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Chemistry 30 Introduction and Review Resources: In-class notes and handouts Objectives: 1. Recognize and interpret domestic hazard symbols. 2. Recognize and interpret WHMIS symbols. 3. Recognize various pieces of lab equipment. 4. Identify the location of safety devices. 5. List safety procedures. 6. List the important base and derived units of the metric system (2-2). 7. List the important prefixes in the metric system (2-2). 8. Perform metric conversions (2-3). 9. Use and perform calculations in scientific notation (2-3). 10. Discuss uncertainty (2-3). 11. Compare and contrast between accuracy and precision (2-3). 12. Use significant digits in calculations (2-3). 13. Use both the Stock and Classical systems to name ionic compounds (chapter 7). 14. Use the numbering system to name molecular compounds (chapter 7). 15. Be able to name acids ( chapter 15). 16. Be able to name simple organic compounds (chapter 20). 17. Be able to write and balance chemical equations (chapter 8). 18. Be able to identify types of chemical reactions and complete equations based on activity series (chapter 8). Vocabulary: accuracy acute biohazard carcinogen chronic molecular ionic bond alkyne cyclic monoatomic dependent variable exponent flammable independent variable oxidize ionic alkane aromatic hydrocarbon diatomic precision radioactive toxic uncertainty binary covalent bond alkene aliphatic synthesis polyatomic 1 double replacement single replacement decomposition combustion activity series organic inorganic Lab Safety Hazardous Materials in the Home Poison - skull and crossbones symbol - poisons act by entering the body in some way - poisons can enter the body in one of three ways: a) ingestion (eating) b) inhalation (breathing) c) absorption through the skin - poisons can cause anything from mild illness to death, depending on the nature of the poison Corrosive - Flammable - symbolized by a flaming surface - are substances which can burn easily or cause other materials to burn Explosive - symbolized by an exploding bomb - can cause injury or death as a result a blast or because of the materials expelled by the blast (metal shards) - usually are pressurized aerosol containers which may explode when heated Radiation - symbolized by a 3-sided fan - radioactive materials emit high energy atomic particles or high energy radiation (x-rays, gamma rays), or both - found in smoke detectors and involve no danger if kept at a safe distance - symbolized by a hand eaten away to the bones in a beaker of liquid are chemicals which can act on clothing, skin, eyes or internally by drinking or eating can cause symptoms ranging from mild rash to serious skin damage can damage clothing can cause blindness can cause death if ingested Octagon represents that the contents are dangerous. Triangle represents that the container is dangerous. WHMIS Workplace Hazardous Materials Information System WHMIS is a system of warning symbols and information sheets which detail the danger, safe handling and disposal of a variety of chemical substances in Canada. All chemicals handled in Canada must be labeled using the WHMIS system. There are 8 classes of hazards, each with its own symbol. Toxicity 2 Acute Toxicity - refers to a substance which has immediate effects. If you were to eat breathe, or absorb the substance through your skin, the effects would happen in short order. Effects can include illness, organ damage or death. Chronic Toxicity - refers to the effects of a substance through repeated exposure over a long period (weeks, months or years). Effects may be similar to those of acute toxicity, like those which occur as a result of long term alcohol or cigarette abuse. The effects can also include cancer, allergies or chronic diseases (bronchitis, emphysema, cirrhosis of the liver, etc.) Biohazard - refers to an infectious agent (bacteria, virus or some other organism) which may spread disease if improperly handled. In Case of An Accident: Inhaled Poison Remove the patient to fresh air and apply artificial respiration if necessary. Keep the victim warm with blankets. Contact of Poison with Skin or Eyes Flood affected area with