Name _________________________________________________ Period ______________
Mr. Rivas’ Super Awesome Chemistry Final Exam STUDY GUIDE MATTER Fill in the following chart for TYPICAL Solids, Liquids and Gases (not water) State of Matter Particle Shape Density Rank Arrangement Definite or 1 = highest Regular or Indefinite 3 = lowest Irregular
Energy Rank 1 = highest 3 = lowest
Compressibility Yes or No
Gas
Liquid
Solid
Atoms – draw a diagram of an atom of Lithium (mass number 7), labeling the following parts – nucleus, electron cloud, protons, neutrons, electrons
Label as: molecule, atom, compound, substance, homogeneous mixture, heterogeneous mixture ________ 1. ________ 2. ________ 3. ________ 4. ________ 5. ________ 6. ________ 7.
H2O A single particle of pure potassium A sample of sand scooped from a beach Sodium chloride Iced tea The smallest particle of an element that retains all of its chemical properties Chicken soup
PHYSICAL AND CHEMICAL CHANGES AND PROPERTIES Classify the following as a chemical or physical property or chemical or physical change ________ 1. Liquid collects on the side of a cold drink glass ________ 2.
Hexane is flammable
________ 3.
Baking soda decomposes in the cake, releasing CO 2 to make it fluffy.
________ 4.
Copper II sulfate is blue.
________ 5.
Ice is melting
________ 6.
Pure water has a density of 1 g/cm3
STATES OF MATTER Mark the statements as true or false. If false, make changes to correct the statement. 1. Liquids have the greatest amount of kinetic energy among the phases of matter 2.
In general, a solid is the densest phase of matter.
1
3.
Particles in a solid demonstrate rotation, vibration and translation.
4.
The particles of liquids have low amounts of translation unless they are stirred.
5.
When a solid is melting, its temperature is still rising as it melts.
6.
Both solids and liquids can flow.
7.
Translation is the type of kinetic energy in which a particle moves from place to place.
DENSITY CALCULATIONS Solve the following problems involving density calculations. Show all work. ________ 1.
What is the formula for density?
________ 2.
A cube is 17 cm3 in volume with mass of 4.0 grams. Calculate its density.
________ 3.
If a ball has a density of 27.5 g/cm3, determine its volume if its mass is 626 g.
METRIC UNITS Perform the following determinations. ________ 1. 1 kilometer (km) = ________________ meters (m) ________ 2. 1 milligram (mg)= ________________ grams (g) ________ 3. 1 liter (L)
= ________________ milliliters (mL)
________ 4. 2.3 kilograms (kg)= _________________ grams (g) ________ 5. 12 meters (m) = _______________________ centimeters (cm)
ATOMIC STRUCTURE Fill in the following table with the appropriate information Subatomic Particle Charge Location Nucleus or electron cloud Proton
Mass
Order of Size 1= largest 3 = smallest
Neutron Electron
PERIODIC TABLE Answer the following questions What number of valence electrons indicates the most stable arrangement? Which group on the periodic table fits this description without forming ions? Elements on the periodic table are organized in order of increasing ____________.? Elements are identified by the number of ________ in their nucleus. The number of __________ and ___________ may vary.
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ORGANIZATION OF THE PERIODIC TABLE Label the following on the diagram below. Then, label the number of valence electrons inside the space for each group. s-block metals Aluminum group Halogens
p-block non-metals Carbon group Noble Gases
d-block alkali metals Nitrogen group transition metals
f-block alkali earth metals Oxygen Group inner transition metals
Define each of the following trends and indicate them the periodic table above using arrows. Be sure to indicate group and periodic trends. Atomic size Electronegativity Ionization energy
ELECTRON CONFIGURATION Write the full electron configuration for the following elements. Circle the valence electrons. Helium Strontium Barium
Oxygen
Fluorine
Neon
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BONDING Draw the covalent structure for the following chemical formulas. Name the shape of the molecule. 1. PBr3 3. SO3
2.
SO2
4.
