Teachers, notes, videos, classes, textbooks, problems, worksheets, tests, quizzes, reports, experiments, study packets and study partners can only assist YOU to do this. In first year chemistry you have learned a great deal. Most of it applies to IB chemistry learning objectives. The most successful students STUDY AHEAD . The purpose of this study packet is for you to review and relearn topics from first year Chemistry over the summer.
If you intend to take the IB external exam and earn college credit for the course... please do not put this off until the end of the summer.
There are 8 topics. All of these topics should assist your review of first year Chemistry. Work on a topic each week . You will turn in all work completed at the start of next school year. If you are sure you will be testing IB SL Chemistry put particular emphasis in your study of the Energy option , since this will be a major part of the test you take later next year from IB. This option is not something we studied in chemistry year one, it’s the only new topic in this packet.
The rest is to review and refine your understandings.
In this study guide are the major IB objectives that have been covered in first year
Chemistry. They are sometimes different from what you have learned in chemistry year one. There are also other objectives that will be used this coming year. The IB organization has identified command verbs such as, ‘state’, ‘explain’, ‘draw’,…. It is useful to familiarize yourself with command verbs, since we will be using them throughout the year for tests and assignments .
Course textbook: Chemistry: Course Companion By Geoffrey Neuss Second Edition
Oxford University Press. This author writes many of the IB test questions.
IB online: http://ibchem.com/
IB textbook online : http://en.wikibooks.org/wiki/IB_Chemistry
General Chemistry textbook online : http://en.wikibooks.org/wiki/General_Chemistry
From year one chemistry http://www.chem1staab.blogspot.com
and from year two chemistry http://www.chem2staab.blogspot.com.
I have provided an online link for a general chemistry textbook and an IB chemistry textbook. You can also use online tutorials, videos and animations that you find. All the
IB topics are covered by Richard Thornley on his YouTube channel . The IB Chemistry curriculum is online at http://ibchem.com/ and if you click on the syllabus link, then on topics you are learning, you will find useful text material. A word of caution however, if you just copy these without understanding them you are not learning. You should look for a more detailed understanding. You must write, diagram, draw, make tables, solve problems, review, review, review until you know that you know !

We will spend some time on review and questions and answers over this packet and the material from first year Chemistry. You will turn in your work on this packet in for class credit, it is summer homework. You will be introduced to the other members of IB
Chemistry and you will form study partners and/or groups. You will take a test that will count towards your first semester grade on this year one SL material ,
. You can expect multiple choice and short answer questions.
For each topic I have included what IB expects you to be able to do or show you know.
Use the sources above for study. At the bottom of the topic I have included some advice or online links to use. You will turn in all work completed to me at the start of the school year.
1) Apply the mole concept to substances.
The mole concept applies to all kinds of particles: atoms, molecules, ions, electrons, formula units, and so on. The amount of a substance is measured in moles (mol). The approximate value of
Avogardro’s constant (L), 6.02 X 10 23 mol -1 should be known.
2) Determine the number of particles and the amount of a substance in moles.
Practice problems where you convert between the amount of substance (in moles) and the number of atoms, molecules, ions, electrons and formula units. Find practice problems in Chemistry textbooks, or online and do at least 10 different kinds.
3) Define the terms relative atomic mass (Ar) and relative molecular mass (Mr).
4) Calculate the mass of one mole of a species from its formula. The term molar mass (in g mol-1) will be used in IB.
5) Solve problems involving the relationship between the amount of substance in moles, mass and molar mass.
6) Distinguish between the terms empirical formula and molecular formula
7) Determine the empirical formula from the percentage composition or from other experimental data. You may be able to find virtual simulation experiments online to help with this.
8) Determine the molecular formula when given both the empirical formula and experimental data.
9) Deduce chemical equations when all reactants and products are given. Be sure you are aware of the difference between coefficients and subscripts. Given a reaction statement you should be able to write the reaction equation, balanced, with state symbols.
10) Identify the mole ratio of any two species in a chemical equation
11) Apply the state symbols (s),(l), (g) and (aq). You may what to consider when are these symbols necessary in understanding and when are they redundant?

