Atomic Structure and the Periodic Table of the Elements

Periodic Trends

-As goes to right, acid properties increase and basic properties decrease. -As goes to right, atomic radii decreases -As goes to right, ionization energy increases -As goes to right, nonmetallic properties increase

Trend of Atomic Radii

-Atomic radii decrease from left to right across a period in the Periodic Table (until the noble gases) -Atomic radii increase from top to bottom in a group or family

Common Characteristic Properties of Transition Elements

-They often form colored compound -They can have a variety of oxidation states -At least one of their compounds has an incomplete d electron subshell -They are often good catalysts -They are silvery blue at room temperature (except copper and gold) -They are solids at room temperature (except mercury) -They form complex ions -They are often paramagnetic due to unpaired electrons.

Radioactive Decay and Nuclear Charge

-alpha decay: decay particle: a; particle mass: 4; particle charge: +2; change in nucleon number: decreases by 4; change in atomic number: decreases by 2 -beta decay: decay particle: B; particle mass: 0; particle charge: -1; change in nucleon number: no change; change in atomic number: increases by 1 -gamma radiation: decay particle: y; particle mass: 0; particle change: 0; change in nucleon number: no change; change in atomic number: no change -positron emission: decay particle: B+; particle mass: 0; particle change: +1; change in nucleon number: no change; change in atomic number: decreases by 1 -electron capture: decay particle: e-; particle mass: 0; particle charge: -1; change in nucleon number: no change; change in atomic number: decreases by 1

measuring atomic radius

-metals: done by measuring the distance between two nuclei in the solid state and dividing this distance by 2. Such measurements can be made with x-ray diffraction -nonmetallic elements: exists in pure form as a molecule, such as chlorine, measurements can be made of the distance between nuclei for two atoms covalently bonded together.

Dalton's Five Basic Principles

1.) All matter is made up of very small, discrete particles called atoms. 2.) All atoms of an element are alike in weight, and this weight is different from that of any other atom. 3.) Atoms cannot be subdivided, created, or destroyed 4.) Atoms of different elements combine in simple whole-number ratios to form chemical compounds. 5.) In chemical reactions, atoms are combined, separated, or rearranged

3 methods of detection of alpha, beta, and gamma rays

1.) photographic plate 2.) scintillation counter 3.) geiger counter

nonmetals

13-17

noble gases

18

Heavy Metals

3-12. brittle: 3-7. ductile: 8-11. Low melting: 12

transmutation

A conversion of an element to a new element, due to a change in the number of protons. Can be produced artifically by bombarding the nuclei of a substance with various particles from a particle accelerator, such as the cyclotron.

Scintillation counter

A fluorescent screen (e.g., ZnS) will show the presence of electrons and X-rays, as already mentioned. If the screen is viewed with a magnifying eye-piece, small flashes of light, called scintillations, will be observed. By observing the scintillations, one not only can detect the presence of alpha particles, but also can actually count them.

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 CO2 in the atmosphere contains carbon-14. The stable isotope of carbon is carbon-12. Carbon-14 is a beta emitter and decays to form nitrogen 14. In any living organism, the ratio of carbon-14 to carbon-12 is the same as in the atmosphere because of the constant interchange of materials between organism and surroundings. When an organism dies, this interaction stops, and the carbon-14 gradually decays to nitrogen. By comparing the relative amounts of cabron-14 and carbon-12 in the remains, the age of the organism can be established.

quantum mechanics or wave mechanics

A more mathematical model developed in the 1920s concerning the behavior of electrons inside the atom.

electronegativity

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. Electronegativity decreases down a group and increases across a period. The inert gases can be ignored. The lower the electronegativity number, the more electropositive an element is said to be. The most electronegative element is in the upper right corner-F, fluorine. The most electropositive is in the lower left corner of the chart- Fr, francium.

nuclear chain reaction

A reaction in which an initial step leads to a succession of repeating steps that continues indefinitely. Nuclear chain reactions are used in nuclear reactors and nuclear bombs

John Newlands

A scientist who in 1863 proposed the idea of repeating octaves of properties in the development of a systematic pattern for the elements

deuterium

Accounts for 0.015% of Earth's hydrogen atoms. Each deuterium atom has a nucleus containing one proton and one neutron

groups

Also known as families, the vertical columns on the Periodic Table

Ionization Energy Trend Down Table

Among these elements, the energy needed gradually declines. This can be explained by considering the distance of the involved energy level from the positively charged nucleus. With each succeeding noble gas, a more distant p orbital is involved, therefore making it easier to remove an electron from the positive attraction of the nucleus. As more energy levels are added to the atomic structure as the atomic number increases, the additional negative fields associated with the additional electrons screen out some of the positive attraction of the nucleus.

