Download CHEM 400 - El Camino College

CHEM 101
E X AM 2 I F O R M AT I O N
SPRING 2016
GENERAL INFORMATION
Exam Time, Place, Format, and Rules
The exam will take place on Wednesday, 05/25, in room SCI 314 during the regular class time. (We will go
over a couple of topics from Chapter 11 and will do the lab check out before the exam.) The exam will have
two parts: multiple choice questions and open-end questions. You will need to bring a Scantron form No.
882-E. Partial credit will be given for open-end questions only. Open-end questions will be of two
categories: problems and short essay questions. They will be similar to those from your homework
assignments, worksheets, and quizzes. You are expected to give a logical, well organized, and complete
solution for each problem. The answer only responses to problems will not receive much credit. You do not
need to show your work for conversions involving metric prefixes. Your responses to essay question should
be short, but complete; they should demonstrate your knowledge of specific information as it was presented
and discussed in lecture and/or the textbook. Your random personal thoughts on a subject will generally get
little credit. During the exam, you are not allowed to use any dictionaries and you are not allowed to use
any electronic devices except for your personal electronic calculator. You cannot share a calculator with a
classmate or use your cell phone in place of a calculator. Your cell phone cannot be in your hands at any
time during the exam.
General Tips on Preparation
1. Review your own lecture notes or, in case your are not a good note taker or missed a lecture, borrow
them from a classmate. On the exam you should expect questions on any material covered in lecture
and lab including instructor’s demonstrations. You will not be tested on any topics that are covered in
the current chapters of the textbook, but have not been discussed in lecture or lab.
2. Complete all the worksheets from lecture and lab (even if you have not submitted them for grading) and
review your graded quizzes. Problems on the exam will be similar to problems from the lecture and lab
worksheets, homework assignments, and quizzes. You should be able to redo all the problems on
MasteringChemistry for practice even if you already submitted the answers.
3. Review the exam outline below and identify areas that you need to work on. Review the related
textbook material. Solve a few problems from each chapter that are similar to the problems we solved in
class or to the problems you solved on MasteringChemistry. A problem is considered to be solved if
your answer matches closely the answer from the textbook or a worksheet. The answers to the textbook
in-chapter problems are in Appendix IV; the answers to the textbook end-of-chapter odd-numbered
problems are in Appendix III.
Textbook Chapters (look through the exam outline for specific sections)
Chapter 5 (summary on pages 235 – 236); Chapter 6 (summary on pages 284 – 286);
Chapter 7 (summary on pages 328 – 329); Chapter 8 (summary on pages 372 – 373);
Chapter 9 (summary on pages 416 – 417); Chapter 10 (summary on pages 472 – 473);
Chapter 11 (summary on pages 533 - 534); Chapter 18: section 18.2 only (balancing redox reactions).
Look through/study the sections to avoid surprises on the day of the exam! Solve as many recommended
problems from the textbook as possible!
Laboratory Work to Know
Exp. 6 “Analysis of Al/Zn alloy”; Exp. 7 “Enthalpy of Reaction”; Exp. 10 “Periodic Properties”; Exp. 9/16
“Conductometric Titration / Gravimetric Analysis”; and Exp. 18 “Redox Titrations”
EXAM OUTLINE
GASES
Textbook sections: 5.1 – 5.7 (concepts and calculations); 5.8 – 5.10 (no calculations except for Graham’s
Law); 11.5 (vapor pressure; omit the Clausius-Clapeyron).
How do properties of gases differ from properties of liquids and solids on macroscopic level (appearance,
density, compressibility) and microscopic level (distances between molecules and molecular motion)?
Be familiar with basic ideas of kinetic molecular theory. Be able to distinguish between ideal gas and real
gas. Know how to interpret temperature and pressure of gasses in terms molecular motion. Remember that T
is directly proportional to the average kinetic energy (K.E.avg) of gas molecules, not to the average velocity
(vavg) ! T ~ K.E.avg =
2
mv avg
2
, where m is the mass of one molecule of gas. At the same temperature larger
molecules on average move slower than smaller ones.
Pressure (force per unit area). Common units pressure that you should know how to use: atm, mmHg, and
torr. (The SI unit of pressure is pascal that is abbreviated as Pa, but you are not responsible for remembering
its relation to other common units.)
