Download Chemistry Review Fill in the blank

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Chemistry Review
Scientific Method
1. Know the steps to the scientific method.
2. Know your safety rules!
a. How do you safely dilute an acid?
b. What is the proper safety equipment used in every lab?
c. What two things are important to remember when heating a test tube?
3. Chemistry is the study of __________________ and the ____________________ it undergoes.
4. Name the piece of equipment you would use to measure each of the following in the lab and the unit:
a. Mass
b. Length
c. Temperature
d. Volume
5. Qualitative data is information that _____________________________________________.
6. Quantitative data is information that _____________________________________________.
7. Exothermic reaction ___________________ energy and energy is shown on the _______________ side.
8. Endothermic reaction __________________ energy and energy is shown on the _______________ side.
9. Independent variable: the variable that you plan to ______________________.
10. Dependent variable: the variable that changes in response to the changes in the independent variable.
11. A control is a ___________________ for comparison.
12. You are asked to study the effect of temperature on the volume of a balloon. The balloon’s size
increases as it is warmed. What is the independent variable? Dependent variable? What factor is held
constant? How would you construct a control?
Unit 1
1. Mass: ______________________________________________________________________________
2. Matter: _____________________________________________________________________________
3. Element: ____________________________________________________________________________
4. Atom: ______________________________________________________________________________
5. Compound: __________________________________________________________________________
6. Molecule: ___________________________________________________________________________
7. States of matter:
Phase
1. Solid
2. Liquid
3. Gas
Shape/Volume
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8. Physical change (phase change or change in state)
9. Physical property (change in state) vs. chemical property (can undergo a change into another
substance)
10. Chemical change
Indicators
Examples
1.
2.
3.
4.
11. How do you know you have produced a certain gas?
a. Oxygen
b. Hydrogen
c. Carbon dioxide
12. Classification of matter
a. Pure Substance: element (cannot be decomposed) vs. compound (can be decomposed by
chemical means)
b. Mixtures—homogenous (solutions, uniformed compositions, all the same) vs. heterogeneous
(not uniformed)
Unit 2
1. Metric conversion (King Henry Died Unusually Drinking Chocolate Milk)
2. Volume (amount of space occupied by an object)
a. L x w x h (for regular shaped objects…like cubes)
b. Water displacement (for irregular shaped objects)
c. Graduated Cylinder (for liquids)
3. Density = mass/volume (remember the triangle)
a. What is the density of a substance that has a mass of 22.5g and a volume of 5.0mL? What is
this substance?
b. What is the density of a 3.03cm x 10.cm x 2.5cm substance that has a mass of 50.0g? What
is this substance?
c. What is the mass of a piece of copper with a volume of 20mL?
d. A graduated cylinder is filled with water to 50.mL. 20g of pellets are added, raising the
water level to 60mL. What is the density?
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Unit 3
1. History of the atom:
Democritus
Thomson
Rutherford Quantum Theory (Modern)
2. Democritus  Gave us the name of the atom (atomos)…hard solid sphere
3. Dalton’s Atomic Theory (isotopes/atomic bomb)
1.
2.
3.
4.
5.
4. Discovery of the electron - _______________ (Thomson)
5. Discovery of the nucleus - __________________________________(Rutherford)
Discovered that the atom was:
a. ________________________________
b. ________________________________
6. Know the locations, charges, and relative mass of each of the particles found in the atom
Charge
Relative Mass
Location
Proton
Electron
Neutron
7. Atomic number = __________________________________
8. Atomic mass = __________________________________
9. Know how to identify the atomic number, mass number, number of electrons, protons, and neutrons of
various atoms
Electrons Protons Mass Number Atomic Number Neutrons
238
U
16
O-2
23
Na+1
10. Isotopes—Same element!!! Different Masses!!! Different Neutrons!!! They differ in the number of
___________ not ___________ (which always must stay the same as they indicate the atomic number!)
Ex.
