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  • #3349

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    Organic Chemistry – Concepts

    1. Empirical formula is the simplest formula that shows the ratio of each kind of atom in a molecule. e.g. C2H5 is the empirical formula for C4H10

    2. Molecular formula shows the actual number of each kind of atoms in a molecule. e.g. C4H10

    3. Structural formula shows how the atoms are connected to each other in a molecule. e.g. CH3CH2CH2CH3

    4. Displayed/full formula shows all the bonds and relative placing of all the atoms in a molecule.

    5. Homologous series are compounds have the same general formula and functional group and each homologue differs from its neighbor by a fixed group of atoms (e.g.–CH2). As we go down a homologous series, the chemical properties remain unchanged but there is a gradual change in physical properties. Examples of homologous series are alkanes, alkenes, alcohols…..

    6. Structural isomerism refers to compounds with the same molecular formula but different structural formula. E.g. CH3COOCH3 and C2H5COOH

    7. Stereoisomerism refers to compounds that have the same molecular formula but with different spatial arrangements.

    • Geometric isomers have same carbon skeleton with double bonds restricting free rotation. For geometric isomerism to exist, there must be two different groups of atoms bonded to each side of the C=C bond.

    • Optical isomers are non-superimposable mirror images of each other (enantiomers). Isomers have at least one chiral C atom, i.e. there are four different groups attached and have no plane of symmetry. An equal proportion of enantiomers forms a racemic mixture which is optically inactive.

    8. The primary structure of a protein shows the exact order (or unique sequence) of the -amino acids held by peptide/amide linkages along the polypeptide chain. The primary structure determines what the protein is, how it folds and its function.

    9. The secondary structure refers to the detailed configurations of the polypeptide chain. In a protein molecule, the long chain of amino acid units may be coiled into an -helix or folded into a -pleated sheet. Both structures are stabilized by hydrogen bonds between the N-H group of one amino acid residue and the C=O group of another along the main chain.

    10. The tertiary structure of the protein refers to the overall 3-dimensional shape of the entire protein involving folding or coiling of the chains. It shows how protein molecules are arranged in relation to each other.

    There are four types of R group interactions which hold the tertiary structure in its shape.

     van der Waals’ forces (induced dipole-induced dipole bonding) exist when non-polar R groups (e.g. alkyl or aryl groups) come close together. They are usually found on the inside of globular proteins where, because they are hydrophobic, they do not interfere with solubility.

     hydrogen bonding between polar groups (e.g.. –CH2OH, -COOH and –NH2 groups).

     ionic bonding eg. –COO-, -NH3+, and >NH2+.

     disulfide linkages eg. –SH or –CH2-S-S-CH2- groups.
    Quaternary structure of proteins refers to the spatial arrangement of its protein subunits. It shows how the individually folded protein subunits are packed together to yield large structures. This only applies to proteins that contain two or more polypeptide chains. The individual polypeptide chains are called the subunits. E.g. haemoglobin contains 4 subunits, each containing a haem group.

    11. It is stabilized by the same R-group interactions that stabilise the tertiary structure.

    12. Denaturation is the loss of biological activity of a native protein. When proteins are denatured, the secondary and tertiary structures are disrupted i.e. the R group interactions are broken or destroyed. Note that the primary structure remains unaffected.
    Factors that can lead to denaturation include extremes in pH, temperature, ionic salts, heavy metal compounds, presence of organic solvents etc.

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    #3387

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    TO Master to PERFECTION before A’levels (Part 1)

    Standard Definitions (Don’t Memorize. But appreciate and understand why key terms are important)

    – Relative atomic, isotopic, molecular and formula mass, based on the 12C scale (just give mathematical expression)

