GHW#2: Louisiana Tech University, Inorganic Chemistry-CHEM 281. POGIL Exercises on Chapter 1 & 2. Electronic Configuration and Lewis structures
Why?
In explaining chemical changes atomic structure is very important. It was discovered through many experiments starting with alchemy. The modern atomic theory requires treating electron as waves leading to Schrödinger wave equation. It is essential to understand the significance of this equation, wave function and its interpretation. We can use Bohr model to describe atomic spectrum of hydrogen but requires wave mechanical models to interpret the atomic spectroscopy of multi-electron atoms. The wave shapes of the atomic orbitals are useful describing bonding. The electron configurations of poly-electronic atom and their ions are useful in explaining the bonding and magnetic properties of atoms.
How is bonding in chemical compounds is predicted using the periodic table? How did Lewis come up with a theory to explain bonding? What rules you follow to draw a Lewis elemental symbols and molecular structures? How bond order and partial bond order is calculated? How you calculate formal charge and use to predict stability of Lewis structures.
Learning Objectives
Understand the 1) How atomic structure was discovered, 2) The Schrödinger wave equation and its significance, 3) Wave function and its interpretation,4) Atomic absorption spectroscopy, 5) Shapes of the atomic orbitals, 6) The polyelectron atom and their electron configurations, 7) Ion electron configurations, 8) Magnetic properties of atoms.
Success Criteria
Understand
· The experiments 1) cathode-ray tubes 2) Oil drop experiment, 3) Rutherford's a-particle Experiments,4) Moseley's X-ray experiment, 5) Chadwick's bombardment of 9Be with a-particles and their importance in finding modern atomic structure.
· The wave properties and Electromagnetic radiation (EMR):
· The importance of radiation profiles emitted by hot solid bodies and Plank equation.
· Bohr model in explaining hydrogen emission spectrum.
· Importance of photoelectric effect and particle description of electrons
· Wave-Particle Duality of Matter and Energy and wave technical models
· Heisenberg uncertainty principle.
· Atomic spectroscopy: a) Emission spectroscopy b) Absorption spectroscopy
· Quantum numbers
· Sub-levels and periodic table: Long form and orbital blocks (s, p, d and f)
· Pauli Exclusion Principle, Building-Up Principle, and Hund's rule
· Electronic Configuration: core and valance orbital box formats
· The Schrödinger Wave Equation and Its Significance
· Atomic Orbitals and their relation to wave function, their shapes, number nodes
· Unpaired electrons and paramagnetism.
· Magnetic properties of materials:
Resources
INORGANIC CHEMISTRY by Inorganic Chemistry By Peter Atkins, Tina Overton, Jon Rourke, Mark Weller, Fraser Armstrong, 5th Edition 2010.
Prerequisites: Freshman and organic chemistry
What you already know
1) Atomic spectroscopy: c) Emission spectroscopy d) Absorption spectroscopy
2) Heisenberg uncertainty principle. The position and momentum of a particle cannot be simultaneously measured with arbitrarily high precision.
3) Wave - mechanical model of atom
4) Quantum numbers
In wave mechanics there are rules to obtain Quantum numbers:
n value could be 1, 2, 3, 4, 5, . . . . Etc.
l values depend on n value: can have 0 . . . (n - 1) values
ml values depends on l value: can have -l . , 0 . . . +l values of ml
ms should always be -1/2 or +1/2
5) Sub-levels: s, p, d, and f sub-levels (different l values)
6) The Schrödinger Wave Equation and Its Significance
7) Max Born Interpretation: y2 = atomic orbital ( region finding an electron is highest)
8) Wave Function:
Radial Wave Function (R n, l (r )) and
Angular Wave Function Y l, ml (q, f);
9) Nodes in the yn, l, ml, ms :
Total nodes, n-1
Radial nodes = n -1- l,
and Angular nodes = l
10) Shapes of the Atomic Orbitals: See text book and class notes for chapter 1.
