About this EAMCET Medical Chemistry practice test

EAMCET Medical Chemistry carries the same 40-of-160 weight as Botany and Zoology, yet it's the section most Agriculture and Pharmacy aspirants under-revise, since biology naturally gets more attention. The syllabus spans Physical Chemistry (thermodynamics, equilibrium, electrochemistry, kinetics), Organic Chemistry (hydrocarbons, functional groups, biomolecules, polymers) and Inorganic Chemistry (atomic structure, bonding, s- and p-block elements) — a mix of fast calculations and pure recall, in roughly equal parts. This set covers both halves properly, with a full worked explanation behind every answer so you know the reasoning, not just the letter. Treat this section with the same seriousness as your biology papers, and the marks follow.

EAMCET Medical Chemistry Practice Test sample questions

These starter questions help you launch a chemistry mock test quickly. Swap them with your own worksheet, notebook, or textbook questions any time.

1. 1. According to Bohr's model, what is the energy of an electron in the ground state of a hydrogen atom?

2. 2. What is the maximum number of electrons that can be accommodated in the M shell (n = 3)?

3. 3. What is the hybridization of each carbon atom in ethyne (C2H2)?

4. 4. Which molecule is nonpolar despite having polar bonds: BF3 or NH3?

5. 5. What is the molecular shape and approximate bond angle of ammonia (NH3) according to VSEPR theory?

6. 6. A gas occupies 2 L at 300 K and 1 atm. What volume will it occupy at 600 K at the same pressure, assuming ideal behavior?

7. 7. What is the value of the compressibility factor (Z) for an ideal gas at all temperatures and pressures?

8. 8. For an isothermal process involving an ideal gas, what is the change in internal energy (ΔU)?

9. 9. For the reaction N2(g) + 3H2(g) → 2NH3(g), ΔH = -92 kJ/mol. Is this reaction exothermic or endothermic?

10. 10. According to the second law of thermodynamics, what is the sign of the entropy change of the universe for a spontaneous process?

11. 11. For the reaction A(g) ⇌ B(g), Kc = 4. If the initial concentration of A is 1 mol/L and no B is present initially, what is the equilibrium concentration of B?

12. 12. According to Le Chatelier's principle, in which direction does increasing pressure shift the equilibrium N2(g) + 3H2(g) ⇌ 2NH3(g)?

13. 13. For the reaction 2SO2(g) + O2(g) ⇌ 2SO3(g), what is the relationship between Kp and Kc?

14. 14. What is the standard cell potential (E°cell) of a Daniell cell, given E°(Cu2+/Cu) = +0.34 V and E°(Zn2+/Zn) = -0.76 V?

15. 15. How many faradays of electricity are required to deposit 1 mole of aluminum metal from Al3+ ions?

16. 16. What happens to the electrical conductivity of a metallic conductor as its temperature increases?

17. 17. For a first-order reaction, what is the relationship between half-life (t1/2) and the rate constant (k)?

18. 18. If the rate of a reaction doubles when the concentration of one reactant is doubled, with all else constant, what is the order of the reaction with respect to that reactant?

19. 19. Why do alkali metals impart characteristic colors during a flame test?

20. 20. Which s-block compound is commonly known as "washing soda," and what is its chemical formula?

21. 21. What is the trend in thermal stability of the hydrides of group 15 elements from NH3 to BiH3?

22. 22. What product is formed when propene reacts with HBr in the presence of peroxide?

23. 23. What is the hybridization of each carbon atom in benzene, and how many sigma bonds does each carbon form?

24. 24. Which reagent distinguishes an aldehyde from a ketone, and what is the positive test result for an aldehyde?

25. 25. What product is formed when ethanol is oxidized with acidified potassium dichromate, if the product is distilled off as it forms?

26. 26. What is observed when a carboxylic acid reacts with sodium bicarbonate solution?

27. 27. What type of glycosidic linkage joins the two glucose units in maltose?

28. 28. What type of bonding primarily holds the two strands of a DNA double helix together?

29. 29. What type of polymerization forms Nylon-6,6, and what are its two monomers?

