Welcome to MPCC Chemistry

Atoms - the Building Blocks of Matter

Atomic Models - How They’ve Changed Over Time

The Subatomic Parts of the Atom and the Periodic Table

Subatomic Parts

The Periodic

Atomic Quantities

Average Atomic Mass

Electron Configuration

The Shell Model

The Quantum Model

Welcome to MPCC Chemistry

What is Chemistry?

To get an overview of what you will learn about in this class, watch this short video clip: Introduction to Chemistry Video Clip^[1]. Chemistry is the study of the composition of matter.

For the purposes of this class, here is a copy of a Periodic Table from Sciencegeek.net for you to reference.

Atoms - the Building Blocks of Matter

Atomic Models - How They’ve Changed Over Time

Learning Target 1.1.1: I can explain how the model of the atom has changed over time into our current model.

Goal 1.1.1.a I can describe the contributions of Dalton, Thomson, Rutherford, Chadwick and Bohr toward our current model of the atom.

Goal 1.1.1.b I can describe how each scientist’s contribution was built on earlier knowledge.

The model of the atom has changed a lot over the past 2400 years. The following text, links, and videos will give you a brief overview of the changes that have taken place. Your teacher may assign you one of the following scientists on which to focus your research: Dalton, Thomson, Rutherford, Chadwick, or Bohr. As you follow the links to readings and videos, complete these questions and be ready to report out to your class on your findings.

Matter is anything that has mass and takes up space. In about 400 B.C., a Greek philosopher named Democritus began to think about what makes up matter. He proposed that matter could be cut into smaller and smaller pieces until it could not be any further divided. He called this smallest piece of matter atomos (which means not cuttable). At the time, no one experimented to find out if Democritus’s ideas about these indivisible pieces was correct or not. About 200 years ago, a scientist named John Dalton^[2] did some experiments with Democritus’s idea of atomos and came up with a theory of what he thought atoms were like and how they acted and interacted. Read more about this: John Dalton's Atomic Theory text^[3]

After about another 100 years of experimentation by the developing scientific community, scientists began to find flaws with Dalton’s theories. J.J. Thomson used a cathode ray tube^[4] (Cathode Ray Tube Demo video^[5]) to discover that there were charged parts within the atom with his “plum pudding model” (see Thomson model video^[6] or Thomson model text^[7]); the negatively charged parts are known today as electrons. A more precise measurement of the mass and charge of an electron were done by Robert Millikan several years later (see Millikan Oil Drop Experiment video^[8] or Millikan Oil-Drop Experiment text^[9]). Ernest Rutherford and his associates, went on to discover that the positively charged parts of the atom were contained in a very tiny, very dense nucleus (see Gold Foil Experiment video^[10] or Rutherford’s Atomic Model text^[11]); these positively charged parts are known today as protons. With the knowledge of the nucleus, Niels Bohr modified the atomic model to include the possible locations of the electrons around the nucleus into distinct energy levels or quanta, similar to planets rotating around the sun ^[12] (see Bohr Model flash file (stop at Multi-electron atoms)^[13]or Bohr Model text^[14]). Later, James Chadwick did experiments to prove that there was another particle in the nucleus about the same mass as a proton, but without any charge (see Atomic Structure: Discovery of the Neutron video^[15] or Chadwick discovers neutron text^[16]); these neutral particles are known today as neutrons. The modern day model of the atom is called the Quantum Mechanical Model.

For an interactive experience to see results of experiments under the different models of the atom, see a PhET Simulation: Models of the Hydrogen Atom^[17] 

You can complete this Hydrogen PhET activity worksheet^[18] to compare what you observe in the simulation to the original experiment, or ask your teacher for a hands-on experience of a model of Rutherford’s experiment.

The Subatomic Parts of the Atom and the Periodic Table

Learning Target 1.1.2: I can describe the parts of an atom and how to use the periodic table to determine the numbers of each part for different elements.

Subatomic Parts

Goal: 1.1.2.a I can describe the mass, charge, and location of protons, neutrons and electrons relative to each other

Our modern understanding of the parts of an atom is summarized in Table 1.1 below. Electrons can not be broken down into smaller pieces as far as we know, but protons and neutrons are made of smaller pieces called quarks (for more about quarks: Particle Adventure website^[19], Standard Model video^[20], or Standard Model text^[21]) . For the purposes of explaining how the particles in the table are related to the periodic table of elements, we will not break protons and neutrons into their smaller pieces.

