 electrons in more strongly, giving a smaller radius. So $\ce{Na+}$ has the larger ionic size.
Calculate the oxidation number of i) N in NH4OH ii) Cr in K2Cr2O7.
(i) N in $\ce{NH4OH}$:
Treat as $\ce{NH4+}$ and $\ce{OH-}$. In $\ce{NH4+}$:
$$ \begin{aligned} x+4(+1) &= +1 \ &\Rightarrow x = -3 \end{aligned} $$
N $=-3$.
(ii) Cr in $\ce{K2Cr2O7}$:
$$ \begin{aligned} 2(+1)+2x+7(-2) &= 0 \ 2+2x-14 &= 0 \ 2x &= 12 \ x &= +6 \end{aligned} $$
Cr $=+6$.
Write the relationship between Kp and Kc for the reactions i) PCl5(g) <=> PCl3(g) + Cl2(g) ii) N2(g) + 3H2(g) <=> 2NH3(g).
General relation: $K_p=K_c(RT)^{\Delta n}$, where $\Delta n=$ (gaseous product moles $-$ gaseous reactant moles).
**(i)
$$ \ce{PCl5 <=> PCl3 + Cl2} $$
:**
$$ \begin{aligned} \Delta n &= 2-1 \ &= +1 \end{aligned} $$
so
$$ \begin{aligned} K_p &= K_c(RT)^{1} \ &= K_c,RT \end{aligned} $$
**(ii)
$$ \ce{N2 + 3H2 <=> 2NH3} $$
:**
$$ \begin{aligned} \Delta n &= 2-4 \ &= -2 \end{aligned} $$
so
$$ \begin{aligned} K_p &= K_c(RT)^{-2} \ &= \dfrac{K_c}{(RT)^2} \end{aligned} $$
Which one is the least abundant isotope of hydrogen? Mention its two uses.
The least abundant isotope is tritium, $\ce{^3_1H}$ (radioactive, present only in traces).
Uses: (i) as a fuel in nuclear fusion (thermonuclear) reactions; (ii) as a radioactive tracer and in self-luminous (glow-in-the-dark) paints/dials.
Suggest your ideas to protect the ozone layer from its depletion.
The ozone layer can be protected by reducing the substances that destroy ozone:
(CFCs release Cl atoms in the stratosphere that catalytically destroy ozone:
$$ \ce{Cl + O3 -> ClO + O2} $$
.)
Give a balanced chemical reaction for the preparation of CO in the laboratory. Why is CO a harmful gas?
Laboratory preparation (dehydration of formic acid by warm conc. $\ce{H2SO4}$): $$\ce{HCOOH ->[conc.\ H2SO4] CO(g) + H2O}$$
Why harmful: carbon monoxide combines with the haemoglobin of blood about 200-300 times more strongly than oxygen does, forming stable carboxyhaemoglobin. This blocks haemoglobin from carrying oxygen, so the body tissues are starved of oxygen, which can cause suffocation and death.
Write down the molecular formula and one use of each of the following: i) Boric acid ii) Orthophosphoric acid.
(i) Boric acid: $\ce{H3BO3}$ - used as a mild antiseptic (eye wash) and in glazing/enamels.
(ii) Orthophosphoric acid: $\ce{H3PO4}$ - used in the manufacture of phosphate fertilizers and as an acidulant/rust remover and in soft drinks.
What is meant by slag? Write an example of it.
Slag: a fusible, easily removable waste substance formed during the smelting of a metal when the flux combines with the gangue (earthy impurity) of the ore. Being lighter, molten slag floats on the molten metal and is skimmed off.
Example: in the extraction of iron, the flux limestone gives $\ce{CaO}$, which combines with the silica gangue to form calcium silicate slag: $$\ce{CaO + SiO2 -> CaSiO3}\ \text{(slag)}$$
What happens when i) bleaching powder is treated with dil. H2SO4 ii) gypsum salt is heated above 130 degree C?