12) Calculate theoretical yields from chemical equations. Given a chemical equation and the mass or amount (in moles) of one species, calculate the mass or amount of another species. (g to mol to mol to g)
13) Determine the limiting reactant and the reactant in excess when quantities of reacting substances are given.
14) Solve problems involving theoretical, experimental and percentage yield.
15) Apply Avogadro’s law to calculate reacting volumes of gases.
16) Apply the concept of molar volume at standard temperature and pressure in calculations The molar volume of an ideal gas under standard conditions is 2.24
X 10 -2 m 3 mol -1 (22.4 dm 3 mol -1 or 22.4 L/mol; because dm 3 = L)
17) Solve problems involving the relationship between temperature, pressure and volume for a fixed mass of an ideal gas
18) Solve problems using the ideal gas equation,
19) Analyze graphs relating to the ideal gas equation.
20) Distinguish between the terms solute, solvent, solution and concentration. (g dm-3 and mol dm-3 or g/L and mol/L)
21) Solve problems involving concentration, amount of solute and volume of solution.
For this topic you will need to
… And turn in next fall.
Here are some websites I found on Google search. Make sure you do some gas law stoichiometry. You may
as well… http://academic.evergreen.edu/curricular/matterandmotion/chem_phys%5Cpractic e_problems.htm
and http://www.scienceiscool.org/stoichiometry/problems.html
and http://www.sciencegeek.net/Chemistry/taters/Unit4Stoichiometry.htm
1) State the position of protons, neutrons and electrons in the atom.
2) State the relative masses and relative charges of protons, neutrons and electrons
3) Define the terms mass number (A), atomic number (Z), and isotopes of an element.
4) Deduce the symbol for an isotope given its mass number and atomic number.
5) Calculate the number of protons, neutrons and electrons in atoms and ions from the mass number, atomic number and charge. Use a periodic table. The best online table is a www.ptable.com. You should be familiar with this dynamic online table.
6) Compare the properties of the isotopes of an element.
7) Discuss the uses of radioisotopes. Examples should include C-14 in radiocarbon dating, Cobalt 60 in radiotherapy, and Iodine 131 and 125 as medical tracers.
8) Describe and explain the operation of a mass spectrometer. There is a good video at the YouTube channel wwwRSCorg called Mass Spectrometry MS

9) Describe how the mass spectrometer may be used to determine relative atomic mass using the C-12 scale.
10) Calculate non-integer relative atomic masses and abundance of isotopes from a spectrophotometer data ‘graph.’
11) Describe the electromagnetic spectrum. You should be able to identify the UV, visible and IR regions, and to describe the variation in wavelength, frequency and energy across the spectrum.
12) Distinguish between a continuous spectrum and a line spectrum.
13) Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels. Show how lines in a line spectrum directly related to differences in energy levels. What is convergence? Calculations, knowledge of quantum numbers and historical references will not be assessed in IB.
14) Deduce the full electron arrangement for atoms and ions up to Z = 20. Use a periodic table. An example would be for Z = 17 ( Chlorine) which is electron arrangement 1s2, 2s2, 2p6, 3s2, 3p5.
Remember …you can use YouTube for videos on topics and it may be as effective for you as reading and studying the textbook…
1) Describe the arrangement of elements in the periodic table in order of increasing atomic number. How are the elements arranged in a modern periodic table? What is the periodic law? The history of the periodic table will not be assessed.
2) Distinguish between the terms group and period.
3) Apply the relationship between the electron arrangement of elements and their position in the periodic table up to Z = 20.
4) Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table.
5) Define the terms first ionization energy and electronegativity.
6) Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativites and melting points for the alkali metals (Li to Cs) and the halogens (F to I). Explanations for the first four trends should be given in terms of the balance between the attraction between the nucleus for the electrons and the repulsion between electrons. Explanations based on effective nuclear charge are not required.
7) Describe and explain the trends in atomic radii, ionic radii, first ionization energies, and electronegativites for elements across period 3.
8) Compare the relative electronegativity values of two or more elements based on their positions in the periodic table.
9) Discuss the similarities and differences in the chemical properties of elements of the same group. Cover the following reactions… 1) The first three alkali metals with water. 2) The first three alkali metals with halogens Cl
2
, Br
2
and I
2
. 3)
Halogens (Cl
2
, Br
2
and I
2
) with halide ions (Cl , Br , and I ).

10) Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3 . Equations are required for reactions of
Na
2
O, MgO, P
4
O
10
and SO
3
with water.
Use ptable.com for help with this topic. Be familiar with its use.
1) Describe the ionic bond as the electrostatic attraction between oppositely charged ions.
2) Describe how ions can be formed as a result of electron transfer. Describe formation of cations, anions.
3) Deduce which ions will be formed when elements in groups 1, 2 and 3 lose electrons. (For example Mg 2+ , K + , etc)
4) Deduce which ions will be formed when elements in groups 5, 6 and 7 gain electrons.
5) State that transition elements form more than one ion. Include examples such as Fe 2+ and Fe 3+
6) Predict whether a compound of two elements would be ionic from the position of the elements in the periodic table or from their electronegativity values. You may want to use an electronegativity scale for this. Provide several examples.
7) State the formula of common polyatomic ions formed by non-metals in periods 2 and 3. Examples include NO
3
, OH , SO
4
2, CO
3
2, PO
4
3, NH
4
+ , and
HCO
3
- So you should see that the command verb ‘state’ in this case means you should have memorized these ions, their charges and symbols and names.
8) Describe the lattice structure of ionic compounds. You should be able to describe the structure of sodium chloride as an example of an ionic lattice.
9) Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei.
Single and multiple bonds should be considered. Examples include O
2
(ethyne).
, N
2
, CO
2
, HCN, C
2
H
4
(ethene), and C
2
H
2
10) Describe how the covalent bond is formed as a result of electron sharing.
Dative (coordinate) bonds are required. Examples include CO , NH
4
+ and H
3
O +
11) Deduce the Lewis (electron dot) structures of molecules and ions for up to four electron pairs on each atom. A pair of electrons can be represented by dots, an x, a combination of dots and x’s, or by a line. Note Cl-Cl is not a Lewis structure, but Cl to Cl with a total of 14 dots in pairs showing would be ok.
Provide several examples.
12) State and explain the relationship between the number of bonds, bond length and bond strength. The comparison should include the bond lengths and bond strengths of: Two carbon atoms joined by single, double and triple bonds; and the carbon atom and the two oxygen atoms in the carboxyl group of carboxylic acid
(Google it if you don’t know what that is.)
13) Predict whether a compound of two elements would be covalent from the position of the elements in the periodic table or from their electronegativity values.

14) Predict the shape and bond angles for species with four, three and two negative charge centers on the central atom using the valence shell electron pair repulsion theory (VSEPR). Examples should include CH
4
, NH
3
, H
2
O,
NH
4
+ , C
2
H
4
, SO
2
, C
2
H
2
and CO
2
15) Predict whether or not a molecule is polar from its molecular shape and bond polarities. Include several examples. You might try to Google it and use some of the previous examples.
16) Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C 60 fullerene which is commonly called Bucky balls)
17) Describe the structure of and bonding in silicon and silicon dioxide.
18) Describe the types of intermolecular forces (attractions between molecules that have temporary dipoles, permanent dipoles or hydrogen bonding) and explain how they arise from the structural features of molecules. The term van der Waals’ forces can be used to describe the interaction between non-polar molecules.
19) Describe and explain how intermolecular forces affect the boiling points of substances. Compare hydrogen bonding for HF and Cl; H
2
O and H
2
S; NH
3
and
PH
3
; CH
3
OCH
3 and CH
3
CH
2
OH; CH
3
CH
2
CH
3
, CH
3
CHO and CH
3
CH
2
OH.
20) Describe the metallic bond as the electrostatic attraction between a lattice of positive ions and delocalized electrons.
21) Explain the electrical conductivity and malleability of metals.
22) Compare and explain the properties of substances resulting from different types of bonding. Examples should include melting and boiling points, volatility, electrical conductivity and solubility in non-polar and polar substances.
1) Outline the characteristics of chemical and physical systems in a state of equilibrium. Make a drawing of a conical beaker with H
2
O in it and explain what is going on in the equilibrium in the closed system. Make a sketch of a graph of concentration over time for a reaction going to equilibrium and show reactants and products on the graph.
2) Deduce the equilibrium constant expression (K c
) from the equation for a homogenous reaction. Give several examples from different equations.
Consider gases, liquids and aqueous solutions.
3) Deduce the extent of a reaction from the magnitude of the equilibrium constant. When the equilibrium constant is much less than or much greater than
1, what does that mean exactly?
4) Apply Le Chatelier’s principle to predict the qualitative effects of changes of temperature, pressure and concentration on the position of equilibrium and on the value of the equilibrium constant. Explain how change in concentration, in temperature, in pressure and how adding a catalyst affects the position of equilibrium. You do not need to state Le Chatelier’s principal in IB testing later, but it may help you to understand this assessment statement.
5) State and explain the effect of a catalyst on an equilibrium reaction.

