Chapter 1 Chemistry Review Key
1) Chemistry Review
Key Terms: empirical knowledge theoretical knowledge law of conservation of mass coefficient chemical amount mole Key Concepts:
Write chemical equations when given reactants and products (1.5, 1.6)
Write balanced chemical equations (2.2, 2.3)
Interpret balanced chemical equations in terms of chemical amount (in moles) (2.3)
Convert between chemical amount and mass (2.4)
Classify chemical reactions (2.5, 2.6)
Predict the solubility of elements and ionic and molecular compounds in water (2.6)
Predict products for chemical reactions (2.5, 2.6)
Memorized molecular compounds pg. 34
Sample Questions:
Chapter One Review pg. 40 #11-19, 25 Chapter Two Review pg. 66 #5-7, 9
Unit Review pg. 68 #1-8, 11-25 …show more content…
Are You Ready pg. 74 #2-8
2) Chemical Bonding
Key Terms: structural formula valence electron orbital valence orbital bonding electron lone pair octet rule
Lewis formula electronegativity covalent bond ionic bond bonding capacity empirical formula molecular formula structural formula
VSEPR theory polar molecule nonpolar molecule nonpolar covalent bond polar covalent bond bond dipole dipole–dipole force hydrogen bond crystal lattice covalent network intermolecular force
London force
Key Concepts:
Define valence electron, electronegativity, and ionic bond (3.1, 3.3)
Use the periodic table and Lewis structures to support and explain ionic bonding theory (3.1)
Explain how an ionic bond results from the simultaneous attraction of oppositely charged ions (3.1)
Relate electron pairing to covalent bonds (3.1, 3.2)
Draw electron-dot diagrams (Lewis symbols/formulas) of atoms, molecules and ions, writing structural formulas for molecular substances & using Lewis structures to predict bonding in molecules (3.2)
Apply VSEPR theory to predict molecular shapes (3.3)
Illustrate, by drawing, the structure of simple molecular substances (3.2)
Explain intermolecular forces, London forces, dipole–dipole attractions, and hydrogen bonding (3.4)
Relate properties of substances to the predicted intermolecular bonding in the substance (3.4, 3.5) Determine the polarity of a molecule based on structural shapes and unequal charge distribution (3.3) describe bonding as a continuum ranging from electron transfer to equal sharing of electrons. (3.3, 3.4)
Sample Questions:
Unit Review pg. 137 #3,4,6-17, 25,26,28, 32,39,41,43,47-51
3) Gases
Key Terms: pressure atmospheric pressure
STP
SATP
Boyle’s law absolute zero absolute temperature scale
Charles’ law combined gas law law of combining volumes
Avogadro’s theory molar volume ideal gas ideal gas law universal gas constant (R)
Key Concepts:
Convert between the Celsius and absolute (kelvin) temperature scales (4.1, 4.4)
Describe the behaviour of real and ideal gases in terms of kinetic molecular theory (4.2, 4.4)
Explain the law of combining volumes (4.2)
Illustrate how Boyle’s, Charles’, and combined gas laws are related to the ideal gas law (4.4)
Perform calculations using Boyle’s, Charles’ and the Combined Gas Law (4.4)
Perform calculations based on the ideal gas law under STP, SATP, and other conditions (4.4)
Sample Questions: Pg. 159 #20-23, Pg. 166 #7, Pg. 171 #5-7, Pg. 176 #5, 9 Unit Review pg. 181 #1-10, 12-18, 20, 21, 26-29
4) Solutions, Acids and Bases
Key Terms: solute solvent electrolyte non-electrolyte dissociation ionization amount concentration standard solution stock solution saturated solution solubility dynamic equilibrium hydronium ion pH/pOH acid−base indicator neutralization strong and weak acid strong and weak base monoprotic acid or base polyprotic acid or base
Key Concepts:
Explain the difference between dissociation and ionization (5.2)
Differentiate between electrolytes and non-electrolytes (5.1, 5.2)
Understand the major entities present when any substance is in a water environment (5.2: Tables 2,3)
Express concentration in various ways (%, ppm, amount concentration) (5.3)
Perform calculations involving concentration, chemical amount, volume and/or mass (5.3)
Use dissociation equations to calculate ion concentration (5.3)
Describe the procedures/calculations needed for preparing solutions from a pure solid or dilution (5.4)
Define solubility and identify the factors that affect it (5.5)
Explain a saturated solution in terms of equilibrium (5.5)
Recall the empirical definitions of acidic, basic, and neutral solutions determined by using indicators, pH, and electrical conductivity (6.1)
Calculate H3O+(aq) and OH−(aq) concentrations, pH, and pOH of acid and base solutions based on log expressions (6.2)
Use appropriate SI units to communicate the concentration of solutions and express pH and concentration to the correct number of significant digits (6.2)
Solutions, Acids and Bases (continued)
Compare magnitude changes in pH and pOH with changes in concentration for acids and bases (6.2)
