6.1 Development of the Modern Periodic Table Objectives: 1. Describe the major advancements in development of the periodic table 2. Describe the organization of the elements on the periodic table 3. Classify elements by group, period, block, metallic characteristics, family name, phase, series name 4. Characterize metals, nonmetals and metalloids. History of the Periodic Table s Development 18 th Century 23 elements 19 th Century 70 elements John Newlands English Chemist (1837-1898) 1864 Newlands arranged elements by atomic mass and noticed that their properties repeated every 8 elements (Law of Octaves) Meyer, Mendeleev, & Moseley Lothar Meyer and Dmitri Mendeleev related chemical properties to the atomic masses of elements. Mendeleev organized his table of elements by atomic mass and arranged elements with similar properties so that they lined up in columns. This was the first periodic table. Mendeleev s periodic table left blanks where he predicted elements were yet to be discovered (Sc, Ga, Ge) He also predicted the properties of those elements. Henry Moseley English Chemist (1887-1915) Working in Rutherford s lab he correlates the positive charges from the nucleus with an integer resulting in the atomic number. Determines that the order of the elements should be based on this fundamental property rather than the property of mass
Periodic Law When elements are arranged in order based on increasing atomic number, their physical and chemical properties show a periodic repetition. The Modern Periodic Table Reading the Periodic Table Different periodic tables include different information. Typically, you find: Atomic number (an integer) Atomic symbol (case matters) Atomic mass (to a certain number of sig figs)
Periodic Nomenclature Columns are called groups or families o 18 columns in standard periodic table o Traditionally numbered I-VIII, followed by A or B o Modern number system is 1-18 Rows area called periods o 7 periods in standard periodic table
The two rows at the bottom of the periodic table are called the inner transition elements or metals. The inner transition metals include the lanthanide series (4f) and actinide series (5f). Main group or representative elements o Elements whose number ends with an A o The s and p blocks. o o Transition elements o Elements whose number ends with B o The d block Inner transition elements o The elements at the bottom of the periodic table o The f block Metals, Non-metals and Metalloids The Stair Step o To the right of the P.T. there is a heavy stair step line that gives us a good idea if an element is metallic, non-metallic or semimetallic.
Metals o Elements to the left of the stair step are metals o Largest group of elements o Typically solid at room temperature o Have luster o Good conductors of heat and electricity o Malleable and ductile Non-Metals o Elements to the right of the stair step are non-metals o Most are gas at room temperature o No metallic luster o Poor conductors of heat & electricity o Neither malleable nor ductile Metalloids/Semi-metals o Elements along the stair step line are metalloids o Share properties of metals and non-metals. o Have metallic properties o Include semi-conductors Special Group Names Alkali metals (group IA) Alkaline-earth metals (group IIA) Halogens (group VIIA) Noble Gases (group VIIIA) Other families are named based on the element that appears at the top of the column (e.g. Group 14 (4A) = carbon family) Hydrogen is often placed by itself, in Group 1A, Group 17 or both 1 & 17 because of its unique properties. 6.2 Classification of the Elements Objectives: 1. Determine the valence electrons for elements 2. Describe general properties of elements in each block Organizing the Elements by Electron Configuration Valence Electrons Electrons in the outermost energy level of an atom Determine the chemical properties of the element Elements are grouped on the periodic table based on having the same number of valence electrons
Li 1s 2 2s 1 Na 1s 2 2s 2 2p 6 3s 1 K 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 Rb 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 5s 1 Cs 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 5s 2 5p 6 6s 1 There may be no more than 8 valence electrons for any one atom Valence Electrons & Period The period number of an element indicates the energy level of the valence electrons for that element Valence Electrons and Group Number The Roman numerals on the periodic table for the main-group elements show the number of valence electrons available for bonding The number of valence electrons for the transition elements is technically 2, but in reality the electrons from the lower s sublevel mingle with the d sublevel electrons to create a variety of bonding possibilities. The s-, p-, d-, & f-block Elements s-block elements Groups IA & IIA (plus Helium) Will have either 1 or 2 valence electrons p-block elements Groups IIIA XIIIA Will have 3 8 valence electrons Includes the noble gases o Incredibly stable because the have a maximum number of valence electrons o Rarely react chemically (larger noble gases may be forced to react) d-block elements Transition metals Will usually have 2 electrons in outermost level (s-sublevel) and a partially filled set of d orbitals Have similar chemical & physical properties f-block elements Inner transition metals Will usually have 2 electrons in outermost level (s-sublevel) and a partially filled set of f orbitals Have similar chemical & physical properties d- and f- sublevels Have many different possible distributions of electrons as these large sublevels shift electrons among similar energy sublevels and many orbitals
6.3 Periodic Trends Objectives: 1. Describe the nature of periodic trends 2. Define atomic radius, ion radius, ionization energy, electronegativity and electron 3. Relate ionization energy to electron configurations (box orbital really) 4. Compare elements by atomic radius, ion radius, ionization energy, electronegativity and electron Certain properties of elements change in a predictable way as you move through the periodic table. These predictable changes are called periodic trends. Periodic trends relate to the attraction of the nucleus to electrons. Atomic Radius Atomic radius conceptually distance from the center of the nucleus to the outermost electron pragmatically distance between nuclei of adjacent atoms Rules for estimating relative atomic radii Atoms get larger down a group. The principal quantum number increases which means that higher and higher energy levels are being used to store electrons. Atoms get smaller moving left to right across the periodic table. Moving across a period, additional electrons occupy the same principal energy level (they don t get farther from the center), but because additional protons are present in
the nucleus, there is greater pull from the positively charged nucleus on the negatively charged electrons Ionic Radius Ionic radius size of an ion of a particular element Rules for estimating relative ionic sizes When an atom loses electrons, it becomes smaller. Electrons are shed from the outermost energy level making it smaller. Since the size of the negative charge decreases, the positively charged nucleus can pull harder on the electrons and get them closer to the nucleus. When an atom gains electrons, it becomes larger. Additional electrons are added at the fringe of the atom. Repulsive forces between additional electrons make them spread out. Attraction from the positively charged nucleus is spread thinner to more electrons and is not as strong. Ionization Energy Ionization energy energy required to strip off electrons (create an ion) Basically ionization energy is how well an atom holds its electrons This is directly related to atomic radius. The greater the atomic radius, the less attraction and therefore the less energy required to strip electrons off.
First ionization energy energy required to strip off highest energy electron Ionization energy is measured in joules/atom or joules/mole Rules for estimating relative first ionization energies Ionization energy increases as you move left to right across the periodic table Ionization energy increases as you move up the periodic table Successive ionization energies Successive ionization energies are the energies required to strip off successive electrons (i.e., energy to strip 2 nd, 3 rd, etc.) These successive ionization energies have patterns that tell us how stable certain electron configurations are. Noble gas configurations are very stable. When there are as many electrons as in a noble gas, stripping one off will be very difficult. (High ionization energy.) Electronegativity Electronegativity ability to attract electrons in a chemical bond. Atoms with high electronegativity will tend to pull on electrons harder in a chemical bond. When electrons are pulled to one side, that side develops a more negative charge. The other side develops a more positive charge.
Summary of Periodic Trends Trend/Property Across a Period (L R) Down a Group Atomic Number Increases Increases Atomic Mass Increases Increases Atomic Radius Decreases Increases Ionic Radius Depends Increases Ionization Energy Increases Decreases Electronegativity Increases Decreases