84257 - Chemistry

Course Unit Page

Academic Year 2017/2018

Learning outcomes

Describe the general characteristics of elements, molecules, and chemical reactions relevant to biology.

Course contents

The course contents are structured in:

1. The Structure of Atoms and The Periodic Table

  • Composition of the atom (electrons, protons and neutrons), atomic number, mass number, isotopes.
  • Atomic structure and atomic models: Democritus, Thompson and Crookes tubes, Rutherford and the gold foil experiment, Bohr and atomic spectra, double-slit experiment, De Broglie, Heisenberg, Schrodinger and the hydrogen atom model.
  • Electronic configuration of polyelectronic atoms and ions: electronic spin, Pauli principle, Hund rules, effective nuclear charge.
  • The periodic table and trends (atomic size, ionization energy, electronic affinity), electronegativity, ion size.

2. Chemical Bonding, Molecular Structures and Intermolecular Interactions

  • Chemical bonding: atoms vs. molecules, ionic and covalent compounds, valence and core electrons.
  • Ionic compounds: ionic crystal lattice, lattice energy and physical properties.
  • Covalent compounds and Lewis structures: single, double and triple bonds, dative bond, formal charges, resonance structures, the octet rule and its exceptions; properties of chemical bonds (bond order, bond length and bond energy).
  • Molecular geometry according to the VSEPR theory; Valence bond theory and hybrid atomic orbitals.
  • Non-polar and polar covalent bonds, polarity and molecular dipole moment.
  • Intermolecular forces: ion-ion interaction, ion-permanen dipole, permanent-permanent dipoles, permanent-induced dipoles, fluctuating dipole-induced dipole (London dispersion force) interactions and hydrogen bond.

3. Nomenclature of Inorganic Compounds, Chemical Reactions and Stoichiometry

  • Valence, oxidation number, real charges vs. formal charges.
  • Nomenclature of inorganic compounds (traditional, IUPAC and Stock): metals vs. non-metal elements, oxides, acids, hydroxides, hydrides and salts.
  • The Mole concept: definitions and examples.
  • Molecular mass vs. molar mass, percent composition of a chemical compound, empirical vs. molecular formulas.
  • Chemical equation and its balancing : stoichiometry, limiting reactant and excess reactant, reaction yield.
  • Stoichiometry exercises: examples and practice.

4. States of Matter

  • Gases and their properties: Boyle's law, Charles's law, Avogadro's hypothesis, the ideal gas law and the kinetic-molecular theory of gases, gas mixtures and partial pressures, Dalton's law.
  • Liquids and intermolecular forces: vapour pressure, phase diagrams and phase changes, heating and cooling curves of pure substances.
  • The solid state: molecular solids, covalent/network solids, ionic compounds and metals; crystalline and amorphous solids.

5. Solutions and Colligative Properties

  • Solutions: the solution process, aqueous solutions, solubility and equilibrium, temperature effects and pressure effects on solubility in aqueous solutions, Henry's law.
  • Units of Concentration: molarity, mole fraction, molality, mass/volume and mass/mass percent, ppm.
  • Colligative properties of solutions (with non-volatile solutes): Raoult's law, boiling point elevation, freezing point depression, osmotic pressure and osmolarity, van 't Hoff factor.

6. Thermodynamics and Thermochemistry

  • Exothermic and endothermic processes; energy, heat and work; the system, the environment and the universe; open, closed and isolated systems; specific heat capacity and latent heat in phase changes.
  • First principle of thermodynamics: internal energy, enthalpy, Hess's law and standard enthalpy of formation, standard enthalpy change of a reaction.
  • Thermodynamics and spontaneous processes: entropy (thermodynamic and statistical definitions), second law of thermodynamics, standard entropy and standard entropy change of a reaction.
  • Third law of thermodynamics: free energy, free energy and spontaneous processes, temperature effect.

7. Chemical Kinetics

  • Chemical kinetics: reaction rate and kinetic equations; concentration effect (zero order and first order kinetics), temperature effect (collision theory), contact area and catalysts.
  • Effective activation energy and impact, catalyst role and homogeneous and heterogeneous catalysis (hints).
  • Reaction mechanisms and molecularity.

8. Chemical Equilibrium

  • Chemical equilibrium: kinetic considerations and dynamic aspects.
  • Equilibrium constant: homogeneous and heterogeneous equilibria, equilibrium constant expression, equilibrium composition calculation, reaction quotient.
  • Effect of a perturbation on the chemical equilibrium and Le Chatelier's principle: effect of change in concentration, temperature, pressure or volume, adding a catalyst.
  • Exercises on chemical equilibrium: gas-phase reactions, reaction in solution or with solid reactants.

9. Acid-Base Equilibria

  • Acid-base equilibria in aqueous solutions: definition of acid and base according to Arrhenius, strong/weak acids and bases, mono and polyprotic acids, hydronium ion.
  • Definition of acid and base according to Bronsted and Lowry, conjugate acid-base pairs, amphiprotic substances.
  • Autoionization of water and Kw; definition of pH and pOH.
  • Exercises on acid-base equilibrium: acid and base dissociation constants (Ka and Kb); pH determination of solutions with strong acids or bases; pH determination of solutions with weak acids or bases; pH determination of solutions with salts.
  • Reactions between acids and bases: acid-base titrations and buffer solutions and the common ion effect.

10. Redox Reactions and Fundamentals of Electrochemistry

  • Oxidation-reduction reactions: electron exchange; balancing molecular redox reactions; balancing ionic redox reactions in acid and basic solutions.
  • Redox reactions and the Daniel cell: anode, cathode and salt bridge; standard notation for electrochemical cells; reactive and inert electrodes; electromotive force and standard reduction potentials.
  • Electrolysis: basic concepts and comparison between galvanic and electrolytic cells.

