Atomic Structure Study Guide

CHAPTER 1- Atom Basics

Strand 1: The Structure and Properties of Atoms

Chapter Outline

• Atomic Structure (Chem.1.1)
• Isotopes and Decay (Chem.1.2)
• Half-life (Chem.1.3)
• Nuclear Reactions (Chem.1.4)
• Periodic Table (Chem.1.5)

 

Atoms have substructures of their own including a small central nucleus containing protons and neutrons surrounded by a larger region containing electrons. The strong nuclear interaction provides the primary force that holds nuclei together. Without it, the electromagnetic forces between protons would make all nuclei other than hydrogen unstable.

Processes of fusion, fission, and radioactive decay of unstable nuclei involve changes in nuclear binding energies. Elements are placed in columns and rows on the periodic table to reflect their common and repeating properties.

 

 

1.1 Atomic Structure (Chem.1.1)

Explore this Phenomenon

The images are different models of the atom throughout history. Look for patterns in the structure and function of these models of atoms. Then answer the following questions:

1. Why are there differences between the models?
2. How are the models similar?
3. Can you predict the order these models were created?

 

Standard Chem.1.1

Obtain, evaluate, and communicate information regarding the structure of the atom on the basis of experimental evidence. Emphasize the relationship between proton number and element identity, isotopes, and electrons in atoms. Examples of experimental evidence could include the gold foil experiment, cathode ray tube, or atomic spectrum data. (PS1.A)

Throughout history the model of the atom has changed. As you read this section, look for how different evidence is used to draw conclusions about the structure of atoms.

 

What Are Atoms?

Atoms are the building blocks of matter. They are the smallest particles of an element that still have the element's properties. Elements, in turn, are pure substances—such as nickel, hydrogen, and helium—that make up all kinds of matter. All the atoms of a given element are identical in that they have the same number of protons, one of the building blocks of atoms. They are also different from the atoms of all other elements, as atoms of different elements have a different number of protons.

 

Size of Atoms

Unlike bricks, atoms are extremely small. The radius of an atom is well under 1 nanometer, which is one-billionth of a meter. If a size that small is hard to imagine, consider this: trillions of atoms would fit inside the period at the end of this sentence.

 

Subatomic Particles

Although atoms are very tiny, they consist of even smaller particles. Three main types of particles that make up all atoms are:

• protons, which have a positive electric charge.
• electrons, which have a negative electric charge.
• neutrons, which are neutral in electric charge.

The model below shows how these particles are arranged in an atom. The particular atom represented by the model is helium, but the particles of all atoms are arranged in the same way. At the center of the atom is a dense area called the nucleus , where all the protons and neutrons are clustered closely together. The electrons constantly move around the nucleus. Helium has two protons and two neutrons in its nucleus and two electrons moving around the nucleus. Atoms of other elements have different numbers of subatomic particles, but the number of protons always equals the number of electrons. This makes atoms neutral in charge because the positive and negative charges "cancel out."

 Model of a Helium Atom     

Early Ideas of Atoms

All matter in the universe is made of atoms (basic unit of matter). All modern scientists accept the concept of the atom, but when the concept of the atom was first proposed about 2,500 years ago, ancient philosophers laughed at the idea. It has always been difficult to convince people of the existence of things that are too small to see. We will spend some time considering the evidence (observations) that convince scientists of the existence of atoms.

 

Democritus and the Greek Philosophers

One of the first people to propose “atoms” was a man known as Democritus. As an alternative to the beliefs of many Greek philosophers, he suggested that atomos, or atomon—tiny, indivisible, solid objects - make up all matter in the universe. Sadly, it took over two millennia before the theory of atomos (or “atoms”, as they’re known today) was fully appreciated.

While it must be assumed that many more scientists, philosophers, and others studied the composition of matter after Democritus, a major leap forward in our understanding of the composition of matter took place in the 1800s with the work of the British scientist John Dalton.