11) Periodic table: Long form and orbital blocks (s, p, d and f)
12) Building-Up Principle
13) Exception to Building Up Principle: Transition Metal Elements
Electronic Configuration of d-block and f-block elements d5 or d10 and f7 or f14 are stable
14) Electronic Configuration: core and valance orbital box formats
15) Pauli Exclusion Principle: No two electrons in an atom can have same four Q. N. s.
16) Atomic Orbitals: Wave shapes shown by electrons in a sublevel. s, p, d, f orbitals.
17) Hund's rule: filling p, d, f orbitals: Electrons will spread over all available (3ps, 5ds, 7fs) and try to have parallel spins.
18) Unpaired electrons: paramagnetism and diamagnetism
New concepts
1) Ion Electron Configurations: Electron configurations can also be written for ions by starting with the ground-state configuration for the atom. For cations, remove a number the outermost electrons equal to the charge. For anions, add a number of outermost electrons equal to the charge. Transition metals remove s-electrons before d-electrons
2) Magnetic Properties of Atoms:
Diamanetism: paired electrons.
Paramagnetism: unpaired electrons.
Ferromagnetism: cooperativity of unpaired electron among different atoms.
Supermagnetic Materials: Magnets from superconductors.
Covalent Bonding and the Periodic Table
When you look at the inert gas elements of group 8 they only exist as monoatomic gases, and do not in general react with other elements. In contrast, gaseous elements in other groups exist as diatomic molecules (H2, N2, O2, F2 & Cl2), and all but nitrogen are quite reactive to produce compounds where they achieve close valence shell electron configurations. Non metallic elements in the periodic table share electrons to form compounds or molecules with shared electron pairs between two atoms: covalent bond.
a) Covalent bonds: Two atoms in molecule or ion in which each atom shares one valence electron with the other atom to from a single bond that keeps the two atoms together.
E.g. In fluorine molecule, F2 each atom F atom has 7 valence electrons and shares one valence electrons to form a covalent bonding pair. Bonding pair is represented by a line.
b) Coordinate Covalent bonds (Lewis/acid-Base adduct): Two atoms in molecule or ion in which one atom shares two valence electrons with the other atom to from a bond that keeps the two atoms together.
E.g. In ammonia ion fluorine molecule, NH4+, N atom has a lone pair of electrons and shares with the H+ ion to from a single bond that keeps the two atoms together. Coordinate covalent bonding pair is represented by a arrow pointing to atom receiving the electron.
Lewis Theory
Lewis Theory of Covalent Bonding: The idea that the stability of noble gas electron configurations and the realization of the connection of reactivity of elements to achieve octet of valance electrons through formation of electron-pair bonds with other atoms of main group elements except duet for hydrogen.
Lewis Electron-Dot (Line) Formulas: A simple way of writing out a formula that shows the disposition of the shared valance electron pairs between the different atoms in a molecule or ion.
Lewis structure
Covalent Bonds in Lewis Structure of Molecules and Polyatomic Ions: Lewis electron dot structures are representations of the distribution of electrons in molecules and polyatomic ions. A Lewis structure can be drawn for a molecule or ion by following three steps:
1. Calculate the number of valence electron pairs (including charges, if any)
2. Pick the central atom: Lowest EN (electronegative) atom, largest atom, and/or atom forming most bonds is usually central atom.
3. Connect central atom to all terminal atoms with a single bond.
4. Fill octet to terminal atoms
5. Fill octet to central atom: make double triple bonds if necessary.
6. Check that total number of valance electron pairs in the structure match with the lines.
Bond pairs: An electron pair shared by two atoms in a bond.
Lone pair: An electron pair found solely on a single atom.
Multiple Covalent Bonds in Lewis Structures: For every pair of electrons shared between two atoms, a single covalent bond is formed. Some atoms can share multiple pairs of electrons, forming multiple covalent bonds.
E.g.. For example, oxygen, O2 (which has six valence electrons) needs two electrons to complete its valence shell. In O2, two pairs of electrons are shared, forming two covalent bonds called a double bond.
