Blog Archives

October 3 - 26, 2017

9/25/2017

Housekeeping: I want to check your understanding of chemical bonding. So take out the practice problems from last week and we will go through them. Then we will review for your quiz. If you are still struggling with nomenclature and formula writing, now is the time to speak up.

Agenda:
1. Chemical bonding review
2. VSEPR Theory

Content Review:
Text: Chapters 10 & 11
Links: Chemical Equations The Mole

Student Missions:
Mission 1: Getting in Shape!
Mission Objectives: You should be able to...
1. Determine the structure of simple molecules using VSEPR Theory.
2. Predict the molecular geometry for simple compounds.
3. List and describe the three kinds of intermolecular forces.

I already have these lessons prepared for another class. Click on the links.
VSEPR Theory
Intermolecular Forces

More on intermolecular forces: Chemical bonds are stronger than intermolecular forces and more energy is required for their formation.

Temporary dipoles occurs when electrons are dense on one end of a molecule. All molecules can become temporary dipoles. Permanent dipoles are molecules which have a clear negative and positive end (two poles) resulting from polar bonds.

Water molecules have high intermolecular forces resulting from hydrogen bonding, which is the strongest IMF.

IMFs can determine boiling points and states of matter. Read about it here and here.

Mission 2: The Five Families.
Mission Objective. You should be able to...
1. Identify the five kinds of chemical equations
2. Balance chemical equations

I don't feel the need to get redundant when we can just go here. It is expected that you will be able to identify/classify any equation that is given to you based on what the reactants look like. So pay attention!

Download this handout and practice writing equations. You are turning formulas into words, which we have been doing since last week.

Mission 3: Count 'em Up!!
Mission Objective. You should be able to...
1. Balance chemical equations.
2. Define and understand the significance of "delta" in a chemical equation.

The basic rule about balancing equations is this: The number of atoms on the reactant side should equal the number of atoms on the product side. The easiest way to do this is to identify individual atoms and count. Add coefficients to the beginning of chemical formulas to make the math work. Coefficients go AT THE BEGINNING of a formula; nowhere else. For polyatomic ions, if they appear on both sides of the equation, you can keep them together as a unit. It makes it easier to count that way.

In chemical equations, a delta (triangle) is used over the arrow to denote that heat is used so that the reaction can take place. You are expected to know this.

Here is an adorable conceptual video to help you understand how and why equations must be balanced. Hint: we cannot go around breaking laws and stuff! The second equation gives you five simple rules for balancing equations.

Mission 4: The Mole.
Mission Objectives. You should be able to...
1. Define the mole
2. Complete mole conversions using mass and molar mass.

The mole is the SI unit for counting particles. One mole of any substance contains 6.02 * 10^23 particles. This number is known as Avogadro's Number, named after Italian scientist Amedeo Avogadro.

"Particles" is a generic term for atoms, formula units, and molecules. Avogadro's number of atoms clearly references an element, Avogadro's number of formula units references an ionic compound, and Avogadro's number of molecules references a covalent compound.

The molar mass of a substance is either the element's atomic mass on the periodic table, or the sum of the atomic masses of a formula unit or a compound.

For example: Hydrogen's molar mass is 1.01 g/mol. But water, which is H2O, is 18.02 g/mol. The 18.02 comes from adding the two hydrogens (2.02) with oxygen (rounded up to 16.00). So whenever you are dealing with the molar mass of a compound, you must add up the atomic mass of all of the atoms in that compound.

Sample mole problems:

1) How many moles are in 25 grams of water?
2) How many grams are in 4.5 moles of Li2O?
3) How many molecules are in 23 moles of oxygen?
4) How many moles are in 3.4 x 10^23 molecules of H2SO4?
5) How many molecules are in 25 grams of NH3?
6) How many grams are in 8.2 x 10^22 molecules of N2I6?

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September 18 - 29, 2017

9/18/2017

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Housekeeping: We are now moving into chemical bonding.

Agenda:
1. Discussion of ion formation.
2. How ionic bonds are formed.
3. How covalent bonds are formed.
4. Determination of polarity.

Content Review:
Links: Chemical Bonding This page contains PowerPoints that cover ionic and covalent bonding and a summary table.

For those of you that prefer PowerPoints, here they are:
Ionic Compounds Covalent Compounds Chemical Names & Formulas
Textbook Readings: Chapter 7, 8 & 9, respectively.

Student Missions:
Mission 1: Ions, Ions, Ions!
Mission Objectives. You should be able to...
1. Determine which elements form cations and which elements form anions.
2. Contemplate the structure of an ionic bond based on valence electrons.
3. Predict the name of an ionic compound and a covalent compound.

Recall that the group number on the periodic table represents the number of valence electrons in that group. Also recall that group 18, the noble gases, have stable octets. The objective for any element that is not a noble gas is to lose or gain electrons to have a noble gas configuration. Stability is key.

