Review for Final Sheet - AP:

Save this PDF as:
 WORD  PNG  TXT  JPG

Size: px
Start display at page:

Download "Review for Final Sheet - AP:"

Transcription

1 Answer these questions on a separate sheet of paper: Unit 1 Topics Lab skills Reading a graph Percent error What is chemistry? Classifying matter picture Phases of matter picture Phase change diagram Gases what s in between the molecules? Heating curve Cooling curve Physical vs. chemical Law of conversation of matter graph Accuracy and precision Significant figures Scientific notation Unit conversions Metric conversions Density Density graph Nomenclature Review for Final Sheet - AP: Questions 1. What is chemistry the study of? a. Chemistry is the study of matter b. Matter is anything with mass and volume c. Chemistry is the study of anything with mass and volume 2. What is matter? What are some examples of matter? a. Matter is anything with mass and volume b. Chemistry is the study of anything with mass and volume 3. What is the scientific process? What are the different steps of the scientific process? a. The scientific method is a body of techniques for investigating phenomena, acquiring new knowledge, or correcting and integrating previous knowledge. To be termed scientific, a method of inquiry is commonly based on empirical or measurable evidence subject to specific principles of reasoning. i. Ask a Question ii. Do Background Research iii. Construct a Hypothesis iv. Test Your Hypothesis by Doing an Experiment v. Analyze Your Data and Draw a Conclusion vi. Communicate Your Results 4. What are the different types of observations that you can gather in an experiment? a. Qualitative observations involved with qualities (no numbers) b. Quantitative observations involved with quantities (numbers) 5. Explain what types of lab equipment is used when? a. Measurement biuret, graduated cylinder b. Holding liquids beakers c. Measuring temperature thermometers d. Measuring mass balance 6. What is a physical change? What are some examples of physical changes? a. Physical can be observed without changing the chemical composition (color change by using your pencil on your notebook paper); changes the appearance of the item, but it is still the same substance before and after the change

2 7. What is a chemical change? What are some examples of chemical changes? a. Chemical changes the chemical composition of the substance (ie. fire); what you end with is different than what you start with 8. What are the three states of matter? Give characteristics of each state? a. Solid, liquid, gases b. Particles in a: i. gas are well separated with no regular arrangement. ii. liquid are close together with no regular arrangement. iii. solid are tightly packed, usually in a regular pattern. c. Particles in a: i. gas vibrate and move freely at high speeds. ii. liquid vibrate, move about, and slide past each other. iii. solid vibrate (jiggle) but generally do not move from place to place. Liquids and solids are often referred to as condensed phases because the particles are very close together. The following table summarizes properties of gases, liquids, and solids and identifies the microscopic behavior responsible for each property Some Characteristics of Gases, Liquids and Solids and the Microscopic Explanation for the Behavior gas liquid solid assumes the shape and volume of its container particles can move past one another compressible lots of free space between particles flows easily particles can move past one another assumes the shape of the part of the container which it occupies particles can move/slide past one another not easily compressible little free space between particles flows easily particles can move/slide past one another retains a fixed volume and shape rigid - particles locked into place not easily compressible little free space between particles does not flow easily rigid - particles cannot move/slide past one another 9. Explain why a phase change is a physical change, not a chemical change. a. The chemical composition has not been changed, just the distance between the molecules 10. What is a pure substance? Give an example. a. Pure substance = compound or element b. A pure substance is just one type of molecule or compound c. oxygen 11. What is an element? Give an example. a. An element is just one type of atom b. Carbon 12. What is a mixture? Give an example. a. A mixture is made up of two or more different molecules/atoms that are not chemically bonded b. An alloy 13. List the different ways that mixtures can be separated. a. Magnetism, chromatography, distillation, centrifuge, decanting, filtration 14. What is a homogeneous mixture? Give an example. a. Homogenous the same throughout; two substances that are mixed together in a uniform manner 15. What is a heterogeneous mixture? Give an example. a. Heterogeneous different throughout; two substances that are mixed together but not in a uniform manner 16. What are the different units of the metric system? You will need to memorize the conversion factors. a ml = 1 L b. 100 cm = 1 m c g = 1 kg

3 17. Explain how to do factor label conversion problems. Example: convert 5 m to cm. 5 m x 100 cm / 1 m = 500 cm 18. How do you convert from 100 ml of water to L of water? 100 ml x 1 L/1000 ml = 0.1 L 19. How do you convert from 50 meters to kilometers? 50 m x 1 km/1000 m = 0.05 km 20. How do you convert from molecules to moles? Provide an example. a. 1 mol = x molecules (use this as a conversion factor) b. 5 mol x 6.02x10 23 molecules/1 mol = x molecules 21. What are sig. figs? What are the rules for sig. figs? What are the rounding rules for adding and subtracting? Multiplying and dividing? a. Sig figs are a universal way of evaluating the precision and accuracy in a measurement 1. All non zero numbers are significant (meaning they count as sig figs) 613 has three sig figs has six sig figs 2. Zeros located between non-zero digits are significant (they count) 5004 has four sig figs 602 has three sig figs has 16 sig figs! 3. Trailing zeros (those at the end) are significant only if the number contains a decimal point; otherwise they are insignificant (they don t count) has four sig figs has six sig figs has two sig figs unless you re given additional information in the problem 4. Zeros to left of the first nonzero digit are insignificant (they don t count); they are only placeholders! 0.52 three sig figs has two sig figs also has two sig figs! a. What are the rules for adding/subtracting? i. Round to the number with the lower number of decimal places b. What are the rules for multiplying/dividing? ii. Round to the number with the lower number of sig. figs. 22. What is precision and accuracy? Identify if a set of numbers is precise or accurate. a. Accuracy tells me how correct you were; how close were you to the number you were supposed to get; getting 100% on a test b. Precision requires multiple trials how close together several trials are; getting an 80% on every test in the class 23. Which set of student s data is more precise: a , 10.24, 10.23, b. 10, 11, 12, 9 i. A 24. How do you convert from degrees Celsius to degrees Kelvin? Convert 24 o C to K. a. Add to C b = K

4 Unit 2/7 Topics Using the Periodic table Regions of the periodic table Patterns of the periodic table History of the atom Properties of metals, nonmetals, metalloids Subatomic particles Average atomic mass Picture ion Isotope notation Picture isotope Atomic diagrams Quantum theory Electron configuration Picture of s, p, d, f Noble gases full s and p electron clouds Orbital diagrams Valence electrons Lewis dot diagrams Periodic trends Nuclear chemistry Questions 1. The elements on the periodic table are arranged according to what? a. Atomic number (number of protons) 2. Define group. How many groups are there on the periodic table? a. The columns on the periodic table b Define period. How many periods are there on the periodic table? a. The rows on the periodic table b Where are the metals located? a. Left of the stair step 5. Where are the nonmetals located? a. Right of the stair step 6. Which group is called the alkali metals? The alkaline earth metals? The noble gases? The Halogens? What are the charges for each of these groups? a. 1, 2, 18, 17 b. +1, +2, 0, Where are the transition metals located? a. Groups 3-12 and under the stair step 8. What are the three subatomic particles? What are their charges? a. Protons, neutrons, electrons b. +1, 0, What is an ion? An isotope? What subatomic particle is different with each of these? a. An ion is an element with a charge, positive (+) or negative ( ) i. Ions are formed when the number of electrons are changed ii. Have same number of protons, but different number of electrons iii. This means that they have a charge ( a positive or negative sign and number after the element symbol) iv. electron b. Isotopes - Atoms that have the same number of protons, but the number of neutrons have changed i. These are what are commonly used in nuclear chemistry ii. You can tell you have an isotope because the mass number you are using is different than what is listed on the periodic table or because the neutrons are different than what it should be from the periodic table iii. neutron