Formal Charge
Formal charge of atoms in Lewis structure is a number assigned to it according to the number of valence electrons of the atom and the number of electrons around it. The formal charge of an atom is equal to the number of valence electrons, Nv.e, subtract half of the bonding electrons, ½ Nb.e,. and . subtract the number of unshared electrons, Nus.e. (Formal charge = Nv.e. - ½ Nb.e. - Nus.e.)
Bond Length and Bond Energy
Bond Energy order: single=1 < double=2 < triple=3
Bond length: single (1pair) > double (2 pairs) > triple (3 pairs)
Bond lengths from periodic trends in atomic radii
Bond length is proportional to the sun of atomic radii forming the bond.
Atomic radii trend: Li> Be> B> C> N> O> F
Bond Length Order: Li-H> Be-H> B-H> C-H> N-H> O-H> F-H
Resonance Structures
The Lewis structure some molecules and ions have a double bond, that could be placed between alternating positions. For example, carbonate ion, CO32-, has three oxygen atoms that can from a double bond creating three possibilities, and they have the same types of bonds and electron positions, but they are not identical. There individual Lewis structures are resonance structures for carbonate ion.
Resonance is a common feature of many molecules and ions of interest in organic chemistry.
Partial Bond Order
Bond order is the term used to distinguish between the number of bonds that exist between two atoms. Under some conditions when there are resonance structures, atoms may find themselves with an average bond that appears to be half-way between a single and a double. In that case of CO32-, the bond order would be 4/3= 1.33. As a result, the average appearance of the bond over a period of time will be the bond order.
GHW#2: CHEM 281 Name:______
Key Questions (relatively simple to answer using the Focus Information)
1) Describe the following concepts used in getting electronic configuration of multielectron atoms:
a) Building-Up Principle:
b) Electronic Configuration: core and valance orbital box formats
c) Pauli Exclusion Principle:
d) Hund's rule:
e) Exception to Building Up Principle: Transition Metal Elements
2) Explain concisely why nitrogen has three electrons in different p orbitals with parallel spins rather than the other possible arrangements.
3) Write noble gas core ground state electron configuration for atoms of
a) Calcium:
b) Iron:
c) Silver:
4) Write noble gas core ground state electron configuration for ions of
a) potassium(+):
b) scandium (3+):
c) copper (2+):
5) Predict the common charge of the Cu ions. Explain your reasoning in terms of electron configurations.
6) Determine the number of unpaired electrons in atoms of
a) nitrogen;
b) aluminum:
c) Iron (+3):
7) Write the electron configuration expected for element 113 and the configurations for the two cations that it is most likely to from.
8) In the text set of orbitals after the f orbitals are g orbitals. How many g orbitals would there be? What would be the lowest principle quantum number n that would process g orbitals? Deduce the atomic number of the first element at which g orbitals would begin to be filled on the basis of the patterns of the d and f orbitals.
9) Describe the following atomic properties and their tern in the periodic table:
a) Atomic size:
b) Effective nuclear charge:
c) Ionization potential:
d) Electron Affinity:
10) Describe the types of bonding seen among elements in the periodic table.
11) How the periodic table is used in predicting bonding in elements and compounds?
12) Predict the bonding in following elements and compounds:
- O2:
- SCl2:
- SiO2:
- CuZn:
- Na:
- SiCl4:
- KCl:
13) How you distinguish between covalent molecules and network covalent molecules?
14) Draw Lewis dot-symbols and predict number of covalent bonds they will make
- B
b. C
15) Construct electron-dot diagrams for:
a) Ammonium ion:
b) Carbon tetrachloride:
c) Silicon hexafluoride(2-) ion:
d) Pentafluorosulfate(IV) ion, SF5-:
Resonance structures, partial bond order and formal charge
16) For the nitrate ion:
- Construct an electron-dot diagram:
- Draw the possible resonance structures:
- Estimate the average nitrogen-oxygen bond order:
- Draw a partial bond representation of the ion:
17) The boron trifluoride molecule, BF3, is depicted as having three single bonds and an electron-deficient central boron atom. Use the concept of formal charge to suggest why a structure involving a double bond to one of the fluorine, which would provide an octet to the boron, is not favored.