and so it begins n.
Skip this Video
Loading SlideShow in 5 Seconds..
And so it begins… PowerPoint Presentation
Download Presentation
And so it begins…

And so it begins…

111 Views Download Presentation
Download Presentation

And so it begins…

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

  1. And so it begins… The goal of AP Chemistry is to pass the AP exam on May 6th. Everything we cover throughout the year could potentially be on the exam and some things will definitely be on the exam. Read the relevant sections in the book beforehand! These PowerPoint notes are organized in as logical an order as possible, as detailed as possible, and as easy to follow as possible. However, you will be very well-served by underlining or highlighting and writing extra notes in the margin. Also, do NOT throw any of these notes away until after the exam and even then they might be helpful in college as a reference (even if you AP out of freshman chemistry, there is a lot of stuff in this course that can help in other college chemistry classes). Most topics in this class are related to each other which helps when learning new material and reviewing. But it also means that if you get behind, it can be difficult to catch back up. So if at any time you feel that you may be getting lost, LET ME KNOW AS SOON AS POSSIBLE!! AP chemistry can be a little difficult sometimes but it can also be a lot of fun. READ THE BOOK!!

  2. Unit 1 – Measurement and Chemical Nomenclature Chapter 1 – everything except Section 1.9 AND except converting between oF and oC but converting oC to K and back is required Chapter 2 – Mostly good but you do not need to know historical details. Chapter 22 – pages 997-1003, 1005-1006, 1010-1015 – Organic nomenclature, isomers, and functional groups are important but NOT organic reactions.

  3. Physical Properties and Experimental Design • extensiveproperty – _____________________ on amount of substance present • ex. • intensiveproperty – _____________________ on amount of substance present • ex. • independent variable – what is ________________ • dependent variable – what is _________________ • control – standard to which the results are compared

  4. Temperature (pg 19) Temperature is a measure of the kinetic energy (energy of motion) of particles. Particles that are moving faster have more energy and thus higher temperature. Kelvin is the absolute temperature scale meaning that it is NEVER negative. To convert from degrees Celsius to Kelvin, add 273. K = oC + 273 On the multiple choice section of the AP exam, there is often a question or two that has answer choices that use oC when it should be in K. That is a good way to eliminate obviously wrong answer choices. 0 K is _________________________ which is the level of energy where particle motion is at its minimum (particles never completely stop moving). It can’t move less than the minimum so that’s as low as the temperature can reach.

  5. Density and Avogadro’s Number(pgs 24, 82) density – ratio of mass to volume of a substance units are any mass over any volume g/mL, g/cm3, kg/L, g/L, lb/gal, lb/ft3, etc. Avogadro’s number – number of things in a mole 6.022 x 1023 mol-1 mol-1 means 1/mol. The “1” in 1/mol is a placeholder for the “things”. For example, there are 6.022 x 1023 atoms in one mole of atoms. This can also be stated as 6.022 x 1023 atoms/mol or atoms mol-1.

  6. Change In chemistry and most other sciences, a change in a value is denoted by the Greek letter __________, D. In all cases, change isthe final value minus the initial value. We use this a whole lot in AP. Dsomething = somethingf – somethingi where f is final and i is initial The temperature of a beaker of water goes from 45 oC to 30. oC. What is the value of DT?

  7. Units (pg 9)

  8. Units (pg 9)

  9. Units (pg 9) There are 1012 grams in 1 teragram (1012 g = 1 Tg). There are 103 meters in 1 kilometer (103 m = 1 km). According to the tables, there are 10-9 meters in 1 nanometer (10-9 m = 1 nm). Since it may be easier to use positive exponents, you can flip that over and around: There are 109 nanometers in 1 meter (109 nm = 1 m). kilo, milli, nano, centi, and occasionally micro are used most often on the AP exam and are the only ones that you have to know.

  10. Significant Figures (pg 13-16) Significant figures are used to report data as accurately as possible. For example, if your bathroom scale says that you weigh 145 lbs, you could not say that you weigh 145.348191 lbs because the scale is not that accurate – it does not weigh to the millionth of a pound. On the free response questions on the AP exam (and on assignments in this class), you must be within one significant figure - high or low - of the correct number. For example, if the answer should have three sig figs, your response can have either two, three, or four. Any more or less than the minimum results in a one-time one-point deduction for that entire question. If you are in doubt as to how many sig figs to use, three sig figs are usually good enough because there will rarely be more than four or fewer than two sig figs in a problem (bag of tricks).

