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ETEA Medical Test Chemistry Syllabus

The Chemistry syllabus for ETEA Medical Test in full: 12 chapters and 39 topics, every one listed below, plus 50 Chemistry flashcards with answers.

12Chapters
39Topics
50Flashcards

Chemistry syllabus: all 12 chapters

Expand any chapter to see its topics and sub-topics.

  1. Fundamentals of Chemistry

     3 topics
    • Atomic Mass and Mole Concept
    • Stoichiometry
    • Empirical and Molecular Formula

  2. Atomic Structure

     4 topics
    • Subatomic Particles
    • Quantum Numbers and Orbitals
    • Electronic Configuration
    • Periodic Trends

  3. Chemical Bonding

     4 topics
    • Ionic and Covalent Bonds
    • Hybridization
    • VSEPR Theory and Molecular Geometry
    • Hydrogen Bonding

  4. States of Matter

     3 topics
    • Gas Laws
    • Liquids and Intermolecular Forces
    • Solids and Crystal Lattice

  5. Chemical Equilibrium

     3 topics
    • Law of Mass Action
    • Le Chatelier's Principle
    • Solubility Product

  6. Acids, Bases and Salts

     3 topics
    • Theories of Acids and Bases
    • pH and Buffers
    • Salt Hydrolysis

  7. Thermochemistry

     3 topics
    • Enthalpy and Hess's Law
    • First Law of Thermodynamics
    • Spontaneity and Free Energy

  8. Electrochemistry

     3 topics
    • Redox Reactions
    • Electrolysis
    • Galvanic Cells and Electrode Potential

  9. Chemical Kinetics

     3 topics
    • Rate of Reaction
    • Order of Reaction
    • Factors Affecting Reaction Rate

  10. Fundamentals of Organic Chemistry

     3 topics
    • Hydrocarbons
    • Functional Groups and Isomerism
    • Alkyl Halides

  11. Organic Functional Groups

     4 topics
    • Alcohols and Phenols
    • Aldehydes and Ketones
    • Carboxylic Acids and Derivatives
    • Amines

  12. Biochemistry and Industrial Chemistry

     3 topics
    • Carbohydrates and Proteins
    • Lipids and Nucleic Acids
    • Macromolecules and Polymers

ETEA Medical Test Chemistry flashcards — 50 questions with answers

  1. Define one mole in terms of the number of particles (Avogadro's number).

    One mole is the amount of substance containing 6.022 x 10^23 particles (Avogadro's number).

  2. What is the relationship between the relative atomic mass of an element and the mass of one mole of its atoms?

    The relative atomic mass expressed in grams equals the mass of one mole (the molar mass) of that element's atoms.

  3. How is the number of moles calculated from a given mass and molar mass?

    Number of moles = given mass (g) / molar mass (g/mol).

  4. What volume does one mole of any ideal gas occupy at STP?

    22.4 dm3 (22.4 litres) at STP (0 C and 1 atm).

  5. How do you calculate the number of particles from the number of moles?

    Number of particles = number of moles x 6.022 x 10^23 (Avogadro's number).

  6. Define stoichiometry.

    Stoichiometry is the quantitative study of the relationships (mass, mole, and volume ratios) between reactants and products in a balanced chemical equation.

  7. In stoichiometry, what is a limiting reactant?

    The limiting reactant is the reactant that is completely consumed first, thereby limiting the amount of product formed.

  8. How is percentage yield calculated?

    Percentage yield = (actual yield / theoretical yield) x 100.

  9. What is the mole ratio used for in stoichiometric calculations?

    The mole ratio, taken from the coefficients of the balanced equation, is used to convert moles of one substance into moles of another.

  10. Define empirical formula.

    The empirical formula is the simplest whole-number ratio of the atoms of each element in a compound.

  11. Define molecular formula.

    The molecular formula shows the actual number of atoms of each element present in one molecule of a compound.

  12. How is the molecular formula obtained from the empirical formula?

    Molecular formula = empirical formula x n, where n = molar mass / empirical formula mass.

  13. Outline the steps to determine an empirical formula from percentage composition.

    Divide each element's percentage by its atomic mass to get moles, then divide all values by the smallest to get the simplest whole-number ratio.

  14. State the relative charge and relative mass of a proton.

    A proton has a relative charge of +1 and a relative mass of 1 amu.

  15. State the relative charge and relative mass of a neutron.

    A neutron has a relative charge of 0 (neutral) and a relative mass of 1 amu.

  16. State the relative charge and relative mass of an electron.

    An electron has a relative charge of -1 and a relative mass of about 1/1836 amu (negligible).

  17. What does the atomic number (Z) of an element represent?

    The atomic number equals the number of protons in the nucleus (and the number of electrons in a neutral atom).

  18. Define isotopes.

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

  19. What does the principal quantum number (n) describe?

    The principal quantum number (n) describes the main energy level (shell) and the relative size/energy of the orbital; it has values 1, 2, 3, ...

  20. What does the azimuthal (angular momentum) quantum number (l) determine?

    It determines the shape of the orbital (subshell); l = 0 (s), 1 (p), 2 (d), 3 (f), ranging from 0 to n-1.

Show the other 30 Chemistry flashcards

  1. What does the magnetic quantum number (m_l) specify?

    It specifies the orientation of the orbital in space, with values ranging from -l to +l, including 0.

  2. What are the two possible values of the spin quantum number (m_s)?

    +1/2 and -1/2, representing the two opposite spin directions of an electron.

