Blog Archives

August 24 - 31, 2017

8/23/2017

Housekeeping: We are now in Chapter 03: Periodicity. This chapter is very short, so we can knock it out this week and take the exam over it and the atom within the next two weeks.

Agenda:
1. Discussion of the history of the periodic table
2. Periodic Trends activity

Content Review:
The Periodic Table Periodicity
Textbook:

Student Missions:
Mission 1: As Usual, HIStory.
Mission Objectives: You should be able to...
1. Track the development of the periodic table from Antoine Lavoisier to Henry Moseley.
2. Explain why Mendeleev's work was so significant.

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. Define, describe and explain periodic trends across periods and down groups.

...because 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.

You can review these basic trends on the periodicity webpage. Be sure you know what each one represents.

Your textbook goes into some detail about trends in electron affinity, metallic/nonmetallic character, and properties of metal and nonmetal oxides.

Electron affinity is defined as the energy required to detach an electron from the single charged negative ion in the gas phase. This process is exothermic. (Ionization energy is endothermic) The more negative the electron affinity value, the greater the attraction of the ion for the electron. You will need to reference the data booklet to see the values for the elements.

Metallic character decreases across a period and increases down a group. It follows the same trend as atomic radius. Metals have low IEs and tend to lose electrons. Therefore they can be oxidized. Nonmetals have high IEs and tend to gain electrons. Therefore they can be reduced.

Metal oxides are basic; they react with water to form metal hydroxides (see p. 86). Nonmetal oxides are acidic; they react with water to form acid solutions.

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August 14 - 25, 2017

8/6/2017

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Housekeeping: There will be some time this week where we will discuss your IA topics, and a period where you guys will be allowed class time to work on your atomic theory timeline cards. I'm still debating on whether to quiz this week or next week. I'm leaning towards next week.

Agenda:
Review: Mass Spectrometry & Quantized Energy
Electron Configuration

Content Review:
Atomic Theory Structure Electron Behavior Electron Configuration

Student Missions
Mission 1: Review.
Mass spectrometers are designed to determine the relative atomic masses of elements. There are five stages in the process: vaporization, ionization, acceleration, deflection and detection. Samples are heated and vaporized to produce gaseous atoms and molecules and then bombarded by high energy electrons, generating positively charged species. he positive ions are accelerated into an electric field and deflected by a magnetic field that is perpendicular to their path (recall the video).

The degree of deflection depends on the mass to charge ratio (m/Z). The species with the smallest mass (m) and highest charge (Z) will be deflected the most. Particles with no charge are not deflected. The detector then detects species of a particular m/z ratio. The ions hit the counter and the electric signal is detected. It is a plot of relative abundance (of each isotope) versus mass number (A). Peak height indicates the relative abundance of the isotope.

Study the worked example on pages 48-49 in your text.

The electromagnetic spectrum is a spectrum of wavelengths that comprise the various types of electromagnetic radiation. Energy is inversely proportional to wavelength. High energy radiation has very small wavelengths and low energy radiation have very long wavelengths. SI unit for energy is the joule (J), wavelength is measured in meters (m), and frequency is measured in hertz (Hz).

When white light is heated and/or passed through a prism, a continuous spectrum is produced. If a pure gaseous element is subjected to an electrical discharged, the gas emits radiation. The result is an emission spectrum. If a cloud of cold gas is placed between a hot metal and a detector, an absorption spectrum is produced. The gaseous atoms absorb certain wavelengths (which show up on the emission spectrum). See below.

The lines in an emission spectrum have specific wavelengths. Each corresponds to a discrete amount of energy. This is called quantization. A photon (discrete packet of light) is a quantum of radiation and relates wavelength and energy. Look at the relationship between speed of light and Planck's constant (page 52). As a result, Bohr used this information to describe electron behavior in the hydrogen atom. Bohr's model works only for hydrogen. The reason it doesn't work for other elements is because (1) it assumes that the positions of the electrons are fixed, (2) it assumes that energy levels are circular, and (3) it suggests an incorrect size scale for atoms.

Mission 2: SUPERMODELS!
Mission Objectives. You should be able to...
1. Describe Heisenberg's and Schrodinger's contributions to atomic theory.
2. Explain how the quantum mechanical model (QMM) works.

The Bohr model is the standard, and even though it is outdated, it still provides a simplistic understanding of how electrons behave when energized. Research done by Werner Heisenberg & Erwin Schrodinger in the 1920s & 1930s produced a different understanding of the atom.

The Bohr model fails with elements beyond hydrogen because it assumes that the electron's trajectory can be precisely described. This is not true because any attempt to measure an electron's position disturbs its motion or velocity. This principle is called Heisenberg's Uncertainty Principle: an electron's position and velocity cannot be known simultaneously because manipulating one affects the other. The best that can be hoped for is a probability of where the electron is likely to be.

Schrodinger came along and, using the principles of wave/particle theory, proposed that a wave equation can be used to describe electron behavior in the same way light behavior is described. The equation works for hydrogen and elements beyond hydrogen, and the solutions to the equation are known as atomic orbitals. These are regions of space around the nucleus in which there's a 90% chance of finding an electron. Orbital shape depends on the energy of the electron.

Make note of the chart that she draws to denote the relationship between principal energy level, sublevel, & number of electrons held. This information corresponds to pages 76-77 and is much easier to understand. You are expected to know the shapes and names of s and p atomic orbitals.

Mission 3: There Are Always Rules.
Mission Objectives. You should be able to...
1. Describe the arrangement of electrons in the quantum mechanical model.
2. Draw orbital diagrams and write electron configurations for elements 1 - 36.