PO43-
Indicate the correct word(s) to complete each sentence. ________ 1. Ionic bonds are formed between metals and (other metals / non-metals) ________ 2. Covalent bonds are formed between non-metals and (other non-metals / metals) ________ 3. Metals (do / do not) form bonds with other metals ________ 4. The transition metals (lose / gain) electrons to form ions. ________ 5. When comparing degree of polarity between two bonds, the bond with the greatest polarity has the (largest / smallest ) difference between the two atoms in the bond.
ELECTRON DOT STRUCTURES Draw the electron dot structure of the valence electrons for the following elements. Indicate which, if any, fulfill the “octet rule” when neutral. If they do not fulfill the octet rule when neutral, indicate the number of electrons gained or lost for the atom to fulfill the rule. Sodium Silicon Chlorine Magnesium
Phosphorus
Aluminum
Sulfur
Argon
ION FORMATION Fill in the following table regarding ions: Element Sodium
Group Name
Ion Type cation/anion
Charge # and +/-
Metal / Non-Metal
Selenium Boron Phosphorus Calcium Fluorine Radon Phosphorus
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BONDING Fill in the following table regarding ionic bonding: Cation
Charge
Anion
Calcium
Sulfur
Strontium
Chlorine
Aluminum
Iodine
Cesium
Nitrogen
Magnesium
Oxygen
Boron
Phosphorus
Sodium
Fluorine
Charge
Balanced Formula
Chemical Name
Formula and Charge
Balanced Formula
Chemical Name
POLYATOMIC IONS Fill in the following chart: Cation
Formula and Charge
Anion
Calcium
Hydroxide
Strontium
Nitrate
Ammonium
Acetate
Cesium
Chlorate
Magnesium
Cyanide
Boron
Sulfate
Sodium
Nitrite
BALANCING EQUATIONS AND REACTION SYMBOLS Balance all equations. Label reaction type for all reactions. 1. N2(g) + H2(g) 2.
KClO3(s)
∆
KCl(aq) +
3.
NaCl(s) +
F2(g)
NH3(g) O2(g)
NaF(aq) +
Cl2(g)
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4.
AgNO3(aq) +
MgCl2(aq)
AgCl↓ +
5.
C3H8(l) +
O2(g)
CO2(g)
+
Mg(NO3)2(aq)
H2O(l)
Answer the following questions based upon the formulas above: 1. What does the subscript “(g) “indicate in equation 1? 2. What does the subscript “(s)” indicate in equation 2? 3. Does equation 3 take place in water? How do you know? 4. Write the formula for the precipitate in equation 4. Rewrite its name in word format. 5. Write the 3 criteria for identifying a combustion reaction. 6. What does “∆” indicate in equation 2?
REACTIONS AND STOICHIOMETRY: Perform the following stoichiometric calculations. Make sure you have a balanced equation! CH4 + O2 CO2 + H2O How many grams of O2 are needed to react with 9.02 g of CH4?
2. N2 + H2 NH3 How many grams of NH3 are formed when 11.7 g of H2 react?
3. C + O2 CO How many grams of carbon are needed to produce 0.98 g of CO?
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GASES Name and describe the three types of particle motion. Indicate the phases of matter that exhibit each.
Describe the “Kinetic Theory of Gases” and list the three assumptions associated with it.
Define “phase change.” Define these types of phase changes. Give an example of where you might experience each. Melting Vaporization Freezing or Crystallization
Condensation
Evaporation
Sublimation
GAS LAWS: Solve the following problems using the gas laws
1.
A gas with volume of 5.3 L at 17°C decreases to -3°C. Find the new volume.
2.
A 2.1 L sample of gas a standard pressure is moved to a 5.2 L container. Find the new pressure.
3.
An aerosol can contains gas at 15 atm and 25°C. Find the pressure inside the can at 100.°C
4.
Find the number of moles of gas in a container at the following conditions – 6.0 atm, 3.6 L, and 118°C.
5.
A gas at 742 mmHg and -18°C occupies 1.7 L. Find the volume of the gas at STP.
6.
Define absolute zero. Give the temperature for it in both °C and K.
7.
A mixture of four gases is stored at 1187.3 mmHg. H2 is 127 mmHg, O2 is 350 mmHg and CO2 is 642 mmHg. Find the partial pressure of the Helium in the mixture.
8.
Under what conditions does a gas stop acting “ideal” and begin acting “real?”
