Expert answer:chemistry

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EXAM 2 REVIEW PACKET – CHEM 151 – 006/009 – Fall 2017 – Due before exam on D2L
Name: __________________________________________
UA NetID: ___________________________________@email.arizona.edu
Define or explain each of the following terms, and show how they relate to the other terms in
each group. Build a concept map using these terms to show their relationship to one another.
Protons
Neutrons
Electrons
Mass spectrometry
Bohr model of the atom
Quantized energy levels
PE and KE of an electron
Absorption and emission
Infrared spectroscopy
Photoelectron spectroscopy
Atomic radius and ionization energy
PE and KE of bond formation
Resonance structures
VSEPR and molecular geometry
What do the red and the green arrows represent?
What is the energy of transition between n=1 and n=3?
If photons of 341 nm were measured for one of the
transitions represented by a green arrow, which transition
would that correspond to (A or B)? Explain why.
What would be the wavelength of photons from the other transition? (between 100 – 900 nm)
An unknown compound has the following elemental
analysis: 54.53% C, 9.15% H, 36.32% O by mass.
Determine the molecular formula of the compound.
Which of the six proposed line structures (below) are
in agreement with the IR data? (hint: there are two)
O
O
OH
O
O
O
O
OH
O
O
Re-draw the selected line structures as full-atom Lewis structures below.
Draw the full-atom structures of two fragments from the MS: 15 m/z and 29 m/z
Are they consistent with both structures drawn above? (hint: they should be)
Explain how the MS peaks at 43 and 45 m/z rule out one of the two selected structures.
Write out the electron configuration for these atoms:
C
O
O2Mg
Mg2+
Draw the photoelectron spectra (PES) of neutral
carbon and oxygen atoms relative to each other.
Which atom (carbon or oxygen) would have the smaller atomic radius? Explain why.
Which atom (carbon or oxygen) would have the higher first ionization energy? Explain why.
Do the PES spectra that you have drawn agree with the explanations above?
(If not, re-draw the spectra to show the differences in IE and atomic radius)J
Based on the PES spectra drawn for carbon and oxygen atoms, explain why oxygen generally
makes two bonds while carbon generally makes four bonds.
Draw the Lewis structure of carbon monoxide (CO). Propose a counter-argument to the
previous explanation for how many bonds each atom should have.
Oxygen and sulfur both have 6
valence electrons, but sulfur has
more diverse bonding patterns.
While oxygen generally makes 2
bonds, sulfur can make 2, 4, or 6
bonds due to its size.
By forming additional bonds, the sulfur atom is in violation of the “octet rule”. Explain what this
means in terms of valence shell capacity and potential energy. (hint: electron repulsion)
Explain how the formation of covalent bonds through delocalization of valence electrons
lowers the total energy of the molecule. (hint: why is bond formation a stabilizing effect?)
Using the arguments above, explain why sulfur has “hypervalent” bonding (i.e. bonding that
exceeds the capacity of the valence shell), while oxygen strictly obeys the “octet rule”.
Benzene is often drawn with alternating C–C
(single) and C=C (double) bonds. This drawing
implies that the electrons making up each of
these double bonds are localized to the two
carbons between which they are drawn. If this
were the case, we would be able to measure
two different lengths for the carbon-carbon
bonds in benzene (154 pm and 134 pm).
However, experimental observations suggest a different structure for benzene where the
carbon-carbon bond lengths are the same within the molecule, implying that the electrons
originally drawn in the double bonds are delocalized among all six the carbons.
Explain how the kinetic energy of the electrons changes when the electrons that would be
localized in the C=C bonds are delocalized over all of the carbons in the benzene molecule.
The observed bond length in benzene (140 pm) is shorter than predicted (144 pm) by
averaging the bond lengths of three C–C bonds and three C=C bonds. Since atomic radius is
half of the bond length, the effective atomic radius of the carbon atoms in benzene is also
smaller than predicted. Explain how delocalization of electrons in benzene decreases the
atomic radius of carbon atoms. (hint: think of re-distribution of electron density)
Explain how the potential energy of the benzene molecule changes if the atomic radius of the
carbon atoms decreases.
…
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