U0.LP0 - Summer Assignment - DUE August 16, 2019  

 

Name (First, Last) ________________________________________________________________ 

 

 

Welcome to AP Chemistry! We will hit the ground running, starting the first day of school and not letting up until the exam in mid-May. This assignment is paper-based (you do not need a computer to complete it) and is review material from general chemistry. It is due on the first day of school. If at any point you need help on this assignment, please feel free to contact me by email (jmason@striveprep.org) or by phone (619-913-0022). I will do my best to respond in a timely manner. Please use the checklist below to keep track of the items you must complete this summer. ❏ Familiarize yourself with the periodic table. You should be able to quickly recognize the symbols

(be able to name them) and their positions on the periodic table of the first 20 elements. There is a periodic table on the second page for your reference.

❏ Memorize the attached list of sixty familiar ions. You should be able to write the symbol of the

ion when given the name, and vice versa. Included in this packet are (1) naming patterns to help your memorization (2) a list of ions by name (3) ion flashcards that you should cut out and practice with (4) practice ion quizzes

● You should be able to complete the ions quizzes in less that 4 minutes (basically, you need to know then better than your multiplication tables).

● These are very similar to the ion quiz you will have in the first week of school.

❏ Complete the review problems. Please show your work when possible and write neatly. You should be familiar with this material, but the assignment may still take a while (especially if you haven’t done any chemistry in a while). Do not put this off until the last minute -- the more you review this summer, the more prepared you will be to succeed in AP Chemistry.

I am looking forward to starting the school year strong. Have a wonderful summer break.

Ion Memorization Naming Patterns Almost all of the polyatomic ions follow naming patterns. Once you learn those, you will only need to memorize a few outliers that don’t follow any rules. If you have internet access, I also highly recommend watching some youtube videos on memorizing polyatomic ions, there are some really great ones out there. Below are some guidelines that will help you start naming polyatomic ions.

Rule Explanation Example

The suffix (ending) “-ide” is the smallest.

An ion with the suffix “-ide” has no oxygens.

Bromide is Br-

The prefix (beginning) “hypo-” is smaller than “-ite”

An ion with the prefix “hypo-” will have one less oxygen than its counterpart with an “-ite” suffix.

Hypobromite is BrO- , while Bromite is BrO2-

The suffix “-ate” is bigger that “-ite”

An ion with the suffix “-ate” will have one more oxygen than its counterpart with an “-ite” ending

Bromate is BrO3- , while Bromite is BrO2-

The prefix “per-” is bigger than the “-ate” suffix

An ion with the prefix “-per” will have one more oxygen than its counterpart with an “-ate” suffix.

Perbromate is BrO4- , while Bromate is BrO3-

Now, all you have to memorize is how many oxygens bond with each element at the “-ate” level. From the “-ate” ion, you can move up or down depending on the suffix or prefix you are given. For example, if you need to know the ion hypochlorite, and you have chlorate (ClO3

-) memorized, all you need to do is remove two oxygens (remove one for “-ite”, and another for “hypo-”) to get ClO-.

But, say you reallllllly hate memorizing. Don’t worry, the internet has many tricks up its sleeve! If you’re having trouble memorizing how many oxygens go with each “-ate” ions, you can use the sentence below to help.

Nick the Camel ate Clam Supper in Phoenix with Brad

Nitrate Carbonate Chlorate Sulfate Phosphate Bromate

NO3- CO3

2- ClO3- SO4

2- PO43- BrO4

-

1. Each underlined letter represents the first element in the polyatomic ion. 2. The number of consonants (non-vowels) equals the number of oxygen atoms in the ion. 3. The number of vowels (a,e,i,o,u) equals the number of negative charges on the ions. Sometimes, ions will have hydrogens bonded to them as well. Hydrogens add one positive charge because they have a +1 charge (H+). For example, Carbonate (CO3

2-) becomes Hydrogen Carbonate (HCO3-).

