Category Archives: Chemistry Lesson

Outline for a High School Chemistry Class

Need to complete a high school chemistry class with a wet lab? We did and that is how this blog was born. To help you get started, we have put together an overview of a high school chemistry course with links to posts about the lessons and labs.

I. Getting a High School chemistry lab kit.

We chose the The Home Scientist for our hands-on chemistry lab kit.

Of those offered, we used the CK01A Standard/Honors Home School Chemistry Laboratory Kit

The kit comes with most of the equipment and chemicals you need for a high school chemistry wet lab that you can do at home. It also comes with a free .pdf manual to download with complete, extensive lab instructions. (Link on this page). You will have to supply some materials, which are common household items for the most part. You will also need a place to store the chemicals away from small children and pets.

II. Choosing a High School Chemistry Textbook:

The text we used was Introductory Chemistry (4th Edition) by Nivaldo J. Tro. The book the textbook reading assignments in the lessons refer to that text. We chose this text over the Essentials version because it has three additional chapters at the end which could be used as reference material. We used chapters 1-16.

As typical with the textbook industry, there is now a newer edition.

Introductory Chemistry (5th Edition) by Nivaldo J. Tro

We liked the tone of the text and the great illustrations.

Another text that is commonly used in high school chemistry classrooms is:

Chemistry by Steven S. Zumdahl and Susan A. Zumdahl

Want more information? Don’t forget to check the links in the Online Chemistry Textbooks page and our Choosing a Chemistry Textbook post from the beginning of the class, which has more options.

III. High School Chemistry Class Outline

The following are links to the blog posts we used for the class. I split some of the chapters in the Tro text, so we ended up meeting for 22 weeks.

I have left out the sidebars, which were posts about some more information on questions the students found interesting.

Keeping a Chemistry Laboratory Notebook

Lesson 1: Introduction to Chemistry

Lab 1 Density of Liquids: Soft Drinks and Water

Lesson 2: Measurement and Problem Solving

Lab 2 Density of Solids and Measurement Challenge

Lesson 3: Matter

Lab 3: Separation of Mixtures

Lesson 4: Energy

Lab 4: Heat Capacity

Lesson 5: Atoms and Elements

Lab 5: Chemistry Unleashed

Lesson 6: Molecules and Compounds

Lab 6: What we did

Lesson 7: Calculating Chemical Composition

Lab 7: Finding Moles and Molecules

Lesson 8: Chemical Reactions Part 1

Lab 8: From Topic I, Recrystallization and Salting Out

Sidebar: Lab 8 Update

Lesson 9: Classifying Chemical Reactions

Lab 9: Topic III, Classifying Chemical Reactions

Lesson 10: Quantities in Chemical Reactions

Lab 10: Double Displacement Reactions

Lesson 11: The Electromagnetic Spectrum

Lab 11: Photochemistry

Lesson 12: Electrons, Atom Models, and the Periodic Table

(Note: For the lab, we did models of atom orbitals using Model Magic modeling clay. Let me know in the comments if you would like details.)

Lesson 13: Chemical Bonding

Lab 13: Conductance of Ionic and Molecular Solutes

Lesson 14: Gases

Lab 14: Gas Properties and Laws

Lesson 15: Properties of Liquids and Solids

Lab 15: Viscosity and Other Physical Properties of Liquids

Lesson 16: Solutions

Lab 16: Solubility and Solutions

Lesson 17: Acids and Bases

Lab 17: Investigating pH
Lesson 18: Rates of Chemical Reactions

Lab 18: Chemical Kinetics

Lesson 19: More About Rates of Chemical Reactions

Lab 19: Effect of Catalysts on Reactions

Lesson 20: Oxidation and Reduction

Lab 20: Sweet Redox Reactions for National Chemistry Week

Lab 21: Electrochemistry

For Lab 22, we had a review with activities and tasks from throughout the course.

outline-for-high-shcool-chemistry-course-homeschool

If you would like to know more about any of the materials or coursework, please feel free to leave questions in the comments.

Lesson 20: Oxidation and Reduction

Can you believe we are on the last chapter of the textbook already? Our chemistry lessons are flying by.

Textbook Reading: Chapter 16, Oxidation and Reduction on pp. 577- 603, paying particular attention to the Rules for Assigning Oxidation States on page 581.

Oxidation and reduction are fundamental to chemistry, but can be a bit complicated to understand. This is probably why the author introduces them in the final chapter after you have a good understanding of other concepts.