A
ZX
Ex. 20F
Z means ____________
A means ____________
Atomic Number ____________
Mass Number ____________
11. Atomic Mass Unit (amu)—relative scale based on ___________________________________
12. Average Atomic Mass--________________________________________________________________
13. Molar Mass conversions
a. What is the molar mass of CO2?
b. How many grams are in 3.0 moles of H2SO4?
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c. How many molecules are in 64 grams of O2?
d. How many moles are in 84.2 grams of CO2?
e. How many moles are in 3.04x1023 molecules of H2?
f. How many atoms are in 3.46 moles of carbon?
g. How many grams are in 4.59x1025 particles of NaCl?
Unit 4
1. Electromagnetic radiation
2. c = wavelength ( ___ ) x frequency ( ____ )
3. E = Planck’s constant ( h ) x frequency ( ____ )
4. How are wavelength, frequency, and energy related?
Longest wavelength
Shortest frequency
Least Energy
Shortest wavelength
Highest frequency
Most Energy
5. Wave/Particle duality of electrons means __________________________________________________.
6. Quanta: ____________________________________________________________________________.
7. Energy Levels (____). The larger the energy level, the ___________________ away from the nucleus.
** Highest occupied energy level corresponds to ______________________________.
Ex.
Br2
O2
8. Ground State _______________________ vs. Excited State ________________________________
9. Hydrogen Line Emission Spectrum—Emitted as electrons fell from an _________________ state to a
__________________ state.
10. Bohr’s Model of the Atom (Good for the hydrogen atom only!)
a. Electrons ________________ nucleus only in fixed energy ranges called orbits.
b. Electrons can neither gain nor lose energy in an orbit, but they can move to a different orbit by
gaining or losing energy.
c. Lowest energy orbit is closet to the nucleus
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11. Questions using Bohr Model of the atom: What is the energy which is emitted when the electron falls
from the 6th energy level to the 3rd?
12. Flame tests (metals emitting colored light)
13. Emission and absorption of photons
a. Emission: ______________________
b. Absorption: ________________________
14. Quantum theory…basis for how electrons are arranged
15. Principal Quantum Number (main energy level) symbol  _______
16. Angular momentum Quantum Number (shape) symbol  ________
a. s
b. p
c. d
17. Magnetic Quantum Number (orientation)
a. s
b. p
c. d
18. Spin Quantum number—each orbital can hold ________ electrons ( _____________ )
19. Where are the s, p, d, and f blocks located on the periodic table?
20. Hund’s Rule
21. Pauli Exclusion Principle:
22. Heisenberg Uncertainty Principle:
23. Orbital Notation for bromine
24. Electron configuration for bromine
25. Noble gas configuration for bromine
26. How many valence electrons do these have? What element is it? How many electrons will it gain/lose?
a. 1s22s22p63s23p64s23d6
b. [Ne] 3s23p4
***To be stable the “p” and “d” orbitals must be ½ filled or complete filled.
27. Number of orbitals = n2
Ex. How many orbitals can the 4th energy level have?
28. Number of electrons = 2n2
Ex. How many electrons can the 6th energy level hold?
Unit 5
1. Mendeleev: _________________________________________.
2. Mosely: _________________________________________.
3. Periods ( ______________________ ) vs. Groups ( ___________________________________ )
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4. The periodic law: elements in the same ______________ have similar ____________________, the
same number of ___________________________, and the same ______________________________.
5. Cations are _________________ charged ions. They are formed by __________________ electrons.
6. Anions are _________________ charged ions. They are formed by __________________ electrons.
7. The periodic table is arranged by ______________________________________.
8. The periodic table is mostly __________________________.
9. Metals:
a. Where?
b. Gain or lose electrons?
c. Positive or negative ions?
d. Properties?
10. Non-metals:
a. Where?
b. Gain or lose electrons?
c. Positive or negative ions?
d. Properties?
11. Metalloids:
a. Where?
b. Which elements?
c. Properties?
12. Representative Elements (Main Group Elements)
a. Where?
13. Alkali Metals
a. Where?
b. Gain or lose electrons? (How many?)
c. Charge/Oxidation number?
d. Properties
14. Alkaline Earth Metals
a. Where?
b. Gain or lose electrons? (How many?)
c. Charge/Oxidation number?