  11. Measuring Significant Figures(pg 13-16) When taking measurements in lab, the correct number of significant figures is all numbers you know for certain plus one estimate. For example, you know that the gray thing below is between ______ and ______ cm long. Those are the numbers you know for certain. The estimate is the next number. Call it _____ cm. This ruler can measure to ______ sig figs.

  12. Significant Figures Rules (pg 13-16) • All nonzero digits are significant. • Zeroes at the right end of a number AND to the right of the decimal are significant. • ex. 23.80, 0.060 • Zeroes between significant figures are significant. • ex. 409, 90050, 30.0 • In scientific notation, the rules above apply to everything except the “x 10n” term. • ex. 4.0 x 104, 2.03 x 10-5 • Counting numbers and conversion factors are exact and thus have infinite significant figures. • 30 people in the room, a carbon atom has 6 protons, 1000 L = 264.17 gal, 1 inch = 2.54 cm

  13. Adding and Subtracting Sig Figs (pg 13-16) Three steps: • With one number on top of the other, line up the decimal points. • Add or subtract the numbers. • Start from the first column with both digits being significant and go to the right. The last sig fig for the answer is the last column in which all digits in the column are significant.

  14. Adding and Subtracting Sig Figs (pg 13-16) 1020 – 437 0.00310 + 0.052

  15. Multiplying and Dividing Sig Figs (pg 13-16) • Multiply or divide the numbers. • The answer has the same number of sig figs as the number with the fewest number of sig figs. 302 x 40 = 25 / 8740 =

  16. Rounding For long problems, do not round until the final answer – it’s best to keep all numbers in your calculator or at least write more digits than you will need then round at the very end. For example, the following calculation is done by keeping the extra digits then again by rounding to sig fig every step of the way. 23 x 2 = 46 46 / 709 = 0.06488 = 0.06 23 x 2 = 46 = 50 50 / 709 = 0.0705 = 0.07 The answer has one sig fig because of the 2 but note that 0.06 is not the same as 0.07. The point is that you should not round or worry about sig figs until the final answer.

  17. Unit Conversions (pg 16-19) 1000 L = 264.17 gal Convert 55 gallons to liters. Convert 709 nanometers to meters

  18. Unit Conversions (pg 16-19) 1000L = 264.17 gal, 1 kg = 2.20 lb Convert 13.53 g/mL to lb/gal.

  19. Dimensional Analysis You may not know how to do a problem but if you know what units that the answer has, then you may be able to get the right answer anyway. For example, what is the density of lead if 3.14 cubic centimeters contains 0.172 mol of lead?

  20. Dimensional Analysis A single photon capable of breaking a Cl-Cl bond needs a minimum of 4.018 x 10-19 J of energy. Calculate the energy, in kJ mol-1, of the Cl-Cl bond.

  21. Diatomic Molecules Several nonmetals in their pure forms exist as molecules of two atoms. For example, hydrogen is H2 and bromine is Br2. They do NOT exist naturally as H and Br by themselves. Molecules like this are diatomic because they have _________________. Remember the following word and capitalization to help you out: The formulas for these seven elements MUST be diatomic unless the problem says something like “atomic chlorine” which is just one chlorine, Cl, or if they are part of a compound such as NaCl. Anything else is Cl2.

  22. Next up… The last part of this first unit is naming compounds. We will cover how to name all compounds from ionic and covalent to acids, hydrates, and finally organic compounds. It is a lot all at once but you should be familiar with this from past years (with the possible exception of organic nomenclature). We will be going through it quickly but it is VERY important that you fully understand ALL of it since it will show up on pretty much every topic from here on out. There is a fair amount of memorization but there are patterns that will help – more on that later.

