TENTATIVE COURSE SCHEDULE:
The following list of chapters (from Tro, Chemistry: A Molecular Approach) is only intended to provide a rough outline of the course content. Your instructor will provide you with a more specific lecture schedule including exam dates.
Chapter 1: Matter, Measurement, and Problem Solving
Chapter 2: Atoms and Elements
Chapter 3: Molecules and Compounds
Chapter 4: Chemical Reactions and Chemical Quantities
Chapter 5: Introduction to Solutions and Aqueous Reactions
Chapter 7: Thermochemistry
Chapter 8: The Quantum-Mechanical Model of the Atom
Chapter 9: Periodic Properties of the Elements
Chapter 10: Chemical Bonding I: The Lewis Model
Chapter 11: Chemical Bonding II: Molecular Shapes, VB and MO Theory
Chapter 6: Gases
Chapter 12: Liquids, Solids, and Intermolecular Forces
Chapter 13 (part): Solids (13.2 – 13.6)
Jane Bush college chem
Week
Competencies
8/15
Write the name and symbol for selected elements.
Lab procedures
Density
Quiz on Element symbols
8/22
Solve problems involving density.
Precision/Accuracy
Measurements
Density Quiz
8/29
State the basic units of measurement for length, mass, volume and temperature in the SI system.
Give the numerical equivalent of selected SI prefixes.
Express numerical answers to the correct number of significant figures.
Convert temperatures between Fahrenheit, Celsius and Kelvin scales.
Measurement quiz
9/5
Distinguish among elements, compounds and mixtures.
Distinguish between physical and chemical properties and physical and chemical changes.
9/12
Recall the basic ideas in Dalton's atomic theory.
Summarize the experiments of J.J. Thomson, Robert Millikan and Ernest Rutherford that characterized the structure of the atom.
E.
Describe atoms in terms of electrons, protons and neutrons.
Given the isotopic masses and fractional abundances for a naturally occurring element, calculate its atomic weight
Identify the following areas of the periodic table: metals, nonmetals and metalloids; main groups, transition metals, inner transition metals; alkali metals, alkaline earth metals, halogens and noble gases.
Quiz atomic structure
9/19
Solve problems using dimensional analysis, including conversion of units.
Calculate the molecular mass of a compound from its formula.
Solve problems relating the mass of a compound to the number of moles of a compound.
Solve problems relating the mass of a compound to the number of molecules.
9/26
Calculate the percent composition of a compound from its formula.
Determine the empirical formula of a compound from its percent composition.
Determine the molecular formula of a compound from its empirical formula and molecular mass.
10/3
Test unit 1
10/10
Compare and contrast molecular compounds and ionic compounds.
10/17
Write the name and symbol for selected polyatomic ions.
Write names and formulas for the following types of compounds: acids.
Write names and formulas for the following types of compounds binary molecular compounds
Write names and formulas for the following types of compounds: ionic compounds
Polyatomic Ion quiz
10/24
Write the name and symbol for selected polyatomic ions.
Write names and formulas for the following types of compounds: acids.
Write names and formulas for the following types of compounds binary molecular compounds
Write names and formulas for the following types of compounds: ionic compounds
Naming quiz
10/31
Write a balanced chemical equation given the reactants and products.
Solve problems relating grams and moles of substances in balanced chemical equations.
11/7
Recognize the limiting reagent in a reaction and do calculations with limiting reagent.
Calculate theoretical yield and percent yield when actual yield is given
Predict the product of the combustion reactions of hydrocarbons and simple compounds having C, H and O.
Identify chemical reactions by type: combination, decomposition, combustion.
Stoich quiz
11/14
Explain how to make solutions of given concentration.
Distinguish among strong, weak and nonelectrolytes in solution.
List the common acids and bases and classify each as a strong or weak electrolyte.
Explain how to dilute solutions to a specified volume or concentration.
11/28
Write balanced complete and net ionic equations.
Electrolytes and Molarity quiz
Solve solution stoichiometry problems.
12/5
Write balanced complete and net ionic equations.
Solve solution stoichiometry problems
12/12
.
Assign oxidation numbers to atoms in molecules and ions.
Recognize oxidation-reduction reactions and identify oxidizing and reducing agents.
Balance simple oxidation-reduction reactions by the half-reaction method.
12/19
Test 2
1 /2
Solve solution stoichiometry
1/9
Titrations
1/16
Convert between torr, mm Hg, standard atmosphere and Pascal.
Demonstrate an understanding of the gas laws (Charles', Boyle's ideal, etc.) by working problems with them.
List the points of the kinetic molecular theory and describe how this theory explains the common gas laws.
1/23
Convert between torr, mm Hg, standard atmosphere and Pascal.
Demonstrate an understanding of the gas laws (Charles', Boyle's ideal, etc.) by working problems with them.
List the points of the kinetic molecular theory and describe how this theory explains the common gas laws.
Combined gas law, PV=nrt
A. Work stoichiometry problems involving gases and the gas laws.
B. Describe how the relative rates of diffusion and effusion of two gases depend on their molar masses (Graham's law).
C. Describe how a real gas differs from an ideal gas.
Stoichiometry of gases
1/30
specific heat qmcat lab and calc
Recognize and illustrate the law of conservation of energy.
Distinguish between a system and its surroundings and describe the energy changes in a system and its surroundings during a given reaction.
Solve calorimetry and heat capacity problems. breaking and forming bonds
enthalpy of reaction
State the first law of thermodynamics
Solve problems involving enthalpies for physical and chemical changes.
2/6
. Determine the enthalpy of reaction using bond energies.
Calculate standard enthalpies of reaction from standard enthalpies of formation.