5 10. When looking at the periodic table, which number is the atomic mass number and which is the atomic number? Is it the atomic number or the mass number that determines which element you are looking at? a. Atomic mass number decimal b. Atomic number whole number c. Atomic number 11. If you have an element that has 5 protons, 5 electrons, and 6 neutrons, what is its atomic number? What is its atomic mass number? Which element is it? a. Atomic number = 5 b. Atomic mass = 11 c. boron 12. Remember that elements in the same group are similar. So, which elements would be similar to S? a. The rest of column How many energy levels (rings) surround an atom of Mg? Be? Cs? a. 3 b. 2 c Be able to draw a Bohr diagram. Do one for P. 15. What are the four quantum numbers? What letters are they represented by? Know what numeric values the four quantum numbers can have. a. There are a total of four quantum numbers: the principal quantum number (n), the orbital angular momentum quantum number (l), the magnetic quantum number (m l), and the electron spin quantum number (m s). b. n = row number (1-7) c. l = shape (0 for s, 1 for p, 2 for d, 3 for f) d. ml = + - l e. ms = spin (+1/2 or ½) 16. Which orbital is sphere shaped? S, p, f, d? Which is dumbbell shaped? a. S b. p 17. How many electrons can exhibit a single orbital? (Maximum number) a Write the electron configuration for Mg. For S. For Ar. a. Mg = 1s 2 2s 2 2p 6 3s 2 b. S = 1s 2 2s 2 2p 6 c. Ar = 1s 2 2s 2 2p 6 3s 2 3p How many electrons can a single orbital hold? a How many electrons can energy level 2 hold? a How many electrons can energy level 3 hold? a How many electrons can energy level 4 hold? 5? 6? 7? a. 32 b. 18

6 c. 18 d Which energy level has the highest amount of energy? Which one has the least? a. 7 b How many valence electrons does Al have? V? O? a. 3 b. 2 c How many dots are in the Lewis Dot diagram for Ge? N? a. Ge with 4 dots. Ge is in group 14 b. N with 5 dots. N is in group What happens to atomic radius as I move from left to right across the periodic table? a. decreases 27. What happens to atomic radius as I move from top to bottom down the periodic table? a. increases 28. What is the difference between a continuous, emission, and absorption spectra? 29. Be able to calculate wavelength and frequency. Explain how to solve for each of these. a. λν = c b. λ = c/ν c. ν = c/λ 30. How do you solve for the energy of a photon? If a laser produces light at 695 nm, what is the energy of the photon? a. E = hv b. 695 nm x 1m/1 x 10 9 nm = 6.95 x 10-7 m c. ν = c/λ = x 10 8 m/s / 6.95 x10-7 m = 4.3 x10 14 s -1 d. E = (6.626 x10-34 Js)( 4.3 x10 14 s -1 ) = 2.8 x J

7 Unit 3 Topics Organic vs. inorganic Types of bonding (ionic, covalent, metal, polyatomic) Ionic compound/bonding picture Covalent compound/bonding picture Metallic bonding - picture Properties of ionic, covalent, and polyatomic compounds Electrical attraction between ions in an ionic compound Ionic compound criss-cross Naming Elements Diatomics Ions Ionic compounds Transition metal ionic compounds Polyatomic compounds Covalent compounds Acids Hydrates The mole Molar mass Percent composition Empirical formulas Molecular formulas Balancing equations Law of conservation of matter balancing equations Balancing equations particle representation Evidence of chemical change Writing word equations Types of reactions Predicting products Activity series predicting products of chemical reactions Stoichiometry BCA table Mole to mole Mole to mass Mass to mole Mass to mass Volume Molecules Using density Stoichiometry with particle representations Percent yield Limiting vs. excess reaction with pictures Questions: 1. What is a compound? Explain why elements combine. Explain how many valence electrons the elements want to have. a. Two or more different atoms bonded together b. To reach a noble gas configuration c What are the three types of bonding? a. Ionic Bonding (Metal + Nonmetal) b. Covalent Bonding (Nonmetal + Nonmetal) c. Polyatomic Bonding (Both Ionic and Covalent Characteristics > 3 or more elements including metal and nonmetals) d. Metallic Bonding (One type of metal atom)

8 3. What are the charges of N? O? B? Li? How do you determine the charges of elements? a. -3 b. -2 c. +3 d. -1 e. Location on periodic table 4. Define ionic compounds. Be able to write ionic compounds using the criss cross method. Be able to write the formulas for ionic compounds or write the names. a. Bond between a Metal (Cation) and a Nonmetal (Anion) b. Involves the TRANSFER OF ELECTRONS c. Interaction between ions (atoms with a charge) d. Compounds held together by electrical charges e. Bond gets stronger as magnitude of charges increase f. Na + and Cl - weaker than Mg 2+ and O 2- g. Ionic compounds commonly known as salts h. Ex: NaCl Table Salt i. Solid at room temperature j. High melting and boiling points k. Due to strong electrostatic interaction between (+) and (-) charges l. Conduct electricity when molten m. When dissolved in water, they are strong electrolytes n. Ionic compounds do not exist as individual molecules, but rather as a group of molecules o. Basic structure of an ionic compound is a lattice-work of oppositely charged particles i. Know how to write ionic formulas: ii. Ca + F a. Use your t chart i. CaF 2 iii. Rubidium and Oxygen a. Rb + O = Rb 2O b. Remember metals with a positive charge are always written first in the molecule i. What compound would be formed when A 2+ bonds with B -3? 1. A 3B 2 iv. Know how to name ionic compounds: v. NaBr a. Sodium bromide vi. K 2O a. Potassium oxide vii. Know how to write ionic formulas with polyatomics: viii. Cs + NO 3 a. Cs(NO 3) ix. Aluminum Acetate a. Al + C 2H 3O 2 i. Al(C 2H 3O 2) 3 x. Know how to name ionic compounds with polyatomics: xi. Li 2CO 3 a. Lithium carbonate xii. Be 3(PO 4) 2 a. Beryllium phosphate xiii. xiv. Know how to write ionic formulas with transition metals: Copper (II) Nitrate a. The roman numeral means that charge of the metal b. Use your t chart i. Cu(NO 3) 2

9 xv. Iron (III) Oxide a. Fe 2O 3 xvi. Know how to name ionic compounds with transition metals: xvii. HgCl 2 a. Mercury (II) chloride xviii. CoSO 3 a. Cobalt (II) sulfite b. What do the numbers inside the parentheses after a transition metal mean? i. They are the charge of the transition metal xix. What are the charges for Ag, Zn, and Cd? i. Ag = +1 ii. Zn = +2 iii. Cd = Define covalent compounds. Be able to write the names/formulas for covalent compounds. a. A.K.A. Molecular Bonds b. Nonmetal + Nonmetal c. NOT composed of ions! d. Involves SHARING OF ELECTRONS e. Involves multiple bonds f. Single, Double, or Triple g. Shown by dashes to indicate shared pairs of e- h. Relatively low melting points and boiling points i. DO NOT conduct electricity when molten j. When dissolved in water, can be strong electrolytes, weak electrolytes, or nonelectrolytes k. Covalent compounds exist as individual molecules that have a definite composition l. They are held together by attractions of SHARED ELECTRONS 6. Define polyatomic compounds. Be able to write polyatomic ionic compounds using the criss cross method. Be able to name the compounds. i. Polyatomic ions look like molecules but they have an overall charge ii. The group of elements, together, have too many electrons or too little electrons iii. Therefore they have a positive or negative charge iv. Characteristics of Polyatomics v. Polyatomic ions are usually a result of a covalent bond between 2 elements vi. Ex: O + H OH - vii. Because of its charge, it then forms an ionic bond with other elements or compounds viii. Ex: OH - + Na + NaOH

