chapter 4 atom structure john dalton (1766- 1844), an english schoolteacher and chemist, studied the...

Chapter 4

ATOM STRUCTURE

• John Dalton (1766-1844), an English schoolteacher and chemist, studied the results of experiments by Lavoisier, Proust, and many other scientists.

Dalton’s Atomic Theory

• The following statements are the main points of Dalton’s atomic theory.

Dalton’s Atomic Theory

1. All matter is made up of indestructible atoms.

2. All atoms of one element are exactly

alike, but are different from atoms of other elements.

3. In a chemical reaction, atoms combine in small whole number ratios.

• Because of Dalton’s atomic theory, most scientists in the 1800s believed that the atom was like a tiny solid ball that could not be broken up into parts.

The Electron

• In 1897, a British physicist, J.J. Thomson, discovered that this solid-ball model was not accurate.

• Thomson’s experiments used a vacuum tube.

• A vacuum tube has had all gases pumped out of it.

The Electron

• At each end of the tube is a metal piece called an electrode, which is connected through the glass to a metal terminal outside the tube.

• These electrodes become electrically charged when they are connected to a high-voltage electrical source.

• When the electrodes are charged, rays travel in the tube from the negative electrode, which is the cathode, to the positive electrode, the anode.

Cathode-Ray Tube

• Because these rays originate at the cathode, they are called cathode rays.

• Thomson found that the rays bent toward a positively charged plate and away from a negatively charged plate.

Cathode-Ray Tube

• He knew that objects with like charges repel each other, and objects with unlike charges attract each other.

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• These electrons had to come from the matter (atoms) of the negative electrode.

Cathode-Ray Tube

• Thomson concluded that cathode rays are made up of invisible, negatively charged particles referred to as electrons.

• From Thomson’s experiments, scientists had to conclude that atoms were not just neutral spheres, but somehow were composed of electrically charged particles.

Cathode-Ray Tube

• Reason should tell you that there must be a lot more to the atom than electrons.

• Matter is not negatively charged, so atoms can’t be negatively charged either.

• If atoms contained extremely light, negatively charged particles, then they must also contain positively charged particles—probably with a much greater mass than electrons.

Cathode-Ray Tube

• In 1886, scientists discovered that a cathode-ray tube emitted rays not only from the cathode but also from the positively charged anode.

Protons

• These rays travel in a direction opposite to that of cathode rays.

• Like cathode rays, they are deflected by electrical and magnetic fields, but in directions opposite to the way cathode rays are deflected.

Protons

• Thomson was able to show that these rays had a positive electrical charge.

• Years later, scientists determined that the rays were composed of positively charged subatomic particles called protons.

• At this point, it seemed that atoms were made up of equal numbers of electrons and protons.

Protons

• However, in 1910, Thomson discovered that neon consistedof atoms of two different masses.

Protons

• Today, chemists know that neon consists of three naturally occurring isotopes.

• Atoms of an element that are chemically alike but differ in mass are called isotopes of the element.

• The third was too scarce for Thomson to detect.

Neutrons

• Calculations showed that such a particle should have a mass equal to that of a proton but no electrical charge.

• Because of the discovery of isotopes, scientists hypothesized that atoms contained still a third type of particle that explained these differences in mass.

• The existence of this neutral particle, called a neutron, was confirmed in the early 1930s.

Rutherford’s Gold Foil Experiment

• In 1909, a team of scientists led by Ernest Rutherford in England carried out the first of several important experiments that revealed an arrangement far different from the cookie-dough model of the atom.

Rutherford’s Gold Foil Experiment

• The experimenters set up a lead-shielded box containing radioactive polonium, which emitted a beam of positively charged subatomic particles through a small hole.

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• The sheet of gold foil was surrounded by a screen coated with zinc sulfide, which glows when struck by the positively charged particles of the beam.

• Today, we know that the particles of the beam consisted of clusters containing two protons and two neutrons and are called alpha particles.

Rutherford’s Gold Foil Experiment

The Gold Foil Experiment

The Nuclear Model of the Atom

• Because most of the particles passed through the foil, they concluded that the atom is nearly all empty space.

