Download Academic Chemistry Final Exam Review

Name ____________________________________________________
Period _____
Academic Chemistry Final Exam Review
Matter – Chapter 2
Important Concepts:
o Classification of Matter: Pure Substances (elements, compounds) vs. Mixtures (homogeneous, heterogeneous)
o States of matter (description)
o Physical vs. chemical changes/properties
Review questions:
1. Classify the following as chemical or physical changes (PC or CC)
_____ freezing of milk
_____ drawing (shaping) of copper into wire
_____ frying an Egg
_____ rusting of an iron nail
2. What are similarities and differences between elements and compounds?
3. Describe what happens when a compound undergoes a change in physical state (from solid to liquid).
Is this a chemical process or a physical process?
Chemical Foundations, elements, atoms, and ions – Chapter 4
Important Concepts:
o Element symbols (calculate electrons, protons, neutrons)
o Structure of the atom (old ideas vs. modern atomic theory)
o Scientists who impacted Atomic Structure (Democritus, Dalton, Rutherford, Thomson, Bohr)
o Isotopes – calculations and symbols
o Ions – (polyatomic and single atom) (charges and relate to atomic structure)
o Parts of the Periodic Table – groups, families, metals, nonmetals – and characteristics
o Average atomic mass determination (mass on the periodic table)
Review Problems:
1.
Element
(atom)
Symbol
Atomic
Number
# of protons
17
sodium
# of neutrons
Mass #
18
Na
# of electrons
Charge
18
23
oxide ion
16
79
11
118
1+
22
11
2. According to the table above, which two particles are isotopes of one another? ____________________________
Why? ______________________________________________________________________________________
1
3. Describe the location on the periodic table for each of the following:
a. Metals __________________________________________________________________________________
b. Nonmetals _______________________________________________________________________________
c. Metalloids _______________________________________________________________________________
d. Noble Gases _____________________________________________________________________________
4. Describe the impact of each of the following scientists. Include contributions, specific experiments, discoveries,
theories and/or proposed models. Know the details!
a. Democritus ____________________________________________________________________________
______________________________________________________________________________________
b. John Dalton ____________________________________________________________________________
________________________________________________________________________________________
________________________________________________________________________________________
c. Ernest Rutherford _________________________________________________________________________
________________________________________________________________________________________
________________________________________________________________________________________
d. JJ Thomson ______________________________________________________________________________
________________________________________________________________________________________
________________________________________________________________________________________
e. Niels Bohr _______________________________________________________________________________
________________________________________________________________________________________
________________________________________________________________________________________
Nomenclature – Chapter 9
Important Concepts:
o Type I naming – binary ionic compounds
o Type II naming – binary ionic with transition metals
o Type III naming – nonmetal to nonmetal (covalent)
Review Questions:
1. Write formulas for the following compounds
a. lead (II) oxide ____________________
○ Acids
○ Polyatomic ions
d. magnesium chloride ______________
b. sodium phosphate __________________
e. sulfuric acid _____________________
c. mercury (I) chloride ________________
f. hydrochloric acid _________________
2. Name the following compounds:
a. CaF2 ____________________________
d. CuBr ___________________________
b. Al(NO3)3 _________________________
e. Fe2O3 ___________________________
c. AgI ____________________________
f.
PbO2 ____________________________
2
g. CO _____________________________
j.
h. CO2 ____________________________
k. H2SO4 __________________________
i.
l.
N2O ____________________________
Measurement and Calculations – Chapter 3
Important Concepts:
○ Accuracy and Precision
o Units of mass, volume, length, density, etc.
o Metric relationships
BF3 ____________________________
HBr ____________________________
○ Temperature Conversions
○ Sig Figs rules (add/subtract, multiply/divide)
Review Questions:
1. Complete the following statements:
a. EXAMPLE: The unit abbreviation “m” stands for ____meter___ and is a unit of ___length___.
b. The unit abbreviation “g” stands for ________________ and is a unit of _______________.
c. The unit abbreviation “mL” stands for _____________________ and is a unit of ________________.
d. The unit abbreviation “K” stands for __________________ and is a unit of _____________________.
e. The unit abbreviation “J” stands for _________________ and is a unit of ____________________.
f. The unit abbreviation “g/mL” stands for ______________________ and is a unit of ________________.
g. The unit abbreviation “atm” stands for _____________________ and is a unit of __________________.
2. A group of students is determining the boiling point of a compound. The following measurements were recorded:
110.2°C, 110.5°C, 109.9°C, and 110.0°C. The accepted value is 110.1°C. Describe the accuracy and precision of
these measurements.
3. Perform the following conversions:
a. 75.5°C = ____________K
b. 323 K = ____________°C
d. 1.50 atm = ______________ torr
c. 275 mL = ___________L
e. 782 torr = ________________ atm
Chemical Composition – Chapter 10
Important Concepts:
o The mole - relationship to Avogadro’s number and mass
o Molar Mass
○ Empirical formulas calculation
○ Molecular formulas from empirical formulas
Review Problems:
1. Calculate the number of atoms of aluminum in 10.0g of aluminum.
2. How many moles of C10H6O3 are present in a 1.56 g sample?
3. A compound used as an additive for gasoline to help prevent engine knock shows the following percentage
composition: 24.27 % C, 4.07% H, and 71.65% Cl. The molar mass of the compound is known to be 98.96 g.