10 Metals: Bonding that takes place between only one kind of metal atom Structure is similar to ionic compound structure Lattice-work structure Only one kind of atom 7. Know how to determine how many atoms of each type of element are present in a molecule. a. subscripts 8. What are the naming rules for elements? How do we name them? a. Look at periodic table 9. What are the 7 diatomics? What are the naming rules for diatomics? How do we name them? Hydrogen (H 2) Nitrogen (N 2) Oxygen (O 2) Fluorine (F 2) Chlorine (Cl 2) Iodine (I 2) Bromine (Br 2) 10. What is a binary ionic compound? What are the naming rules for them? Give an example. a. Ionic compound made up of one metal and one nonmetal b. Name metal first as it appears on pt c. Name nonmetal second, change the ending of the name to -ide 11. What is a transition compound? What are the naming rules for them? Give an example. a. Metal with more than one oxidation state b. Use parentheses to represent the charge of the metal 12. What is a polyatomic ionic compound? What are the naming rules for them? Give an example. a. Compound with three or more elements b. Compound made up of a polyatomic ion c. Don t change the ending of the polyatomic name 13. What is a covalent compound? What are the naming rules for them? Give an example. a. Two nonmetals bonded together b. Name first element first and include a prefix if there is a subscript other than one c. Name second element second and change the ending to ide always include a prefix 14. What is an acid? What are the two types? What are the naming rules for each of them? Give an example. a. A compound with hydrogen that is not water b. Binary hydrogen and an element c. Ternary hydrogen and a polyatomic d. If binary hydro ic acid e. If ternary based off of polyatomic i. ate ic acid ii. ite ous acid

11 15. What is a hydrate? What are the naming rules for them? Give an example. a. An ionic compound that has bonded to water b. Is a solid that you can dehydrate (remove the water from) c. Name the ionic compound using the ionic naming rules, then use the greek prefixes to name the water (hydrate) 16. How do you convert from grams to moles? a. Use molar mass b. # grams /1 x 1 mol / molar mass g = moles 17. How to compute percent composition of an atom in a compound. a. Find the molar mass for the compound b. Divide the mass of each element (times their subscript) by the molar mass, then multiply by 100 c. Example: What is the percent of hydrogen in C 6H 8(OH) 6? i. C = 6 x 12 = 72 ii. H = x 1 = 14 iii. O = 6 x 16 = 96 iv. MM = 182 v. H = 14/182 x 100 = 7.6% 18. How to compute the empirical formula of a compound based on its percent composition. a. Divide the percent of each element by the elements masses b. Divide each answer by the smallest ratio c. Use your answers as the subscripts unless there is a.5 d. If.5, multiply all subscripts by 2 to get to a whole number e. Example: You have a compound containing 83% rubidium, 16% oxygen, and 1% hydrogen. What is the empirical formula for the molecule? i. 83% Rb/ 85 = / = 1 ii. 16% O/16 = 1 / = 1.02 iii. 1% H/1 = 1 / = 1.02 iv. = RbOH

12 19. Know how to use the empirical formula to determine the molecular formula. a. Find the empirical formula if not given i. Divide the percent of each element by the elements masses ii. Divide each answer by the smallest ratio iii. Use your answers as the subscripts unless there is a.5 iv. If.5, multiply all subscripts by 2 to get to a whole number b. Find molecular formula i. Find mass of empirical formula ii. Divide mass given by the mass of the empirical formula iii. Multiple the empirical formula by the answer to the division c. You have a substance with 15.75% hydrogen and the rest is carbon. The molecule has a mass of What is the molecular formula of the molecule? i. Hint: Use the percentages to get the empirical formula, then get the mass of the empirical formula, compare to the mass of the molecule, and multiply the empirical formula by the difference between the mass of the molecule and the mass of the empirical formula % H/1 = / = % C/12 = / = 1 = H 2C Mass = /14 = = H 2C 20. What is a mole? How many molecules make up a mole? What is it used for? a. A mole is the amount of pure substance containing the same number of chemical units as there are atoms in exactly 12 grams of carbon-12 (i.e., X ). b. 1 mole = 6.02 x molecules c. Used as a conversion factor, with stoichiometry, used to make the molecule and atom easier to work with 21. What is a chemical reaction? a. A reaction where the chemical composition changes 22. What are the 5 types of reactions? a. Synthesis (combination), decomposition, single replacement, double replacement, combustion 23. What is the law of conservation of mass? How does it apply to balanced equations? a. The law of conservation of mass states that matter is not created or destroyed during a reaction b. This is the reason why we balance equations 24. Be able to balance equations and tell me what type of reaction is occurring. 1. 1CH4 + 1O2 1CO2 + 2H2 a. Single replacement - anionic 2. 1Ba + 2H2O 1Ba(OH)2 + 1H2 a. single replacement - cationic 25. What side of a reaction is called the reactants? What side is called the products? a. Left b. right 26. What does (s), (l), (g), and (aq) mean? a. S solid b. L liquid c. G gas d. Aq aqueous dissolved in water 27. Predict the products of the following reactions: a. Then balance 1. Synthesis: a. 2Li +1Cl2 2LiCl b. 1Mg + 1F2 1MgF2 c. Rb + O2 Rb2O 2. Decomposition: a. 1CO2 1C + 1O2 b. 2LiH 2Li + 1H2

13 c. 2H2O 2H2 + 1O2 3. Single Displacement: a. 1MgBr2 + 1Cl2 1MgCl2 + 1Br2 b. 2Al + 1Fe2O3 1Al2O3 + 2Fe c. Silver + zinc chloride no reaction 4. Double displacement: a. 1NaOH + 1KCl 1NaCl + 1KOH b. Silver nitrate + zinc chloride - 2AgNO3 + 1ZnCl2 2 AgCl + 1Zn(NO3)2 28. You should be familiar with how to use stoichiometry for a balanced equation. a. Convert to moles if not in moles b. Solve using BCA or Factor Label method c. Convert to grams if desired d. If a limiting vs excess reactant problem, determine which reactant is the limiting i. Convert to moles if not in moles ii. Divide the amount given in moles by the coefficients whichever is the smaller ratio is the limiting e. Use the limiting reactant to solve for the amount of product formed f. Convert to grams if desired BCA tables Stoichiometry: Predicting Amounts in Reactions Stoichiometry is the process of determining how much product is made or how much reactant is needed during a chemical reaction. As we know, in chemical reactions atoms are conserved. We show this in a balanced chemical equation. The balanced chemical equation tells us two things: 1. Which substances begin with (reactants) and end with (products) during the rearrangement process. 2. The ratio of particles involved. This ratio can be seen either as a ratio of individual particles OR as a ratio of moles. In the lab it is only practical to work with moles of substances rather than individual atoms or molecules, and so we interpret our equations as a ratio of moles, or a mole ratio. Example: 2 Mg + 1O2 --> 2 MgO means for every 2 moles of Mg burned, 1 mole of O2 is required to produce 2 moles of MgO, or a ratio of Mg : 1 mole O2 : 2 moles MgO 2 moles We can use this mole ratio relationship to make predictions about how much we need of something, or how much we can make from what we have. Making Predictions In every reaction, there are three stages we need to consider to make good predictions: 1. Before: What we have before the reaction takes place. 2. Change: How much of each substance actually changes during the reaction 3. After: How much of each substance is present after the reaction is complete. Some good organization can help us in making good predictions. We have an organizational table that can help us track the Before-Change-After for a reaction. Below is an example of a problem involving a chemical reaction. Sample Problem: Hydrogen sulfide gas, which smells like rotten eggs, burns in air to produce sulfur dioxide and water. How many moles of oxygen gas would be needed to completely burn 2.4 moles of hydrogen sulfide? Step 1- Write and Balance the equation (describe the reaction and its mole ratio) 2 H2S + 3 O2 --> 2 SO2 + 2 H2O Before: Change

