SCH4U Chemistry Review: Fundamentals

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1 SCH4U Chemistry Review: Fundamentals Particle Theory of Matter Matter is anything that has mass and takes up space. Anything around us and in the entire universe can be classified as either matter or energy. The Particle Theory of Matter states that: 1. Matter is made up of tiny particles (Atoms & Molecules) 2. Particles of Matter are in constant motion. 3. Particles of Matter are held together by very strong electric forces 4. There are empty spaces between the particles of matter that are very large compared to the particles themselves. 5. Each substance has unique particles that are different from the particles of other substances 6. Temperature affects the speed of the particles. The higher the temperature, the faster the speed of the particles. Atomic Theory Elements are substances that cannot be chemically broken down into simpler substances (although they can be chemically combined to create compounds). Examples of elements include iron, oxygen and chlorine. This property of elements is explained by atomic theory which states that elements consist of indivisible particles called atoms. Atoms can combine with each other to form molecules, which are the particles that make up compounds. One of the strongest pieces of evidence for atomic theory is that substances only combine with each other in specific ratios, which corresponds to the number and type of atoms that their particles are made of. Atoms are made of 3 subatomic particles: electrons, protons and neutrons. The protons and neutrons are located in a tiny nucleus in the center of the atom, and the electrons move around the nucleus at a great distance. The Periodic Table of Elements organizes all the elements according to the number of protons, which is called the atomic number. Ions An ion is an atom or molecule where the total number of electrons is not equal to the total number of protons, giving it a net positive or negative electrical charge. An anion (pronounced an-eye-on ) is an ion with more electrons than protons, giving it a net negative charge (since electrons are negatively charged and protons are positively charged). Conversely, a cation (pronounced cateye-on ) is an ion with more protons than electrons. An ion consisting of a single atom is a monatomic ion. If it consists of two or more atoms, it is a polyatomic ion. Polyatomic ions containing oxygen, such as carbonate and sulfate, are called oxyanions. When writing the chemical formula for an ion, its charge is written as a superscript '+' or ' ' following a number indicating the difference between the number of protons and the number of electrons. The number is omitted if it is equal to 1. For example, the sodium cation is written as Na +, the '+' indicating that it has one less electron than it has protons. The sulfate anion is written as SO4 2, the '2 ' indicating that it has two more electrons than it has protons. Since elements in the same column for main group elements have the same number of valence electrons, then those elements will all form ions with the same charge. For example, all alkali metals form ions with a charge of 1+ while halogens will form ions with a charge of 1-. For the transition metal elements, many of them can form more than one type of ion. For example, iron can form a stable ion with a charge of 2+ and it can also form an ion with a charge of 3+. These elements are sometimes referred to as multi-valent. Not all transition metals are multi-valent, however. For example, silver only forms an ion with a charge of 1+. Some common transition metals and their charges are shown below.

2 It is very useful to memorize the most common polyatomic ions and their charges. Here is a table of common polyatomic ions: Table 3 - Some transition metals and their ions Table 4 Some common polyatomic ions and their ions Element Charge Name Charge Formula Iron +2/+3 Ammonium +1 NH4 + Copper +1/+2 Bromate -1 BrO3 - Lead +2/+4 Chlorate -1 ClO3 - Tin +2/+4 Hydrogen carbonate -1 HCO3 - Mercury +1/+2 Hydroxide -1 OH - Nickel +2/+3 Iodate -1 IO3 - Chromium +2/+3/+6 Nitrate -1 NO3 - Cobalt +2/+3 Carbonate -2 CO3 2- Manganese +2/+3/+4/+7 Chromate -2 CrO4 2- Zinc +2 Sulfate -2 SO4 2- Silver +1 Phosphate -3 PO4 3- Balancing Equations TIPS FOR BALANCING EQUATIONS Be sure that each molecular formula is written correctly and each compound is neutral Mentally (or physically) count or tally how many of each type of atom is present on each side of the equation. Begin by balancing elements that are only found in one substance on each side of the equation. Balance oxygen and hydrogen LAST they usually balance out at the end or perhaps only the number of water molecules will need to be adjusted. If there is an odd number of an element on one side and an even number on the other, the odd number will need to be evened out ---for example, use a coefficient of 2 for that substance. If there are polyatomic ions (e.g. SO4 2- ) that remain together as a unit during the reaction, count the polyatomic ion as a unit When tallying, be sure to adjust the count for each and every element that an added coefficient affects. Combustion reactions that don t seem to balance will often come out better if a coefficient of 2 is used for the hydrocarbon. Predicting and Balancing Reactions Being able to predict the products of a chemical reaction comes with experience. Use the following information to help you with predicting the chemical reactions below: 1. Carbonates decompose to carbon dioxide and the oxide of the metal ion 2. Chlorates decompose to oxygen and the chloride of the metal ion 3. Non-metal oxides form acidic solutions when dissolved in aqueous solutions 4. Metal oxides form basic solutions when dissolved in aqueous solutions 5. Electrolysis is a method of decomposing ionic compounds 6. Peroxides decompose to form oxygen and an oxide Stoichiometry Problems Steps to Solving a stoichiometry Problem: 1. Write a complete BALANCED equation.