• To explain the results of the experiment, Rutherford’s team proposed a new model of the atom.

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The Nuclear Model of the Atom

• Because so few particles were deflected, they proposed that the atom has a small, dense, positively charged central core, called a nucleus.

The Nuclear Model of the Atom• The new model of the atom as pictured by

Rutherford’s group in 1911 is shown below.

Atomic Numbers

• It is the number of protons that determines the identity of an element, as well as many of its chemical and physical properties.

• The atomic number of an element is the number of protons in the nucleus of an atom of that element.

Atomic Numbers

• Therefore, the atomic number of an element also tells the number of electrons in a neutral atom of that element.

• Because atoms have no overall electrical charge, an atom must have as many electrons as there are protons in its nucleus.

• The sum of the protons and neutrons in the nucleus is the mass number of that particular atom.

Masses

• The mass of a neutron is almost the same as the mass of a proton.

Masses

• Isotopes of an element have different mass numbers because they have different numbers of neutrons, but they all have the same atomic number.

Atomic Mass

• In order to have a simpler way of comparing the masses of individual atoms, chemists have devised a different unit of mass called an atomic mass unit, which is given the symbol u.

• An atom of the carbon-12 isotope contains six protons and six neutrons and has a mass number of 12.

Atomic Mass

• Chemists have defined the carbon-12 atom as having a mass of 12 atomic mass units.

• Therefore, 1 u = 1/12 the mass of a carbon-12 atom.

• 1 u is approximately the mass of a single proton or neutron.

Information in the Periodic Table

• The number at the bottom of each box is the average atomic mass of that element.

• This number is the weighted average mass of all the naturally occurring isotopes of that element.

Question 1

How does the atomic number of an element differ from the element’s mass number?

Answer

The atomic number of an element is the number of protons in the nucleus. The mass number is the sum of the number of protons and neutrons.

Calculating Atomic Mass

Calculating Atomic Mass

• Copper exists as a mixture of two isotopes.

• The lighter isotope (Cu-63), with 29 protons and 34 neutrons, makes up 69.17% of copper atoms.

• The heavier isotope (Cu-65), with 29 protons and 36 neutrons, constitutes the remaining 30.83% of copper atoms.

Calculating Atomic Mass

• The atomic mass of Cu-63 is 62.930 amu, and the atomic mass of Cu-65 is 64.928 amu.

• Use the data above to compute the atomic mass of copper.

Calculating Atomic Mass• First, calculate the contribution of each

isotope to the average atomic mass, being sure to convert each percent to a fractional abundance.

Calculating Atomic Mass

• The average atomic mass of the element is the sum of the mass contributions of each isotope.

Question 2

The table on the next slide shows the five isotopes of germanium found in nature, the abundance of each isotope, and the atomic mass of each isotope.

Calculate the atomic mass of germanium.

Answer

72.591 amu

• Nuclear chemistry is the study of the structure of atomic nuclei and the changes they undergo.

Nuclear Radiation

• In 1895, Wilhelm Roentgen (1845–1923) found that invisible rays were emitted when electrons bombarded the surface of certain materials.

The Discovery of Radioactivity

• The emitted rays were discovered because they caused photographic plates to darken. Roentgen named these invisible high-energy emissions X rays.

• As is true in many fields, Roentgen’s discovery of X rays created excitement within the scientific community and stimulated further research.

The Discovery of Radioactivity

• At that time, French physicist Henri Becquerel (1852–1908) was studying minerals that emit light after being exposed to sunlight, a phenomenon called phosphorescence.

The Discovery of Radioactivity

• Building on Roentgen’s work, Becquerel wanted to determine whether phosphorescent minerals also emitted X rays.

The Discovery of Radioactivity

• Becquerel accidentally discovered that phosphorescent uranium salts—even whennot exposed to light—produced spontaneous emissions that darkened photographic plates.

• Marie Curie (1867–1934) and her husband Pierre (1859–1906) took Becquerel’s mineral sample (called pitchblende) and isolated the components emitting the rays.

The Discovery of Radioactivity

The Discovery of Radioactivity

• They concluded that the darkening of the photographic plates was due to rays emitted specifically from the uranium atoms present in the mineral sample.