2/13
hess
Calculate enthalpy changes using Hess' law and measured enthalpies of reaction.
Enthalpy stoichiometry
2/20
Solve problems relating frequency, wavelength and energy of electromagnetic radiation.
Explain the essential feature of Planck's quantum theory.
Discuss how line spectra give evidence of energy quantization.
econfig
Describe the wave mechanical model of the atom.
Describe s and p orbitals and recognize d orbitals.
Write a set of quantum numbers for any particular electron.
Write the electron configuration of elements up to atomic number 57.
2/27
Write electron configurations for ions of main group and transition elements.
Relate position on the periodic table to electron configuration and quantum numbers.
per trends
Apply Periodic Trends in atomic radii to predict relative size of an atom.
Predict relative first ionization energies from periodic trends.
Relate position on the periodic table to electron configuration and quantum numbers.
Write a set of quantum numbers for any particular electron.
3/6
Test 3
3/20
Explain the observed changes in value of the successive ionization energies for a given atom. Describe the general differences in chemical reactivity between metals and nonmetals.
Describe the periodic trends in metallic and nonmetallic behavior.
Predict the relative size of anions and cations formed from an atom.
3/27
Determine the number of valence electrons for an atom and write its Lewis symbol.
Draw Lewis structures for atoms, ions and covalent compounds, recognizing when multiple bonds, resonance structures, expanded valence shells, incomplete valence shells and odd electrons are needed.
Relate the number of electron pairs in the valence shell of an atom in a molecule to the geometrical arrangement around that atom.
Describe the bonding in a double and triple bond.
Predict molecular geometry using the VSEPR model.
Describe sp, sp2 and sp3 hybrid orbitals.
4/3
Draw a phase diagram of a substance given proper data and use a phase diagram to predict the phases present at a given temperature and pressure.
Given heating/cooling curves, calculate the heat associated when a given substance changes from one condition to another.
Compare and contrast gases, liquids and solids.
Employ the kinetic molecular model to explain the differences between the gas, liquid and solid states
4/10
Qualitatively explain the relationship between intermolecular forces and properties of liquids and solids.
Predict whether a molecule can have a net dipole moment from the molecular shape and the electronegativity of the atoms involved.
Describe covalent bonding using valence bond theory
Recognize where dipole-dipole forces, hydrogen bonding and London dispersion forces are important.
4/17
Compare and contrast gases, liquids and solids.
Employ the kinetic molecular model to explain the differences between the gas, liquid and solid states.
Recognize where dipole-dipole forces, hydrogen bonding and London dispersion forces are important.
Qualitatively explain the relationship between intermolecular forces and properties of liquids and solids.
Draw a phase diagram of a substance given proper data and use a phase diagram to predict the phases present at a given temperature and pressure.
Given heating/cooling curves, calculate the heat associated when a given substance changes from one condition to another.
4/24
Test 4
5/1
Compare and contrast crystalline and amorphous solids.
Categorize crystalline solids as ionic, molecular, covalent network and metallic solids.
5/8
Final
► See the following pages for specific learning objectives and important terms for each chapter.
CHEM 124 Objectives for Tro 5th Edition
CHAPTER 1 OBJECTIVES
1. Review the scientific method.
2. Be able to differentiate among the three states of matter, and distinguish between element, compound, and mixtures.
3. Distinguish between physical and chemical properties, and physical and chemical changes.
4. State the basic units of measurement for length, mass, volume, and temperature in the SI system.
5. Convert temperatures between Celsius and Kelvin scales.
6. Give the numerical meaning of the selected SI prefixes in Table 1.2: pico, nano, micro, milli, centi, deci, kilo, mega, giga
7. Work problems involving density.
8. Express numerical answers to the correct number of significant figures.
9. Use dimensional analysis and conversion factors to solve problems, including conversion of units.
CHAPTER 1 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
atom
molecule
macroscopic level
atomic and molecular level
hypothesis
experiment
law
theory
matter
solid
liquid
gas
pure substance
element
compound
heterogeneous mixture
homogeneous mixture (solution)
physical change
chemical change (chemical reaction)
chemical property
physical property
International System (SI) of Units (base units)
mass
Kelvin (K) scale
Celsius scale
liter (L)
density
intensive property
extensive property
significant figures
defined quantity (exact numbers)
accuracy
precision
dimensional analysis
conversion factor
scientific notation
CHAPTER 2 OBJECTIVES
1. Recall the basic postulates in Dalton’s atomic theory.
2. Summarize the experiments of J.J. Thomson, Robert Millikan, and Ernest Rutherford that characterized the structure of the atom.
3. Describe atoms in terms of electrons, protons, and neutrons.
4. Identify the following areas of the periodic table: metals, nonmetals and metalloids; main group elements, alkali metals, alkaline earth metals, halogens, noble gases, transition elements, lanthanides, and actinides.
5. Write the name and symbol for any of the 48 selected elements provided by your instructor.
6. Describe the formation of cations and anions and relate the periodic table position of elements to the type and charge of ions usually formed.
7. Given the isotopic masses (in amu) and fractional abundances for a naturally occurring element, calculate its atomic weight.
8.. Given the mass of an element, calculate the number of moles and atoms. Given the number of moles of an element, calculate the mass and atoms.
CHAPTER 2 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
law of conservation of mass
law of definite proportions (aka law of constant composition)
electron
radioactivity
nucleus
proton
neutron
atomic mass unit (amu)
atomic number (Z)
isotope
isotope abundance
percent abundance
fractional abundance
mass number (A)
ion
monatomic ion (instructor will define)
cation
anion
metal
nonmetal
metalloid
periodic table
groups (families)
periods
atomic weight (atomic mass)
mass spectrometer
mole (mol)
Avogadro’s number (NA)
molar mass (of elements)
CHAPTER 3 OBJECTIVES
1.. Compare and contrast molecular compounds and ionic compounds.