Last, some metals don’t bond with oxygens often, and instead form different ions. Their names are easy, because the name includes the charge in roman numerals. For example, Copper (II) is Cu2+ and Lead (IV) is Pb4+. The only confusing aspect of metal-only ions is that their names don’t always match their symbol (Lead and Pb, or Tin and Sn). This is because the symbols are based off of their latin names, which are noted on the flashcards on the next pages. Last Tips 1. Cut out and USE the flashcards. Seriously. They will help. 2. Start by sorting the ions into similar groups, this will help you recognize them faster. Once you have them down, reshuffle into one large pile and work through those. 3. Make sure you practice going from name to formula AND formula to name. You will need to do both by the first day of school and also on the AP exam.

Acetate

Aluminum

Ammonium

Barium

Bromate

Bromide

Bromite

Calcium

Carbonate

Chlorate

Chloride

Chlorite

Chromate

Copper (I) or

Cuprous

Copper (II) or

Cupric

Cyanide

Dichromate

Fluoride

Hydrogen

Hydrogen Carbonate or

Bicarbonate

Ba2+

NH4+

Al3+

C2H3O2-

or CH3COO-

Ca2+

BrO2-

Br-

BrO3-

ClO2-

Cl-

ClO3-

CO32-

CN-

Cu2+

Cu+

CrO42-

HCO3-

H+

F-

Cr2O72-

Hydrogen Sulfate or

Bisulfate

Hydronium

Hydroxide

Hydrobromide

Hypochlorite

Hypoiodite

Iodate

Iodide

Iodite

Iron (II) or

Ferrous

Iron (III) or

Ferric

Lead (II) or

Plumbous

Lead (IV) or

Plumbic

Lithium

Magnesium

Mercury (I) or

Mercurous

Mercury (II) or

Mercuric

Nickel

Nitrate

Nitrite

BrO- or

OBr-

OH-

H3O+

HSO4-

I-

IO3-

IO- or OI-

ClO- or

OCl-

Pb2+

Fe3+

Fe2+

IO2-

Hg22+

Mg2+

Li+

Pb4+

NO2-

NO3-

Ni2+

Hg2+

Oxide

Perbromate

Perchlorate

Periodate

Permanganate

Peroxide

Phosphate

Phosphite

Potassium

Silver

Sodium

Strontium

Sulfate

Sulfide

Sulfite

Thiocyanate

Thiosulfate

Tin (II) or

Stannous

Tin (IV) or

Stannic

Zinc

IO4-

ClO4-

BrO4-

O2-

PO33-

PO43-

O22-

MnO4-

Sr2+

Na+

Ag+

K+

SCN-

SO32-

S2-

SO4

2-

Zn2+

Sn4+

Sn2+

S2O32-

Ion Naming Practice Quizzes

Try to do the “quizzes” below without looking at your flashcards or notes. These are very similar to the ion naming quiz you will have within the first week of classes!

Name Formula

Chlorate

Fluoride

Iron (III)

Strontium

Ammonium

Bisulfate

Hydrogen Carbonate

Nitrite

Nickel

Iodite

Mercury (I)

Chromate

Sulfate

Permanganate

Perbromate

Strontium

Sulfate

Thiosulfate

Name Formula

BrO4-

Ag+

S2O32-

SCN-

IO4-

OCl-

BrO-

Fe2+

NO2-

HSO4-

Cr2O72-

ClO2-

BrO2-

MnO4-

O22-

H3O+

ClO3-

Pb2+

Review Problems The following are problems that students entering AP Chemistry are expected to solve and answer without difficulty. You may use a calculator and the periodic table on the second page. Show your work and report your answers in the correct significant figures for all problems. Math Review 1. Calculate the following answers, being sure to use the correct number of significant figures.

a. 0.003 + 0.18 = b. 1.003 - 3x10-2 = c. 0.300 x 4 = d. 2.9 / 0.0002 =

2. Convert the following units. Report your answer in the correct number of significant figures (hint: use scientific notation as needed).