  3. How many orbitals and electrons can the p subshell hold?

    The p subshell has 3 orbitals and can hold a maximum of 6 electrons.

  4. How many orbitals and electrons can the d subshell hold?

    The d subshell has 5 orbitals and can hold a maximum of 10 electrons.

  5. State the Aufbau principle.

    Electrons fill orbitals starting from the lowest available energy level before occupying higher ones.

  6. State the Pauli exclusion principle.

    No two electrons in an atom can have the same set of all four quantum numbers; an orbital holds at most two electrons with opposite spins.

  7. State Hund's rule.

    Electrons occupy degenerate orbitals singly with parallel spins before any orbital is doubly occupied.

  8. Write the ground-state electronic configuration of chlorine (Z = 17).

    1s2 2s2 2p6 3s2 3p5.

  9. Why does chromium (Z = 24) have the configuration [Ar] 3d5 4s1 instead of 3d4 4s2?

    A half-filled 3d5 subshell with 4s1 gives extra stability due to symmetrical distribution and exchange energy.

  10. Define ionization energy.

    Ionization energy is the minimum energy required to remove the most loosely bound electron from one mole of gaseous atoms.

  11. How does atomic radius vary across a period and down a group?

    Atomic radius decreases across a period (left to right) and increases down a group.

  12. How does ionization energy vary across a period and down a group?

    Ionization energy increases across a period and decreases down a group.

  13. How does electronegativity vary across a period and down a group?

    Electronegativity increases across a period and decreases down a group; fluorine is the most electronegative element.

  14. Define electron affinity.

    Electron affinity is the energy change when one mole of gaseous atoms gains one mole of electrons to form gaseous anions.

  15. Define an ionic bond.

    An ionic bond is the electrostatic force of attraction between oppositely charged ions formed by the complete transfer of electrons from a metal to a non-metal.

  16. Define a covalent bond.

    A covalent bond is the bond formed by the mutual sharing of electron pairs between two non-metal atoms.

  17. Distinguish between a polar and a non-polar covalent bond.

    A polar covalent bond forms between atoms of different electronegativities (unequal sharing); a non-polar covalent bond forms between identical atoms (equal sharing).

  18. What is a coordinate (dative) covalent bond?

    A covalent bond in which both shared electrons are donated by one atom, e.g. in the ammonium ion (NH4+).

  19. What type of hybridization and geometry are associated with sp3 hybridization?

    sp3 hybridization gives 4 hybrid orbitals with tetrahedral geometry and bond angles of 109.5 degrees (e.g. CH4).

  20. What geometry and bond angle result from sp2 hybridization?

    sp2 hybridization gives 3 hybrid orbitals with trigonal planar geometry and bond angles of 120 degrees (e.g. BF3, ethene).

  21. What geometry and bond angle result from sp hybridization?

    sp hybridization gives 2 hybrid orbitals with linear geometry and a bond angle of 180 degrees (e.g. BeCl2, ethyne).

  22. How many sigma and pi bonds are present in a carbon-carbon double bond?

    A double bond contains 1 sigma bond and 1 pi bond.

  23. What is the main postulate of VSEPR theory?

    Electron pairs around a central atom arrange themselves to be as far apart as possible to minimize repulsion, determining the molecule's shape.

  24. Order the repulsion strengths between lone pairs (lp) and bond pairs (bp) in VSEPR theory.

    lp-lp > lp-bp > bp-bp.

  25. Why is the bond angle in water (104.5 degrees) less than in methane (109.5 degrees)?

    Water has two lone pairs on oxygen whose stronger repulsion compresses the H-O-H bond angle below the tetrahedral value.

  26. What is the molecular geometry of ammonia (NH3) and its bond angle?

    Trigonal pyramidal with a bond angle of about 107 degrees (one lone pair on nitrogen).

  27. Define hydrogen bonding.

    Hydrogen bonding is a strong dipole-dipole attraction between a hydrogen atom bonded to a highly electronegative atom (F, O, or N) and a lone pair on another electronegative atom.

  28. Why does water have an unusually high boiling point compared to H2S?

    Water molecules form extensive hydrogen bonds (O-H...O), requiring more energy to break than the weaker forces in H2S, raising its boiling point.

  29. State the combined gas law equation.

    P1V1/T1 = P2V2/T2, where temperature is in kelvin.

  30. State the ideal gas equation and identify each term.

    PV = nRT, where P = pressure, V = volume, n = moles, R = gas constant, and T = absolute temperature in kelvin.

ETEA Medical Test Chemistry FAQ

What is in the ETEA Medical Test Chemistry syllabus?

ETEA Medical Test Chemistry is split into 12 chapters: Fundamentals of Chemistry, Atomic Structure, Chemical Bonding, States of Matter, Chemical Equilibrium, Acids, Bases and Salts, Thermochemistry, Electrochemistry, Chemical Kinetics, Fundamentals of Organic Chemistry, Organic Functional Groups and Biochemistry and Industrial Chemistry. Together they contain 39 topics.

How many chapters are in ETEA Medical Test Chemistry?

12 chapters, containing 39 topics. The largest chapter is Atomic Structure, with 4 topics.

Are there flashcards for ETEA Medical Test Chemistry?

Yes — a 50-card Chemistry deck, printed in full in the flashcards section of this page and available in the Examius app with spaced repetition. The free plan covers one exam of your choice, including its flashcards; Examius Pro, a one-time purchase, unlocks every other exam and deck.