Explain how the use of indicators, pH meters or pH paper can be used to measure [H3O+(aq)] (6.3)
Use the modified Arrhenius theory to define acids as substances that produce H3O+(aq) in aqueous solutions and bases as substances that produce OH−(aq) in aqueous solutions, and recognize that the definitions are limited (6.4)
Define neutralization as a reaction between hydronium and hydroxide ions (6.4)
Differentiate between strong acids and bases and weak acids and bases, qualitatively, using the modified Arrhenius (reaction with water) theory and dissociation (6.5)
Compare the reaction with water (ionization) of monoprotic with that of polyprotic acids and bases (6.5)
Sample Questions: Pg. 202 #8 Pg. 214 #3-9 Pg. 219 #3-5 Chapter 5 Review pg. 231 #2-13, 24-26, 32
Chapter 6 Review pg. 263 #1-11, 14 Unit Review pg. 265-268 #3-11, 23-25, 30, 31,
5) Stoichiometry and Chemical Analysis
Key Terms: net ionic equation spectator ion limiting reagent excess reagent theoretical yield gravimetric stoichiometry percent yield gas stoichiometry solution stoichiometry colorimetry gravimetric analysis titration analysis titration titrant sample equivalence point endpoint Key Concepts:
Identify limitations and assumptions about chemical reactions (7.1)
Write balanced ionic and net ionic equations, including identification of spectator ions, for reactions taking place in aqueous solutions (7.1)
Recognize limiting and excess reagents in chemical reactions (7.1, 7.2, 7.3, 7.4)
Calculate quantities of reactants and/or products involved in chemical reactions using gravimetric, solution, or gas stoichiometry (7.2, 7.3, 7.4)
Define predicted (theoretical) & experimental (actual) yields (7.2, 7.3)
Identify sources of experimental uncertainty in experiments (7.2, 7.3, 7.4)
Contrast quantitative and qualitative chemical analysis (8.1)
Use the stoichiometric method to calculate quantities of substances in chemical reactions (8.2, 8.3, 8.4)
Identify and calculate limiting and excess reagents in chemical reactions (8.3)
Identify the equivalence point on a strong acid–strong base titration curve, and differentiate between an indicator endpoint and a reaction equivalence point (8.4, 8.5)
Describe the function and choice of indicators in acid–base titrations (8.4, 8.5)
Sample Questions: Chapter 7 Review pg.
309 #2-15, 19-25 Chapter 8 Review pg. 347 # 1-17,
19
Unit Review pg. 349 #1-5, 7-11, 14-17, 26, 27, 31-33
Key Terms: empirical knowledge theoretical knowledge law of conservation of mass coefficient chemical amount mole Key Concepts:
Write chemical equations when given reactants and products (1.5, 1.6)
Write balanced chemical equations (2.2, 2.3)
Interpret balanced chemical equations in terms of chemical amount (in moles) (2.3)
Convert between chemical amount and mass (2.4)
Classify chemical reactions (2.5, 2.6)
Predict the solubility of elements and ionic and molecular compounds in water (2.6)
Predict products for chemical reactions (2.5, 2.6)
Memorized molecular compounds pg. 34
Sample Questions:
Chapter One Review pg. 40 #11-19, 25 Chapter Two Review pg. 66 #5-7, 9
Unit Review pg. 68 #1-8, 11-25 …show more content…
Are You Ready pg. 74 #2-8
2) Chemical Bonding
Key Terms: structural formula valence electron orbital valence orbital bonding electron lone pair octet rule
Lewis formula electronegativity covalent bond ionic bond bonding capacity empirical formula molecular formula structural formula
VSEPR theory polar molecule nonpolar molecule nonpolar covalent bond polar covalent bond bond dipole dipole–dipole force hydrogen bond crystal lattice covalent network intermolecular force
London force
Key Concepts:
Define valence electron, electronegativity, and ionic bond (3.1, 3.3)
Use the periodic table and Lewis structures to support and explain ionic bonding theory (3.1)
Explain how an ionic bond results from the simultaneous attraction of oppositely charged ions (3.1)
Relate electron pairing to covalent bonds (3.1, 3.2)
Draw electron-dot diagrams (Lewis symbols/formulas) of atoms, molecules and ions, writing structural formulas for molecular substances & using Lewis structures to predict bonding in molecules (3.2)
Apply VSEPR theory to predict molecular shapes (3.3)
Illustrate, by drawing, the structure of simple molecular substances (3.2)
Explain intermolecular forces, London forces, dipole–dipole attractions, and hydrogen bonding (3.4)
Relate properties of substances to the predicted intermolecular bonding in the substance (3.4, 3.5) Determine the polarity of a molecule based on structural shapes and unequal charge distribution (3.3) describe bonding as a continuum ranging from electron transfer to equal sharing of electrons. (3.3, 3.4)
Sample Questions:
Unit Review pg. 137 #3,4,6-17, 25,26,28, 32,39,41,43,47-51
3) Gases
Key Terms: pressure atmospheric pressure