Table 1.1

For an interactive activity to check your understanding of the properties of protons, neutrons, and electrons use this pne summary SMART Notebook file or use this Properties of pne Sorting Activity printable version.

The Periodic Table

Goal 1.1.2.b I can explain how things are arranged in the periodic table of elements.

Around 1870, before these subatomic particles were discovered, Dmitri Mendeleev started to look for patterns in how the known elements behaved. Scientists at the time knew of about 60 unique elements and their atomic weights compared to hydrogen. He put the elements in order of increasing atomic weights and found that there was a repetition in how the elements acted and reacted with other elements. He grouped these elements with similar properties together and found patterns that repeated. He was also able to predict that there would be new elements discovered to fit into the places where there were gaps in his table. Because it was based on patterns that periodically repeated, this became the first version of what we know today as the Periodic Table of Elements.

“I began to look about and write down the elements with their atomic weights and typical properties, analogous elements and like atomic weights on separate cards, and this soon convinced me that the properties of elements are in periodic dependence upon their atomic weights.”

--Mendeleev, Principles of Chemistry, 1905, Vol. II ^[22]

Complete this Stick Person Challenge activity using these Stick People Cards^[23] to get an idea of the process that Mendeleev used to organize his table.

As the table continued to develop, there were some spots where the patterns did not always line up when they were in order of increasing weight. In 1913, Henry Moseley used X-rays to examine atoms of different elements and apply what Rutherford and Chadwick had discovered about the nucleus. He used the results of his experiments to revise Mendeleev’s table to be arranged in order of increasing number of protons instead of by weight.

Complete this Grouping Elements into Families activity.

To learn more about the properties of different elements, you can visit the Periodic Table of Videos from the University of Nottingham^[24].

The 18 columns of the table are now called families or groups and the 7 rows of the table are called periods. The first 13 groups of the table are good conductors of heat and electricity and the elements in this part of the table are known as metals. Group 1 is a set of very reactive metals called the alkali metals. Group 2 is also a set of reactive metals called the alkaline Earth metals. Groups 3-12 are called Transition Metals. There is a stair-step line that includes boron, silicon, germanium, arsenic, antimony, and tellurium that are known as metalloids. To the right of the metalloids, the elements are poor conductors of heat and electricity and elements in this part of the table are known as nonmetals. Group 17 is a set of very reactive nonmetals called the Halogens. Group 18 is a set of unreactive nonmetals called the Noble Gases. The reasons for the various regions of the tables behaving as they do will be discussed in the next lesson.

Atomic Quantities

Goal: 1.1.2.c I can use the periodic table to calculate the protons, neutrons, and electrons in an atom of an element.

Goal: 1.1.2.d I can explain how isotopes and ions affect the numbers of neutrons and electrons that atoms of a given element may have may have.

Atomic Number: Each element is defined by its unique number of protons. Every atom of a particular element will have exactly the same number of protons. The atomic number is equal to the number of protons in the nucleus of an atom. Therefore, each element has a unique atomic number. The atomic number is sometimes referred to as “Z”.The atomic number is found with the name and/or symbol of the element on the Periodic Table of Elements.

Mass Number: Equal to the mass of a single atom of an element.When the mass of a single atom is calculated, each proton and neutron contribute a mass of 1 atomic mass unit (amu). The mass of the electrons is so small in comparison to the protons and neutrons that they are not included in the mass number. The mass number, therefore, is equal to the number of the protons and neutrons in the nucleus of an atom. When atoms of the same element have different mass numbers, they are called isotopes of each other. The mass number is sometimes referred to as “A”. Some isotopes have unstable nuclei which will decay radioactively^[25]. A specific isotope can be identified in two different ways. One way is called the hyphen notation and has the name of the element followed by a hyphen and the mass number. The other way is called the nuclide notation and has the symbol of the element with the mass number superscripted in front of it and the atomic number subscripted in front of it.

Hyphen Notation         and          Nuclide Notation

Not all isotopes are radioactive. Whether an isotope is radioactive is related to the number of neutrons it has compared to its protons. When an isotope is radioactive, half of a sample will radioactively decay into a different atom in a certain amount of time. The time for half of a sample to radioactively decay is called that isotope’s half-life. As each half life of time passes, another half of the sample radioactively decays. For a model of how the half-life of radioactive isotopes, you can try this Calculating the Half-Life of Twizzlers and M&Ms^[26] lab activity

For an interactive activity to process the difference between Atomic Number and Mass Number: use this Atomic Number vs Mass Number SMART Notebook file or use this Atomic Number vs Mass Number Sorting Activity printable version.

