But There Is Still Work to Do

AP Chemistry

But there is still work to do…

Informational Websites

Below you will find some very helpful website that I advise you to work on in your spare time. These sites are interactive and will provide you with some tutorial support during the summer.

AP Chemistry Wiki Pages (This one will be very helpful. All of the class notes and instructional videos will be found here.)

Dimensional Analysis

Tutorial:

Ionic Nomenclature

Tutorial:

Electron Configuration

Tutorial:

Balancing Chemical

Equations:

Redox Reaction Tutorials:

Scientific Notation, Significant Figures, and the

Factor-Label Method of Solving Problems

Scientific Notation

Scientific notation (Chang, p. 21) is a type of exponential notation in which only one digit is kept to the left of the decimal point. Example: 8.4050 x 10-8.

Significant Figures

It is reasonable that a calculated result can be no more precise than the least precise piece of information that went into the calculation. Thus it is common practice to write numbers in scientific notation with only the last place containing any uncertainty. When we do this we are keeping only the “significant figures” (Chang, pgs 23-24).

To determine the number of significant figures in a number, you read the number from left to right and count all digits starting with the first non-zero digit. Do not count the exponential part.

Thus the number 0.002050 contains 4 significant figures and is written in scientific notation as 2.050 x 10-3. The trailing zeros in a non-decimal number such as 1200 may or may not be significant: the number may be written as 1.2 x 102, 1.20 x 102 or 1.200 x 102 depending on whether it has 2, 3, or 4 significant figures.

Significant Figures in Derived Quantities

When doing calculations, you should use all the digits allowed by your calculator in all intermediate steps. Then in the final step, round off your answer to the appropriate number of “significant figures” such that only the last decimal place contains any uncertainty. You do this by following the rules:

•When adding or subtracting, first express all numbers with the same exponent. Then the number of decimal places in the answer should be equal to the number of decimal places in the number with the fewest decimal places.

•In multiplication or division, the number of significant figures in the answer should be the same as that in the factor with the fewest significant figures.

When using these rules, assume that exact numbers have an infinite number of significant figures (or decimal places). For example, there are exactly 12 inches in one foot.

Solving Problems Using Dimensional Analysis: The Factor-Label Method

Units may be used as a guide in solving problems. First decide what units you need for your answer. Then determine what units you are given in the problem, and what conversion factors will take you from the given units to the desired units. If the units cancel out properly, chances are that you are doing the right thing! The basic set up is

Conversion factors are added until the new units are the same as the units desired. Each conversion factor has a denominator equivalent to the numerator but in different units.

Assignment # 2

1. Carry out the following mathematical operations and express your answer in scientific notation using the proper number of significant figures.

(a) (4.28 x 10-4) + (3.564 x 10-2)

(b) (0.00950) x (8.501 x 107) 3.1425 x 10-11

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2. Use the factor-label method to solve the following problem. Show your work, and give your answer in scientific notation using the proper number of significant figures.

The calorie (1 cal = 4.184 J) is a unit of energy. The burning of a sample of gasoline produces 4.0 x 102 kJ of heat. Convert this energy to calories. (103 J = 1 kJ.)

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3. Use the factor-label method to solve the following problem. Show your work, and give your answer in scientific notation using the proper number of significant figures.

The distance from the sun to the earth is 93 million miles. How many minutes does it take light from the sun to reach earth?

Useful information: 1 km = 0.6214 mile, c = speed of light = 3.00 x 108 m/s)

Atomic Mass, Moles, and the Periodic Table

Atomic Mass and Molar Mass

Isotopic masses cannot be obtained by summing the masses of the elementary particles (neutrons, protons, and electrons) from which the isotope is formed. This process would give masses slightly too large, since mass is lost when the neutrons and protons come together to form the nucleus.

Atomic masses (also called atomic weights) are thus assigned relative to the mass of a particular carbon isotope, , which is assigned the mass of 12 amu exactly. Likewise 1 mole of has a mass of exactly 12 g. Atomic masses and molar masses of other isotopes are calculated based on their mass relative to that of Carbon-12.