(i) Bleaching powder + dilute $\ce{H2SO4}$ (chlorine is liberated - large excess of acid): $$\ce{CaOCl2 + H2SO4 -> CaSO4 + H2O + Cl2(g)}$$
(ii) Gypsum heated above ~130 C: it loses most of its water of crystallisation and changes to plaster of paris (calcium sulphate hemihydrate); on stronger heating (above ~200 C) it becomes anhydrous dead-burnt plaster: $$\ce{2(CaSO4.2H2O) ->[~130,^\circ C] (CaSO4)2.H2O + 3H2O}$$
Write down the IUPAC name of the following organic compounds: i) CH3-O-CH2-CH2-OCH3 ii) CHCl3.
(i) $\ce{CH3-O-CH2-CH2-O-CH3}$: an ethane chain bearing a methoxy group on each carbon, 1,2-dimethoxyethane.
(ii) $\ce{CHCl3}$: trichloromethane (common name chloroform).
Write reactions for the preparation of methane from i) CH3MgBr ii) CH3COONa.
(i) From methyl magnesium bromide (hydrolysis of the Grignard reagent): $$\ce{CH3MgBr + H2O -> CH4 + Mg(OH)Br}$$
(ii) From sodium ethanoate (decarboxylation with soda lime, NaOH + CaO, on heating): $$\ce{CH3COONa + NaOH ->[CaO][\Delta] CH4 + Na2CO3}$$
Identify the major products A and B in the given reaction sequence (reaction diagram involving a benzene derivative) and write their IUPAC names.
(The exact reagents in the printed reaction diagram are not fully legible in the source paper, so the specific structures cannot be reproduced verbatim. )
General approach for a benzene-based sequence: benzene ($\ce{C6H6}$) typically undergoes electrophilic substitution. For example: $$ \begin{aligned} \ce{C6H6 + Cl2 ->[FeCl3] C6H5Cl}\ (A,\ \text{chlorobenzene}) \ \ce{C6H6 + HNO3 ->[conc.\ H2SO4] C6H5NO2}\ (B,\ \text{nitrobenzene}) \end{aligned} $$ Identify A and B by applying the substitution/directing rules to the reagents shown in the original diagram.
Give the functional isomers of ethanoic acid and propanone with their structural formulae.
Ethanoic acid, $\ce{CH3COOH}$ ($\ce{C2H4O2}$): its functional isomer is the ester methyl methanoate, $\ce{H-CO-O-CH3}$ ($\ce{HCOOCH3}$).
Propanone, $\ce{CH3-CO-CH3}$ ($\ce{C3H6O}$): its functional isomer is the aldehyde propanal, $\ce{CH3-CH2-CHO}$.
What are the observations of Rutherford's alpha-ray scattering experiment? Write down the conclusions drawn from it about the structure of the atom.
Observations (a thin gold foil bombarded with $\alpha$-particles):
Conclusions:
Give an example to show that oxidation and reduction go simultaneously. Balance the following equation by the oxidation number or ion-electron method: HNO3 + H2S -> SO2 + NO2 + H2O.
Redox is simultaneous: in every redox reaction, when one species is oxidised (loses electrons) another is reduced (gains those electrons). For example, in
$$ \ce{Zn + CuSO4 -> ZnSO4 + Cu} $$
, Zn is oxidised while $\ce{Cu^{2+}}$ is reduced at the same time.
Balancing
$$ \ce{HNO3 + H2S -> SO2 + NO2 + H2O} $$
: S goes $-2\to+4$ (loses $6e^-$); N goes $+5\to+4$ (gains $1e^-$). So 6 N atoms per 1 S: $$\ce{6HNO3 + H2S -> SO2 + 6NO2 + 4H2O}$$ Check: H
$$ \begin{aligned} 6+2 &= 8 \ &= 8 \end{aligned} $$
N $6=6$; S $1=1$; O $18=2+12+4$. Balanced. ($\ce{HNO3}$ is the oxidant, $\ce{H2S}$ the reductant.)