9.
What volume does one mole of any gas occupy at STP?
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CONVERSIONS Perform the following conversions: 1. 26C to K
5. 23.5 atm to mmHg
2.
-52C to K
6. 89 kPa to Torr
3.
89 K to C
7. 768 mmHg to atm
4.
304 K to C
8. 11.2 atm to kPa
WATER AND SOLUTIONS Draw a picture of a water molecule. Indicate the poles on this polar molecule.
Write a definition and example for each of the following. Soluble
Supersaturated
Saturated
Molarity
Unsaturated
Molality
You have a solution of NaCl dissolved in distilled water. Which is the solute and which is the solvent? How many ions will the NaCl split into?
Describe three ways to increase the rate of dissolving a substance.
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MOLARITY AND MOLALITY Answer the following: Write the general formula for calculating the molarity of a substance.
Write the general formula for calculating the molality of a substance.
A solution is made by dissolving 5.0 moles of NaCl in 2.0 Liters of dH2O. Calculate the molarity of the solution. A solution is made by dissolving 62. g of KCl in 0.50 L of dH2O. Determine the molarity of the solution.
A solution is made by dissolving 92.0 g of BaCl2 in 5.6 kg of dH2O. Determine the molality of the solution.
MOLAR MASS AND PERCENT COMPOSITION Find the molar mass for the following formulas: BaCl2
CuSO4
NaNO3
Ba(OH)2
Determine the percent composition of the following: S in CuSO4
NO3 in NaNO3
H2O in CuSO4 * 5 H2O
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THERMOCHEMISTRY Perform the following conversions: 41.2 kJ = ? J
710F = ? 0C
816 J = ? cal
1320C = ? K
0.0482 kJ = ? cal
255 K = ? 0C
1. How much heat is required to raise the temperature of 23.5 g of copper from 250K to 298K?
2. What is the specific heat (in J/g0C) of an unknown substance if adding 0.124 kJ of heat to 100g of the substance causes the temperature to increase by 120C?
3. How much heat is absorbed by 100.0 mL of water if its temperature increased by 13.2 0C? 4. If 15.00g of water at an initial temperature of 19 0C absorbs 2.03 kJ of heat, what will its final temperature be?
5. If you place 50.00g of a metal at a temperature of 98.00C into 100.0mL of water at 22.30C, and stir the mixture, what is the specific heat of the metal if the final Temperature of the mixture is 29.70C? (answer in J/g0C)
REACTION RATES AND EQUILIBRIUM Describe and give examples of three ways you can speed up the rate of a reaction.
Explain Le Châtelier’s Principle.
How does Le Châtelier’s Principle explain what happens to vapor pressure when a sealed bottle of soda is heated?
Think of a reversible reaction in which heat is a product of the forward reaction. Now imagine that heat is added to the system. Which way will the system shift and why?
Write the Keq equation for the following equilibrium reaction. aA + bB cC + dD
What does it mean if Keq < 1. How about > 1?
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What does it mean if Q > Keq? How about < Keq?
Write an equilibrium constant expression for the following chemical reactions: 1. 2 HIg) H2(g) + I2(g)
2.
2 BrCl2(g) Cl2(g) + Br2(g)
3.
Fe3O4(s) + 4 H2(g) 3 Fe(s) + 4 H2O(g)
Answer the questions based upon the following reaction. 2 BrCl(g) Cl2(g) + Br2(g) 1. Write Keq for the equation.
2. If [Cl2] = 4.00 mol/L and [Br2] = 4.00 mol/L, find the concentration of BrCl is Keq = 11.1.
ACIDS AND BASES Explain the Arrhenius model of acids and bases.
Explain the Brønsted-Lowry Model of acids and bases.
Explain the Lewis Model of acids and bases.
Explain the difference between monoprotic acids and polyprotic acids.
Explain the difference between molar mass (gram formula mass) and gram equivalent mass.
Calculate the pH of solutions having the following ion concentrations. a. b. c. d.
