Trends on the Periodic Table Chemical Periodicity Trends on the Periodic Table

Development of the Periodic Table 70 elements discovered by the mid-1800s No system developed to organize them Dimitri Mendeleev put the elements on cards Grouped in columns by increasing atomic mass Arranged the columns by similarities in the elements

Published in 1869 Predicted the missing elements

Mendeleev Left blanks on the table Predicted the elements that would be there based on the trends Eventually the elements were discovered Found to have the properties predicted

The Modern Periodic Table Henry Moseley, British physicist (1900’s) Arranged the elements by atomic number instead of mass Properties change as you move left to right, from element to element Properties repeat in the next row. Similar chemical properties in groups

Periodic Law When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties Electron configurations in groups are similar in their valence electron configurations

Group A Representative Elements All have electrons in the s, or s and p orbitals Groups number indicates number of valence electrons Except Helium (group 8 with 2 electrons)

Group IA (group 1) Alkali Metals Soft, silvery color Too reactive to be found pure or free form in nature Hydrogen in NOT an alkali metal s1

Group IIA (group 2) Alkaline Earth Metals Harder, denser and stronger than alkali metals Higher melting points Less reactive than alkali metals Still too reactive to be found pure or free form in nature s2

Group VIII (group 18) Noble Gases 8 valence electrons (He =2) Inert (nonreactive), all gases Do not form ions Rarely in compounds Argon is most abundant (1% of air) s2p6

Group VIIA (group 17) Halogens Very reactive; react with metals to form salts Non-metals 7 valence electrons s2p5

Hydrogen Class by itself Because it contains just one proton and one electron Behaves unlike any other element Reacts with many other elements Most common element in the universe 1s1

Group B: Transition Metals Usually only two valence electrons Groups 3-12 Called the d block Can lose one, two, or three valence electrons depending on the element that it react with Form compounds and solutions that are brightly colored Some are relatively inactive; gold, silver, platinum

Inner Transition Elements Called the Rare Earth Metals F-block Lanthanide series; soft malleable metals with high lustre and conductivity

Metals

Metals Inner Transition Metals

Nonmetals

Metalloids Elements touching the stair step Boron is a metaloid Aluminum (acts more like a metal) Have properties of metals and non-metals Silicon and Germanium used for computer chips and solar cells

Metalloids

Classifying Elements by Electron Configuration Periodic table useful for identifying properties of elements Electron plays the most significant role in chemical reactions and properties of elements There is a relationship between electron configuration, placement on periodic table and chemical and physical properties

Four Classes Noble Gases Representative Elements (main group) Transition Elements Inner transition elements

Periodic Table Blocks

Periodic Trends Atomic Radius Electronegativity (Electron Affinity) Ionization Energy Density Melting Point Boiling Point

Concepts & Terms Electron Shielding Effective Nuclear Charge Increasing Principal Energy Level Lowest Energy Level (closest to the nucleus) Increasing mass (increasing atomic number) Increasing/decreasing volume (electron cloud)

Atomic Radius: Group Trends Increases as you move down a group Why? Increase in protons, but increase in principal energy levels (further away from the nucleus) Electron shielding by the electrons in the energy levels closed to the nucleus = reduced effective nuclear charge

Atomic Radius: Periodic Trends Atomic Radius DECREASES across a period Why? Principal Energy level remains the same (same electron shielding) Nucleus is becoming more positive thus increasing the effective nuclear charge

Atomic Radius

Trends in Ionization Energy When an atom gains or loses an electron, it becomes an ion. The energy required to overcome the attraction of the nuclear charge and remove an electron from a gaseous atom is called the ionization energy. Removing one electron results in the formation of a positive ion with a 1+ charge. Na(g) Na+(g) + e-

The energy required to remove this first outermost electron is called the first ionization energy. To remove the outermost electron from the gaseous 1+ ion requires an amount of energy called the second ionization energy, and so forth.

Trends in Ionization Energy

Ionization energies Large increase in ionization energy between the first and second ionization energies in sodium Easy to remove the first electron Much harder to remove the second electron Second electron is in the next lowest energy level Group 2A large increase for the third electron (change in energy level)

Ionization Energy: Group Trends Ionization Energy decreases as you move down a group Why? The size of the atom increases, energy level increases Outermost electron is farther from the nucleus Outermost electron in most easily removed.

Ionization Energy: Periodic Trends First Ionization Energy increases as you move across the rows from left to right Why? The nuclear charge increases with increasing numbers of protons The shielding effect is constant as the electrons are in the same principal energy level

Trends in First Ionization Energy

Cation vs. Anion

Trends in Electronegativity Electronegativity is the tendency for an atom to attract shared electrons to itself The atom becomes more negative Electronegativity decreases as you move down a group Electronegativity increases as you move across a row from left to right

The metallic elements at the far left of the periodic table have low electronegativities. By contrast, the nonmetallic elements at the far right (excluding the noble gases), have high electronegativities. The electronegativity of cesium, a metal, the least electronegative element, is 0.7; the electronegativity of fluorine, a nonmetal, the most electronegative element, is 4.0. Because fluorine has such a strong tendency to attract electrons, when it is chemically combined to any other element it either attracts the shared electrons or forms a negative ion. In contrast, cesium has the least tendency to attract electrons.

Electron Affinity

Periodic Trends

Periodic Table Puzzle Activity Key Hair = Valence Electrons Fingers = Protons/ Atomic Number Arms = Period/energy levels Pattern = group Size = Radius (in groups only, does not work for a period) Smile = electronegativity, electron affinity