AP Chemsitry Summer Review Packet 2016 · There are some sources that will aid you in reviewing for the summer: 1. Check out the AP ... is to work with everyone ... 2016 AP Chemistry

2016 AP Chemistry Summer Review Packet

Congratulations on being a member of the AP Chemistry family! I hope that you are planning on having a fun filled summer but I also hope you are getting ready for a challenging year. AP Chemistry is an extremely fast paced course that covers a substantial amount of material. It is equivalent to a freshman college chemistry course with rigorous math, considerable independent work and a strong laboratory component. The end goal is to prepare everyone to take the AP chemistry exam on May 1, 2017. AP Chemistry is a difficult class, but with determination and perseverance, you will surely succeed! The AP Chemistry curriculum is extremely content-intensive and will require you to study every day, read and take notes on the sections that are going to be covered before lecture, know and understand key terms, form study groups outside school hours, and apply your knowledge to a whole new situation. There are some sources that will aid you in reviewing for the summer:

1. Check out the AP chemistry website. This website is formatted similar to last year. Spend some time going through the different parts especially the calendar and document pages. You can find the website at: http://tinyurl.com/NansenAPChem. Please go online and review the summer review packet. This is a set of concepts that you should be familiar with BEFORE coming to class the first day of school.

2. Purchasing the textbook. For the 2016-2017 school year I will be using 12th edition of Chemistry: The Central Science by Brown, Lemay, Bursten, and Murphy (ISBN: 0321696727). Amazon and Abebooks are a great place to find cheap textbooks. YOUR SUMMER ASSIGNMENT IS TO READ CHAPTERS 1-3 AND TAKE NOTES (IF NEEDED). My hope is that you start AP Chemistry with 80% of the knowledge that you left Chemistry with. Chapters 1-3 are review.

3. I would suggest that you purchase a study book as well. If you purchase a book make sure it’s the most recent edition that reflects the most current 2014 exam rewrite. A couple of good books are:

a. AP Chemistry a Crash Course b. Cracking the AP Chemistry Exam 2016 edition c. 5 Steps To A 5: AP Chemistry, Cracking the AP Chemistry Exam d. Barron's AP Chemistry

4. GET ORGANIZED. I would organize your papers, labs, notes, etc from last year so you have them ready for reference in case you need it. Also, get your notebooks, binders, etc ready to go. I am going to have everyone have a dedicated lab notebook so think about buying a composition notebook or notebook you can keep just for labs.

5. Last, e-mail me questions about problems in the packet, concerns, issues or fears that you have. I will be more than willing to work with you. My e-mail is [email protected]

Goal Setting: I firmly believe that setting a goal for this class. Additionally, I think it is wise to write it down somewhere that you see it everyday. Having a written goal provides you with the motivation to do well in the class. The goal could be short term or long term. Regardless, having this will help you succeed. My personal goal for this class is to work with everyone to try to get them to score a 3, 4 or 5 on the AP exam. I know the ways that I am going to help but this is a goal that I am going to try to accomplish. What is your goal? I am looking forward to working with you next year. Have a great summer and I will see you in the fall. Sincerely, Melissa Nansen


Things you should know/review for the class 1. Simple Topics

a. Know Symbols for Elements b. Rules for Significant Figures c. Diatomic Elements d. Common Polyatomic Ions and Their Charges e. Strong Acids and Bases and Strong Electrolytes f. Prefixes for Molecular Nomenclature g. Metric System Prefixes h. Base Units of the Metric System i. Solubility Rules j. Periodic Trends k. Basic pH Information l. Acid Nomenclature m. Electron Configurations Using Periodic Table n. Safety Rules o. Name and Proper Use of Laboratory Equipment. p. Assigning Oxidation Numbers

2. You should be able to write formulas and/or name compounds for:

a. Binary Ionic Compounds. b. Binary Molecular Compounds c. Ternary Ionic Compounds d. Acids and Bases

3. Be able to complete equations for reactions involving:

a. Combination b. Decomposition c. Single-replacement d. Metathesis (precipitation) e. Acid-base (neutralization) f. Combustion g. Net ionic and Redox reactions.