During many chemical reactions, electrons are moved around. When an atom loses electrons, chemists say it has been oxidized. When an atom gains electrons during a reaction, it is said to be reduced.

The first question you might have is why is gaining electrons called reduction. Actually, you have to go back 500 years or so to the foundations of chemistry through alchemy to find out the answer.

Hematite_macle (Image in the public domain from Wikimedia)

When early metal workers melted iron ore, the iron oxides in the ore released oxygen gas and the process resulted in pure iron. Because mass was lost (there was less iron produced that ore used), it was called a reduction.

The chemical equation:

2 Fe2O3 ->  4 Fe (solid) + 3 O2 (gas)

If we consider what is happening to the iron atoms, we realize the Fe ions are gaining electrons to become pure iron metal.

4 Fe+2 + electrons –> 4 Fe0

A few hundred years later, early chemists realized that oxygen was the gas being released and that adding oxygen to metals caused the formation of metal oxides. This led to the idea that the reverse of reduction involved gaining of oxygen, or oxidation.

We now know that oxygen doesn’t have to be involved in oxidation-reduction reactions and that reduction is actually the gain of electrons by atoms, but we are stuck with those historical, relatively-inaccurate names.

All is not lost, however, because chemistry students have come up with some tricks to remembering the terms.

1. This is a common way to remember the terms:

leo-lion(Lion image by Petr Kratochvil at publicdomainpictures.net)

LEO = Loss of Electrons is Oxidation

GER = Gain of Electrons is Reduction

2. Another version is OILRIG , which is short for Oxidation is Loss (of electrons) and Reduction is Gain (of electrons).

3. You may also think of a gain of electrons as increasing a negative charge, or in other words, reducing the charge of the ion to a smaller number.

Why look at Reduction-Oxidation or Redox?

How easily a metal loses electrons will help predict how reactive it is and its behavior when mixed with other elements. Redox states help chemists figure out the likelihood certain reactions will occur.

Another use is to figure out the amount of certain substances in samples by redox titration, similar to acid-base titration. For example, it is possible use a redox titration to find out the amount of vitamin C (ascorbic acid) in different types of food or juice.

Let’s finish up with a discussion of Redox Reactions by Mr. Anderson of Bozeman Science.

Please let me know if you have any questions or comments about oxidation-reduction.

Lesson 19: More About Rates of Chemical Reactions

Last week was an introduction to rates of chemical reactions. This week we finish the chapter with Le Châtelier’s principle and speeding up reactions using catalysts.

Textbook Reading: Finish Chapter 15, pp. 546-566.

Le Châtelier’s principle helps predict the effect of disturbances to equilibrium in reversible reactions.

So, how do you pronounce Le Châtelier? Mr. Anderson at Bozeman Science has the answer:

The activation energy is the amount of energy that must be added to a system for two substances to react to form products. Some reactions don’t need much energy to proceed, like our sodium bicarbonate and acetic acid reaction last week. Others, like mixing hydrogen and oxygen gas to form water, take a lot of energy for the product to be formed.

activation-energy

In our lab we are going to take a look at how adding a catalyst can allow reactions to go forward with less added energy. The catalyst is not part of the reaction, that is it doesn’t end up in the product, but does allow reactions to happen that might not otherwise occur. Thus, catalysts can speed up reactions.

Please let me know if you have any questions.

Lesson 18: Rates of Chemical Reactions

Reactions are where it is at in chemistry.

Textbook Reading:  First part of Chapter 15, sections 15.1-15.6, pp. 531-top of 546.

What controls the rate of chemical reactions?

According to the collision theory, when more atoms and molecules in a system are undergoing collisions, there is a higher chance they will bump together in such a way that a reaction can occur. For a reaction to occur particles must collide with enough energy and in the correct orientation. Even if correctly oriented, not all collisions successfully produce products because not all particles have minimum energy needed for the reaction to occur, or activation energy.

Hank Green of Crash Course Chemistry has a concrete way of explaining collision theory, activation energy, and reaction rate: demolition derby!

Another useful (but not too serious) analogy for increasing the rate of reaction can be found in this TED video: How to speed up chemical reactions (and get a date) by Aaron Sams.

So in summary, to speed up chemical reactions you need to:

1. Shove the reactants together by increasing the pressure.
2. Increase the concentration of the reactants or number of particles.
3. Heat the mixture to increase the energy level and number of collisions.
4. Increase the surface area of the reactants.
5. Add a catalyst to lower the activation energy level to what is available in the system.

