BioMedical Admissions Test (BMAT) Section 2: Scientific Knowledge and Applications — Chemistry Syllabus

Section 2: Scientific Knowledge and Applications — Chemistry syllabus: all 4 chapters

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

1. 1

   Atomic Structure and the Periodic Table

   3 topics

   • Structure of the atom
     • Protons, neutrons and electrons
     • Mass number, atomic number and isotopes
   • Electron configuration
     • Energy levels and shells
     • Relating configuration to group and period
   • Periodic trends
     • Groups, periods and reactivity
     • Group 1, Group 7 and Group 0 properties
2. 2

   Bonding and Structure

   4 topics

   • Ionic bonding
     • Formation and properties of ionic compounds
   • Covalent bonding
     • Simple molecules and giant covalent structures
   • Metallic bonding
     • The electron sea model and conductivity
   • Structure and properties
     • Relating bonding to melting point, conductivity and solubility
3. 3

   Chemical Reactions and Quantitative Chemistry

   3 topics

   • Balancing equations
     • Conservation of mass
     • State symbols and ionic equations
   • The mole and stoichiometry
     • Relative atomic and formula mass
     • Moles, mass and Avogadro's constant
     • Reacting mass and limiting reagent calculations
   • Concentration and yield
     • Concentration of solutions
     • Percentage yield and atom economy
4. 4

   Types of Reaction

   5 topics

   • Acids, bases and salts
     • The pH scale and neutralisation
     • Reactions of acids with metals, carbonates and bases
   • Redox and reactivity
     • Oxidation and reduction in terms of electrons
     • The reactivity series and displacement
   • Electrolysis
     • Electrolysis of molten and aqueous compounds
     • Products at electrodes
   • Rates and energetics
     • Factors affecting rate of reaction
     • Exothermic and endothermic reactions
   • Organic chemistry basics
     • Hydrocarbons and homologous series
     • Combustion and cracking

BioMedical Admissions Test (BMAT) Section 2: Scientific Knowledge and Applications — Chemistry flashcards — 50 questions with answers

1. What three subatomic particles make up an atom, and what are their relative charges and masses?

   Protons (charge $+1$, mass $1$), neutrons (charge $0$, mass $1$) and electrons (charge $-1$, mass $\approx \frac{1}{1836}$). Protons and neutrons sit in the nucleus; electrons occupy shells around it.

2. Define atomic number ($Z$) and mass number ($A$).

   Atomic number $Z$ = number of protons in the nucleus (also defines the element). Mass number $A$ = number of protons + number of neutrons. Number of neutrons $= A - Z$.

3. What are isotopes?

   Isotopes are atoms of the same element (same number of protons, same $Z$) but with different numbers of neutrons, so they have different mass numbers $A$. They have identical chemical properties but slightly different physical properties.

4. How is relative atomic mass ($A_r$) calculated from isotopic abundances?

   $$A_r = \frac{\sum (\text{isotope mass} \times \text{\% abundance})}{100}$$ It is the weighted mean mass of an atom relative to $\frac{1}{12}$ of a carbon-12 atom.

5. State the maximum number of electrons each of the first three electron shells can hold.

   Shell 1 (n=1): 2 electrons; shell 2 (n=2): 8 electrons; shell 3 (n=3): 8 (in basic GCSE/BMAT model) up to 18. General rule: $2n^{2}$.

6. Write the full electron configuration of a calcium atom ($Z=20$) using sub-shell notation.

   $\text{Ca}: 1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{6}\,4s^{2}$. The $4s$ sub-shell fills before $3d$.

7. In what order do the $4s$ and $3d$ sub-shells fill, and what is the maximum capacity of $s$, $p$ and $d$ sub-shells?

   $4s$ fills before $3d$ (lower energy). Capacities: $s$ holds 2, $p$ holds 6, $d$ holds 10 electrons.

8. How does atomic radius change across a period and down a group, and why?

   Across a period: decreases (more protons / increasing nuclear charge pull electrons in the same shell closer). Down a group: increases (more occupied shells and greater shielding outweigh increased nuclear charge).

9. Define first ionisation energy and state its general trends across a period and down a group.

   First ionisation energy is the energy needed to remove one mole of electrons from one mole of gaseous atoms: $\ce{X(g) -> X^+(g) + e^-}$. It increases across a period (greater nuclear charge) and decreases down a group (more shielding, electron further out).

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

    Electronegativity (an atom's ability to attract a bonding pair of electrons) increases across a period and decreases down a group. Fluorine is the most electronegative element.

11. What are the characteristic properties of Group 1 (alkali metals) and how does reactivity change down the group?

    Soft, low-density metals with one outer electron; react with water to give an alkaline hydroxide and hydrogen. Reactivity increases down the group because the outer electron is lost more easily.

12. What are the characteristic properties of Group 7 (halogens) and how does reactivity change down the group?

    Diatomic non-metals ($\ce{F2, Cl2, Br2, I2}$) with seven outer electrons; they gain one electron to form $-1$ ions. Reactivity decreases down the group because the incoming electron is gained less easily.

13. Why are the Group 0 (noble) gases unreactive?

    They have a full outer shell of electrons (a stable octet, or 2 for helium), so they have little tendency to gain, lose or share electrons.

14. Describe how ionic bonding forms and the type of elements involved.

    Ionic bonding forms when a metal transfers electrons to a non-metal, producing oppositely charged ions held together by strong electrostatic attraction. E.g. $\ce{Na -> Na+ + e-}$ and $\ce{Cl + e- -> Cl-}$ giving $\ce{NaCl}$.

15. Why do ionic compounds have high melting points and conduct electricity only when molten or in solution?

    Strong electrostatic forces throughout the giant lattice require much energy to break, giving high melting points. Ions are fixed in the solid (no conduction) but become free to move when molten or dissolved, allowing conduction.

16. Describe covalent bonding and give an example of a double bond.

    A covalent bond is a shared pair of electrons between two non-metal atoms. A double bond is two shared pairs, e.g. in oxygen $\ce{O=O}$ ($\ce{O2}$) or carbon dioxide $\ce{O=C=O}$.

17. Contrast the properties of simple molecular substances with giant covalent (macromolecular) structures.

    Simple molecular (e.g. $\ce{H2O}$, $\ce{CO2}$): low melting/boiling points (weak intermolecular forces), non-conducting. Giant covalent (e.g. diamond, $\ce{SiO2}$): very high melting points (strong covalent network), generally non-conducting (except graphite).

18. Describe the bonding and structure of metals (metallic bonding).

    A lattice of positive metal ions in a 'sea' of delocalised outer electrons; the electrostatic attraction between ions and delocalised electrons is the metallic bond.

19. Explain why metals are good conductors of electricity and are malleable.

    Delocalised electrons are free to move and carry charge, giving electrical (and thermal) conductivity. Layers of ions can slide over one another while the bonding remains, so metals are malleable and ductile.

20. Compare the structures of diamond and graphite, both made only of carbon.

    Diamond: each C bonded to 4 others in a rigid 3D tetrahedral lattice — very hard, non-conducting. Graphite: each C bonded to 3 others in layers of hexagons with delocalised electrons between layers — soft (layers slide) and electrically conducting.

BioMedical Admissions Test (BMAT) Section 2: Scientific Knowledge and Applications — Chemistry FAQ