CHCA AP Chemistry Summer Assignment

Transcription

1 CHCA AP Chemistry Summer Assignment The summer assignment for AP Chemistry is quite simple (but not easy). You need to master the formulas, charges, and names of the common ions. On the first day of the school year, you will be given a quiz on these ions. You will be asked to: write the names of these ions when given the formula and charge write the formula and charge when given the names I have included several resources in this packet. First, there is a list of the ions that you must know on the first day. This list also has, on the back, some suggestions for making the process of memorization easier. For instance, many of you will remember that most of the monatomic ions have charges that are directly related to their placement on the periodic table. There are naming patterns that greatly simplify the learning of the polyatomic ions as well. Also included is a copy of the periodic table used in AP Chemistry. Notice that this is not the table used in first year chemistry. The AP table is the same that the College Board allows you to use on the AP Chemistry test. Notice that it has the symbols of the elements but not the written names. You need to take that fact into consideration when studying for the afore mentioned quiz! I have included a sheet of flashcards for the polyatomic ions that you must learn. I strongly suggest that you cut them out and begin memorizing them immediately. Use the hints on the common ions sheet to help you reduce the amount of memorizing that you must do. Do not let the fact that there are no flashcards for monatomic ions suggest to you that the monatomic ions are not important. They are every bit as important as the polyatomic ions. If you have trouble identifying the charge of monatomic ions (or the naming system) then I suggest that you make yourself some flashcards for those as well. Doubtless, there will be some students who will procrastinate and try to do all of this studying just before the start of school. Those students may even cram well enough to do well on the initial quiz. However, they will quickly forget the ions, and struggle every time that these formulas are used in lecture, homework, quizzes, tests and labs. All research on human memory shows us that frequent, short periods of study, spread over long periods of time will produce much greater retention than long periods of study of a short period of time. I could wait and throw these at you on the first day of school, but I don t think that would be fair to you. Use every modality possible as you try to learn these speak them, write them, visualize them. Additionally, if you haven t had chemistry since sophomore year (or if you had CP Chemistry), or you just want to review to get ahead, please download Chapters 1 3 Notes off of my website. Work through the problems (answers are provided for most). We will cover these 3 chapters in 3 weeks! I look forward to seeing you all at the beginning of the next school year. If you need to contact me during the summer, you can me and I will get back to you quickly. Most of you know that I have a website, mrgansle.site11.com, and there are some review activities in the AP Chemistry section that will give you an idea whether you have adequately mastered the summer assignment. Have a blessed summer, Mr. Gansle paul.gansle@chca oh.org

2 Common Ions and Their Charges A mastery of the common ions, their formulas and their charges, is essential to success in AP Chemistry. You are expected to know all of these ions on the first day of class, when I will give you a quiz on them. You will always be allowed a periodic table, which makes indentifying the ions on the left automatic. For tips on learning these ions, see the opposite side of this page. From the table: Ions to Memorize Cations Name Cations Name H + Hydrogen Ag + Silver Li + Lithium Zn 2+ Zinc Na + Sodium 2+ Hg 2 Mercury(I) K + Potassium + NH 4 Ammonium Rb + Rubidium Cs + Cesium Be 2+ Beryllium Anions Name Mg 2+ Magnesium NO 2 Nitrite Ca 2+ Calcium NO 3 Nitrate Ba 2+ Barium 2 SO 3 Sulfite Sr 2+ Strontium 2 SO 4 Sulfate Al 3+ Aluminum HSO 4 Hydrogen sulfate (bisulfate) OH Hydroxide Anions Name CN Cyanide H Hydride 3 PO 4 Phosphate F Fluoride 2 HPO 4 Hydrogen phosphate Cl Chloride H 2 PO 4 Dihydrogen phosphate Br Bromide NCS Thiocyanate I Iodide 2 CO 3 Carbonate O 2 Oxide HCO 3 Hydrogen carbonate (bicarbonate) S 2 Sulfide ClO Hypochlorite Se 2 Selenide ClO 2 Chlorite N 3 Nitride ClO 3 Chlorate P 3 Phosphide ClO 4 Perchlorate As 3 Arsenide BrO Hypobromite Type II Cations Name BrO 2 Bromite Fe 3+ Iron(III) BrO 3 Bromate Fe 2+ Iron(II) BrO 4 Perbromate Cu 2+ Copper(II) IO Hypoiodite Cu + Copper(I) IO 2 iodite Co 3+ Cobalt(III) IO 3 iodate Co 2+ Cobalt(II) IO 4 Periodate Sn 4+ Tin(IV) C 2 H 3 O 2 Acetate Sn 2+ Tin(II) MnO 4 Permanganate Pb 4+ Lead(IV) 2 Cr 2 O 7 Dichromate Pb 2+ Lead(II) 2 CrO 4 Chromate Hg 2+ Mercury(II) 2 O 2 Peroxide 2 C 2 O 4 Oxalate NH 2 Amide 3 BO 3 Borate 2 S 2 O 3 Thiosulfate

