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IB Chemistry HL -1 Juniors

Welcome to the IB Program and IB Chemistry HL! This summer assignment is designed as a review of the major concepts from Honors Chemistry that will be developed further in IB Chemistry HL. In addition to developing your knowledge of advanced chemical concepts, you will also explore chemical relationships through designing and performing labs. You will not be required to turn in this assignment; however, you will take a formative quiz over these review topics during the second week of school. Please take the time to carefully go through these topics, review the concepts, and memorize the necessary material, or you will not be prepared for this course.

Supply List: Please purchase the following supplies prior to the first day of class.

  • Graphing calculator
  • 1.5 inch three-ring binder
  • Notebook paper
  • Blue or black pens
  • I also encourage you to purchase a study book for the IB Chemistry Exam. I recommend (however, you may purchase a different one): IB Chemistry: Study Guide: Oxford IB Diploma Program, 2014 edition (Paperback), by Geoffrey Neuss, Publication Date:October 29, 2014, ISBN-10:0198393539, ISBN-13:978-0198393535 **Be sure to get the 2014 edition (first examinations 2016)-the older version is VERY different!**

Internet Sources:

Khan Academy Chemistry

ChemGuide – UK (very helpful throughout the year, as it is specific to IB Chemistry):

Chemistry Lectures

Content to Memorize: This information was all required in Honors Chemistry – make sure you still know it.

  • Polyatomic Ions: You must know the formula and charge for each
  • Charges of Groups 1, 2, 3, 5, 6, and 7 on the periodic table
  • Strong Acids (as required for IB): HI, HBr, HCl,HClO 3, HClO 4 , HCl, HNO 3 , H 2 SO 4
  • Strong Bases (as required for IB): NaOH, KOH, RbOH, Ba(OH) 2
  • Metric (SI Units):
Property Unit Symbol
Mass kilogram kg
Time second d
Temperature Kelvin K
Volume cubic meter m3
Pressure Pascal Pa

Vocabulary: Use a trusted source on the internet and/or your Honors Chemistry notebook to define the following terms in your own words, and give examples as appropriate. Create
a note card for each term.

  1. Element
  2. Atom
  3. Compound
  4. Proton
  5. Neutron
  6. Electron
  7. Isotope
  8. Ion
  9. Relative atomic mass
  10. Period (on the periodic table)
  11. Group (on the periodic table)
  12. Transition elements
  13. Alkali metals
  14. Alkaline earth metals
  15. Halogens
  16. Noble gases
  17. Ionization energy
  18. Atomic radius
  19. Electronegativity
  20. Melting point
  21. Boiling point
  22. Ionic bond
  23. Covalent bond
  24. Cation
  25. Anion
  26. Conductivity
  27. Avogadro’s constant (number)
  28. Molecular Formula
  29. Empirical Formula
  30. Sublimation
  31. Reactants
  32. Products
  33. Solute
  34. Solvent
  35. Solution
  36. Precipitate (as in chemistry)
  37. Molarity
  38. Aqueous
  39. Saturated
  40. Unsaturated
  41. STP
  42. Enthalpy
  43. Exothermic reaction
  44. Endothermic reaction
  45. Hess’s Law
  46. Kinetic Molecular Theory
  47. Strong Acid
  48. Weak Acid
  49. Bronsted-Lowry Acid
  50. Bronsted-Lowry Base
  51. pH
  52. Neutralization reaction

Quantitative Chemistry: Solve the following problems, showing all of your work. Include units and the appropriate number of significant figures in your answers. If you are struggling, refer to your honors chemistry notebook and use the internet sources provided above to review.

  1. Convert 4,672,000,000 into scientific notation.
  2. Convert 0.000005210 into scientific notation.
  3. Convert 50.0 g to milligrams.
  4. Convert 150. mL to liters.
  5. How many significant figures are in the number 4.0070 x 10 12 ?
  6. An object has a mass of 40.1g and occupies a volume of 8.20 mL. What is the density of the object?
  7. Calculate the percent yield if 28.0g of MgCl 2 is produced, but 32.0g of MgCl 2 should have been produced.
  8. How many atoms are in 52.4g of nickel?
  9. 6.00g of water contains how many moles of water?
  10. What is the molar mass of methane?
  11. How many hydrogen atoms are in 3.0 moles of ethanol, C 2 H 5 OH?
  12. What is the empirical formula of glucose, C 6 H 12 O 6 ?
  13. A compound with an empirical formula of CH 2 has a molecular mass of 42.09. What is its molecular formula?
  14. A compound of nickel has a mass composition of 37.9% nickel, 20.7% sulfur, and 41.4% oxygen. What is its empirical formula?
  15. Aluminum and iron(III) oxide react to form iron and aluminum oxide. What mass of iron is produced from the reaction of 21.4g of aluminum and 91.3g of iron(III) oxide? What is the limiting reactant? What is the excess reactant?
  16. What volume of nitrogen forms when 100. g of ammonia, NH 3 , decomposes completely into its elements at STP?
  17. A helium party balloon has a volume of 12.0L. At room temperature (25°C) the internal pressure is 1.05atm. Calculate the number of moles of helium in the balloon.
  18. The gas left in a used aerosol can is at a pressure of 1.00atm at 27.0°C. If this can is thrown into a fire, what is the pressure of the gas when its temperature reaches 927 °C?
  19. The volume of a gas is 20.0L at 275K and 92.1kPa. Find its volume at STP.
  20. A solution is made by dissolving 17.0g of lithium iodide in enough water to make 387mL of solution. What is the molarity of the solution?
  21. What volume of 18.0M sulfuric acid is required to prepare 16.5L of 0.126M sulfuric acid?
  22. Calculate the [OH – ] in a solution that has a pH of 3.70.
  23. A solution has a pH of 4.37. What is the [H + ] in the solution?
  24. A 15.6g sample of ethanol absorbs 868J of heat. If the initial temperature of the ethanol was 21.5°C, what is the final temperature of the ethanol?

Concept Review: Answer the conceptual questions below. If you don’t remember a topic, use the internet sources provided above to review.

  1. Determine the number of protons, neutrons, and electrons for each:
    1. Sulfur
    2. Chloride
    3. Calcium ion
  2. Which is larger? Ca or Ca +2 Why?
  3. Which is larger? F or F -1 Why?
  4. Why is sodium larger than chlorine?
  5. Why is fluorine smaller than iodine?
  6. Why does it take less energy to remove an electron from Potassium than from Bromine?
  7. List the following elements in order from smallest to largest electronegativity: Magnesium, Sulfur, Francium
  8. Write full length electron configurations for Na, Al, and Cl 1-
  9. Draw dot diagrams for Nitrogen and Fluorine.
  10. Draw the Lewis structures for NH 3 and CO 2 .
  11. Write and balance chemical equations for:
    1. The combustion of methane
    2. The single replacement reaction of zinc and hydrochloric acid
    3. The neutralization reaction of sulfuric acid and sodium hydroxide
    4. The decomposition of dinitrogen pentoxide.
  12. Endothermic reactions have a (negative/positive) ∆H, and the products are (more/less) stable than the reactants.
  13. Exothermic reactions have a (negative/positive) ∆H, and the products are (more/less) stable than the reactants.
  14. Why are aqueous solutions of ionic compounds considered electrolytes?

I will be checking email periodically (not daily) throughout the summer if you have questions – scheroj@cfbisd.edu

Have a wonderful, safe summer! Mrs. Schero