Charge of an Atom: Individual atoms of a pure element are neutral in charge. This means that if an atom is composed of positively charged protons and negatively charged electrons, the numbers of protons must be equal to the number of electrons for the charges to cancel to zero. Most of the time, the way the electrons are arranged in an atom of a pure element is not very stable and atoms may give electrons to or take electrons from another atom to improve the stability of the arrangement. When this happens, the atom becomes charged and the resulting particle is called an ion. If an ion has more electrons than protons, it is negatively charged and if an ion has more protons than electrons, it is positively charged. See Calculating the Charge of an Atom video^[27]. In your notes, write down the two types of particles you need to know the numbers of to calculate the charge of an atom and how you decide if an atom is positively charged, negatively charged, or neutral.

For an interactive activity to practice using the information on the periodic table to identify the numbers of protons, neutrons, and electrons, you can click on this HTML 5 version of PhET Atom Builder^[28]

Practice working with Atomic Number, Mass Number and the number of protons, neutrons, & electrons in atoms of various element: Calculating PNE Worksheet

Practice working with Isotope Notations and Ions: Nuclear Atom Worksheet^[29]

Average Atomic Mass

Goal 1.1.2.e I can calculate the average atomic mass of an element as a weighted average of the isotopes.

Because there are atoms of the same element (same atomic number) with different numbers of neutrons (different mass number), the average atomic mass will be different than the mass of a single atom of an element. Using a mass spectrometer^[30] scientists are able to figure out the natural abundance of all of the different isotopes that exist in a sample of an element and then calculate a weighted average. This weighted average is recorded on the periodic table. When you do calculations involving large numbers of atoms of an element in the next unit, you will use the average atomic mass rather than the mass of a single atom (mass number).

For additional information: Average Atomic Mass text^[31] 

For an interactive simulation to learn more about this: Isotopes and Atomic Mass PhET Simulation^[32]

Complete this activity sheet^[33] while you do the simulation

Complete this hands-on Isotopes of M&Mium lab activity to help understand how a mass spectrometer works to help scientists calculate the average atomic mass of an element.

Watch Atomic Mass: Introduction video^[34] and How to Calculate Atomic Mass Practice Problems video^[35]. As you watch the video, write down the pieces of information that you will need to calculate a weighted average mass.

Ask your teacher for a practice sheet on Average Atomic Mass.

For a video summary of what is discussed in this section, see Atoms and the Periodic Table video^[36]. While you watch the video, write down any questions you still have about atoms and the periodic table.

Electron Configuration

Learning Target 1.1.3: I can explain how electrons are arranged around the nucleus.

Goal: 1.1.3.a I can explain how electrons are divided into different energy levels

The Shell Model

In the Bohr model, electrons are arranged in distinct energy levels orbiting the nucleus at fixed distances. This is sometimes called the Shell Model. Each energy level holds a maximum number of electrons. Bohr formulated this model from data collected from the quantized emission of photons of specific energies & frequencies (seen as different colors of light) from the hydrogen atom. (See Bohr Model flash file (stop at Multi-electron atoms))^[37]While this model is not quite right, it is a good way to start looking at the arrangement of electrons into energy levels around the nucleus.

In the shell model, the first level of electrons can hold 2 electrons, the second level can hold 8 electrons, the third level can hold 18, and the fourth level can hold 32. However, the third level and beyond do not completely fill before moving to the next level.

Here is a diagram showing how the first 36 electrons fill:

To practice this, use this Drawing Models of Atoms Worksheet

Goal 1.1.3.b I can determine the number of valence electrons for each element in the main block groups of the periodic table.

Looking at your diagrams, you can easily see how many electrons each model has in its highest energy level. These electrons in the shell furthest from the nucleus are called valence electrons. The electrons in the inner energy levels are called core electrons.

Compare your drawings for F and Cl. How many valence electrons in each? What group are they in?

Compare your drawings for H and Na. How many valence electrons in each?

Now draw Li and K. How do those drawings compare to H and Na? What group are all of these in?