Erwin Schrödinger

An Austrian physicist, who in conjunction with Heisenberg, agreed with the de Broglie concept that the electron is bound to the nucleus in a manner similar to a standing wave.

Henry Moseley

An English scientist who first determined the atomic numbers of the elements through the use of X-rays. Stated that the properties of elements are a periodic function of their atomic numbers, thus changing the basis of the periodic law from atomic weight to atomic number.

spectroscope

An instrument which examines the light emitted by energized atoms. It contains a prism or diffraction grating which disperse the light to allow and examination of the spectra

Noble Gas Notation

Another method of simplifying the electron distribution to the orbitals. In this method you represent all of the lower filled orbitals up to the closest noble gas. By enclosing its symbol in brackets, it represents all of the complete noble gas configuration. Then the remaining orbitals are written in the usual way.

mass spectroscopy

Another tool used to identify specific atomic structures. Based on the concept that differences in mass cause differences in the degree of bending that occurs in a beam of ions passing through a magnetic field. It separates isotopes of the same element based on differences in their mass. The intensity on the photographic plate indicates the amount of each particular isotope.

excited state

Any level higher than the ground state

Nucleus

Area of the atom in which most of the mass of the atom was located and which was positively charged

History of Radiation

At the same time advances in atomic theory were occurring, scientists were noticing phenomena associated with emissions from the nucleus of atoms in the form of X-rays. While Roentgen announced the discovery of X-Rays, Becquerel was exploring the phosphorescence of some materials. Becquerel's work received little attention until early in 1898, when Marie and Pierre Curie entered the picture. Searching for the source of the intense radiation in uranium ore, Marie and Pierre Curie used tons of it to isolate very small quantities of two new elements, radium and polonium, both radioactive. Along with Becquerel, the Curies shared the Nobel Prize in physics in 1903.

isotopes

Atoms of the same element that have different masses. The isotopes of a particular element all have the same number of protons and electrons but different numbers of neutrons. Most elements consist of mixtures of isotopes. Tin, for example, has ten stable isotopes, the most of any element. 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.

atomic spectra

Based on the release of energy when an electron moves from an excited state to the ground state. When electrons drop to energy level one, ultraviolet, when electrons drop to energy level two, visible light, when electrons drop to energy level 3, infrared

wave-mechanical model

Complex equations, the solutions of which give specific wave functions calls orbitals. The electron does not move in a circular orbit in this model. Rather, the orbital is a 3 dimensional region around the nucleus that indicates the probable location of an electron but gives no information about its pathway. Made up of four quantum numbers, which give the position with respect to the nucleus, the shape of the orbital, its spatial orientation, and the spin of the electron in the orbital

Electron Configuration Exceptions

Cr: 1s2 2s2 2p6 3s2 3p6 3d5 4s1 Cu: 1s2 2s2 2p6 3s2 3p6 3d10 4s1

photons

Discrete radiant energy, released when energy is released in the "allowed" values

spectra

Distinct colored lines which result when light is dispersed through a prism. Since only particular energy jumps are available in each type of atom, each element has its own unique emission spectra made up of only the lines of specific wavelength that correspond to its atomic structure. Spectral lines can be used in the identification of unknown specimens.

Atomic Radii in Groups

For a group of elements, the atoms of each successive member have 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. Therefore, the atomic radius increases down a group.

alkali metals

Group 1 metals, which are active metals which react with water to form bases

Light Metals

Groups 1 and 2

covalent radius

Half of the distance between the nuclei for two atoms covalently bonded together

Nuclear fission reactions

Have been in use since the 1940s. The first atomic bombs used in 1945 were nuclear fission bombs. Since that time, many countries, including our own, have put nuclear fission power plants into use to provide a new energy source for electrical energy. Basically a nuclear fission reaction is the splitting of a heavy nucleus into two or more lighter nuclei.