There are three rules for writing electron configurations and drawing orbital diagrams. They are listed here. Take a look and make sure you understand each one. Corresponding pages in the textbook: 56 - 58.

This is the shorthand version of all this information: When describing the arrangement of an element's electrons, it is important to understand that (1) you must begin at the ground state, which is the energy level closest to the nucleus (n = 1) and work your way up, (2) there are four sublevels (s p d f) within each principal energy level; each sublevel has a particular shape called an orbital, and (3) each orbital can hold a maximum of two electrons spinning in opposite directions.

This model of the atom, compiled from the research of Schrodinger, Heisenberg, and a few others, is called the Quantum Mechanical Model. It is a mathematical model explained by four quantum numbers.

We will work our way through the first 36 elements. Make sure you have your periodic table on hand.

Homework: page 62 - 63.

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August 7 - 11, 2017

8/6/2017

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Housekeeping: Good afternoon. Welcome to the Mad!Lab!

Agenda:
Welcome
Introduction to the course
Class procedures
Housekeeping (5 min)
Warmup & review (10 min)
Missions (20-25 min)
Wrap-up/homework (5 min)
Navigating the website (Where to find…?)

Content Review:
Atomic Theory   Structure   Electron Behavior   Electron Configuration

Student Missions
Mission 1: Itty Bitty Things.
Mission Objectives: You should be able to...
1. Describe the structure of an atom.
2. Describe how radioisotopes are used in medicine.

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. Almost every element has isotopes, and what is represented on the periodic table is an average of all of the isotopes of a particular element (which is why the atomic mass is a decimal). In a later mission, you will learn how to use mass spectrum data to calculate relative atomic mass (RAM).

This is basic structural information about the atom; what we know for sure. We also suspect something else: that subatomic particles are made up of quarks and all particles have anti-particles that, when they collide, release energy in the form of gamma rays.

Mr. Andrew Weng breaks down the entirety of section 2.1. He's not the most exciting lecturer, but he's comprehensive.

The atomic number, Z: This is the number of protons in the nucleus. Neutral atoms have the same number of protons and electrons.

The mass number, A: This is the number of protons and neutrons in the nucleus. Nucleons are the collective name for the subatomic particles in the nucleus.

The element symbol, X: This is the symbol of the element that is on the periodic table.

Radioisotopes (radioactive isotopes) are used in nuclear medicine for diagnostics, treatment and research. They are also used as tracers in biochemical and pharmaceutical research and as chemical clocks in geological and archaeological dating.

Mission 2: Back That Mass Up!!!
Mission Objectives: You should be able to...
1. Define "relative atomic mass".
2. Explain how a mass spectrometer works.

Relative atomic mass is defined as the ratio of the average mass of an atom to the unified atomic mass unit (1 amu). The book defines 1 amu as 1.6605402 * 10^-27.

A mass spectrometer can be used to measure the mass of individual atoms and the relative atomic mass of an element. Because the mass of an atom is ridiculously small (duh!), we use relative values according to an agreed-upon standard. Carbon-12 is the standard by which all atoms are measured (so I'm sure you can imagine that makes carbon a snarky element who doesn't get invited to the good parties).

There are five stages in this process: vaporization, ionization, acceleration, deflection and detection. You should be able to explain what happens at each stage.

Paul Anderson of Bozeman Science goes into detail about mass spectrometry. We will start the video at 1:45, unless you want to hear about John Dalton's contributions to chemistry (we will get to this shortly).

When mass spec data is produced, it is in the form of a mass spectrum. The X-axis shows the mass/charge ratio of the different ions on the C-12 scale, which is considered almost equal to their mass. The Y-axis shows the relative abundance of the ions.

Image courtesy of chemguide.co.uk

Mission 3: Color Me BADD!!
Mission Objectives. You should be able to...
1. Explain the phenomenon of emission spectra.
2. Describe the line emission spectrum of hydrogen and its relationship to the Bohr Model.
3. Determine how an atom is structured based on its energy level(s).

The electromagnetic spectrum is a visualization of electromagnetic radiation. It ranges from low energy radio waves to high energy gamma rays. All eMag waves travel at the same speed (c) but are distinguished by their different wavelengths (greek letter lambda). Different colors of visible light (the sliver of the spectrum we can see) have different wavelengths. There is a full eMag spectrum in your Data Booklet.

The above image from the Pearson text (p. 71) shows the changing wavelength of the eMag spectrum in meters (m). On the left side, we have high energy gamma waves, which have the shortest wavelengths, and on the right side, we have radio waves, which have the longest wavelengths. The sliver of color between UV and IR is what our eye sees, magnified as the bar of color directly above.

Wavelength is measured from the crest (or trough) of one wave to the crest (or trough) of another. The number of waves that pass a given point at any time is called frequency. The relationship between frequency and wavelength is inverse: the shorter the wavelength, the higher the frequency.

When eMag radiation is passed through atoms, some of the radiation is absorbed and used to excite the atoms from low energy levels to high energy levels. The mass spectrometer analyses the transmitted radiation relative to the incident radiation and an absorption spectrum is produced. Absorption spectra show colors with black lines and the black lines represent where the energy is absorbed.

Emission spectra result when a high voltage is applied to the gas. Emission spectra show black with a few colored lines. The colored lines are the same as those that are missing from the absorption spectrum. They read like bar codes to identify unknown elements.

Mr. Anderson is on the case. SL can stop watching around 3 minutes in, but HL must watch the entire video.

Let's play around with average atomic mass. Download this handout. We will have to make adjustments because clearly, we don't have M&Ms, Skittles or Reese's Pieces.

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G11 IB Chemistry Mad!Lab!

August 24 - 31, 2017

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