Dalton studied the weights of various elements and compounds . He noticed that matter always combined in fixed mathematical ratios based on weight , or volume in the case of gases. Chemical compounds always contain the same proportion of elements by mass, regardless of amount. Dalton also observed that there could be more than one combination of two elements.

 

Dalton's Atomic Theory (1804)

From his experiments and observations, as well as the work from peers of his time, Dalton proposed a new theory of the atom known as Dalton's atomic theory. The general tenets of this theory were as follows:

• All matter is composed of extremely small particles called atoms.
• Atoms of a given element are identical in size, mass, and other properties. Atoms of different elements differ in size, mass, and other properties.
• Atoms cannot be subdivided, created, or destroyed.
• Atoms of different elements can combine in simple whole number ratios to form chemical compounds.
• In chemical reactions , atoms are combined, separated, or rearranged.

 

Thomson’s Plum Pudding Model

In the mid-1800s, scientists were beginning to realize that the study of chemistry and the study of electricity were actually related. First, a man named Michael Faraday showed how passing electricity through mixtures of different chemicals could cause chemical reactions. Shortly after that, the scientists found that by forcing electricity through a tube filled with gas, the electricity made the gas glow! Scientists didn’t understand the relationship between chemicals and electricity until a British physicist named J. J. Thomson began experimenting with what is known as a cathode ray tube.

Thomson’s experiment with cathode rays found that the ray moved away from negatively charged plates and toward positively charged plates. What does this say about the charge of the ray?

J.J. Thomson did some rather complex experiments with magnets, and used his results to prove that cathode rays were not only negatively charged, but also had mass. Remember that anything with mass is part of what we call matter. In other words, these cathode rays must be the result of negatively charged “matter” flowing from the cathode to the anode. But there was a problem. According to J. J. Thomson’s measurements, either these cathode rays had a ridiculously high charge, or else had very, very little mass – much less mass than the smallest known atom.

How was this possible? How could the matter making up cathode rays be smaller than an atom if atoms were indivisible? J. J. Thomson made a radical proposal: maybe atoms are divisible. J. J. Thomson suggested that the small, negatively charged particles making up the cathode ray were actually pieces of atoms. He called these pieces “corpuscles”, although today we know them as electrons. Thanks to his clever experiments and careful reasoning, J. J. Thomson is credited with the discovery of the electron.

Now imagine what would happen if atoms were made entirely of electrons. First of all, electrons are very, very small; in fact, electrons are about 2,000 times smaller than the smallest known atom, so every atom would have to contain a whole lot of electrons. But there’s another, even bigger problem: electrons are negatively charged. Therefore, if atoms were made entirely out of electrons, atoms would be negatively charged themselves and that would mean all matter was negatively charged as well. Of course, matter isn’t negatively charged. In fact, most matter is what we call neutral – it has no charge at all. If matter is composed of atoms, and atoms are composed of negative electrons, how can matter be neutral?

The only possible explanation is that atoms consist of more than just electrons. Atoms must also contain some type of positively charged material that balances the negative charge on the electrons. Negative and positive charges of equal size cancel each other out, just like negative and positive numbers of equal size. What do you get if you add +1 and -1? You get 0, or  nothing. That’s true of numbers, and that’s also true of charges. If an atom contains an electron with a -1 charge, but also some form of material with a +1 charge, overall the atom must have a(+1)+(−1)=0 charge – in other words, the atom must be neutral, or have no charge at all.

Thomson’s plum pudding model was much like a chocolate chip cookie. Notice how the chocolate chips represent the negatively charged electrons, while the positive charge is spread throughout the entire “batter”.

When Thomson discovered the negative electron, he realized that atoms had to contain positive material as well – otherwise they wouldn’t be neutral overall. As a result, Thomson formulated what’s known as the “plum pudding” model for the atom. According to the “plum pudding” model, the negative electrons were like pieces of fruit and the positive material was like the batter or pudding. This made a lot of sense given Thomson’s experiments and observations. Thomson had been able to isolate electrons using a cathode ray tube; however, he had never managed to isolate positive particles.

