Humboldt State University ® Department of Chemistry

Richard A. Paselk

SA 560B; 826-5719

Hour Exam 3 Study Guide

Spring 2014 version


Homework since Exam II (see suggested problems on Schedule). Review Quizzes (see Keys on my Moodle site), particularly those on materials since exam II. Review nomenclature so that you can read questions and understand them. Review concepts from exams I & II that we have used in looking at additional systems etc.

Be able to convert numbers to scientific notation and use numbers expressed in scientific notation; do all calculations with proper significant figures; make all conversions within the metric system (SI).

Thermochemistry (Chapter 6)

Understand/define exothermic and endothermic. What are exothermic and endothermic reactions. What is heat? work? energy? enthalpy? DeltaE = q + w. DeltaH= DeltaE - w. When are enthalpy and energy equal? Be able to solve heat capacity problems, calorimeter problems. Understand thermochemical equations (what happens to DeltaH if the reaction is rewritten in reverse? What is a formation reaction? What are: standard state, standard heat of formation, standard enthalpy of formation? Deltaf . Know Hess's Law. Be able to solve heat/enthalpy of formation problems. Module 11(Thermochemistry)

Electrons, Atoms and Periodicity (Chapter 7)

What are De Broglie waves? = h/mv. When is this equation used? What is the relationship between m & f? m & ? Be able to rearrange and to solve problems with this equation. Module 12 (Energy & Light) What does the Schrödinger equation describe? What are nodes? Be able to apply the concept of nodes to standing waves (as on a string) and to electron distributions. What does refer to for electrons and atoms? Be able to draw and to interpret cross-sectional diagrams of electron distributions in atoms for s & p orbitals. Be able to draw and interpret probability density ( vs. r) plots for s & p orbitals for any value of n. How is the number of nodes related to n? How is a node represented on a vs r plot? Remember there is always one node at infinity. Atomic Orbitals supplement

What is an orbital? How many electrons can one orbital accommodate? What is a shell? Know number and letter designations (n= 1,2,3.. or K,L,M,...).What is a subshell? How are shells and subshells related to energy? Be familiar with energy and filling diagrams. What is spin? How does it effect orbital filling? What is paramagnetism? when does it occur?

Be able to give electronic configurations using all three of the conventions we have discussed. Know which orbitals are being filled for different regions of the periodic table. Know what the four quantum numbers are and be able to assign quantum numbers to any electron in an atom, or to give an orbital designation for a set of quantum numbers. What information does each of the quantum numbers convey? (energy, orbital shape etc.) What is the Pauli Exclusion Principle and why is it important? How are nodes related to quantum numbers? (n = # nodes, l = angular nodes,...) Module 13(Quantum Numbers, Electronic Structure of Atoms & Periodicity)

Discuss/Understand the various periodic properties we have discussed (atomic radii, ionization energies, electron affinities, electronegativities, stoichiometric ratios, metallic vs. non-metallic properties) and know the trends.

How do these periodicity's relate to electronic configuration? Be able to explain trends, etc. in the plots of these properties. Be able to describe the properties of elements in grps I-A, II-A, VII-A, & VIII-A. Be able to name each of these groups of elements. Where are the transition metals? the lanthanide's? the actinides?

Chemical Bonding (Chapter 8)

What is electronegativity? How is it used? Be able to use electronegativity (both quantitatively and qualitatively-hi,lo rules) to predict the ionic/covalent nature of a bond.

What is an ionic bond? How is it formed? What forces maintain it? Anion. Cation. Ion pair. Be able to apply the octet rule to guess the ionic forms of the various "A" group elements. Why are outer electron shell octets so stable? What is a covalent bond? How do ionic and covalent bonds differ? How is it possible for an ionic bond to be "strong" and unstable while a covalent bond may be stable while not being as "strong"? How are electrons arranged in each?

What is a Lewis Structure? What is it intended to show? For which elements are Lewis structures most useful? (A groups) Be able to draw Lewis structures for all the A-group elements in their atomic and predicted ionic forms. Kernel vs. inert gas core.

Be able to draw a correct Lewis structure for any covalent molecule made up of atoms from the "A": groups. (Remember-you use Electronegativity to determine covalent or ionic first!) When do you need to use multiple bonds? Resonance? Are multiple bonds real? What does resonance represent in the real molecule? Be able to make proper resonance based sets of Lewis structures. For which A-grp elements does the octet rule not hold? (H to B) When is valence expansion required? Be able to draw correct expanded valence shell Lewis structures. Which orbitals are involved in expanded valence shells? What is the outermost shell of an atom? Does the outer-most shell of an atom need to have any electrons in it? Polar vs. non-polar bonding. Be able to assign Formal charges to each atom in a molecule. What are Formal charges useful for? Be able to predict chemical stoichiometry using Lewis structures, and balancing.