  23. Ionic Compounds (pg 53-55) An ionic compound is one in which valence electrons are transferred from one atom (forming a positively charged ion) to another atom (forming a negatively charged ion) and the oppositely charged ions form a bond. positive ion: negative ion: Ionic compounds are usually metal-nonmetal but they don’t have to be. You must memorize the 66 ions (name, symbols, and charges) on the “Stuff I Should Know…” sheet. The following table is helpful for a few of them and you have to know the group names also. The provided flashcards have only 60 of the ions. The other six are bisulfide, bisulfite, chromium, nitride, oxalate, and phosphide. These six will appear on the third ion quiz.

  24. Ion Charges (pg 55)

  25. Ions Also to help with memorizing the ions, many are related to each other by a prefix (bi-, hypo-, per-) or a suffix (-ite, -ate). The prefix bi- means an H+ is added to the front of the ion. For example, carbonate is __________ so bicarbonate is __________. Note that the charge of bicarbonate is __________ because of the +1 coming from the H+. Sulfide is __________ so bisulfide is __________. Note also that “bi” does NOT mean that there are two of something. That’s “di” which we will get to later.

  26. Ions For the polyatomic ions that contain a halogen, the prefixes hypo- and per- go along with the suffixes -ite and -ate and all have to do with the number of oxygens. All have a -1 charge regardless of the number of oxygens. • hypo-something-ite has one oxygen • something-ite has two oxygens • something-ate has three oxygens • per-something-ate has four oxygens So if you know one of them then you know the rest. For example, bromate is ______ so write the formulas for the rest of the ions below: bromite hypobromite perbromate You can use these hints or you can straight memorize the ions. It is completely up to you.

  27. Forming Ionic Compounds(pg 57–67) Cross the numbers of the charges to make subscripts. The cation is always first and the anion is always second. Ca Cl Fe3+ O S Mn

  28. Forming Ionic Compounds(pg 57–67) Subscripts that come in with a polyatomic ion DO NOT CHANGE! You may have to put additional subscripts outside of parentheses. Na CO3 Al IO3 NH4 SO4

  29. Naming Ionic Compounds(pg 57–67) Naming a binary ionic compound (contains only two elements) that does NOT contain Fe, Cu, Sn, or Hg is three steps: • name of the cation • name of the anion • change the end of the anion to “ide” NaCl Mg3N2 ZnF2

  30. Naming Ionic Compounds(pg 57–67) Naming an ionic compound with a polyatomic ion is the easiest of them all. Nothing changes.* • name of the cation • name of the polyatomic ion Ca(NO3)2 NaHSO3 K2CO3 *The only exception are compounds containing ammonium, NH4+. Since ammonium is a cation, you just write the word “ammonium” then the anion follows the rules of binary ionic compounds. For example, NH4Cl is ___________________________.

  31. Naming Ionic Compounds (pg 57–67) For compounds containing Cu, Fe, Sn, or Hg, the same rules apply but since each of these metals can have different charges (which you must memorize) simply stating “iron oxide” is not good enough because there are two possible iron oxide compounds: FeO or Fe2O3. We distinguish between the two using a Roman numeral that is the charge of the metal. In FeO, Fe is ____ so FeO is In Fe2O3, Fe is ____ so Fe2O3 is In Sn(NO3)4, Sn is ____ so Sn(NO3)4 is

  32. Naming Ionic Compounds (pg 57–67) Another fairly common method of naming compounds with Cu, Fe, Sn, or Hg is by using the Latin names. The ion with the lowest charge ends in “ous” and the ion with the highest charge ends in “ic”. copper: iron: tin: mercury: Mercury is a little weird obviously. A good way to remember copper and iron is the phrase, “Cu one and two, Fe two and three.”

  33. Naming Ionic Compounds (pg 57–67) Write the names of the following: Al2(CO3)3 HgO Ag2S CuSO4 Write the formulas of the following: stannous chloride potassium dichromate ammonium oxalate calcium perchlorate iron(III) sulfite

  34. Naming Covalent Compounds (pg 57–67) A covalent compound (also called a molecule) is one in which electrons are shared between atoms. These involve nonmetals or metalloids since metals do not bond covalently. Naming molecules involves prefixes which correspond to the number of atoms of a particular element. Naming molecules is three steps: • use prefix on first element name (note that “mono” is never used on the first element) • use prefix on second element • change end of second element to “ide”