14 After Step 2: Fill in the Before line with the Given information; mark what you must Find on the table (with units) 2 H2S + 3 O2 --> 2 SO2 + 2 H2O Before: 2.4moles xs moles 0 moles 0 moles Change moles After NOTE: Assume reactants you don t have amounts for are present with more than enough available (excess, or xs ) for the reaction to be completed. Step 3: Use ratio of coefficients to determine the Change made 2 H2S + 3 O2 --> 2 SO2 + 2 H2O Before: 2.4moles xs moles 0 moles 0 moles Change: 2.4 moles 3.6 moles +2.4 moles +2.4 moles After NOTE: Reactants are consumed/decrease (-), products accumulate/increase (+) Step 4: Complete the table for what remains After the reaction is complete 2 H2S + 3 O2 --> 2 SO2 + 2 H2O Before: 2.4 moles xs moles 0 moles 0 moles Change: 2.4 moles 3.6 moles +2.4 moles +2.4 moles After: 0 moles xs moles 2.4 moles 2.4 moles In this case, desired answer is in moles Answer (in moles): 3.6 moles O2 are needed to burn 2.4 moles H2S. If mass is required, convert moles to grams in the usual way 3.6moles O2 * 32 grams = 115 grams O2 1 mole Answer (in grams): 115 grams O2 are needed to burn 2.4 moles H2S. i. You have the following equation: 2ZnS + 3O2 2ZnO + 2SO2 1. If you have 6 moles of ZnS, how many moles of ZnO will be produced? (Show all of your work) 6 Mols Zns / 1 x 2 mols ZnO / 2 mols ZnS = 6 mols ZnO 2. If you have 9 moles of O2, how many moles of SO2 will be produced? (Show all of your work) 9 mols O 2 / 1 x 2 mols SO 2 / 3 mols O 2 = 6 mols SO 2 3. If you have 65 g of ZnO produced, how many moles of O2 were used in the reaction? 65 g ZnO / 1 x 1 mol ZnO / 81 g x 3 mols O 2/2 mols ZnO = 1.2 mols O 2 4. If you have 150 g of oxygen, how many grams of sulfur dioxide is produced? 150 g O 2/1 x 1 mol O 2/32 g x 2 mols SO 2/3 mols O 2 x 64 g SO 2/1mol = 200 g SO What side of a reaction is called the reactants? The products? a. Left b. Right 30. What is percent yield? How do you solve for it? a. The percent yield looks at how much product is produced by a reaction compared to the expected amount due to the stoichiometry b. In chemistry, the reaction yield is the amount of product produced by a chemical reaction. The theoretical yield is the maximum amount of product that can be produced in a perfectly efficient reaction. In reality, most

15 reactions are not perfectly efficient - the reaction's actual yield is usually less than the theoretical yield. To express the efficiency of a reaction, you can calculate the percent yield using this formula: %yield = (actual yield/theoretical yield) x 100. i. Actual yield is from the experiment ii. Theoretical yield is the answer from the stoichiometry problem Theoretical yield: Amount of substance you expect to produce in grams. This value is calculated from your BCA table. The 3.6 moles of O2 was taken from the BCA table, the conversion to grams is no different than the grams to mole conversions we normally do. Theoretical yield is just a fancy word for how much stuff should I make. If mass is required, convert moles to grams in the usual way 2.4moles SO2 * 64 grams = grams O2 1 mole Answer (in grams): grams O2 are made when burning 2.4 moles H2S. Percent Yield Percent yield is used to determine how good you are at performing labs. (don't worry your grade will not be determined by the quality of your percent yield). To determine your percent yield you take the amount of product you actually got (actual yield) and divide it by the amount of product you should have gotten (theoretical yield). Lets look at our theoretical yield for oxygen in the above problem. We should have made gram s of sulfur dioxide. Lets say our lab only produced grams of product. To find our percent yield: Actual Yield X 100 = X 100 = 65.3% Yield Theoretical Yield Our boss may have something to say about that... Usually you would like your percent yield to be 90% or better, but it really depends on what kind of reaction you are doing and to how good of a percent yield you can realistically expect.

16 Unit 4 Topics: Aqueous solutions Molarity Molality Precipitation reactions Solubility rules Redox reactions Ionic equation Net ionic equations Questions: 1. What is the definition of molarity? a. The concentration of a solute in solution can be expressed in terms of molarity. molarity (M) = moles of solute Liter of solution b. The symbol [ ] is commonly used to represent the molarity of a substance in solution. 2. What is the definition of molality? a. Similar to molarity except mols per kg b. Abbreviated as m c. The concentration of a solute in solution can be expressed in terms of molality. molality (m) = moles of solute kg of solution 3. Be able to use molarity and molality for calculations and conversions. a. Molarity b. What is the molarity of a solution made by dissolving 2.5 g of NaCl in enough water to make 125 ml of solution? c. 2.5 g NaCl x 1 mole NaCl g NaCl = mole d. molarity = mole NaCl L = 0.34 M NaCl e. Conversions: i. What mass of K 3PO 4 is required to prepare 4.00 Liters of 1.50 M solution? L 1.50 mol 212 g x x = 1272 g K 3PO 4 i. 1 1 L 1 mol 3. Dissolve 1272 g K 3PO 4 in enough water to make 4.00 L of solution. f. Molality i moles NaOH ii. 40 g NaOH iii. 1 L solution v. = 5 g of NaOH dissolved in 1 L of Water iv. 1 mole NaOH g. Dilutions with molarity: i. Dilution process of adding more solvent to solution to prepare a more dilute solution ii. Concentrated solution large amount of solute iii. Dilute solution small amount of solute iv. M 1V 1 = M 2V 2 4. What are precipitation reactions? Be able to use the solubility rules for precipitation reactions. a. A precipitation reaction is a reaction which results in the formation of an insoluble product, or precipitate. b. A precipitation is an insoluble solid that separates form the solution. c. Precipitations reactions usually involve ionic compounds. They result from double displacement reactions.

17 5. Be able to write net ionic equations. Be able to identify and cancel spectator ions. a. A chemical equation for a reaction involving ions in which only those species that actually react are included. i. Atom balance- There must be the same number of atoms on both sides. ii. Charge balance- There must be the same total charge on both sides b. Example: Molecular Equation i. Pb(NO 3) 2 (aq) + 2 NaI (aq) PbI 2 (s) + 2 NaNO 3 (aq) c. Steps: 1 break each aqueous item into its ions leave solids and liquids together i. Pb +2 (aq) + NO 3 - (aq) + Na + (aq) + I - (aq) PbI 2 (s) + Na + (aq) + NO 3 - (aq) d. Steps: 2 use coefficients to balance the number of ions and the charges on each side of the arrow e. Ionic Equation i. Pb +2 (aq) + 2 NO 3 - (aq) + 2 Na + (aq) + 2 I - (aq) PbI 2 (s) + 2 Na + (aq) + 2 NO 3 - (aq) f. Steps: 3 cancel out ions that appear in identical forms among both the reactants and products of an ionic equation g. Ionic Equation i. Pb +2 (aq) + 2 NO 3 - (aq) + 2 Na + (aq) + 2 I - (aq) PbI 2 (s) + 2 Na + (aq) + 2 NO 3 - (aq) h. Net ionic equation i. Pb +2 (aq) + 2 I - (aq) PbI 2 (s) i. What was cancelled out are called spectator ions i. they are present during a reaction, but do not play a role in the reaction 6. How many chloride ions are in 200 ml of a 2.5 M solution of calcium chloride? a. 200 ml x 1L/1000 ml x 2.5 mol / 1L x 2 atoms(subscript) = 1 mol of atoms 7. Know how to solve and write redox reactions in acidic and basic solutions. i. Acidic solution: 1) Separate the given reaction into 2 half reactions. 2) Balance all atoms except O, H. 3) Balance O atoms by adding H 2O.s. 4) Balance H atoms by adding H+ ions. 5) Balance the charge by adding electrons (e-) to the more positive side. 6) Multiply each reaction by an integer so that the # electrons gained = # electrons lost. 7) Combine the half reactions and simplify by canceling common species. You should have H+ left in your reaction. ii. Basic solution: 1) Do steps 1-7 above. 2) Note the # of H+ ions. Add this number of OH- ions to both sides.

