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# 4.2: The Nuclear Model of the Atom

Difficulty Level: At Grade Created by: CK-12

## Lesson Objectives

• Distinguish between the three main subatomic particles.
• Understand the contributions of J. J. Thomson, Robert Millikan, and Ernest Rutherford to atomic theory.
• Describe the structure of the nuclear atom.

## Lesson Vocabulary

• atomic model
• cathode ray
• cathode ray tube
• electron
• neutron
• nucleus
• proton

## Introduction

Dalton’s atomic theory represented an improvement over the idea of Democritus because the theory was based on experimental findings and the scientific method. However, his theory did have its shortcomings. He believed that atoms were indivisible, meaning that the atom was the smallest possible component of matter. Further investigations in the late 1800s proved that atoms can indeed be broken down into smaller particles. It is the unique number and arrangement of these subatomic particles that makes atoms of one element different from those of every other element. The three fundamental particles are called the proton, the neutron, and the electron.

## Discovery of the Electron

The first discovery of a subatomic particle resulted from experiments into the nature of the relationship between electricity and matter.

### Cathode Rays

In 1897, English physicist J. J. Thomson (1856-1940) experimented with a device called a cathode ray tube, in which an electric current was passed through gases at low pressure. A cathode ray tube (Figure below) consists of a sealed glass tube fitted at both ends with metal disks called electrodes. The electrodes are then connected to a source of electricity. One electrode, called the anode, becomes positively charged while the other electrode, called the cathode, becomes negatively charged. Once this happens, a beam called a cathode ray travels from the cathode to the anode.

To produce a cathode ray, a glass tube filled with a low-pressure gas is connected to a power source. The green beam, or cathode ray, moves from the cathode to the anode.

Investigations were carried out to determine the nature of the cathode ray. The results of two further experiments supported the hypothesis that the cathode ray consisted of a stream of particles.

1. When an object was placed between the cathode and the opposite end of the tube, it cast a shadow on the glass.
2. A cathode ray tube was constructed with a small metal rail between the two electrodes. Attached to the rail was a paddle wheel capable of rotating along the rail. Upon connecting the cathode ray tube to a power source, the wheel rotated from the cathode towards the anode. This suggested that the cathode ray was made of particles that must have mass.

In order to determine if the cathode ray consisted of charged particles, Thomson used magnets and charged plates to deflect the ray (Figure below). His findings are summarized below.

1. Cathode rays were deflected by a magnetic field in the same manner as a wire carrying an electric current, which was known to be negatively charged.
2. Cathode rays were deflected away from a negatively charged metal plate and towards a positively charged plate.

The deflection of a cathode ray by a magnet. From the direction and extent of the deflection, Thomson was able to determine the charge-to-mass ratio of the electron.

Thomson knew that opposite charges attract one another, while like charges repel one another. Together, the results of the cathode ray tube experiments showed that cathode rays are actually streams of tiny negatively charged particles moving at very high speeds. While Thomson originally called these particles corpuscles, they were later named electrons.

Thomson conducted further experiments which allowed him to calculate the charge-to-mass ratio (e/me) of the electron. In units of coulombs to grams, e/me = 1.8 × 108 C/g. He found that this value was a constant and did not depend on the gas used in the cathode ray tube or on the metal used as the electrodes. He concluded that electrons were negatively charged subatomic particles present in atoms of all elements.

Watch a video of a cathode ray tube experiment at www.dlt.ncssm.edu/core/Chapter3-Atomic_Str_Part1/cathode-rm-lg.htm.

### Charge and Mass of the Electron

American physicist Robert Millikan (1868-1953) carried out a series of experiments between 1908 and 1917 that allowed him to determine the charge of a single electron. Millikan’s experiment was called the oil drop experiment (Figure below).

Millikan’s oil drop experiment: Oil drops that are sprayed into the main chamber fall through a tiny hole into an electric field, after which they can be viewed through a microscope. This experiment allowed Millikan to determine the charge of the electron.