4. Use dimensional analysis to solve problems/Stoichiometric Calculations using balanced equation


SIMPLE TOPICS The Elements Know the name and symbol of elements. Know family names or groups. Find your Periodic Table that we colored last year. This is a great resource. You may even want to make a new copy

Significant Figures Rules for counting significant figures are summarized below. 1. Zeros within a number are always significant. Both 4308 and 40.05 contain four significant figures. 2. Zeros that do nothing but set the decimal point are not significant. Thus, 470,000 has two significant figures. 3. Trailing zeros that aren't needed to hold the decimal point are significant. For example, 4.00 has three significant figures. 4. Remember the Pacific and Atlantic Rule (present vs. absent)


Sig Fig continued: When measurements are added or subtracted, the answer can contain no more decimal places than the least accurate measurement. *FEWEST DECIMALS* 150.0 g H2O + 0.507 g salt 150.5 g solution When measurements are multiplied or divided, the answer can contain no more significant figures than the least accurate measurement. *FEWEST SIG FIGS* Example: To illustrate this rule, let's calculate the cost of the copper in an old penny that is pure copper. Let's assume that the penny has a mass of 2.531 grams, that it is essentially pure copper, and that the price of copper is 67 cents per pound. We can start by from grams to pounds. Solution: There are four significant figures in both the mass of the penny (2.531) and the number of grams in a pound (453.6). But there are only two significant figures in the price of copper, so the final answer can only have two significant figures so 0.37 cents. Know the 7 Diatomic Elements Remember: Start a 7 and make a 7 or BrINClHOF Hydrogen (H2); Nitrogen (N2); Oxygen (O2); Fluorine (F2); Chlorine (Cl2); Bromine (Br2); Iodine (I2) Know the name, symbol for, and charge of Common Polyatomic Ions The polyatomic ions are arranged in related groups according, mostly, to the periodic table. In memorizing these ions look for similarities with in families. The suffix -ate is the common ion in a group of similar ions, X-ite means one oxygen less, hypo-X-ite two oxygens less, per-X-ate one oxygen more. The prefix thio- usually means an oxygen atom has been replaced with a sulfur atom (not always but usually).


Positive Ions (Cations) 1+ Ammonium (NH4

+) Cesium (Cs+) Copper (I) or Cuprous (Cu+) Hydrogen (H+) Lithium (Li+) Potassium (K+) Silver (Ag+) Sodium (Na+) 2+ Barium (Ba2+) Cadmium (Cd2+) Calcium (Ca2+) Chromium (II) or Chromous (Cr2+) Cobalt (II) or Cobaltous (Co2+) Copper (II) or Cupric (Cu2+) Iron (II) or Ferrous (Fe2+) Lead (II) or Plumbous (Pb2+) Magnesium (Mg2+) Manganese (II) or Manganous (Mn2+) Mercury (I) or Mercurous (Hg2

2+)

Mercury (II) or Mercuric (Hg2+) Strontium (Sr2+) Nickel (II) (Ni2+) Tin (II) or Stannous (Sn2+) Zinc (Zn2+) 3+ Aluminum (Al3+) Chromium (III) or Chromic (Cr3+) Iron (III) or Ferric (Fe3+) Negative Ions (Anions) 1- Acetate (C2H3O2-) Bromide (Br-) Chlorate (ClO3

-) Chloride (Cl-) Cyanide (CN-) Dihydrogen phosphate (H2PO4

-) Fluoride (F-) Hydride (H-) Hydrogen carbonate or bicarbonate (HCO3

-) Hydrogen sulfite or bisulfite (HSO3

-)

Hydroxide (OH-) Iodide (I-) Nitrate (NO3

-) Nitrite (NO2

-) Perchlorate (ClO4

-) Permanganate (MnO4

-) Thiocyanate (SCN-) 2- Carbonate (CO3

2-) Chromate (CrO4

2-) Dichromate (Cr2O7

2-) Hydrogen phosphate (HPO4

2-) Oxide (O2-) Peroxide (O2

2-) Sulfate (SO4

2-) Sulfite (SO4

2-) Sulfide (S2-) 3- Arsenate (AsO4

3-) Phosphate (PO4

3-)

