CP Chemistry Final Exam Study Guide June 2013
In general, you should review the concepts in the Study Guide at the end of each chapter. This will help you determine if you have mastered the contents of the chapter. Also, you should review the worksheets distributed during the semester, and you should try to solve the types of problems encountered during the terms. If you kept any exams, look at them as well
Chapter 1
Recognize the importance of safety when working in a chemistry laboratory, and know the most common safety rules we practiced during the year.
Understand the metric system (units of 10), including the common prefixes used in chemistry. (See page 20.)
Recognize the basic metric measurements for length, volume and mass (meters, liters, grams).
Know about the International Systems Units (SI ).
Measurement: certain and estimated digits; understand the reasons for differences in results. Recognize the difference between precision and accuracy.
Be able to convert units from small to large, and from large to small.
Examples: meters to centimeters
grams to kilograms
liters to milliliters
Significant figures “sig figs”
Atlantic – Pacific Rule
If there is a decimal present (Pacific) begin counting sigfigs from the left starting with the first nonzero number.
If there is no decimal present (decimal absent, Atlantic) start from the right with the first nonzero number.
Be able to solve problems using the correct number of significant figures. Know the rules to determine significant figures in addition and multiplication, and for subtraction and division.
Addition: number of sigfigs in the answer will depend on the number with the highest place value (tenths, hundredth, units, tens etc.) of uncertainty.
Multiplication: number of sigfigs in the answer depends on the number with the least number of sigfigs.
Know how to write numbers in correct scientific notation.
Study vocabulary terms found at the end of the chapter on page 49.
Chapter 2
Recognize examples of Kinetic Energy, Potential Energy & Radiant Energy.
Laws of Conservation of Energy – see textbook page 57
Temperature Scales (Celsius and Kelvin) and how to convert them. What is Absolute Zero?
States of Matter (Solid, Liquid, Gas and Plasma)
Physical and chemical change; The Law of Conservation of Matter.
Elements, Compounds and Mixtures
Know the definitions of each of these terms, how to identify each type.
Chapter 3
Early models of the Atom (Use the summary sheets distributed earlier. Recognize the important scientists who contributed to our understanding of atomic structure throughout history.)
Modern Atomic Theory: Protons, Electrons, Neutrons, Nucleus (What are they, where are they located, and what are their relative size and charges?)
Identify: Atomic Numbers, Ions, Isotopes, Atomic Mass.
Section 3-4: This deals with radioactivity, so you should know the types (alpha, beta, gamma), their properties and health hazards. You should also be able to balance nuclear equations when given a nucleus and/or radioactivity type.
Chapter 4
Know the four characteristics of an electromagnetic wave. (p126)
Know the key components of Quantum Theory (Planck’s Theory, dual properties of radiant energy).
Know what a line spectrum is vs. a continuous spectrum. How does an atom emit light? How is this related to Quantum Theory?
Know about the Bohr Model of the atom and Heisenberg’s Uncertainty Principle.
Recognize that the quantum-mechanical model of the atom deals with the probability of finding electrons within certain regions of space (atomic orbitals) having specific shapes and energies. (Electron orbitals: s,p,d,f)
Review Electron Configuration diagrams, and the rules for filling orbitals (Aufbau Principle, Pauli Exclusion Principle & Hund’s Rule).
Chapters 5-6
History of the Periodic Table & contributions of the following people:
¨ Moseley
¨ Mendeleev
¨ How to read/use the modern Periodic Table
Periodic Law: When elements are arranged in order of increasing atomic number, their physical and chemical properties show a periodic pattern. Be able to use the general trends of the following:
§ Electron affinity & electronegativity
§ First & Successive Ionization Energies
§ Atomic Radius & Ionic Radius
§ Ionic Charge & Ion Size
Be able to read and use the Periodic Table. (What is in each ‘square’ of the Table?)
Groups (the columns or families) and Periods (the rows)
q Names of the common Groups (Alkali, Alkaline Earth, Halogen, Noble Gas, Transition elements), and know their general properties.
q Reactivity trends in the Periodic Table (Atomic and ionic size, ionization energy, electronegativity
q General properties of Metals, Nonmetals and Metalloids (or Semimetals)
Locations of the Sub-level blocks on the Periodic Table
v s,p,d,f-blocks
Electron Configuration
Ø Be able to follow the Aufbau, Pauli and Hund rules to write the long form and the noble gas abbreviations of electron configuration diagrams for simple atoms, and how these relate to the periodic trends.
Chapter 7
Ionic Bonding
Ionic Bonding and Ionic Compounds (How do they form?)
Octet Rule
Drawing simple Lewis Dot Diagrams following the rules
Cations and Anions (including Monatomic and Polyatomic Ions)
Empirical (simplest) formulas
Writing and Naming Ionic formulas Correctly:
¨ Total + and - charges need to be equal
¨ Learn the common ions and their charges, including the polyatomic ions (See the handout.)
¨ Be able to name Ionic Compounds
¨ Use the criss-cross method to write correct formulas of ionic compounds
¨ Know the common acid names and formulas
Covalent Bonding
Draw molecular formulas
Draw structural formulas
Recognize Single, Double and Triple Bonds
Predicting Polar and Non-Polar Bonds using the electronegativity difference between atoms (p241)
Be able to write formulas for Molecular Compounds from names, and to name formulas
Know the common numerical prefixes (p 246) used to name these compounds.
Chapter 8
Know about polarity and how molecular shape determines polarity. What is the role of electronegativity? Why is polarity important in affecting molecular interactions? (Water is an excellent example. Why?)