Elements in G1 will LOSE their valence electrons to form ions with a +1 charge. Elements in G2 will LOSE theirs to form ions with a +2 charge. Elements in G13 will lose theirs to form ions with a +3 charge. Any time an element LOSES a valence electron, a cation is formed. Cations are positively charged ions. Metallic elements form cations.

Conversely, elements in G15 will GAIN 3 electrons to form ions with a -3 charge. Elements in G16 will GAIN 2 electrons to form ions with a -2 charge, and elements in G17 will GAIN 1 electron to form an ion with a -1 charge. Any time an element GAINS a valence electron, an anion is formed. Anions are negatively charged ions. Nonmetals form anions.

This sweet little 10-second interactive shows what happens to the size of atoms when ions are formed. Make note of which elements will get smaller and which will get bigger.

For more on ion formation, see what David Weightman has to say about it. It's three minutes long.

Ionic compounds are formed when a cation bonds with an anion through the transfer of valence electrons. Another way of saying this is that an ionic compound forms when a metal bonds with a nonmetal.

For example, the sodium atom will transfer its one valence electron to chlorine, an atom with seven valence electrons. Sodium does not need it's valence electron, and chlorine craves to have one more valence electron, so the ionic bond is easily formed between them. This pattern holds true for any G1 element bonding with a G17 element. The pattern also works for G2 and G16 elements, and G13 and G15 elements. I call these perfect pairings, because they result in 1:1 compounds.

Of course, nothing is perfect because any metal can bond with any nonmetal. So there are compounds with ratios greater than 1:1. It's not difficult because it's all about balancing charges. Tyler DeWitt talks you through it. It would be helpful to have your periodic table handy.

Practice with this interactive.

Mission 2: Sharing is Caring!
Mission Objectives. You should be able to...
1. Determine when covalent bonds are formed and which elements form them.
2. Predict the name of an ionic compound and a covalent compound.

Covalent bonds are the result of valence electrons being shared in order to achieve stable octets. If you think back...way back...to the short lesson on Lewis Dot, you'll recall that certain elements contained a certain number of dots that represented their valence electrons.

Anyhoo, if you take a look at groups 15, 16 & 17 (the nonmetals), this is where covalent bonding takes place. These elements can bond with themselves to share a pair (or two, or three) of electrons to form stable octets. Take a look below. Fluorine will share a pair of electrons with another fluorine and a single covalent bond is formed (represented by a line).

Group 16 elements will share two pairs of electrons (double covalent bond) and group 15 elements will share three pairs of electrons (triple covalent bond). Again, the goal is stability. The Octet Rule rules all.

There are seven elements that are found in nature covalently bonded to each other. They are called diatomic elements. A useful memory aid is HONI Bring Fried Clams. So when you're writing these elements in chemical equations, you need to make sure you write them correctly.

Coordinate covalent bonds are bonds where one atom contributes both bonding electrons. You demonstrate them in structural formulas using an arrow that shows where the electrons came from. Once formed, the coordinate covalent bond is just like any other bond. Carbon monoxide is a good example of this.

Image courtesy of mrdchemcwiki.wikispaces.com

Polyatomic ions are a group of atoms that are tightly bound that behaves as a unit and has either a positive or negative charge. I gave you an ion reference chart, which contains a comprehensive list of polyatomic ions. You need to print a copy and keep it with your periodic table. Don't lose it.

Get familiar with covalent bonding using this interactive.

We will practice writing and naming ionic and covalent compounds.

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September 19 - 27, 2017

9/12/2017

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Housekeeping: Looking at the rest of this month, here's what I have scheduled:
This week, periodicity. Next week, there needs to be an exam over atomic structure and periodicity. The best (and only) date to do it is September 20th. The week after that, you guys will do a Meet the Elements lab on September 29, 2017.

The cereal box/cookie box is due on Tuesday, September 19.

Agenda:
1. Project Update
2. Electron Behavior
3. Discussion and analysis of periodic trends

Content Review:
Links: The Periodic Table Periodic Trends
Textbook Readings: Chapter 6.

Student Missions:
Mission 1: What We Already Know.
Mission Objectives. You should be able to...
1. Explain how Mendeleev organized the first periodic table.
2. Describe the ways in which the periodic table is organized.
3. Name the families of the periodic table.

The PT has a long and sordid (I wish) history and is the result of the work of many scientists. Dmitri Mendeleev gets the most credit because he found an organizational system that worked better than others and could make predictions about future elements. The Royal Society of Chemistry provides an in-depth review which you should read outside of class. However, Mendeleev's version of the PT didn't quite work (which element was the one to cause him problems?) because he ordered the elements by increasing atomic mass. We now know, thanks to Henry Moseley, that elements are ordered by increasing atomic number.