    – Mole in terms of the Avogadro constant
    – VSEPR (2 assumptions)
    – Basic assumptions of the kinetic theory as applied to an ideal gas
    – Standard enthalpies (11 of them)
    – Hess’ Law
    – Entropy
    – Standard electrode potential and standard cell potential
    – Dynamic Equilibrium, LCP
    – Strong and weak acids and bases
    – Kc, KP, Ka, Kb, Kw, KSP,pH etc. (m. expression)
    – Rate of reaction; rate equation; order of reaction; rate constant; (m. expression)
    – Half life of a reaction
    – Rate-determining step
    – Activation energy
    – Catalysts
    – Transition metal, ligands, complex, coordination number
    – Proteins 1o,2o,3o structure, Denaturation

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    #3403

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    FAQ – Group II

    For Gp II metals (e.g. magnesium, calcium, barium)

    1. Reactivity (with water or other substances) increases down the group

    Explanation:
    • Increase in shielding effect outweighs increase in nuclear charge down the group due to increase
    in number of electron shells.
    • Hence ionization energy decreases
    • Metals are able to form ions more easily.

    For Gp II ionic compounds (e.g. carbonates, nitrates and hydroxides),

    2. Melting point decreases down the group

    Explanation:
    • Increase in cation size
    • leads to decrease in magnitude of lattice energy.
    • Hence, ionic bonds are weaker and more easily broken.

    3. Thermal decomposition temperature increases down the group

    Explanation:
    • Increase in cation size,
    • leads to decrease in charge density of cation.
    • Hence, electron cloud of anion is less distorted and hence less easily decomposed.

    Note: Quote ionic radius values from Data Booklet to explain this. Calculate charge density if necessary.

    4. Solubility of sulfates decreases down the group (not in syllabus but
    may still be tested)

    Explanation:
    • Increase in cation size
    • leads to significant decrease in magnitude of hydration energy (as compared to slight decrease in magnitude of lattice energy)
    • Hence, enthalpy of solution (= LE – ΔHhyd) is less exothermic.

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    #3447

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    FAQ – Group VII

    For Group VII halogens (e.g. Cl2, Br2, I2)

    1. Volatility decreases down the group (i.e. boiling point increases)

    Explanation:
    • Increase in electron cloud size hence more easily polarised
    • leads to stronger dispersion forces between molecules

    2. Oxidising power decreases down the group

    X2 + 2e ⇔ 2X−

    Explanation:
    • Decrease in effective nuclear charge, hence electron affinity decreases
    • Hence less easily accepts electrons (i.e. less easily reduced)
    • Eθ value decreases (Quote from Data Booklet)

    Two important proofs of this are:
    • Any halide (e.g. I−) can be displaced by the halogen (e.g. Br2) above it.
    • Reaction with sodium thiosulfate (explain change in oxidation state of S)

    3. Reactivity with H2 decreases down the group
    X2 + H2 → 2HX

    Explanation:
    • Atomic size of X increases
    • Hence H – X bond becomes longer and weaker
    • Product formed is less and less stable, hence reactivity decreases.

    Note: Quote bond energy values to explain this, NOT Eθ values

    For exam based questions with solutions please contact @9863 9633

    #3491

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    FAQ – Gases

    Gases

    Difference between assumptions and conditions

    • 2 assumptions of ideal gas
    o Negligible intermolecular forces
    o Negligible particle volume compared to volume of container

    • 2 conditions at which a gas acts most ideally
    o High temperature
    o Low pressure

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    #3531

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    Colour Summary

    Flame Test Colors
    Li Deep red
    Na Yellow
    K Violet
    Mg Bright white
    Ca Orange-red
    Sr Red
    Ba Green
    Cu Blue-green
    P Pale blue-green
    S Blue
    Fe Gold
    Pb Blue-white
    Zn Blue-green

    Aqueous Ion Colors
    Cu+ Green
    Cu2+ Blue
    [CuCl4]2- Yellow
    Cu(NH3)4 2+ Dark Blue; produced when ammonia is added to Cu2+ solutions

    Fe2+ yellow-green (depending on the anion)
    Fe3+ orange-red (depending on the anion)
    FeSCN]2+ Red-brown, Wine-red to dark orange

    Co2+ Pink
    CoCl42- Blue (Co2+ with HCl will form a CoCl4 2- complex that is blue)