water, for at least 5 minutes. Remove contaminated clothing. DO NOT attempt to use chemical antidote. Swallowed Poison If the person is conscious and able to swallow, immediately dilute the poison by giving the victim 2 to 4 cups of milk or water. Swallowed Corrosives DO NOT INDUCE VOMITING. Give milk and water. If vomiting occurs naturally, hold head below hips to avoid choking. (Note: Corrosives include drain cleaners, toilet bowl cleaners, ammonia, oven cleaners, turpentine, kerosene, furniture polish, gasoline, pine oil and bleaches. Safety in the Chemistry Lab You must know the location of safety equipment in the lab: - fire extinguisher absorbent - eye wash fire blanket - sinks Lab Procedures and Rules 1. No eating or drinking in the lab. 2. Treat all chemicals as if they were hazardous: - never taste chemicals. wash hands after chemicals have been handled. wear eye protection when instructed. 3. Never perform unauthorized experiments. 4. Report all accidents immediately. Do not attempt to clean it up until checking with the teacher. 3 5. If you get a chemical solution in your eye, do not wait for the teacher; go to the eyewash station immediately and wash the eye for at least 5 minutes. 6. If you get chemicals on your clothes, wash the clothes thoroughly. 7. Do not wear loose clothing during a lab. Tie long hair back. 8. Do not sit on the lab bench; you do not know how clean it is. 9. Clean all equipment thoroughly and put it back where it belongs after a lab. 10. Follow directions concerning the safe disposal of chemicals and solutions. 11. Clean your lab station thoroughly after a lab. Metric System SI BASE UNITS Quantity length mass volume temperature time amount of matter electric current Base Unit Symbol metre gram litre kelvin second mole ampere m g L K s mol A SI DERIVED UNITS Quantity Name of Unit Symbol density kilogram per cubic metre kg · m-3 Expressed in Terms of SI Base Units kg · m-3 (kg/m3) force Newton N kg · m · s-2 (kg · m / s2) pressure Pascal Pa N · m-2 heat energy Joule J N·m 4 (N / m2) METRIC PREFIXES This table lists the metric prefixes which are significant: Prefix Symbol Factor by which Unit is Multiplied Exponential Notation 1012 tera T 1 000 000 000 000 giga G 1 000 000 000 mega M kilo k 1 000 103 hecto h 100 102 deca da 10 101 1 100 109 106 1 000 000 THE BASE UNIT deci d 0.1 10-1 centi c 0.01 10-2 milli m 0.001 10-3 micro 0.000 001 10-6 nano n 0.000 000 001 10-9 pico p 0.000 000 000 001 10-12 SCIENTIFIC NOTATION The rules for putting numbers into scientific notation are simple. They work on the assumption that all numbers contain a decimal point. For instance, the number 125 can be visualized as 125.000; 31 can be visualized as 31.000. There are two basic rules: 1) For numbers larger than 1, scientific notation is determined by counting the number of times the decimal place must be moved to the left, leaving just one number to the left of the decimal. The number of times the decimal must be moved is the power of 10 (the exponent). Example: 3000 = 3000.0 = 3 x 103 454 000 = 454 000.0 = 4.54 x 105 3 860 000 = 3 860 000.0 = 3.86 x 106 602 000 000 000 000 000 000 000 = 6.02 x 10 23 5 2) For numbers smaller than 1, scientific notation is determined by counting the number of times the decimal place must be moved to the right, leaving just one number to the left of the decimal. The number of times the decimal must be moved is the exponent, but it is a negative number. 0.068 = 6.8 x 10-2 0.000 049 3 = 4.93 x 10-5 0.000 000 002 41 = 2.41 x 10 -9 Example: If the decimal does not have to be moved, the exponent is zero. For example, the number 1.23 written in scientific notation is 1.23 x 100. IF A NUMBER IS LARGER THAN 9999 OR SMALLER THAN 0.001 IT MUST BE WRITTEN IN SCIENTIFIC NOTATION. Between these extremes you may use either decimal or scientific notation. Metric Conversion Using unitary rates - this method involves multiplication using conversion factors. eg. - 15 mm = m to solve, set up a ratio such that the units of the unknown are the units desired: 15 mm x m = mm m (m/mm is the conversion factor; when the units are cancelled out only 'm' is left) - next fill in the units of the conversion factor. These