Determine the empirical formula and the molecular formula of the compound.
3
Chemical Reactions, an introduction – Chapter 11 & 20
Important Concepts:
o Writing a chemical equation
○ State symbols
○ Balancing equations
Review Questions:
1. Balance the following reactions:
a. ___K(s) + ___H2O(l) → ___H2(g) + ___KOH(aq)
b. ___NH3(g) + ___O2(g) → ___NO(g) + ___H2O(l)
c.
___SiO2(s) + ___HF(aq) → ___SiF4(g) + ___H2O(l)
2. Write the following reactions. Remember to balance and include states of matter.
a. When solid ammonium nitrate is heated, it produces nitrogen gas, oxygen gas, and water vapor.
b. Gaseous nitrogen monoxide decomposes to produce dinitrogen monoxide gas and nitrogen dioxide gas.
Reactions in Aqueous Solutions - (Chapter 11 & 20 continued)
Important Concepts:
o Recognizing types of reactions (solid formation, water, gas, redox)
o Precipitation reactions – Solubility Rules
o Acid and base/Neutralization reactions (salt + water products)
o Predicting products of all reaction types - decomposition, synthesis, precipitation, neutralization, single
replacement (Activity Series), combustion
Review Questions:
1. Write the balanced chemical equation with states of matter to predict the products formed when the following
aqueous solutions undergo double displacement reactions. For each reaction, identify the type of double
displacement reaction (Acid-Base or Precipitation) and identify the spectator ions by symbol and charge.
a. HNO3 and Ba(OH)2
Reaction type: ______________________________
Spectator Ions: ________________
b. Na2SO4 and Pb(NO3)2
Reaction type: ______________________________
Spectator Ions: ________________
2. Predict the products and balance the following reactions. Identify the type of reaction in the space provided
(Single displacement, Synthesis, Decomposition, Combustion, or No Reaction).
a. K (s) + Cl2 (g) →
Reaction Type: _____________________________
b. Fe2O3 (s) + Al (s) →
Reaction Type: _____________________________
c. CH4 (g) + O2(g) →
Reaction Type: _____________________________
4
d. Cu (s) + HCl (aq) →
Reaction Type: _____________________________
e. H2O (l) →
Reaction Type: _____________________________
Chemical Quantities and Stoichiometry – Chapter 10 & 12
Important Concepts:
o Mole to mole ratios
o Stoichiometry calculations with and without Limiting Reactant
o Percent yield (% = [actual/theoretical] x 100)
Review Questions:
1. If 3 moles of chlorine gas are reacted with excess sodium iodide, how many moles of I2 will be formed? (balance
first!)
Cl2 + NaI → NaCl + I2
2. If 72.4 g of Magnesium metal are reacted with excess O2 gas, how many moles of magnesium oxide are formed?
Mg + O2  MgO
3. Lithium nitride is prepared by the reaction of lithium metal and nitrogen gas. Calculate the mass of lithium
nitride formed when 56.0g of nitrogen gas and 56.0g of lithium are reacted in the unbalanced equation below.
(Hint: you must determine the limiting reactant)
Li (s) + N2 (g) → Li3N (s)
a. Calculate the mass left over of the excess reactant.
b. What is the percent yield if 85.2 g of lithium nitride is experimentally produced?
5
Modern atomic theory – Chapter 5 and Chapter 6
Important Concepts:
o Emission of light by atoms (flame testing)
o Ground state vs. Excited state
o Quantum mechanical model vs. Bohr model of the atom
o Electron Configuration (energy levels, sublevels, orbitals)
o Orbital diagrams (Hund’s Rule, Aufbau Principle, Pauli Exclusion Principle)
o Periodic Trends (electronegativity, atomic radius, ionization energy, metallic character)
o Valence Electrons
Review Questions:
1. Write the electron configuration and the orbital diagram for Iron (Fe).
2. Write the Noble gas configuration for Nickel (Ni).
3. Explain Hund’s Rule.
4. Explain the Aufbau Principle.
5. Explain the Pauli Exclusion Principle.
6. Describe the periodic trends associated with each of the following properties (in groups and periods):
a. Electronegativity – _____________________________________________________________________
b. Atomic radius – _______________________________________________________________________
c. Ionization energy – _____________________________________________________________________
d. Metallic character - ____________________________________________________________________
7. Which of the periodic trends from question #6 above both generally decrease when the elements in Period 4 on
the Periodic Table are considered in order from left to right?