    Cr3+ Violet (Cr(NO3)3 to Green (CrCl3)
    CrO4 2- Yellow
    Cr2O7 2- Orange

    Ni2+ Green

    Mn2+ Pink
    MnO4 – Purple (Mn w/ +7 oxidation state is purple)
    MnO4 2- Green

    Pb3+ blue-green (Pb2+ and Pb4+ are colorless)

    V2+ violet
    V3+ blue-green

    Ti(H2O)6 3+ Purple

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    #3556

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    TO Master to PERFECTION before A’levels – Part 1

    Standard Definitions (Don’t Memorize. But appreciate and understand why key terms are important)

    – Relative atomic, isotopic, molecular and formula mass, based on the 12C scale (just give mathematical expression)
    – Mole in terms of the Avogadro constant
    – VSEPR (2 assumptions)
    – Basic assumptions of the kinetic theory as applied to an ideal gas
    – Standard enthalpies (11 of them)
    – Hess’ Law
    – Entropy
    – Standard electrode potential and standard cell potential
    – Dynamic Equilibrium, LCP
    – Strong and weak acids and bases
    – Kc, KP, Ka, Kb, Kw, KSP,pH etc. (m. expression)
    – Rate of reaction; rate equation; order of reaction; rate constant; (m. expression)
    – Half life of a reaction
    – Rate-determining step
    – Activation energy
    – Catalysts
    – Transition metal, ligands, complex, coordination number
    – Proteins 1o,2o,3o structure, Denaturation

    Standard Explanations (must be concise (save time), accurate and complete) – You must know this so well you have are absolutely confident of reproducing them under stressful exam conditions.

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    #3594

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    TO Master to PERFECTION before A’levels – Part 2

    1. Atomic Structure • Ionisation Energy (Trend across the period + 2 anomalies, Down the gp, successive IE,*TM)

    Remark
    • Predicting position from successive IE
    Refer to AMS, Gases and Atomic Structure Revision Notes

    2. Bonding
    • Boiling point/melting point
    • Volatility
    • Electrical conductivity
    • Solubility

    3. Energetics (entropy)
    • Discuss the effects on the entropy of a chemical system by the
    following:
    (i) change in temperature
    (ii) change in phase
    (iii) change in the number of particles (especially for gaseous
    systems)
    (iv) mixing of particles

    Remark
    • Predict the effect of temperature change on the spontaneity of a
    reaction, given standard enthalpy and entropy changes disorderliness”/”ways to arrange particles” are key words.
    Refer to entropy lect notes.

    Using Gibbs equation ΔG=ΔHTΔS.
    Make sure you are comfortable with putting your thoughts into words.

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    #3642

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    TO Master to PERFECTION before A’levels – Part 3

    4. Electrochemistry
    • Predict qualitatively how the value of an electrode potential varies with the concentration of the aqueous ion
    *Could be a disguised change e.g.adding NaOH (aq) to ppt. Mn+.

    Refer to our consolidated list last term

    5. Chem Eqm
    • Apply LCP to deduce qualitatively (from appropriate information)
    the effects of changes in concentration, pressure or temperature, on a system at equilibrium.
    • Deduce whether changes in concentration, pressure or temperature or the presence of a catalyst affect the value of the equilibrium constant for a reaction

    E.g. Given Kc ↑ when temp ↓, predict if reaction is exo/endo.
    6. Ionic Eqm
    • Explain the choice of suitable indicators for acid-base titrations,given appropriate data
    • Explain how buffer solutions (i)control pH (ii) describe and explain
    their uses, including the role of H2CO3/HCO3– in controlling pH in
    blood

    3 main points.
    With equations with SINGLEHEADED arrows

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    #3664

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    Glossary of Terms

    1. State
    A concise answer with little or no supporting argument, e.g. a numerical answer that can be obtained ‘by inspection’ is required.

    2. List
    A number of points, generally each of one word, with no elaboration is required. Where a given number of points is specified, this should not be exceeded

    3. Explain
    Reasoning or some reference to theory is required (depending on the context)

    4. Describe
    State in words (using diagrams where appropriate) the main points of the topic. It is often used with reference either to particular phenomena (where answers should include reference to observations associated) or to particular experiments.