How would you detect the presence of nitrogen in an organic compound by the Lassaigne test method?
Lassaigne's test for nitrogen:
Deduce the relationship between molecular mass and vapour density of a gas using Avogadro's hypothesis. An oxide of nitrogen contains half of its own volume of nitrogen and has a vapour density equal to 15. Determine the molecular formula of the nitrogen oxide.
Molecular mass and vapour density: by Avogadro's hypothesis, equal volumes of gases (same T, P) contain equal numbers of molecules. Vapour density (VD) is the ratio of the mass of a volume of gas to the mass of the same volume of hydrogen:
$$ \begin{aligned} \text{VD} &= \dfrac{\text{mass of }n\text{ molecules of gas}}{\text{mass of }n\text{ molecules of }\ce{H2}} \ &= \dfrac{M}{2} \ &\Rightarrow \boxed{M=2\times\text{VD}} \end{aligned} $$
Molecular formula:
$$ \begin{aligned} M &= 2\times15 \ &= 30\ \text{g/mol} \end{aligned} $$
"Contains half its own volume of nitrogen" means 1 volume of the oxide gives $\tfrac12$ volume of $\ce{N2}$, i.e. one molecule of oxide contains one N atom (mass 14).
$$ \begin{aligned} \text{Remaining mass} &= 30-14 \ &= 16 \ &= \text{one O atom} \end{aligned} $$
So the molecular formula is $\ce{NO}$ (nitric oxide).
Why can't HBr be prepared by using NaBr and conc. H2SO4? Give a balanced chemical reaction for the preparation of HCl. How would you test the presence of Cl- or Br- ions in aqueous solution?
Why not HBr from NaBr + conc. $\ce{H2SO4}$: conc. sulphuric acid is a strong oxidising agent. Although it first liberates HBr, it then oxidises the HBr to bromine, so pure HBr cannot be obtained: $$\ce{2HBr + H2SO4(conc) -> Br2 + SO2 + 2H2O}$$ (A non-oxidising acid such as syrupy $\ce{H3PO4}$ is used instead.)
Preparation of HCl (chloride is not oxidised by $\ce{H2SO4}$): $$\ce{NaCl + H2SO4 ->[\Delta] NaHSO4 + HCl(g)}$$
Test for $\ce{Cl-}$ / $\ce{Br-}$ (add dilute $\ce{HNO3}$, then $\ce{AgNO3}$):
$$ \ce{Ag+ + Cl- -> AgCl(s)} $$
.
$$ \ce{Ag+ + Br- -> AgBr(s)} $$
.
Describe the Kipp's apparatus method for the manufacture of H2S. Give a reaction to show that H2S is a reducing agent.
Kipp's apparatus for $\ce{H2S}$: Kipp's apparatus is a three-bulbed glass generator used to make a gas on demand at room temperature. Lumps of iron(II) sulphide ($\ce{FeS}$) are placed in the middle bulb, and dilute sulphuric (or hydrochloric) acid is poured in the top bulb. When the tap is opened, acid rises to the FeS and produces $\ce{H2S}$: $$\ce{FeS + H2SO4(dil) -> FeSO4 + H2S(g)}$$ When the tap is closed, the gas pressure pushes the acid back down away from the FeS, stopping the reaction automatically.
$\ce{H2S}$ as a reducing agent (it reduces chlorine, itself giving sulphur): $$\ce{H2S + Cl2 -> 2HCl + S(s)}$$
Write short notes on i) Hydrometallurgy ii) Zone refining.
i) Hydrometallurgy: extraction of a metal by dissolving its ore in a suitable aqueous reagent (leaching) to form a soluble salt/complex, then recovering the metal from solution by displacement with a more reactive metal or by electrolysis. Example - silver:
$$ \ce{Ag2S + 4NaCN -> 2Na[Ag(CN)2] + Na2S} $$
, then
$$ \ce{2Na[Ag(CN)2] + Zn -> Na2[Zn(CN)4] + 2Ag} $$
.
ii) Zone refining: a method to obtain very pure metals (e.g. Si, Ge) based on the fact that impurities are more soluble in the molten than in the solid metal. A circular heater is moved slowly along a rod of the metal; as the molten zone travels, impurities are carried along with it to one end, leaving behind ultra-pure metal. The impure end is finally cut off.