We’ll be trying some of these methods in lab.

 

Edit: I’m adding this video from Mr. Anderson at Bozeman Science, too. He talks about the equilibrium constants.

Please let me know if you have any questions.

Lesson 17: Acids and Bases

Acids and bases are everywhere. They are in our food, household products, even in our own bodies! They are relevant and relatable.

Textbook Reading:  Chapter 14, Acids and Bases, pp. 487-521.

People have known for centuries that acids:

  • Taste sour (like lemons)
  • Dissolve/corrode metals
  • Turn blue litmus paper red

On the other hand, bases:

  • Taste bitter (like caffeine in coffee)
  • Feel slippery
  • Turn red litmus paper blue

In addition to these criteria, chemists have been refining and honing their definitions of acids and bases.

In 1884, Svante Arrhenius from Sweden realized an acid is a material that can release a proton or hydrogen ion (H +) when in aqueous solution and a base releases hydroxide ions.

Because this definition did not apply to non aqueous solutions, other scientists continued to refine the definition until two scientists in 1923 came up with a definition that would work for any situation. The Brønsted-Lowry definition says that an acid is a proton donor and a base is a proton acceptor.  Thus ammonia, which has no hydroxide group, still acts as a base by accepting a proton.

Some molecules, such as water, can act either as an acid or as a base according to this definition. Molecules that can act as an acid or a base are called “amphoteric.”

Some acids are defined as “strong” and some as “weak,” generally based on how much they ionize. Strong bases dissociate completely, whereas weak bases only dissociate slightly.

This short video from TED explains these terms.

pH Scale

Chemists measure how acidic or basic a substance is using the pH scale. Although no one knows for sure how the name came to be, it is acceptable to think of pH as the “power” of hydronium ions, thus how many hydronium ions are present. (A hydronium ion is a water molecule with an extra proton – H3O+. Although often used interchangeably with hydrogen ion, hydronium is more technically correct.) Thus, pH = -log [H3O+], where the brackets mean “concentration of.”

When the concentration of hydronium ions is high (pH less than 7), the substance is said to be acidic. If the pH is =7, then the substance is neutral and if the pH greater than 7, then the substance is basic.

Although the scale is often labelled from 1-14 or 0-14, there are really no limits to the ends. Substances have been found with a negative pH, but it turns out that it is very difficult to measure the hydronium ion concentration in the negative range.

pH-scale

Bozeman Science has a particularly clear overview of pH in this video:

Be sure to let me know if you have any questions about these materials.

Lesson 16: Solutions

Now that we have a better understanding of liquids, it is time to revisit solutions and solubility. I’m sure you will be pleased to learn you will finally be finding out what the 6M in 6M HCl means.

Textbook Reading: Chapter 13, pp. 447-477. Molarity starts on page 457.

Definitions review:

We already touched a few of these concepts in an earlier lab, so it should be review.

The solute is the smallest part of a solution, or the substance being dissolved.

The solvent is the larger part, or the part doing the dissolving.

A solution is a solute dissolved in a solvent.

The text mentions temperature can have different impacts on solubility.

You probably have direct experience trying to add sugar to a cold drink versus a hot drink. Isn’t it easier to stir in sugar when the liquid is warmer? Solids in general are more soluble at high temperatures and less soluble at low temperature.

SolubilityVsTemperature(image pubic domain)

As you can see from the graph, however, not all salts follow this general trend.

Gases

Think about how quickly a warm soda goes flat. Gases in solution react just the opposite of solids. In general, gases are more likely to stay in solution at low temperatures than high ones.

Molarity

Now it is time to learn how to quantify solutions. One way is to calculate the molarity, which is the number of moles of solute per liter of solution.

molarity-equation

More review:
Do you remember what a mole of something is?

(A mole of a substance is simply Avogadro’s number or 6.022 x 1023 items of that substance.)

Do you remember how to calculate the molar mass from Lesson 7?

(Use your periodic table to find all the atomic mass units (amus) for the atoms in the molecule and then add them together and convert to grams.)

Mr. Causey walks us through the process of calculating molarity and making dilutions in this video:

Note: Some of you skip these videos, but because molarity is such a big part of doing chemistry, I really recommend you spend the time with this one.

In addition, PhET has an awesome interactive about molarity. Be sure to click the “show values” box to really see what is going on.

Please let me know if you have any questions about your readings or this lesson.

Lesson 15: Properties of Liquids and Solids

Intermolecular forces rule when it comes to liquids and solids.