    5. Discuss
    A critical account of the points involved in the topic should be provided.

    6. Outline
    Be concise i.e. restrict the answer to giving the essentials only.

    7. Predict
    Make a logical connection between other pieces of information. Such information may be wholly given in the
    question or may depend on answers extracted in an early part of the question.

    8. Deduce
    Used in a similar way as predict except that some supporting statement is required, e.g. reference to a law/principle, or the necessary reasoning is to be included in the answer

    9. Comment
    It is an open-ended instruction, inviting one to recall or infer points of interest relevant to the context of the question, taking account of the number of marks available.

    10. Suggest
    It is used in two contexts, i.e either to imply that there is no unique answer (e.g. in chemistry, two or more substances may satisfy the given conditions describing an ‘unknown’), or to imply that candidates are
    expected to apply their general knowledge to a ‘novel’ situation, one that may be formally ‘not in syllabus’.

    11. Find
    Can be interpreted as calculate, measure, determine etc

    12. Calculate
    A numerical answer is required. In general working should be shown.
    Note:The misuse of units and/or significant figures is liable to penalty.

    13. Determine
    It implies that the quantity cannot be measured directly but is obtained by calculation, substituting measured and known values of other quantities into a standard formula.

    14. Sketch
    When applied to graph work, the shapes and/or position o the curve need only be qualitatively correct but some quantitative aspects (e.g. passing through the origin, having an intercept at a particular value) may be looked for.

    In diagrams, a simple and freehand drawing is acceptable but care should be taken over proportions and the clear exposition of important details.

    15. Construct
    Often used in relation to chemical equations where one is expected to write a balanced chemical equation,not by factual recall but by analogy or by using information in the question.

    16. Compare
    Both similarities and differences between things or concepts should be provided.

    17. Classify
    Group things based on common characteristics.

    18. Recognise
    Often used to identify facts, characteristics or concepts that are critical (relevant/appropriate) o the understanding of the situation, event, process or phenomenon.

    For exam based questions with solutions please contact @9863 9633

    #3712

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    Power Revision

    A Level Chemistry – 2 hrs Each Lesson

    1. Atoms Molecules and Stoichiometry – 2 lessons
    2. Chemical Bonding – 2 lessons
    3. Chemical Energetics- 2 lessons
    4. Reaction Kinetics – 2 lessons
    5. Chemical Equilibrium – 2 lessons
    6. Ionic Equilibrium – 2 lessons
    7. Introduction Organic/Alkanes/Alkenes – 2 lesssons
    8. Arenes – 1 lessons
    9. Halogen Derivatives – 2 lesons
    10. Hydroxy Compounds – 2 lessons
    11. Carbonyl Compounds – 2 lessons
    12. Carboxylic Acids and Derivatives – 2 lessons

    For enquirers please contact HP 98639633 or Hp 96790479

    #3730

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    Atoms, Molecules and Stoichiometry Part 1

    1 Define the following terms:

    (a) Isotopes are atoms of the same element with the same number of
    protons but different number of neutrons.

    (b) Relative isotopic mass is the number of times the isotope is heavier than 1/12 the mass of an atom of carbon-12.

    (c) Relative atomic mass of an element is the average mass of its atoms in the isotopic mixture, compared to 1/12 the mass of an atom of carbon-
    12.Symbol: Ar

    (d) Relative molecular mass of a molecule is the average mass of its
    molecule compared to 1/12 the mass of an atom of carbon-12.
    Symbol: Mr

    (e) Avogadro’s Law states that equal volumes of all gases under the same
    conditions of temperature and pressure contain the same number of
    molecules/atoms (valid for ideal gases or gases under ideal-like
    conditions).

    (f) Molecular formula gives the actual number of the number of atoms of
    each element present in the compound.

    2 Concept : Limiting Reagent and Dilution

    The process of obtaining iodine from oil field brines involves the following three step process. In the first reaction, 250 g of sodium iodide (NaI) is reacted with 340 g of silver nitrate (AgNO3).