STP
SATP
Boyle’s law absolute zero absolute temperature scale
Charles’ law combined gas law law of combining volumes
Avogadro’s theory molar volume ideal gas ideal gas law universal gas constant (R)
Key Concepts:
Convert between the Celsius and absolute (kelvin) temperature scales (4.1, 4.4)
Describe the behaviour of real and ideal gases in terms of kinetic molecular theory (4.2, 4.4)
Explain the law of combining volumes (4.2)
Illustrate how Boyle’s, Charles’, and combined gas laws are related to the ideal gas law (4.4)
Perform calculations using Boyle’s, Charles’ and the Combined Gas Law (4.4)
Perform calculations based on the ideal gas law under STP, SATP, and other conditions (4.4)
Sample Questions: Pg. 159 #20-23, Pg. 166 #7, Pg. 171 #5-7, Pg. 176 #5, 9 Unit Review pg. 181 #1-10, 12-18, 20, 21, 26-29
4) Solutions, Acids and Bases
Key Terms: solute solvent electrolyte non-electrolyte dissociation ionization amount concentration standard solution stock solution saturated solution solubility dynamic equilibrium hydronium ion pH/pOH acid−base indicator neutralization strong and weak acid strong and weak base monoprotic acid or base polyprotic acid or base
Key Concepts:
Explain the difference between dissociation and ionization (5.2)
Differentiate between electrolytes and non-electrolytes (5.1, 5.2)
Understand the major entities present when any substance is in a water environment (5.2: Tables 2,3)
Express concentration in various ways (%, ppm, amount concentration) (5.3)
Perform calculations involving concentration, chemical amount, volume and/or mass (5.3)
Use dissociation equations to calculate ion concentration (5.3)
Describe the procedures/calculations needed for preparing solutions from a pure solid or dilution (5.4)
Define solubility and identify the factors that affect it (5.5)
Explain a saturated solution in terms of equilibrium (5.5)
Recall the empirical definitions of acidic, basic, and neutral solutions determined by using indicators, pH, and electrical conductivity (6.1)
Calculate H3O+(aq) and OH−(aq) concentrations, pH, and pOH of acid and base solutions based on log expressions (6.2)
Use appropriate SI units to communicate the concentration of solutions and express pH and concentration to the correct number of significant digits (6.2)
Solutions, Acids and Bases (continued)
Compare magnitude changes in pH and pOH with changes in concentration for acids and bases (6.2)
Explain how the use of indicators, pH meters or pH paper can be used to measure [H3O+(aq)] (6.3)
Use the modified Arrhenius theory to define acids as substances that produce H3O+(aq) in aqueous solutions and bases as substances that produce OH−(aq) in aqueous solutions, and recognize that the definitions are limited (6.4)
Define neutralization as a reaction between hydronium and hydroxide ions (6.4)
Differentiate between strong acids and bases and weak acids and bases, qualitatively, using the modified Arrhenius (reaction with water) theory and dissociation (6.5)
Compare the reaction with water (ionization) of monoprotic with that of polyprotic acids and bases (6.5)
Sample Questions: Pg. 202 #8 Pg. 214 #3-9 Pg. 219 #3-5 Chapter 5 Review pg. 231 #2-13, 24-26, 32
Chapter 6 Review pg. 263 #1-11, 14 Unit Review pg. 265-268 #3-11, 23-25, 30, 31,
5) Stoichiometry and Chemical Analysis
Key Terms: net ionic equation spectator ion limiting reagent excess reagent theoretical yield gravimetric stoichiometry percent yield gas stoichiometry solution stoichiometry colorimetry gravimetric analysis titration analysis titration titrant sample equivalence point endpoint Key Concepts:
Identify limitations and assumptions about chemical reactions (7.1)
Write balanced ionic and net ionic equations, including identification of spectator ions, for reactions taking place in aqueous solutions (7.1)
Recognize limiting and excess reagents in chemical reactions (7.1, 7.2, 7.3, 7.4)
Calculate quantities of reactants and/or products involved in chemical reactions using gravimetric, solution, or gas stoichiometry (7.2, 7.3, 7.4)
Define predicted (theoretical) & experimental (actual) yields (7.2, 7.3)
Identify sources of experimental uncertainty in experiments (7.2, 7.3, 7.4)
Contrast quantitative and qualitative chemical analysis (8.1)
Use the stoichiometric method to calculate quantities of substances in chemical reactions (8.2, 8.3, 8.4)
Identify and calculate limiting and excess reagents in chemical reactions (8.3)
Identify the equivalence point on a strong acid–strong base titration curve, and differentiate between an indicator endpoint and a reaction equivalence point (8.4, 8.5)
Describe the function and choice of indicators in acid–base titrations (8.4, 8.5)
Sample Questions: Chapter 7 Review pg.
309 #2-15, 19-25 Chapter 8 Review pg. 347 # 1-17,
19
Unit Review pg. 349 #1-5, 7-11, 14-17, 26, 27, 31-33