15. Transition Metals
a. Where?
b. Gain or lose electrons? (How many?)
c. Charge/Oxidation number?
d. Properties
e. What do you use when naming?
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16. Halogens
a. Where?
b. Gain or lose electrons? (How many?)
c. Charge?
d. Properties
17. Noble Gases
a. Where?
b. Gain or lose electrons? (How many?)
c. Properties
18. Reactivity:
a. Metals more reactive closer to _______________
b. Nonmetals more reactive closer to ____________________
c. Group that is not reactive or inert? ____________________
19. Metallic Activity Trend
a. Definition
b. Increases towards ___________________
c. Least metallic group is ____________________
20. Valence Electrons—correspond to the ___________________________________ of the periodic table.
21. Atomic Radius (Covalent Radius) trend
a. Definition
b. Greatest atomic radius? __________________
c. Smallest atomic radius? __________________
d. Reasoning
e. Put the following in order from largest to smallest atomic radius—Ca, K, Cu, Se?
22. Ionization Energy
a. Definition
b. Greatest ionization energy? __________________
c. Smallest ionization energy? __________________
d. Reasoning
e. Put the following in order from lowest to highest ionization energy N, Bi, P, Sb?
23. Electronegativity
a. Definition
b. Greatest electronegativity? __________________
c. Smallest electronegativity? __________________
d. Reasoning
e. Put the following in order from greatest to lowest electronegativity Mg, Ra, Be, Ca?
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24. Ionic Radii
a. Definition
b. Cations
c. Anions
d. Reasoning
25. Always remember Trends will always be _________________...either at the very _________ or
______________... all the way to the ______________ or all the way to the ________________
Unit 6
1. Balance electrons are responsible for the element’s __________________________________________.
2. Why do atoms bond? _________________________________________________________________.
3. Metallic bonds involve only ________________. They are characterized by
____________________________ also called _______________________________________.
4. Ionic compounds involve ____________ and _______________. They are characterized by the
__________________ of electrons.
5. Covalent molecules involve _____________________. They are characterized by the ______________
of electrons.
6. Type of bonding and electronegativity differences.
Type
Electronegative Difference
Example
7. Diatomic molecules (magic 7…HOFBrINCl)
a. Make sure you know the seven.
b. All have _________________________________ bonding.
8. Octet rule states all atoms ____________________________________________________________.
9. Electron Dot Notation (valence electrons shown as dots only!)
a. C
b. Cl
10. Lewis structure
a. CCl4
b. CO2
c. NH3
(want – have)/2 = number of bonds
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11. The number of bonds between two atoms is related to the strength and length to the bond.
a. Single Bond…example F2
b. Double Bond…example O2
c. Triple Bond…example N2
12. Molecular Geometry and VSEPR…main shapes
Shape
# of Atoms Attached
a. Linear
Example
b. Bent
c. Trigonal pyramidal
d. Tetrahedral
13. How do you know if the molecule is polar or nonpolar?
a. Polar _________________________ and ____________________ lone pairs around center.
b. Nonpolar _________________________ and ____________________ lone pairs around center.
14. Intermolecular Forces (These are NOT bonds…they are FORCES of attraction!!)
Type
Strength/ Boiling Point
a. London Dispersion (van der Waals)
Example
b. Dipole-Dipole
c. Hydrogen Bonding
15. Properties of Bonds
Ionic Compounds
1.
Covalent Molecules
Metallic Bonds
2.
3.
4.
5.
6.