Rutherford’s Nuclear Model

Everything about Thomson’s experiments suggested the “plum pudding” model was correct – but according to the scientific method, any new theory or model should be tested by further experimentation and observation. In the case of the “plum pudding” model, it would take a man named Ernest Rutherford to prove it inaccurate.

Disproving Thomson’s “plum pudding” model began with the discovery that an element known as uranium emits positively charged particles called alpha particles as it undergoes radioactive decay. Radioactive decay occurs when one element decomposes into another element. It only happens with a few very unstable elements.

Ernest Rutherford was fascinated by all aspects of alpha particles. For the most part, though, he seemed to view alpha particles as tiny bullets that he could use to fire at all kinds of different materials. One experiment in particular, however, surprised Rutherford and everyone else.

Rutherford found that when he fired alpha particles at a very thin piece of gold foil, an interesting thing happened (A). Almost all of the alpha particles went straight through the foil as if they’d hit nothing at all. This was what he expected to happen. If Thomson’s model was accurate, there was nothing hard enough for these small particles to hit that would cause any change in their motion.

 

Every so often, though, one of the alpha particles would be deflected slightly as if it had bounced off of something hard. Even less often, Rutherford observed alpha particles bouncing straight back at the “gun” from which they had been fired! (B) It was as if these alpha particles had hit a wall “head-on” and had ricocheted right back in the direction that they had come from, indicating they were hitting a very small, very dense particle in the atom.

Rutherford thought that these experimental results were rather odd. Rutherford described firing alpha particles at gold foil like shooting a high-powered rifle at tissue paper. Would you ever expect the bullets to hit the tissue paper and bounce back at you? Of course not! The bullets would break through the tissue paper and keep on going, almost as if they’d hit nothing at all. Therefore, the fact that most alpha particles passed through didn’t shock him. On the other hand, how could he explain the alpha particles that got deflected? Furthermore, how could he explain the alpha particles that bounced right back as if they’d hit a wall?

Rutherford concluded that the only way to explain his results was to assume that the positive matter forming the gold atoms was not, in fact, distributed like the batter in plum pudding, but rather, was concentrated in one spot, forming a small positively charged particle somewhere in the center of the gold atom. We now call this clump of positively charged mass the nucleus - (the small, dense, positively charged center of the atoms). According to Rutherford, the presence of a nucleus explained his experiments, because it implied that most alpha particles passed through the gold foil without hitting anything at all. Once in a while, though, the alpha particles would actually collide with a gold nucleus, causing the alpha particles to be deflected, or even to bounce right back in the direction they came from.

While Rutherford’s discovery of the positively charged atomic nucleus offered insight into the structure of the atom, it also led to some questions. According to the “plum pudding” model, electrons were like plums embedded in the positive “batter” of the atom. Rutherford’s model, though, suggested that the positive charge wasn’t distributed like batter, but rather, was concentrated into a tiny particle at the center of the atom, while most of the rest of the atom was empty space. What did that mean for the electrons? If they weren’t embedded in the positive material, exactly what were they doing? And how were they held in the atom? Rutherford suggested that the electrons might be circling or “orbiting” the positively charged nucleus as some type of negatively charged cloud, but at the time, there wasn’t much evidence to suggest exactly how the electrons were held in the atom.

 

Rutherford suggested that electrons surround a central nucleus.

 

Despite the problems and questions associated with Rutherford’s experiments, his work with alpha particles definitely seemed to point to the existence of an atomic “nucleus.” Between J. J. Thomson, who discovered the electron, and Rutherford, who suggested that the positive charges in an atom were concentrated at the atom’s center, the 1890s and early 1900s saw huge steps in understanding the atom at the subatomic level. It was clear that an atom contains negatively charged electrons and a nucleus containing positive charges. In the next section, we’ll look more carefully at the structure of the nucleus, and we’ll learn that while the atom is made up of positive and negative particles, it also contains neutral particles that neither Thomson, nor Rutherford, were able to detect with their experiments.