Gas laws: mathematical relationships between macroscopic parameters (amount of substance, mol;
temperature, K; pressure; volume) describing the ideal gas. You are encouraged, but do not have to remember
each law by name as long as you can write the appropriate relationship when solving gas law problems.
Remember how the ideal gas law formula PV=nRT allows you to deduce equations such as P1×V1 = P2×V2
(n=const, T=const).
Remember the value of R with 4 significant digits and the appropriate units: R = 0.08206
atm ⋅ L
. As long
mol ⋅ K
as you remember that 1 atm = 760 torr or 760 mmHg (exactly), you can multiply the value of R given above
by 760
torr
torr ⋅ L
to obtain R = 62.36
.
atm
mol ⋅ K
Be able to distinguish between densities and molar volumes of liquids, solids and gases. Equal numbers of
molecules of gases occupy the same volume under the same conditions of temperature and pressure (the
Avogadro’s law). The Avogadro’s law is a gas law and it is not applicable to liquids and solids. The molar
volume of 22.4 L/mol at STP (0°C and 1 atm) is a good approximation for all gases, but cannot be applied to
liquids and solids. The molar volume of water at 25ºC and 1 atm is about 0.018 L/mol (the volume of the
ideal gas under the same conditions is 24.5 L/mol).
Know how the densities of gases depend on temperature, pressure and identity of a particular gas (its molar
mass).
Know how to use the ideal gas law (PV=nRT) to convert between the amount of gas in moles and the volume
(an essential skill in solving stoichiometry problems where a reactant or a product is a gas). In solving
stoichiometry problems always rely on a properly written and properly balanced chemical equation. (Do not
assume that all substances react in one-to-one mole ratio. Do not simply equate the moles of gas with the
moles of a nongaseous reactant or product: use the mole ratio from the balanced chemical equation!!!)
If three out of the four parameters describing the ideal gas (P, V, n, T) are given, know how to use the ideal
gas law formula (PV=nRT) to find the forth.
Be familiar with the concept of partial pressure of gases in a mixture and the law of summation of partial
pressures of gases (Dalton’s) law. Know how to use the Dalton’s law to calculate the partial pressure of the
gas collected over water if the vapor pressure of water is given in the problem or a reference table “Vapor
Pressure of Water vs. Temperature” is available).
Liquid-vapor equilibrium: vapor pressure depends on temperature, but not on the volume of the system. What
is common and what is different between boiling and vaporization? Vapor pressure vs. temperature tables (a
blue laminated handout for water) and graphs. Boiling point as a function of pressure. Normal boiling point.
THERMOCHEMISTRY
Textbook sections: 6.2 – 6.9 (no PV work calculations such as in Example 6.4); 9.10 (bond energies); 11.7
(∆Hfus and ∆Hvap).
Have a basic understanding of the subject and significance of thermochemistry and chemical
thermodynamics.
Have a basic understanding of the concepts of energy, internal energy, heat and work. J (a derived SI unit)
and cal as units of energy, work, and heat. Know the new definition of calorie (1 cal is exactly 4.184 J) and
the old one (the amount of energy required to increase the temperature of one gram of liquid water by one
degree Celsius).
Distinguish between heat at constant volume (qv = ΔE) and heat at constant pressure (qp = ΔH). Concept of
enthalpy. Why do chemists prefer to use enthalpy rather that internal energy?
Know how calorimetry can be used to determine specific heats of substances and heats of reactions. What are
the advantages and limitations of a coffee-cup calorimeter? What are other types calorimeters that you know?
Be able to compare and contrast the terms in each of the following pairs: system and surroundings,
temperature and heat, heat and specific heat, heat capacity and molar heat capacity, specific heat and molar
heat capacity, heat and work, internal energy and enthalpy, chemical equation and thermochemical equation,
endothermic and exothermic.
Be able to understand and use the formula: q = c×m×ΔT. You are recommended to memorize the specific
heat of H2O(l): 1.00 cal/(g∙°C) and 4.184 J/(g∙°C).
Know what thermochemical equations and heats of reactions are. The ΔH that accompanies a
thermochemical equations has a unit kJ (per reaction as described by the equation with the coefficients
representing moles of each reactant and product), not kJ per mole of a specific substance.