Mission 2: It's PERIODIC Because...
Mission Objective. You should be able to...
1. Explain the relationships between atomic number and ionization energy, atomic radius, electronegativity, and ionic radii.

...the organization of elements via atomic number showed properties (with a magical number of 8, of course) that were...wait for it...periodic! Elements in the same columns tend to behave the same way. We now know this is because of their valence electrons.

Because of periodicity, trends in elements can be predicted. We discuss four types of trends: ionization energy, electronegativity, atomic radius and ionic radius.

Atomic Radius (AR)

As atoms gain protons, their sizes decrease because of the increased positive pull of the nucleus on the valence electrons. Think in terms of a game of tug of war; with the nucleus on one end and the valence electrons on the other. Adding more protons pulls the valence electrons closer. Atomic number increases from left to right, so atoms decrease in size from left to right. Atomic number increases from top to bottom, so atoms increase in size from top to bottom.

Therefore, atomic radii decrease across a period and increase down a group.

Ionic Radius (IR)

Metals form cations easily, and nonmetals form anions easily. Cations are always smaller than the neutral atoms from which they form, because the stripping of the valence electrons increases the attraction from the nucleus. As a result, the nucleus pulls the remaining electrons closer. Anions are always larger than the neutral atoms from which they form, because adding electrons decreases the nuclear pull.

Therefore, just like atomic radii, ionic radii decrease across a period and increase down a group.

Ionization Energy (IE)

When an atom gains or loses an electron, it becomes an ion. Positively charged ions are called cations and negatively charged ions are called anions. The energy required to overcome the attraction of the nuclear charge and remove an electron is called ionization energy.

Removing one electron results in the formation of a cation with a +1 charge (and the energy required to do so is called the first ionization energy). Removing two electrons (which is more difficult) results in a cation with a +2 charge (second ionization energy). As the ionization energy increases, it becomes more and more difficult to remove electrons. Larger atoms' valence electrons are further away from the nucleus, and therefore have lower IEs.

Similarly, adding electrons increases the negative charge. Adding one electron results in the formation of an anion with a -1 charge. Adding two electrons results in an anion with a -2 charge, and so on and so forth.

Therefore, ionization energy increases across a period and decreases down a group.

Electronegativity (EN)

The electronegativity of an element is the tendency of its atoms to attract electrons when they are chemically combined with another element. Noble gases do not have electronegativity values because they do not readily form compounds; they are inert. Electronegativity follows the same pattern as ionization energy and the opposite pattern of atomic radius.

Therefore, ionization energy increases across a period and decreases down a group.

This summary image is courtesy of WilenskyChemistry.

Mission 3: Going Backwards.
Mission Objectives. You should be able to...
1. Briefly describe the arrangement of electrons in an atom.

As you're learning about atomic theory and the development of the atom, you should discover that most of the models do not specify the arrangement of electrons. JJ Thomson acknowledged their existence, but his model left a lot to be desired in terms of arrangement. Nagaoka's model is similar to Bohr's model in that the electrons orbit the positively charged nucleus. Bohr's model works for hydrogen, but for successively higher atomic numbers, the model fails. Bohr predicted that electrons orbit the nucleus in energy levels similar to how planets orbit the sun, a two-dimensional orbit.

However, we now know that electrons orbit the nucleus in 3-D regions called orbitals. The currently accepted atomic model, the quantum mechanical model, is a mathematical model that is based on the probability of locating an electron. 90% of the time, electrons can be found in these orbitals.

Below you have the s and p orbitals.

Electrons obey three rules in their arrangement. You can read about them here. As atomic number increases, electrons fill energy levels and sublevels in an orderly fashion. The simplest sublevel is "S." The above image shows the S sublevel/orbital on the left. The image on the right is the P sublevel. You can look up D and F sublevels/orbitals. They're really funky-looking. As energy level increases, the size of the sublevel/orbital increases. Electrons in the outermost energy levels are called valence electrons. These electrons are the ones that determine chemical behavior.

Below are a series of videos that goes into detail about the behavior of electrons. We will watch each video and analyze the information. You're required to be familiar with the first 20 elements (H through Ca).

Homework: Review the videos and take a look at this practice test. See if you can answer the questions. The test comes with an answer key so you can check your work. If there is a reason why you feel you can't answer a question, then that is a question you must ask me during class.

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September 5 - 15, 2017

9/4/2017

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Housekeeping: You have a quiz tomorrow over macromolecules. 20 MC and one or two short answer. You'll have all of 30 minutes to complete the quiz. We will review before you take said quiz.

You also owe me two labs: the macromolecule investigation and the separating mixtures lab. The macromolecules lab is due tomorrow (you should have been done with this last week). The separating mixtures lab will be due Friday.

You will finish the separating mixtures lab today in class and then we will go over what you will submit.

