How To Understand The Periodic Table

Transcription

1 Periodic Table and Periodic Trends Periodic Table The first reasonably successful attempt to organize the elements was made by Dimitri Mendeleevin He had the idea of arranging elements in order of increasing atomic mass, and, most importantly, found that elements with similar chemical and physical properties occured periodically. He placed these similar elements under each other in columns. 1

2 In 1914, Henry Moseley determined that a better arrangement was in order of increasing atomic number, giving us the periodic table we have today. We can define the periodic table as an arrangement of elements in order of increasing atomic number placing those with similar chemical and physical properties in columns Periodic Table 2

3 Classification Three broad categories of elements called Metals Non-metals Metalloids To separate the metals and non-metals we draw a stairstep line to the left of and below B, Si, As, Te, and At Classification 3

4 Classification Metals solids at room temperature (except Hg) metallic luster malleable and ductile good conductors of heat and electricity Non-metals gases or solids at room temperature (except Br2) variety of color and appearance brittle solids insulators (poor conductors) Classification Metalloids intermediate in properties between metals and non-metals solids at room temperature many have more that one structure (one metallic, the other non-metallic) some are semi-conductors 4

5 Groups Vertical columns are called groups. Elements within a group have similar chemical and physical properties. Groups are designated at the top by the numbers 1-8 and by the letters A and B. (Note: group labeling is somewhat arbitrary, so watch out for other designations, particularly with A and B.) A group elements-representative or main group elements B group elements- Transition elements Periods Horizontal rows are called periods. Periods are designated by the numbers on the left in the periodic table. The two long rows placed just below the main body of the table are the inner transition elements. Elements are the Lanthanide Series Elements are the Actinide Series 5

6 Ionization Energy 6

7 Ionization Energy The ionization energy of an atom is the amount of energy requiredto remove an electron from the gaseous form of that atom or ion. 1 st ionization energy -The energy required to remove the highest energy electron from a neutral gaseous atom. For Example: Na (g) Na + (g) + e- I 1 = 496 kj/mole Ionization Energy Notice that the ionization energy is positive. This is because it requires energy to remove an electron. 1 st ionization energy decreases down a group. This is because the highest energy electrons are, on average, fartherfrom the nucleus. As the principal quantum number increases, the size of the orbital increases and the electron is easier to remove. 1 st ionization energy increases across a period. 7

8 Ionization Energy Electron Affinity 8

9 Electron Affinity Electron Affinity is the energy associated with the addition of an electon to a gaseous atom. Example: Cl (g) + e - Cl - (g) E.A. = -349 kj/mole Notice the sign on the energy is negative. This is because energy is usually releasedin this process, as apposed to ionization energy, which requires energy. A more negative electron affinity corresponds to a greaterattraction for an electron. (An unbound electron has an energy of zero.) Electron Affinity Electron affinity becomes less negative down a group. As the principal quantum number increases, the size of the orbital increases and the affinity for the electron is less. The change is small and there are many exceptions. Electron affinity decreases or increases across a period depending on electronic configuration. 9

10 Atomic Radius Atomic Radius In general the size of the atom depends on how far the outermost valence electron is from the nucleus. With this in mind we understand two general trends... Size increases down a group: The increasing principle quantum number of the valence orbitals means larger orbitals and an increase in atomic size. 10

11 Atomic Radius Size generally decreases across a period from left to right: To understand this trend it is first important to realize that the more strongly attracted the outermost valence electron is to the nucleus then the smaller the atom will be. Atomic Radius 11

12 Electron Negativity Electron Negativity Electronegativity-the ability of an atom in a molecule to attract electrons to itself. 12

13 Metallic Character Metallic Character Generally, the more metallic character an element has, the more basic its oxide will be. Likewise, the more non-metallic character an element has, the more acidic its oxide will be. The metallic character of an element can be determined by its position on the periodic table. Reference 2/6/

14 Electron Configurations Orbital information is often described in a notation called an electron configuration The electron configurationof an atomic species (neutral or ionic) allows us to understand the shape and energy of its electrons It gives us a better understanding of its bonding ability, magnetism and other chemical properties Orbitalsare definitely not orbits. They are electron cloudscharacterized by values for n, l, m l, and, m s. Orbitals are described as regions (energy levels) of high probability for finding an electron. 14

15 Electron Configurations Notation Represents the principal energy level Shows the number of electrons in the orbital 1s 2 Indicates the shape of the orbital The following table shows the possiblenumber of electrons that can occupy each orbital in a given subshell 15

16 Rules The Aufbau Principlestates that energy levels must be filled from the lowest to the highest and you may not move on to the next level unless the previous level is full. Use the periodic table as a guide Pauli Exclusion Principlestates that no two electrons (fermions) in an atom can have identical quantum numbers. Only two electrons may occupy any one orbital and they must have opposite spins Hund s Rule: every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin Exceptions to the rules Rules There are a few exceptions to the rules listed above when filling electron configurations. A half-full s orbital and a d subshell with 5 or 10 is more stable than following the Aufbau Principle. Cr, Mo, W:s 1 d 5 Cu, Ag, Au: s 1 d 10 16

17 Orbital Diagrams This information is also described using orbital diagrams. Arrows are added to an orbital diagram to show the distribution of electrons in the possible orbitals and the relative spin of each electron _ _ 1s 2s 2p Writing Electron Configurations Determine the number of electrons in the atom from its atomic number Following the rules, add electrons to the sublevels in correct order of occupation 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p s = 2, p = 6, d = 10, f =14 To check your complete electron configuration, look to see whether the location of the last electron added corresponds to the element s position on the periodic table Madelung s Rule 17

18 Drawing Orbital Diagrams Draw a line for each orbital of each sublevel mentioned in the complete electron configuration. Draw one line for each s sublevel, three for each p sublevel, 5 for each d sublevel and 7 for each f sublevel. Label each sublevel _ _ 1s For orbitals containing two electrons, draw one arrow up and one arrow down to indicate the electrons opposite spins. A_Multi-electron_Atoms/Electron_Configuration 18