Here are drawings for B and Al:

How many valence electrons in each? What group are these in? Can you see any relationship between the number of valence electrons and the group number?

To learn more about this, watch Valence Electrons and the Periodic Table video^[38]. As you watch this video, write down two things you learned and one question you still have.

When you look at the elements in the same column or group, they have the same number of valence electrons. There are 18 columns on the periodic table and they are numbered as Groups 1-18. However, columns 1, 2, and 13-18 are called the main block groups and can follow an alternative numbering system of 1A, 2A, 3A, 4A, 5A, 6A, 7A, and 8A. In this numbering system, the family number matches the number of valence electrons. See the diagram below for the trend.

This model also helps us easily draw Dot Diagrams or Lewis Dot Structures. A Lewis Dot Structure is the symbol for the element surrounded by dots to represent the valence electrons of a particular element, based on its’ family. A further discussion of Lewis Dot Diagrams and practice drawing them is in Unit 2.

Some of the families or groups of the periodic table have special names. Group 1 is called the Alkali metals because when they are added to water, they make an alkaline (basic) solution. Group 2 is called the Alkaline Earth metals. Group 17 or 7A is called the Halogen family. Group 18 or 8A is called the Noble Gas family because it is unreactive with most other elements.

Goal 1.1.3.c I can explain how electrons are divided into different types of orbitals & suborbitals.

The Quantum Model

The current model of the atom is built on the idea that the electrons have locations within an area called the electron cloud which surrounds the nucleus. Although he was right that the area around the nucleus holds electrons of different energies, Bohr’s idea that each electron had a specific orbit like planets going around the sun wasn’t quite right. Erwin Schrödinger developed a mathematical model that showed that electrons did not move in a specific orbit, but in regions where the electrons of a specific energy and orbit are most likely to be found (electron cloud text^[39]).

How scientists convert a cloud of electron density into a balloon (CK-12 Foundation)^[40]

To extend your learning, watch Quantum Mechanical Model video ^[41] for more information or link to Quantum Mechanical Model text^[42]

 Experimental data has shown that when the density of an electron cloud is imaged, most of the electrons are found to be within various predictable regions. ^[43] The different shaped regions are thought to be associated with electrons in different orbitals. There are thought to be seven different energy levels and four different types of orbitals in the different energy levels. As the energy of the electrons increase, they are able to travel further away from the positively charged nucleus. There are also more possible orbital shapes as the energy levels increase. Most of volume of an atom is caused by its electron cloud, which is mostly empty space. An analogy that is often used to relate the sizes of the nucleus and the atomic diameter (which is caused by the electrons) is imagine the nucleus is the size of a marble, the outer edge of the electron cloud would be as big as a football stadium.

The first energy level has only one type of orbital, s. The second energy level has s and p orbitals. The third level has s, p and d orbitals. and the fourth level has s, p,d, and f orbitals. If there were bigger elements, there could be other types of orbitals theoretically.

s Orbitals^[44]

This orbital is spherical in shape and can hold one pair of electrons:

p Orbitals

There are three sub-orbitals designated as p orbitals and they have dumbbell shapes. Each of the p orbitals has a different orientation in three-dimensional space. Each sub-orbital can hold up to two electrons.

d Orbitals

There are a total of five sub-orbitals of d orbitals . Note that all five of the orbitals have specific three-dimensional orientations. Again, each sub-orbital can hold a maximum of two electrons.

f Orbitals

The most complex set of orbitals are the f orbitals . There a total of seven different sub-orbital shapes. Again, note the specific orientations of the different f sub-orbitals which each hold a maximum of two electrons.

There are three rules that govern how the electrons fill into these orbitals and sub-orbitals. These are the Aufbau principle, the Pauli Exclusion principle and Hund’s rule. The Aufbau principle^[45] says that electrons will fill the lowest energy orbitals first before moving to a higher energy orbital. The Pauli Exclusion principle^[46] says that when two electrons are in the same sub-orbital, they can not have the same magnetic spin. And Hund’s rule^[47] says that electrons will fill one electron into each sub-orbital level first before adding a second electron to the sub-orbital level.

Try this analogy to electron configuration: Hog Hilton^[48]

See Electron Configuration video^[49] summary. Make note of any questions you have after watching the video.

Practice writing electron configurations online^[50]

See how well you understand this concept with Electron Configuration and the Periodic Table Practice^[51]