units are unitary rates; they have a value of 1. In this case 1 m is equal to 1000 mm. You must be very familiar with the value of the prefixes to make this work. 15 mm x - 1 m = 1000 mm m now finish the calculation: (15 x 1 )( mm x m ) = 0.015 m 100 mm Make sure you record the examples in class. Assignment: 1) 16 kg = ? g 2) 632 cm = ? km 3) 2.18 x 105 μN = ? mN 4) 5) 6) 0.036 s = ? ns 7120 Mg = ? Tg 8.88 x 10-10 kL = ? mL 6 UNCERTAINTY This is a fundamental issue in science. A scientist knows nothing for certain; there are no laws in science. A theory can never be proven; it may only be disproven. Every measurement done in science has some amount of uncertainty. For instance, if you take the mass of a substance the scale may read 46.58 g. The last digit is rounded off, so the mass could be as low as 46.575 g and as high as 46.584 g. A more sensitive scale could be used to reduce the uncertainty, but it will always be there. Two terms associated with uncertainty are accuracy and precision. They are defined as follows: Accuracy - how close a measured or calculated value is to a known or real value. Precision - how close many repeated measurements are to each other. Scientists repeat measurements as a way to reduce uncertainty. If a number of measurements are very close to one another, they have good precision and the scientist can be assured that the average of the measurements is likely close to the actual value (accurate) SIGNIFICANT DIGITS Refers to the certain digits in any measurement. This is related to issues of uncertainty; the results of a calculation can be no more precise than the least precise measurement which goes into it. To determine the number of significant digits you must be able to handle zeros and their relationship to the non-zero digits in a number. Note the following rules: 1. All non-zero numbers are considered significant; that is, they are counted 123 has 3 significant digits; 1267 has 4 s.d. 2. There are two situations where zeros are significant: i) Zeros between two non-zero numbers are considered significant. 102 has 3 s.d.; 10203 has 5 s.d.; 1002 has 4 s.d. ii) A zero at the end of a decimal number is significant. 12.00 has 4 s.d.; 0.010 has 2 s.d.; 1200.000 has 7 decimal places Note: in the last example, all the zeros are significant because they are between s.d. 3. In any other situation zeros are not considered significant: i) For a number larger than 1, a zero between the decimal and the first non-zero number is not significant. 120 has 2 s.d.; 10200 has 3 s.d.; 130 000 000 has 2 s.d. ii). For a number smaller than 1, a zero between the decimal and the first non-zero number is not significant. 0.0012 has 2 s.d.; 0.02102 has 4 s.d.; 0.000 000 001 has 1 s.d. 7 Exact numbers - are numbers that are defined (conversion factors in the metric system) or numbers which result from counting objects (like the coefficients used to balance chemical equations). Exact numbers are said to have an infinite number of significant digits. Rounding off - is necessary when a number from a calculation must have the number of significant digits reduced. The rules for rounding are as follows: o if the digit following the last digit to be kept is greater than 5, the last digit is increased by 1 e.g. o 123.5 123.44 rounded to 4 s.d. is now 123.4 if the digit following the last digit to be kept is equal to 5, followed by a nonzero digit, the last digit is increased by 1 e.g. o rounded to 4 s.d. is now if the digit following the last digit to be kept is less than 5, the last digit stays the same e.g. o 123.46 123.452 rounded to 4 s.d. is now 123.5 if the digit following the last digit to be kept is equal to 5, and not followed by a nonzero digit , the last digit is increased by 1 only if it produces an even number 123.45 123.55 rounded to 4 s.d. is now rounded to 4 s.d. is now 123.4 123.6 Rule for addition and subtraction Add or subtract and then round-off so that the answer is no more precise than the least precise number in the calculation.. Rule for multiplication and division Multiply or divide and then round-off so that the answer has no more significant digits than the number with the fewest significant digits in the calculation. Remember that any exact numbers do not enter into the determination of least significant digits. For long, multi-step calculations: Do not round off the number at each step in your calculator; keep the entire number, with all its decimal places and use it in the next step. The danger is that you introduce ROUNDING ERROR. Record the examples in class. 