Chapter 5 The Periodic Law

Vocabulary Section 1 Periodic Law Periodic Table Lanthanide Actinide Section 2 Alkali Metals Alkaline-Earth Metals Transition Metals Main-Group Elements Halogens Section 3 Atomic Radius Ion Ionization Ionization Energy Electron Affinity Cation Anion Valence electrons Electronegativity

History of the Periodic Table Mendeleev—arranged the elements in order of increasing atomic mass Created a table in which elements with similar properties were grouped together In this way, properties of undiscovered elements could be predicted Moseley—arranged atoms according to increasing atomic number Gave us the periodic law—the physical and chemical properties of the elements are periodic functions of their atomic numbers Produced a table in which elements with similar properties fell into the same column, or group.

Electron Configuration and the Periodic Table Section 2 Electron Configuration and the Periodic Table

S-Block Elements Group 1—alkali metals Group 2—alkaline earth metals Silvery appearance; Soft enough to cut with a knife Extremely reactive, found as compounds in nature Contains 1 valence electron (outermost energy level) Group 2—alkaline earth metals Harder than alkali metals Although less abundant, still found as compounds in the Earth’s crust Contains 2 valence electrons

P-Block Elements 6 Groups Boron Carbon Group 13 Always found combined with other elements in nature Contains 3 valence electrons Carbon Group 14 Group varies in many properties because? Contains 4 valence electrons Allotropes – forms of an element in the same state of matter, but have different physical and chemical properties Example: diamond and graphite are both made of carbon and are both solids, they differ in their properties because they have different structures – they are isotopes.

Allotropes of Carbon Example: diamond and graphite are both made of carbon and are both solids, they differ in their properties because they have different structures – they are isotopes.

P-Block Elements Continued Nitrogen Group 15 Properties vary as in the carbon group Contains 5 valence electrons Oxygen Group 16 Act as nonmetals, meaning they tend to gain electrons Contain 6 valence electrons

P-Block Elements Continued Halogens Group 17 Form compounds with metals (salts) A salt is a combination of a metal and a halogen Reactive nonmetals that tend to gain 1 electron Contains 7 valence electrons Noble Gases Group 18 Colorless Unreactive Contains 8 valence electrons – except helium which has 2 (filled orbitals)

d-Block Elements Known as Transition metals Share properties such as conductivity, luster, malleability, atomic size, electronegativity, ionization energy, melting points, and boiling points May have multiple oxidation states Tend to lose electrons to become positively charged ions

f-Block Elements: Inner Transition Metals Lanthanide series (4f) Silvery metals Similar reactivity like Group 2 (alkaline-earth metal) elements Actinide series (5f) Radioactive elements Only four exist in nature; the rest are synthetic elements that are created in particle accelerators

Bell Ringer For the following elements, name the block and group in which these elements are located on the periodic table; and whether high or low reactivity [He] 2s2 2p3 [Xe] 4f14 5d10 6s1 [Rn] 7s1 [Ar] 3d10 4s2 4p5

Section 3 Periodic Trends

Atomic Radii What is it? Period Trend Group Trend One-half the distance between the nuclei of identical atoms that are bonded together Period Trend Gradual decrease in atomic radii from left to right Electrons pulled closer by the increasing positive charge of the nucleus Example: Lithium is larger than Fluorine Group Trend Gradual increase down a group As electrons occupy sublevels in higher energy levels farther from the nucleus, the sizes of the atoms increase Positive charge of the nucleus is overridden by sheer number of electrons Iodine is larger than fluorine

Ionization Energy (IE) What is it? The energy required to remove one electron from the outer shell of an atom A + energy  A+ + e- Ion—an atom or group of elements that has a positive or negative charge Periodic Trend Ionization Energy increase from left to right across a period Caused by increasing positive charge of nucleus Example: It takes more energy to remove an electron from fluorine than from lithium Group Trend Decreases from top to bottom down groups Increase in distance from the nucleus = decrease in attraction between the nucleus and electrons Example: It takes less energy (easier) to remove an outer electron from Rubidium than from lithium

Removing electrons from positive ions Second ionization energy—energy required to remove additional electrons from positive ions. Requires much more energy than first ionization energy. With each additional electron removed, the effect of nuclear charge increases on the electrons, making them increasingly more difficult to remove from the outer shell.

Recall Rank in order of increasing atomic radius: Mn, Ca, Cu, Br Li, Rb, H, K Rank in order of increasing ionization energy: Al, P, S, Si Sb, P, As, N

Electron Affinity What is it? Period Trend Group Trend The difference in energy when an electron is added to a neutral atom. Most atoms release energy (F and Cl). A + e-  A- + energy Represented by a negative number. Some require energy to gain an electron (N). (must be forced) A + e- + energy  A- Represented by a positive number. Period Trend Increases from left to right (halogens gain electrons most readily) Caused by increase in effective nuclear charge; want to be noble gases. Group Trend Generally decreases down the group. Caused by dominant increasing atomic radius which decreases effective nuclear charge.

Ionic Radii What is it?? Periodic Trend Group Trend The radii of the most common ions of the elements. Positive ions—cations (loss of electrons) A+ Negative ions—anions (gain of electrons) A- Periodic Trend Cation radii—decrease from left to right Electron cloud shrinks due to increasing nuclear charge and the removal of higher- energy level electrons. Anion radii—increase from left to right Electron cloud increases due to a decrease in nuclear charge as electrons are added, and an increase in repulsion due to increased number of negatively-charged electrons. Group Trend Gradual increase in ionic radii as you go down a group.

Valence Electrons Chemical compounds are form because electrons are lost, gained or shared Only the outer electrons are involved in forming compounds Valence electrons – electrons available to be lost, gained or shared in the formation of chemical compounds Ex: an electron lost from the 3s sublevel of Na would form Na+

Fluorine is the most electronegative atom Electronegativity What is it?? The ability of an atom in a compound to attract electrons from another atom in the compound. Periodic Trend Increases from left to right Due to increasing effective nuclear charge Group Trend Decreases from top to bottom Fluorine is the most electronegative atom

http://www.rsc.org/periodic-table/trends