How is ammonia manufactured by Haber's process? What happens when ammonia is passed through i) copper sulphate solution ii) heated copper oxide? Give reactions for the conversion of ammonia to nitric acid.
Haber's process: purified nitrogen (from air) and hydrogen (from water gas) in a $1:3$ ratio are compressed to about $200\ \text{atm}$ and passed at about $450,^\circ\text{C}$ over a finely divided iron catalyst (with molybdenum promoter): $$\ce{N2(g) + 3H2(g) <=>[Fe][450,^\circ C] 2NH3(g)}\quad \Delta H=-ve$$ The reaction is exothermic and reduces gaseous moles, so high pressure and moderately low temperature favour $\ce{NH3}$; unreacted gases are recycled.
i) $\ce{NH3}$ into $\ce{CuSO4}$ solution: a deep-blue complex forms: $$\ce{CuSO4 + 4NH3 -> [Cu(NH3)4]SO4}$$
ii) $\ce{NH3}$ over heated CuO: ammonia reduces it to copper: $$\ce{3CuO + 2NH3 ->[\Delta] 3Cu + N2 + 3H2O}$$
Ammonia to nitric acid (Ostwald): $$ \begin{aligned} \ce{4NH3 + 5O2 ->[Pt/Rh] 4NO + 6H2O} \ \ce{2NO + O2 -> 2NO2} \ \ce{4NO2 + O2 + 2H2O -> 4HNO3} \end{aligned} $$
Write the principle and process along with a self-explanatory diagram for the manufacture of sodium by Down's method. Why is a sodium fire in the laboratory not extinguished by adding water? Mention any two advantages of Down's process. How is sodium converted into washing soda?
Down's process (principle): electrolysis of fused sodium chloride (with $\ce{CaCl2}$/$\ce{KF}$ added to lower the melting point to ~$600,^\circ\text{C}$). $$ \begin{aligned} \text{Cathode (iron): }\ce{Na+ + e- -> Na(l)} \ \text{Anode (graphite): }\ce{2Cl- -> Cl2(g) + 2e-} \end{aligned} $$ Molten sodium collects at the cathode and is drawn off; chlorine is a valuable by-product. A steel gauze diaphragm keeps Na and $\ce{Cl2}$ apart.
Sodium fire not put out with water: sodium reacts violently with water, giving flammable hydrogen and much heat,
$$ \ce{2Na + 2H2O -> 2NaOH + H2(g)} $$
; the hydrogen catches fire and the fire spreads instead of being quenched. (Dry sand is used.)
Two advantages of Down's process: (i) it gives very pure sodium; (ii) chlorine is obtained as a useful by-product (and there is no danger from caustic products, and the process is continuous).
Sodium to washing soda: sodium first reacts with water to give caustic soda,
$$\ce{2Na + 2H2O -> 2NaOH + H2}$$
the caustic soda then absorbs carbon dioxide to give sodium carbonate,
$$\ce{2NaOH + CO2 -> Na2CO3 + H2O}$$
and crystallisation finally gives the decahydrate, washing soda,
$$\ce{Na2CO3 + 10H2O -> Na2CO3.10H2O}$$
State Boyle's law and Charle's law. Derive the relation PV = nRT. An evacuated glass vessel weighs 50 g when empty, 148 g when filled with a liquid of density 0.98 g/ml, and 50.5 g when filled with an ideal gas at 760 mmHg at 27 degree C. Determine the molecular mass of the gas. How many molecules are present in 12.3 g of the gas?
Boyle's law: at constant temperature, $V\propto\dfrac{1}{P}$ (for a fixed mass of gas).