J.J. Thomson and Cathode Ray Experiment

He is credited with the discovery of the electron as the first subatomic particle in 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 electrical and magnetic fields. Therefore, he concluded that cathode rays are made up of very small, negatively charged particles. 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. He was awarded a Nobel Prize in 1906 for his accomplishment

gamma rays

High-energy radiations similar to x-rays. electromagnetic radiation identical with light; high energy, no charge 1.) Beta particles and gamma rays are usually emitted together; after a beta is emitted a gamma ray follows 2.) Arrangement in nucleus is unknown. Same velocity as visible light. 3.) Range: no specific range 4.) Shielding needed: about 13 cm of lead 5.) Interactions: weak of itself; gives energy to electrons, which then perform the ionization

periods

Horizontal rows of the periodic table. 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 3 periods are short. 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.

John Dalton

In 1805, he 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.

Dimitri I. Mendeleev

In 1869, proposed a table containing 17 columns and is usually given credit for the first period table since he arranged elements in groups according to their atomic weights and properties. It is interesting to note that Lothar Meyer proposed a similar arrangement about the same time. 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 (such as Fe, Co, Ni), and a second period of seven elements. The first and second periods of seven across were later distinguished by use of the letters A and B attached to the group symbols, which were Roman numerals. This nomenclature of periods (IA, IIA, etc.) has been revised in the present Periodic Table. His table had the elements arranged by atomic weights with recurring properties in a periodic manner. Where atomic weight placement disagreed with the properties that should occur in a particular spot in the table, mendeleev gave preference to the element with the correct properties. He even predicted elements for places that were not yet occupied in the table. These predictions proved to be amazingly accurate and led to wide acceptance of his table.

Niels Bohr

In 1913 the Danish physicist Niels Bohr published a theory explaining the line spectrum of hydrogen. He proposed a planetary model that quantized the energy of electrons to specific orbits.

Bohr Model of the Atom

In 1913, Niels Bohr (Denmark) 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 concentrically around the nucleus, and each level is designated by the number: 1, 2, 3,... The closer to the nucleus, the less energy an electron needs in one of these levels, but it has to gain energy to go from one level to another that is farther away from the nucleus. Because of its simplicity and general ability to explain chemical change, the Bohr model still has some usefulness today. Bohr's electron distribution t principal energy levels has the formula 2n^2. 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.

G. N. Lewis

In 1916, devised the electron dot notation, which may be used in place of the electron configuration notation. The electron notation shows only the chemical symbol surrounded by dots to represent the electrons in the incomplete outer level. 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.

nodes

Indicate regions of zero probability

parent nuclide

Initial element in a decay reaction

Robert Millikan and Oil Drop Experiment

It was an American scientist who in 1909 was able to measure the charge on an electron using the apparatus. Oil droplets were sprayed into the chamber and, in the process, became randomly charged by gaining or losing electrons. The electric field was adjusted so that a negatively charged drop would move slowly upward in front of the grid in the telescope. Knowing the rate at which the drop was rising the strength of the field, and the mass of the drop, Millikan was able to calculate the charge on the drop. Combining the information with the results of Thomson, he could calculate a value for the mass of a single electron. Eventually, this number was found to be 9.11X10^-28 gram.

Late 19th Century Discoveries

J.J. Thompson discovered the electron beam in a cathode ray tube in 1897. Soon afterward, Henri Bacquerel 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. More and more physicists turned their attention to the structure of the atom.

most active metal

Lower left corner

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.

Neutron

Nuclear particle that would be neutral and account for missing mass. Discovered by James Chadwick in 1932. n^0, symbol: 1 0 n, actual mass: 1.675 X 10^-24, relative mass compared to proton: 1, discovery: J.C. Chadwick 1932

nucleons

Particles (protons and neutrons) shown in the nucleus in models

Ernest Rutherford and Gold Foil Experiment

Performed a gold foil experiment that had tremendous implications for atomic structure in England, 1911. Using alpha particles confirmed that there was mostly empty space between the nucleus and electrons. Alpha particles (helium nuclei) 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 2 years of studying the results, Rutherford finally came up with an explanation. He reasoned that the rebounded 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. If the nucleus were the size of a marble, then the atom would be about the size of a football field. 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 new nuclear particle that would be neutral and would account for the missing mass.