Unit 7
1. Binary Compound: ______________ elements. 2nd element ends in _____________.
2. Types of Naming
a. Ionic Bonding—between _______________ and __________________
i. Sodium Chloride
ii. Mg3N2
b. Covalent Bonding—between _______________ and _____________ (use _____________)
i. CCl4
ii. Dinitrogen pentoxide
c. Transition metals—between ________________ and _____________ (use _________________)
i. Iron (II) oxide
ii. Manganese (III) oxide
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d. Polyatomic Naming—involves at least one polyatomic…you have more than two capital
letters…look up all polyatomics in reference table!!)
i. Calcium hydroxide
ii. FeSO4
e. Acids—Hydrogen is in front…Still cross charges!!!
i. Two elements—Hydro____________ic
ii. More than two elements—depends on polyatomic
1. If polyatomic ends in –ate…acid ends in –ic
2. If polyatomic ends in –ite…acid ends in –ous
iii. Nitric
iv. Acetic
v. Hydrochloric
vi. Sulfurous
3. Percent composition by mass…Part/Whole x100
a. Calculate the percent by mass of water in silver sulfate pentahydrate
4. Empirical Formula ( _________________________ ) vs. Molecular Formula ( ___________________ )
a. A compound contains 32.38% Na, 22.65% S, and 44.99% O. What is the empirical formula?
b. A molecule has an empirical formula of CH4 and a molar mass of 48, what is the molecular
formula?
Unit 8
1. Balance equations follow the Law of Conservation of Mass, which states ________________________.
2. Ionic equations are written to show what actually happens.
a. Spectator Ions are ions which _________________________________________________.
b. Ionic equations show everything but __________________________________________.
3. Predict the products, balance and label what type of reaction each of the following are
a. ___ C3H8 + ___ O2 
type: ____________________________
b. ___ NaCl 
type: ____________________________
c. ___ H2 + ___ O2 
type: ____________________________
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d. ___ NaF(aq) + ___Pb(NO3)2(aq)
type: ____________________________
e. ___Ag + ___Ca(NO3)2 
type: ____________________________
f. ___HgO + Δ 
type: ____________________________
g. ___ SO2 + ___ H2O 
type: ____________________________
h. ___NaI(aq) + ___ Pb(C2H3O2)2(aq)
type: ____________________________
Unit 9
1. What do the following symbols mean?
a. (s)
b. (l)
c. (g)
d. (aq)
2. What is a mole fraction? _______________________________________________________________
3. Use the reaction to answer the following questions… (remember to be able to use the foldable you must
make it after the test has begun!)
___ N2(g) + ___ H2(g)  ___ NH3(g)
a. How many moles of nitrogen gas are required to produce 20 moles of ammonia (NH3)?
b. How many grams of hydrogen gas (H2) are required to fully react with 10 moles of N2?
c. What is the total number of moles of H2 used to produce 34 grams of NH3?
d. How many grams of H2 are required to react with 28 grams of N2?
e. How many liters of nitrogen are used to form 3.45g of ammonia?
f. How many grams of ammonia are formed when 5.67L of hydrogen react with excess nitrogen?
g. How many molecules of hydrogen are needed to fully react with 34.6moles of nitrogen?
h. How many liters of nitrogen are needed to make 4.3 moles of ammionia?
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i. How many moles of ammonia can 4.56x10 molecules of nitrogen make?
j. How many grams of hydrogen are needed to fully react with 4.56x1023 molecules of nitrogen?
k. How many liters of nitrogen are needed to fully react with 3.56x1024 molecules of hydrogen?
l. How many molecules of ammonia are made when 3.56x1024 molecules of hydrogen react?
m. How many molecules of ammonia can 56.7g of hydrogen produce?
n. What is the total number of moles of H2 used to produce 34 liters of NH3?
o. What is the total amount of grams of N2 needed to completely react with 64L of H2?
p. What is the total number of molecules of NH3 produced from 54.7 liters of nitrogen gas?
Unit 10
1. Kinetic Molecular Theory—based on _____________________!!!
2. Five characteristics of ideal gases:
1.
2.
3.
4.
5.
3. Ideal Gas ( ______________________________ ) vs. Real Gas ( ______________________________ )
4. Solubility of gases _______________ when temperature increases and ________________ when
pressure increases.