    NaI + AgNO3 —> AgI + NaNO3
    2AgI + Fe —> FeI2 + 2Ag
    2FeI2 + 3Cl2 —> 2FeCl3 + 2I2

    (a) Determine the limiting reagent in the first reaction.
    No of mol NaI available = +25023 127 = 1.6667
    No of mol AgNO3 available = ++340108 14 48 = 2.00
    NaI ≡ AgNO3 ⇒ NaI is the limiting reagent

    (b) Calculate the mass of iodine crystals that can be obtained from the whole process.
    NaI ≡ AgI ≡ ½ FeI2 ≡ ½ I2
    No mol of I2 obtained = ½ ( 1.6667 ) = 0.83335
    Mass of I2 obtained = 0.83335 x 2(127) = 212 g

    (c) The iodine crystals obtained was then dissolved in 50.0 m3 of organic solvent, trichloromethane. Calculate the concentration of iodine solution in g dm-3.

    Concentration of I2 = 50000212 = 0.00423 g dm-3

    (d) In a further experiment, a certain volume of organic solvent, trichloromethane was added to lower the concentration of the iodine solution calculated in (c) to 1.40 X 10-5 mol dm-3. Calculate the volume of trichloromethane added to achieve this effect.

    CoVo = CdVd ⇒ (0.00423254)(50000) = (1.40 x 10-5) (Vd)
    Vd = 59480 dm3
    ∴ Volume of trichloromethane = 59480 – 50000 ≈ 9.48 x 103 dm3

    (e) State the assumption made in determining your answers in parts (c) and (d).

    Iodine is soluble in organic solvent / trichloromethane

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    #3753

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    Atoms, Molecules and Stoichiometry Part 2

    3. (a) Define the term mole.

    The mole is defined as the amount of substance that contains the same
    number of particles as there are atoms in 12 g of pure carbon-12.

    (b) Why is the phrase “the mass of one mole of oxygen” ambiguous?

    It is because the statement can either refer to one mole of oxygen atoms or one mole of oxygen molecules (O2). The statement can also refer to 16O or 18O isotopes.

    (c) A meteorological balloon of 2 m diameter has a volume of 4.19 m3. It floats since it is given an upthrust equal to the mass of air it displaces.

    Calculate:
    (i) the mass of hydrogen in the balloon,

    No of moles of hydrogen gas = 4190 / 23 = 182.2
    Mass of hydrogen = 18.2 x 2 = 364 g

    (ii) the mass of air it displaces,

    Volume of air displaced = 4190 dm^3
    There should be 182.2 mol of air.
    Mass of air displaced = 182.2 x 29 = 5280 g

    (iii) the load the balloon can carry for it just to lift off from the ground.

    Upthrust = mass of air displaced = mass of hydrogen + load
    Load = 5283.8 – 364. = 4919.4 = 4.92 kg

    4. The reaction of silicon tetrachloride with moist ethoxyethane produces two oxochlorides with the formulae Si2OCl6 and Si3O2Cl8. When 0.10 g of one of these oxochlorides completely reacted with water, all of its chlorine was converted into chloride ions, and produced 0.303 g of silver chloride precipitate when an excess of aqueous silver nitrate was added.
    Deduce the identity of the oxochloride.

    AgCl Ξ Cl
    No of mol of Cl in 0.303 g of AgCl
    = 0.303/(108+35.5)
    = 2.11 x 10-3

    Relative molecular mass of Si2OCl6
    = 2×28.1+16+6×35.5
    = 285.2

    Relative molecular mass of Si3O2Cl8
    = 3×28.1+2×16+8X35.5
    = 400.3

    No of mole of Cl in 0.1g of Si2OCl6
    = (0.1/285.2) x 6
    = 2.10 x 10-3

    No of mole of Cl in 0.1g of Si3O2Cl8
    = (0.1/400.3) x 6
    = 2.00 x 10-3

    The oxochloride is Si2OCl6.