18 3) Combine OH- and H+ to form H 2O and simplify the reaction. iii. *If the equation is correctly balanced, the number of each type of atom and the net charge must be equal for both sides. 8. What does reduction mean? Oxidation? What is a reducing agent? And oxidizing agent? a. Reduction gaining electrons b. Oxidation losing electrons c. Reducing agent the one that is oxidized d. Oxidizing agent the one that is reduced 9. What are acids? Properties? i. Acidic solutions have a sour taste. ii. Vinegar, lemon juice, and soda are acidic. iii. An acid is a species that produces H + ions in water solution. iv. Properties of an acid v. A substance that produces excess of H + ions in water vi. Almost all acids contain hydrogen vii. Sour taste viii. React with carbonate in rocks to produce carbon dioxide gas as they dissolve the rock ix. React with some metals to produce hydrogen gas x. Aqueous acid solutions conduct electricity xi. ph = 0-6 xii. Arrhenius: Acids produce H + ions in solution, bases produce OH - ions. xiii. Brønsted Lowry: Acids are proton (H + ) donors, bases are proton acceptors. i. HCl + H 2O Cl - + H 3O + ii. acid base 10. What are bases? Properties? a. Basic solutions have a slippery feeling b. Ammonia, most detergents and cleaning agents are basic. c. A base is a species that produces OH - ions in water solution.

19 d. Properties of bases e. A substance that forms excess OH - ions in water f. Most contain hydroxide (OH - ) g. Bitter taste h. Feel slippery i. ph is between 8 and 14 j. React with acids to form salts 11. What are the three ways acids can be classified? Provide examples for each way. a. Strong vs. weak b. Monoprotic vs. polyprotic c. Binary vs. ternary 12. What are the strong acids you were asked to know? a. HCl b. HClO 4 c. HBr d. HI e. HNO 3 f. H 2SO What are the weak acids you were asked to know? a. HCN b. HF c. HC 2H 3O 2 d. HNO 2 e. H 2SO What are the strong bases you were asked to know? a. LiOH b. NaOH c. KOH d. Ca(OH) 2 e. Ba(OH) 2 f. Sr(OH) What are the weak bases you were asked to know? a. NH 3 b. Na 2CO What is ph? a. The acidity or basicity of a solution can be described in terms of its [H + ] b. The ph is a number that is derived from the concentration of hydronium ions [H 3O + ] in a solution c. [H 3O + ] and the [H + ] are the same d. Another word for concentration is molarity (mols per liter) e. ph = negative logarithm (to the base of 10) of the hydrogen concentration 17. What is the formula for ph? a. ph = - log [H3O + ] 18. Be able to solve for ph or [H] using the ph formula. a. ph = - log [H3O + ] b. 10 -ph = [H + ] 19. What is poh? i. Calculation of the concentration of a base ii. Tells how much OH- is present iii. Tells us how basic a solution is iv. Cannot be used directly to tell you ph 20. How can poh be used to determine ph? i. no 21. What is the formula for poh? i. p[oh] = - log [OH - ] 22. What is a conjugate base? Provide an example. i. Conjugate base: ii. The ion or molecule that is formed when an acid gives up (loses) a hydrogen ion iii. B - is the conjugate base of HB 23. What is a conjugate acid? Provide an example.

20 i. Conjugate acid: ii. The ion or molecule that is formed when a base accepts (gains) a hydrogen ion iii. HA is the conjugate acid of A - iv. In weak acid/base reactions, hydrogen ions are very similar to naked protons they have a positive charge, so if you lose one you end up as a negative ion and if you gain one you end up gaining a positive charge 24. What is titration? i. Lab techniques used to find the concentration of an acid or base solution ii. Definition: 1. Procedure in which a solution of known concentration is used to determine the concentration of a second unknown solution iii. Known concentration titration standard iv. Indicator any substance in a solution that changes its color as it reacts with either an acid or a base 25. What is the formula for titration? i. MaVa = MbVb 26. How can titration be used? i. Titration measuring the volume of a standard solution (solution of known concentration) required to react with a measured amount of sample ii. Objective determine the point at which the reaction is complete (called the equivalence point) iii. Reached when the number of moles of OH- added is exactly equal to the number of moles of acid originally present 27. Why is it important to understand acids and bases? i. Acids and bases are all around you ii. Acids and bases are in the body 1. Your body is buffered (your blood s ph must stay relatively the same; a change of 0.5 or more would cause your body to stop working/functioning) 2. Amino acids are weak acids that are subunits of proteins 3. Nucleic acids are made from acid and base components 4. Fatty acids are subunits of lipids 5. Antacid tablets iii. Acids and bases affect the environment around us

21 Unit 5 Topics: Kinetic theory of gases Gas volume vs. pressure changes Gases temperatures with vectors (pictures) Gas laws Charles s Boyle s Combined Ideal Questions: 1. What are the gas laws? 1. Laws for the gas relationships 2. Charles s 3. Boyle s 4. Combined 5. Ideal 2. What is Charles law? What is the formula? What does it tell us? 1. Charles' law (also known as the law of volumes) is an experimental gas law that describes how gases tend to expand when heated. 2. Shows that temperature and volume do the same (both increase, decrease) 3. What happens to the volume of a gas as the temperature increases? 1. increases 4. What is Boyle s law? What is the formula? What does it tell us? 1. Boyle's law (sometimes referred to as the Boyle Mariotte law, or Mariotte's law) is an experimental gas law that describes how the pressure of a gas tends to increase as the volume of a gas decreases. 2. Shows that pressure and volume are opposites of each other What happens to the pressure of a gas as the volume increases? 1. decreases 6. What is the combined gas law? What is the formula? What does it tell us? 1. The combined gas law is a gas law that combines Charles's law, Boyle's law, and Gay-Lussac's law. There is no official founder for this law; it is merely an amalgamation of the three previously discovered laws. These laws each relate one thermodynamic variable to another mathematically while holding everything else constant. 2. P1V1T2 = P2V2T1 7. How are pressure, volume, and temperature related? 1. Pressure and volume are opposites 2. Pressure and temperature do the same 3. Volume and temperature do the same 8. What is the ideal gas law? What is the formula? What does it tell us? 1. The ideal gas law is the equation of state of a hypothetical ideal gas. It is a good approximation to the behavior of many gases under many conditions, although it has several limitations What is dalton s gas law? What is the formula? What does it tell us? 1. In chemistry and physics, Dalton's law (also called Dalton's law of partial pressures) states that in a mixture of non-reacting gases, the total pressure exerted is equal to the sum of the partial pressures of the individual gases or 10. What is graham s gas law? What is the formula? What does it tell us? 1. Graham found experimentally that the rate of effusion of a gas is inversely proportional to the square root of the mass of its particles

22 Unit 6 Topics: Thermochemistry Temperature vs. heat Endothermic vs. exothermic Spontaneous vs. nonspontaneous Energy diagrams Activation energy diagrams 1 st law of thermodynamics 2 nd law of thermodynamics Enthalpy Entropy Gibb s free energy Specific heat Questions: 1. What is the definition of thermochemistry? a. Deals with the changes in heat energy that accompany a chemical reaction 2. What is the difference between temperature and heat? a. Temperature i. A measure of the average kinetic energy of random motion of the particles in a substance b. Heat i. A measure of the total amount of energy transferred from an object of high temperature to one of low temperature 3. What is the definition of an exothermic reaction? Endothermic? What are some examples? How do each of these reactions feel to you? a. Exothermic i. A physical or chemical change in which energy is released by a system to its surroundings 1. Example: Rain forms or liquid water freezes (decrease the internal energy of the water molecules) 2. Most common more energy is released when forming the new bonds than was used when breaking the old bonds 3. Feels hot to you b. Endothermic i. A physical or chemical change in which a system absorbs energy from its surroundings 1. Example: Melting or boiling (increase the internal energy of the water molecules) 2. Less likely to occur require special conditions 3. Feels cold to you 4. Which type of reaction is more likely to occur endothermic or exothermic? a. exothermic 5. What does a reaction-energy profile look like? What is a reaction-energy profile used for? Be able to draw one for an endothermic reaction (lose energy) and an exothermic reaction (gain energy).

