Chapter 2 Atomic Structure and the Periodic Table of Elements

Chapter 2: Atomic Structure and the Periodic Table of Elements - - 1. 2. 3. 4. 5. - - The idea of small, invisible particles being the building blocks of matter can be traced back more than 2,000 years to the Greek philosophers Demoscritus and Leucippus. These particles, considered to be so small and indestructible that they could not be divided into smaller particles, were called atoms, the Greek word for indivisible. The English word atom comes from this Greek word. This early concept of atoms was not based upon experimental evidence but was simply a result of thinking and reasoning on the part of the philosophers. It was not until the eighteenth century that experimental evidence in favor of the atomic hypothesis began to accumulate. Finally, around 1805, John Dalton proposed some basic assumptions about atoms based on what was known through scientific experimentation and observation at that time. These assumptions are very closely related to what scientists presently know about atoms. For this reason, Dalton is often referred to as the father of modern atomic theory. Some of these ideas: All matter is made up of very small, discrete particles called atoms. All atoms of an element are alike in weight, and this weight is different from that of any other kind of atom. Atoms cannot be subdivided, created, or destroyed. Atoms of different elements combine in simple whole-number ratios to form chemical compounds. In chemical reactions, atoms are combined, separated, or rearranged. By the second half of the 1800s, many scientists believed that all the major discoveries related to the elements had been made. The only thing left for young scientists was to do was to refine what was already known. This came to a surprising halt when J.J. Thompson discovered the electron beam in a cathode ray tube in 1897. Soon afterward, Henri Becquerel announced his work with radioactivity, and Marie Curie and her husband, Pierre, set about trying to isolate the source of radioactivity in their laboratory in France. During the late nineteenth and early twentieth centuries, more and more physicists turned their attention to the structure of the atom. In 1913 the Danish physicist Niels Bohr published a theory explaining the spectrum line of hydrogen. He proposed a planetary model that quantized the energy of electrons to specific orbits. The work of Louis de Broglie and others in the 1920s and 1930s showed that quantum theory described a more probabilistic model of where the electrons could be found that resulted in the theory of orbitals. Basic Electric Charges The discovery of the electron as the first subatomic particle is credited to J.J. Thomson (England, 1897). He used an evacuated tube connected to a spark coil. As the voltage across the tube was increased, a beam became visible. This was referred to as a cathode ray. - - - - Thomson found that the beam was deflected by both electric and magnetic fields. Therefore, he concluded that cathode rays are made up of very small, negatively charged particles, which became known as electrons. Further experimentation led Thomson to find the ratio of the electrical charge of the electron to its mass. This was a major step toward understanding the nature of the particle. - Ernest Rutherford (England, 1911) performed a gold foil experiment that had tremendous implications for atomic structure. - Alpha particles passed through the foil with few deflections. - However, some deflections were almost directly back toward the source. This was unexpected and suggested an atomic model with mostly empty space between a nucleus, in which most of the mass of the atom was located and which was positively charged, and the electrons that defined the volume of the atom. After two years of studying the results, Rutherford finally came up with an explanation. He reasoned that the rebound alpha particles must have experienced some powerful force within the atom. And he assumed this force must occupy a very small amount of space, because so few alpha particles had been deflected. He concluded that the force must be a densely packed bundle of matter with a positive charge. He called this positive bundle the nucleus. He further discovered that the volume of a nucleus was very small compared with the total volume of an atom. The electrons, he suggested, surrounded the positively charged nucleus like planets around the sun, even though he could not explain their motion. Further experiments showed that the nucleus was made up of still smaller particles called protons. Rutherford realized, however, that protons, by themselves, could not account for the entire mass of the nucleus. He predicted the existence of a nuclear particle that would be neutral and would account for the missing mass. In 1932, James Chadwick discovered this particle, the neutron. Bohr Model of the Atom - In 1913, Niels Bohr proposed his model of the atom. This pictured the atom as having a dense, positively charged nucleus and negatively charged electrons in specific spherical orbits, also called energy levels or shells around this nucleus. - These energy levels are arranged in circles around the nucleus, and each energy level is designated by a number: 1, 2, 3, ...The closer to the nucleus, the less energy an electron needs, but it has to gain energy to go from one level to another that is father away from the nucleus. Principal Energy Level Max Number of Electrons (2𝑛2 ) 1 2 2 8 3 18 4 32 5 50 Components of Atomic Structure - - - Particle Charge Electron Symbol Actual Mass Relative Mass Compared to Proton Discovery −(𝑒 − ) 9.109 × 10−28 g 1/1,837 J.J. Thomson-1897 Proton +(𝑝+ ) 1.673 × 10−24 g 1 Early 1900s Neutron 0(𝑛0 ) 1.675 × 10−24 g 1 J.C. Chadwick1932 The number of protons in the nucleus of an atom determines the atomic number. All atoms of the same element have the same number of protons and therefore the same atomic number; atoms of different elements have different atomic numbers. Thus, the