Presentation on theme: "Chapter 5 Periodic Table Mendeleev noticed that when the elements were arranged in order of increasing atomic mass, certain similarities in their chemical."— Presentation transcript:
Chapter 5 Periodic Table
Mendeleev noticed that when the elements were arranged in order of increasing atomic mass, certain similarities in their chemical properties appeared at regular intervals. Repeating patterns are referred to as periodic. Mendeleev created a table in which elements with similar properties were grouped together—a periodic table of the elements. Mendeleev and Chemical Periodicity
Mendeleev’s Periodic Table Dmitri Mendeleev Refer to pg 124 Text – For Elements out of Place.
Properties of Some Elements Predicted By Mendeleev
Moseley and the Periodic Law In 1911, the English scientist 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. What does this mean?
Periodicity of Atomic Numbers
Noble Gases – Group 18 or VIII – Unreactive gases with eight valence electrons. Lanthanides – Atomic # added to the periodic table in Actinides – Atomic #
The Modern Periodic Table The Periodic Table is an arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same column, or group. Period or series - A horizontal row across the periodic table (7 total) Group or family – A vertical column on the periodic table (18 total)
A Spiral Periodic Table
The Periodic Table Period Group or family Period Group or Family
Never found pure in nature Easily lose valence electron React violently with water In their pure states,they have a soft silvery appearance and can be cut with a knife React with halogens to form salts Group configuration = ns 1 n arrangement of the elements in The Properties of Group I: the Alkali Metals
Group 2 of the periodic table 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. Group configuration ns2 Radium was widely used in self-luminous clock and watch hands, until too many watch factory workers had died of it. This antique watch is still quite radioactive and will remain for thousands of years.
Halogens – Group 17 or VII React vigorously with most metals to form salts. Fluorine and chlorine are gases at room temp., bromine is a reddish liquid, and iodine is a dark purple solid. Astantine is a synthetic element prepared in very small quantities. Group configuration ns 2 np 5
Transition Metals “Group B Metals” Many exceptions to Aufbau exist in this area. Why? Why was the mad hatter mad?
Properties of Metals Metals are good conductors of heat and electricity Metals are malleable Metals are ductile Metals have high tensile strength Metals have luster Located to the left of the zigzag line on the periodic table
Examples of Metals Potassium, K reacts with water and must be stored in kerosene Zinc, Zn, is more stable than potassium Copper, Cu, is a relatively soft metal, and a very good electrical conductor. Mercury, Hg, is the only metal that exists as a liquid at room temperature
Properties of Nonmetals Carbon, the graphite in “pencil lead” is a great example of a nonmetallic element. Nonmetals are poor conductors of heat and electricity Nonmetals tend to be brittle Many nonmetals are gases at room temperature Located to the right of the zigzag line on the periodic table
Examples of Nonmetals Sulfur, S, was once known as “brimstone” Microspheres of phosphorus, P, a reactive nonmetal Graphite is not the only pure form of carbon, C. Diamond is also carbon; the color comes from impurities caught within the crystal structure
Properties of Metalloids Metalloids straddle the border between metals and nonmetals on the periodic table. They have properties of both metals and nonmetals. Metalloids are more brittle than metals, less brittle than most nonmetallic solids Metalloids are semiconductors of electricity Some metalloids possess metallic luster
Silicon, Si – A Metalloid Silicon has metallic luster Silicon is brittle like a nonmetal Silicon is a semiconductor of electricity Other metalloids include: Boron, B Germanium, Ge Arsenic, As Antimony, Sb Tellurium, Te
f bloc k Blocks of the Periodic Table The s and p blocks = the main group elements. The d block = the transition metals. The f block = the inner transition metals. d bd lock s block S P f
Sample Problem A a. Without looking at the periodic table, identify the group, period, and block in which the element that has the electron configuration [Xe]6s 2 is located. b. Without looking at the periodic table, write the electron configuration for the Group 1 element in the third period. Is this element likely to be more reactive or less reactive than the element described in (a)? Periods and Blocks of the Periodic Table
Solution for Sample Problem A It is in the sixth period, as indicated by the highest principal quantum number in its configuration, 6. The element is in the s block. The element is in Group 2, as indicated by the group configuration of ns 2.
