 Lesson 4: Atomic Structure

OBJECTIVES

• Define the word atom

• Describe the four points of Dalton's atomic theory of matter

• Identify and describe the two kinds of electrical charges

• Describe how particles with the same charge affect each other

• Describe how particles of different charges affect each other

• Discuss how atoms are related to electricity

• Explain how cathode rays and radioactivity are related to atomic structure

• Explain how Rutherford's experiment showed the existence of the nucleus

• Describe alpha, beta, and gamma radiation

• Name and describe the three subatomic particles

• Determine the number of protons, neutrons, and electrons in an atom or ion

• Explain how an ion differs from an atom

• Determine number of subatomic particles from the symbol of a nuclide

• Identify isotopes

• Calculate average atomic mass from relative abundance

• Describe the changes that accompany nuclear reactions

• Perform calculations with half-lives

Atomic Structure, Atomic Number, Mass Number

The atom is the basic unit of matter. It is the smallest particle of an element that still has the characteristics of that element. Every atom has a positively charged central nucleus, composed of positively charged protons (1+) and uncharged neutrons. The nucleus is surrounded by a number of negatively charged electrons (1-). Like magnets, particles of opposite charge are attracted to one another, and particles with the same charge are repelled from one another. The structure of the atom that we are familiar with today has a long history to which many people contributed. The discovery of these particles and their arrangement, relative masses, and charges required many experiments.

The word atom comes from the Greek word atomos, meaning "indivisible," and indeed the idea that all matter is composed of tiny, indivisible particles goes back to ancient Greece. The first scientist in modern times to revive this idea was a British scientist named John Dalton. In 1803, after observing that elements in a given compound always exist in the same proportions, Dalton formulated his atomic theory of matter. The theory had four main points:

• Each element is composed of extremely small particles called atoms

• All atoms of a given element are identical; the atoms of different elements are different and have different properties (including different masses)

• Atoms of an element are not changed into different types of atoms by chemical reactions; atoms are neither created nor destroyed in chemical reactions

• Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

Michael Faraday was an English chemist who, in 1832-33, carried out a series of experiments attempting to use electricity to isolate elements from known compounds. His work led him to discover that the amount of electricity applied to a sample compound is related to the amount of an element that is isolated. He concluded that the structure of the atom must play a role in this.

Building on Faraday's experiments, English physicist J. J. Thompson discovered that electricity consisted of tiny subatomic particles called electrons. In 1897, Thompson was the first person to measure the mass-to-charge ratio of a single electron. He found that the particles he observed were common to all atoms. Thomson also knew that the negative charge he measured must be balanced by a particle of positive charge, since there was no overall charge on an atom. Thomson devised his own model of the atom, which is often described as the blueberry muffin model: he pictured an atom as a uniform ball of positive charge, with the electrons scattered uniformly throughout, like blueberries in a muffin.

The American physicist R. A. Millikan determined the unit of charge using oil drops and an electrical field. By calculating the velocity with which a charged drop was deflected from the electrical field, Millikan was able to calculate the charge on the drop. The mass could also be calculated. By performing many experiments, Millikan was able to calculate the charge on an individual electron and the mass of an individual electron.

The true arrangement of particles in the atom was discovered by Ernest Rutherford, another English physicist. Rutherford performed his famous platinum and gold foil experiments with a college student who was named Ernest Marsden. They determined that some positively charged helium nuclei, alpha particles (He2+) were deflected by platinum foil, at angles greater than 90 degrees. Sometimes the particles were deflected almost 180 degrees, back the way they came! For an alpha particle to be deflected that much, it must pass very close to a group of particles that also had a positive charge. The group of particles would have to have a strong positive charge, concentrated in a relatively small area. That is what the nucleus of an atom is – a relatively small area with a concentrated positive charge.

The individual protons and neutrons were more difficult to isolate and identify. At this time, it was known through Rutherford’s and others experiments that while the proton and electron had the same charge, that the mass of the proton was much greater. In any atom, the number of protons and electrons are the same. The existence of an uncharged particle was discovered by comparing the number of positive charges (number of protons) in an atom with the mass. The number of mass units was almost twice the number of protons. Where was the extra mass coming from?

James Chadwick’s experiments with beryllium and positively charged alpha particles revealed the presence of the neutron. By bombarding beryllium with alpha particles, Chadwick was able to observe that an uncharged particle was given off, which had the same mass as a proton. Chadwick used a specific type of beryllium, beryllium-9. The number 9 is the mass number, which indicates the total number if protons and neutrons in an atom. The number of protons in an atom of a specific element always identifies the element. The number of neutrons in atoms of the same element may be different. When neutrons are different in atoms of the same element, they are called isotopes. The element neon has two common isotopes, neon-20 and neon-21. Both isotopes have 10 protons, since neon is element number 10 in the Periodic Table. By subtracting the number of protons from the mass number, we see that neon-20 has 10 neutrons, and neon-21 has 11 neutrons. Beryllium-9 has 4 protons and 3 neutrons.

