Ch. 14: Chemical Periodicity Standard: Matter consists of atoms that have internal structures that dictate their chemical and physical behavior. Targets: Describe the arrangement of elements in the periodic table in order of increasing atomic number. Distinguish between the terms group and period. Apply the relationship between the electron arrangement of elements and their position in the periodic table. Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table. Discuss the similarities and differences in the chemical properties of elements in the same group. Describe and explain the group and periodic trends in atomic radii, first ionization energies and electronegativities.

Development of the Periodic Table Describe the arrangement of elements in the periodic table in order of increasing atomic number. Development of the Periodic Table Johan Dobereiner Grouped similar elements into groups of 3 (triads) such as chlorine, bromine, and iodine. (1817-1829). John Newlands Found every eighth element (arranged by atomic weight) showed similar properties. Law of Octaves (1863). Dmitri Mendeleev Arranged elements by similar properties but left blanks for undiscovered elements (1869).

Development of the Periodic Table Describe the arrangement of elements in the periodic table in order of increasing atomic number. Distinguish between the terms group and period. Development of the Periodic Table Henry Mosley Arranged the elements by increasing atomic number instead of mass (1913) Glen Seaborg Discovered the transuranium elements (93-102) and added the actinide and lanthanide series (1945) Elements are arranged by increasing atomic number into periods (rows) and groups or families (columns)

Arrangement of the Periodic Table Describe the arrangement of elements in the periodic table in order of increasing atomic number. Arrangement of the Periodic Table Metals Left side of the periodic table (except hydrogen). High electrical conductivity, high luster, ductile, malleable Alkali metals: Group 1 Alkaline earth metals: Group 2 Transition metals: Group 3-12, lanthanide & actinide series

Arrangement of the Periodic Table Describe the arrangement of elements in the periodic table in order of increasing atomic number. Arrangement of the Periodic Table Nonmetals Right side of the periodic table Poor conductors, nonlustrous, nonmalleable, dull Halogens: Group 17 Noble gases: Group 18

Arrangement of the Periodic Table Describe the arrangement of elements in the periodic table in order of increasing atomic number. Arrangement of the Periodic Table Metalloids Between metals and nonmetals Properties intermediate between metals and nonmetals

Arrangement of the Periodic Table: pg 392-393 Apply the relationship between the electron arrangement of elements and their position in the periodic table. Arrangement of the Periodic Table: pg 392-393 Noble Gases: Outermost s and p sublevels are filled. Group 18 Ending configuration is s2p6 (except He) Eight valence electrons (except He) Row number equals highest energy level See Figure 14.5 pg 395 Pg 396: practice 1,2; Pg 396: 3,4,5

Arrangement of the Periodic Table: pg 392-393 Apply the relationship between the electron arrangement of elements and their position in the periodic table. Arrangement of the Periodic Table: pg 392-393 Representative Elements: Outermost s and p sublevels are partially filled. Group 1-2 and 13-18 1 (s1); 2 (s2); 13 (s2p1); 14 (s2p2)… Group number equals valence electrons Row number equals highest energy level Transition Metals Filling the d & f sublevels Go to Chemistry: periodicty Trends Practice 1 See Figure 14.5 pg 395 Pg 396: practice 1,2; Pg 396: 3,4,5

Shortcut Electron Configuration Apply the relationship between the electron arrangement of elements and their position in the periodic table. Shortcut Electron Configuration Based on the electron configuration of the noble gases. He ends in 1s2; Ne ends in 2p6; Ar ends in 3p6; Kr ends in 4p6; etc. Write the electron configuration and orbital filling diagram for Se Se has 34 electrons Go back to the previous noble gas: Ar (18 electrons). Begin the configuration with [Ar] which accounts for 18 electrons and then begin with 4s2. Continue until you reach 34 electrons [Ar]4s23d104p4 [Ar] __ __ __ __ __ __ __ __ __ 4s 3d 4p

GO TO NOBLE GAS CONFIGURATION PRACTICE Handout Apply the relationship between the electron arrangement of elements and their position in the periodic table. Shortcut Electron Configuration Write the electron configuration and orbital filling diagram for Au Au has 79 electrons Go back to the previous noble gas: Xe (54 electrons). Begin the configuration with [Xe] which accounts for 54 electrons and then begin with 6s2. Continue until you reach 79 electrons [Xe]6s24f145d9 [Xe] __ __ __ __ __ __ __ __ __ __ __ __ __ 6s 4f 5d GO TO NOBLE GAS CONFIGURATION PRACTICE Handout

Shortcut Electron Configuration Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the periodic table. Shortcut Electron Configuration Electron dot diagrams Group 1: 1 dot X Group 15: 5 dots X Group 2: 2 dots X Group 16: 6 dots X Group 13: 3 dots X Group 17: 7 dots X Group 14: 4 dots X Group 18: 8 dots (except He) X GO TO WHITE BOARD LEWIS DOT PRACTICE and Noble gas configuration practice.

Trends in the Periodic Table GO TO PERIOD TRENDS GRAPH ACTIVITY GO TO REACTIVITY LAB

Periodic Trend Definitions Atomic Radius: half the internuclear distance between two atoms of the same element (pm)

Periodic Trend Definitions Electronegativity: a measure of the tendency of an atom in a molecule to attract a pair of shared electrons towards itself

trends are easier to understand if you comprehend the following the ability of an atom to “hang on to” or attract its valence electrons is the result of two opposing forces the attraction between the electron and the nucleus the repulsions between the electron in question and all the other electrons in the atom (often referred to the shielding effect) the net resulting force of these two is referred to effective nuclear charge

This is a simple, yet very good picture. Do you understand it?

3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Atomic Radii The radius of an atom, measured in pm (picometers) Periodic trend (Period 3 Trend) Atomic size decreases as you move across a period. The increase in nuclear charge increases the attraction to the outer shell so the outer energy level progressively becomes closer to the nucleus Group trend for Alkali metals & Halogens Atomic size increases as you move down a group of the periodic table. Adding higher energy levels

Atomic Radii

3.2.2 Describe and explain the trends in atomic radii, electronegativity's and reactivity. Tendency for the atoms of the element to attract electrons when they are chemically combined with atoms of another element. Helps predict the type of bonding (ionic/covalent). Periodic Trend (Period 3 Trend) Increases as you move from left to right across a period. Nonmetals have a greater attraction for electrons than metals & there is a greater nuclear charge that can attract electrons Group trend for Alkali metals & Halogens Generally decreases as you move down a group in the periodic table. For metals, the lower the inner nuclear force the more reactive. For nonmetals, the higher the inner nuclear force the more reactive.

Electronegativity

3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Reactivity The relative capacity of an atom, molecule or radical to undergo a chemical reaction with another atom, molecule or radical. Don’t worry about the periodic trend!!! Group trend for Alkali metals Increases as you move down group 1 in the periodic table Since alkali metals are more likely to lose an electron, the ones with the lowest inner nuclear force are the most reactive since they require the least amount of energy to lose a valence electron. Group trend for Halogens Decreases as you move down group 17 in the periodic table Since halogens are more likely to gain an electron, the ones with the greatest electronegativity are the most reactive since they are most effective at gaining a valence electron.

Discuss the similarities and differences in the chemical properties of elements in the same group. Group 1: Alkali Metals Have 1 valence electron Shiny, silvery, soft metals React with water & halogens Oxidize easily (lose electrons) Reactivity increases down the group Group 17: Halogens Have 7 valence electrons Colored gas (F2, Cl2); liquid (Br2); Solid (I2) Oxidizer (gain electrons) Reactivity decreases down the group