3 Tips for Learning the Ions From the Table These are ions can be organized into two groups. 1. Their place on the table suggests the charge on the ion, since the neutral atom gains or loses a predictable number of electrons in order to obtain a noble gas configuration. This was a focus in first year chemistry, so if you are unsure what this means, get help BEFORE the start of the year. a. All Group 1 Elements (alkali metals) lose one electron to form an ion with a 1+ charge b. All Group 2 Elements (alkaline earth metals) lose two electrons to form an ion with a 2+ charge c. Group 13 metals like aluminum lose three electrons to form an ion with a 3+ charge d. All Group 17 Elements (halogens) gain one electron to form an ion with a 1 charge e. All Group 16 nonmetals gain two electrons to form an ion with a 2 charge f. All Group 15 nonmetals gain three electrons to form an ion with a 3 charge Notice that cations keep their name (sodium ion, calcium ion) while anions get an ide ending (chloride ion, oxide ion). 2. Metals that can form more than one ion will have their positive charge denoted by a roman numeral in parenthesis immediately next to the name of the Polyatomic Anions Most of the work on memorization occurs with these ions, but there are a number of patterns that can greatly reduce the amount of memorizing that one must do. 1. ate anions have one more oxygen then the ite ion, but the same charge. If you memorize the ate ions, then you should be able to derive the formula for the ite ion and viceversa. a. sulfate is SO 4 2, so sulfite has the same charge but one less oxygen (SO 3 2 ) b. nitrate is NO 3, so nitrite has the same charge but one less oxygen (NO 2 ) 2. If you know that a sufate ion is SO 2 4 then to get the formula for hydrogen sulfate ion, you add a hydrogen ion to the front of the formula. Since a hydrogen ion has a 1+ charge, the net charge on the new ion is less negative by one. a. Example: 3 PO 4 2 HPO 4 H 2 PO 4 phosphate hydrogen phosphate dihydrogen phosphate 3. Learn the hypochlorite chlorite chlorate perchlorate series, and you also know the series containing iodite/iodate as well as bromite/bromate. a. The relationship between the ite and ate ion is predictable, as always. Learn one and you know the other. b. The prefix hypo means under or too little (think hypodermic, hypothermic or hypoglycemia ) i. Hypochlorite is under chlorite, meaning it has one less oxygen c. The prefix hyper means above or too much (think hyperkinetic ) i. the prefix per is derived from hyper so perchlorate (hyperchlorate) has one more oxygen than chlorate. d. Notice how this sequence increases in oxygen while retaining the same charge: ClO ClO 2 ClO 3 ClO 4 hypochlorite chlorite chlorate perchlorate

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5 Sulfite Sulfate Hydrogen sulfate Phosphate Dihydrogen Phosphate Hydrogen Phosphate Nitrite Nitrate Ammonium Thiocyanate Carbonate Hydrogen carbonate Borate Chromate Dichromate Permanganate Oxalate Amide Hydroxide Cyanide Acetate Peroxide Hypochlorite Chlorite Chlorate Perchlorate Thiosulfate

6 HSO 4 SO 4 2 SO 3 2 HPO 4 2 H 2 PO 4 PO 4 3 NH 4 + NO 3 NO 2 HCO 3 2 CO NCS 3 SCN Cr 2 O 7 2 CrO 4 2 BO 3 3 NH 2 C 2 O 4 2 MnO 4 C 2 H 3 O 2 CH 3 COO CN OH ClO 2 ClO O 2 2 S 2 O 3 2 ClO 4 ClO 3

7 South Pasadena AP Chemistry 1 Matter and Measurement 1. Matter Normally exists in 3 physical states: Liquid Fixed volume, Fluid; Takes on the shape of lower part of container; welldefined surface Solid Rigid Shape; very little volume change as temperature and pressure change Gas Volume expands to fill the container; volume varies according to temperature and pressure Kinetic Molecular Theory The idea that matter consists of molecules or atoms that are in constant, random motion. Kinetic Energy = Energy of motion; higher temperature = more motion Macroscopic seen with the eyes. Microscopic seen with a microscope Particulate or Submicroscopic Structures at the atomic level (what we think about) Mixtures Heterogeneous Mixture A mixture where the properties of the mixture vary throughout. (Like an Oatmeal cookie, the different components are visible) Homogeneous Mixture Also called a solution, where the components mix at a molecular level; different properties of the mixtures are unnoticeable. Purification The separation of a mixture into its components. (techniques: distillation, filtration, & chromatography) 2. Elements A substance that cannot be decomposed further by chemical means Names given by symbols: Example: Helium = He, Gold = Au, Aluminum = Al [Keep for Reference] B L U F F E R S G U I D E 3. Physical Properties Properties of a lone sample (ex. mass, volume, boiling temp, melting temp, conductivity, etc.) Density is the physical property that relates the mass of an object to its volume Density = Mass / Volume Extensive Property Properties, like mass and volume, that depend on the amount of substance Intensive Properties Properties like color and density; independent of the amount of substance Temperature how hot a substance is; physical properties (like density) vary with temp Celsius 0 C for freezing point of water and 100 C for melting point of water. Kelvin same scale as Celsius; 0 C = 273 K; 0 K = no motion; Celsius o = Kelvin 4. Chemical Properties How substance interacts with other substances. Ex. forms gas with acid; burns in air, etc. 5. Physical and Chemical Change Physical Change where the identity of all the substances remain unchanged (melting, boiling, grinding, pounding into sheets, etc.) Chemical Change (Reaction) atoms rearrange to convert one substance into another Chemical Equation A representation of the chemical reaction taking place Example: P 4 + 6Cl 2 4PCl 3 6. Measurements/Calculations Accuracy how close to a true value ; quantified by percent error. Precision how close measurements are to each other. Measured by significant figures or ± notation. [I assume you know metric system.] Dimensional Analysis use of a conversion factor to change units (ex: metric conversions, mass & volume, time units, etc.)