Mass number = number of protons + number of neutrons.

Each element is identified by a unique atomic number, which is shown on the Periodic Table. The atomic number is equal to the number of protons in each atom of the element. The number of protons and electrons in a neutral atom of an element is always the same.

How does this work? Let’s complete the chart for an example, oxygen-16. The mass number is the number after the dash, 16.

 Symbol Atomic Number Number of Protons Number of Electrons Mass Number Number of Neutrons O 8 8 8 16 8

For an atom of Ca-40, the mass number is the number after the dash, 40. The atomic number can be obtained by looking for Ca on the Periodic

Table. The atomic number is 20, the number of protons.

Mass number – number of protons = number of neutrons.

Number of neutrons equals 20. If  no charge is specified, the number of electrons is the same as the number of protons. The number of electrons is therefore 20.

 Symbol Atomic Number Number of Protons Number of Electrons Mass Number Number of Neutrons Ca 20 20 20 40 20

We can follow this process for any isotope for which a mass number is designated. If we do not know the mass number, we can use the Periodic Table to obtain an average atomic mass, which is usually a decimal, then round it to the nearest whole number. For instance, the element sulfur, S, has an average atomic mass on the Periodic Table of 32.07. By rounding this number to the nearest whole number, 32, we now have the mass number for a common isotope of sulfur. We can obtain the number of protons from the atomic number on the Periodic Table, and the number of electrons will be the same if no charge is indicated.

Ions and Ion Formation

The number of electrons can vary if charged particles called ions are formed. How can we determine the number of protons and electrons in an ion? Remember that the atomic number always gives the number of protons. We can find the atomic number on the Periodic Table, by looking in the same box as the symbol for an element. Each proton has a 1+ charge. An atom of an element has the same number of protons and electrons, and that means the total charge adds to zero. Atoms of an element are often said to be neutral for this reason.

Example: (Use the formula: positive charge + negative charge = net charge)

The element sodium, Na, has atomic number 11, meaning 11 protons and 11 electrons: (11+) + (11-) = 0. There is zero net (overall) charge, the atom is neutral; it is not an ion.

An atom of chlorine has atomic number 17, 17 protons and 17 electrons: (17+) + (17-) = 0. Zero net charge, this is a neutral atom.

Let’s try this for ions. An ion is an atom that carries a charge either through gaining or losing an electron. A cation is an atom that loses an electron and has a positive charge. An anion is an atom that gains an electron and has a negative charge. We do not observe negative ions by themselves, they are balanced by positive ions. Using an everyday substance as our example, we can determine the number of protons and electrons for a positive ion, Na1+, and a negative ion, Cl1-. These ions are found in table salt, sodium chloride, NaCl. The ions are formed from the elements sodium, Na, and chlorine, Cl.

The sodium ion, Na1+, has atomic number 11. This means 11 protons, 11 positive charges. The total charge is 1+, overall. That means that there is one “extra” or unbalanced charge. What about the other 10 positive charges? They are still there in the nucleus, and their charge is balanced by 10 electrons. Use the formula positive charge + negative charge = net charge.

We already know that 11 protons = 11+, and net charge = 1+ ; so substitute into the equation:

(11+) + number of negative charges = 1+

The math on this one is easy! There are 10 negative charges (10-), which means 10 electrons. The sodium ion, Na1+ , has 10 electrons; a neutral atom of sodium has 11 electrons. The sodium atom loses one electron to form the positive ion. Where does the electron go? Let’s take a look at the negative ion in the NaCl compound.

For the chloride ion, Cl1-, the process is the same. The atomic number is 17, indicating 17 protons. The net charge is 1-. How many of those negative charges are balanced with a positive charge?

Remember: Positive charge + negative charge = net charge

(17+) + number of negative charges = 1-

number of negative charges = 18- , and number of negative charges equals number of electrons, 18!

A neutral chlorine atom has 17 electrons. The chloride ion has 18 electrons. A chlorine atom has to gain an electron to form the negative ion. Therefore, ions are formed by the loss or gain of electrons. This process works for any ion, for which you know the symbol and charge. The symbol allows you to determine the atomic number by looking at a Periodic Table.

Calculating the Average Atomic Mass

The average atomic mass of an element is determined by averaging the natural abundance of its isotopes.