When tiny drops of oil were sprayed into a chamber, the oil drops picked up a static charge and were suspended between two charged plates. Millikan was able to observe the motion of the oil drops with a microscope and found that the drops lined up in a specific way between the plates, based on the number of electric charges that they had acquired. From the data gathered in this experiment, he was able to accurately determine the charge of an individual electron. Then, using Thomson’s previous measurement of an electron's charge-to-mass ratio, he was also able to calculate the mass of a single electron.

Charge of one electron=1.602×1019 CMass of one electron=9.11×1028 g\begin{align*} & \text{Charge of one electron} = -1.602 \times 10^{-19} \ \text{C} \\ & \text{Mass of one electron} = 9.11 \times 10^{-28} \ \text{g} \end{align*}

The incredibly small mass of the electron was found to be approximately 1/1840 the mass of a hydrogen atom, so scientists realized that atoms must also contain other, far more massive particles. Additionally, at least one of these particles must carry a positive charge, because complete atoms are electrically neutral.

Participate in a simulation of Millikan's oil drop experiment at this site: http://www.dlt.ncssm.edu/core/Chapter3-Atomic_Str_Part1/Chapter3-Animations/OilDrop.htm.

### Protons and Neutrons

If cathode rays are electrons that are given off by the metal atoms of the cathode, then what remains of the atoms that have lost those electrons? We know several basic things about electrical charges. They are carried by particles of matter. Millikan’s experiment showed that they exist as whole-number multiples of a single basic unit. Atoms have no overall electrical charge, meaning that each and every atom contains an exactly equal number of positively and negatively charged particles. A hydrogen atom, the simplest kind of atom, contains only one electron. When that electron is removed, a positively charged particle should remain.

In 1886, Eugene Goldstein (1850-1930) discovered evidence for the existence of this positively charged particle. Using a cathode ray tube with holes in the cathode, he noticed that there were rays traveling in the opposite direction from the cathode rays. He called these canal rays and showed that they were composed of positively charged particles. The proton is a positively charged subatomic particle that is present in all atoms. The mass of the proton is about 1840 times the mass of the electron.

In 1932, English physicist James Chadwick (1891-1974) discovered a third subatomic particle. The neutron is a subatomic particle with no electrical charge and a mass that is approximately the same as the mass of a proton. Table below summarizes the properties of the three fundamental subatomic particles.

Properties of Subatomic Particles
Particle Symbol Relative Electrical Charge Relative Mass (amu) Actual Mass (g)
Electron e- 1- 1/1840 9.11 × 10-28
Proton p+ 1+ 1 1.67 × 10-24
Neutron n0 0 1 1.67 × 10-24

1 amu (atomic mass unit) = 1.66 × 10-24 g

## Discovery of the Atomic Nucleus

The next step after the discovery of subatomic particles was to figure out how these particles were arranged in the atom. This is a difficult task because of the incredibly small size of the atom. Therefore, scientists set out to design a model of what they believed the atom could look like. The goal of each atomic model was to accurately represent all of the experimental evidence about atoms in the simplest way possible. Following the discovery of the electron, J. J. Thomson developed what became known as the “plum pudding” model (Figure below). In this model, the electrons were suspended in a uniform lump of positive charge like blueberries in a muffin. This model of the atom soon gave way, however, to a new model developed by New Zealander Ernest Rutherford (1871-1937).

In Thomson’s plum pudding model of the atom, electrons are embedded in a uniform sphere of positive charge. Plum pudding is an English dessert that is similar to a blueberry muffin.

In 1911, Rutherford and coworkers Hans Geiger and Ernest Marsden initiated a series of groundbreaking experiments that would completely change the accepted model of the atom. The experimental setup is shown in Figure below. When they bombarded very thin sheets of gold foil with fast moving alpha particles, they got some unexpected results. An alpha particle is a type of positively charged particle whose mass is about four times that of a hydrogen atom. It occurs naturally as a product of radioactive decay.

(A) The experimental setup for Rutherford’s gold foil experiment: A radioactive element that emitted alpha particles was directed toward a thin sheet of gold foil, which was surrounded by a screen that would allow detection of the deflected particles. (B) According to the plum pudding model (top) all of the alpha particles should have passed through the gold foil with little or no deflection. Rutherford found that a small percentage of alpha particles were deflected at large angles, which could be better explained by an atom that contained a very small, dense, positively-charged nucleus (bottom).