  35. Naming Covalent Compounds (pg 57–67) N2O5 CO2 carbon monoxide tetraphosphorus decasulfide sulfur hexalflouride

  36. Naming Acids (pg 57–67) A binary acid contains hydrogen and one other element. These three steps are also used for HCN even though it is not a binary acid: • use the prefix “hydro” • state the name of the other element (“hydro” and the element should be one word) • change the end of the element to “ic” and add “acid” HI HCl HBr

  37. Naming Acids (pg 57–67) An oxyacid contains oxygen usually in the form of a polyatomic ion. Do NOT use hydro for these. • If the polyatomic ion ends in “ite”, change “ite” to “ous” and add “acid”. • If the polyatomic ion ends in “ate”, change “ate” to “ic” and add “acid”. HNO3phosphorous acid periodic acid

  38. Hydrates A hydrate is a compound (usually ionic) with water molecules bonded to it. To denote how many water molecules, use the covalent prefixes. CuSO4•5H2O Pb(NO3)2•9H2O

  39. Organic Chemistry(pg 997-1003, 1005) We will cover the basics of organic nomenclature. alkane – all single bonds ex. methane, ethane, propane, octane alkene – contains double bond ex. ethene (ethylene), propene alkyne – contains triple bond ex. ethyne (acetylene)

  40. Saturated and Unsaturated (pg 997-1003, 1005) hydrocarbon – containshydrogen and carbon only carbohydrate – contains hydrogen, carbon, and oxygen only Alkanes are saturated hydrocarbons since they contain the maximum number of hydrogens. Alkenes and alkynes are unsaturated hydrocarbons since they can react with hydrogen gas to become saturated.

  41. Organic Nomenclature (pg 997-1003, 1005) These prefixes correspond to the number of carbons in the longest chain of carbons. meth 1 eth 2 prop 3 but 4 pent 5 hex 6 hept 7 oct 8 non 9 dec 10

  42. Formulas(pg 997-1003, 1005) All alkanes (methane, ethane, propane, etc.) use the following mathematical formula to figure out the chemical formula. CnH2n+2 ex. octane

  43. Formulas(pg 997-1003, 1005) All simple alkenes (only one double bond) use the following mathematical formula to figure out the chemical formula. CnH2n ex. pentene

  44. Formulas(pg 997-1003, 1005) All simple alkynes (only one triple bond) use the following mathematical formula to figure out the chemical formula. CnH2n-2 ex. hexyne

  45. Functional Groups(pg 1010-1015) Functional groups are special groups (kind of like polyatomic ions) on organic molecules. The eight most common are on the “Stuff I Should Know…” sheet. “R” is simply an abbreviation for the rest of the molecule that is not the functional group. Be able to recognize these functional groups given the structure or formula. An alcohol is simply –OH attached to a carbon instead of –H. The alcohol above can be generically written as R–OH where R is everything except the OH. To name an alcohol, add “ol” to the end of the formula name. The hydrocarbon above is ___________ so the alcohol is ____________.

  46. Functional Groups (pg 1010-1015) An aldehyde is COH in which both the oxygen and hydrogen are attached to the carbon and the oxygen is double bonded to the carbon. R can be anything else that is not the aldehyde. For example, consider the following: R

  47. Functional Groups (pg 1010-1015) A ketone is a double bonded oxygen, =O, attached to a carbon that is between two other carbons. To name a ketone, add “one” (as in phone) to the end of the formula name. The hydrocarbon above is ____________ so the ketone is ______________.

  48. Functional Groups (pg 1010-1015) An ether is an oxygen in a chain of carbons. R – O – R The formula on the right is _____________________ which appears every so often on the AP exam. We’ll get to how to name it in a little bit.

  49. Functional Groups (pg 1010-1015) Carboxylic acid is –COOH replacing a carbon and everything attached to it. Note that the hydrogen in the carboxylic acid group is attached to the single bonded oxygen. A carboxylic acid is just like any other acid – it donates the hydrogen and lowers the pH of the solution. Generically a carboxylic acid can be written as the following: R R

  50. Functional Groups (pg 1010-1015) An ester is similar to a carboxylic acid except that the ester is within a chain of carbons.