Unit 4: Periodic Table

Classifying the Elements  Demitri Mendeleev (1869)- Russian Scientist Concluded that physical and chemical properties of elements appear in regular intervals when listed by increasing atomic mass Not current classification

Classifying the Elements HHenry Moseley (1913) –English Scientist Used x-rays to identify atomic number He concluded that physical and chemical properties of elements were listed by increasing atomic # Modern classification system

Modern Periodic Law CChemical similarities of elements are due to the # of valence electrons AAll elements in the same group have the same # of valence electrons, therefore they have the same properties

Arrangement of the Table 11. Periods (horizontal rows) Labeled on the left hand side of the periodic table TThere are currently 7 periods The number of the period IIndicates the number of energy levels IIndicates in which energy level the valence electrons are found Ex. Elements in period 3 have electrons in 3 energy levels with the valence electrons in the 3 rd energy level TThe number of valence electron increases from left to right Properties of elements change as you move from left to right (different # of valence electrons)

22. Groups (vertical columns/families) 1. Group 1- Alkali metals 2. Group 2- Alkaline earth metals 3. Groups 3-12 transition metals (form colored solutions) 4. Group 17-Halogens 5. Group 18-Noble (inert) gases EElements in groups show similar properties because they have the same number of valence electrons

Properties of Elements AAtomic Radii (size of an atom) 1. Trends in a period RRadius decreases from left to right IIn each period, metals have larger radii than nonmetals TThe # of valence electrons and protons increase from left to right (# of energy levels stay the same) An increase in the # of protons (nuclear charge) results in a greater attraction for valence electrons This creates a smaller radius

Atomic Radius 2. Trends in a Group Radius increases from top to bottom As you go down a group tthe number of protons increase bbut the # of inner energy levels also increases These inner energy levels shield the nucleus from the valence electrons RReduces attractive forces between protons and valence electrons IIncreases the size of the radius

Ionic Radius AAtoms can lose or gain electrons to become ions TThis result causes a change in the size of the atom Metals tend to lose their valence and become positive ions tthe radius decreases in size MMetal atoms have a larger radius than their corresponding ion Nonmetals tend to gain electrons and become negative ions TThe radius increases in size NNonmetal atoms have a smaller radius than their corresponding ions

Na vs Na +1

Cl vs Cl -1

Electronegativity  The ability of an atom to attract electrons when bonded to another atom Range from 0.7 (Francium) to 4.0 (Fluorine) The higher the value, the stronger the pull for electrons Noble gases usually do not have values since they do not want to gain or lose electrons  Very stable on their own (monoatomic)  Form very few stable compounds

Electronegativity Trends PPeriod Electronegativity increases from left to right Metals tend to have low EN values Nonmetals tend to have high EN values GGroup EN values decrease down a group The highest EN values are found at the top

Ionization Energy  The amount of energy needed to remove the most loosely bound electron from an atom  Atoms with more than one electron have more than one ionization energy (regents exam only focuses on first ionization energy)

Trends in Ionization Energy PPeriod Ionization energy increase from left to right BBecause the nuclear charge is increasing and valence electrons are more strongly attracted, more energy is needed to remove them from an atom Group IIonization energy decreases down a group Because the number of energy levels increase, valence electrons are farther from the nucleus It is easier to remove electrons when they are farther away from the positive charge of the nucleus (shielding)

Reactivity  Some elements can be found uncombined in nature Example: Noble gases (monoatomic)  Other elements are so reactive, they are never found uncombined in nature Example: elements of groups 1,2, and 17  Atoms react when they gain, lose, or share electrons Groups 1 and 2 contain most reactive metals, since they lose electrons easily Group 17 contains most reactive nonmetals since they gain electrons easily

Metallic vs. Nonmetallic Characteristics  Metallic properties Increase down a group Ex. Group 14 from top to bottom nonmetals-> semi-metals-> metals Decrease across a period metals-> semi-metals-> nonmetals

Staircase TThe thickened black line (staircase) separates metals (left of line) from nonmetals (right of line) SSemi-metals (metalloids) are found adjacent to the staircase B-Boron, Si-Silicon, Ge-Germanium, As- Arsenic, Sb-Antimony, Te-Tellurium

Metals  Make up about 75 percent of all elements  The most reactive metallic element is Fr in the lower left corner of the table  All are solids at STP with the exception of Hg (mercury – liquid metal)

Metals  Metallic properties 1. Low ionization energy Easy to pull off electrons form positive ions 2. Low electronegativity Do not attract electrons easily 3. Smaller ionic radii than atomic radii due to loss of electrons 4. Malleable Hammered into shape Metallic Properties 5. Ductile Drawn into wire 6. Good conductors of heat and electricity 7. Have luster (shine)

Nonmetals  Can exist in the solid, liquid, and gas phase at STP 1. H 2,O 2, F 2, N 2, Cl 2, are diatomic gases 2. Group 18 elements are all gases 3. Br 2 is diatomic and the only liquid nonmetal at STP 4. All other nonmetals are solids including I 2 (diatomic)  Group 17 contains the most reactive nonmetals (upper right)  Fluorine individually is the most reactive nonmetal

Nonmetallic Properties  High ionization energy 1. Hard to pull off electrons 2. Usually gain to become negatively charged High electronegativity 1. Attract electrons easily Larger ionic radii than atomic radii due to gain of electrons Brittle Poor conductors of heat and electricity Lack luster

Semi-Metallic Properties  Exhibit properties of both metals and nonmetals  Can be brittle  Semi- conductors of heat and electricity