8. Explain how an atom in the ground state becomes “excited”.
9. What happens when an excited electron “drops” back down into its ground state?
10. How many valence electrons are in an atom of calcium (in the ground state)? ______
How many valence electrons are in an atom of chlorine (in the ground state)? ______
6
Chemical Bonding – Chapter 8
Important Concepts:
o Types of bonds (ionic, covalent, polar, nonpolar)
o Electronegativity differences
o Intermolecular forces
o Lewis Dot structures (ionic and covalent) (multiple bonds) (N.A.S)
o Molecular Geometry (shape): tetrahedral, linear, etc…
o Molecular polarity and Bond polarity
Review Questions:
1. What type of bonding is found in Noble Gases? Why?
2. Using electronegativity values found on the back of your periodic table, determine if the following bonds are
nonpolar covalent, ionic, or polar covalent
a. S-S ______________________
b. S-O ______________________
c. S-H ______________________
d. S-K ______________________
3. Draw the Lewis Structure for O2. How many electrons are being shared? What is the shape of this molecule?
4. Draw the Lewis Structure for CO2. How many electrons are being shared? What is the shape of this molecule?
5. Draw the Lewis Structure for CCl4. How many electrons are being shared? What is the shape of this molecule?
Gases – Chapter 14
Important Concepts: (formulas will be given to you)
o Pressure units
○ Kinetic Molecular Theory
o Boyle’s Law
○ Gas Stoichiometry
o Charles’s Law
○ Molar volume (22.4 L = 1 mole)
o Ideal Gas Law (PV = nRT)
○ STP
o Combined Gas Law
Review Questions:
1. An engine cylinder’s volume is 0.75 L and the air-fuel mixture inside it is 1.00 atm. Calculate the pressure of the
gas in the cylinder after the piston compresses the mix to 0.075 L.
2. A sample of gas at 15oC and 1.00 atm has a volume of 2.58 L. The temperature is then raised to 38oC at 1.00 atm.
What is the new volume of the gas?
7
3. What volume is occupied by 0.250 mol of carbon dioxide gas at 25oC and 371 torr?
4. A sample of diborane gas (B2H2) a substance that bursts into flames when exposed to air, has a pressure of
0.454 atm at a temperature of -15oC and a volume of 3.48L. If conditions are changes to those of STP, what will
be the new volume of the sample?
5. Calculate the volume of hydrogen produced at 1.50 atm and 19oC by the reaction of 26.5g of zinc with excess
hydrochloric acid according to the balanced equation: (gas stoichiometry)
Zn +2HCl → ZnCl2 +H2
6. What are the statements of the Kinetic Molecular Theory?
Solutions – Chapter 16
Important Concepts:
o Dissolving (describe)
o Like dissolves like (molecular polarity)
o Factors that affect dissolving
○ Mass percent (% = 100 x [mass solute/mass solution])
○ Molarity (M = moles solute/liters solution)
○ Solution Stoichiometry
Review Questions:
1. Chemists often say “like dissolves like”. What does this statement mean?
2. What are the 3 factors that influence the rate of dissolving?
3. A solution is prepared by mixing 1.00 g of ethanol, C2H5OH, with 100.0 g of water. Calculate the mass percent
of ethanol in this solution.
8
4. Calculate the molarity of a 1.25 L solution that contains 1.56 g of hydrochloric acid, HCl.
5. How much solid K2Cr2O7 (Molar mass = 294.20 g) must be weighed to make 1.00 L of a 0.200 M solution?
6. What volume of 12 M HCl must be taken to prepare 0.75 L of 0.25M HCl?
7. Hydrochloric acid reacts with calcium hydroxide in an acid-base reaction. If 15.0 mL of a 3.0 M HCl solution
reacts with an excess amount of calcium hydroxide, what mass of the salt product can we predict to form?
(solution stoichiometry)
HCl + NaOH  NaCl + H2O
8. What are we able to calculate using titration?
Nuclear Chemistry – Chapter 25
Important Concepts:
o Types of Decay
o Fission and Fusion
o Half Life
Review Questions:
1. What are the 3 types of radioactive decay?
2. What is fission? What is fusion? Give an example of each.
3. How long would it take for 640.0g of Bismuth-213 to decay down to 40.0g if it has a half-life of 19.7 minutes?
Thermochemistry and Phase Changes– Chapter 17
Important Concepts:
o Endothermic and exothermic reactions
o Enthalpy
o
Heating/Cooling Curves
○ Calorimetry (heat lost = heat gained)
○ Heats of Fusion and Vaporization
○ Phase Changes
Review Questions:
1. What happens during an exothermic reaction? What is the sign of the H value?
9
2. What happens during an endothermic reaction? What is the sign of the H value?
3. What is a change in state from a liquid to a gas called?
4. What is sublimation?
5. Use this heating curve to answer the questions below.
a.
What is the melting point (temperature) of the substance above?
b.
What is the freezing point (temperature) of the substance above?
c.
Does it take longer for the substance to melt or boil (vaporize)?
d.
In what state does the substance exist at room temperature?
6. How much energy would it take to completely melt a 23.6 g sample of ice at 0°C? The heat of fusion for H2O is
6.02 kJ/mol.
7. How much energy would be produced if 2.50 moles of methane were reacted with excess O2 gas in the reaction
below. The ∆H of the reaction is -883.0 kJ of energy (per mole of methane reacted).
CH4 (g) + 2 O2 (g) → CO2 (g) + 2 H2O (g)
H = -883.0 kJ
10