30. 30. What is the primary chemical cause of ozone layer depletion in the stratosphere?

Answer key & quick explanations

Short answers for the sample questions above. Use this to self-check before generating a fresh AI-built mock test.

1. 1. According to Bohr's model, what is the energy of an electron in the ground state of a hydrogen atom?

   -13.6 eV

   Bohr's formula gives En = -13.6/n² eV for hydrogen-like species. Substituting n = 1 for the ground state yields -13.6 eV. The negative sign simply reflects that the electron is bound to the nucleus rather than free.

2. 2. What is the maximum number of electrons that can be accommodated in the M shell (n = 3)?

   18

   The formula 2n² gives the maximum electron capacity of a shell, so for n = 3 this is 2(9) = 18. This matches the combined capacity of the 3s, 3p, and 3d subshells (2 + 6 + 10 electrons).

3. 3. What is the hybridization of each carbon atom in ethyne (C2H2)?

   sp hybridization

   Each carbon forms one sigma bond to hydrogen and one sigma bond to the other carbon using two sp hybrid orbitals, leaving two unhybridized p orbitals to form the two pi bonds of the triple bond. This arrangement produces the molecule's linear, 180° geometry.

4. 4. Which molecule is nonpolar despite having polar bonds: BF3 or NH3?

   BF3

   BF3 has a trigonal planar shape with three identical B-F bond dipoles arranged symmetrically, so their vectors cancel and the net dipole moment is zero. NH3, by contrast, is pyramidal with a lone pair, so its bond dipoles do not cancel and it remains polar.

5. 5. What is the molecular shape and approximate bond angle of ammonia (NH3) according to VSEPR theory?

   Trigonal pyramidal, about 107°

   Nitrogen in NH3 has three bonding pairs and one lone pair, giving a distorted tetrahedral electron geometry. Because lone pair-bond pair repulsion exceeds bond pair-bond pair repulsion, the H-N-H angle is compressed from the ideal 109.5° to roughly 107°.

6. 6. A gas occupies 2 L at 300 K and 1 atm. What volume will it occupy at 600 K at the same pressure, assuming ideal behavior?

   4 L

   At constant pressure, Charles's Law states V1/T1 = V2/T2, so V2 = 2 L × (600 K/300 K) = 4 L. Doubling the absolute temperature at fixed pressure simply doubles the gas volume.

7. 7. What is the value of the compressibility factor (Z) for an ideal gas at all temperatures and pressures?

   Z = 1

   Z is defined as PV/nRT, and for an ideal gas this ratio equals exactly 1 under all conditions by definition. Real gases deviate from Z = 1, typically dropping below it at moderate pressures and rising above it at very high pressures, due to intermolecular forces and finite molecular volume.

8. 8. For an isothermal process involving an ideal gas, what is the change in internal energy (ΔU)?

   Zero

   Internal energy of an ideal gas is a function of temperature alone. Since temperature stays constant throughout an isothermal process, ΔU = 0, and any heat absorbed by the gas is entirely converted into work.

9. 9. For the reaction N2(g) + 3H2(g) → 2NH3(g), ΔH = -92 kJ/mol. Is this reaction exothermic or endothermic?

   Exothermic

   A negative sign on ΔH means the system releases heat to its surroundings, which is the defining feature of an exothermic reaction. Here, 92 kJ of energy is given out per mole of reaction as ammonia is formed.

10. 10. According to the second law of thermodynamics, what is the sign of the entropy change of the universe for a spontaneous process?

    Positive (ΔS_universe > 0)

    The second law requires that the combined entropy of a system and its surroundings must increase for any spontaneous process to occur. A process where ΔS_universe would be negative simply cannot happen on its own.

11. 11. For the reaction A(g) ⇌ B(g), Kc = 4. If the initial concentration of A is 1 mol/L and no B is present initially, what is the equilibrium concentration of B?

    0.8 mol/L

    Letting x be the concentration of B formed, Kc = x/(1-x) = 4 leads to 4 - 4x = x, so x = 0.8 mol/L. Correspondingly, the leftover concentration of A at equilibrium is 0.2 mol/L.