6) Apply the concepts of kinetics and equilibrium to industrial processes.
Explain the Haber process and Contact process. Explain how change in concentration, in temperature, in pressure and how adding a catalyst affects these processes.
1) Define acids and bases according to the Bronsted-Lowry and Lewis theories.
Use balanced reaction equations for examples. Write definitions of acids and bases in both theories and explain what is being transferred in each one, and in what direction.
2) Deduce whether or not a species could act as a Bronsted-Lowry and/or Lewis acid or base. Explain when a chemical species could be a Lewis acid but not a
Bronsted-Lowry acid.
3) Deduce the formula of the conjugate acid (or base) of any Bronsted-Lowry base (or acid). Use example equations to show these formulas, labeled. Make sure you clearly show the location of the proton transferred.
4) Outline the characteristic properties of acids and bases in aqueous solution.
Bases that are not hydroxides, such as ammonia, soluble carbonates and hydrogencarbonates should be included. Alkalis are bases that dissolve in water.
What are the effects of indicators on the reactions of acids with bases, metals and carbonates?
5) Distinguish between strong and weak acids and bases in terms of the extent of dissociation, reaction with water and electrical conductivity.
6) State whether a given acid or base is strong or weak. Give at least two examples of strong and weak acids and bases. Make a table.
7) Distinguish between strong and weak acids and bases, and determine relative strengths of acids and bases using experimental data. How can pH be used to do this? How can conductivity be used to do this?
8) Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale.
9) Identify which of two or more aqueous solutions are more acidic or alkaline using pH values. What would be the pH in an aqueous solution that is alkaline?
Acidic? Neutral?
10) State that each change of one pH unit represents a 10-fold change in the hydrogen ion concentration. Make a scale that represents the relationship between pH and powers of 10.
11) Deduce changes in hydrogen ion concentration when the pH of a solution changes by more than one pH unit. Explain how pH changes are related to hydrogen ion concentration in solutions.
1) Describe and give examples of random uncertainties and systematic errors.
2) Distinguish between precision and accuracy. Is is possible for a measurement to have great precision yet be inaccurate (for example, if the top of a meniscus is read in a pipette or a measuring cylinder; how does that differ)?

3) Describe how the effects of random uncertainties may be reduced. How are random uncertainties, but not systematic errors reduced by repeated readings?
Explain this.
4) State random uncertainty as an uncertainty range (+/-). Use measurements from different instruments, such as from a ruler, a burette, and an electronic balance as examples.
5) State the results of calculations to the appropriate number of significant figures. How does the number of significant figures differ when a measurement is added/subtracted verses when it is multiplied/divided?
1) State the essential idea : Societies are completely dependent on energy resources. The quantity of energy is conserved in any conversion but the quality is degraded.
2) Define energy and state a few of the units it can be measured in.
3) Define the difference between renewable and non-renewable energy and discuss the sources of each type. ( The energy of fossil fuels originates from solar energy which has been stored by chemical processes over time. These abundant resources are nonrenewable but provide large amounts of energy due to the nature of chemical bonds in hydrocarbons.)
4) Write the equations for energy density, for the specific energy, and for the efficiency of an energy transfer (expressed as a percentage).
5) Go to oilprice.com and find and state the order for the main energy sources currently used in the world. Top energy source to bottom energy source…
6) Define narratively and model with chemical structures the following: crude oil, petroleum, coal, natural gas.
7) Explain ‘cracking’ and ‘catalytic reforming.’
8) Explain octane rating.
9) Explain carbon footprint.
10) Make a table or diagram of the relative advantages and disadvantages of fossil fuels including coal, oil and natural gas.
11) The fusion of hydrogen nuclei in the sun is the source of much of the energy needed for life on earth. There are many technological challenges in replicating this process on earth but it would offer a rich source of energy. Fission involves the splitting of a large unstable nucleus into smaller stable nuclei. Widespread use of nuclear fission for energy production would lead to a reduction in greenhouse gas emissions, but it creates radioactive waste and there are serious storage and disposal issues involved with it’s use.
Explain nuclear fission and fusion using narrative, diagrams and models.
12) Solar energy can be converted to chemical energy in photosynthesis. Visible light can be absorbed by molecules that have a conjugated structure with an extended system of alternating single and multiple bonds. Light can be absorbed by chlorophyll and other pigments. Photosynthesis converts light energy into chemical energy. What is the chemical equation to model this process?
13) Fermentation of glucose produces ethanol, which can be used as a biofuel. What is the chemical equation for this process? What are the advantages of biofuels? What are the disadvantages?

14) Why are vegetable oils (which have similar energy content to diesel fuel) not used in internal combustion engines?
15) How is transesterification used to produce oils with lower viscosity that can be used in diesel engines?
16) The claims of cold fusion were dismissed as the results were not reproducible.
Discuss if it is always possible to obtain replicable results in natural sciences.
17) Gases in the atmosphere that are produced by human activities are changing the climate as they are upsetting the balance between radiation entering and leaving the atmosphere. Some people question the reality of climate change, and question the motives of scientists who have ‘exaggerated’ the problem. However, most in the scientific community agree that the are two likely effects due to global warming. 1) changes in agriculture and biodistribution as the climate changes. 2) rising sea levels due to thermal expansion and melting polar ice caps and glaciers. Create a table of the sources, relative abundance and effects of different greenhouse gases.
18) Explain some different approaches to the control of carbon dioxide emissions.
19) Explain the molecular mechanisms by which greenhouse gases absorb infrared radiation.
20) Discuss the evidence for the relationship between the increased concentration of gases and global warming.