Ch. 5 Notes (1) PDF

Section 1 History of the Periodic Section 1 History of the Periodic Chapter 5 Table Chapter 5 Table Objectives...

Section 1 History of the Periodic Section 1 History of the Periodic Chapter 5 Table Chapter 5 Table Objectives Mendeleev and Chemical Periodicity Explain the roles of Mendeleev and Moseley in the Mendeleev noticed that elements arranged in order of development of the periodic table. increasing atomic mass had similarities in their chemical properties appear at regular intervals. Describe the modern periodic table. Repeating patterns are referred to as periodic. Explain how the periodic law can be used to predict the physical and chemical properties of elements. Mendeleev created a table in which elements with similar properties were grouped together—a periodic Describe how the elements belonging to a group table of the elements. of the periodic table are interrelated in terms of atomic number. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved. Section 1 History of the Periodic Section 1 History of the Periodic Chapter 5 Table Chapter 5 Table Mendeleev and Chemical Periodicity, Properties of Some Elements Predicted By Mendeleev continued After Mendeleev placed all the known elements in his periodic table, several empty spaces were left. In 1871 Mendeleev predicted the existence and properties of elements that would fill three of the spaces. By 1886, all three of these elements had been discovered. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved.  Section 1 History of the Periodic Section 1 History of the Periodic Chapter 5 Table Chapter 5 Table Moseley and the Periodic Law Periodicity of Atomic Numbers Henry Moseley discovered that the elements fit into patterns better when they were arranged according to atomic number, rather than atomic weight. The Periodic Law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers. In other words properties of elements will repeat at regular intervals based on the elements atomic numbers. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved. Section 1 History of the Periodic Section 2 Electron Configuration Chapter 5 Table Chapter 5 and the Periodic Table The Modern Periodic Table Objectives The Periodic Table is an arrangement of the Explain the relationship between electrons in elements in order of their atomic numbers so that sublevels and the length of each period of the elements with similar properties fall in the same periodic table. column, or group. Locate and name the four blocks of the periodic table. Explain the reasons for these names. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved.  Section 2 Electron Configuration Section 2 Electron Configuration Chapter 5 and the Periodic Table Chapter 5 and the Periodic Table Objectives, continued Periods and Blocks of the Periodic Table Elements are arranged vertically in the periodic table Discuss the relationship between group in groups that share similar chemical properties. configurations and group numbers. Elements are also organized horizontally in rows, Describe the locations in the periodic table and the or periods. general properties of the alkali metals, the alkaline- The length of each period is determined by the earth metals, the halogens, and the noble gases. number of electrons that can occupy the sublevels (blocks) being filled in that period. The periodic table is divided into four blocks, the s, p, d, and f blocks. The name of each block is determined by the electron sublevel being filled in that block. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved. Section 2 Electron Configuration Section 2 Electron Configuration Chapter 5 and the Periodic Table Chapter 5 and the Periodic Table Relationship Between Periodicity and Periods and Blocks of the Periodic Table, Electron Configurations continued The elements of Group 1 of the periodic table are known as the alkali metals. lithium, sodium, potassium, rubidium, cesium, and francium In their pure state, all of the alkali metals have a silvery appearance and are soft enough to cut with a knife. The elements of Group 2 of the periodic table are called the alkaline-earth metals. beryllium, magnesium, calcium, strontium, barium, and radium Group 2 metals are less reactive than the alkali metals, but are still too reactive to be found in nature in pure form. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved.  Section 2 Electron Configuration Section 2 Electron Configuration Chapter 5 and the Periodic Table Chapter 5 and the Periodic Table Periods and Blocks of the Periodic Table, Periods and Blocks of the Periodic Table continued The d sublevel first appears when n = 3. Hydrogen has an electron configuration of 1s1, but despite the ns1 configuration, it does not share the The 3d sublevel is slightly higher in energy than the 4s sublevel, same properties as the elements of Group 1. so these are filled in the order 4s first, then 3d. Hydrogen is a unique element. Like the Group 2 elements, helium has an ns2 group configuration. Yet it is part of Group 18. Because its highest occupied energy level is filled by two electrons, helium possesses special The d-block elements are metals with typical metallic properties and are often referred to as transition chemical stability. Just like the Noble Gases. elements. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved. Section 2 Electron Configuration Section 2 Electron Configuration Chapter 5 and the Periodic Table Chapter 5 and the Periodic Table Periods and Blocks of the Periodic Table, Periods and Blocks of the Periodic Table, continued continued The elements of Group 17 are known as the halogens. The p-block elements consist of all the elements of fluorine, chlorine, bromine, iodine, and astatine Groups 13–18 except helium. The halogens are the most reactive nonmetals. The p-block elements together with the s-block elements are called the main-group elements. They react vigorously with most metals to form examples of the type of compound known as salts. The properties of elements of the p block vary greatly. At its right-hand end, the p block includes all of the The metalloids, or semiconducting elements, are nonmetals except hydrogen and helium. located between nonmetals and metals in the p block. All six of the metalloids are also in the p block. The metals of the p block are generally harder and denser than the s-block alkaline-earth metals, but At the left-hand side and bottom of the block, there are eight softer and less dense than the d-block metals. p-block metals. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved.  Section 2 Electron Configuration Section 2 Electron Configuration Chapter 5 and the Periodic Table Chapter 5 and the Periodic Table Periods and Blocks of the Periodic Table, continued In the periodic table, the f-block elements are wedged between Groups 3 and 4 in the sixth and seventh periods. Their position reflects the fact that they involve the filling of the 4f sublevel. The first row of the f block, the lanthanides, are shiny metals similar in reactivity to the Group 2 alkaline metals. The second row of the f block, the actinides, are between actinium and rutherfordium. The actinides are all radioactive. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved. Section 3 Electron Configuration Section 3 Electron Configuration Chapter 5 and Periodic Properties Chapter 5 and Periodic Properties Objectives Atomic Radii Define atomic and ionic radii, ionization energy, To compare different atomic radii, they must be electron affinity, and electronegativity. measured under specified conditions. Compare the periodic trends of atomic radii, Atomic radius may be defined as one-half the ionization energy, and electronegativity, and state the distance between the nuclei of identical atoms that reasons for these variations. are bonded together. Define valence electrons, and state how many are present in atoms of each main-group element. Compare the atomic radii, ionization energies, and electronegativities of the d-block elements with those of the main-group elements. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved.  Section 3 Electron Configuration Section 3 Electron Configuration Chapter 5 and Periodic Properties Chapter 5 and Periodic Properties Atomic Radii, continued Periodic Atoms tend to be smaller the farther to the right they Trends of Radii are found across a period. The trend to smaller atoms across a period is caused by the increasing positive charge of the nucleus, which attracts electrons toward the nucleus. Atoms tend to be larger the farther down in a group they are found. The trend to larger atoms down a group is caused by the increasing size of the electron cloud around an atom as the number electron sublevels increases. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved. Section 3 Electron Configuration Section 3 Electron Configuration Chapter 5 and Periodic Properties Chapter 5 and Periodic Properties Atomic Radii, continued Atomic Radii, continued Sample Problem E Sample Problem E Solution Of the elements magnesium, Mg, chlorine, Cl, sodium, Sodium has the largest atomic radius Na, and phosphorus, P, which has the largest atomic radius? Explain your answer in terms of trends of the All of the elements are in the third period. Of the four, periodic table. sodium has the lowest atomic number and is the first element in the period. Atomic radii decrease across a period. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved.  Section 3 Electron Configuration Section 3 Electron Configuration Chapter 5 and Periodic Properties Chapter 5 and Periodic Properties Ionic Radii Ionic Radii, continued A positive ion is known as a cation. Cationic and anionic radii decrease across a period. The formation of a cation by the loss of one or more electrons always leads to a decrease in atomic radius. The electron cloud shrinks due to the increasing The electron cloud becomes smaller. nuclear charge acting on the electrons in the The remaining electrons are drawn closer to the same main energy level. nucleus by its unbalanced positive charge. The outer electrons in both cations and anions are in A negative ion is known as an anion. higher energy levels as one reads down a group. The formation of an anion by the addition of one or more electrons always leads to an increase in There is a gradual increase of ionic radii down atomic radius. a group. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved. Section 3 Electron Configuration Section 3 Electron Configuration Chapter 5 and Periodic Properties Chapter 5 and Periodic Properties Ionization Energy Ionization Energy, continued An ion is an atom or group of bonded atoms that In general, ionization energies of the main-group has a positive or negative charge (gains or loses elements increase across each period. an electron). This increase is caused by increasing nuclear charge. Sodium (Na), for example, easily loses an A higher charge more strongly attracts electrons in the same electron to form Na+. energy level. Among the main-group elements, ionization Any process that results in the formation of an energies generally decrease down the groups. ion is referred to as ionization. Electrons removed from atoms of each succeeding element in a group are in higher energy levels, farther from the The energy required to remove one electron from nucleus. a neutral atom of an element is the ionization The electrons are removed more easily. energy, IE (or first ionization energy, IE1). Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved.  Section 3 Electron Configuration Section 3 Electron Configuration Chapter 5 and Periodic Properties Chapter 5 and Periodic Properties Ionization Energy, continued Ionization Energy, continued Periodic trends in ionization energy are shown in the graph below. Sample Problem F Consider two main-group elements, A and B. Element A has a first ionization energy of 419 kJ/mol. Element B has a first ionization energy of 1000 kJ/mol. Decide if each element is more likely to be in the s block or p block. Which element is more likely to form a positive ion? Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved. Section 3 Electron Configuration Section 3 Electron Configuration Chapter 5 and Periodic Properties Chapter 5 and Periodic Properties Ionization Energy, continued Electron Affinity Sample Problem F Solution The energy change that occurs when an electron is Element A has a very low ionization energy, which means that acquired by a neutral atom is called the atom’s atoms of A lose electrons easily. electron affinity. Element A is most likely to be an s-block metal because ionization energies increase across the periods. Electron affinity generally increases across periods. Element B has a very high ionization energy which means that atoms of B have difficulty losing electrons. Increasing nuclear charge along the same sublevel attracts electrons more strongly Element B would most likely lie at the end of a period in the p block. Electron affinity generally decreases down groups. Element A is more likely to form a positive ion because it has The larger an atom’s electron cloud is, the farther a much lower ionization energy than element B does. away its outer electrons are from its nucleus. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved.  Section 3 Electron Configuration Section 3 Electron Configuration Chapter 5 and Periodic Properties Chapter 5 and Periodic Properties Valence Electrons Electronegativity Chemical compounds form because electrons are Valence electrons hold atoms together in chemical lost, gained, or shared between atoms. compounds. The electrons that interact in this manner are In many compounds, the negative charge of the those in the highest energy levels. valence electrons is concentrated closer to one atom than to another. The electrons available to be lost, gained, or shared in the formation of chemical compounds Electronegativity is a measure of the ability of an are referred to as valence electrons. atom in a chemical compound to attract electrons from another atom in the compound. Valence electrons are often located in incompletely filled main-energy levels. Electronegativities tend to increase across example: the electron lost from the 3s sublevel of Na to periods, and decrease or remain about the same form Na+ is a valence electron. down a group. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved. Section 3 Electron Configuration Section 3 Electron Configuration Chapter 5 and Periodic Properties Chapter 5 and Periodic Properties Electronegativity, continued Electronegativity, continued Sample Problem G Sample Problem G Solution Of the elements gallium, Ga, bromine, Br, and All of these elements are in the fourth period. calcium, Ca, which has the highest electronegativity? Explain your answer in terms of periodic trends. Bromine has the highest atomic number and is farthest to the right in the period. Bromine should have the highest electronegativity because electronegativity increases across the periods. Chapter menu Resources Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved. Copyright © by Holt, Rinehart and Winston. All rights reserved. End of Chapter 5 Show Chapter menu Resources Copyright © by Holt, Rinehart and Winston. All rights reserved.