a. 0.3 g = _______________________________ mg b. 1μL = ________________________________ dL c. 4.5 cm = ______________________________ m d. 3x103 kg = ____________________________ mg

Matter, Atomic Structure, and Bonding 1. Select the best definition of matter.

a. Something of importance b. Anything that takes up space, and has mass and volume. c. An atom’s weight in mg

2. Which of the following can be identified as matter? Select all that apply.

a. CO2

b. H2O c. C6H12O6 (glucose) d. A bicycle

3. Match the following instruments to the physical property that they measure.

Graduated Cylinder Beaker Mass Scale Volume Balance

4. Match the following properties to the correct SI units that they are measured in.

Grams (g) Cubic Centimeters (cm3) Mass Litres (L) Milliliters (mL) Volume Micrograms (μg) Grams per Cubic Centimeter (g/cm3) Density

5. A block of gold has the following dimensions: 1.0 cm by 2.0 cm by 0.5cm. Remember to report your answers in correct significant figures for all calculations.

a. Find the volume of the block (Volume = lwh)

b. The mass of the block is 19.3g. Calculate the density of the block. (Density = m/V)

c. Say you wanted to create a crown from this bar of gold. You melted the bar, poured it into a mold, and then let it cool. Will the density of the gold change? Why or why not?

6. Match the following properties to their correct category (intensive or extensive).

Volume Melting Point Intensive Property Density Extensive Property Mass

7. Give an example of the following phases. I have completed the first as an example. a. Solid (s) _______________________Ice_______________________

b. Liquid (l) ________________________________________________

c. Gas (g) ________________________________________________

d. Aqueous (aq) _________________________________________________

8. Define the following terms.

a. Chemical Change

b. Physical Change

c. The Law of Conservation of Mass

d. Isotope

e. Compound

f. Atomic Mass

g. Atomic Number

9. Label the diagram below by filling in the blanks with the following words: Atomic Number, Chemical Symbol, Chemical Name, Average Atomic Mass. 10. Using the diagram to the right, answer the following questions.

a. # Protons = b. # Electrons = c. # Neutrons =

11. Identify and label the following directly on the periodic table on page 2.

a. Alkali Metals e. Noble Gases b. Alkaline Earth Metals f. Metals c. Halogens g. Non-Metals d. Metalloids h. Transition Metals

12. Lithium is made up of two isotopes, lithium-6 and lithium-7. The isotopes are found in the abundance percentages of 7.6% and 92.4%, respectively. Calculate the average atomic mass of lithium. Hint: for each isotope, convert the percent abundance to a decimal (80% → 0.80), then multiply this by the isotope’ mass number. Then, add the products. 13. Define the following terms.

a. Radioactive Isotope

b. Nuclear Reactions

c. Radioactive Decay

d. Alpha Decay

e. Alpha Particle (α)

f. Beta Decay

g. Beta Particle (β)

h. Half-Life

i. Gamma Rays (γ) 14. Which of the following forms of radiation does not change the identity of the radioactive atom?

a. Alpha Decay b. Beta Decay c. Gamma Rays

15. Write the equation representing the alpha decay of radium-222. What is the daughter isotope? (Ra atomic number = 88) 16. Label the electron shells on the diagram of a sodium atom below, using the words core electrons and valence electrons. 17. In order to have a very stable atom or compound, there must be _______ (#) valence electrons. 18. Fill in the blanks with gain and lose. In an ionic compound, metals ________ electrons while nonmetals ________ electrons. 19. Write the full electron configuration for sulfur (S). 20. Write the electron configuration for cobalt (Co). (You may use noble gas shorthand if you prefer) 21. Define the following types of chemical bonds, then rank their strength.

a. Ionic Bond

b. Covalent Bond

c. Hydrogen Bond

d. Van Der Waals

Molecular Structure and Properties 1. Define an isomer. 2. Which of the molecules to the right show the correct bonding tendencies? ___________________ 3. Describe the Octet Rule. 4. Explain why nitrogen bonds with hydrogen to form NH3 , but not NH2 or NH4 . Use lewis dot structures to support your answer. 5. List 3 non-metal elements that can combine with 1 fluorine atom to satisfy the octet rule. (Hint: Use the periodic table to find charges of the elements)

1. 2. 3.