2. Write the name and formula for the polyatomic ions included in the list provided by your instructor.
3.. From formulas write names and from names write formulas for the following types of compounds: ionic compounds and binary molecular compounds.
4.. Write names from formulas and formulas from names for acids and bases
5.. Given a formula, calculate the molar mass and convert between mass, number of moles and number of atoms, molecules or formula units.
6. Find an empirical formula from percent composition. When molar mass is available, also find molecular formula. Calculate percent composition from formula.
CHAPTER 3 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
ionic bond
ionic compound
covalent bond
molecular compound
chemical formula
empirical formula
molecular formula
structural formula
formula unit
polyatomic ion
binary compound
hydrated compounds (hydrates)
acids
molar mass (of compounds)
mass percent composition (mass percent)
CHAPTER 4 OBJECTIVES
1. Given the formulas of reactants and products, balance a chemical equation.
2. Predict the products of the combustion reactions of hydrocarbons and simple compounds having C, H, and O.
3. Do calculations with grams and moles of substances in balanced equations.
4. Recognize the limiting reactant in a reaction and do calculations with limiting reactant.
5. Calculate theoretical yield and calculate percent yield when actual yield is given.
.... CHAPTER 4 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
chemical equation
reactant
product
stoichiometry
stoichiometric coefficient
limiting reactant
theoretical yield
actual yield
percent yield
CHAPTER 5 OBJECTIVES
1.. Explain how to make solutions of given concentration.
2. Explain how to dilute solutions to a specified volume or concentration.
3. Work solution stoichiometry problems.
4. Distinguish between strong, weak, and nonelectrolytes in solution. Classify individual acids and bases as a strong or weak electrolyte.
5. Use a given set of solubility rules to predict precipitations.
6. Write and balance complete and net ionic equations.
7. Identify the acid-base species in the reaction using the Arrhenius and Bronsted-Lowry methods. (Bronsted-Lowry found in Ch 17)
8. Assign oxidation numbers to atoms in molecules or ions and balance simple oxidation-reduction reactions by charge and mass.
9. Be able to distinguish between precipitation, acid-base, gas-evolving and redox reactions.
CHAPTER 5 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
solution
solvent
solute
molarity
dilution
electrolyte
strong electrolyte
nonelectrolyte
strong acid
weak acid
weak electrolyte
strong base (chapter 17)
weak base (chapter 17)
precipitate reactions (exchange reactions)
precipitate
molecular equation
complete ionic equation
spectator ion
net ionic equation
Acid-Base reaction (neutralization reaction)
Arrhenius acid
Arrhenius base
salt
Bronsted-Lowry acid and base (Ch 17)
Conjugate acid and base (Ch 17)
titration
equivalence point
acid-base indicator
gas-evolution reaction
oxidation-reduction reaction (redox reaction)
oxidation (oxidized)
reduction (reduced)
oxidation state (oxidation number)
oxidizing agent
reducing agent
half-reaction
CHAPTER 7 OBJECTIVES
1. Recognize and illustrate the law of conservation of energy.
2. Distinguish between a system and its surroundings and describe the energy changes in a system and its surroundings during a given reaction.
3. Work problems involving enthalpy for chemical changes.
4. Work specific heat capacity and calorimetry problems.
5. Calculate enthalpy changes using Hess’s law and measured enthalpies of reaction.
6. Use standard molar enthalpies of formation to calculate enthalpies of reaction.
CHAPTER 7 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
thermodynamics
energy
heat
kinetic energy
potential energy
law of conservation of energy (first law of thermodynamics)
system
surroundings
joule (J)
calorie (cal)
internal energy
state function
thermal equilibrium
specific heat capacity (specific heat)
enthalpy
exothermic process
endothermic process
calorimeter
Hess’s Law
standard state
standard molar enthalpy of formation (standard heat of formation)
CHAPTER 8 OBJECTIVES
1. Work problems relating frequency, wavelength, and energy of electromagnetic radiation.
2. Discuss how atomic line spectra give evidence of energy quantization.
3. Describe the wave mechanical model of the atom.
4. Write a set of quantum numbers for any particular electron.
5. Describe the s and p orbitals, and recognize d orbitals.
6. Describe the scientific contributions of Planck, Einstein, de Broglie, Bohr, Schrödinger, and Heisenberg.
CHAPTER 8 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
electromagnetic radiation (electromagnetic spectrum)
wavelength (λ)
frequency (ν)
photoelectric effect
Planck’s constant
photons
line emission spectrum
continuous spectrum
de Broglie equation
uncertainty principle
quantum mechanics (wave mechanics)
particle-wave duality
energy levels
principle quantum number (n)
angular momentum quantum number (l)
magnetic quantum number (ml)
spin quantum number (ms)
CHAPTER 9 OBJECTIVES
1. Describe the scientific contribution of Mendeleev and Pauli.
2. Write the electron configuration of elements that have the expected configuration.
3. Relate position on the periodic table to electron configuration and quantum numbers.
4. Write the electron configuration for ions that have the expected configuration.
5. Describe and explain Periodic Trends (atomic radii, first ionization energy, electron affinity and metallic/nonmetallic behavior).