Textbook Reading: Chapter 12, pages 411-437.

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Let’s start out with an introduction to liquids and solids by Bozeman Science.


 

Intermolecular forces

Remember the forces that hold atoms together to form molecules that we learned about in chapter 10: covalent bonds, polar covalent bonds and ionic bonds? Now we are going to find out about forces between molecules or intermolecular forces, which is what solids and liquids are all about.

The three types of intermolecular forces are dispersion forces (also called London forces), dipole-dipole forces and hydrogen bonds.

Dispersion forces were first recognized by Fritz London, which is why they are often called London forces. They are found between all molecules (both polar and nonpolar) and a formed due to temporary unequal sharing of electrons. In larger atoms or molecules, the valence electrons are, on average, farther from the nuclei than in smaller ones. The electrons are thus held less tightly and can more easily form temporary dipoles. Because of this, larger and heavier molecules exhibit stronger dispersion forces than smaller ones.

This quick video shows how the electrons move around an atom, but the forces work in a similar way around molecules as well.

Dipole-dipole forces occur between molecules that are permanently dipolar. You can figure out whether a molecule is dipolar by examining the differences in electronegativity between the atoms and also by examining the molecule’s shape.

dipole-force-graphic

Hydrogen bonds are the strongest type of intermolecular force. They form in the special case of hydrogen bonded to fluorine, nitrogen or oxygen. Although very specific, hydrogen bonds are not uncommon and in fact form important links between molecules in our DNA.

Base_pair_GC.svg

Hydrogen bonding between guanine and cytosine in DNA

(Illustration public domain from Wikimedia)

Intermolecular forces determine the physical properties of liquids and solids, such as surface tension, viscosity, melting point, boiling point, volatility, etc. We will examine some of these properties in lab.

This final video goes in depth about physical properties and intermolecular forces. Stick with it and you will find out why Kevlar is so strong!

Please feel free to contact me if you have any questions about this week’s lesson.

Lesson 14: Gases

Remember the states of matter? This week we learn more about gases, next week we’re on to liquids and solids.

Textbook Reading: Chapter 11, pages 359-399.

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Kinetic Molecular Theory

The kinetic molecular theory reveals the properties of gases relative to liquids or solids. It assumes that the molecules are very small relative to the distance between them, that the molecules are in constant and random motion, and that they frequently collide with each other and with the walls of any container without interacting. The average kinetic energy of gas molecules depends on temperature.

Translational_motion

(Translational motion- Gif by A.L. Greg at Wikipedia CC BY-SA 3.0) Caption:  “The temperature of an ideal monatomic gas is a measure of the average kinetic energy of its atoms. The size of helium atoms relative to their spacing is shown to scale under 1950 atmospheres of pressure. The atoms have a certain, average speed, slowed down here two trillion fold from room temperature.”

The Simple Gas Laws

Our understanding of how gases behave started with the work of some early scientists. They are now named Boyle’s Law, Charles’s Law, and Avogadro’s Law.

gas-laws

The ABC’s of gas: Avogadro, Boyle, Charles by Brian Bennett for TED

Later, it was realized these laws were related and the same man who gave us the periodic table, Mendeleev, put them all together in the Ideal Gas Law.

PV=nRT

Note: The Ideal Gas Law ignores some factors about real gases, such as the particles do have mass and that they do interact with each other and their surroundings. Corrections have to be made to adjust for these deviations.

Mr. Anderson at Bozeman Science gives a detailed overview of gases, some cool animations and a way to estimate absolute zero!

Please feel free to contact me if you have any questions about this week’s lesson.

Lesson 13: Chemical Bonding

Time to put those atom models to use and make molecules.

Textbook Reading: Chapter 10, pages 325-349.

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Lewis Dot Diagrams and Structures

Last week we learned that the outer electrons, called valence electrons, of an atom are the ones involved in bonding and chemical reactions.  Lewis structures consist of the element’s symbol surrounded by dots to indicate the valence electrons.

lewis-stuctures

There are a few rules for creating Lewis structures:

1. Find the number of valence electrons for a given element using the group numbers of the periodic table. There will never be more than 8.

2. Represent the valence electrons by placing dots on four sides around the symbol for the element.

3. Start filling with single dots. If there are 5 or more valence electrons, pair them after the four single dots have been placed. Exception:  Helium has a single pair of two dots because that is its stable configuration (see illustration above).