23 6. What does the 1 st law of thermodynamics say? a. Energy must be conserved b. We can add energy or take away energy from materials c. The Internal Energy considers all energy within the material - kinetic & potential d. The Change in this Energy: e. E = E final - E initial 7. What is enthalpy? What is the definition and what is it used for? What is the symbol for enthalpy? What is the formula for the change of enthalpy during a reaction? What does a positive change in enthalpy mean? What does a negative change in enthalpy mean? What are the units for enthalpy? Make sure you can solve problems for the change in enthalpy if you are given the enthalpy for each molecule in the reaction. (What is the equation for the change in entropy?) a. Heat transferred from a system under constant pressure b. Enthalpy of a reaction is the heat transferred into or out of the materials during a chemical change. c. Heat energy is measured in a quantity called enthalpy d. State property e. It is represented as H i. The change in heat energy that accompanies a chemical reaction is represented as H f. H = H (products) - H (reactants) i. Positive H values indicate an endothermic reaction energy is absorbed ii. Negative H values indicate an exothermic reaction energy is given off (released) 1. The difference in enthalpy between the reactants and the products equals the amount of heat change during the reaction 8. Define activation energy. What is the symbol for activation energy? What happens if the activation energy is not met for a reaction? Be able to label a drawing of a reaction pathway showing activation energy and the change in enthalpy for the reaction. a. Activation energy: i. The minimum amount of energy that must be supplied to a system to start a chemical change ii. Illustrated using the activation energy diagrams

24 9. What is the definition of entropy? What is the symbol used for entropy? Know how to finish this statement, the higher the entropy (what?) a. Can be thought of as a measure of the amount of disorder or randomness in a solution b. If everything is ordered and neat low entropy system c. State property d. Gas has higher entropy than liquid, Liquid has higher entropy than solid, aqueous has higher entropy than solid, and solid is the lowest e. Increasing the temperature of a substance increases its entropy ( see figure 17.2) f. Represented by the letter S g. Units are J / K (Joules / Kelvin) h. Equation: i. S = S (products) - S (reactants) j. What does entropy tell us? i. Positive spontaneous, more disordered ii. Negative non-spontaneous, less disordered k. The more disordered 10. Know what increases the entropy of a reaction (what is entropy effect by). As we increase (what), we increase entropy? a. Disorder b. Temperature i. In order for anything to happen spontaneously anywhere in the universe, the entropy of the universe must increase ii. The universe prefers more randomness rather than less (Likes things disordered) iii. Entropy becomes more important as temperature increases

25 11. Which of the physical states (gas, liquid, solid) has the highest entropy? Be able to explain this in a short answer response. a. Gas most disorder, most energy 12. Be able to solve for the change in entropy for a system if you are given the entropy for each molecule in the reaction. (What is the equation for the change in entropy)? What are the units of entropy? a. S = S (products) - S (reactants) b. What does entropy tell us? i. Positive spontaneous, more disordered ii. Negative non-spontaneous, less disordered 13. What does the second law of thermodynamics say? a. In order for anything to happen spontaneously anywhere in the universe, the entropy of the universe must increase b. The universe prefers more randomness rather than less (Likes things disordered) c. Entropy becomes more important as temperature increases 14. Can the effects of entropy by overcome? (Meaning, even though the universe wants to increase the entropy of a reaction, can we have reactions that decrease the entropy?) Be able to answer this question in a short answer response. a. Yes, if the enthalpy is negative (exothermic and spontaneous) and there is low temperature 15. What is the trend between enthalpy & entropy and spontaneity? Is the reaction spontaneous or nonspontaneous when a. Enthalpy is negative and entropy is positive i. Spontaneous at all temperatures b. Enthalpy is negative and entropy is negative i. Spontaneous at low temperatures only c. Enthalpy is positive and entropy is positive i. Spontaneous at high temperatures only d. Enthalpy is positive and entropy is negative i. Never spontaneous 16. What is free energy? What letter is it represented by? What is the equation for free energy? Is the reaction spontaneous or nonspontaneous when Gibb s free energy is positive and when Gibb s free energy is negative? What are the units for free energy? i. Free energy is a combination of entropy and enthalpy 1. Uses temperature to determine whether entropy or enthalpy can cause a change to occur ii. Free energy is the quantity of energy that is available or stored to do useful work or to cause a change iii. Represented by the letter G iv. Equation is (Gibbs-Helmholtz equation) 1. G = H - T ( S) v. Positive 1. The change is not spontaneous 2. This does not mean that it cannot happen, it just will not occur spontaneously (naturally, on its own) 3. The reverse is spontaneous vi. Negative 1. The change is spontaneous 2. The reaction will occur naturally on its own without any outside help vii The system is at equilibrium

26 2. There is no tendency for the reaction to occur in either direction 17. What is the definition of specific heat? What is symbol for specific heat? What is the change in heat using specific heat? a. The amount of heat required to raise the temperature of a one gram of a substance one degree Celsius b. Each pure substance has a unique specific heat i. This means that it can be used to identify an unknown substance c. A substance with low specific heat means that it requires a small amount of energy to feel hot i. Example: firework, explosive d. The higher the specific heat, the more energy that is required to heat the substance i. Example: ice, earth, oceans ii.

27 Unit 8 Topics: Ionic, covalent, metallic, polyatomic bonds and compounds Bonding Shapes of molecules Ionic Covalent Lewis dot structures Electron group (parent structures) Molecular shapes Polarity of bonds Polarity of molecules Intermolecular forces Questions: 1. What is the difference between an ionic and covalent bond? a. Ionic metal and non-metal i. Transfer electrons b. Covalent two nonmetals i. Share electrons ii. Stronger bond 2. Why elements bond? a. To achieve a noble gas configuration 3. How many electrons do the atoms want in their valence shell? a. Most want 8 but remember that there are exceptions 4. What happens to energy in regards to forming and breaking bonds? a. Energy is required to break bonds b. Energy is released when forming bonds 5. What are the three types of bonding? What types of molecules are bonded together with each type? a. Ionic Bonding (Metal + Nonmetal) b. Covalent Bonding (Nonmetal + Nonmetal) c. Polyatomic Bonding (Both Ionic and Covalent Characteristics > 3 or more elements including metal and nonmetals) d. Metallic Bonding (One type of metal atom) 6. What is an ionic compound? What is it made out of? What happens to the electrons in an ionic compound? Be able to draw a Lewis Dot structure for ionic bonding. a. Bond between a Metal (Cation) and a Nonmetal (Anion) b. Involves the TRANSFER OF ELECTRONS c. Interaction between ions (atoms with a charge) d. Compounds held together by electrical charges e. Bond gets stronger as magnitude of charges increase f. Na + and Cl - weaker than Mg 2+ and O 2- g. Ionic compounds commonly known as salts h. Ex: NaCl Table Salt i. Solid at room temperature j. High melting and boiling points k. Due to strong electrostatic interaction between (+) and (-) charges l. Conduct electricity when molten m. When dissolved in water, they are strong electrolytes n. Ionic compounds do not exist as individual molecules, but rather as a group of molecules o. Basic structure of an ionic compound is a lattice-work of oppositely charged particles p. Step One: Write the correct formula for the compound q. Step Two: List each element separate, making sure to evenly distribute the anion and cation with each other r. Step Three: Draw the Lewis Dot diagrams of each of the individual elements s. Step Four: Use arrows to show how the electrons move from the cation(s) to the anion(s)