6. Predict the relative sizes of anions and cations formed from an atom.
7. Explain the observed changes in value of the successive ionization energies for a given atom.
CHAPTER 9 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
electron configuration of neutral atoms
Pauli exclusion principle
orbital box diagram
aufbau (“building up”) principle
Hund’s rule
valence electrons
core electrons
noble gas notation
effective nuclear charge
atomic radius (atomic size)
electron configuration of ions
paramagnetic
diamagnetic
ion size (ionic radius)
first ionization energy
electron affinity
isoelectronic
CHAPTER 10 OBJECTIVES
1. Determine the number of valence electrons for an atom and write its Lewis symbol.
2... Use electronegativity differences between bonding atoms to classify bonds as polar covalent, nonpolar covalent, or ionic.
3. . Recognize when the octet rule applies to the arrangement of electrons in the valence shell.
4... Draw Lewis structures for atoms, ions, and covalent compounds. Recognize when multiple bonds, resonance structures, expanded valence shells, deficient electrons, and odd electrons are needed.
5... Be able to calculate the enthalpy of reaction using bond energies.
CHAPTER 10 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
chemical bond
ionic bond
covalent bond
Lewis electron dot symbol
octet rule
bond pair
lone pair (nonbonding electrons)
single bond
double bond
triple bond
bond polarity
polar covalent bond
electronegativity
dipole moment
Lewis electron dot structures (Lewis structures)
resonance (delocalized bonding)
resonance hybrid
formal charge
free radical
coordinate covalent bond
bond energy
bond length
CHAPTER 11 OBJECTIVES
1. Predict electron geometry and molecular geometry using the VSEPR model.
2. Predict whether a molecule can have a net dipole moment from the molecular shape and electronegativity of the atoms involved.
3. Use valence bond (VB) theory to describe the orbital used to form covalent bonds.
4. Be able to describe sp, sp2, and sp3 hybrid orbitals.
5. Describe the bonding in a molecule or ion that contains double and/or triple bonds.
6. Describe MO theory for bonding and antibonding orbitals.
7. Determine bond orders and relate them to relative bond strength.
8. Relate MO theory to concepts such as the structural, energetic, spectroscopic and magnetic properties of molecules.
CHAPTER 11 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
valence shell electron-pair repulsion (VSEPR) model
electron geometry
linear
trigonal planar
tetrahedral
trigonal-bipyramidal
octahedral
bond angle
molecular geometry
bent
trigonal pyramidal
simple perspective drawing
valence bond theory
hybrid orbitals
sp hybrid orbital
sp2 hybrid orbital
sp3 hybrid orbital
sp3d hybrid orbital (instructor’s option)
sp3d2 hybrid orbital (instructor’s option)
sigma bond (σ bond)
pi bond (π bond)
bonding/antibonding orbitals
bond order
CHAPTER 6 OBJECTIVES
1. Describe the general characteristics of gases compared to solids and liquids.
2. Convert between mm Hg, torr, standard atmosphere, and Pascal.
3. Demonstrate an understanding of the gas laws (Boyle’s, Charles’, Avogadro’s, Ideal, etc.) by working problems with them.
4. Work stoichiometry problems involving gases and the gas laws.
5. List the points of the kinetic molecular theory and describe how this theory explains the common gas laws.
6. Describe how the relative rates of diffusion and effusion of two gases depends on their molar masses (Graham’s Law).
7. Describe how a real gas differs from an ideal gas.
CHAPTER 6 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
pressure
millimeters of mercury (mm Hg)
pascal (Pa)
atmosphere (atm)
barometer
combined gas law (instructor’s option)
ideal gas constant (R)
standard molar volume
standard temperature and pressure (STP)
partial pressure
Dalton’s law of partial pressures
mole fraction
kinetic-molecular theory of gases (kinetic theory)
root-mean-square (rms) speed
diffusion
effusion
Graham’s law of effusion
significance of van der Waals equation
CHAPTER 12 OBJECTIVES
1. Compare and contrast crystalline and amorphous solids.
2. Use the kinetic molecular model to explain the differences between the gas, liquid, and solid states.
3. Recognize where dipole-dipole forces, London dispersion forces, and hydrogen bonding are important.
3. Qualitatively explain the relationship between intermolecular forces and properties of liquids and solids.
4. Work problems involving enthalpy for physical changes.
5. Interpreting heating/cooling curves.
6. Draw a phase diagram of a substance given proper data; also use a phase diagram to predict which phases are present at a given temperature and pressure.
CHAPTER 12 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
crystalline solid
amorphous solid
change of state
intermolecular forces
London dispersion forces (induced dipole - induced dipole)
dipole-dipole forces
hydrogen bonding
ion-dipole forces
surface tension
viscosity
capillary action
adhesive forces
cohesive forces
vaporization (evaporation)
condensation
vapor pressure
boiling
boiling point
heat of vaporization
sublimation
melting
melting point
heat of fusion
phase diagram
triple point
CHAPTER 13 (13.4 – 13.6)
1. Categorize solids as ionic, metallic, molecular, network covalent, or amorphous.
CHAPTER 13 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
molecular solid
ionic solid
metallic solid
network covalent solid
The following list of chapters (from Tro, Chemistry: A Molecular Approach) is only intended to provide a rough outline of the course content. Your instructor will provide you with a more specific lecture schedule including exam dates.
Chapter 1: Matter, Measurement, and Problem Solving
Chapter 2: Atoms and Elements
Chapter 3: Molecules and Compounds
Chapter 4: Chemical Reactions and Chemical Quantities
Chapter 5: Introduction to Solutions and Aqueous Reactions
Chapter 7: Thermochemistry
Chapter 8: The Quantum-Mechanical Model of the Atom
Chapter 9: Periodic Properties of the Elements
Chapter 10: Chemical Bonding I: The Lewis Model
Chapter 11: Chemical Bonding II: Molecular Shapes, VB and MO Theory
Chapter 6: Gases
Chapter 12: Liquids, Solids, and Intermolecular Forces
Chapter 13 (part): Solids (13.2 – 13.6)
Jane Bush college chem
Week
Competencies
8/15
Write the name and symbol for selected elements.