Charles's law: at constant pressure, $V\propto T$ (T in kelvin).
Derivation of $PV=nRT$: combining Boyle's, Charles's and Avogadro's ($V\propto n$) laws, $V\propto\dfrac{nT}{P}$, so $V=\dfrac{RnT}{P}$, giving $\boxed{PV=nRT}$.
Numerical:
$$ \begin{aligned} \text{volume of vessel} &= \text{volume of liquid} \ &= \dfrac{148-50}{0.98} \ &= \dfrac{98}{0.98} \ &= 100\ \text{ml} \ &= 0.1\ \text{L} \ \text{mass of gas} &= 50.5-50 \ &= 0.5\ \text{g} \end{aligned} $$
$T=300\ \text{K}$, $P=1\ \text{atm}$.
$$ \begin{aligned} n &= \dfrac{PV}{RT} \ &= \dfrac{1\times0.1}{0.0821\times300} \ &= 4.06\times10^{-3}\ \text{mol} \ M &= \dfrac{\text{mass}}{n} \ &= \dfrac{0.5}{4.06\times10^{-3}} \ &= \mathbf{123\ g/mol} \end{aligned} $$
Molecules in $12.3\ \text{g}$:
$$ \begin{aligned} n &= \dfrac{12.3}{123} \ &= 0.10\ \text{mol} \ \text{molecules} &= 0.10\times6.022\times10^{23} \ &= \mathbf{6.02\times10^{22}} \end{aligned} $$
Write short notes on any two: i) Characteristics of homologous series ii) Laboratory preparation of ethyne iii) Mendeleev's periodic table and its anomalies iv) Law of mass action and equilibrium constant.
i) Characteristics of homologous series: members share one general formula and the same functional group; successive members differ by $\ce{-CH2-}$ ($14\ \text{u}$); physical properties change gradually with molar mass; all members have similar chemical properties and similar general methods of preparation.
ii) Laboratory preparation of ethyne: water is added to calcium carbide, evolving ethyne, collected over water: $$\ce{CaC2 + 2H2O -> Ca(OH)2 + C2H2(g)}$$
iii) Mendeleev's periodic table and anomalies: Mendeleev arranged elements by increasing atomic mass (properties a periodic function of atomic mass). Anomalies: no fixed place for hydrogen; isotopes not accommodated; anomalous pairs where a heavier element precedes a lighter one (Ar-K, Co-Ni, Te-I); dissimilar elements sometimes grouped together; lanthanides could not be placed.
iv) Law of mass action and equilibrium constant: the rate of a reaction is proportional to the product of the active masses (molar concentrations) of the reactants. Applying it to a reversible reaction
$$ \ce{aA + bB <=> cC + dD} $$
at equilibrium gives the equilibrium constant $K_c=\dfrac{[\ce{C}]^c[\ce{D}]^d}{[\ce{A}]^a[\ce{B}]^b}$, a constant at a given temperature.
From the equation, mol gives mol , so .
Gay-Lussac's law of gaseous volumes: when gases react, they do so in volumes that bear a simple whole-number ratio to one another and to the volumes of the gaseous products, provided all volumes are measured at the same temperature and pressure.
Example:
; one volume of hydrogen and one volume of chlorine give two volumes of hydrogen chloride ().
Hygroscopic substance: absorbs moisture from the air but does not dissolve in it (stays solid), used as a drying agent. Example: conc. , , silica gel.
Deliquescent substance: absorbs so much moisture from the air that it dissolves in the absorbed water, forming a solution. Example: , , .
Magnetic quantum number (): it gives the orientation (spatial direction) of an orbital in space relative to the applied magnetic field, and hence the number of orbitals in a subshell ( ranges from to ).
For a 3p orbital: , and can be . The orbital corresponds to (one of the set).
Valence shell means the full configuration is .
so atomic number (calcium).