Agenda:
1. Housekeeping
2. Finish the lab
3. Discussion of write up/submissions
4. Atomic Theory & Structure (I)

Content Review:
Weebly Links: Matter
Weebly Links: The Atom   Atomic Theory  Nuclear Chemistry   Electrons
Link: Substances & Mixtures Elements & Compounds   Matter, Elements & Atoms

Student Missions:
Mission 1: Itty Bitty Things...Rack 'Em Up!!!
Mission Objectives. You should be able to...
1. Describe the structure of the atom.
2. Explain how atoms differ.

All neutral atoms contain the same number of protons & electrons. The number of protons determines the element's identity.  For instance, 8 protons = oxygen. 17 protons = chlorine. 20 protons = calcium. This does not change.

Electrons determine chemical behavior.  Valence electrons (electrons in the outermost energy levels) are significant in this regard, because the number of valence electrons determine how an element behaves in certain conditions. Elements with an octet (8 valence electrons) are unusually stable and do not combine to form compounds (noble gases have an octet, with the exception of helium).

Neutrons determine isotopes.  They do not affect the charge or the element's identity. However, they do affect the mass of the nucleus. Several elements have multiple isotopes.

Mission 2: Elements And Their Cousins.
Mission Objectives. You should be able to...
1. Calculate the atomic mass of an element.
2. Identify isotopes of various elements.

Atoms with different numbers of neutrons are called isotopes. For instance, carbon has fifteen known isotopes, but only three are commonly referenced: carbon-12, carbon-13, and carbon-14.

All three isotopes have six protons (carbon's atomic number), but C-12 has 6 neutrons, C-13 has 7 neutrons, and C-14 has 8 neutrons.

Isotopes are usually written in a form called standard notation, which includes the mass number (A), the atomic number (Z) and the symbol of the element (X).

Have you ever wondered why, if atomic numbers represent the number of protons (and theoretically, neutrons) in the nucleus, and they're assigned masses of 1 each, then why are the given atomic masses on the periodic table decimals? In other words, if the mass number (protons + neutrons) are whole numbers, then why is the atomic number a decimal?

Well, that's because the average atomic mass is the average of all isotopes for any one element. In order to calculate the AAM, you need to know the percent abundance of each isotope (this should be a percentage, which you turn into a decimal), and the atomic mass of each isotope. DO NOT ROUND THESE NUMBERS. You will do your rounding at the very end.

Multiply the percent abundance (now a decimal) by its atomic mass. Do this for each isotope. Then add all the isotopes together. Your answer should equal the value that is listed on the periodic table.

ChemWiki went hard by demonstrating how to calculate AAM with pictures.

Wanna practice? Of course you do!

Here's a short video showing how it's done.

Let's Practice! I have a handout (or two) that you need to complete. It's important that you get the hang of how to do these calculations. I smell a quiz coming next week...maybe.

We are going to do the Candium Lab, where you determine the average atomic mass of a bag of candy. This lab is your first write-up, and I've included the template for you to follow. It is important that you follow the instructions and ask questions when you aren't sure whether something should be included. You'll need a calculator.

Mission 3: Atomic Theory. How did we go from "nothing" to "something" to "something really complex?"
Mission Objectives. You should be able to...
1. Trace the development of atomic theory from ancient times to now.
2. Explain why models have to be revised and refined.
3. Describe how our understanding of the atom has changed over time.

What we know about the atom has changed over time because of technological advances, obviously. The book does not provide enough historical context, so you can add to your knowledge base by reading up on atomic theory. Ask yourself, "How is the ability to understand a fact of nature is limited by our ability to observe it?"

Dalton's atomic theory has four parts. Name them. What makes his theory so significant for its time (early 1800s)?

You will track the development and refinement of atomic theory and resultant models from Democritus to Bohr. I will put you into groups and assign you a theorist/model. There will be pictures required. You will present your research on _________________.

Regardless of your assigned model, be able to answer this question:  Why is the Bohr model still used as a reference even though it is no longer the accepted model? What is the currently accepted model?

Some links to help you are below.
1.  BBC Bitesize
2.  Modern Atomic Theory
3.  Early Atomic Theory
4.  Atomic Theory Timeline
5.  The Quantum Mechanical Model

Questions you should be able to answer:
1. Who is your scientist/theorist and what were their scientific qualifications? No more than 3-4 sentences explaining this.
2. What specific particle did they study, or did they study the atom in general?
3. What kinds of experiments did they do? Describe in detail.
4. What was the name of their atomic model?
5. Provide a picture of the model.
6. Why did the model eventually have to be replaced, if it was replaced? In other words, models change because they can no longer sufficiently explain certain phenomena. Why did your assigned atomic model have to be replaced?

Homework:  Begin working on your presentations. Grading Rubric.

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