8 Multiplication in Scientific Notation The rule is that the integers (the first number in each scientific notation) in each number are multiplied, with the resulting number becoming the integer in the answer. The exponents (the powers of 10) are added; the answer becomes the exponent in the answer. For example: (3.4 x 102)(2.0 x 103) Problem: Solution: 1. Multiply the integers 3.4 x 2.0 = 6.8 2. Add the exponents 102 x 103 = 102+3 = 105 3. Combine 6.8 x 105 If the integer in the solution has more than one digit to the left of the decimal point the scientific notation must be corrected: (6.0 x 103)(2.5 x 107) 15.0 x 1010 = 1.50 x 1011 Problem: Solution: Division in Scientific Notation Division in s/n is similar to multiplication, except that in the case of division the integers are divided, and the exponent of the second number is subtracted from the first. For example: Problem: Solution: (3.4 x 102) / (2.0 x 103) 1. Divide the integers 3.4 / 2.0 = 1.7 2. Subtract the exponents 102 / 103 = 102-3 = 10-1 3. Combine 1.7 x 10-1 If the integer in the solution has less than one digit to the left of the decimal point the scientific notation must be corrected: Problem: Solution: (6.0 x 103)(8.0 x 107) 0.75 x 10-4 = 7.50 x 10-5 9 Addition and Subtraction in Scientific Notation These two operations are slightly different from multiplication and division. Numbers that are expressed in s/n can only be added or subtracted if the exponents are the same. If the exponents are the same the integers are added or subtracted and the exponent is carried into the solution. For example: Addition: (2.07 x 106) + (1.30 x 106) = 3.37 x 106 Subtraction: (2.07 x 106) - (1.30 x 106) = 0.77 x 106 = 7.70 x 105 If the exponents are different there are two methods which may be followed to solve the problems: 1) Move the decimal of one number to change the exponent. Example: = = 2) (2.75 x 103) + (3.20 x 102) (2.75 x 103) + (0.32 x 103) 3.07 x 103 Convert s/n numbers back to normal notation. Add or subtract as required, then convert the answer back to s/n format. Example: = = = (5.75 x 104) - (2.37 x 103) 57 500 - 2 370 55 130 5.51 x 104 NOMENCLATURE There are two types of compounds; inorganic and organic. There are two types of inorganic compounds; molecular and ionic. Molecular compounds are bound together almost exclusively with covalent bonds and are compounds composed of non-metallic elements. Ionic compounds are bound together by ionic bonds and are compounds composed of metallic and non-metallic ions. Organic substances are comprised almost exclusively of hydrogen and carbon (hydrocarbons). Further, most compounds we will deal with are binary compounds, compounds made up of two parts. Thus when naming a compound one must consider the two parts involved, as well as whether the compound is molecular or ionic. The naming system used hinges on whether the compound is molecular or ionic. Finally, naming is done using two principles: 1) 2) The name must be as simple as possible; nothing unnecessary is added to a name. The name must be unique to a substance; two different structures cannot have the same name. Elements Are substances made up of one kind of atom. Most elements on the Periodic Table are monoatomic. Eight elements are diatomic (H2, N2, O2, F2, Cl2, Br2, I2 and At2). Two elements are polyatomic (S8, P4). 