(i) : a linear molecule; the two nitrogens are triple/double bonded and the terminal N bonds to O, with resonance (total 16 valence electrons).
(ii) (perchloric acid): the central Cl is bonded to four O atoms - three by double bonds and one single bond: skeleton (Cl in the state).
The dipole moment () measures the separation of positive and negative charge in a molecule; it is a vector.
Thus the magnitude of the net dipole moment tells us whether, and how strongly, a molecule is polar.
is larger than .
Both ions are isoelectronic (each has 10 electrons, the neon configuration). But has 11 protons while has 12 protons. The greater nuclear charge of pulls the same 10 electrons in more strongly, giving a smaller radius. So has the larger ionic size.
(i) N in :
Treat as and . In :
N .
(ii) Cr in :
Cr .
General relation: , where (gaseous product moles gaseous reactant moles).
**(i)
:**
so
**(ii)
:**
so
The least abundant isotope is tritium, (radioactive, present only in traces).
Uses: (i) as a fuel in nuclear fusion (thermonuclear) reactions; (ii) as a radioactive tracer and in self-luminous (glow-in-the-dark) paints/dials.
The ozone layer can be protected by reducing the substances that destroy ozone:
(CFCs release Cl atoms in the stratosphere that catalytically destroy ozone:
.)
Laboratory preparation (dehydration of formic acid by warm conc. ):
Why harmful: carbon monoxide combines with the haemoglobin of blood about 200-300 times more strongly than oxygen does, forming stable carboxyhaemoglobin. This blocks haemoglobin from carrying oxygen, so the body tissues are starved of oxygen, which can cause suffocation and death.
(i) Boric acid: - used as a mild antiseptic (eye wash) and in glazing/enamels.
(ii) Orthophosphoric acid: - used in the manufacture of phosphate fertilizers and as an acidulant/rust remover and in soft drinks.
Slag: a fusible, easily removable waste substance formed during the smelting of a metal when the flux combines with the gangue (earthy impurity) of the ore. Being lighter, molten slag floats on the molten metal and is skimmed off.
Example: in the extraction of iron, the flux limestone gives , which combines with the silica gangue to form calcium silicate slag:
(i) Bleaching powder + dilute (chlorine is liberated - large excess of acid):
(ii) Gypsum heated above ~130 C: it loses most of its water of crystallisation and changes to plaster of paris (calcium sulphate hemihydrate); on stronger heating (above ~200 C) it becomes anhydrous dead-burnt plaster:
(i) : an ethane chain bearing a methoxy group on each carbon, 1,2-dimethoxyethane.
(ii) : trichloromethane (common name chloroform).
(i) From methyl magnesium bromide (hydrolysis of the Grignard reagent):
(ii) From sodium ethanoate (decarboxylation with soda lime, NaOH + CaO, on heating):
(The exact reagents in the printed reaction diagram are not fully legible in the source paper, so the specific structures cannot be reproduced verbatim. )
General approach for a benzene-based sequence: benzene () typically undergoes electrophilic substitution. For example:
Identify A and B by applying the substitution/directing rules to the reagents shown in the original diagram.
Ethanoic acid, (): its functional isomer is the ester methyl methanoate, ().
Propanone, (): its functional isomer is the aldehyde propanal, .
Observations (a thin gold foil bombarded with -particles):
Conclusions:
Redox is simultaneous: in every redox reaction, when one species is oxidised (loses electrons) another is reduced (gains those electrons). For example, in
, Zn is oxidised while is reduced at the same time.
Balancing
: S goes (loses ); N goes (gains ). So 6 N atoms per 1 S: Check: H
N ; S ; O . Balanced. ( is the oxidant, the reductant.)
Lassaigne's test for nitrogen:
Molecular mass and vapour density: by Avogadro's hypothesis, equal volumes of gases (same T, P) contain equal numbers of molecules. Vapour density (VD) is the ratio of the mass of a volume of gas to the mass of the same volume of hydrogen:
Molecular formula:
"Contains half its own volume of nitrogen" means 1 volume of the oxide gives volume of , i.e. one molecule of oxide contains one N atom (mass 14).