Lewis Structure

Picture with the element, showing only the valence electrons as dots in an electron dot notation. Shows the atomic symbol to represent the nucleus and inner shell electrons. It shows dots to represent the valence electrons.

Proton

Positive particle in the nucleus. p^+, symbol: 1 1 H, actual mass: 1.673X10^-24 g, discovery: early 1900s

alpha particle

Positively charged particles of helium nuclei (4 2 He), 2+. 1.) Ejection reduces the atomic number by 2, the atomic weight by 4 amu 2.) High energy, relative velocity 3.) Range: about 5 cm in air 4.) Shielding needed: stopped by the thickness of a sheet of paper, skin 5.) Interactions: produces about 100,000 ionizations per centimeter; repelled by the positively charged nucleus; attracts electrons, but does not capture them until its speed is much reduced

Pauli Exclusion Principle

Postulate of Wolfgang Pauli, which states that in a given atom no two electrons can have the same set of four quantum numbers. Therefore, each orbital can hold only two electrons. No two electrons can have the same four quantum numbers

magnetic quantum number

Represented by m sub l. s=1 space-oriented orbital, p=3 space-oriented orbitals, d=5 space-oriented orbitals, f=7 space-oriented orbitals. The value of m can equal -l,...0...+l. The number of spatial orientations of orbitals.

Spin Quantum Number

Represented by m sub s. + spin or - spin. The value of m= +1/2 or -1/2. Describes the spin in either of two possible directions. Each orbital can be filled by only two electrons with opposite spins.

Angular Momentum Quantum Number

Represented by the letter l, refers to the shape of the orbital. The number of possible shapes is limited by the principal quantum number. The first energy level has only one possible shape, the s orbital because n=1 and the limit of l=(n-1)=0. The second has two possible shapes, the s and the p. s, p, d, f (in order of increasing energy). The value of l can be = 0, 1,..., n-1. l=0 indicates spherical-shaped s orbital. l=1 indicates a dumbbell-shaped p orbital. l=2 indicates a five orbital orientation d orbital

Atomic Radii in Periods

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 because of the additional protons pulls the electrons more tightly around the nucleus. This attraction more than balances the repulsion between the added electrons and the other electrons, and the radius is generally reduced. The inert gas at the end of the period has a slight increase in radius because of the electron repulsion in the filled outer principal energy level.

Uncertainty Principle

Stated by Werner Heisenberg in 1927. States that it is impossible to know both the precise location and precise velocity of a subatomic particle at the same time.

beta particle

Streams of high-speed electrons. Negatively charged, 1- 1.) Ejected when a neutron decays into a proton and an electron 2.) High velocity, low energy 3.) Range: about 12 m 4.) Shielding needed: stopped by 1 cm of aluminum or thickness of average book. 5.) Interactions: weak because of high velocity, but produces about 100 ionizations per centimeter

Max Planck

Suggested in his quantum theory of light that light has both particlelike properties and wavelike characteristics.

valence of the atom

The absolute number of electrons gained, lost or borrowed

First Ionization Energy

The amount of energy supplied to one outer electron to remove it from its atom.

classical mechanics

The branch of physics that deals with the motion of bodies under the influence of forces. Used by Bohr to calculate the orbits of the hydrogen atom, but could not explain the ability of electrons to stay in only certain energy levels without the loss of energy or why a change of energy occurred only when an electron "jumped"from one energy level to another and why the electron could not exist at any energy level between these levels.

valence electron

The electrons found in the outermost energy level

Transition Elements

The elements involved with the filling of a d sublevel with electrons after two electrons are in the s sublevel of the next principal energy level. The first examples of these are the elements between calcium, atomic number 20, and gallium, atomic number 31. Their electron configurations are the same in the 1s, 2s, 2p, 3s, and 3p sublevels. With the variable number of electrons available for bonding, it is not surprising that transition elements can exhibit variable oxidation numbers.