Lab procedures
Density
Quiz on Element symbols
8/22
Solve problems involving density.
Precision/Accuracy
Measurements
Density Quiz
8/29
State the basic units of measurement for length, mass, volume and temperature in the SI system.
Give the numerical equivalent of selected SI prefixes.
Express numerical answers to the correct number of significant figures.
Convert temperatures between Fahrenheit, Celsius and Kelvin scales.
Measurement quiz
9/5
Distinguish among elements, compounds and mixtures.
Distinguish between physical and chemical properties and physical and chemical changes.
9/12
Recall the basic ideas in Dalton's atomic theory.
Summarize the experiments of J.J. Thomson, Robert Millikan and Ernest Rutherford that characterized the structure of the atom.
E.
Describe atoms in terms of electrons, protons and neutrons.
Given the isotopic masses and fractional abundances for a naturally occurring element, calculate its atomic weight
Identify the following areas of the periodic table: metals, nonmetals and metalloids; main groups, transition metals, inner transition metals; alkali metals, alkaline earth metals, halogens and noble gases.
Quiz atomic structure
9/19
Solve problems using dimensional analysis, including conversion of units.
Calculate the molecular mass of a compound from its formula.
Solve problems relating the mass of a compound to the number of moles of a compound.
Solve problems relating the mass of a compound to the number of molecules.
9/26
Calculate the percent composition of a compound from its formula.
Determine the empirical formula of a compound from its percent composition.
Determine the molecular formula of a compound from its empirical formula and molecular mass.
10/3
Test unit 1
10/10
Compare and contrast molecular compounds and ionic compounds.
10/17
Write the name and symbol for selected polyatomic ions.
Write names and formulas for the following types of compounds: acids.
Write names and formulas for the following types of compounds binary molecular compounds
Write names and formulas for the following types of compounds: ionic compounds
Polyatomic Ion quiz
10/24
Write the name and symbol for selected polyatomic ions.
Write names and formulas for the following types of compounds: acids.
Write names and formulas for the following types of compounds binary molecular compounds
Write names and formulas for the following types of compounds: ionic compounds
Naming quiz
10/31
Write a balanced chemical equation given the reactants and products.
Solve problems relating grams and moles of substances in balanced chemical equations.
11/7
Recognize the limiting reagent in a reaction and do calculations with limiting reagent.
Calculate theoretical yield and percent yield when actual yield is given
Predict the product of the combustion reactions of hydrocarbons and simple compounds having C, H and O.
Identify chemical reactions by type: combination, decomposition, combustion.
Stoich quiz
11/14
Explain how to make solutions of given concentration.
Distinguish among strong, weak and nonelectrolytes in solution.
List the common acids and bases and classify each as a strong or weak electrolyte.
Explain how to dilute solutions to a specified volume or concentration.
11/28
Write balanced complete and net ionic equations.
Electrolytes and Molarity quiz
Solve solution stoichiometry problems.
12/5
Write balanced complete and net ionic equations.
Solve solution stoichiometry problems
12/12
.
Assign oxidation numbers to atoms in molecules and ions.
Recognize oxidation-reduction reactions and identify oxidizing and reducing agents.
Balance simple oxidation-reduction reactions by the half-reaction method.
12/19
Test 2
1 /2
Solve solution stoichiometry
1/9
Titrations
1/16
Convert between torr, mm Hg, standard atmosphere and Pascal.
Demonstrate an understanding of the gas laws (Charles', Boyle's ideal, etc.) by working problems with them.
List the points of the kinetic molecular theory and describe how this theory explains the common gas laws.
1/23
Convert between torr, mm Hg, standard atmosphere and Pascal.
Demonstrate an understanding of the gas laws (Charles', Boyle's ideal, etc.) by working problems with them.
List the points of the kinetic molecular theory and describe how this theory explains the common gas laws.
Combined gas law, PV=nrt
A. Work stoichiometry problems involving gases and the gas laws.
B. Describe how the relative rates of diffusion and effusion of two gases depend on their molar masses (Graham's law).
C. Describe how a real gas differs from an ideal gas.
Stoichiometry of gases
1/30
specific heat qmcat lab and calc
Recognize and illustrate the law of conservation of energy.
Distinguish between a system and its surroundings and describe the energy changes in a system and its surroundings during a given reaction.
Solve calorimetry and heat capacity problems. breaking and forming bonds
enthalpy of reaction
State the first law of thermodynamics
Solve problems involving enthalpies for physical and chemical changes.
2/6
. Determine the enthalpy of reaction using bond energies.
Calculate standard enthalpies of reaction from standard enthalpies of formation.
2/13
hess
Calculate enthalpy changes using Hess' law and measured enthalpies of reaction.
Enthalpy stoichiometry
2/20
Solve problems relating frequency, wavelength and energy of electromagnetic radiation.
Explain the essential feature of Planck's quantum theory.
Discuss how line spectra give evidence of energy quantization.
econfig
Describe the wave mechanical model of the atom.
Describe s and p orbitals and recognize d orbitals.
Write a set of quantum numbers for any particular electron.
Write the electron configuration of elements up to atomic number 57.
2/27
Write electron configurations for ions of main group and transition elements.
Relate position on the periodic table to electron configuration and quantum numbers.
per trends
Apply Periodic Trends in atomic radii to predict relative size of an atom.
Predict relative first ionization energies from periodic trends.
Relate position on the periodic table to electron configuration and quantum numbers.
Write a set of quantum numbers for any particular electron.