Be able to give a written statement of Hess’s law of heat summation and know how to apply in solving
problems. What is the significance of Hess’s law in thermochemistry?
Be able to give a written definition of the term enthalpy of formation (ΔH°f). Be able to write a
thermochemical equation for a reaction of formation for a particular substance provided you know the heat of
formation. Be able to give examples of: substances for which ΔH°f is set up to be zero; substances for which
ΔH°f can be obtained directly from a calorimetry experiment; and substances for which ΔH°f can only be
calculated through Hess’s law.
Know how to use the table of enthalpies of formation of selected substances (a blue laminated handout) tto
calculate heats of reactions.
Know how to calculate the enthalpy of formation of an organic substance if the value of its heat of
combustion is given.
Know what the term bond energy means. Know how to use the values of bond energies to calculate the ∆H for
a reaction that occur in gaseous phase.
WAVES AND PHOTONS
Textbook section: 7.2
Be able to clearly explain what sound and light have in common and what the principle difference between
them is.
Know the characteristics of waves, velocity, frequency, and wavelength, and the equation that ties them all.
Know the value of the speed of light with 3 significant digits.
Know that diffraction and interference are phenomena characteristic of any type of waves. What does the
double slit experiment prove?
Be familiar with electromagnetic (EM) spectrum. Know what is common between visible light, X-rays, and
radio waves (all are types of EM radiation) and what distinguishes them from each other (wavelength,
frequency, energy of individual photons).
What is quantum? What is photon? Know the Planck relation (Ephoton = hν) and the value (with 4 significant
digits) and the unit of the Planck constant (h = 6.626×10‒34 J·s).
Be able to find the energy of photon if the wavelength of the radiation is given and the other way around.
Remember that chemists like to express everything per one mole. You should know how to get the value of
energy of one mole photon when the energy of a single photon is known and the other way around.
What is photoelectric effect? What are the threshold frequency and wavelength? What is the best explanation
of photoelectric effect: in terms of waves or particles (photons)? Why?
ATOMIC EMISSION SPECTRA AND THE BOHR MODEL OF THE ATOM
Textbook sections: 7.3 (all); 7.5 (pages 318 – 321).
What is meant by continuous and line spectrum? To which category do the spectra of the sunlight and an
incandescent light bulb belong to? What is the appearance of spectra coming from excited (energized) atoms?
What are the applications of the spectrum analysis atomic emission spectra (analysis of atomic emission
spectra)?
What are the main ideas (postulates) of the Bohr model of the atom? How does it explain the emission
spectrum of hydrogen? What are the limitations of the Bohr model?
The Rydberg equation and it relation to the Bohr model: which of the two was proposed first and which
explains which?
QUANTUM MECHANICAL MODEL OF THE ATOM AND THE ELECTRON CONFIGURATIONS
OF ATOMS AND MONATOMIC IONS
Textbook sections: 7.4 – 7.6 (skim reading; study pictures); 8.3 – 8.5; 8.7.
Be able to outline the principle differences between the Bohr model and the approach pioneered by
Schrodinger in describing the behavior of electrons in the atom: electron as a particle vs. electron as a wave;
discrete energy states postulated vs. derived; orbits vs. orbitals. Can the Bohr model be applied to atoms with
more than one electron or to molecules? What about the Schrodinger mechanics?
What are the results of solving the Schrodinger equation for hydrogen atom (ψ functions, |ψ|2, orbitals)? Be
able to sketch 1s, 2s, 2px , 2py , 2pz orbitals. What are the nodal surfaces (if any) each of those orbitals have?
What is the principle difference between the orbital diagrams for hydrogen atom ant many-electron atoms?
How electrons are arranged in many-electron atoms? How many orbitals are in an s, p, d, and f sublevel?
What is the maximum electron occupancy of each sublevel?
Be able to write a complete or a condensed (using the noble notation) electron configuration of an atom or
monatomic ion in its ground state using a blank periodic table as the only reference.
What are the valence electrons? Where are they located in case of s, p, and d elements?
THE PERIODIC PROPERTIES OF ELEMENTS
Textbook sections: 2.7; 8.2; 8.6 – 8.9.
Be able to give a statement of the periodic law in modern terms. What kind of elements are aligned in each
group of the periodic system? Are they similar, identical, or dissimilar in their properties? Do they generally
show more similarities in their chemical or physical properties?