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    #3776

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    Atoms, Molecules and Stoichiometry Part 3

    Concept: Determining Formula of compound

    A metal hydroxide, M(OH)n, is one of the products formed in an electrochemical cell used to power golf trolleys. 50.0 cm3 of solution containing 0.028 mol of M(OH)n in 1 dm3 requires 21.00 cm3 of sulfuric acid for complete neutralisation. The sulfuric acid contains 0.2 g of hydrogen ions in 1 dm3.

    (a) Calculate the number of moles of sulfuric acid that will react with 1 mole of M(OH)n.

    H2SO4 ≡ 2H+

    No. of moles of H+ in 1 dm3 = 0.2/1.0 = 0.200

    [H2SO4] = 0.2002= 0.100 mol dm-3

    No. of moles of H2SO4 in 21.10 cm3 = (21.00×0.100)/1000 = 2.10 x 10-3

    No. of moles of M(OH)n in 50.0 cm3 = (50.00×0.028)/1000 = 1.40 x 10-3

    No. of moles of H2SO4/No. of moles of M(OH)n
    = 2.10 x 10-3/1.40 x 10-3
    = 1.5

    1 mol of M(OH)n reacts with 1.50 mol of H2SO4

    (b) Hence, determine the value of n.

    2M(OH)n (aq) + nH2SO4 (aq) → M2(SO4)n (aq) + 2nH2O (l)
    Mole ratio: 2M(OH)n ≡ nH2SO4
    M(OH)n ≡ (n/2)H2SO4
    n/2 = 3/2
    n = 3

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    #3788

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    Atoms, Molecules and Stoichiometry Part 4

    Concept: Back titration and calculation of Percentage Purity

    The chemical used for detecting proteins, biuret reagent, H2NCONHCONH2, can be formed by heating urea, (NH2)2CO.

    2(NH2)2CO → H2NCONHCONH2 + NH3 — (1)

    3.88 g of impure sample of urea (NH2)2CO was heated strongly above its melting point. The ammonia liberated was absorbed in 32.0 cm3 of 2.00 mol dm-3 sulfuric acid. The resulting solution was made up to 500 cm3 with distilled water. 25.0 cm3 of the solution required 25.50 cm3 of 0.20 mol dm-3 sodium hydroxide for neutralization, using methyl orange as an indicator.
    (a) Calculate the percentage purity of urea in the sample. [4]

    2NH3 + H2SO4 (excess) → (NH4)2SO4 — (2)

    H2SO4 (remaining) + 2NaOH → Na2SO4 + 2H2O — (3)

    No. of moles of NaOH in 25.50 cm3 = 25.50 x (0.20/)1000 = 5.10 x 10-3

    From (3) 2NaOH Ξ H2SO4 (remaining)

    No. of moles of H2SO4 in 25.0 cm3 = (5.10 x 10-3)/2 = 2.55 x 10-3

    No. of moles of H2SO4 in 500 cm3 = 2.55 x 10-3 x (500/25) = 0.0510

    Initial no. of moles of H2SO4 in 32.0 cm3 = 2.00 x (32/1000) = 0.0640

    No. of moles of H2SO4 reacted with NH3 = 0.0640 – 0.0510 = 0.0130

    From (2) and (1) 2NH3 Ξ H2SO4 (reacted) Ξ 4(NH2)2CO

    Mass of urea, (NH2)2CO produced = 0.0130 x 4 x 60 = 3.12 g

    Percentage purity of urea, (NH2)2CO = (3.12/3.88) × 100 = 80.4 %

    (b) Determine the maximum number of hydrogen atoms present in the biuret reagent after the impure sample of urea was heated. [2]

    From (2) and (1) 2NH3 Ξ H2SO4 (reacted) Ξ 2H2NCONHCONH2

    No. of moles of H2NCONHCONH2 = 0.0130 x 2 = 0.0260

    H2NCONHCONH2 Ξ 5H

    No. of H atoms in biuret reagent = 0.0260 x 5 x 6.02 x 10^23
    = 7.83 x 10^22

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