Chapter 5 Periodic Law

Section 5-1 History of the Periodic Table

The 1800’s – A Time of Chemical Discovery In the 1800’s, new elements were being discovered at a fast rate. Scientists realized they needed a way to classify them. Also needed a method to determine atomic mass. In 1860, a conference determined how to find relative atomic mass.

Dmitri Mendeleev (1834 – 1907) Used new atomic mass values in a textbook. Attempted to organize elements according to their properties.

Mendeleev discovered periodic trends When he arranged elements in order of increasing atomic mass, similarities in properties appeared at regular intervals. This pattern is referred to as periodic.

The First Periodic Table Mendeleev grouped elements with similar properties together, to create the first periodic table. A copy is found on page 124.

Unique Characteristics of Mendeleev’s Table Mendeleev left gaps in his periodic table. These can be seen at elements 45,68, 70. Mendeleev predicted both existence and properties of these elements.

Filling the Gaps By 1886, all three of these elements were discovered. These are now known as Sc, Ga, and Ge. The properties of these three are very similar to what Mendeleev predicted.

Mendeleev’s Periodic Table Worked. As a result of his successful predictions, the scientific community was forced to take Mendeleev seriously.

Mendeleev’s Periodic Table Worked Mendeleev became known as the founder of periodic law.

Criticisms of Mendeleev’s Table Mendeleev did not exactly follow the order of atomic mass. Notice that Te, with an atomic mass of 128 comes before I with an atomic mass of 127. The scientific community disregarded Mendeleev’s work because he changed his rules to fit his table.

Henry Moseley – (1887 – 1915) Studied atomic emission spectra. Found a pattern in them. Led him to discover atomic number. Mendeleev’s periodic table was in perfect order of atomic number.

Periodic Law Moseley’s discovery changed the definition of periodic law. The periodic law states that when elements are arranged in order of increasing atomic number, elements with similar properties will appear at regular intervals.

The Modern Periodic Table Over 40 new elements have been discovered since Mendeleev. Most of these can be grouped by properties with other elements, but some form three new groups.

Three New Groups of Elements Noble Gases – Group VIII A. These are difficult to observe because they are odorless, colorless, and very unreactive. Lanthanides – the 14 elements from 58 to 71. Actinides – the 14 elements from 90 to 103.

Periodicity All group A elements show periodicity. The pattern of periodicity in group A elements is: 8, 8, 18, 18, 32. These are the numbers of elements between elements with similar properties. The periodic table is divided into four sections.

Section 5-2 Electron Configuration and the Periodic Table

The s - block S – block elements are chemically reactive metals. All group 1 elements end in s 1. Group IA elements are called alkali metals. They are more reactive than Group IIA. All have a silvery appearance and can be cut with a knife. All react strongly with water, and can pull moisture out of the air. Are usually stored in oil, and are rarely found in a pure form in nature.

Group IIA elements are called alkaline earth metals. All of their electron configurations end in s 2. Harder, denser, and stronger than alkali metals. Higher melting points, and are less reactive, but still too reactive to be found in nature.

The s - block Helium and Hydrogen are special cases. Hydrogen ends in s 1, but is not an alkali metal. It is a unique element that is considered to be in its own family, and its properties do not resemble those of any known group.

The s - block Helium ends in s 2. However, it is more closely related to elements in group VIIIA. Since the first energy level has only an s orbital, and two electrons fill it, helium has a full outer energy level. This gives it the same stability and unreactivity as a noble gas.

The d - block More complex electron configurations. All metals, called transition metals. They are lustrous, and good conductors. These are more typical metals. Palladium, Platinum, and Gold are among the least reactive elements. All transition metals are less reactive than the s – block.

The p – block, Groups IIIA - VIIIA Contains mostly nonmetals, a few metals, and all metalloids. Group VIIA are the halogens, which are the most reactive nonmetals. They all form salts with metals. F and Cl are gases, Br is the only liquid nonmetal, and I is a solid. As is synthetic.

The p – block, Groups IIIA - VIIIA The metalloids are found along the stair step line, and are all brittle solids. They are also good semiconductors. The metals in the p – block are harder and denser than s – block metals, but softer than transition metals. They are also stable in air.