 

The Quest for the Neutron

Clues are generally considered to involve the presence of something – a footprint, a piece of fabric, a bloodstain, something tangible that we can measure directly. The discoveries of the electron and the proton were accomplished with the help of those kinds of clues. Cathode ray tube experiments showed both the negatively charged electrons emitted by the cathode and a positively charged proton (also emitted by the cathode). The neutron was initially found not by a direct observation , but by noting what was not found.

Research has shown the properties of the electron and the proton . Scientists learned that approximately 1837 electrons weighed the same as one proton. There was evidence to suggest that electrons went around the heavy nucleus composed of protons. Charge was balanced with equal numbers of electrons and protons which make up an electrically neutral atom . But there was a problem with this model – the atomic number (number of protons) did not match the atomic weight . In fact, the atomic number was usually about half the atomic weight. This indicated that something else must be present. That something must weigh about the same as a proton, but could not have a charge – this new particle had to be electrically neutral.

In 1920, Ernest Rutherford tried to explain this phenomenon. He proposed that the "extra" particles were combinations of protons and electrons in the nucleus . These new particles would have a mass very similar to a proton , but would be electrically neutral since the positive charge of the proton and the negative charge of the electron would cancel each other out.

In 1930, German researchers bombarded the element beryllium with alpha particles (helium nuclei containing two protons and two neutrons with a charge of +2). The particles produced in this process had strong penetrating power, which suggested they were fairly large. In addition, they were not affected by a magnetic field , so they were electrically neutral. The French husband-wife research team of Frederic and Irene Joliot-Curie used these new "rays" to bombard paraffin, which was rich in protons. The unknown particles produced a large emission of protons from the paraffin.

The English physicist James Chadwick (1891-1974) repeated these experiments and studied the energy of these particles. By measuring velocities, he was able to show that the new particle has essentially the same mass as a proton. So we now have a third subatomic particle with a mass equal to that of a proton, but with no charge. This particle is called the neutron. Chadwick won the Nobel Prize in Physics in 1935 for his research.

Even though electrons, protons, and neutrons are all types of subatomic particles, they are not all the same mass. When you compare the masses of electrons, protons and neutrons, what you find is that electrons have an extremely small mass, compared to either protons or neutrons. On the other hand, the masses of protons and neutrons are fairly similar, although technically, the mass of a neutron is slightly larger than the mass of a proton. Because protons and neutrons are so much more massive than electrons, almost all of the mass of any atom comes from the nucleus, which contains all of the neutrons and protons.

 

Particle	Relative Mass (amu)	Electric Charge	Location
Electron	1/1840	+1	Outside the nucleus
Proton	1	+1	Nucleus
Neutron	1	0	Nucleus

The previous table shown gives the properties and locations of electrons, protons, and neutrons. The third column shows the masses of the three subatomic particles in grams. The second column shows the masses of the three subatomic particles in “atomic mass units”. Atomic mass units (amu) - (one-twelfth the mass of a carbon-12 atom) are useful because the mass of a proton and the mass of a neutron are almost exactly 1.0 in this unit system.

In addition to mass, another important property of subatomic particles is their charge. You already know that neutrons are neutral, and thus have no charge at all. Therefore, we say that neutrons have a charge of zero. What about electrons and protons? You know that electrons are negatively charged and protons are positively charged, but what’s amazing is that the positive charge on a proton is exactly equal in magnitude (magnitude means “absolute value” or “size when you ignore positive and negative signs”) to the negative charge on an electron. The third column in the table shows the charges of the three subatomic particles. Notice that the charge on the proton and the charge on the electron have the same magnitude.

