Jane Bush Chemistry Fun
  • Chemistry is Fun
    • About
    • Contact
    • schedule
  • Honors Chem
    • Resources
    • honors chem expectations
    • honors chem sequence
  • College Chem
    • JCCC information
    • Resources
    • expectations
    • course competencies
    • proposed schedule
  • Photos from Class
  • resources
    • Periodic table
  • Blog
​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
 
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
 

 
 
Powered by Create your own unique website with customizable templates.