Know what covalent, metallic, ionic, and van der Waals radii are.
What are the periodic trends in atomic radii? Be able to explain them in terms of electron configurations of
atoms and the effective nuclear charge.
How do sizes of atoms and the corresponding cations compare? How do sizes of atoms and the corresponding
anions compare?
What is the first, second, third etc. ionization energies? How do they compare to each other? Do ionization
energies have positive or negative values? Why? Why are there occasional gigantic jumps in the successive
ionizations energies for a particular atom?
What are the general trends in the first ionization energies across the period and down the group? Which
element has the highest first ionization energy? Which one has the lowest?
What is electron affinity? Do electron affinities usually have positive or negative values? Why? What are
the periodic trends in electron affinities?
What are the periodic trends in electronegativity? Remember that electronegativity is the property of an atom
in a chemical bond unlike ionization energy and electron affinity that are a properties of an isolated atom.
Be familiar with general chemical reactivity patterns of elements in the periodic table. What is the most
reactive nonmetal? What is the least reactive nonmetal? What is the most reactive metal? How does metallic
character of the elements changes across the period and down the group? Why is the border between metals
and nonmetals is a diagonal and not horizontal or vertical?
Know the family names for elements of groups IA, IIA, VIIA, and VIIIA.
What are the periodic trends in oxidation numbers: (i) within a group; (ii) across the period (for main-group
elements)?
Be able to give a couple of examples of reactions that are characteristic for all group IA and all group IIA
elements: reactions with halogens and reactions with water.
Be able to give a couple of examples of reactions that are characteristic for all group VIIA elements: reactions
with hydrogen and reactions with group IA and group II A metals.
Be able to correctly predict the product of: (i) a reaction of an oxide of a group IA or group IIA metal with
water (a hydroxide is the product); (ii) a reaction of an oxide of a nonmetal with water (an oxyacid is the
product)
BASIC IDEAS OF CHEMICAL BONDING
Textbook sections: 9.2 (all); 9.3 (all); 9.4 (the first and the last subsections only); 9.5 – 9.11; Fig. 10.6 on
page 443.
Know what the following characteristics of chemical bonds mean: bond length, bond energy, bond polarity,
bond order. Concept of electronegativity. Difference in electronegativities of two atoms bonded to each other
and the polarity of the bond.
Be familiar with the basic ideas of Lewis theory: the octet rule for monatomic ions, the octet rule for
molecules, covalent bond as a shared electron pair, double bond as two shared electron pairs, triple bond as
three shared electron pairs. What are the limitations of the octet rule? Is it universal for all elements in all the
molecules and ions they form?
Be familiar with a physical interpretation of the shared electron pair (compare the distribution of electron
density in each of the two isolated hydrogen atoms and in a hydrogen molecule). Be aware that there two
kinds of shared electron pairs in double and triple bonds: σ and π types.
What is bond polarity? How is it estimated based on the relative electronegativities of the atoms in a bond?
Be able to write a Lewis symbol for an atom, monatomic cation, or monatomic anion. Remember that dots
are for the electrons of the valence shell only. (The Lewis symbol for a monatomic cation may contain no dots
at all.)
Be able to describe the formation of a binary ionic compound (monatomic cation + monatomic anion) in terms
of loss/gain of electrons and the octet rule.
Be able to draw a Lewis representation of the formula unit of any ionic compound as a group (no dashes
between ions!!!) of Lewis structures for cations and anions written in the ratio resulting in a zero net charge
for the entire group.
Know a quick and easy way of drawing a Lewis structure for a simple molecule starting with Lewis symbols
for the individual atoms (“electron handshaking”).
Be able to apply the procedure (“pool-and-distribute”) for writing Lewis structures based on combining all
valence electrons and distributing them between atoms as bonding pairs (dashes) or lone pairs (pairs of dots)
in order to obtain Lewis structures for any molecule or ion.