Know the Strong Acids and Bases The 7 Strong Acids

• HNO3 nitric acid • H2SO4 sulfuric acid • HBr hydrobromic acid • HI hydroiodic acid • HCl hydrochloric acid • HClO3 chloric acid • HClO4 perchloric acid

The 8 Strong Bases

• LiOH lithium hydroxide • NaOH sodium hydroxide • KOH potassium hydroxide • RbOH rubidium hydroxide • CsOH cesium hydroxide • Ca(OH)2 calcium hydroxide • Sr(OH)2 strontium hydroxide • Ba(OH)2 barium hydroxide

Strong Electrolytes: Strong electrolytes consist of substances that ionize completely (approx. 100%) in water. These substances are strong acids, strong bases, and soluble ionic salts. Weak Electrolytes: Substances that ionize approx. 1-5% in water form weak electrolytes. These substances are weak acids, weak bases, and slightly soluble ionic salts. Non-Electrolytes: Substances that dissolve in water but are considered polar covalent, like table sugar and some alcohols, are considered non- electrolytes.


Solubility Rules 1 All compounds containing alkali metal cations and the ammonium ion are soluble. 2 All compounds containing NO3

-, ClO4-, ClO3

-, and C2H3O2- anions are soluble.

3 All chlorides, bromides, and iodides are soluble except those containing Ag+, Pb2+, or Hg22+.

4 All sulfates are soluble except those containing Hg22+, Pb2+, Sr2+, Ca2+, or Ba2+.

5 All hydroxides are insoluble except compounds of the alkali metals, Ca2+, Sr2+, and Ba2+. 6 All compounds containing PO4

3-, S2-, CO32-, and SO3

2- ions are insoluble except those that also contain alkali metals or NH4

+. Chemical Naming Binary Covalent: compound composed of two nonmetals.

Such a compound is named by using a Greek prefix designating the number of atoms for the elements in the formula. Note that the Greek prefix "mono-" is not used with the first element, just the second. Also, end the name of the second element in "ide".

Ionic Naming: Metal and Nonmetal

Name the cation first and indicate charge if necessary (transition metals) Then name the anion, noting that nonmetals change their ending to –ide Common ions keep ending –ate or –ite

Acid Naming: Hydrogen and anion

Binary: Since all these acids have the same cation, H+, we don't need to name the cation. The acid name comes from the root name of the anion name. The prefix hydro- and the suffix -ic are then added to the root name of the anion.

• HCl, which contains the anion chloride, is called hydrochloric acid. • HCN, which contains the anion cyanide, is called hydrocyanic acid.

Ternary or Oxyacids (anion contains the element oxygen): Since all these acids have the same cation, H+, we don't need to name the cation. The acid name comes from the root name of the oxyanion name or the central element of the oxyanion. Suffixes are used based on the ending of the original name of the oxyanion. If the name of the polyatomic anion ended with -ate, change it to -ic for the acid and if it ended with -ite, change it to -ous in the acid.

• HNO3, which contains the polyatomic ion nitrate, is called nitric acid. • HNO2, which contains the polyatomic ion nitrite, is called nitrous acid.