Chapter 9
Types of Chemical reactions and examples:
Combustion (Needs O2 as a reactant)
Direct combination (A + B à AB) (One product is formed)
Decomposition reaction (AB à A + B) (Reactant breaks down into components)
Single replacement (A + BC à AC + B) (One element replaces another)
Double replacement (AB + CD à AD + CB) (Partners swap places)
Note that electrons are transferred in all but double replacement reactions (just ions are present).
Know the difference between a reactant and a product, and where each appears in an equation.
Be able to explain what a chemical reaction is and why it occurs.
Be able to write balanced chemical equations (change the coefficients), which means you will need to know ionic charges and be able to correctly write compound formulas. (See Chapter 7.)
Chapter 10
Define “mole” and describe why/how it is used.
Identify Avogadro's number (6.022 X 1023 ‘things’)
Identify Atomic Mass from Periodic Table
Calculate formula mass (in amu) and molar mass (in grams) of a molecule from its formula and the atomic mass of individual elements. (Review pages 319 and 320.)
Be able to convert:
Moles to mass & mass to moles (using the molar mass in grams)
Moles to particles & particles to moles (using Avogadro’s number)
Moles to volume & volume to moles (This is only for a gas at STP using the molar volume, 22.4 liters/mole.) (STP is Standard Temperature, O°C, and Pressure, 1 atmosphere.)
Use the Factor Label Method to set up and solve problems!
Determine percent composition of a molecule from atomic mass and molar mass.
Find the empirical formula and molecular formula from known lab data.
(Practice the problems on your worksheets so you remember how to do them. See the text and your notes for examples.)
Chapter 11
Be able to interpret balanced chemical equations.
Be able to solve:
Mole-mole problems (from chemical equations, formulas and molar quantities)
Mass-mass problems (same)
Mass-volume problems (same)
Volume-volume problems (same)
Identify ‘limiting reactants’ and solve ‘limiting reactant’ problems.
Determine the percent yield of a chemical reaction [Actual Yield divided by Expected Yield) X 100%].
Chapter 12
Recognize that thermochemistry is the study of the heat effects in chemical reactions.
Exothermic reactions release heat (term appears on product side of the equation), but endothermic reactions absorb heat (term appears on reactant side of equation).
Know the concept of enthalpy (ΔH), the heat absorbed or released in a chemical reaction at constant pressure (that is, without pressure volume work, the PV term).
Recognize that the Standard Enthalpy Change, ΔH°, for a reaction is measured when the reactants and products are in their ‘standard states. Exothermic reaction have negative ΔH° (loss of heat from the system), and endothermic reactions have positive ΔH° (gain of heat by the system). Enthaply changes may be calculated from ΔH° and the number of moles.
Calculate enthalpy changes for a chemical reaction. Determine specific heat capacity and heat capacity. Use q = m X C X ΔT and qrxn = -qsurr correctly.
According to Kinetic Theory, particles of matter have kinetic energy, and heat is the transfer of kinetic energy from a hotter object to a colder object. This is in contrast to the old Caloric Theory which thought that heat was a fluid called caloric.
Recognize the difference between joules, calories and Kilocalories (Calories) and interconvert them.
Chapter 13
List the properties of gases, and how the Kinetic-Molecular Theory explains these properties.
Know the difference between elastic and inelastic collisions, which helps explain gas behavior.
List the postulates of the Kinetic-Molecular Theory. (See page 423 and handout.)
Describing gases requires that we know the amount (moles, n), volume, temperature and pressure. The gas laws contain these terms, but the Combined Gas Law is the easiest to remember and to use:
P1V1/T1 = P2V2/T2
[See the handout for a review of Boyle’s Law (P&V), Charles’ Law (V&T), Gay-Lussac’s Law (P&T), Avogadros’ Law (V-n at constant T,P.]
Always ‘go Kelvin’ when solving gas law problems!
The Ideal Gas Equation, PV = nRT, describes the physical behavior of an ideal gas, one that satisfies all of the postulates of Kinetic-Molecular Theory.
Real gases have volume and exert force on each other, so their behavior is not ideal. But many gases behave like ideal gas at high temperature and low pressure because these conditions minimize gas molecule interactions.
Dalton’s Law of Partial Pressures indicates that gases behave independently of each other in a mixture, so PT= pa + pb + pc + … (Total pressure is the sum of the partial pressures.)
Some gases, like hydrogen and helium, have low density, so they have ‘lifting power.’
Gases diffuse, and the way to measure this is to do an effusion experiment where gas particles are allowed to slowly pass through a tiny opening. Smaller gases move the fastest (effusion), and larger ones move slower.
Chapter 14
As materials are heated, they pass through several phases—solid liquid and gas.
Energy is required for phase transitions. Energy added to a substance at the melting/freezing point or boiling/condensation point will not increase the temperature of the material until the phase transition is complete. All energy absorbed by the material at these temperatures goes to break inter-molecular bonds.
Apart from phase changes, any energy absorbed or released from a substance goes to increase the kinetic energy of the particles, resulting in increased particle speed, measured as increased temperature.
At a given temperature, all particles, regardless of phase, have the same average particle speed.
Differences in melting point, boiling point or heat capacity of substances is related to differences in the relative strengths of intermolecular bonds, which takes us back to chapter 7, bond types.
Students should be able to represent the chapter 14 concepts, as outlined here, in the form of a heating curve, energy transfer diagrams, or as algebraic formulas (q=mCDT , DHvaporization= q/amount)
AS YOU CAN SEE, WE’VE COVERED A LOT THIS YEAR. DON’T BE AFRAID TO USE YOUR TEXTBOOK, ESPECIALLY THE REVIEW SECTIONS AT THE END OF EACH CHAPTER. NOW IS THE TIME TO GET STARTED. STUDY HARD!