12. 12. According to Le Chatelier's principle, in which direction does increasing pressure shift the equilibrium N2(g) + 3H2(g) ⇌ 2NH3(g)?

    Toward the product side (forward direction, favoring NH3)

    The reactant side has 4 moles of gas while the product side has only 2 moles, so raising the pressure pushes the equilibrium toward the side with fewer gas molecules. This favors greater NH3 formation, which is why the industrial Haber process is run at high pressure.

13. 13. For the reaction 2SO2(g) + O2(g) ⇌ 2SO3(g), what is the relationship between Kp and Kc?

    Kp = Kc(RT)^-1

    The relation Kp = Kc(RT)^Δn applies, where Δn is moles of gaseous product minus gaseous reactant: 2 - 3 = -1. Substituting gives Kp = Kc(RT)^-1, so Kp is numerically smaller than Kc at any temperature above absolute zero.

14. 14. What is the standard cell potential (E°cell) of a Daniell cell, given E°(Cu2+/Cu) = +0.34 V and E°(Zn2+/Zn) = -0.76 V?

    1.10 V

    Cell potential is calculated as E°cathode minus E°anode. Copper is reduced at the cathode and zinc is oxidized at the anode, so E°cell = 0.34 - (-0.76) = 1.10 V, and the positive value confirms the reaction proceeds spontaneously as written.

15. 15. How many faradays of electricity are required to deposit 1 mole of aluminum metal from Al3+ ions?

    3 faradays

    The reduction half-reaction Al3+ + 3e- → Al shows that three moles of electrons are needed per mole of aluminum deposited. Since one faraday corresponds to one mole of electrons, depositing 1 mole of Al requires 3 faradays, roughly 289,500 coulombs of charge.

16. 16. What happens to the electrical conductivity of a metallic conductor as its temperature increases?

    It decreases

    Heating a metal increases the thermal vibration of its lattice ions, which scatters the freely moving conduction electrons more frequently and raises resistance. This behavior is opposite to that of electrolytic solutions and semiconductors, whose conductivity rises with temperature.

17. 17. For a first-order reaction, what is the relationship between half-life (t1/2) and the rate constant (k)?

    t1/2 = 0.693/k

    Integrating the first-order rate law shows that t1/2 equals ln2 divided by k, which numerically is 0.693/k. A distinctive feature of first-order kinetics is that this half-life value does not depend on the starting concentration of the reactant.

18. 18. If the rate of a reaction doubles when the concentration of one reactant is doubled, with all else constant, what is the order of the reaction with respect to that reactant?

    First order

    Writing rate = k[A]^n and doubling [A] to get a doubled rate gives 2 = 2^n, which is only satisfied when n = 1. This directly identifies the reaction as first order in that particular reactant.

19. 19. Why do alkali metals impart characteristic colors during a flame test?

    The heat of the flame excites the loosely held valence electron, which emits visible light of a specific wavelength when it falls back to the ground state

    Alkali metals have a single outer electron that is easily promoted to a higher energy level by the thermal energy of a flame. As this electron relaxes back down, it releases the absorbed energy as a photon of a fixed wavelength, giving sodium its yellow flame and potassium its lilac flame.

20. 20. Which s-block compound is commonly known as "washing soda," and what is its chemical formula?

    Sodium carbonate decahydrate, Na2CO3·10H2O

    Washing soda is the everyday name for the hydrated form of sodium carbonate, industrially manufactured by the Solvay process. It finds wide use as a water softener and as a raw material in the glass and detergent industries.

21. 21. What is the trend in thermal stability of the hydrides of group 15 elements from NH3 to BiH3?

    Thermal stability decreases down the group (NH3 is most stable, BiH3 least stable)

    Moving down group 15 from nitrogen to bismuth, the central atom grows larger and the E-H bond length increases, which weakens the bond dissociation enthalpy. Consequently NH3 withstands heating well, while BiH3 decomposes into its elements very readily.

22. 22. What product is formed when propene reacts with HBr in the presence of peroxide?

    1-bromopropane

    Peroxides promote a free-radical chain mechanism in which the bromine atom ends up on the terminal, less-substituted carbon, giving 1-bromopropane rather than the Markovnikov product. This reversal is called the peroxide or Kharasch effect and is observed only with HBr, not with HCl or HI.