• Marie Curie named the process by which materials give off such rays radioactivity; the rays and particles emitted by a radioactive source are called radiation.

• As you may recall, isotopes are atoms of the same element that have different numbers of neutrons.

Types of Radiation

• Isotopes of atoms with unstable nuclei are called radioisotopes.

Types of Radiation

• These unstable nuclei emit radiation to attain more stable atomic configurations in a process called radioactive decay.

• During radioactive decay, unstable atoms lose energy by emitting one of several types of radiation.

Types of Radiation

• The three most common types of radiation are alpha (α), beta (β), and gamma (γ).

Types of Radiation

• The three most common types of radiation are alpha (α), beta (β), and gamma (γ).

Types of Radiation• Ernest Rutherford (1871–1937), whom you

know of because of his famous gold foilexperiment that helped define modern atomic structure, identified alpha, beta, and gamma radiation when studying the effects of an electric field on the emissions from a radioactive source.

Types of Radiation• The effect of an electric field on three types

of radiation is shown. • Positively charged alpha particles are

deflected toward the negatively charged plate.

Types of Radiation

• Negatively charged beta particles are deflected toward the positively charged plate.

Types of Radiation

• Beta particles undergo greater deflection because they have considerably less mass than alpha particles.

Types of Radiation

• Gamma rays, which have no electrical charge, are not deflected.

Types of Radiation

• The charge of an alpha particle is 2+ due to the presence of the two protons.

• An alpha particle (α) has the same composition as a helium nucleus—two protons and two neutrons—and is therefore given the symbol .

Types of Radiation

• Alpha radiation consists of a stream of alpha particles.

• Radium-226, an atom whose nucleus contains 88 protons and 138 neutrons, undergoes alpha decay by emitting an alpha particle.

Types of Radiation

• Notice that after the decay, the resulting atom has an atomic number of 86, a mass number of 222, and is no longer radium.

• The newly formed radioisotope is radon-222.

Types of Radiation

• The particles involved are balanced. That is, the sum of the mass numbers (superscripts) and the sum of the atomic numbers (subscripts) on each side of the arrow are equal.

Types of Radiation

• Because of their mass and charge, alpha particles are relatively slow-moving compared with other types of radiation.

• Thus, alpha particles are not very penetrating—a single sheet of paper stops alpha particles.

Types of Radiation

• A beta particle is a very-fast moving electron that has been emitted from a neutron of an unstable nucleus.

• Beta particles are represented by the symbol. The zero superscript indicates the

insignificant mass of an electron in comparison with the mass of a nucleus.

Types of Radiation

• The –1 subscript denotes the negative charge of the particle.

• Beta radiation consists of a stream of fast-moving electrons.

Types of Radiation

• An example of the beta decay process is the decay of iodine-131 into xenon-131 by beta-particle emission.

Types of Radiation• Note that the mass number of the product

nucleus is the same as that of the original nucleus (they are both 131), but its atomic number has increased by 1 (54 instead of 53).

Types of Radiation

• This change in atomic number, and thus, change in identity, occurs because the electron emitted during the beta decay has been removed from a neutron, leaving behind a proton.

Types of Radiation

• Because beta particles are both lightweight and fast moving, they have greater penetrating power than alpha particles.

• A thin metal foil is required to stop beta particles.

Types of Radiation

• As you can see from the symbol, both the subscript and superscript are zero.

• Gamma rays are high-energy (short wavelength) electromagnetic radiation. They are denoted by the symbol .

Types of Radiation

• Thus, the emission of gamma rays does not change the atomic number or mass number of a nucleus.

• Gamma rays almost always accompany alpha and beta radiation, as they account for most of the energy loss that occurs as a nucleus decays.

Types of Radiation

• For example, gamma rays accompany the alpha-decay reaction of uranium-238.

• The 2 in front of the γ symbol indicates that two gamma rays of different frequencies are emitted.

• Because gamma rays have no effect on mass number or atomic number, it is customary to omit them from nuclear equations.

Radioactive Decay

• It may surprise you to learn that of all the known isotopes, only about 17% are stable and don’t decay spontaneously.