Masses of “average” atoms are found by summing isotopic masses, weighting each isotopic mass by its abundance (see Chang, p. 76). Thus one “average” C atom has a mass of 12.01 amu, and the mass of 1 mole of “average” carbon atoms has a mass of 12.01 g. These average masses are what are given on the periodic chart.

What is a Mole?

Since atoms and molecules are so tiny, it is convenient to talk about a large number of them at a time. The chemical counting unit is known as the mole. A mole is defined as the amount of substance that contains as many elementary entities (atoms, molecules, or other particles) as there are atoms in exactly 12 g of the isotope. It has been found experimentally that

This value is known as Avogadro’s number. Just like 1 dozen of anything always contains 12 items, 1 mole of anything always contains 6.022 x 1023 items.

Molecular Masses and Compound Masses

Molecular masses are found by summing atomic masses (see Chang, pp. 81-82). They are often called molecular weights. Thus the mass of 1 mole of water, H2O, would be 2 x (molar mass of H) plus 1x (molar mass of O) or [(2 x 1.008 g) + (1 x 16.00 g)] = 18.02 g.

Ionic compounds such as NaCl do not contain molecules. Their formulas give the relative numbers of each kind of atom in the sample. What we mean by the molar mass (or the molecular weight) of an ionic compound is really the formula weight. The formula weight is the sum of the atomic masses in the formula.

Percent Composition of Compounds

The percent composition by mass is the percent by mass of each element in a compound. If there are n moles of an element per mole of compound, the percent by mass of the element is calculated using the equation,

The sum of the % compositions of all elements in a compound is 100%.

Assignment # 3

Exercises

1. The atomic mass scale gives masses in atomic mass units (amu) relative to the mass of carbon-12.

(a) What is the mass of one 12C atom in atomic mass units (amu)? ______

(b) What is the mass of an average C atom in atomic mass units (amu)? ______

(d) What is the mass of an average Br atom in amu? ______

2. The molar mass scale gives masses in grams (g) relative to the mass of 12C.

(a) What is the mass in grams of 1 mole (mol) of 12C? ______

(b) What is the mass in grams of 1 mole (mol) of carbon? ______

(d) What is the mass in grams of 1 mole (mol) of Na? ______

3. How many 12C atoms are present in a mole of 12C ?

4. Cinnamic alcohol is used mainly in perfumery, particularly in soaps and cosmetics. Its molecular formula is C9H10O.

(a) Calculate the percent composition by mass of C, H, and O in cinnamic alcohol.

(b) How many molecules of cinnamic alcohol are contained in a sample of mass 0.469 g?

Assignment # 4

Naming Inorganic Compounds

To name a compound you must first decide whether the substance is an ionic or molecular compound. Ionic compounds are easily recognized since they usually contain both metallic and non-metallic elements. The most common exception to this rule are ionic compounds containing the ammonium ion, NH4+, such as (NH4)2CO3 or NH4Br which contain no metal ions. Molecular compounds typically contain only non-metallic atoms (and metalloids).

Conventions for naming ionic compounds are given. To successfully follow the rules, however, you must be first learn the names of common ions. Names of ionic compounds do not give the number of each type of ion in the formula: the chemist is supposed to be able to figure that out from his/her knowledge of ion charges and the requirement that salts be neutral (and thus have a sum of zero for the ion charges in the formula).

Binary compounds of the non-metals are named following the guidelines given in Chang on pp. 62-64. Note that when naming these molecular compounds, the number of atoms of a given type is commonly indicated with a prefix (di-, tri-, tetra, etc.).

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Exercises

1. Complete the following chart of corresponding ion names and formulas.

Cation Name / Formula / Anion Name / Formula
(1) potassium ion / (11) nitrate ion
(2) / Fe3+ / (12) / H2PO4-
(3) ammonium ion / (13) hydrogen carbonate (or bicarbonate) ion
(4) / Ba2+ / (14) / MnO4-
(5) silver ion / (15) perchlorate ion
(6) / Cu2+ / (16) / S2-
(7) zinc ion / (17) acetate ion
(8) / Co2+ / (18) dichromate ion
(9) hydrogen ion / (19) / CO32-
(10) chromium(III) ion / (20) sulfite ion

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2. Complete the following chart of corresponding compound names and formulas. Circle the names of all non-ionic (i.e., molecular) compounds.