Negative and positive charges of equal magnitude cancel each other out. This means that the negative charge on an electron perfectly balances the positive charge on the proton. In other words, a neutral atom must have exactly one electron for every proton. If a neutral atom has 1 proton, it must have 1 electron. If a neutral atom has 2 protons, it must have 2 electrons. If a neutral atom has 10 protons, it must have 10 electrons.

 

Atomic Number and Mass Number

If all atoms contain protons, neutrons, and electrons what makes each one unique? Scientists can distinguish between different elements by counting the number of protons. If an atom has only one proton, we know it’s a hydrogen atom. An atom with two protons is always a helium atom. If scientists count four protons in an atom, they know it’s a beryllium atom. An atom with three protons is a lithium atom, an atom with five protons is a boron atom, an atom with six protons is a carbon atom... the list goes on.

 

     

It is sometimes difficult to distinguish one element from another. Each element however, does have a unique number of protons. Sulfur has 16 protons. Silicon has 14 protons, and gold has 79 protons.

 

Since an atom of one element can be distinguished from an atom of another element by the number of protons in its nucleus, scientists are always interested in this number, and how this number differs between different elements. Therefore, scientists give this number a special name. An element’s atomic number (the number of protons in the nucleus of an atom) is a whole number usually written above the chemical symbol of each element. The modern periodic table is based on the atomic number of elements.

 

 

Of course, since neutral atoms have one electron for every proton, an element’s atomic number also tells you how many electrons are in a neutral atom of that element. For example, hydrogen has an atomic number of 1. This means that an atom of hydrogen has one proton, and, if it’s neutral, one electron as well. Gold, on the other hand, has an atomic number of 79, which means that an atom of gold has 79 protons, and, if it’s neutral, and 79 electrons as well.

The mass number of an atom is the total number of protons and neutrons in its nucleus. Why do you think that the “mass number” includes protons and neutrons, but not electrons? You know that most of the mass of an atom is concentrated in its nucleus. The mass of an atom depends on the number of protons and neutrons. You have already learned that the mass of an electron is very, very small compared to the mass of either a proton or a neutron (like the mass of a penny compared to the mass of a bowling ball). Counting the number of protons and neutrons tells scientists about the total mass of an atom.

Mass number = (number of protons) + (number of neutrons)

An atom’s mass number is very easy to calculate provided you know the number of protons and neutrons in an atom.

 

Atomic Number

 

Example

What is the mass number of an atom of helium that contains 2 neutrons?

Solution:

(Number of protons) = 2

(Number of neutrons) = 2

Mass number = (number of protons) + (number of neutrons) 2+2 = 4

 

There are two main ways in which scientists frequently show the mass number of an atom they are interested in. It is important to note that the mass number is not given on the periodic table. These two ways include writing a nuclear symbol or by giving the name of the element with the mass number written.

To write a nuclear symbol, the mass number is placed at the upper left (superscript) of the chemical symbol and the atomic number is placed at the lower left (subscript) of the symbol. The complete nuclear symbol for helium-4 is drawn above.

The following nuclear symbols are for a nickel nucleus with 31 neutrons and a uranium nucleus with 146 neutrons.

      

In the nickel nucleus represented above, the atomic number 28 indicates the nucleus contains 28 protons, and therefore, it must contain 31 neutrons in order to have a mass number of 59. The uranium nucleus has 92 protons as do all uranium nuclei and this particular uranium nucleus has 146 neutrons.

The other way of representing these nuclei would be Nickel-59 and Uranium-238 where 59 and 238 are the mass numbers of the two atoms, respectively. Note that the mass numbers (not the number of neutrons) is given to the side of the name.

 

 

Putting It Together

Let us revisit this phenomenon:

Look for patterns in the structure and function of these models of atoms.

1. What is the correct order of the models?
2. What evidence caused the atomic models to change over time?
3. What are the three subatomic particles? What charge does each one have?