The f – block elements, the Lanthanides and the Actinides Found at the bottom of the periodic table. These are called Rare Earth Metals. Lanthanides are shiny and relatively reactive, like group IIA. Actinides are all radioactive. The first four are found on earth, all others are man – made.

Arrangement of the Periodic Table The periodic table is arranged so that the period number is equal to the highest occupied energy level. The length of a period also equals the number of electrons in an energy level. The number of electrons in the highest energy level for all group A elements is equal to the group number.

Section 5-3 Electron Configuration and Periodic Properties

Atomic Radius One half the distance between the nuclei of two identical atoms that are bonded together. What happens to Atomic Radius as you travel from left to right across a period?

Atomic Radius It decreases. As you move across the period you add more electrons to the same energy level, but you also add more protons to the nucleus. More protons exerts a stronger attractive force and pulls the electrons in more tightly. What happens to atomic radius as you go down a group?

Atomic Radius It goes up. While you add more protons to the nucleus, which exert a stronger pull, you also add more electrons, but this time, you add them in higher energy levels. These higher energy levels are farther from the nucleus, making the radius larger. Fr has the largest radius, He the smallest.

Ionization Energy The energy required to remove an electron from a neutral atom. An ion is an atom or group of atoms with a positive or negative charge. As we move from left to right across a period, does ionization energy go up or down?

It goes up. Electrons are held more tightly by the nucleus, and so, are harder to remove. Generally, you can only remove electrons from Groups IA, IIA, and IIIA. In some cases, they can also be removed from IVA.

Ionization Energy Some elements, particularly those in groups IIA and IIIA, can lose more than one electron. The energy required to do this is called the second or third ionization energy. These are always higher than the first ionization energy. This is because as each electron is removed, the others are held more tightly by the nucleus.

Ionization Energy As a rule, Group IA loses one electron, Group IIA loses two electrons, and Group IIIA loses three electrons. The transition metals also lose electrons, but the number varies due to the overlap of orbitals. What happens to Ionization Energy as you go down a group?

Ionization Energy It goes down. As you go down a column, you add energy levels to each atom. The energy levels put the outer electrons farther away from the nucleus. Electrons that are farther from the nucleus are easier to remove.

Electron Affinity The energy change that occurs when a neutral atom gains an electron. Most atoms release energy as they gain electrons. As you move from left to right on the periodic table, what happens to electron affinity?

Electron Affinity It goes up. In most cases, only the elements in Group VA, VIA, and VIIA gain electrons, and so they have the highest electron affinity. This is because they have more protons in their nuclei, and so attract electrons more strongly.

Electron Affinity Some atoms can gain more than one electron. Generally, atoms in Group VA gain three electrons, atoms in group VIA gain two electrons, and atoms in group VIIA gain one electron. The second or third electron affinities are usually lower than the first. This is because you have the same number of protons pulling on more electrons so they pull more weakly.

Electron Affinity Group VIIIA, the noble gases, have no electron affinity. They have a full sublevel, and do not need them. The same is true of group IIA. The d-block and f-block have unpredictable electron affinities.

What happens to electron affinity as you move down a group? It goes down. The larger size of the atoms keeps the electrons farther from the nucleus, and therefore they are not attracted as strongly.

Ionic Radius Ionic Radius is determined by the type of ion formed. A positive ion is a cation. These tend to be smaller than the atom they came from. These are formed by Group IA, IIA, and IIIA, and the transition metals.

Ionic Radius A negative ion is an anion, and they tend to be larger than the atom they came from. These are formed by Group VA, VIA, and VIIA. Ionic Radius tends to decrease from left to right across a period. It tends to increase going down a column.

Ionic Radius As a rule, metals tend to form cations, and nonmetals form anions. Group IVA and VIIIA do not gain or lose electrons, and so, do not form ions.

Electronegativity A measure of the ability of an atom in a compound to attract electrons. Tend to increase going from left to right across a period, and decrease going down a column. Noble gases have no electronegativity, since they do not form compounds.

Electronegativity The most electronegative element is F, fluorine. The least electronegative element is Fr, francium.

Valence Electrons Any electron in the highest occupied energy level. Eight valence electrons provides the maximum stability an atom can have. This is referred to as the octet rule. The number of valence electrons in each group A element is equal to its group number.