[H+] = 1.0 x 10-2 M [H+] = 3.0 x 10-6 M [OH-] = 8.2 x 10-6 M [OH-] = 4.0 x 10-3 M
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Write balanced equations for the following acid-based neutralization reactions. a. nitric acid and cesium hydroxide b. hydrobromic acid and calcium hydroxide c. sulfuric acid and potassium hydroxide d. acetic acid and ammonium hydroxide Answer the following questions regarding titration reactions. a. Determine the molarity of a CsOH solution if 30.0 mL of the solution is neutralized by 26.4 mL of a 0.250 M HBr solution.
b. Find the molarity of a nitric acid solution if 43.33 mL of a 0.1000 M KOH solution is needed to neutralize 20.00 mL of a solution of unknown concentration.
c. Find the concentration of a household ammonia cleaning solution if 49.90 mL of 0.5900 M HCl is needed to neutralize 25.00 mL of the cleaning solution.
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BIG IDEAS Intro Unit 1. Chemists use the metric system when recording their measurements. 2. Dimensional analysis must be used to convert between measurement units. 3. Chemicals have both physical and chemical properties that can be used to tell them apart, and also to separate components of a mixture. 4. When it comes to evaluating a solute/solvent relationship, remember the phrase “like dissolves like.” Atoms/Periodic Table Units 1. Ionic compounds form when atoms gain or lose electrons. Metals lose electrons, nonmetals gain them. 2. The number of electrons gained or lost can be predicted with an understanding of the octet rule and the number of valence electrons an atom contains. 3. The periodic table is organized by electron configuration, and elements are classified as metals, nonmetals or metalloids based on physical and chemical properties. 4. Atoms have a central nucleus that contains the protons and neutrons, and electrons are found in orbitals located outside the nucleus. 5. Electron configurations are used to describe locations of electrons within levels, sublevels and orbitals. 6. Metal reactivity is summarized in the Metal Activity Series. Formula Writing/Moles/Chemical Reactions Units 1. Conversions between moles, particles, and mass are widely used in chemistry 2. Chemical reactions must be balanced with coefficients to obey the law of conservation of mass/atoms. 3. The quantity 6.02 x 1023 is called a mole, and the periodic table lists the mass of 1 mole of each element’s atoms, called the molar mass of the element. 4. Coefficients are interpreted as either particle, mole, or (if everything is a gas) liter ratios. 5. Stoichiometry is the name given to the process of using coefficients as mole ratios to convert between chemicals in a reaction equation. How much is used? How much is made? These questions can be answered with a stoichiometric calculation. Gases Unit: 1. 4 variables describe a gas’ behavior: P, V, n and T. 2. Boyle, Charles and Gay-Lussac developed mathematical gas laws that are used to predict and calculate changes in gas properties. 3. At STP, one mole of any gas occupies a volume of 22.4 L. This is called the molar volume of a gas. Solutions Unit: 1. Molarity and molarity are two units used to express the concentration of a solution quantitatively. 2. Concentrated and dilute are relative terms for expressing the concentration of a solution qualitatively. 3. Solutes and solvents are the two components of a solution. Thermochemistry Unit: 1. The collision theory explains that gas-phase chemical reactions occur when molecules collide with sufficient kinetic energy and proper orientation. 2. Heat, represented by q, is energy that transfers from one object to another because of a temperature difference between the objects. 3. An endothermic process absorbs heat from the surroundings. An exothermic process releases heat to the surroundings. 4. Calorimetry is the measurement of the heat flow into or out of a system for chemical and physical processes. The insulated device used to measure the absorption or release of heat in chemical or physical processes is called a calorimeter. 5. You can calculate the heat absorbed or released by the surroundings (qsurr) using q = m x Cp x ΔT. Equilibrium Unit: 1. Le Châtelier’s principle describes what happens when stress is applied to a system in dynamic equilibrium. 2. A catalyst lowers the activation energy of a reaction. 3. Keq is used to determine whether reactants or products are favored in a reaction. 4. Q is used to determine whether a reaction will shift forward or backward. Acids and Bases Unit: 1. The Arrhenius, Brønsted-Lowry, and Lewis definitions of acids and bases are all based on different criteria. 2. pH is based on the concentration of H+ ions in solution. 3. An imbalance of H3O+ and OH- ions causes solutions to be either acidic or basic. 4. Titrations are used to determine the concentration of a solution using a solution of known concentration.
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