28 t. End result: The cation(s) do not have any valence electrons, while the anion(s) have complete shell of valence electrons 7. What is a covalent compound? What is it made out of? What happens to the electrons in a covalent compound? Be able to draw a Lewis Dot structure for covalent bonding. a. A.K.A. Molecular Bonds b. Nonmetal + Nonmetal c. NOT composed of ions! d. Involves SHARING OF ELECTRONS e. Involves multiple bonds f. Single, Double, or Triple g. Shown by dashes to indicate shared pairs of e- h. Relatively low melting points and boiling points i. DO NOT conduct electricity when molten j. When dissolved in water, can be strong electrolytes, weak electrolytes, or nonelectrolytes k. Covalent compounds exist as individual molecules that have a definite composition l. They are held together by attractions of SHARED ELECTRONS Step 1: Write down the formula for your compound Step 2: evenly distribute the symbols of the different elements in your compound Step 3: determine how many valence electrons each of the elements in your compound has. Add these together to determine the total number of valence electrons that you have to have in your diagram Step 4: determine how many bonds each of the elements probably wants to make Step 5: Use bonds to connect the elements Step 6: Add lone pairs of electrons to the elements to fulfill the octet rule Step 7: Check that you have the correct number of valence electrons being used in your diagram Note if you have too many valence electrons being shown in your diagram, erase your lone pairs of electrons and add a double or triple bond to your elements Notes if you have too few valence electrons being shown, erase your bonds and instead of using a double or triple bond, use a single or double bond 8. What is a polyatomic compound? What type of bonding occurs there? a. Polyatomic ions look like molecules but they have an overall charge b. The group of elements, together, have too many electrons or too little electrons c. Therefore they have a positive or negative charge d. Characteristics of Polyatomics e. Polyatomic ions are usually a result of a covalent bond between 2 elements f. Ex: O + H OH - g. Because of its charge, it then forms an ionic bond with other elements or compounds h. Ex: OH - + Na + NaOH 9. What is a resonance structure? What are the two compounds that we looked at that had resonance structure? Be able to identify resonance structures when looking at them. a. The concept of resonance is involved whenever a single Lewis structure does not adequately reflect the properties of a substance. Resonance forms differ only in the distribution of e -.

29 10. What are the physical properties of an ionic bond? What are the physical properties of a covalent bond? i. Bond between a Metal (Cation) and a Nonmetal (Anion) ii. Involves the TRANSFER OF ELECTRONS iii. Interaction between ions (atoms with a charge) iv. Compounds held together by electrical charges v. Bond gets stronger as magnitude of charges increase vi. Na + and Cl - weaker than Mg 2+ and O 2- vii. Ionic compounds commonly known as salts viii. Ex: NaCl Table Salt ix. Solid at room temperature x. High melting and boiling points xi. Due to strong electrostatic interaction between (+) and (-) charges xii. Conduct electricity when molten xiii. When dissolved in water, they are strong electrolytes xiv. Ionic compounds do not exist as individual molecules, but rather as a group of molecules xv. Basic structure of an ionic compound is a lattice-work of oppositely charged particles I. A.K.A. Molecular Bonds II. Nonmetal + Nonmetal III. NOT composed of ions! IV. Involves SHARING OF ELECTRONS V. Involves multiple bonds VI. Single, Double, or Triple VII. Shown by dashes to indicate shared pairs of e- VIII. Relatively low melting points and boiling points IX. DO NOT conduct electricity when molten X. When dissolved in water, can be strong electrolytes, weak electrolytes, or nonelectrolytes XI. Covalent compounds exist as individual molecules that have a definite composition XII. They are held together by attractions of SHARED ELECTRONS 11. What is the molecular structure of an ionic bond? a. 3D cube 12. Be able to determine the molecular structure of a covalent compound based on how many electron groups are present and how many lone pairs are present. 13. Number of electron groups n/a only 2 elements total Number of lone pairs n/a varies Electron geometry (Parent Shape) Linear Molecular Geometry (Molecular Shape) Only 2 elements Linear Examples N2 2 0 Linear CO2 Linear

30 3 0 Trigonal planar Trigonal Planar BF3 3 1 Trigonal planar Bent NO2 4 0 Tetrahedral Tetrahedral CCl4 4 1 Tetrahedral Pyramidal NH3 4 2 Tetrahedral Bent H2O Number of electron groups Number of lone pairs Electron geometry (Parent Shape) Molecular Geometry (Molecular Shape) Examples 5 0 Trigonal PCl5 Trigonal 6 0 Bipyramidal Octahedral Bipyramidal SF6 6 2 Octahedral Octahedral Square Planar XeF4 6 1 Octahedral See-Saw SF4

31 5 2 Trigonal T-Shaped ClF3 Bipyramidal 14. What is electronegativity? What is a polar covalent bond? Which elements are the most electronegative? a. Polarity takes a closer look at the trend of Electronegativity i. The ability of an atom to attract a pair of electrons to itself b. Refers to a separation of charge i. Can refer to a bond ii. Can refer to an entire molecule c. F, O, N most electronegative d. Polar bond is where the difference in the electronegativity is between.5 and 1.67 there is an uneven distribution of electrons in the bond or molecules 15. What is a nonpolar covalent bond? Be able to determine if the compound is a nonpolar covalent bond, a polar covalent bond, or an ionic bond by looking at the change in electronegativity. a. True covalent bond b. When two nonmetal atoms are bonded and they have the same electronegativity value c. They share the electrons equally d. No partial charge builds!! H H 16. Based on geometry, determine if a compound is polar or nonpolar. Be able to give examples of polar and nonpolar compounds a. Need Polar Bonds b. Need Lack of Symmetry c. Bent & Pyramids are Polar d. Unshared electrons usually make Polar e. Dissimilar atoms usually make polar. i. Step 1: Draw the Lewis Dot structure for the molecule ii. Step 2: Determine what electron-group, then what molecular shape the molecule has iii. Step 3: Identify where the positive and negative ends are for each bond iv. Step 4: Draw in the arrows for the polarity and see if the polar bonds cancel in the molecule f. If forces cancel out nonpolar, if they do not polar g. Linear, trigonal planar, and tetrahedral tend to be nonpolar but can be polar depending on forces h. Bent and pyramidal tend to be polar, but could be nonpolar depending on forces 17. What is the hybridization of the various molecular shapes? a. Linear 2sp b. Trigonal Planar 3sp 2 c. Tetrahedral 4sp 3 d. Trigonal Bipyramidal 5sp 3 d

32 e. Octahedral 6sp 3 d What are the different intermolecular forces? Which molecules have the different types? a. See next answer b. London all molecules c. Dipole polar molecules d. Hbond H with F, N, O 19. What are the general properties related to intramolecular or intermolecular forces? Which type of force is stronger? Intermolecular Forces Intramolecular forces (bonding forces) exist within molecules and influence the chemical properties. Intermolecular forces exist between molecules and influence the physical properties. We can think of H 2O in its three forms, ice, water and steam. In all three cases, the bond angles are the same, the dipole moment is the same, the molecular shape is the same and the hybridization of the oxygen is the same. However, the physical properties of H 2O are very different in the three states. As solid ice, H 2O possesses a definite shape and volume. It is incompressible. Liquid water possesses a definite volume, but will assume the shape of its container. It is slightly compressible. Steam will assume both the shape and volume of its container and is extremely compressible. Intermolecular forces (IMF) are the forces which cause real gases to deviate from ideal gas behavior. They are also responsible for the formation of the condensed phases, solids and liquids. The IMF govern the motion of molecules as well. In the gaseous phase, molecules are in random and constant motion. Each gas molecule moves independently of the others. In liquids, the molecules slide past each other freely. In solids, the molecules vibrate about fixed positions. Heating Curves The transitions between the phases, phase changes, can be viewed in terms of a Heating Curve, like the one shown below, for water. It is a plot of time versus temperature. The time axis represents the addition of heat as a function of time. The longer the time span, the more heat has been added to the system. In this Heating Curve, we are starting with ice at -20 o C.