5. Diffusion (__________________________) vs. Effusion (____________________________________)
6. ____________ particles move faster than _______________ particles
7. Average Kinetic energy is directly proportional to ___________________________.
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8. Boyles’ Law: relates ___________________ and ___________________
a. Formula:
b. Relationship:
c. Graph:
d. A 400mL sample of gas at a pressure of 400torr is reduced to 150torr at a constant temperature.
What is the new volume of the gas?
9. Charles’ Law: relates ___________________ and ___________________
a. Formula:
b. Relationship:
c. Graph:
d. A 5.0L container of gas at 100oC is raised to a temperature of 200oC. What is the new
temperature?
10. Gay-Lussac’s Law: relates ___________________ and ___________________
a. Formula:
b. Relationship:
c. Graph:
d. A tire was at a pressure of 4.0atm and a temperature of 25oC. If the temperature is raised to
30oC, what is the new pressure?
11. Combined Gas Law
a. Formula:
b. A 10.L container of nitrogen gas is at a temperature of 100oC, 0.80atm. If the pressure is
decreased to 0.68atm and volume is increased to 15L, what is the new temperature?
12. Ideal Gas Law
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a. Formula:
b. How many grams of nitrogen gas are present in a 25.0L container at a pressure of 0.75atm and a
temperature of 80oC?
13. Molar Volume: ___________ L = 1 mole at ___________
a. How many moles are in 6.6L of nitrogen at STP?
b. What volume is occupied by 2.0 moles of nitrogen gas at STP?
14. Stoichiometry of gases (Liters to Liters)
a. Given the equation ___ C5H12 + ___ O2  ___ CO2 + ___ H2O. How many liters of water are
produced when 5.0L of C5H12 is burned?
15. Avogadro’s Law
a. Formula
b. Relationship
c. Graph
16. Dalton’s Law of Partial Pressure
a. Formula
b. A sample of nitrogen gas is collected over water. The pressure of the water is 21.1mmHg. What
is the pressure of the nitrogen gas if the atmospheric pressure is 785mmHg?
17. Vapor pressure of water is a function of temperature. ____________ temperature causes the vapor
pressure to ________________.
Unit 12
1. Know the differences between the states of matter.
Gases
Particles
Density
Fluidity
Compressibility
Liquids
Solids
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2. At a phase change, physical equilibrium occurs…both phases exist at the same time.
3. Energy and changes of state: Know the phase changes and if the change is endothermic and exothermic.
a. Freezing
b. Melting
c. Condensation
d. Evaporation
e. Sublimation
f. Deposition
4. Remember a phase change is a physical change!
5. Be able to read the phase diagram, identify regions, phases, and phase changes.
a. Critical Point: _________________________________________________________
b. Triple Point: __________________________________________________________
6. Molar Heat of Fusion: The energy required to change one mole of a ____________ to a ____________.
7. Molar Heat of Vaporization: The energy required to change one mole of a _________ to a __________.
a. How much heat energy is absorbed when 72g of ice melts at STP?
b. How many grams of water are used to release 5000J when turned into vapor?
8. Freezing and Cooling Curves. Know the different regions, phases, and phase changes.
9. Phase changes can occur with a change of _______________________ and/or ____________________.
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Unit 13
1. Solutions
a. __________________ is the dissolving substance
b. __________________ is the dissolving medium
c. __________________ can conduct electricity
2. Solution Equilibrium
a. Unsaturated
i. Definition:
ii. Dot is __________________ the line
b. Saturated
i. Definition:
ii. Dot is __________________ the line
c. Super-saturated
i. Definition
ii. Dot is __________________ the line
3. Solubility—know how to read the solubility curve and how to use solubility rules located in your
reference tables.
a. How many grams of sodium nitrate can be dissolved in 100g of water at 40oC?
b. At what temperature can 22 grams of potassium chlorate be dissolved in 100g of water?
c. What salt is the most soluble at 70oC?
d. What salt is the least soluble at 30oC?
e. What do the dotted lines mean?
4. Molarity
a. Formula
b. What is the molarity of a solution containing 2.0 grams of NaCl in 30mL of water?
c. How many grams must be dissolved to make 2.5L a 5.0M solution of HCl?