3/6
Test 3
3/20
Explain the observed changes in value of the successive ionization energies for a given atom. Describe the general differences in chemical reactivity between metals and nonmetals.
Describe the periodic trends in metallic and nonmetallic behavior.
Predict the relative size of anions and cations formed from an atom.
3/27
Determine the number of valence electrons for an atom and write its Lewis symbol.
Draw Lewis structures for atoms, ions and covalent compounds, recognizing when multiple bonds, resonance structures, expanded valence shells, incomplete valence shells and odd electrons are needed.
Relate the number of electron pairs in the valence shell of an atom in a molecule to the geometrical arrangement around that atom.
Describe the bonding in a double and triple bond.
Predict molecular geometry using the VSEPR model.
Describe sp, sp2 and sp3 hybrid orbitals.
4/3
Draw a phase diagram of a substance given proper data and use a phase diagram to predict the phases present at a given temperature and pressure.
Given heating/cooling curves, calculate the heat associated when a given substance changes from one condition to another.
Compare and contrast gases, liquids and solids.
Employ the kinetic molecular model to explain the differences between the gas, liquid and solid states
4/10
Qualitatively explain the relationship between intermolecular forces and properties of liquids and solids.
Predict whether a molecule can have a net dipole moment from the molecular shape and the electronegativity of the atoms involved.
Describe covalent bonding using valence bond theory
Recognize where dipole-dipole forces, hydrogen bonding and London dispersion forces are important.
4/17
Compare and contrast gases, liquids and solids.
Employ the kinetic molecular model to explain the differences between the gas, liquid and solid states.
Recognize where dipole-dipole forces, hydrogen bonding and London dispersion forces are important.
Qualitatively explain the relationship between intermolecular forces and properties of liquids and solids.
Draw a phase diagram of a substance given proper data and use a phase diagram to predict the phases present at a given temperature and pressure.
Given heating/cooling curves, calculate the heat associated when a given substance changes from one condition to another.
4/24
Test 4
5/1
Compare and contrast crystalline and amorphous solids.
Categorize crystalline solids as ionic, molecular, covalent network and metallic solids.
5/8
Final
► See the following pages for specific learning objectives and important terms for each chapter.
CHEM 124 Objectives for Tro 5th Edition
CHAPTER 1 OBJECTIVES
1. Review the scientific method.
2. Be able to differentiate among the three states of matter, and distinguish between element, compound, and mixtures.
3. Distinguish between physical and chemical properties, and physical and chemical changes.
4. State the basic units of measurement for length, mass, volume, and temperature in the SI system.
5. Convert temperatures between Celsius and Kelvin scales.
6. Give the numerical meaning of the selected SI prefixes in Table 1.2: pico, nano, micro, milli, centi, deci, kilo, mega, giga
7. Work problems involving density.
8. Express numerical answers to the correct number of significant figures.
9. Use dimensional analysis and conversion factors to solve problems, including conversion of units.
CHAPTER 1 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
atom
molecule
macroscopic level
atomic and molecular level
hypothesis
experiment
law
theory
matter
solid
liquid
gas
pure substance
element
compound
heterogeneous mixture
homogeneous mixture (solution)
physical change
chemical change (chemical reaction)
chemical property
physical property
International System (SI) of Units (base units)
mass
Kelvin (K) scale
Celsius scale
liter (L)
density
intensive property
extensive property
significant figures
defined quantity (exact numbers)
accuracy
precision
dimensional analysis
conversion factor
scientific notation
CHAPTER 2 OBJECTIVES
1. Recall the basic postulates in Dalton’s atomic theory.
2. Summarize the experiments of J.J. Thomson, Robert Millikan, and Ernest Rutherford that characterized the structure of the atom.
3. Describe atoms in terms of electrons, protons, and neutrons.
4. Identify the following areas of the periodic table: metals, nonmetals and metalloids; main group elements, alkali metals, alkaline earth metals, halogens, noble gases, transition elements, lanthanides, and actinides.
5. Write the name and symbol for any of the 48 selected elements provided by your instructor.
6. Describe the formation of cations and anions and relate the periodic table position of elements to the type and charge of ions usually formed.
7. Given the isotopic masses (in amu) and fractional abundances for a naturally occurring element, calculate its atomic weight.
8.. Given the mass of an element, calculate the number of moles and atoms. Given the number of moles of an element, calculate the mass and atoms.
CHAPTER 2 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
law of conservation of mass
law of definite proportions (aka law of constant composition)
electron
radioactivity
nucleus
proton
neutron
atomic mass unit (amu)
atomic number (Z)
isotope
isotope abundance
percent abundance
fractional abundance
mass number (A)
ion
monatomic ion (instructor will define)
cation
anion
metal
nonmetal
metalloid
periodic table
groups (families)
periods
atomic weight (atomic mass)
mass spectrometer
mole (mol)
Avogadro’s number (NA)
molar mass (of elements)
CHAPTER 3 OBJECTIVES
1.. Compare and contrast molecular compounds and ionic compounds.
2. Write the name and formula for the polyatomic ions included in the list provided by your instructor.
3.. From formulas write names and from names write formulas for the following types of compounds: ionic compounds and binary molecular compounds.
4.. Write names from formulas and formulas from names for acids and bases
5.. Given a formula, calculate the molar mass and convert between mass, number of moles and number of atoms, molecules or formula units.
6. Find an empirical formula from percent composition. When molar mass is available, also find molecular formula. Calculate percent composition from formula.