6. Define the following.

a. Polar Molecule

b. Nonpolar Molecule

c. Intermolecular Force

d. Electronegativity

7. Complete the table below by creating 3-carbon molecules that contain the following functional groups. I have completed the first two as an example.

Functional Group Name Functional Group Structure Molecule

No Functional Group

No Functional Group

Carboxyl

Ester

Ketone

Hydroxyl

Amine

6. Name the following molecular shapes as bent, pyramidal, tetrahedral, trigonal planar, and linear.

Structure Name

Phase Changes and Behavior of Gases 1. Define the following.

a. Melting Point

b. Boiling Point

2. Convert -40°C to °F (F = C + 32).59

3. Convert -40°F to Kelvin (K = C + 273) 4. Fill in the blank. When temperature increases, the average speed of gas particles _________________ . 5. What do scientists hypothesise would happen to the motion of gas particles at absolute zero? 6. An expandable cylinder contains 10 mL of air at 295 K. What will the volume of gas be if the temperature is increased to 550 K? (Hint: Use Charles’ Law V = kT. First, find the proportionality constant (k), then use that to find the new volume) 7. Match the following terms to the correct phase change. Sublimation (g) → (l) Melting (l) → (g) Vaporization (s) → (l) Freezing (l) → (s) Deposition (s) → (g) Condensation (g) → (s) 8. Gas pressure and volume are __________________________ to each other.

a. Proportional b. Inversely Proportional c. No Correlation

Review: Gas Laws

Law Equation Constants

Charles’s Law V = kT P and amount of gas does not change

Gay-Lussac’s Law P = kT V and amount of gas do not change

Boyle’s Law P = k (1/V) T and amount of gas do not change

● k is a proportionality constant. k is different for every gas and every gas law, and must be experimentally determined or calculated with a known set of volume, pressure, or temperature.

9. A balloon is filled with air to a volume of 1.5 L. The temperature of the air is 20°C. The balloon is placed in the refrigerator until the air in the balloon is 10°C. Calculate the new volume of the balloon by following the steps below.

a. Will the volume increase or decrease? Explain how you know.

b. Fill in the blanks below. Initial Conditions Final Conditions P1 = __________ P2 = __________ T1 = __________ T2 = __________ V1 = __________ V2 = ?

c. Which condition is staying constant? Which Gas Law should be used?

d. Convert the temperatures to Kelvin.

e. Determine the value of k (use the initial conditions).

f. Use k to calculate V2 Review: Avogadro’s Law

● If two gas samples have the same pressure, temperature, and volume, they contain the same number of molecules, even when comparing different gases.

10. The box to the right contains CO2 gas. a. How many molecules are in the box? ___________ b. How many atoms are in the box? ___________

11. Two balloons, one filled with helium and one filled with carbon dioxide rest at identical pressures, temperatures, and volumes.

a. The balloon containing carbon dioxide has a greater mass than the one containing helium. Explain why, using Avogadro’s Law.

b. The number of atoms in the helium balloon is ______ than the number of atoms in the carbon dioxide balloon.

i. Greater than ii. Less than

iii. Equal to Review: The Ideal Gas Law

● Relates pressure (P), volume (V), temperature (T), and number of moles (n) for any gas sample. ● Includes the Universal Gas Constant, R, which is the same for all gases. ● PV = nRT, where R = 0.082 (L*atm/mol*K) ● Since R is defined in units of L*atm/mol*K, you can only use this value of R (without

conversions) if your gas sample is measured in litres, atmospheres, kelvins, and moles. 12. How many moles of hydrogen (H2) gas are contained in a volume of 2L at 40.0°C and 1.5atm? Stoichiometry, Solution Chemistry, Acids and Bases 1. When an ice cube melts, which of the following quantities will change?

a. The number of atoms b. The mass c. The volume d. The density e. All of the above

2. Balance the following chemical equations by writing the correct coefficients in the blanks. If there should be only 1 molecule, you may leave the space blank.