Compound Name / Formula / Compound Name / Formula
(1) silver nitrate / (11) sodium hydrogen phosphate
(2) / Ni(CH3CO2)2 / (12) / SO3
(3) ammonium sulfate / (13) potassium permanganate
(4) / P2O5 / (14) / Al2S3
(5) sodium oxide / (15) cobalt(III) sulfate
(6) / NH4NO3 / (16) / Ag2CrO4
(7) nitrogen trichloride / (17) / SrF2
(8) / NaHCO3 / (18) sulfur hexafluoride
(9) iron(II) acetate / (19) / NH3
(10) carbon tetrachloride / (20) / LiClO3

Assignment # 5

Empirical and Molecular Formulas

The empirical formula of a compound gives the simplest whole number ratio of different types of atoms in the compound. All salt formulas are empirical formulas. On the other hand, the molecular formula of a compound may or may not be the same as its empirical formula. For example, the molecular formula of butane is C4H10 while its empirical formula is C2H5. The molecular formula gives the true number of each kind of atom in a molecule.

Empirical formulas may be easily determined from experimental data.

Usually you must first determine how many grams of each type of atom are in the compound. If percent composition data is given, assume that you have 100.0 g of the compound; then the number of grams of each element is equal to the percentage for that element.

The next task is convert the grams of each element to moles of the element. Be sure to keep at least three significant figures in your answers.

The final step is to write the molar amounts of each element as subscripts in the formula. Then divide all molar subscripts by the smallest value in the set. At this point, the subscripts may all be very close to whole numbers; if so, you are finished. If one (or more) of the subscripts is not close to a whole number, multiply all molar subscripts by the simple factor which makes all subscripts whole numbers.

Once the empirical formula is determined, the molecular formula is easily found if the molar mass (molecular weight) of the molecule is also known. You first calculate the molar mass of the empirical formula. Then you divide the molar mass of the molecule by the molar mass of the empirical formula. The division should give a simple whole number. That number is the factor by which all subscripts in the empirical formula must be multiplied to obtain the molecular formula.

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Exercises

1. The molecular formula of the antifreeze ethylene glycol is C2H6O2. What is the empirical formula?

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2. A well-known reagent in analytical chemistry, dimethylglyoxime, has the empirical formula C2H4NO. If its molar mass is 116.1 g/mol, what is the molecular formula of the compound?

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3. Nitrogen and oxygen form an extensive series of oxides with the general formula NxOy. One of them is a blue solid that comes apart, reversibly, in the gas phase. It contains 36.84% N. What is the empirical formula of this oxide?

______

4. A sample of indium chloride weighing 0.5000 g is found to contain 0.2404 g of chlorine. What is the empirical formula of the indium compound?

Types of Chemical Reactions

One skill that chemists learn over time is that of writing and balancing equations. The first task is deciding what type of reaction is taking place. In this chapter we study three types:

Precipitation Reactions: In these reactions two soluble salts usually react to form to an insoluble salt (the precipitate!) and a soluble salt. The cations of the reacting salts exchange anions. See Chang, Table 4.2, p. 119 for solubility guidelines.
Acid-Base Reactions: Most commonly an acid of the type HX or H2X reacts with a basic hydroxide to form a salt plus water. Alternatively, the acid may react with ammonia (NH3) to form an ammonium salt (but no water). These are proton transfer reactions in which H+ (the proton) is transferred from the acid to the base.
Oxidation-Reduction Reactions: These are reactions in which one type of atom increases in oxidation number (is oxidized) and another type of atom decreases in oxidation number (is reduced). A large number of oxidation-reduction (redox) reactions contain one or more reactants or products, which are pure elements.

Note that hydroxides can react with acids in acid-base reactions, and also with other salts in precipitation reactions.