10 Common Substances These are compounds that are know by names other than their systematic names. You must be familiar with their formulas, systematic anmes and common names. Here is the list: Formula Chemical Name Common Name H2 O hydrogen oxide water NaCl sodium chloride table salt HCl hydrogen chloride hydrochloric acid HNO3 hydrogen nitrate nitric acid H2SO4 hydrogen sulfate sulfuric acid H3PO4 hydrogen phosphate phosphoric acid CH3COOH acetic acid vinegar CaSO4 calcium sulfate gypsum (dry wall) NH3 nitrogen trihydride ammonia H2 O2 hydrogen peroxide hydrogen peroxide C2H5OH ethanol drinking alcohol CH3OH methanol wood alcohol CH4 methane natural gas O3 ozone ozone C12H22O11 sucrose table sugar KCl potassium chloride potash NaOH sodium hydroxide lye, caustic soda CaO calcium oxide lime If you are asked to name one of these formulas, give the name written here in bold. 11 Naming of ionic compounds Ionic compounds can be named using either of two similar methods. In order to name ionic compounds one must first understand how ionic compounds come to be. Ionic compounds are formed as a result of cations and anions which are attracted to each other because of their opposite charge and join in ratios which result in a net charge of zero. Thus Na 1+ and Cl1- form the ionic molecule NaCl. Ca2+ and F1- form the ionic molecule CaF2. Mg2+ and O2- form the ionic molecule MgO. Naming of ionic compounds is a way of clearly stating their chemical formula. The name which a compound receives gives enough information to determine exactly its chemical formula. The method is simple; the name of the cation comes first, the anion is second. Tables 4 and 5 of the package of tables gives the name of common cations and anions, including polyatomic cations and anions. When giving a chemical formula, it is important to identify both the ions involved and their charge. Knowing the charge of the ions is very important. For instance, in the compound FeCl 2 the iron is iron(II), or ferrous. Thus the compound name is either iron(II) chloride or ferrous chloride. There is a difference between iron(II) chloride and iron(III) chloride. You are required to use the Stock system in this class, but you should be able to recognize Classical nomenclature when it comes around. Stock Method - Classical Method - for naming ionic compounds (metal + non-metal) uses the oxidation number to identify the charge on the cation (iron II, iron III) for those cations which has more than one possible charge for naming ionic compounds uses archaic names to identify the charge on the cation (ferrous = Fe2+, ferric = Fe3+) To use either method for naming compounds: 1. Identify the cation and the anion. eg. 2. Find the name of the cation and the anion. eg. 3. Na3PO4 = Na1+ , PO43- Na1+ = sodium, PO43- = phosphate Put the two together. eg. sodium phosphate To use either method to determine the formula of a compound from the given name: 1. Use the tables to find out the formulas for the respective ions. eg. ammonium nitrate = NH41+ , NO31- 2. Combine the ions in a ratio which will result in a net charge of zero. eg. 1+, 1-; combine in a 1:1 ratio; NH4NO3 12 Naming of molecular compounds Molecular compounds use a prefix system of naming. Prefixes are used which indicate the number of atoms of each element present in the compound. The prefixes are as follows: mono = 1 di = 2 tri = 3 tetra = 4 penta = 5 hexa = 6 hepta = 7 octa = 8 nona = 9 deca = 10 Thus the compound N2O4 is named dinitrogen tetraoxide. The only exception to the numbering is when the first element in the compound is present as a single atom, such as CO2. The prefix "mono" is omitted and the compound name is carbon dioxide. This exception is more visible in CO, or carbon monoxide. "Mono" is omitted on the carbon, but appears on the oxygen. To use the method for naming compounds: 1. Determine the identity of the elements and the number of each element. eg. 2. Find the appropriate prefixes. eg. 3. N2O5 = 2 nitrogen, 5 oxygen 2 = di-, 5 = penta- Put the name together with the appropriate prefixes. eg. dinitrogen pentaoxide To use the method to determine the formula of the compound from the name: 1. Identify the prefixes and their meaning. eg. 2. tetraphosphorus decaoxide = 4 P, 10 O Put the formula together. eg. P4O10 Waters of hydration Some molecules exist in nature in relationship with water. Water just seems to be attracted to the molecules and it is possible to determine the number of water molecules relative to the number of molecules of the substance. These water molecules are called waters of hydration. For instance CaSO 4 · 2 H2O is the formula for gypsum, used in dry wall. There are two waters of hydration for every gypsum molecule. The proper name for this molecule is calcium sulfate dihydrate. The number of water molecules is indicated using the numbering system from molecular compounds. 