_____ C2H6(g) + _____ O2(g) → _____ CO2(g) + _____ H2O (l)

_____ KClO3(s) → _____ KCl(s) +_____ O2(g)

_____ KBr(aq) + _____ AgNO3(aq) → _____ KNO3(aq) + _____ AgBr(s)

_____ P4(s) + _____ H2(g) → _____ PH3(g)

3. Fill in the table of the types of reactions below.

Name General Equation Example

Combination Reaction A + B → C

A → B + C

Single _________________

_______________________

Ca(s) +2HCl(aq) → CaCl2(aq) +H2(g)

Double _________________

_______________________

4. Which has more mass?

a. 1 mol of hydrogen or 1 mol of carbon: ____________________________________ b. 1 mol of copper or 1 mol of gold: ________________________________________ c. 5 mol of carbon or 1 mol of gold: ________________________________________

(Hint: convert both amounts from moles to grams using the element’s molar mass from the periodic table) 5. What is the mass of 5 mol of iron (III) oxide, Fe2O3? 6. Which has more moles of oxygen atoms, 153g of BaO or 169g of BaO2? 7. Define the following.

a. Solvent

b. Solute

c. Homogeneous Mixture

d. Heterogeneous Mixture

e. Saturated Solution

8. A beaker of salt water has a 1.5M molarity. Half of the solution is poured out of the beaker. Will the molarity change? Why or why not? 9. Determine the molarity of 0.10 mol of NaCl in 1.0L of solution. (Molarity = moles of solute / litres of solution) 10. How many moles of sodium cations (Na+) are dissolved in a 1.0L sample of 9.0 M Na2SO4? 11. How many grams of solute do you need to make a 1 L solution of the following? (Hint: use the molarity equation)

a. 0.50 M NaCl

b. 2.0 M CaCl3 12. Fill in the blanks with the words donor and acceptor.

Acids are proton _______________ , while bases are proton _______________ .

13. Fill in the table below with the molarity, proton concentration [H+], and pH of the following solutions. (pH = -log [H+])

Molarity [H+] pH

1.0 M HCl 0

1.0 x 10-1 M

0.00001 M HCl

14. Fill in the table below with the molarity, OH- concentration [OH-], [H+], and pH of the following solutions. ([H+][OH-] = 1x10-14)

Molarity [OH-] [H+] pH

Water (no solute) 7

1.0 x 10-5 M

1.0 x 10-14 M

15. Ms. Mason has a solution with a pH of 3. She wants it to be basic, and decides to dilute the solution with water. Will this allow her to create a basic solution? Why or why not? 16. A neutralization reaction occurs when an acid and a base react and neutralize each other. Fill in the blanks below to complete the general equation for a neutralization reaction.

Acid + Base → ____________ + _____________

17. Which of the following reactions would create a neutralization reaction that forms sodium nitrate (NaNO3) as a product? (Hint: Write out the products of each reaction)

a. Mg(NO3)2 + NaOH b. HNO3 + NaOH c. HNO3 + NaCl

Explain how you chose your answer.

18. Suppose you have 500g of AlCl3 and an unlimited supply of NaOH. How many grams of Al(OH)3 and NaCl can you make? Follow the steps below to calculate the maximum yield.

a. Balance the equation below.