CHAPTER 3 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
ionic bond
ionic compound
covalent bond
molecular compound
chemical formula
empirical formula
molecular formula
structural formula
formula unit
polyatomic ion
binary compound
hydrated compounds (hydrates)
acids
molar mass (of compounds)
mass percent composition (mass percent)
CHAPTER 4 OBJECTIVES
1. Given the formulas of reactants and products, balance a chemical equation.
2. Predict the products of the combustion reactions of hydrocarbons and simple compounds having C, H, and O.
3. Do calculations with grams and moles of substances in balanced equations.
4. Recognize the limiting reactant in a reaction and do calculations with limiting reactant.
5. Calculate theoretical yield and calculate percent yield when actual yield is given.
.... CHAPTER 4 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
chemical equation
reactant
product
stoichiometry
stoichiometric coefficient
limiting reactant
theoretical yield
actual yield
percent yield
CHAPTER 5 OBJECTIVES
1.. Explain how to make solutions of given concentration.
2. Explain how to dilute solutions to a specified volume or concentration.
3. Work solution stoichiometry problems.
4. Distinguish between strong, weak, and nonelectrolytes in solution. Classify individual acids and bases as a strong or weak electrolyte.
5. Use a given set of solubility rules to predict precipitations.
6. Write and balance complete and net ionic equations.
7. Identify the acid-base species in the reaction using the Arrhenius and Bronsted-Lowry methods. (Bronsted-Lowry found in Ch 17)
8. Assign oxidation numbers to atoms in molecules or ions and balance simple oxidation-reduction reactions by charge and mass.
9. Be able to distinguish between precipitation, acid-base, gas-evolving and redox reactions.
CHAPTER 5 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
solution
solvent
solute
molarity
dilution
electrolyte
strong electrolyte
nonelectrolyte
strong acid
weak acid
weak electrolyte
strong base (chapter 17)
weak base (chapter 17)
precipitate reactions (exchange reactions)
precipitate
molecular equation
complete ionic equation
spectator ion
net ionic equation
Acid-Base reaction (neutralization reaction)
Arrhenius acid
Arrhenius base
salt
Bronsted-Lowry acid and base (Ch 17)
Conjugate acid and base (Ch 17)
titration
equivalence point
acid-base indicator
gas-evolution reaction
oxidation-reduction reaction (redox reaction)
oxidation (oxidized)
reduction (reduced)
oxidation state (oxidation number)
oxidizing agent
reducing agent
half-reaction
CHAPTER 7 OBJECTIVES
1. Recognize and illustrate the law of conservation of energy.
2. Distinguish between a system and its surroundings and describe the energy changes in a system and its surroundings during a given reaction.
3. Work problems involving enthalpy for chemical changes.
4. Work specific heat capacity and calorimetry problems.
5. Calculate enthalpy changes using Hess’s law and measured enthalpies of reaction.
6. Use standard molar enthalpies of formation to calculate enthalpies of reaction.
CHAPTER 7 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
thermodynamics
energy
heat
kinetic energy
potential energy
law of conservation of energy (first law of thermodynamics)
system
surroundings
joule (J)
calorie (cal)
internal energy
state function
thermal equilibrium
specific heat capacity (specific heat)
enthalpy
exothermic process
endothermic process
calorimeter
Hess’s Law
standard state
standard molar enthalpy of formation (standard heat of formation)
CHAPTER 8 OBJECTIVES
1. Work problems relating frequency, wavelength, and energy of electromagnetic radiation.
2. Discuss how atomic line spectra give evidence of energy quantization.
3. Describe the wave mechanical model of the atom.
4. Write a set of quantum numbers for any particular electron.
5. Describe the s and p orbitals, and recognize d orbitals.
6. Describe the scientific contributions of Planck, Einstein, de Broglie, Bohr, Schrödinger, and Heisenberg.
CHAPTER 8 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
electromagnetic radiation (electromagnetic spectrum)
wavelength (λ)
frequency (ν)
photoelectric effect
Planck’s constant
photons
line emission spectrum
continuous spectrum
de Broglie equation
uncertainty principle
quantum mechanics (wave mechanics)
particle-wave duality
energy levels
principle quantum number (n)
angular momentum quantum number (l)
magnetic quantum number (ml)
spin quantum number (ms)
CHAPTER 9 OBJECTIVES
1. Describe the scientific contribution of Mendeleev and Pauli.
2. Write the electron configuration of elements that have the expected configuration.
3. Relate position on the periodic table to electron configuration and quantum numbers.
4. Write the electron configuration for ions that have the expected configuration.
5. Describe and explain Periodic Trends (atomic radii, first ionization energy, electron affinity and metallic/nonmetallic behavior).
6. Predict the relative sizes of anions and cations formed from an atom.
7. Explain the observed changes in value of the successive ionization energies for a given atom.
CHAPTER 9 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
electron configuration of neutral atoms
Pauli exclusion principle
orbital box diagram
aufbau (“building up”) principle
Hund’s rule
valence electrons
core electrons
noble gas notation
effective nuclear charge
atomic radius (atomic size)
electron configuration of ions
paramagnetic
diamagnetic
ion size (ionic radius)
first ionization energy
electron affinity
isoelectronic
CHAPTER 10 OBJECTIVES
1. Determine the number of valence electrons for an atom and write its Lewis symbol.
2... Use electronegativity differences between bonding atoms to classify bonds as polar covalent, nonpolar covalent, or ionic.
3. . Recognize when the octet rule applies to the arrangement of electrons in the valence shell.
4... Draw Lewis structures for atoms, ions, and covalent compounds. Recognize when multiple bonds, resonance structures, expanded valence shells, deficient electrons, and odd electrons are needed.
5... Be able to calculate the enthalpy of reaction using bond energies.