23. 23. What is the hybridization of each carbon atom in benzene, and how many sigma bonds does each carbon form?

    sp2 hybridization; 3 sigma bonds per carbon

    Every carbon in the benzene ring is sp2 hybridized and forms sigma bonds to two neighboring ring carbons and to one hydrogen atom, totaling three sigma bonds. The leftover unhybridized p orbital on each carbon overlaps sideways to create the delocalized pi system responsible for aromatic stability.

24. 24. Which reagent distinguishes an aldehyde from a ketone, and what is the positive test result for an aldehyde?

    Tollens' reagent; deposition of a bright silver mirror

    Tollens' reagent, ammoniacal silver nitrate, oxidizes the aldehyde group to a carboxylate while the silver ion is itself reduced to metallic silver that coats the test tube as a mirror. Ketones have no hydrogen on the carbonyl carbon to be oxidized this way, so they give a negative result.

25. 25. What product is formed when ethanol is oxidized with acidified potassium dichromate, if the product is distilled off as it forms?

    Acetaldehyde (ethanal)

    Acidified K2Cr2O7 first oxidizes the primary alcohol ethanol to the aldehyde stage. Removing the volatile aldehyde by distillation as soon as it forms prevents its further oxidation to acetic acid, so acetaldehyde is isolated as the main product.

26. 26. What is observed when a carboxylic acid reacts with sodium bicarbonate solution?

    Brisk effervescence due to the release of carbon dioxide gas

    Carboxylic acids react with NaHCO3 to form the sodium salt of the acid, water, and carbon dioxide, which escapes as visible bubbling. This simple test is useful because phenols, though also weakly acidic, do not react with sodium bicarbonate to release CO2.

27. 27. What type of glycosidic linkage joins the two glucose units in maltose?

    α-1,4-glycosidic linkage

    Maltose consists of two α-D-glucose units connected through an α-1,4 bond between the C1 anomeric carbon of one unit and the C4 carbon of the other. Because one anomeric carbon remains free, maltose still behaves as a reducing sugar.

28. 28. What type of bonding primarily holds the two strands of a DNA double helix together?

    Hydrogen bonding between complementary base pairs

    The antiparallel strands of DNA are linked through hydrogen bonds formed between specific base pairs: adenine with thymine through two hydrogen bonds, and guanine with cytosine through three. This selective complementary pairing underlies the accuracy of DNA replication.

29. 29. What type of polymerization forms Nylon-6,6, and what are its two monomers?

    Condensation polymerization; hexamethylenediamine and adipic acid

    Nylon-6,6 arises from condensation polymerization between hexamethylenediamine and adipic acid, with water eliminated as each amide linkage forms. Its name reflects the six carbon atoms present in each of the two monomer units.

30. 30. What is the primary chemical cause of ozone layer depletion in the stratosphere?

    Chlorine free radicals released from chlorofluorocarbons (CFCs)

    CFCs are chemically inert enough to survive the journey to the stratosphere, where ultraviolet radiation breaks them apart and releases chlorine free radicals. Each such radical can catalytically destroy thousands of ozone molecules in a chain reaction before it is finally deactivated.

Curriculum Mapping & Learning Guide

Use this breakdown to identify which skills each question tests and guide post-test review.

EAMCET Medical Chemistry units covered

1. Chapter 1: Atomic Structure
2. Chapter 2: Chemical Bonding and Molecular Structure
3. Chapter 3: States of Matter
4. Chapter 4: Thermodynamics
5. Chapter 5: Chemical Equilibrium
6. Chapter 6: Electrochemistry
7. Chapter 7: Chemical Kinetics
8. Chapter 8: s-Block and p-Block Elements
9. Chapter 9: Hydrocarbons
10. Chapter 10: Organic Functional Groups (Alcohols, Aldehydes, Ketones, Acids)
11. Chapter 11: Biomolecules (Carbohydrates, Proteins, Nucleic Acids)
12. Chapter 12: Polymers and Environmental Chemistry

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