_____ AlCl3(aq) + _____ NaOH(aq) → _____ Al(OH)3(s) + _____ NaCl(aq)

b. Calculate the molar masses of the reactants and products. AlCl3 = Al(OH)3 = NaOH = NaCl =

c. Convert the limiting reagent from grams to moles.

d. Use mole ratios to determine the moles created of each product.

e. Use the molar mass to convert from moles to grams.

f. Suppose you ran this reaction and got 200g of Al(OH)3. Determine the percent yield. (Percent yield = actual yield / theoretical yield *100% , where actual yield is the experimental amount, and theoretical yield is what you calculated above)

Energy and Thermodynamics 1. Define and give and example of the following.

a. Exothermic Reaction

b. Endothermic Reaction 2. Fill in the blanks with the words high and low. Heat is a transfer of energy between two objects due to a difference in temperature. Heat always transfers from a ________ temperature to a _______ temperature. 3. Describe the connection between temperature and kinetic energy. 4. Match the following terms to their definitions

Conduction A warm substance changes location, like warm air rising. Convection Electromagnetic waves carry energy from an energy source, like the sun. Radiation A substance transfers heat to another substance it is in contact with, like

accidentally touching a hot stove.

Review: Laws of Thermodynamics

First Law of Thermodynamics Energy is always conserved, it cannot be created or destroyed

Second Law of Thermodynamics Entropy of a system tends to increase over time (energy tends to “spread out” and become more disordered)

5. Thermal energy is defined as the total amount of energy within the particles in a sample. Below, two containers of water are at the same temperature. Which sample has more thermal energy? Explain.

Sample 1 Sample 2

6. The heat required to raise the temperature of 1g of a substance by 1°C is called the specific heat capacity. Which of the following are true statements about heat capacity? Select all that apply.

a. Substances with low specific heat capacity can heat up and cool down easily (with a small energy transfer).

b. Substances with a high specific heat capacity can heat up and cool down easily (with a small energy transfer).

c. Water has a higher specific heat than aluminum foil. d. Water has a lower specific heat than aluminum foil.

7. What is a combustion reaction? 8. When a reaction is exothermic, the heat of reaction (ΔH) is

a. Positive b. Negative

Review: Heat Transfer

● Heat transfer to any substance can be calculated using the equation q = mCpΔT, where m is the mass, Cp is the substance’s specific heat capacity, and ΔT is the temperature change.

9. You have 15g of menthol at 75°C, but you want it to be at 20°C. How much energy (in calories) must be removed from the sample? (Menthol Cp = 0.58 cal/g°C) 10. Below is a diagram showing the heating curve of water. Fill in all of the labels using the words solid, liquid, gas, boiling point, and melting point. .

11. When the water in the diagram above is being heated, a constant source of energy is being transferred to the substance. However, the temperature is constant (not increasing) at two points. Explain why. 12. What is bond energy? 13. Calculate the net energy change for the combustion of 1 mol of methane with 2 mol of oxygen by following the steps below.

CH4 + 2O2 → CO2 +2H2O

Bond C - H O = O C = O O - H

Bond Energy (kJ/mol) 413 495 799 467

a. Calculate the energy required to break all of the bonds in both reactants. (Hint: This process requires energy, will that make this total positive or negative?)

b. Calculate the energy released from the formation of all the new bonds. (Hint: This process requires energy, will that make this total positive or negative?)

c. Calculate the net exchange of energy (Net Energy Exchange = Energy Input + Energy Output).

d. Is this reaction exothermic or endothermic? Explain. 14. Explain the picture to the right. Include brief definitions of potential and kinetic energy.

15. Which of the following energy diagrams show an exothermic reaction? Which shows an endothermic reaction? Explain how you know, using the words heat of reaction (ΔH), negative, and positive. 16. Label the activation energy on both diagrams above. Which reaction is more energetically favorable? Explain. 17. Which of the following actions can speed up reaction rates? Select all that apply.

a. Increasing the temperature b. Decreasing the temperature c. Increasing the concentration or pressure d. Decreasing the concentration or pressure e. Adding energy (like microwave radiation or light) f. Adding a catalyst g. Mixing the reactants

18. Select the 2 TRUE statements about catalysts below.

a. A catalyst is consumed or changed during a reaction. b. A catalyst is not consumed or changed during a reaction. c. A catalyst works by lowering the activation energy of a reaction. d. A catalyst works by making the heat of reaction (ΔH) negative.

19. The energy diagram below shows a reaction without a catalyst. Draw a new line on the diagram to represent the same reaction with a catalyst.