Lesson 2 – Atomic Theory


Objective:

  • to show how evidence is used to modify theories in the scientific method

Timeline:

  • 1/2 class

Background Information:

As far back as the Greeks, we have been wondering what makes up the world around us. How far could we divide the stuff we can touch, feel, measure and weigh – matter?  Was there a point where matter became indivisible? Or, could we keep dividing it again and again, never reaching the limit? These questions started scientists and philosophers on the quest to find the nature of matter. Along the way, they developed the atomic theory.

Atomic Theory

In the early 1800’s, scientist John Dalton started experimenting with gases. His experiments led him to believe that everything was made from small balls of matter, called atoms. These atoms could not be broken down further.  He theorized that elements were made of only one type of atom and that all atoms of an element weighed the exact same amount. Compounds were formed when elements mixed, and always formed in specific ratios.

The Dalton (Billiard Ball) Model

Dalton's Model

However, in science, a theory is only as good as the evidence that backs it up. Only a few years later another scientist, J.J. Thomson, discovered that atoms could be broken down further. He used a cathode ray tube. These are the tubes that were used to create pictures in old TV’s. He found that the cathode ray (the beam of particles) could be bent by a charged object. The ray moved towards positively charged things, and away form negatively charged things. This meant, the beam was made of negatively charged particles. He called these electrons. He decided that atoms were a sphere of positive mass, like a pudding, with negative electrons stuck in. The electrons could be dislodged, like in a cathode ray tube.

The Thomson (Raisin Bun) Model

Thomson's Model

Only a few years later, another scientist found new evidence about the structure of an atom. Ernest Rutherford was experimenting with a new discovery – that of radioactive materials. Rutherford and his students took a sheet of very thin gold, and fired particles at it. These particles, called alpha particles, are ejected from radioactive substances. They are positively charged, and don’t have much mass. Imagine: throwing a tennis ball against a wall – you would expect it to bounce off the wall all the time. Rutherford expected that the alpha particles would bounce back off the gold sheet in the same manner. However, most passed straight through the sheet! That was extremely weird and unexpected. A few alpha particles did bounce back, but were deflected at odd angles.

He concluded from these results:

1. atoms must be mostly space. Think about that – the chair you’re sitting on, your desk, your cup of coffee – is mostly empty space.

2. There was a strong positive lump in the middle of an atom (think of the alpha particles being reflected, positives repulse positives).

The Rutherford (Atomic) Model

This is the model we most commonly use today. There is a nucleus (the central ball) made up of positive protons.

Circling around the nucleus are negative electrons.

In an atom, the number of protons = the number of electrons. This means, the positive and negative charges are balanced. Atoms are neutral.

Rutherford and his students noticed something else odd about the atoms they were studying. According to Dalton, atoms of the same type (ex, all oxygen atoms) should have the same mass. However, this was not true. Some atoms were heavier than others. Rutherford’s student proposed that there was a third particle present, a neutron. Neutrons wouldn’t affect the atom in any way, except give it more mass. We call these isotopes.

Ex. Helium

Helium Atom Helium has 2 protons, so we give it a mass of 2.

Since atoms are neutral, it also has to have 2 electrons. Electrons are so small though, that we don’t count them in the mass.

Helium should have a mass of 2.

When we measure it though, it turns out there is a mass of 4. There are 2 neutrons in the nucleus that give Helium more mass!

Activity: Practice drawing atoms and using nuclear notation.

Use a periodic table to help you with this activity. Draw out the first 10 atoms on the periodic table (Hydrogen to Neon). But first, where do you find the information you need?

  • The atomic number (the non-decimal number in the top corner of a periodic table) tells you how many protons an atom has. It also tells you how many electrons it has (remember, atoms are neutral).
  • The mass number (the decimal number) is the sum of all the protons and neutrons. Ex. In our helium atom, 2 protons + 2 neutrons = a mass of 4.
  • To find the number of neutrons, subtract! Mass – protons = neutrons
  • Electrons exist in orbits. Only a certain amount of electrons can be in 1 orbit. To find out how many electrons can fit in an orbit, count across the row of the periodic table. The number of elements = the number of electrons that can exist there. Each time you go down a row, we say there is a new energy level. Ex. The first row on the periodic table has only 2 elements, hydrogen and helium, that means only 2 electrons can stay on the first level. The second row has 8 elements, which means 8 electrons can fit on the second level.

Ex. Sodium

Sodium dot diagram

Sodium is on the 3rd level down, so it has 3 energy levels of electrons.

2 are on the first level, 8 on the second, and 1 on the third to equal 11.

In the middle, I drew a nucleus, and said that P (protons) = 11 and N (neutrons) = 12.

This is called a Lewis dot diagram.

Notice that I put the electrons in pairs, this is mostly to keep track of how many I have. In reality, those electrons would not want to stay together!

Now: Try to draw diagrams for the first ten elements. Remember, Helium has been done for you already, above!

Self Check:

1. Describe a pattern you see in the arrangement of elements in the periodic table. How could this help you remember atomic structure?

2. Complete the following table:

Element Number of Protons Number of Electrons Number of Neutrons
Helium-3 2 2 1
Carbon-12
Cobalt-65
Sulfur-40
Iodine-140

3. Many elements have naturally occurring isotopes. For example, carbon can exist as carbon-12,  carbon-13 and carbon-14. Is carbon still a pure substance? Why or why not?

4. Describe, in paragraph form, how the evidence from Rutherford’s experiment disproved Dalton’s and Thomson’s atomic models. Be specific about the subatomic particles.

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2 Responses to “Lesson 2 – Atomic Theory”

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