13 Organic Nomenclature Organic chemistry refers primarily to a class of compounds called hydrocarbons. These compounds contain two elements, hydrogen and carbon. Other elements, such as oxygen, nitrogen, sulfur, can also be included, but these are minor constituents, if they are present at all. The naming of organic compounds is based on three considerations: 1. 1. Length of the carbon chain 2. Type of bonds which join the carbons 3. Functional groups attached to the carbon chain Length of the carbon chain When naming a carbon compound it is necessary to find the longest, continuous carbon chain in the compound. Part of the name comes as a result of a prefix which indicates the number of carbons in the chain. Following is a list of the prefixes: 2. Number of carbons Prefix 1 2 3 4 5 6 7 8 9 10 methethpropbutpenthexheptoctnondec- Type of bonds which join the carbons. There are three main groups of hydrocarbons, based on the type of bonds which join the carbons: a) Alkanes - the carbon chain is joined exclusively by single bonds. - an alkane is indicated in the name by the suffix -ane. - the general formula for an alkane is CnH2n+2 b) Alkenes - the carbon chain contains at least one double bond. - an alkene is indicated in the name by the suffix -ene. - the general formula for an alkene is CnH2n c) Alkynes - the carbon chain contains at least one triple bond. - an alkyne is indicated in the name by the suffix -yne. - the general formula for an alkyne is CnH2n-2 14 In the case of alkenes and alkynes, it is necessary to indicate where on the carbon chain the double or triple bond is located. This is done by numbering the carbons on the chain, starting from the end closest to the double or triple bond. 3. Functional groups attached to the carbon chain A functional group is any atom or group of atoms other than hydrogen attached to the main carbon chain. Functional groups are identified according to their content, as well as their placement on the carbon chain. Like alkenes and alkynes, it is necessary to indicate where on the carbon chain the functional group is located. This is done by numbering the carbons on the chain, starting from the end closest to the functional group. In a case when there is both a double or triple bond and a functional group, the numbering of the carbon chain is done in favor of the bond, rather than the group. There are three main types of functional groups it is necessary to recognize: a) Hydrocarbon chains If a chain of carbons is off the main chain, the presence of the functional group is named using the prefixes from 1. above, followed by the suffix -yl remembering also to indicated where the functional group is located. For instance, a butane with a two-carbon chain on the number-2-carbon is called 2-ethylbutane. b) Halogens If one of the halogens is attached to the carbon chain, their name is changed slightly: Fluorine Chlorine Bromine Iodine c) becomes becomes becomes becomes FluoroChloroBromoIodo- Hydroxyl groups If an hydroxyl (-OH) group is attached to the carbon chain, the hydrocarbon becomes a new class of organic compounds, called alcohols. The location of the hydroxyl group is indicated by carbon number at the beginning of the name, and the extra suffix -ol is added at the end of the name. For instance, ethane with an hydroxyl or alcohol group attached is now called ethanol. Note: Numbering of the location of special features like multiple bonds or functional groups is not necessary in cases where there is no question where the bond or functional group must be. This applies mainly to small molecules of 1, 2, or 3 carbons. 