3 2. Write all information given in the problem under the appropriate chemical formula in your equation. 3. If given a mass, calculate the number of moles of all species given. 4. Determine the limiting reagent by dividing the number of moles of each compound by its coefficient from the balanced equation. The smallest number determines which is the limiting reagent. 5. Using the number of moles of limiting reagent, determine the number of moles of your unknown species using a mole ratio. The first line of the mole ratio uses the coefficients from the equation and the second line uses the value of your known and an X to represent the unknown value. Cross multiply your mole ratio and solve for X. 6. Determine the mass of your unknown. 7. If asked for the mass of your excess repeat step 5 to determine the moles of excess reagent USED UP. To find what is left over subtract this number from the original number of moles calculated in step Finish the question with a statement. Practice questions Significant Digits 1. Give the correct number of significant figures for 4500, 4500., , Give the answer to the correct number of significant figures: =? 3. Give the answer to the correct number of significant figures: =? 4. Give the answer to the correct number of significant figures: (1.3 x 10 3 )(5.724 x 10 4 ) =? 5. Give the answer to the correct number of significant figures: (6305)/(0.010) =? Balancing Reactions Balance and classify the reaction type for each of the following equations. 1. AgNO3 + CuCl2 AgCl + Cu(NO3)2 2. C8H18 + O2 CO2 + H2O 3. FeS2 + O2 Fe2O3 + SO2 4. (NH4)2SO4 + NaOH NH3 + H2O + Na2SO4 5. Mg + N2 Mg3N2 6. Al4C3 + H2O Al(OH)3 + CH4 7. PbO2 PbO + O2 8. Ca(OH)2 + HNO3 Ca(NO3)2 + H2O 9. C3H8 + O2 CO2 + H2O 10. Al + NaOH + H2O NaAl(OH)4 + H2 11. CuSO4 + Fe Fe2(SO4)3 + Cu 12. Potassium + oxygen potassium oxide 13. Calcium + water calcium hydroxide + hydrogen 14. Copper (I) sulfide + oxygen copper (I) oxide + sulfur dioxide 15. Sodium chloride + lead (II) nitrate sodium nitrate + lead (II) chloride 16. Zinc oxide + hydrochloric acid zinc chloride + water 17. Aluminum + iron (III) oxide aluminum oxide + iron 18. Magnesium oxide + phosphoric acid magnesium phosphate + water 19. Iron (III) sulfate + calcium hydroxide iron (III) hydroxide + calcium sulfate 20. Calcium + chloric acid calcium chlorate + hydrogen Predicting Chemical Reactions For each of the equations below, predict and balance the following synthesis and decomposition reactions. Use abbreviations to indicate the phase of reactants and products where possible (i.e. aq, s, l, g) 1. A sample of calcium carbonate is heated. 2. Sulfur dioxide gas is bubbled through water. 3. Solid potassium oxide is added to a container of carbon dioxide gas. 4. Liquid hydrogen peroxide is warmed. 5. Solid lithium oxide is added to water. 6. Molten aluminum chloride is electrolyzed 7. A pea-sized piece of sodium is added to a container of iodine vapour. 8. A sample of carbonic acid is heated. 9. A sample of potassium chlorate is heated. 10. Solid magnesium oxide is added to sulfur trioxide gas. Stoichiometry Review Questions Answer the following questions. Make sure your final answer has the correct number of significant digits. 1. Determine the number of moles present in 10.0 g of hydrogen gas. 2. Determine the mass of nitrogen tetroxide molecules present in 0.20 mol of N2O4. 3. Calculate the molar mass of a substance if mol of it weighs 3.2 g. Which of the following substances is it? CO2, SO2, N2O4, Cl2 4. How many molecules of ammonia are present in 1.7 g of ammonia gas? 5. Calculate the percentage composition of gold chloride (AuCl4) 6. You are asked to measure out 0.50 mol of calcium chloride. How many grams would you need?