5. Dilution
a. Formula
b. How many liters of 3.00M HF solution must be used to dilute 5.0L of 6.00M HF?
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Unit 14
1. Colligative properties: properties that change when a solute is added.
2. Freezing Point Depression…the freezing point will ______________________ when a solute is added.
3. Boiling Point Elevation…the boiling point will ______________________ when a solute is added.
4. Vapor Pressure…the vapor pressure of a liquid will ______________________ when a solute is added.
Unit 15
1. Equilibrium:
2. ________ forward reaction = ________ reverse reaction
3. Dynamic Equilibrium: ________________ and ________________ are changing but because the
________ are ________ we can’t see any changes.
4. Equilibrium constant is symbolized by the ________ and is found by ________________ over
________ raised to the ________________. ________ and ________ are NOT included.
5. If K is larger than one…more ________
6. If K is smaller than one…more ________
7. Homogeneous Equilibrium:
8. Heterogeneous Equilibrium:
9. Write the equilibrium constant for the following:
a. Ammonia gas, NH3, decomposes into nitrogen gas and hydrogen gas.
b. ____ CO(g) + ____ H2(g)  ____ CH4(g) + ____ H2O(l)
10. For the reaction N2(g) + 3Cl2(g) ↔ 2NCl3(g), an analysis of an equilibrium mixture is preformed at a
certain temperature. It is found that [NCl3] = 0.19M, [N2] = 0.0014M, and [Cl2] = 4.3x10-3M. Calculate
K for the reaction.
11. If the following equilibrium concentrations are found for the reaction N2(g) + 3H2(g) ↔ 2NH3(g), what
is the equilibrium constant?
[NH3] = 0.031M
[N2] =0.085M
[H2] = 0.0031M
12. Tell whether the equilibrium will shift to the right, will shift to the left, or not be affected.
UO2(s) + 4HF(g) ↔ UF4(g) + 2H2O(g)
a. UF4(g) is added.
b. HF(g) is removed.
c. Pressure is increased.
13. Tell whether the equilibrium will shift to the right, will shift to the left, or not be affected.
2SO3(g) + Heat ↔ 2SO2(g) + O2(g)
a. Oxygen gas is added.
b. The pressure is increased by decreasing the volume of the reaction container.
c. The temperature is decreased.
d. Gaseous sulfur dioxide is removed.
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Unit 16
1. Electrolyte
a. Definition
b. Things that are electrolytes:
2. Nonelectrolyte
a. Definition
b. Things that are nonelectrolytes:
3. Concentrated (__________________________) vs. Dilute (_____________________________)
4. Strong Acid/Base (_________________________) vs. Weak Acid/Base(________________________)
5. Make sure you know the different properties of acids and bases.
Acid
Base
pH
Proton
Produce ______ in water
Taste
Phenolphthalein Indicator
Litmus Paper Indicator
Example
Reacts with metal to form:
6. Arrhenius Acid:
Arrhenius Base:
7. Brønsted-Lowry Acid:
Brønsted-Lowry Base:
8. Label the following and state the conjugate acid/base pairs: HCl + NH3  NH4+ + Cl9. pH scale/pH/pOH
a. pH of a solution is the negative of the common logarithm of the hydronium ion concentration
b. pOH of a solution is the negative of the common logarithm of the hydroxide ion concentration
c. Formulas (remember these are in your reference tables)
i. pH = -log[H3O+]
ii. pOH = -log[OH-]
iii. Kw = [H3O+] [OH-] = 1.0x10-14
iv. pH + pOH =14
v. [H3O+] = 10-pH
vi. [OH-] = 10-pOH
d. What is the pH of a substance that has a [H3O+] of 1.0x10-2?
e. What is the pH of a solution whose HCl concentration is 1.0x10-1?
f. What is the pH of a substance that has a [H+] concentration of 1.0x10-4?
g. What is the pOH of a substance that has an [OH-] of 1.0x10-3?
h. What is the pH of a solution whose NaOH concentration is 1.0x10-2?
i. What is the pOH of a substance that has a [OH-] of 1.0x10-3? Is this solution acidic or basic?
j. What is the [H3O+], [OH-], pOH of a solution which has a pH of 4?