15 Nomenclature Flowchart Is the substance an element? Yes Name as an element Yes Name as a common substance Yes It is a molecular compound. Use the numbering system Yes It is an ionic compound. Use Stock system No Is the substance a common substance? No Does the substance contain non-metals only? No Does the substance contain a metal and a non-metal? No It is likely an organic compound. Use organic nomenclature 16 Chemical Formulas and Equations A chemical equation is a form of shorthand which gives an outline of the progress of a chemical reaction: H2O → H2 + O2 REACTANT PRODUCT One very important principle of chemistry affects chemical equations. That principle is the law of conservation of mass. Very simply, it states that matter is neither gained nor lost in a chemical reaction. In a chemical reaction you end up with the same number and type of each atom that you started with. This applies to mass and to atoms. This principle must be reflected in chemical equations. Balancing equations is really very simple. It is similar to the method used to determine the chemical formulas for ionic compounds, as discussed in Unit 3. For instance, using the equation for the decomposition of water, first tally up the number of each kind of atom on each side of the reaction equation: 1 H2O → H=2 O=1 1 H2 + 1 O2 H=2 O=2 Once this is done, the method is very straightforward. The idea is to have the same number of atoms on each side of the reaction equation. To do this, it is necessary to determine if there are fewer atoms on one side of equation, compared with the other side. For instance, in this case the product side has one more oxygen than the reactant side. In order to balance the equation for oxygen, the number of oxygen atoms on the reactant side must be increased; the only way to do that is to increase the number of water molecules on the reactant side: 2 H2O → H=4 O=2 1 H2 + 1 O2 H=2 O=2 As you can see, increasing the number of water molecules balances the equation in terms of oxygen atoms, but now the reactant side has too many hydrogen atoms; we must now balance the equation in terms of hydrogen. The only way to do this is to increase the number of hydrogen atoms on the product side, by increasing the number of hydrogen molecules: 2 H2O H=4 O=2 → 2 H2 + 1 O2 H=4 O=2 Now the equation is balanced. The total number of each type of atom on the reactant side is equal to that on the product side, and the law of conservation of mass is satisfied. Now for another example we use a reaction between lead nitrate (Pb(NO 3)2) and sodium iodide (NaI), both of which produce clear solutions. When these solution react a solid yellow compound is produced called lead iodide (PbI2), as well as another clear solution which contains sodium nitrate (NaNO 3). The reaction equation follows: Pb(NO3)2 + NaI → PbI2 + NaNO3 17 Step 1: Tally up the number of each atom: Pb(NO3)2 + NaI → PbI2 + NaNO3 Pb = 1 Na = 1 I=1 N=2 O=6 Pb = 1 Na = 1 I=2 N=1 O=3 Step 2: Begin balancing for each atom: This equation is already balanced for lead and sodium; the next on the list is iodine. Pb(NO3)2 + 2 NaI → PbI2 + NaNO3 Pb = 1 Na = 2 I=2 N=2 O=6 Pb = 1 Na = 1 I=2 N=1 O=3 Note that the only way to increase the number of iodine atoms on the reactant side is to add another NaI molecule, thus increasing both the number of sodium and iodine atoms. This causes the number of sodium atoms to become unbalanced, so the number of sodium atoms on the product side must be increased, by adding another sodium nitrate molecule: Pb(NO3)2 + 2 NaI → PbI2 + 2 NaNO3 Pb = 1 Na = 2 I=2 N=2 O=6 Pb = 1 Na = 2 I=2 N=2 O=6 Other information can also be gained or given from a reaction. It is sometimes useful to know the state a compound is in for a chemical reaction. By "state" we mean whether the compound is a liquid, solid, or a gas. This information is given by the use of subscripts following each compound. The list of subscripts is as follows: (s) (l) (g) (aq) (ppt) - solid liquid gas aqueous (meaning that the compound is dissolved in water.) precipitate (meaning that the reaction produces a solid which falls out of solution.) Using the two previous examples, these subscripts can be quite valuable: 2 H2O(l) 2 H2(g) + 1 O2(g) Pb(NO3)2(aq) + 2 NaI(aq) PbI2(ppt) + 2 NaNO3(aq) 18 Types of Chemical Reactions There are five basic types of chemical reactions: 1) Synthesis - the combination of two or more substances to form a compound. A + B → AB another way to look at it is: element + element compound 2) Decomposition - One substance breaks down to form two or more simpler substances. AB → A + B another way to look at it is: compound element + element 3) Combustion - involves the burning of a hydrocarbon in the presence of oxygen to form carbon dioxide and water. CxHy + O2 → CO2 + H2O 4) Single Replacement - reactions occur when one element is replaced by another in a compound A + BC → AC + B another way to look at it is: element + compound element + compound 5) Double Replacement - reactions occur when the elements in a solution of reacting compounds exchange places, or replace each other. AB + XY → AY + XB another way to look at it is: compound + compound compound + compound . 