33 As we add heat, we raise the temperature of the ice. The heat that we are adding is increasing the Kinetic Energy of the system (KE = ½ mu 2 ) which is proportional to the absolute temperature (K) of the system. In the solid phase, the allowed motions are in vibrational movements within the molecules. In the case of water, the O-H bonds are stretching and bending. The bond lengths and angles are oscillating around the predicted values. The amount of heat required to raise the temperature of the ice is determined by the heat capacity of ice, the heat required to change the temperature of 1 gram of ice by 1 o C. The heat capacity of each phase of each substance is unique, and depends on the chemical nature of the substance. When the temperature reaches 0 o C, the melting point of ice, further addition of heat does not change the temperature. At this phase transition temperature, the added energy goes to changing the Potential Energy of the system. This is the energy associated with the IMF, which are holding the H 2O molecules in the solid state. It is coulombic in nature, arising from the attraction of charged species. In the case of H 2O, it is the attraction between the partial positive charges on the H and the partial negative charges on the O. As we discussed earlier in the semester, these are hydrogen bonds, holding the water molecules in the crystalline structure of ice. At the phase transition temperature, 0 o C, all of the ice will be converted to liquid water. As soon as the phase change is complete (s l), addition of heat will then lead to an increase in temperature of the liquid water. The increase in temperature is, again, an increase in the KE of the system. The movement of the water molecules will increase in the liquid phase. There is still some degree of hydrogen bonding between molecules, but they are no longer in fixed positions in a crystal lattice. There is a second phase transition at 100 o C. At this temperature, the water, at 100 o C, is converted to steam at 100 o C. The remaining hydrogen bonds are broken, and all of the water molecules are now moving independently of each other, with no remaining hydrogen bonding. The liquid water is converted to steam. As soon as this happens, addition of heat raises the temperature of the steam and increases the average kinetic energy of the gas molecules, as predicted by the Molecular Kinetic Theory. Strength of IMF The heat of fusion (heat required to melt a solid) and heat of vaporization (heat required to vaporize a liquid) are determined by the strength of the Intermolecular Forces. Substances with high IMF will have higher melting and boiling points. It will require more energy to break the IMF.

34 Most IMF are weaker than chemical bonds. To break the O-H bond in H 2O requires 935 kj/mol. To break the IMF in ice (heat of fusion) requires 6.02 kj/mol and to break the remaining IMF in the vaporization of water requires 40.7 kj/mol. All IMF are electrostatic in nature, the interaction of positive and negative charges. The strength of the IMF will, then, depend on the magnitude of these charges. Ionic bonding The strongest IMF is ionic bonding. These are the bonds between metals and non-metals, involving ions. Examples are NaCl and NH 4OH. Coulomb's law states that the potential energy, E is proportional to the amplitude of the charges, Q 1 and Q 2, divided by the distance squared, d 2. In salts, there are full positive charges on the cations, which have lost electrons, and full negative charges on the anions, which have gained electrons. One of the defining features of salts is their extremely high melting points. A large amount of energy is required to separate the positive and negative ions from their positions in the crystalline lattice. Dipole-dipole These are the interactions that exist between neutral, but polar substances. They involve the attraction of partial positive and partial negative charges present in polar compounds. We can compare the boiling points of two polar compounds, of similar size, chloromethane, CH 3Cl and cyanomethane, CH 3CN. b.p. = 249 K b.p. = 355 K These are both polar compounds, but the dipole moments, are different. CH 3Cl has a dipole moment of 2 D (debye, the units of polarity), and CH 3CN has a dipole moment of 3.9 D. This is because nitrogen is more electronegative than chlorine. This difference in the strength of the IMF leads to a difference in the boiling points of the compounds, CH 3Cl boils at 249 K and CH 3CN boils at 355 K. For similarly sized molecules, the larger the dipole moment ( ) the stronger the IMF. Hydrogen Bonding This is a special case of dipole-dipole interactions. The partial positive charge comes from a hydrogen atom bonded to F, O or N. The partial negative charge comes from a lone pair on O, N or F. Shown below are a number of possibilities for hydrogen bonding (- ---). Because O, N and F are so electronegative, their bonds with H are extremely polar. Because H is small, with no inner core electrons, it has a small, concentrated charge, which can approach the electronegative atom (O, N or F) very closely, making a very strong interaction. The boiling points of hydrogen bonding species are much higher than those of similarly sized molecules that don't exhibit hydrogen bonding, as shown in the plot below.

35 Notice that H 2O has a higher b.p. than HF, which is more polarized. This is because water is unique in being able to form four hydrogen bonds per molecule. Hydrogen bonding determines the structure of solid water (ice) making it less dense than liquid water. This means that ice floats, and lakes freeze from the top down. It is the only substance that is less dense as a solid than as a liquid. London Dispersion Forces These are the IMF that exist between non-polar atoms or molecules. Again, they are due to attractions between opposite charges. In this case, the attractive forces are due to instantaneous dipole moments. The picture below represents a "snapshot" of two helium atoms. At this particular instant, both valence electrons on the atom on the right are on the same side of the nucleus. This will create an instantaneous dipole moment in that helium atom. There will be an attraction between this area of negative charge and the nucleus of an adjacent helium atom. In addition, the electrons on the second atom will be repelled by the electrons on the first, and will also form an "induced" dipole in the second atom. The picture on the right shows the induced dipole moments and the attraction between them. These are the London Dispersion Forces. They are instantaneous electrostatic attractions. They are very weak, and only operate over very short distances. The strength of the London Dispersion Forces (LDF) depends on how easily the electron cloud of an atom or molecule can be distorted or polarized. The further from the nucleus that an electron exists, the more loosely it is held and the more polarizable it will be, leading to a stronger LDF. So, the strength of LDF increases with atomic and molecular size. A comparison of the boiling points of the noble gases demonstrates this effect. Substance atomic number b.p. ( o C) He (with added pressure) Ne Ar Kr Xe Rn The strength of the LDF also varies with the shape of molecules. LDF operate over very short distances. The closer molecules can approach each other, the stronger the LDF. We can compare two different compounds that both have the formula C 5H 10. One is the open chain molecule, pentane, and the other is the branched compound, neo-pentane, both shown below. pentane neo-pentane b.p. = 309 K b.p. = 282 K They possess exactly the same number of electrons and degree of polarizability but the have very different boiling points, indicating very different LDF. If we look at the space filling models, which show the volume occupied by the electrons in the molecules. The pentane molecule is linear in shape. The neo-pentane molecule is more spherical. Shown below are two molecules of each compound, at their closest approach to each other.

36 pentane b.p. = 309 K b.p. = 282 K The linear molecules can approach each other very closely, and along the whole length of the molecule. This increases the strength of the LDF by decreasing the distance between the molecules. The neo-pentane is an almost spherical molecule, that can only make contact over a very limited area. This increases the distance between the partial charges and lowers the strength of the LDF. All atoms and molecules possess LDF, since they all possess electrons. We can compare the boiling points of two polar molecules, HBr and HCl. HCl is more polar, with a larger dipole moment. However, its b.p.(189k) is lower than that of the less polar HBr (206 K). This is due to the LDF. HBr has 36e - and HCl, only 18 e -. Properties of Liquids Surface Tension Water will "bead-up" on a waxed surface. This is due to the "imbalance" of IMF at the surface of a liquid. In the interior of the liquid, the molecules are surrounded by other, similar molecules, but at the surface they are only attracted to the sides and inward. This works to reduce the surface area, attempting to create the shape with the minimal surface area, a sphere. The surface tension is the amount of energy required to stretch or increase the surface area. Molecules with high IMF also have high surface tensions. The surface tension of mercury is even higher, due to the metallic bonds ("sea" of shared electrons) between atoms. The intermolecular attraction between like molecules is called cohesion. The intermolecular attraction between unlike molecules is called adhesion. Both forces are in play when we look at a comparison of the meniscus shapes of water and mercury. Both liquids have strong cohesive forces (hydrogen bonding and metallic bonding) so both will have rounded shapes. However, water also has adhesive forces with glass. Glass is composed of Si and O. The oxygen in glass can hydrogen bond to water. So, the water level will rise near the glass, due to this attraction, giving a concave profile to the water at the surface. This is also the cause of capillary action, the rise of water in narrow tubes. Mercury has no adhesive forces with glass, so the surface minimizes the exposure to glass by rounding up, and is it not drawn up into the tube.