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.

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. ( ). 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).

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)

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 (1A) –Alkaline earth metals: Group 2 (2A) –Transition metals: Group B, lanthanide & actinide series

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 and nonlustrous –Halogens: Group 17 (7A) –Noble gases: Group 18 (0)

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

Apply the relationship between the electron arrangement of elements and their position in the periodic table. Arrangement of the Periodic Table: pg Noble Gases: Outermost s and p sublevels are filled. –Ending configuration is s 2 p 6 (except He) –Eight valence electrons (except He) –Row number equals highest energy level

Apply the relationship between the electron arrangement of elements and their position in the periodic table. Arrangement of the Periodic Table: pg Representative Elements: Outermost s and p sublevels are partially filled. –Group A elements (1A-7A) –1A (s 1 ); 2A (s 2 ); 3A (s 2 p 1 ); 4A (s 2 p 2 )… –Group number equals valence electrons –Row number equals highest energy level Transition Metals –Filling the d & f sublevels

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 1s 2 ; Ne ends in 2p 6 ; Ar ends in 3p 6 ; Kr ends in 4p 6 ; 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 4s 2. Continue until you reach 34 electrons –[Ar]4s 2 3d 10 4p 4 –[Ar] __ __ __ __ __ __ __ __ __ 4s 3d 4p

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 6s 2. Continue until you reach 79 electrons –[Xe]6s 2 4f 14 5d 9 –[Xe] __ __ __ __ __ __ __ __ __ __ __ __ __ 6s 4f 5d GO TO NOBLE GAS CONFIGURATION PRACTICE

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 1A: 1 dotXGroup 5A: 5 dotsX Group 2A: 2 dotsXGroup 6A: 6 dots X Group 3A: 3 dotsXGroup 7A: 7 dotsX Group 4A: 4 dotsXGroup 0: 8 dots (except He)X GO TO LEWIS DOT PRACTICE

Trends in the Periodic Table

13 Periodic Trend Definitions Atomic Radius: half the internuclear distance between two atoms of the same element (pm) Ionic radius: the radius of an ion in the crystalline form of a compound (pm)

14 Periodic Trend Definitions First ionization energy: The energy required to remove one electron from each atom in one mole of gaseous atoms under standard thermodynamic conditions (kJ mol -1 )

15 Periodic Trend Definitions Electronegativity: a measure of the tendency of an atom in a molecule to attract a pair of shared electrons towards itself Melting Point: the temperature at which a solid becomes a liquid at a fixed pressure (degrees Kelvin)

many 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?

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

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20 Atomic Radii

Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Ionic Radii The radius of the ion form of atoms (cations and anions) Positive ions are smaller than their atoms. –Fewer electrons so nucleus attracts remaining electrons more strongly –One fewer energy level since valence electrons removed. Negative ions are larger than their atoms –More electrons so nucleus has less attraction for them –Greater electron-electron repulsion Periodic trend (Period 3 Trend) –Decrease as you move across a period, then spike and decrease again –This 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 –Ions get larger down a group –More energy levels are added

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Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points First Ionization Energies The energy required to remove the first electron from a gaseous atom. Second ionization removes the second electron and so on. Can be used to predict ionic charges. Periodic Trend (Period 3 Trend) –Increases as you move from left to right across a period. –Effect of increasing nuclear charge makes it harder to remove an electron. Group trend for Alkali metals & Halogens –Generally decreases as you move down a group in the periodic table –Since size increases down a group, the outermost electron is farther away from the nucleus and is easier to remove.

25 Filled n=1 shell Filled n=2 shell Filled n=3 shell Filled n=4 shell Filled n=5 shell

Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points Electronegativity 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 number the more reactive. –For nonmetals, the higher the number the more reactive.

27 Electronegativity

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 1 st ionization energy 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 7 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 1A: Alkali Metals Have 1 valence electron Shiny, silvery, soft metals React with water & halogens Oxidize easily (lose electrons) Reactivity increases down the group Group 7A: Halogens Have 7 valence electrons Colored gas (F 2, Cl 2 ); liquid (Br 2 ); Solid (I 2 ) Oxidizer (gain electrons) Reactivity decreases down the group