Bond dissociation energy. Bond length. Understand/be able to explain the energetics of ionic reactions. Be able to do a Born-Haber calculation given appropriate data. (In other words, you need to be able to recognize the required steps.) Understand/be able to calculate reaction enthalpies (DeltaH) given bond energies. How does bond length vary with multiple bonding?

VSEPR Theory

What is VSEPR theory? Be able to use the VSEPR method to determine the geometry of a molecule, including demonstrating the process involved-what assumptions are made with this model. Know the various orientations of electron pairs around a central atom and the various molecular geometry's they predict (handout in Discussion manual, p LN-4). What relations are there between electronic configuration and molecular shape? Does a given electron configuration always give the same molecular shape? Explain. Polarity: Be able to determine whether a given bond is polar or not. Be able to determine whether a given molecule is polar or not. Defend your decision (think charge separation [differences in electronegativity] and geometry). See also VSEPR Theory and Molecular Geometry Module

Covalent Bonding – A Quantum View (Chapter 9)

When atomic orbital sets are filled, or half-filled they become completely symmetrical. Orbitals are orbitals: 1) Only two electrons can be accommodated in any orbital. 2) No two electrons can have the same "address" (the same set of quantum numbers). 3) For a molecules the "address" becomes the molecule over which the electrons are shared rather than the atom. 4) We have conservation of orbitals - a molecule will have the same number of orbitals as the atoms which make up the molecule. 5) For our purposes we can also assume a conservation of orbital energy.

Hybrid Atomic Orbitals

What are hybrid atomic orbitals? What do we mean when we say this is a "localized electron" model? What does it tell us about the electron distribution in molecules? Why are hybrid orbitals popular? (relatively easy to calculate/visualize). What are their limitations? (NOT looking at molecules, so can only be a limited approximation.) How are bonds visualized with hybrid orbitals? Know the five common hybrid orbital sets (sp, sp2, sp2, dsp3, d2sp3) and their geometries (linear, trigonal planar, tetrahedral, trigonal bipyramidal, and octahedral). Be able to describe the molecular geometries we discussed with VSEPR Theory in terms of hybrid orbitals. Does the hybrid orbital set correspond to the electronic or to the molecular geometry?

What are the two basic bond types in molecular orbitals? (Sigma (s) bonds - cylindrically symmetrical around the axis connecting the bonded atoms, and Pi (p) bonds - made up of two lobes with planar symmetry round a plane though the nuclei of the two bonded atoms.) Note that single bonds are always sigma bonds, and in a multiply bonded system the "central" bond is a sigma bond. The "second" and "third" bond of multiply bonded atoms are pi bonds. For systems with two pi bonds the bond panes are perpendicular to each other and the two bonds form a symmetrical cylinder around the sigma bond. What is resonance and why is it needed in hybrid orbital theory? How is this a "failure" of the theory? Why does hybrid orbital theory not give information about electronic energy levels in molecules?

Be able to describe the hybridization of atoms in molecules given their structures, and be able to specify the bond types (sigma or pi) between atoms in a structure. (Sample question: Give the hybridization for each atom and the bond types for: CH3CCCH2COH.)

Molecular Orbital Model of Bonding

How does MO theory differ from hybrid orbital theory? What advantages does it have over hybrid orbital theory? Disadvantages? Be able to draw energy diagrams. Know and be able to explain the molecular energy diagrams we discussed in lecture. How is bonding described in MO theory? Are bond shapes the same (sigma & pi)? Be able to determine when molecules are diamagnetic or paramagnetic and explain why. Be able to give the bond order of a diatomic molecule. Define bond order. (Sample questions: 1) Draw a Molecular Orbital Energy-level diagram for NO. Include all of the valence electrons in your diagram, and determine the bond order and whether the molecule should be diamagnetic or paramagnetic based on your diagram. 2) Draw a molecular orbital energy-level diagram for NO showing the original atomic energy levels and the MO energy-levels. )

You will be provided with a Periodic Table and Useful Constants/Equations etc.


Syllabus / Schedule

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© R A Paselk

Last modified 6 March 2013