Download Chemistry FINAL: CONTENT Review Packet

Chemistry FINAL: CONTENT Review Packet
Name:___________________________________________Period:_____Date:____________________
Classification of Matter & Chemical/ Physical Changes
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2.
3.
4.
5.
6.
7.
8.
9.
_________________ are substances that are made up of two or more elements which are chemically combined
_______________________is made from two or more substances that are physically combined
The ability to do work is known as ________________
________________________ are substances that are made up of only one type of atom
____________________________ is anything that has both mass and volume
______________________________is a solid, produced by a reaction that separates a solution.
___________________________ is a form of energy
A ___________________________ consists of a solute in a solvent
List 2 lab safety precautions.
10. What do the following formulas represent: NaCl (aq) and H2 (g)
11. Define what a heterogeneous mixture and a homogeneous mixture are and provide an example of
each.
12. Suppose that during a reaction a chemistry student touches the beaker and observes that it feels COLD. The
student should conclude that the chemical reaction is ______________________.
a) Why?
13. Suppose that during a reaction a chemistry student touches the beaker and observes that it feels HOT. The
student should conclude that the chemical reaction is ______________________.
a) Why?
14. Classify the following as an element, compound, homogeneous mixture or heterogeneous mixture.
a) Table salt _____________________________
b) Aluminum_____________________________
c) Dirt_____________________________
d) Sugar water_____________________________
15. Define the terms chemical and physical change
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16. Classify the following as chemical change or physical change.
a) Iron rusting ____________________________
b) Ice melting____________________________
c) Evaporation____________________________
d) HCl reacting with Mg to create H2 gas____________________________
17. Another name for a homogeneous mixture is a(n) ________________
18. What are four indicators of a chemical reaction?
Measurement & Density
1. What is the difference between a qualitative measurement and a quantitative measurement?
2. For each of the following properties, what would be the unit abbreviation(s)
a. Mass ___________________________________________________
b. Length___________________________________________________
c. Volume___________________________________________________
d. Temperature___________________________________________________
3. What is the difference between accuracy and precision?
4. Try these conversions:
a. 5000 mg = ______________g
b. 1L = ________________mL
5. Write the following numbers in scientific notation.
a. 74,600 km ___________________________
b. 0.0443 km ___________________________
6. Write the following numbers in ordinary notation.
a. 8.5 x 107 km ______________________________
b. 9.67 x 10-4 km ____________________________
7. Define density.
8. Find the density of an unknown solid given that a 5.62g sample occupies 2.35 cm3
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Atomic Structure, Quantum Numbers, and Electron Notation
_________
1. proton
_________
2. atom
_________
_________
3. mass number
4. atomic mass unit
_________
5. electron
_________
_________
6. mole
7. average atomic mass
_________
_________
_________
8. atomic number
9. neutron
10. isotopes
a. the total number of protons and neutrons
in the nucleus of an atom
b. the weighted average mass of the atoms
in a naturally occurring sample of an element
c. 1/12 the mass of a carbon-12 atom
d. the number of protons in the nucleus of an
element
e. atoms with the same number of protons
but different number of neutrons
f. negatively charged subatomic particle
g. the smallest particle of an element that
retains the properties of that element
h. a counting unit; 6.022 x 1023
i. positively charged subatomic particle
j. subatomic particle with no charge
1. Suppose an atom has 27 protons and 32 neutrons. What is the mass number? _________
2. Suppose an atom has 29 protons and 35 neutrons. How many electrons does it have? __________
3. Suppose an isotope has a mass number of 27 and an atomic number of 13.
a. Write the hyphen notation for this isotope. __________________
b. Write the nuclear symbol for this isotope.__________________
4. What is electromagnetic radiation? Provide examples.
5. What is the BEST way to describe the nature of an electron?
6. List and describe the four quantum numbers.
7. One orbital can hold a MAXIMUM of ______ electrons.
8. What it the correct electron configuration notation and orbital notation for Na and Al
Diagonal Rule:
1s2
2s2
3s2
4s2
5s2
6s2
7s2
2p6
3p6
4p6
5p6
6p6
7p6
3d10
4d10
5d10
6d10
4f14
5f14
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Periodic Table
1. ______________________________ are elements that show the properties of both metals and nonmetals.
2. How are elements arranged on the Periodic Table?
3. What is electronegativity?
4. What is the most electronegative element? _______________________
5. Write whether an increase or decrease of a trend occurs as you go across the Periodic Table from left
to right in a period, and when you go from top to bottom in a group.