8 South Pasadena AP Chemistry 2 Atoms and Elements The Development of the Atomic Theory: I can: State the four signpost scientists, their experiments, what they added to the atomic theory, and the name of their model. Define the three theories that Dalton explained in terms of atoms: o Law of Conservation of Matter o Law of Definite/Constant Proportions o Law of Multiple Proportions Give examples and solve calculation problems related to each of the three theories. Sketch a cathode ray tube as demonstrated in class and state how J.J. Thomson s experiments led to the idea that atoms have positive and negative parts, the negative parts are all the same, and the negative parts (called electrons) have a certain charge/mass ratio. Define cathode rays. State the factors that determine how much a moving charged particle will be deflected by an electric or magnetic field. Explain Millikan s oil drop experiment & how it added to the atomic theory. Sketch the setup used by Ernest Rutherford (the goldfoil experiment), show what he observed, and explain how these observations led to the idea that most of the mass of the atom is concentrated into a tiny, amazingly massive, positivelycharged nucleus. Parts of the Atom: State the three particles that make up an atom, their symbol, their charge, their mass, and their location. State the number of protons, neutrons, and electrons in any atom or ion. Explain that isotopes are two atoms with the same atomic number (number of protons) but different mass numbers (number of nucleons protons + neturons). Represent the nucleus with isotopic notation, such as: Rn Recognize when two nuclei are isotopes of each other. Name Period Date / / S T U D Y L I S T Molar Mass Calculations: Calculate the isotopic mass of an atom given the resting mass of protons and neutrons. Explain that a mole of any element is actually made up of various isotopes in a constant percentage abundance. Calculate the average atomic mass of an element using the percent abundance and mass of each isotope. Calculate the percent abundance of isotopes given the average atomic mass and isotopic masses of an element. The Families of the Periodic Table: List the common families of the periodic table and recognize to which family any element belongs. Recognize metals, nonmetals, and metalloids (semimetals) on the periodic table. State and define the terms conductivity, malleability, ductility, and sectility. State some element facts such as which elements are too radioactive to exist, which is the largest nonradioactive element, which element has the greatest density, and which element has the highest melting point. Explain how Dmitri Mendeleev put together the periodic table and why we give him credit for the table even though others were working along the same lines. List the three elements that Mendeleev predicted and where they are located on the periodic table. A Little Nuclear Chemistry: State that Henri Becquerel discovered radioactivity and Marie Curie studied it. List the three Becquerel rays (alpha, beta, and gamma) and state why alpha particles were the perfect tool for Ernest Rutherford to study the structure of atoms. State that the alpha particle is the same as a helium nucleus, a beta particle is a highspeed electron, and a gamma ray is a highenergy form of light.

9 South Pasadena AP Chemistry 3 Chemical Formulas I can: Formulas Look at a formula and state how many elements and atoms are in that compound. Calculate the molecular mass or molar mass of any compound. State that the mass of a molecule is measured in amu s and the mass of a mole is measured in grams. Give examples of empirical formulas, molecular formulas, and structural formulas. Identify a formula as empirical, molecular, or structural. Ionic Compounds I can: List the names and formulas of 60 ions. State whether a compound is an ionic compound or a nonmetal compound. Write the formula of an ionic compound given the two ions or its name. Know when to use parentheses. Name an ionic compound given the formula. Determine the charge on an ion from information in an ionic formula. Nonmetal Compounds aka Molecular Compound Write the formula of a binary nonmetal compound (molecular compound) given its name. Name a binary nonmetal compound (molecular compound) given its formula. Name Period Date / / S T U D Y L I S T Percent Composition Calculate the percent composition (by mass) for any compound. Calculate the empirical formula from percent composition data. Determine the molecular formula of a compound given its empirical formula and molar mass. Hydrates Give examples of hydrates and anhydrous compounds. Calculate the formula of a hydrate from dehydration data. The Mole State the significance of the mole. State the three mole facts for any substance (molar volume, molar mass, Avogadro s number) 1 mole = 22.4 STP (gases only) 1 mole = 6.02 x particles (particles = molecules or atoms) 1 mole = gram molecular mass of chemical Use dimensional analysis to convert between moles, mass, volume, and number of particles for a chemical. Use density as a conversion factor in mole problems. Use gas density to calculate molar mass.