Be aware of concepts of resonance and formal charges (refer to the class discussions) to predict the best
possible Lewis structure. Remember that resonance structures are considered when a single Lewis structure
cannot be given to adequately describe the experimentally observed properties of a molecule or a polyatomic
ion. Remember that resonance structures differ from each by the location of certain electrons within a
“stationary” molecule (polyatomic ion). They are not transformed into each simply by rotating or flipping a
molecule as a whole. They are indeed different electronic structures. Remember that in general formal
charges are hypothetical charges and used as a theoretical tool to predict the best possible electronic structures
of molecules or polyatomic ions. They may or may not be close to real charges on atoms in a molecule or a
polyatomic ion.
Concepts of bond energy, bond length, covalent radii, and van der Waals radii. Use of bond energies in
calculating the ∆H for a reaction that occur in gaseous phase. Use of bond lengths and van der Waals radii to
estimate the dimensions of molecules.
VSEPR MODEL AND SHAPES OF MOLECULES AND POLYATOMIC IONS
Textbook sections: 10.2 – 10.4.
Starting with a correct Lewis structure be able to predict and name the electron geometry and molecular
geometry (shape) of a molecule (polyatomic ion) using the rules of VSEPR theory.
Based on your predictions be able to sketch (draw on a piece of paper) a 3D structure of a molecule or a
polyatomic ion. You should include into your drawing the lone pairs for the central atom. The lone pairs on
the central atom can be shown as clouds with pairs of dots and the lone pairs on the terminal atoms are shown
simply as pairs of dots (check your lecture notes and the answer key to the worksheets on Lewis structures,
shapes and polarities). Bonds that are facing the viewer are shown as wedges and bonds directed away from
the viewer as dashed lines.
Know the ideal bond angles for each of the five VSEPR electron geometries (two clouds, three clouds, four
clouds, five clouds, and six clouds). For a particular molecule or a polyatomic ion be able to predict whether
the bond angles will be ideal or somewhat deviate from the ideal (slightly smaller or slightly larger than the
ideal).
ORGANIC MOLECULES – BASIC IDEAS
Textbook sections: 3.2
Know which compounds are classified as organic, what kind of atoms are commonly present in organic
compounds, what the common bonding patterns of those atoms are, and how formulas of organic compounds
are written. (Formulas of organic compounds are often appear in exercises on intermolecular forces and
physical properties of substances. Organic compounds are also very common in everyday life and are
undoubtedly mentioned in all biology courses for which Chem 101 is a prerequisite.)
BOND POLARITIES AND MOLECULAR POLARITIES, AND INTERMOLECULAR FORCES
Textbook sections: 9.6 (bond polarities); 9.11 (metallic bond); 11.2 (physical states); 11.3 (types of
intermolecular forces); 11.5 (omit the Clausius-Clapeyron equation); 11.9 (properties of water).
Know that in general (in chemistry) polar and nonpolar mean uneven (asymmetrical) and even (symmetrical)
distribution of electrons. The word “pole” is used in chemistry synonymously with the word “center”. A
molecule that is polar has poles, the centers of overall positive charges (atomic nuclei) and overall negative
charges (electrons of the molecule’s electron cloud).
Be able to clearly distinguish between polar bond and polar molecule. A molecule with all very polar bonds
can be nonpolar.
Be able to distinguish between chemical bonds (forces between atoms together within molecules and
polyatomic ions and within ionic crystals) and intermolecular forces (IMFs, forces that hold molecules
together in a liquid or a solid substance). Make sure to clearly distinguish the terms atom, molecule,
substance, element, and compound. Apply the term covalent bond to atoms connected to each other in the
molecule (water molecule has two H-O covalent bonds) and the term intermolecular forces to interactions
between molecules within a sample of a substance (there are hydrogen bonds between water molecules in ice
and in liquid water).
Types of intermolecular forces: dispersion (London) forces, dipole-dipole forces, and hydrogen bonds (Hbonds). Do not confuse H-bonds with N-H or O-H covalent bonds.
What are the factors that contribute most to the strength of IMFs of each type?
Know correlations between the strength of IMFs within a particular substance and some physical properties of
the substance (physical state, boiling point, melting point).
What is the difference between vaporization and boiling? Why the boiling point of liquids depends
significantly on the outside pressure?
CRYSTALLINE SOLIDS
Textbook sections: 11.10 (no calculations with Bragg’s Law); 11.11 (pages 520 – 523 only).
Use of X-ray diffraction in determination of crystal structures.
Unit cells: counting the number of atoms per a cell and calculations of the volume and the dimensions of a
cell.