atomic number identifies the element. The sum of the number of protons and the number of neutrons in the nucleus is called the mass number. In some cases, different types of atoms of the same element have different masses. Protium, deuterium, and tritium are isotopes of hydrogen. Isotopes are atoms of the same element that have different masses. The isotopes of a particular element have the same number of protons and electrons but different number of neutrons. In all three isotopes of hydrogen, the positive charge of the single proton is balanced by the negative charge of the electron. Most elements consist of mixtures of isotopes. The percentage of each isotope in the naturally occurring element on Earth is nearly always the same, no matter where the element is found. The percentage at which each of an element’s isotopes occurs in nature is taken into account when calculating the element’s average atomic mass. Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element. Calculating Average Atomic Mass - - The average atomic mass of an element depends on both the mass and the relative abundance of each of the element’s isotopes. For example, naturally occurring copper consists of 69.17% copper-63, which has an atomic mass of 62.919 598 amu, and 30.83% copper-65, which has an atomic mass of copper can be calculated by multiplying the atomic mass of each isotope by its relative abundance (expressed in decimal form) and adding the results. 0.6917 × 62.919 598 𝑎𝑚𝑢 + 0.3083 × 64.927 793 𝑎𝑚𝑢 = 63.55 𝑎𝑚𝑢 Therefore, the calculated average atomic mass of naturally occurring copper is 63.55 amu. Average atomic masses of the elements listed in the Periodic Table, rounded to one decimal place for use in calculations. Valence Electrons - Each atom attempts to have its outer energy level complete and accomplishes this by borrowing, lending, or sharing its electrons. - The electrons found in the outermost energy level are called valence electrons. - The remainder of the electrons are called core electrons. - The absolute number of electrons gained, lost, or borrowed is referred to as the valence of the atom. - This picture can be simplified to , showing only the valence electrons as dots in an electron dot notation. This is called the Lewis structure of the atom. To complete its outer orbit to eight electrons, chlorine must borrow an electron from another atom. Its valence number then is 1. As stated above, when electrons are gained, we assign a −sign to this number, so the oxidation number of chlorine is −1. Since sodium tends to lose this electron, its oxidation number is +1. Atomic Spectra - The Bohr model was based on a simple postulate. Bohr applied to the hydrogen atom the concept that the electron can exist only in certain energy levels without an energy change but that, when the electron changes its state, it must absorb or emit the exact amount of energy that will bring it from the initial state to the final state. - The ground state is the lowest energy state available to the electron. - The excited state is any level higher than the ground state. Lewis Structure (Electron Dot Notation) - - In 1916, G. N. Lewis devised the electron dot notation, which may be used in place of the electron configuration notation. The electron dot notation shows only the chemical symbol surrounded by dots to represent the electrons in the incomplete outer level. Examples are: The symbol denotes the nucleus and all electrons except the valence electrons. The dots are arranged at the four sides of the symbol and are paired when appropriate. PERIODIC TABLE OF ELEMENTS - The history of the development of a systematic pattern for the elements includes the work of a number of scientists. - Dimitri I. Mendeleev in 1869 proposed a table containing 17 columns and is usually given credit for the first periodic table since he arranged elements in groups according to their atomic weights and properties. - In 1871 Mendeleev rearranged some elements and proposed a table of eight columns, obtained by splitting each of the long periods across into a period of seven elements, an eighth group containing the three central elements, and a second period of seven elements. Mendeleev’s table had the elements arranged by atomic weights with recurring properties in a periodic manner. - The horizontal rows of the periodic table are called periods or rows. There are seven periods, each of which begins with an atom having only one valence electron and ends with a complete outer shell structure of an inert gas. The first three periods are short, consisting of 2, 8, and 8 elements, respectively. Periods 4 and 5 are longer, with 18 each, while period 6 has 32 elements, and period 7 is incomplete with 22 elements, most of which are radioactive and do not occur in nature. The vertical columns of the Periodic Table are called groups or families. The elements in a group exhibit similar or related properties. In 1984 the IUPAC agreed that the groups would be numbered 1 through 18. PROPERTIES RELATED TO THE PERIODIC TABLE - Metals are found on the left of the chart with the most active metal in the lower left corner. - Nonmetals are found on the right side with the most active nonmetal in the upper right corner. - The noble or inert gases are on the far right. - Since the most active metals react with water to form bases, the Group 1 metals are called alkali metals. - As you proceed to the right, the base-forming property decreases and the acid-forming properties increase. - The metals in the first two groups are light metals, and those toward the center are heavy metals. - The elements found along the dark line in the periodic table are called metalloids. These elements have certain characteristics of metals and other characteristics of nonmetals. General summary statements about the Periodic Table: - Acid-forming properties increase from left to right on the table. - Base-forming properties are high on the left side and decrease to the right. - The atomic radii of elements decrease from left to right across a period. - First ionization energy increases from left to right across a period. - Metallic properties are greatest on the left side of the table and decrease to the right. - Nonmetallic properties are greatest on the right side of the table and decrease to the left. Atomic Radii 1. Atomic Radii decreases from left to right across a period in the Periodic Table (until the noble gases). a. Since the number of electrons in the outer principal energy level increases as you go from left to right in each period, the corresponding increase in the nuclear charge b/c of the additional protons pulls the electrons more tightly around the nucleus. 