CHAPTER 10 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
chemical bond
ionic bond
covalent bond
Lewis electron dot symbol
octet rule
bond pair
lone pair (nonbonding electrons)
single bond
double bond
triple bond
bond polarity
polar covalent bond
electronegativity
dipole moment
Lewis electron dot structures (Lewis structures)
resonance (delocalized bonding)
resonance hybrid
formal charge
free radical
coordinate covalent bond
bond energy
bond length
CHAPTER 11 OBJECTIVES
1. Predict electron geometry and molecular geometry using the VSEPR model.
2. Predict whether a molecule can have a net dipole moment from the molecular shape and electronegativity of the atoms involved.
3. Use valence bond (VB) theory to describe the orbital used to form covalent bonds.
4. Be able to describe sp, sp2, and sp3 hybrid orbitals.
5. Describe the bonding in a molecule or ion that contains double and/or triple bonds.
6. Describe MO theory for bonding and antibonding orbitals.
7. Determine bond orders and relate them to relative bond strength.
8. Relate MO theory to concepts such as the structural, energetic, spectroscopic and magnetic properties of molecules.
CHAPTER 11 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
valence shell electron-pair repulsion (VSEPR) model
electron geometry
linear
trigonal planar
tetrahedral
trigonal-bipyramidal
octahedral
bond angle
molecular geometry
bent
trigonal pyramidal
simple perspective drawing
valence bond theory
hybrid orbitals
sp hybrid orbital
sp2 hybrid orbital
sp3 hybrid orbital
sp3d hybrid orbital (instructor’s option)
sp3d2 hybrid orbital (instructor’s option)
sigma bond (σ bond)
pi bond (π bond)
bonding/antibonding orbitals
bond order
CHAPTER 6 OBJECTIVES
1. Describe the general characteristics of gases compared to solids and liquids.
2. Convert between mm Hg, torr, standard atmosphere, and Pascal.
3. Demonstrate an understanding of the gas laws (Boyle’s, Charles’, Avogadro’s, Ideal, etc.) by working problems with them.
4. Work stoichiometry problems involving gases and the gas laws.
5. List the points of the kinetic molecular theory and describe how this theory explains the common gas laws.
6. Describe how the relative rates of diffusion and effusion of two gases depends on their molar masses (Graham’s Law).
7. Describe how a real gas differs from an ideal gas.
CHAPTER 6 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
pressure
millimeters of mercury (mm Hg)
pascal (Pa)
atmosphere (atm)
barometer
combined gas law (instructor’s option)
ideal gas constant (R)
standard molar volume
standard temperature and pressure (STP)
partial pressure
Dalton’s law of partial pressures
mole fraction
kinetic-molecular theory of gases (kinetic theory)
root-mean-square (rms) speed
diffusion
effusion
Graham’s law of effusion
significance of van der Waals equation
CHAPTER 12 OBJECTIVES
1. Compare and contrast crystalline and amorphous solids.
2. Use the kinetic molecular model to explain the differences between the gas, liquid, and solid states.
3. Recognize where dipole-dipole forces, London dispersion forces, and hydrogen bonding are important.
3. Qualitatively explain the relationship between intermolecular forces and properties of liquids and solids.
4. Work problems involving enthalpy for physical changes.
5. Interpreting heating/cooling curves.
6. Draw a phase diagram of a substance given proper data; also use a phase diagram to predict which phases are present at a given temperature and pressure.
CHAPTER 12 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
crystalline solid
amorphous solid
change of state
intermolecular forces
London dispersion forces (induced dipole - induced dipole)
dipole-dipole forces
hydrogen bonding
ion-dipole forces
surface tension
viscosity
capillary action
adhesive forces
cohesive forces
vaporization (evaporation)
condensation
vapor pressure
boiling
boiling point
heat of vaporization
sublimation
melting
melting point
heat of fusion
phase diagram
triple point
CHAPTER 13 (13.4 – 13.6)
1. Categorize solids as ionic, metallic, molecular, network covalent, or amorphous.
CHAPTER 13 TERMS
Define or explain the terms in the listing for this chapter and use these terms in examples.
molecular solid
ionic solid
metallic solid
network covalent solid
Week
Lab
8/16
Golf ball lab/Density
8/23
Precision /accuracy lab
8/30
Significant digits lab
9/6
9/13
Isotope lab
9/20
Mole lab practical
9/27
Percent comp lab
10/4
10/11
10/18
Ionic v covalent lab
10/25
11/1
11/8
Copper/silver lab
11/15
Electrolyte lab
11/29
Sports drink lab
12/6
Solubility lab
12/13
Mystery solutions lab
1/3
1/10
titrations
1/17
Gas law labs
1/24
Calculate gas law constant
1/31
Specific heat lab
2/7
Hess law lab
2/14
2/21
Flame tests
Spectroscopic analysiis
2/28
3/7
3/21
3/28
Molecular models
4/4
4/11
IMF lab
4/18
4/25
5/2
5/9
Lab
8/16
Golf ball lab/Density
8/23
Precision /accuracy lab
8/30
Significant digits lab
9/6
9/13
Isotope lab
9/20
Mole lab practical
9/27
Percent comp lab
10/4
10/11
10/18
Ionic v covalent lab
10/25
11/1
11/8
Copper/silver lab
11/15
Electrolyte lab
11/29
Sports drink lab
12/6
Solubility lab
12/13
Mystery solutions lab
1/3
1/10
titrations
1/17
Gas law labs
1/24
Calculate gas law constant
1/31
Specific heat lab
2/7
Hess law lab
2/14
2/21
Flame tests
Spectroscopic analysiis
2/28
3/7
3/21
3/28
Molecular models
4/4
4/11
IMF lab
4/18
4/25
5/2
5/9