4. Exact location of dots can vary (which side placed on doesn’t matter).

Bozeman Science has an in depth explanation about Lewis structures. (Those who like their chemistry to “pop” will enjoy the first part of this video. The last part about Lewis himself is sad.)

When drawing molecules, the sharing of two electrons is represented by a line connecting the symbols. If the two atoms share two pairs of electrons, then the resulting double bond is shown as two lines.

Be aware that as powerful as it is, there are exceptions to the octet rule!

Valence Shell Electron Pair Repulsion (VSEPR)

Remember how the magnets of the same pole repelled each other, pushing away? The VSEPR model uses that concept to predict the shape of molecules.

Some of the structures we will investigate are:

  • linear
  • trigonal planar
  • tetrahedral
  • trigonal pyramidal
  • bent
  • octahedral

(See the Sidebar about VSEPR post to see these shapes)

Mr. Isaacs at IsaacsTeach has an introduction to molecular geometry that explains why water molecules are bent.

Bozeman Science pulls the Lewis Structures and VSEPR together for a comprehensive overall of drawing molecules. He goes into a bit more depth than is covered in your text.

Electronegativity and Polarity

Another aspect of molecules that determines their shape and chemical properties is the ability of the atoms of an element within a bond to attract electrons, or its electronegativity.

You can read the electronegativity of an atom from a specially designed periodic chart like this one:

Periodic_Table_of_Electronegativity_

(Illustration from UC Davis ChemWiki is licensed under a Creative Commons Attribution-Noncommercial-Share Alike 3.0 United States License.)

Non-polar bonds form when the difference in electronegativity between the two atoms is between 0 and 0.4, polar bonds form when the difference in electronegativity between the two atoms is between 0.4 and 2.0, and ionic bonds form when the difference in electronegativity between the two atoms is greater than 2.0.

Why does this matter? Knowing whether a molecule is polar helps predict its characteristics, such as whether it will be soluble in a given solute.

That’s it for this week. Let me know if you have any questions!

Lesson 12: Electrons, Atom Models, and the Periodic Table

After a hectic week, you will be happy to hear we don’t have a lab to write up in your notebooks for this lesson. That means you don’t have to bring your goggles or gloves either. You do, however, need to really focus on the readings in the textbook because this lesson is critical for your understanding of chemistry.

Textbook Reading: Finish Chapter 9, sections 9.5-9.9 on pages 294-315.

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Supplemental:

Introduction:

Electrons are where all the action is in chemistry. By developing sophisticated and complex models of atoms that explain electron behavior, chemists have a powerful tool to understand molecular structures and chemical reactions.

Our most recent model of the atom comes from the study of quantum mechanics. The theory and math behind the model are pretty complex, but chemists have been pulling out some basic concepts that can be extremely useful. Remember, however, that this is a model and may be modified as our understanding increases.

Electrons:

We have some awesome videos this week to supplement your textbook.

Let’s start out with a TED video about the uncertain location of electrons.

Although this video suggests orbitals are where electron can be found 95% of the time, other representations of orbitals are often where the electrons are found 90% of the time. In any case, you can think of orbitals as boxes or rooms where electrons are found most of the time. You should also remember that each orbital can only hold two electrons.

For the musicians in the class, this Crash Course video has an explanation of electrons that might just “resonate” with you 🙂

 

 

Okay, if you don’t know enough about music to understand his analogy, then we can use a simpler analogy.

Single_electron_orbitals

(Illustration of single electron orbitals from UC Davis ChemWiki, licensed under a Creative Commons Attribution-Noncommercial-Share Alike 3.0 United States License.)

Orbitals are grouped around the nucleus in very specific ways, based on the number of electrons and their energy states. The groupings are given names: shells and subshells. If orbitals are rooms where electrons are found most of the time, then subshells are clusters of rooms. You might think of them as one apartment in a large apartment building, or one set of offices in an office complex.

Continuing the building analogy, the shells would be the different floors of the building. The higher floors have more energy and contain more subshells.

The shells and subshells fill with electrons in a very orderly way, so that we can figure out the electron configuration of atoms of each element in the periodic table simply based on the number of electrons it contains. The electron configuration, although it looks complicated, is simply the arrangement of the electrons.

Bozeman Science has a serious explanation of electron configurations. He relates the configurations to “ionization energy” or how hard it is to pull an electron off an atom.

Isn’t the idea that the p orbitals fill like seats on a bus helpful?

Still unsure what all this means? Don’t worry, we’ll being going over it all in class. Be sure to write down your questions and bring them with you.