Periodic Trends Atomic radius The atomic radius is the distance from the atomic nucleus to the outermost stable electron orbital in an atom that is at equilibrium. The atomic radius tends to decrease as one progresses across a period because the effective nuclear charge increases, thereby attracting the orbiting electrons and lessening the radius. The atomic radius usually increases while going down a group due to the addition of a new energy level (shell). However, diagonally, the number of protons has a larger effect than the sizeable radius. For example, lithium (145 pm) has a smaller atomic radius than magnesium (150 pm). Atomic radii decrease left to right across a period, and also increase top to bottom down a group. Ionization potential The ionization potential (or the ionization energy) is the minimum energy required to remove one electron from each atom in a mole of atoms in the gaseous state. The first ionization energy is the energy required to remove one, the nth ionization energy is the energy required to remove the atom's nth electron, after the (n−1) electrons before it have been removed. Trend-wise, ionization potentials tend to increase while one progresses across a period because the greater number of protons (higher nuclear charge) attract the orbiting electrons more strongly, thereby increasing the energy required to remove one of the electrons. There will be an increase of ionization energy from left to right of a given period and a decrease from top to bottom. Electron affinity The electron affinity of an atom can be described either as the energy gained by an atom when an electron is added to it, or conversely as the energy required to detach an electron from a singly-charged anion. As one progresses from left to right across a period, the electron affinity will increase, due to the larger attraction from the nucleus, and the atom "wanting" the electron more as it reaches maximum stability. Down a group, the electron affinity decreases because of a large increase in the atomic radius, electron-electron repulsion and the shielding effect of inner electrons against the valence electrons of the atom. Electronegativity Electronegativity is a measure of the ability of an atom or molecule to attract pairs of electrons in the context of a chemical bond. The type of bond formed is largely determined by the difference in electronegativity between the atoms involved, using the Pauling scale. Trend-wise, as one moves from left to right across a period in the periodic table, the electronegativity increases due to the stronger attraction that the atoms obtain as the nuclear charge increases. Moving down a group, the electronegativity decreases due to the longer distance between the nucleus and the valence electron shell, thereby decreasing the attraction, making the atom have less of an attraction for electrons or protons. Metallic character Metallic character refers to the chemical properties associated with elements classified as metals. These properties, which arise from the element's ability to lose electrons, are: the displacement of hydrogen from dilute acids; the formation of basic oxides; the formation of ionic chlorides; and their reduction reaction, as in the thermite process. As one moves across a period from left to right in the periodic table, the metallic character decreases, as the atoms are more likely to gain electrons to fill their valence shell rather than to lose them to remove the shell. Down a group, the metallic character increases, due to the lesser attraction from the nucleus to the valence electrons (in turn due to the atomic radius), thereby allowing easier loss of the electrons or protons.


The following diagram represents the trend in the periodic table for atomic size).

The following diagram represents the trend for ionic radii


The following represents the trend in ionization energy.


The following represents trends between Metals, Metalloids and Non-Metals


Electron Configurations Using the Periodic Table.


The primary determinant of an element's chemical properties is its electron configuration, particularly the valence shell electrons. For instance, any atoms with four valence electrons occupying p orbitals will exhibit some similarity. The type of orbital in which the atom's outermost electrons reside determines the "block" to which it belongs. The number of valence shell electrons determines the family, or group, to which the element belongs. The total number of electron shells an atom has determines the period to which it belongs. Each shell is divided into different subshells, which as atomic number increases are filled in roughly this order (the Aufbau principle): Subshell Order: S, F, D, P Period 1 1s

2 2s 2p 3 3s 3p 4 4s 3d 4p 5 5s 4d 5p 6 6s 4f 5d 6p 7 7s 5f 6d 7p

Hence the structure of the table. Since the outermost electrons determine chemical properties, those with the same number of valence electrons are grouped together.


Rules for Assigning Oxidation Numbers 1. The convention is that the cation is written first in a formula, followed by the anion. For example, in NaH, the H is H-; in HCl, the H is H+. 2. The oxidation number of a free element is always 0. The atoms in He and N2, for example, have oxidation numbers of 0. 3. The oxidation number of a monatomic ion equals the charge of the ion. For example, the oxidation number of Na+ is +1; the oxidation number of N3- is -3. 4. The usual oxidation number of hydrogen is +1. The oxidation number of hydrogen is -1 in compounds containing elements that are less electronegative than hydrogen, as in CaH2. 5. The oxidation number of oxygen in compounds is usually -2. Exceptions include OF2, since F is more electronegative than O, and BaO2, due to the structure of the peroxide ion, which is [O-O]2-. 6. The oxidation number of an alkali metal element in a compound is +1. 7. The oxidation number of an alkaline earth metal element in a compound is +2. 8. The oxidation number of a halogen element in a compound is -1, except when that element is combined with one having a higher electronegativity. Fluorine is always -1. The oxidation number of Cl is -1 in HCl, but the oxidation number of Cl is +1 in HOCl. 9. The sum of the oxidation numbers of all of the atoms in a neutral compound is 0. 10. The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion. For example, the sum of the oxidation numbers for SO4

2- is -2.

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AP Chemsitry Summer Review Packet 2016 · There are some sources that will aid you in reviewing for the summer: 1. Check out the AP ... is to work with everyone ... 2016 AP Chemistry