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+
10. As the pH of a solution increase the concentration of [H ] increases by _______________.
a. The concentration of [H+] increases by ________ when the pH increases from 3.0 to 5.0.
11. Neutralization
a. Acid + Base  ________ + ________
b. HCl + NaOH  ________ + ________
i. Net Ionic Equation:
c. HBr + Ca(OH)2  ________ + ________
i. Net Ionic Equation:
12. A strong acid produces a ________________ conjugate ________. A strong base produces a
________________ conjugate ________.
13. Strong Acid/Base Titration
a. nM1V1 = nM2V2
b. In a titration, 25mL of a 2.0M solution of HCl and phenolphthalein indicator is placed in a flask.
NaOH was added to the buret with an initial reading of 10.mL. The NaOH was then titrated into
the HCl solution until it turned pink. The final reading of NaOH was 50.mL. What was the
molarity of the NaOH?
c. Vinegar is mainly made of acetic acid. Determine the volume of 0.832M acetic acid (HC2H3O2)
that
requires 22.7mL of a 0.550M solution of NaOH to reach the equivalence point.
14. Titration Curve for Strong Acid and Strong Base
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Unit 17
1. Specific Heat
a. Definition
b. Formula
c. A 5.0g sample was heated from 200K to 500K and was found to have absorbed 40J of heat.
What is the specific heat capacity of the sample?
d. How much heat is needed to raise the temperature of 5.0g of gold by 25oC?
e. Joule is defined as ______________________________________________________________
2. Heat (__________________________________) vs. Temperature (_____________________________)
3. Kinetics…Collision Theory
a. Particles must ______________
b. Particle must enough ________________
c. Particles must collide at the correct ________________
4. Reaction rate is proportional to the number of effect collisions
a. Increase temperature = __________________ number of collisions = ____________ rate
b. Increase concentration = __________________ number of collisions = ____________ rate
c. Increase pressure = __________________ number of collisions = ____________ rate
d. Increase surface area = __________________ number of collisions = ____________ rate
5. Entropy (____________________________) vs. Enthalpy (_________________________________)
a. Entropy increases with a change in state
b. Solid  Liquid  Gas
(Dissolving increases entropy)
6. Nature likes ________________ entropy and _________________ enthalpy.
7. What are the four ways to effect rate?
a. _______________________
b. _______________________
c. _______________________
d. _______________________
8. For a closed system, energy is neither lost nor gained. The total useful energy of an open system is
declining due to entropy.
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9. Potential Energy Diagram
a. Activation Energy = ___________
b. Enthalpy Change = ___________
c. Potential Energy of products = ___________
d. Potential Energy of reactants = ___________
e. Exothermic or Endothermic = ___________
Unit 18
1. Know the properties of the three different nuclear decay particles:
Alpha
Beta
Symbol
Charge
Mass
Strength
Blocked by
Attracted To…
Gamma
2. Nuclear Fission: ___________________________ of a nucleus into ___________________________
3. Nuclear Fusion: ___________________________ of a nucleus into _____________________________
4. Balancing Nuclear Reactions (remember Law of Conservation of Mass)
a. 3717Cl + 11H  _________
b.
64
30Zn
+ _________  6329Cu + 10n
c.
235
92U
 9438Sr + _________
d. _________  13351Sb + 9841Nb
e.
6
3Li
+ 10n  42He + _________
5. Decay is a ___________________________ event, independent of other energy influences.
6. Half-life is __________________________________________________________________________
a. Element-106(Seaborgium) has a half-life of 0.90 seconds. If one million atoms of it were
prepared, how many atoms would remain after 4.5 seconds?
b. Iron-59 is used in medicine to diagnose blood circulation disorders. The half-life of iron-59 is
44.5 days. How much of a 2.000mg sample will remain after 133.5 days?
c. If the passing of five half-lives leaves 25.0mg of a strontium-90 sample, how much was present
in the beginning?