Sample Problem B An element has the electron configuration [Kr]4d 5 5s 1. Without looking at the periodic table, identify the period, block, and group in which this element is located. Then, consult the periodic table to identify this element and the others in its group.
Sample Problem B Solution The number of the highest occupied energy level is 5, so the element is in the fifth period. There are five electrons in the d sublevel, which means that it is incompletely filled. The d sublevel can hold 10 electrons. Therefore, the element is in the d block. For d-block elements, the number of electrons in the ns sublevel (1) plus the number of electrons in the (n 1)d sublevel (5) equals the group number, 6. This Group 6 element is molybdenum.
Sample Problem C Without looking at the periodic table, write the outer electron configuration for the Group 14 element in the second period. Then, name the element, and identify it as a metal, nonmetal, or metalloid.
Sample Problem C Solution The group number is higher than 12, so the element is in the p block. The total number of electrons in the highest occupied s and p sublevels is therefore equal to the group number minus 10 (14 10 = 4). Two electrons are in the s sublevel, so two electrons must also be present in the 2p sublevel. The outer electron configuration is 2s 2 2p 2. The element is carbon, C, which is a nonmetal.
Sample Problem D Name the block and group in which each of the following elements is located in the periodic table. Then, use the periodic table to name each element. Identify each element as a metal, nonmetal, or metalloid. Finally, describe whether each element has high reactivity or low reactivity. a. [Xe]4f 14 5d 9 6s 1 c. [Ne]3s 2 3p 6 b. [Ne]3s 2 3p 5 d. [Xe]4f 6 6s 2
a.The 4f sublevel is filled with 14 electrons. The 5d sublevel is partially filled with nine electrons. Therefore, this element is in the d block. The element is the transition metal platinum, Pt, which is in Group 10 and has a low reactivity. b. The incompletely filled p sublevel shows that this element is in the p block. A total of seven electrons are in the ns and np sublevels, so this element is in Group 17, the halogens. The element is chlorine, Cl, and is highly reactive. Sample Problem D Solution
c. This element has a noble-gas configuration and thus is in Group 18 in the p block. The element is argon, Ar, which is an unreactive nonmetal and a noble gas. d. The incomplete 4f sublevel shows that the element is in the f block and is a lanthanide. Group numbers are not assigned to the f block. The element is samarium, Sm. All of the lanthanides are reactive metals. Sample Problem D Solution
Summation of Periodic Trends
Half of the distance between nuclei in covalently bonded diatomic molecule "covalent atomic radii” Periodic Trends in Atomic Radius Radius decreases across a period Increased effective nuclear charge due to decreased shielding Radius increases down a group Addition of principal quantum levels Determination of Atomic Radius
Table of Atomic Radii
Compare two cations Lithium configuration Lithium ion configuration
Compare two anions Oxygen configuration Oxygen ion configuration
Increases for successive electrons taken from the same atom Tends to increase across a period Electrons in the same quantum level do not shield as effectively as electrons in inner levels Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove Tends to decrease down a group Outer electrons are farther from the nucleus Ionization Energy - the energy required to remove an electron from an atom
Table of 1 st Ionization Energies
Electron Affinity - The energy change that occurs when an electron is acquired by a neutral atom. Most atoms release energy when they acquire an electron
Electronegativity A measure of the ability of an atom in a chemical compound to attract electrons Electronegativities tend to increase across a period Electronegativities tend to decrease down a group or remain the same
Periodic Table of Electronegativities
Ionic Radii Cations Positively charged ions formed when an atom of a metal loses one or more electrons Smaller than the corresponding atom Anions Negatively charged ions formed when nonmetallic atoms gain one or more electrons Larger than the corresponding atom