19 20 Chemistry 30 - Significant Digits 1. 2. For the following count the significant digits: a) 18.56 m i) 1500 C b) 0.5306 kg j) 0.0062 L c) 0.0128 km k) 2.300 kPa d) 20 apples l) 8.0 J e) 1.03 x 104 N m) 15 000 000 A f) 406.010 mol n) 120 mm g) 0.00920 g o) 500 students h) 90 502 cm p) 100 000 t Round off the following numbers to the given number of significant digits: a) 6.249 mm, 2 s.d. b) 10.98 g, 3 s.d. c) 0.0573 mol, 2 s.d. d) 69.95 km/h, 2 s.d. e) 298.036 cm3, 4 s.d. f) 349.9 A, 3 s.d. g) 9.100 g, 2 s.d. h) 56087250 N, 4 s.d. i) 21.35 m, 3 s.d. j) 450.5 kL, 3 s.d. k) 67.77 mg, 1 s.d. l) 2880 L, 4 s.d. m) 675 J, 2 s.d. 20 21 3. Perform the indicated operations. Give the answer to the correct number of significant digits: a) Add c) Add 8.55 mL 11.6 mL 20.0 mL 480 km 24.07 km d) Subtract 136 g 3.49 g e) Subtract 16.56 mL 6.3 mL f) Subtract 51.08 mol 9.9 mol g) i) k) 4. b) Add 9.54 g 6.578 g 10.02 g 18.4 g/mL x 5.5 mL = 1.0058 t x 1000 kg = 1t 358.6 g 2.02 mol = h) 21.4 g x j) 6.0 g 24.3 g/mol l) 2.64 g = 5.38 g/mL 1 kg 1000 g = = Perform the following calculations. Give your answers to the proper number of significant digits. If the numbers are in scientific notation, give the answer in scientific notation. a) 9.25 + 4.10 - 2.05 = b) 134.8 + 2.05 - 13 = c) 14.896 - 2.42 + 4.60 = d) (3.45 x 10-1) - (4.789 x 10-3) = e) (7.95 x 10-2) + (2.05 x 10-1) = f) 4.18 x 0.051 960 = g) 0.50 ÷ 4.12 = h) (93.30 x 10-2) x (4.612 x 101) = i) (1.981 x 101) ÷ (2.5 x 10-2) = j) ((4.68 x 10-4) x (8.743 x 105)) ÷ (1.04 x 10-2) = 21 22 Chemistry 30 Using Ion Charge to Predict Formulas Determine the most likely chemical formulas for the following combinations of ions: a) K1+, Br1- k) Hg2+, O2- b) Ca2+, Cl1- l) K1+ , PO43- c) Li1+, H1- m) Pb4+, O2- d) Fe3+, OH1- n) NH41+ , SO42- e) Ca2+, OH1- o) Na1+, PO43- f) Zn2+, O2- p) Mn7+, O2- g) Mg2+, NO31- q) Na1+, SeO32- h) Fe2+, O2- r) Na1+, SeO42- i) Fe3+, O2- s) Al3+, SO42- j) Sn4+, F1- t) H1+, S2- 22 23 Chemistry 30 - Nomenclature Name the following substances: 1) NH3 26) CO2 2) Fe(NO3)2 27) SO3 3) Fe(NO3)3 28) PbO 4) CuBr2 29) BaS 5) SnCl4 30) PbO2 6) CuSO3 31) SO2 7) CuSO4 32) BaH2 8) C2H5OH 33) CaSO44 H2O 9) Li2CrO4 34) O2 10) CuBr 35) Al2O3 11) CuCr2O7 36) PCl3 12) NH4NO3 37) N2 13) NH4NO2 38) Fe2O3 14) SnCO3 39) CrCl3 15) Sn(CO3)2 40) Pb(NO3)2 16) Na2SO4 41) Al2(SO4)3 17) H2 42) Cu(NO3)2 18) MgSO3 43) (NH4)3PO4 19) KH 44) Mg(OH)2 20) MgBr2 45) Ca(HCO3)2 21) CO 46) H2S 22) Li2O 47) CaH2 23) P2O56 H2O 48) O3 24) I2 49) XeF8 25) CH4 50) S8 23 24 Write molecular formulas for each of the following: 1) sodium fluoride 26) calcium carbonate 2) potassium carbonate 27) ammonium sulfite 3) aluminum sulfide 28) iron (II) hydroxide 4) calcium bromide 29) copper (II) nitrate 5) aluminum chloride 30) oxygen 6) silver oxide 31) lithium dichromate 7) ammonium sulfide 32) hydrogen nitrate 8) barium hydroxide 33) barium bicarbonate 9) iron (III) iodide 34) nitrogen dioxide 10) sodium sulfate 35) carbon monoxide 11) tin (II) nitrate 36) ammonium carbonate 12) potassium sulfite 37) ammonium dichromate 13) magnesium sulfate 38) aluminum hydroxide 14) lithium phosphate 39) tin (II) nitrate 15) bromine 40) tin (IV) carbonate 16) plumbous oxide 41) mercurous chloride 17) hydrogen 42) mercuric chloride 18) sulfur dioxide 43) zinc perchlorate 19) methanol 44) chromium (III) stearate 20) carbon dioxide 45) chromium (II) acetate 21) sodium chromate 46) barium nitrite dihydrate 22) sulfuric acid 47) iron (II) phosphate 23) sodium hydride 48) iron (III) chlorate 24) stannous chloride 49) cuprous bromide 25) hydrogen selenide 50) ammonium oxalate 24