According to the accepted atomic model, in which an atom’s mass and charge are uniformly distributed throughout the atom, the scientists expected that all of the alpha particles would pass through the gold foil with only a slight deflection or none at all. Surprisingly, while most of the alpha particles were indeed undeflected, a very small percentage (about 1 in 8000 particles) bounced off the gold foil at very large angles. Some were even redirected back toward the source. Nothing had prepared them for this discovery. In a famous quotation, Rutherford exclaimed that it was “as if you had fired a 15-inch [artillery] shell at a piece of tissue paper and it came back and hit you.”

Rutherford needed to come up with an entirely new model of the atom in order to explain his results. Because the vast majority of the alpha particles had passed through the gold, he reasoned that most of the atom was empty space. In contrast, the particles that were highly deflected must have experienced a tremendously powerful force within the atom. He concluded that all of the positive charge and the majority of the mass of the atom must be concentrated in a very small space in the atom’s interior, which he called the nucleus. The nucleus is the tiny, dense, central core of the atom and is composed of protons and neutrons.

Rutherford’s atomic model became known as the nuclear model. In this model, the protons and neutrons, which comprise nearly all of the mass of the atom, are located in a nucleus at the center of the atom. The electrons are distributed around the nucleus and occupy most of the volume of the atom. It is worth emphasizing just how small the nucleus is compared to the rest of the atom. If we could blow up an atom to be the size of a large professional football stadium, the nucleus would be about the size of a marble.

Rutherford’s model proved to be an important step towards a full understanding of the atom. However, it did not completely address the nature of the electrons and the way in which they occupied the vast space around the nucleus. It was not until some years later that a more complete understanding of the electron was achieved. This proved to be the key to understanding the chemical properties of elements.

Watch a simulation of Rutherford's experiment through The Concord Consortium's Molecular Workbench. You will need to download the Molecular Workbench application from http://mw.concord.org/modeler/index.html. After installing Molecular Workbench, open it. You will see a browser window. In the address bar, type: "http://mw2.concord.org/public/student/motionandforce/rutherford.cml." You will be taken to the simulation.

## Lesson Summary

• The three fundamental subatomic particles are the electron, the proton, and the neutron.
• Thomson used the cathode ray tube to discover the electron and determine its negative charge.
• Millikan determined the charge and mass of the electron with the oil-drop experiment.
• Rutherford’s gold foil experiment provided evidence for the atomic nucleus, a small dense core of the atom which contains the positive charge and most of the mass.
• The nuclear model of the atom is one in which the nucleus is composed of protons and neutrons, while electrons are distributed throughout the rest of the space.

## Lesson Review Questions

### Recall

1. What evidence did Thomson have for each statement below?
1. Electrons are negatively charged.
2. Electrons are identical and are present in all atoms.
3. Compare the mass (in amu) and relative charge of a neutron to that of a proton.
4. What evidence did Rutherford have for each statement below?
1. Atoms are mostly empty space.
2. The nucleus of the atom is positively charged.

### Apply Concepts

1. Trace the development of the atomic model from Dalton to Thomson to Rutherford. Explain what experimental findings led each scientist to alter the previous model and how his new model fit with the new evidence.

### Think Critically

1. A hydrogen nucleus, which contains only a single proton, has a diameter of 3.7 × 10−10 cm. The equation for the volume of a sphere is 43πr3\begin{align*}\frac{4}{3} \pi r^3\end{align*}, where r is the radius of the sphere.
1. Assuming that the nucleus is spherical, calculate its volume in cm3.
2. Calculate the density of the hydrogen nucleus in g/cm3.
3. The densest element known is osmium, which has a density of 22.6 g/cm3. Comment on the difference in density between osmium and a hydrogen nucleus.
2. How could Rutherford’s experiment be modified in order to determine the relative sizes of different nuclei?
3. Why was it not possible to detect the existence of the neutron with experiments analogous to those used in the discovery of the electron?
4. All matter is composed of atoms, which are in turn composed mostly of empty space. Why is it not possible to walk through a wall or to put your hand right through your desk?

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