 Element Isotope Mass Number Mass (amu) Fractional Abundance Average Atomic Mass Carbon 12 C 12 12 (exactly) 98.89% 12.01 6 13 C 13 13.003 1.11 6 Chloride 35 Cl 35 34.969 75.53 35.45 17 37 Cl 37 36.966 24.47 17 Silicon 28 Si 28 27.977 92.21 28.09 14 29 Si 29 28.976 4.70 14 30 Si 30 29.974 3.09 14

Example: Use the table above to calculate the average atomic mass of silicon.

Multiply the mass by the percent abundance for each of the three isotopes of silicon. (Change the percentage abundance to a decimal by moving the decimal point two places to the left or by dividing by 100.)

(27.977) (0.9221) = 25.80 amu

(28.976) (0.0470) = 1.362 amu

(29.974) (0.0309) = 0.926 amu

Find the sum of the results. 25.80 + 1.36 + 0.926 = 28.08 amu

Round to three significant figures, 28.1 amu.

The average atomic mass of silicon is 28.1 amu.

The French researcher Antoine Henri Becquerel, made the first observations of natural, spontaneous radioactivity. While studying fluorescent properties of substances, he noticed that photographic plates placed next to certain uranium compounds became darkened, as if they had been exposed to light! This is not what he had intended to study, but this accidental discovery had an enormous impact. One of the students in his lab was Marie Curie, who went on to win the Nobel Prize in physics in 1903 with her husband Pierre Curie.  She was awarded another Nobel Prize in 1911 for chemistry for her work on radium and polonium. Marie Curie is one of only three people to win two Nobel Prizes in science. Her daughter and son-in-law shared a Nobel Prize in chemistry in 1935.

Because of these discoveries, Rutherford and other researchers were aware of the existence of elements that are naturally radioactive. This means that the elements undergo a specific type of change by emitting particles of different sizes and charges, with differing amounts of energy, from the nuclei of atoms. Nuclear decay (radioactive decay) causes the number of protons in the nucleus to change. Since the identity of an element is defined by the number of protons in its nuclei, nuclear decay results in the transmutation of one element to another. This type of change, nuclear change, is very different from ordinary chemical reactions.

Two types of particles are emitted during natural nuclear decay: alpha particles and beta particles. Alpha particles consist of two neutrons and two protons bound together--the equivalent of a positively charged helium ion (He2+). Let’s ignore the charge for a moment, and look at the symbol in a new way. Alpha particles are often written in the following form:

 4 2 He

The 4 in the upper left is the mass number. The 2 in the lower left is the atomic number. The helium ion still has its 2+ charge; it is just not shown in many nuclear reactions. The arrangement of particles in the nucleus and the changes that occur in radioactive decay processes get special attention, and the charge is typically ignored. Charges are most often important in ordinary chemical reactions, where they receive much more consideration.

An example of alpha decay is the naturally occurring decay of uranium-238:

 238 92 U ® 234 90 Th + 4 2 He

An alpha particle also has a 2+ charge, even though we do not see/write the charge in nuclear reactions. Note that the element starting material on the left of the arrow, uranium-238, is not the same as the element products helium-4 or thorium-234.

A beta particle is formed when a neutron from the nucleus breaks down into a proton and an electron. The electron, called a beta particle, is expelled from the nucleus. It is important to remember that a beta particle (β-) is not one of the original set of electrons that are outside the nucleus.

An example of naturally occurring beta decay is thorium-234:

 234 90 Th ® 234 91 Pa + β-

The beta symbol is used to be sure that we realize that this is not an electron from outside the nucleus. We can also recognize the correct mass and charge by using another symbol, e:

 234 90 Th ® 234 91 Pa + 0 -1 e

Note that the element starting material on the left of the arrow, thorium-234, is not the same as the element product protactinium-234.

In addition to alpha and beta particles, nuclear decay also produces gamma rays. Gamma rays are very high energy electromagnetic radiation with essentially no charge or mass. They have more penetrating power than X-rays.

Half-Lives

Not all radioactive elements decay at the same rate. Some artificially created isotopes are very unstable and only exist for a few minutes before breaking down. At the other extreme, the radioactive isotope rubidium-87 decays to strontium-87 extraordinarily slowly; it takes 60 billion years (!) for 50% of a particular sample of 87Rb to decay. The time it takes for one half of a sample of a radioactive isotope to decay is called the isotope's half-life. The half-life of 87Rb is therefore 60 billion years. Carbon-14 has a half-life of 5,730 years, and the half-life of the artificial radioisotope iodine-131 is 0.022 years (about eight days).