Paschen series

The emissions, consisting of infrared, that occur when an electron cascades from a level higher than the first level down to n=3

Lyman series

The emissions, consisting of ultraviolet radiation, that occur when an electron cascades from a level higher than the first level down to n=1

Balmer series

The emissions, consisting of visible light, that occur when an electron cascades from a level higher than the first level down to n=2

photographic plate

The fogging of a photographic emulsion led to the discovery of radioactivity. If this emulsion is viewed under a high-power microscope, it is seen that beta and gamma rays cause the silver bromide grains to develop in a scattered fashion.

radon

The gas produced by spontaneous disintegration

History of Atoms

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 Democritus and Leucippus. These particles were considered to be so small and indestructible that they could not be divided into smaller particles. Atoms is the Greek word for indivisible. The english word atom comes from the 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.

ground state

The lowest energy state available to the electron

protium

The most common type of hydrogen, which accounts for 99.985% of the hydrogen atoms found on Earth. TThe nucleus of a protium atom contains one proton only, and it has one electron moving about it.

atomic number

The number of protons in the nucleus of an atom. 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.

Periodic Law

The periodic table is sorted based on atomic number

Aufbau Principle

The principles that an electron occupies the lowest energy orbital that can receive it.

daughter nuclide

The resulting element in a decay reaction

mass number

The sum of the number of protons and the number of neutrons in the nucleus

tritium

The third form of hydrogen, which is radioactive. It exists in very small amounts in nature, but it can be prepared artificially. Each tritium atom contains one proton, two neutrons, and one electron.

half-life

The time required for half of the atoms of a radioactive nuclide to decay

average atomic mass

The weighted average of the atomic masses of the naturally occurring isotopes of an element. The average atomic mass of an element depends on both the ass and the relative abundance of each of the element's isotopes. Can be calculated by multiplying the atomic mass of each isotope by its relative abundance (expressed in decimal form) and adding the results.

Louis de Broglie

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. A young French physicist who in 1924, suggested that if light can have both wavelike and particlelike characteristics as Planck had suggested that, then perhaps particles can also have wavelike characteristics. In 1927, de Broglie's ideas were verified experimentally when investigators showed that electrons could produce diffraction patterns, a property associated with waves. Diffraction patterns are produced by waves as they pass through small holes or narrow slits.

Geiger counter

This instrument is perhaps the most widely used at the present time for determining individual radiation. Any particle that will produce an ion gives rise to an avalanche of ions, so the type of particle cannot be identified. However, each individual particle can be detected.

Principal Quantum Number

This number refers to average distance of the orbital from the nucleus. 1 is closest to the nucleus and has the least energy. The numbers correspond to the orbits in the Bohr model. They are called energy levels. Represented by the letter n. Refers to the principal energy level

most active nonmetal

Upper right-hand corner

Hund's Rule of Maximum Multiplicity

When there is more than one orbital at a particular energy level, only one electron will fill each orbital until each has one electron. Pairing will occur with the addition of one more electron to each orbital.

relationship with wavelength and frequency

With increased wavelength, decreased frequency

Ionization Energy

With the first electron gone, the removal of succeeding electrons becomes more difficult because of the loss of repulsive effects that were present with a greater number of electrons. It should be noted that the lowest ionization energies are found with the least electronegative. Can be plotted against atomic numbers. The highest ionization energy needed to remove the first electron comes from the outer energy level of the noble gases.

Ionization Energy Trend Across Table

Within a period, the ionization energy generally increases. The lowest occurs when a lone electron occupies the outer s orbital. As the s orbital fills with two electrons at atomic number 4, the added stability of the 2s orbital explain the small peak at 4. At atomic number 5, a lone electron occupies the 2p orbital. This electron can be removed with less energy, and therefore a dip occurs in the graph. With the 2p orbitals filling according to Hund's Rule, with only one electron in each orbital before pairing occurs, again a slightly more stable situation, and therefore, another small peak occurs at atomic number 7. After this peak, a dip and continued increases occur until the 2p orbitals are completely filled with paired electrons at the noble gas.

3 Kinds of Radioactive Emissions

alpha particles, beta particles, and gamma rays

metalloids

amphoteric elements along line/stairs. These elements have certain characteristics of metals and other characteristics of nonmetals. Some examples are boron, silicon, arsenic and tellurium

Visible Light Spectra Wavelengths

red (700 nm), orange (600 nm), yellow, green (500 nm), blue, violet (400 nm)

electrons

very small, negatively charged particles. e^-, symbol: 0 -1 e, actual mass: 9.109X10^28 g, mass compared to proton: 1/1,837, discovery: J.J. Thomson, 1897