4 7. How many sodium ions and chloride ions are there in a g sample of sodium chloride? 8. Write the equation for the combustion of carbon in oxygen. 9. What mass of carbon would be necessary to produce g of carbon dioxide? 10. What mass of oxygen would be consumed in this reaction? 11. Ammonia gas and oxygen react to produce nitrogen dioxide and water. Determine the mass of ammonia that can be consumed in this reaction if there are 20.5 g of pure oxygen available. 12. Sodium hydroxide is a strong base and sulfuric acid is a strong acid. a. Write the equation for the neutralization of sodium hydroxide by sulfuric acid. b. What mass of sodium hydroxide can be neutralized by 1.5 mol of sulfuric acid? 13. Silicon carbide is very hard and is used as an abrasive on sand paper. It is made by heating a mixture of carbon and sand in a furnace. SiO2 + 3C SiC + 2CO a. Calculate the mass of SiO2 that will react with 86 g of carbon b. What mass of silicon carbide would be produced if 6.0 g of silicon dioxide was reacted? Answers Significant Digits 1. 2, 4, 2, , or 5.09 x x x 10 5 Balancing Reactions 1. 2AgNO3 + CuCl2 2AgCl + Cu(NO3)2 (double displacement) 2. 2C8H O2 16CO2 + 18H2O (combustion) 3. 4FeS2 + 11O2 2Fe2O3 + 8SO2 (single displacement) 4. (NH4)2SO4 + 2NaOH 2NH3 + 2H2O + Na2SO4 (double displacement) 5. 3Mg + N2 Mg3N2 (synthesis) 6. Al4C3 + 12H2O 4Al(OH)3 + 3CH4 (double displacement) 7. 2PbO2 2PbO + O2 (decomposition) 8. Ca(OH)2 + 2HNO3 Ca(NO3)2 + 2H2O (double displacement/neutralization) 9. C3H8 + 7O2 3CO2 + 4H2O (combustion) 10. 2Al + 2NaOH + 6H2O 2NaAl(OH)4 + 3H2 (single displacement) 11. 3CuSO4 + 2Fe Fe2(SO4)3 + 3Cu (single displacement) 12. 4K + O2 2K2O (synthesis) 13. Ca + 2H2O Ca(OH)2 + H2 (single displacement) 14. 2Cu2S + 3O2 2Cu2O + 2SO2 (single displacement) 15. 2NaCl + Pb(NO3)2 2NaNO3 + PbCl2 (double displacement) 16. ZnO + 2HCl ZnCl2 + H2O (double displacement) 17. 2Al + Fe2O3 Al2O3 + 2Fe (single displacement) 18. 3MgO + 2H3PO4 Mg3(PO4)2 + 3H2O (double displacement) 19. Fe2(SO4)3 + 3Ca(OH)2 2Fe(OH)3 + 3CaSO4 (double displacement) 20. Ca + 2HClO3 Ca(ClO3)2 + H2 (single displacement) Predicting Chemical Reactions 1. CaCO3 CaO + CO2 (rule 1) 2. SO2 + H2O H2SO3 (rule 3) 3. K2O + CO2 K2CO3 (reverse of rule 1) 4. H2O2 H2O + O2 (rule 6) 5. Li2O + H2O LiOH + H2 (rule 4) 6. AlCl3 Al + Cl2 (rule 5) 7. Na + I2 NaI (synthesis) 8. H2CO3 H2O + CO2 (rule 1) 9. KClO3 O2 + KCl (rule 2) 10. MgO + SO3 MgSO4 (synthesis) Stoichiometry Review Questions Answer the following questions. Make sure your final answer has the correct number of significant digits. 1. Determine the number of moles present in 10.0 g of hydrogen gas. [4.95 mol] 2. Determine the mass of nitrogen tetroxide molecules present in 0.20 mol of N2O4. [18 g] 3. Calculate the molar mass of a substance if mol of it weighs 3.2 g. Which of the following substances is it? CO2, SO2, N2O4, Cl2 [64 g/mol; none of them, but Cl2 is the closest at 70 g/mol; oxygen gas has that molar mass) 4. How many molecules of ammonia are present in 1.7 g of ammonia gas? [ molecules of NH3] 5. Calculate the percentage composition of gold chloride (AuCl4) [Au = 58%; Cl = 42%]

5 6. You are asked to measure out 0.50 mol of calcium chloride. How many grams would you need? [mm CaCl2 = g/ mol; m = 55 g] 7. How many sodium ions and chloride ions are there in a g sample of sodium chloride? [ ions of each] 8. Write the equation for the combustion of carbon in oxygen. [C + O2 CO2] 9. What mass of carbon would be necessary to produce g of carbon dioxide? [ g of C] 10. What mass of oxygen would be consumed in this reaction? [ g of O2] 11. Ammonia gas and oxygen react to produce nitrogen dioxide and water. Determine the mass of ammonia that can be consumed in this reaction if there are 20.5 g of pure oxygen available. [4NH3 + 7O2 4NO2 + 6H2O; 6.24 g of NH3] 12. Sodium hydroxide is a strong base and sulfuric acid is a strong acid. a. Write the equation for the neutralization of sodium hydroxide by sulfuric acid. [2NaOH + H2SO4 Na2SO4 + 2H2O] b. What mass of sodium hydroxide can be neutralized by 1.5 mol of sulfuric acid? [ g of NaOH] 13. Silicon carbide is very hard and is used as an abrasive on sand paper. It is made by heating a mixture of carbon and sand in a furnace. SiO2 + 3C SiC + 2CO a. Calculate the mass of SiO2 that will react with 86 g of carbon [ g of SiO2] b. What mass of silicon carbide would be produced if 6.0 g of silicon dioxide was reacted? [4.0 g of SiC]