Across a Period
Trend
Down a Group
1st Ionization Energy
Atomic radius
Electronegativity
6. What name is given to the following?
a. Group #1 Elements: ________________________
b. Group #2 Elements: ________________________
c. Group #18 Elements: _______________________
d. Group #17 Elements: _______________________
7. Use the elements from part of period three, listed below, to answer the following questions. Write the
symbol of the element described in the blank provided. (1pt.each)
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12
13
14
15
16
17
18
Na
Mg
Al
Si
P
S
Cl
Ar
22.99
24.30
26.98
28.09
30.97
32.07
35.45
39.95
_____ a). an element with 15 electrons
_____ b). the most reactive metal
_____ c). the most reactive nonmetal
_____ d). an element with 12 protons
_____ i). an element in the same group as potassium
(refer to wall chart to find potassium-K )
_____ j). an element with an oxidation number of –2
_____ k). an element with 14 neutrons
_____ e). an element with the atomic number of 17
_____ f). the MOST stable element
_____ h). an element in the halogen group
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Bonding
1. A bond where there is a sharing of electrons:
a) mutual
b) electronegativity
c) covalent
d) ionic
2. A bond where there is transferring of electrons:
a) mutual
b) electronegativity
c) covalent
d) ionic
3. What is the difference between nonpolar covalent and polar covalent bond type?
4. How many valance electrons are found in the following groups?
a. Group 1________
b. Group 2________
c. Group 13________
d. Group 14________
e. Group 15________
f. Group 16________
g. Group 17________
h. Group 18________
5. With respect to bonds formed between the following pairs of atoms:
• Determine the electronegativity difference. SHOW WORK!
• Determine the probable bond type (ionic, polar covalent, or nonpolar covalent).
• Assign partial charges to atoms that are part of a polar covalent bond.
Pairs of
Atoms
Electronegativity
Difference
Probable
Bond Type
Partial Charge (if
polar-covalent)
H and H
S and O
K and Br
6. A molecule is a _______________group of atoms held to together by __________________bonds.
7. What is the difference between a cation and an anion?
8. List the diatomic elements?
9. What is the difference between an electron pair and a lone pair?
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10. What is the molecular geometry of NF3
11. What is the molecular geometry of CH4
12. What is the molecular geometry of CO2
Writing Formulas & Chemical Equations
Name the following:
Write formulas for the following:
Pb3(PO4)2
Zinc chloride
Cr2O3
Aluminum hydroxide
1. What are the four types of chemical reactions? Write a simple equation to represent each type of reaction.
2. The empirical formula is the ______________
a) simplest whole number ratio
b) actual composition
c) actual mole ratio
3. A compound that contains water is called a(n)______________
a) hydrate
b) hydroxide
c) anhydrate
d) none of these
4. A calculation that determines the mass/mass or mass/mole relationship is ____________________
a) quantitative
b) qualitative
c) stoichiometry
d) molar mass
5. The substance that comes through the funnel during the filtering process is the _________________
a) filter residue b) the filter paper
c) the filtrate
d) the decant
6. Heating until you get a matching weight is referred to as ______________________
a) being careful b) following directions c) heating to accuracy d) heating to precision
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7. A hydrate that has been heated until dry is a _______________
a) decomposed b) anhydrous
c) hydrated
d) oxidized
8. An insoluble product that forms from the reaction of two liquids is called __________
a) solute
b) precipitate
c)flakes
d) residue
9. To balance and equation, one uses_____________
a) subscripts
b) roman numerals
c) coefficients d) stoichiometry
10. The law of conservation of mass states that the mass of the reactants must be _________ the mass of the
products
a) greater than
b) less than
c) equal to
d) cannot be determined
11. Tarring (re-zeroing) the scale is taking into consideration the ___________
a) mass of a substance
b)mass of an empty container
c) mass of the reactants
d)mass of the products
12.The symbol that indicates a crystal in a reaction is
a) g
b) s
c) aq d) cr
13.When the equation Al2(SO4)3 + Ca(OH)2 Æ Al(OH)3 + CaSO4 is correctly balanced, the coefficient for
Ca(OH)2 is_____________
a) 3
b) 2
c) 1
d) 4
14. Using the equation in #13, what is the mole ratio between Al2(SO4)3 and CaSO4 ?
a) 2:3
b) 1:2 c) 2:3 d) 1:3
15. A substance that increases the rate of a reaction without itself being changed is a(n) ______.
a) precipitate
b) catalyst
c) inhibitor
d) none of the above
16. What is a reversible reaction? How is a reversible reaction represented?
Kinetic Theory, Solids, Liquids, Gases, Phase Changes
1. Name the 3 parts of the kinetic molecular theory.
2. How does temperature affect the kinetic energy of a substance?
3. List the values for STP.
4. If the volume of a gas is decreased, then gas pressure will_____________.
5. If the volume of a gas is increased, then gas temperature will_____________.
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6. How does evaporation differ from sublimation?