2. Atomic radii increase from top to bottom in a group or family. a. For a group of elements, the atoms of each successive member of another outer principal energy level in the electron configuration and the electrons there are held less tightly by the nucleus. This is because of their increased distance from the nuclear positive charge and the shielding of this positive charge by all the core electrons. Ionic Radius Compared with Atomic Radius - Metals tend to lose electrons in forming positive ions. - With this loss of negative charge, the positive nuclear charge pulls in the remaining electrons closer and thus reduces the ionic radius below that of the atomic radius. - Nonmetals tend to gain electrons in forming negative ions. - With this added negative charge, which increases the inner electron repulsion, the ionic radius is increased beyond the atomic radius. Electronegativity - The electronegativity of an element is a number that measures the relative strength with which the atoms of the element attract valence electrons in a chemical bond. - This electronegativity number is based on an arbitrary scale going from 0 to 4. In general, a value of less than 2 indicates a metal. - The electronegativity decreases down a group and increases across a period. Ionization Energy - Atoms hold their valence electrons, then, with different amounts of energy. - If enough energy is supplied to one outer electron to remove it from its atom, this amount of energy is called the first ionization energy. - With the first electron gone, the removal of succeeding electrons becomes more difficult because of the imbalance between the positive nuclear charge and the remaining electrons. - The lowest ionization energy is found with the least electronegative atom. THE NATURE OF RADIOACTIVE EMISSIONS - Three different types of radioactive emission: alpha particles, beta particles, and gamma rays. - Alpha particles are positively charged particles of helium nuclei, beta particles are streams of high-speed electrons, and gamma rays are high-energy radiations similar to X-rays. DECAY SERIES, TRANSMUTATIONS, AND HALF-LIFE - The nuclei of uranium, radium, and other radioactive elements are continually disintegrating. It should be emphasized that spontaneous disintegration produces the gas known as radon. The time required for half of the atoms of a radioactive nuclide to decay is called its half-life. Ex: - Radium on average, half of the radium nuclei present will have disintegrated to radon in 1,590 years - In another 1,590 years half of the remainder will decay and so on When a radium disintegrates, it loses an alpha particle upon gaining two electrons and becomes a neutral helium atom The remainder becomes radon A conversion of an element to a new element (b/c of a change in the number of protons) is a transmutation. The following uranium-radium disintegration series shows how a radioactive atom may change when it loses each kind of article. Note that an atomic number is shown by a subscript (92U) and the isotopic mass by a subscript (238U). The alpha particle is represented by the Greek symbol, 𝛼, and the beta particle by 𝛽 . Radioactive Decay and Nuclear Charge Type of Decay Decay Particle Particle Mass Particle Charge Change in Nucleon Number Change in Atomic Number Alpha Decay 𝛼 4 2+ Decreases by 4 Decreases by2 Beta Decay 𝛽 0 1- No Change Increases by 1 Gamma Radiation 𝛾 0 0 No Change No Change Positron emission 𝛽+ 0 1+ No Change Decreases by 1 Electron capture 𝑒− 0 1- No Change Decreases by 1 Nuclear Systems for Subatomic Particles Particle Symbols Proton p Neutron n Electron 𝑒 − or 𝛽− Positron 𝑒 + or 𝛽+ Alpha Particle 𝛼 Beta Particle 𝛽or 𝛽− Gamma Ray 𝛾 RADIOACTIVE DATING - A helpful application of radioactive decay is in the determination of the ages of substances such as rocks and relics that have bits of organic material trapped in them. - Because carbon-14 has a half-life of about 5,700 years and occurs in the remains of organic materials, it has been useful in dating these materials. A small percentage of 𝐶𝑂2in the atmosphere contains carbon14. The stable isotope of carbon is carbon-12. Carbon-14 is a beta emitter and decays to form nitrogen14: - In any living organism, the ratio of carbon-14 to carbon-12 is the same as in the atmosphere b/c of the constant interchange between organisms and surroundings. When an organism dies, this interaction stops, and the carbon-14 gradually decays to nitrogen. By comparing the relative amounts of carbon-14 and carbon-12 in the remains, the age of the organism can be established. NUCLEAR REACTIONS Nuclear fission reactions are the splitting of a heavy nucleus into two or more lighter nuclei. - A nuclear chain reaction is a reaction in which an initial step, such as the reaction above, leads to a succession of repeating steps that continues indefinitely. A nuclear fission reaction is the combination of very light nuclei to make a heavier nucleus. Extremely high temperatures and pressures are required in order to overcome the repulsive forces of the two nuclei. The energy released in a nuclear reaction (either fission or fusion) comes from the fractional amount of mass converted into energy. Nuclear changes convert matter into energy. Energy released during nuclear reactions is much greater than the energy released during chemical reactions.