Writing Balanced Ionic Equations

The first step in writing a balanced equation is predicting the products of the reaction as discussed above. Then the steps below are completed in sequence:

Balance the Molecular Equation: In the “molecular” equation, nothing is broken up into ions. Salt formulas are written so that the cation charges exactly balance out the anion charges so that the salt is neutral. Then the equation is balanced for atoms.
Balance the Total Ionic Equation: The first step in writing an ionic equation is to decide what species should be broken up into ions. The rules below should help!

Break up into Ions / Do NOT break up! Leave “as is”!

Strong Acids. HCl, HBr, HI, HNO3, HClO4, and H2SO4 are the most common examples; assume other acids are weak.
Strong Bases. NaOH, KOH, or Ba(OH)2 are the most common examples; assume other bases are weak.
Soluble Salts. Salts of the alkali metals, salts containing the NH4+ ion, the NO3- ion, and other salts as specified in Chang, Table 4.2, p. 119.

Weak Acids. Nearly all acids are weak.
Weak Bases. Nearly all bases are weak.
Insoluble Salts. Most salts are insoluble.
Non-electrolytes or Weak Electrolytes. Examples include H2O, gases, pure elements, hydrocarbons, and alcohols.

Balance the Net Ionic Equation: Identify all spectator ions: these are ions that are identical on both sides of the balanced total ionic equation. Remove the spectator ions from the equation. What remains is the net ionic equation. Finally, simplify the stoichiometric coefficients if all of them are divisible by a common factor.

If all the ions are spectator ions so that nothing is left for your net ionic equation, no reaction has taken place!

Net Ionic Equations Fact IQ Sheet

Net Ionic Equation: A balanced chemical equation in which only the ions/compounds involved in the reaction (rxn) are shown. The ions that are notinvolved in the rxn are called the spectator ions.

How To Write Balanced Net Ionic Equations

1) Balance the equation.

2) Break all the aqueous (aq)* compounds up into ions, but do notbreak up polyatomic ions. The charges on these ions are the same as the ones you use when naming the compounds. (e.g. Na2SO4(aq)→ 2Na+(aq) + SO4-2(aq)) Do not break up solids (s) liquids (l) or gases (g). The atoms that make up the compounds in these phaseswill not break up into ions,(e.g. H2O(l), CO2(g), NaI(s)). NOTE: If you are not given the products, you will need to be able to use the solubility rulesto predict the products. Some professors require you learn these rules, while other professors provide them on exams.

* Aqueous means dissolved in water; ionic compounds (metal/non-metal) dissolve in water by breaking up into ions. This has to do with the fact that water is very polar (has a positive and a negative end) and ions have charges. Polar compounds like to mix with other polar or charged compounds, not with non-polar compounds. Example: Oil and water don't mix. This is because water is polar and oil is non-polar.

3) Cancel the ions that are the same on both sides of the equation.

4) Write down what is left.

Example: Write the net ionic equation for the following equation, which, when balanced, might be called the "whole formula" or overall reaction equation: FeCl3(aq) + Na2CO3(aq)→ Fe2(CO3)3(s) + NaCl(aq)

1) Balance to get the overall reaction equation: 2 FeCl3(aq) + 3 Na2CO3(aq)→ Fe2(CO3)3(s) + 6 NaCl(aq)

2) Complete ionic eq.: 2 Fe3+(aq)+ 6 Cl-(aq)+ 6 Na+(aq) + 3 CO32-(aq)→ Fe2(CO3)3(s)+ 6 Cl-(aq)+ 6 Na+(aq)

3) Cancel spectator ions: 2 Fe3+(aq)+ 6 Cl-(aq)+ 6 Na+(aq) + 3 CO32-(aq)→Fe2(CO3)3(s)+ 6 Cl-(aq)+ 6 Na+(aq)

4) Net ionic Equation: 2 Fe3+(aq) + 3 CO32-(aq)→ Fe2(CO3)3(s)

Assignment # 6

Exercises

For each of the following reactions, complete the chart. Be sure to balance all of your equations.

1. Mg(OH)2(s) + HCl(aq)