7. Complete the chart to name the four phases of matter and compare their volumes, shapes, and average
kinetic energy
Phase
Liquid
Shape
Volume
Average Kinetic Energy
definite
definite
definite
Not definite
Very FAST (violent)
Not definite
Fast
8. What is the difference between a crystalline solid and an amorphous solid?
9. Which of the 4 phases of matter is considered to have the strongest intermolecular forces? ________
a) The weakest intermolecular forces? _______________
10. A liquid will boil when its equilibrium vapor pressure EQUALS____________________________.
11. When a system at equilibrium is disturbed by application of a stress, it attains a new equilibrium position that
minimizes the stress. This is known as ____________________________________________________.
12. What is the triple point on a phase diagram?
Mixtures: Solutions, Suspensions & Colloids
1. Define the following terms: solution, colloid, suspension
2. Suppose 15 g of salt is dissolved in 1000 g of water. The sugar is the __________ and the water is the _______
3. Explain what the phrase “like dissolves like” means and give an example
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4. Explain the difference between an unsaturated, saturated and a supersaturated solution. Define these terms!
5. What 3 things can increase the rate at which a solute will dissolve?
7. How could you identify a colloid?
Acids/Bases/Salts
1. Describe acids and bases according to the theories of Arrhenius, Bronstead-Lowry, and Lewis.
2. Explain how blue and red litmus paper determine if a substance is an acid, base or salt.
3. Explain the process of neutralization using a chemical equation.
4. Define pH and draw the pH scale showing the locations of acids, bases, and neutral substances.
5. What happens during a splint test? How do you know O2, CO2 or H2 is produced?
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Nuclear Chemistry
1. Name and describe each of the 3 types of radiation.
2. What is a half-life?
3. Write the nuclear symbol and the hyphen notation for an isotope of phosphorus with a mass number of 32.
4. What is the difference between fission and fusion?
5. Balance the following nuclear reactions.
Things to study:
o ALL Notes
o Labs
o Review Guides
o Worksheets
o Mrs. Hostetter’s website (review PowerPoints)
The ACCUMULATIVE Final Is Worth 100 points:
o 79 multiple choice questions
o 21 points completion questions from lab
STUDY & GOOD LUCK!!
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Chemistry FINAL: MATH Review
Name:_________________________________________Period:_____Date:_______________
1. Calculate the molar mass of iron (II) phosphate. 358g
2. Find the percent composition of sodium carbonate. %Na= 43%, %C = 11%, %O= 45%
3. Convert 6.25 g of sodium chloride to moles. 0.108mol
4. Convert 0.500 mol of potassium iodide to grams. 83g
5. What is the empirical and molecular formula of a compound with the following percent composition:
P = 26.7 %, N = 12.1 %, Cl = 61.2 %. The molar mass of the compound is 695 g/mol.
EF = PNCl2, MF = P6N6Cl12
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6. How many grams of sodium phosphate are needed to completely react with 25.0 g of barium chloride?
13.205g
7. Suppose 100 mL of solution contains 5.92 g of dissolved calcium chloride. Calculate the molarity. 0.538M
8. Suppose 9.56 g of sodium hydroxide is dissolved in 150. g of water. Calculate the molality.
DID NOT COVER!!
9. Write the electron configuration notation, orbital notation, and electron dot symbol for arsenic.
10. Suppose a gas is collected in a 145 mL flask where the temperature of the gas is 23°C and the pressure of the
gas is 785 mm Hg. Calculate the volume of the gas at STP. 138.132mL
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11. Suppose 25.0 mL of a gas is at 24°C. Calculate the temperature at which the gas will occupy 30.0 mL
assuming constant pressure. 356.4K
12. Suppose the length of a piece of string is measured to be 33.45 cm. Calculate the percent error if the string is
actually 34.50 cm long. 3.04%
13. Suppose 4.25 g of sodium chloride is collected after a chemical reaction. Through stoichiometry calculations it
is determined that the theoretical yield of sodium chloride is 5.00 g. Calculate the percent yield. 85%
14. Suppose an unknown radioactive substance with a mass of 120 g has a half-life of 6 months. How much of the
substance will remain after 2 years? 7.5g
15. What is the volume of 10.5 g of copper? (The density of copper is 8.94 g/cm3.) 1.173cm3
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