11. Fundamentals of Organic Chemistry

  • Carbon chemistry: carbon oxidation states, bonds and hybridization.
  • Hydrocarbons: alkanes and cycloalkanes, alkenes, alkenes and aromatic compounds. Sources, nomenclature and isomerism (structural, cis-trans, stereoisomers and chiral compounds). Physical properties and reactivity of hydrocarbons: combustion, halogenation and alkyl halides, double and triple bond additions, substitutions on the aromatic ring, reductions and oxidations.
  • Alcohols, glycols and ethers: nomenclature, physical properties and reactivity (substitutions, eliminations, condensation and oxidation reactions).
  • Aldehydes and ketones and the carbonyl group. Nomenclature, physical properties and redox reactivity.
  • Carboxylic acids and derivatives (esters, anhydrides and amides) .Nomenclature, physical properties and reactivity.
  • Amines: nomenclature, physical properties and reactivity with carboxylic acids.

Readings/Bibliography

In order to have a clear picture of the overall course contents, it is essential to support the PowerPoint presentations prepared by the teacher (downloadable form AMS Campus) with personal notes taken during the lectures.

It is also equally important to integrate the above-mentioned materials with a "General Chemistry" textbook. Although any university-level book is indicated for having a proper preparation, some particularly good textbooks are:

  • John C. Kotz, Paul M. Treichel, John R. Townsend and David A. Treichel,
    "Chemistry & Chemical Reactivity",
    Cengage Learning (9th ed.), 2015.
    ISBN: 978-1-133-94964-0
  • Katherine J. Denniston, Joseph J. Topping, Danaè R. Quirk Dorr and Robert L. Caret
    "General, Organic, and Biochemistry",
    McGraw-Hill Education (9th ed.), 2017.
    ISBN: 978-1-259-25339-3

The two textbooks listed above include many solved exercises and study questions with answers. Such books are also available in ebook format.

Teaching methods

Traditional lecture-based teaching, using blackboard and chalk, with the support of PowerPoint presentations.

Part of the time during the lectures will be devoted to carry out practical exercises, which will be analogous to those requested during the exam.

Assessment methods

The assessment of the course contents is based on a final examination only; no midterm exams are scheduled. The final examination consists of a single written exam, having a maximum duration of three hours.

During the examination, only the use of ordinary stationery (i.e., pencils, pens, rubber, rulers, etc.) and of a calculator is allowed. The periodic table (which sample copy is available on AMS Campus) and any other supplementary material will be directly attached to the exam papers. During the examination, students are not allowed to leave the classroom, unless they accept to definitively deliver all their exam papers and do not enter the classroom anymore. During the exam, the presence of electronic devices, other than a simple watch and the aforementioned calculator, is strictly forbidden.

In order to take the exam, students must sign up through AlmaEsami within the indicated deadlines. All the activities associated to the “BIOCHEMISTRY LABORATORY” must be completed and approved, before considering to take the exam. During the examination, students must have their ID document and their university badge.

The overall exam score is 34.00 points, equally divided in 17.00 points for the “CHEMISTRY” part and 17.00 points for the “BIOCHEMISTRY” one. Each single written exam will include both the “CHEMISTRY” and the “BIOCHEMISTRY” part and cannot be split.

The exam is passed by achieving at least 9.00 points (out of 17.00) in both parts. The final mark is calculated by summing the points obtained in both the above mentioned parts and rounding it to the nearest integer.

Examples:

  • Chemistry: 10.50 pts + Biochemistry: 8.50 pts
    ---> FAILED
  • Chemistry: 11.00 pts + Biochemistry: 13.75 pts
    ---> 24.75 pts ---> PASSED WITH 25/30.

If a student obtains an overall score of 31.00 points or more, 30L is assigned.

The exam results will be published online on AlmaEsami. If the exam is passed, the associated mark will be officially registered only after the student's confirmation to be sent to the integrated course coordinator (i.e., Prof. Michele Di Foggia, michele.difoggia2@unibo.it), using the student institutional email (e.g., name.surname@studio.unibo.it).

The registration of a successful examination is permitted no later than 6 months from the date of the exam; after this deadline, the result achieved and not confirmed will expire.

If a student decides to take the exam again, regardless of the associated outcome, the result obtained in the previous examination will automatically be cancelled. The exam can be repeated at any exam session.

The "CHEMISTRY" exam part consists of:

  • 8 general multiple choice questions [0.50 points each]
  • 2 inorganic and 2 organic nomenclature exercises [0.25 points each]
  • 1 redox reaction including balancing and spontaneity [1 point]
  • 1 exercise involving thermodynamic and/or kinetic and/or chemical equilibrium [4 points]
  • 1 exercise involving acid-base equilibria, including titrations and buffer solutions [3 points]
  • 1 exercise involving organic chemistry [4 points]

A sample exam, including both the "CHEMISTRY" and "BIOCHEMISTRY" parts, can be found on AMS Campus.

Teaching tools

All the lectures will be supported by PowerPoint presentations, downloadable from AMS Campus before the beginning of each lecture.

A sample of the written exam (with solutions) will also be uploaded on AMS Campus, together with other files containing similar questions and/or exercises.

WARNING: PowerPoint slides are an excellent teaching tool during the lectures and a good support for individual study, but they can not be considered exhaustive for the exam preparation!

Office hours

See the website of Filippo Monti