So the molecular formula is (nitric oxide).
Why not HBr from NaBr + conc. : conc. sulphuric acid is a strong oxidising agent. Although it first liberates HBr, it then oxidises the HBr to bromine, so pure HBr cannot be obtained: (A non-oxidising acid such as syrupy is used instead.)
Preparation of HCl (chloride is not oxidised by ):
Test for / (add dilute , then ):
.
.
Kipp's apparatus for : Kipp's apparatus is a three-bulbed glass generator used to make a gas on demand at room temperature. Lumps of iron(II) sulphide () are placed in the middle bulb, and dilute sulphuric (or hydrochloric) acid is poured in the top bulb. When the tap is opened, acid rises to the FeS and produces : When the tap is closed, the gas pressure pushes the acid back down away from the FeS, stopping the reaction automatically.
as a reducing agent (it reduces chlorine, itself giving sulphur):
i) Hydrometallurgy: extraction of a metal by dissolving its ore in a suitable aqueous reagent (leaching) to form a soluble salt/complex, then recovering the metal from solution by displacement with a more reactive metal or by electrolysis. Example - silver:
, then
.
ii) Zone refining: a method to obtain very pure metals (e.g. Si, Ge) based on the fact that impurities are more soluble in the molten than in the solid metal. A circular heater is moved slowly along a rod of the metal; as the molten zone travels, impurities are carried along with it to one end, leaving behind ultra-pure metal. The impure end is finally cut off.
Haber's process: purified nitrogen (from air) and hydrogen (from water gas) in a ratio are compressed to about and passed at about over a finely divided iron catalyst (with molybdenum promoter): The reaction is exothermic and reduces gaseous moles, so high pressure and moderately low temperature favour ; unreacted gases are recycled.
i) into solution: a deep-blue complex forms:
ii) over heated CuO: ammonia reduces it to copper:
Ammonia to nitric acid (Ostwald):
Down's process (principle): electrolysis of fused sodium chloride (with / added to lower the melting point to ~).
Molten sodium collects at the cathode and is drawn off; chlorine is a valuable by-product. A steel gauze diaphragm keeps Na and apart.
Sodium fire not put out with water: sodium reacts violently with water, giving flammable hydrogen and much heat,
; the hydrogen catches fire and the fire spreads instead of being quenched. (Dry sand is used.)
Two advantages of Down's process: (i) it gives very pure sodium; (ii) chlorine is obtained as a useful by-product (and there is no danger from caustic products, and the process is continuous).
Sodium to washing soda: sodium first reacts with water to give caustic soda,
the caustic soda then absorbs carbon dioxide to give sodium carbonate,
and crystallisation finally gives the decahydrate, washing soda,
Boyle's law: at constant temperature, (for a fixed mass of gas).
Charles's law: at constant pressure, (T in kelvin).
Derivation of : combining Boyle's, Charles's and Avogadro's () laws, , so , giving .
Numerical:
, .
Molecules in :
i) Characteristics of homologous series: members share one general formula and the same functional group; successive members differ by (); physical properties change gradually with molar mass; all members have similar chemical properties and similar general methods of preparation.
ii) Laboratory preparation of ethyne: water is added to calcium carbide, evolving ethyne, collected over water:
iii) Mendeleev's periodic table and anomalies: Mendeleev arranged elements by increasing atomic mass (properties a periodic function of atomic mass). Anomalies: no fixed place for hydrogen; isotopes not accommodated; anomalous pairs where a heavier element precedes a lighter one (Ar-K, Co-Ni, Te-I); dissimilar elements sometimes grouped together; lanthanides could not be placed.
iv) Law of mass action and equilibrium constant: the rate of a reaction is proportional to the product of the active masses (molar concentrations) of the reactants. Applying it to a reversible reaction
at equilibrium gives the equilibrium constant , a constant at a given temperature.