The Period Table – basic info Dmitri Mendeleev noticed that when arranged in order of atomic mass, similar properties appeared at regular intervals Periodic Law: there is a periodic repetition in the Table  elements with the same physical and chemical properties end up in the same column.

The Period Table – basic info In today’s periodic table, the elements are ordered by atomic number, rather than atomic mass Groups (18 columns) elements that are in the same group have similar chemical properties, react similarly, and have the same number of valence electrons Periods ( 7 rows) elements that are in the same period have valence electrons in the same energy level

The Period Table – basic info Elements can be divided into three groups – metals, nonmetals, and metalloids

The Period Table – basic info METALS Majority of the elements Good conductors of heat and electricity Shiny/metallic Malleable – can be hammered into thin sheets Ductile – can be drawn into wires Except mercury, all are solid at room temperature

The Period Table – basic info NONMETALS Properties are more variable Poor conductors of heat and electricity Dull solids or powders Brittle – break easily Some solids (C, Si, P, S, Se, I), one liquid (Br), and some gases (H2, N2, O2, F2, Cl2, and noble gases) at room temperature

The Period Table – basic info METALLOIDS Run diagonally across the table between metals and nonmetals Dull-appearing, brittle solids Semiconductors – conduct electricity better than nonmetals, but not as well as metals Consist of B, Si, Ge, As, Sb, Te, and Po

Important groups to know Group 1 – the Alkali metals All very reactive, including with water Called “alkali” metals because when they react with water, they form a basic (alkaline) solution Group 2 – the Alkaline Earth metals Most are found in the Earth’s crust Not as reactive as the alkali metals, but still fairly reactive; combine readily with oxygen to form oxides (ex. magnesium)

Important groups to know Groups 3-12 – the Transition metals Most are metallic silver solids Group 17 – the Halogens All are diatomic when found in their elemental state (F2, Cl2, Br2, I2) All are very reactive Group 18 – the Noble gases All have filled p-orbitals (except helium) All are colorless, monoatomic gases All are very unreactive

Meet the elements!

Warm-up What do each of the following groups of elements have in common? Be, Ca, Sr, and Ra Na, Si, P, and Ar B, As, Sb, and Ge O, Ne, Xe, N, and H

Properties We will explore the periodic trends of three different properties: atomic radius, ionization energy, and electronegativity Atomic radius is the distance from the center of the nucleus to the outermost electron (this is inherently vague)

Properties Ionization energy is the energy required to take a valence electron and remove it from the atom (in other words, to excite it to the  energy level); “ionization energy” is literally the amount of energy required to make an ion Electronegativity is a measure of how much an atom pulls on the electrons in a bond (or, in other words, how badly the atom wants electrons)

Visualize the Trend – Atomic Radius What happens to the atomic radius as you travel across a period? What happens to the atomic radius as you travel down a group?

Visualize the Trend – Ionization Energy What happens to the ionization energy as you travel across a period? What happens to the ionization energy as you travel down a group?

Visualize the Trend – Electronegativity What happens to the electronegativity as you travel across a period? What happens to the electronegativity as you travel down a group?

WHY?

Nuclear charge increases Valence electrons increase within the same energy level Same number of valence electrons, but in higher energy levels that are further away Nuclear charge increases

Atomic Radius As you go across a period, atomic radius decreases ex: Valance electrons are added into the same energy level, but nuclear charge increases, pulling the n=2 shell closer to the nucleus e e e e e e +3 e +4 e +5 e e e e

Atomic Radius As you go down a group, atomic radius increases ex: Valence electrons are added to higher energy levels, which are further from the nucleus e e e e e e e +1 e +3 e e +11 e e e e e

Ionization Energy As you go across a period, ionization energy increases (meaning that it is more difficult to remove an electron) ex: Valence electrons are in the same energy level, but become harder to remove due to increasing positive charge in the nucleus e e e e e e +3 e +4 e +5 e e e e

Ionization Energy As you go down a group, ionization energy decreases (meaning that it is easier to remove an electron) ex: Valence electrons are added to higher energy levels, which are further from the nucleus, making the electrons easier to remove e e e e +1 e e e e +3 e +11 e e e e e e

Electronegativity As you go across a period, electronegativity increases (meaning that atoms pull electrons toward themselves more) ex: Valence electrons are in the same energy level, distance between the atom and what it is bonded to will be the same; atoms with greater nuclear charge will pull electrons toward themselves more e e e e e e +3 e +4 e +5 e e e e

Electronegativity As you go down a group, electronegativity decreases (meaning that atoms pull electrons toward themselves less) ex: Valence electrons are added to higher energy levels, meaning that the distance between the two bonded atoms will be longer; atoms with longer bonds are less able to pull electrons toward themselves e e e e +1 e e e e +3 e +11 e e e e e e

Ions and Ionic Radius

What is an ion? Atoms can gain or lose electrons from their outer energy levels, forming an ion

Anions If an atom gains electrons, because it has more electrons (negatives) than protons (positives), it will have a net negative charge- These are called Anions Nonmetals form anions because by gaining a couple more electrons they will have a full valence shell of 8 electrons

Cations If an atom loses electrons, because it has less electrons than protons, it will have a net positive charge These are called cations Metals form cations because they have low ionization energies and by losing electrons they have a full inner shell of electrons

Practice What atom is this? Will it lose or gain electrons? How many will it gain/lose? What will its charge be?

More Examples Na Be F Determine how many valence electrons each of the above elements have. Determine if they are a metal or a nonmetal and if they will therefore become a cation or an anion. Determine what ion they will form

More on ions Using the charge and symbol of the ion, you should be able to determine the total number of electrons How many electrons do each of the following have? N3- Li+ Fe3+ Cl-

Ionic Radius When an atom gains or loses electrons, its radius changes If an atom loses electrons, because it usually loses an energy level, its radius decreases Example: Na+

For Cations (which have LOST electron(s)): Atomic radius decreases, so the protons pull the fewer remaining electrons closer to the nucleus For Anions (which have GAINED electron(s)): Atomic radius increases, because the attraction of the protons to any one electron decreases.

 Usually the only electrons that react!!!! Electron dot structures are diagrams that show valence electrons in the atoms of an element as dots.

Electron Dot Structures of Some Group A Elements Period Group 1A 2A 3A 4A 5A 6A 7A 8A 1 2 3 4 This table shows electron dot structures for atoms of some Group A elements. Notice that all the electrons within a given group (with the exception of helium) have the same number of electron dots in their structures. . \

Valence Electrons The Octet Rule The octet rule states that in forming compounds, atoms tend to achieve the electron configuration of a noble gas. An octet is a set of eight. Atoms of each of the noble gases (except helium) have eight electrons in their highest occupied energy levels and the general electron configuration of ns2np6. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved.

Draw an Electron Dot Structure for Sodium (Na) Draw the unabbreviated electron configurations for Sodium

Formation of Cations Group 1A Cations When forming a compound, a sodium atom loses its one valence electron and is left with an octet in what is now its highest occupied energy level. The number of protons in the sodium nucleus is still eleven, so the loss of one unit of negative charge produces a cation with a charge of 1+. e– Na 1s22s22p63s1 Na+ 1s22s22p6 octet

Formation of Cations Group 2A Cations Magnesium (atomic number 12) belongs to Group 2A of the periodic table, so magnesium atoms have two valence electrons. A magnesium atom attains the electron configuration of a neon atom by losing both valence electrons and producing a magnesium cation with a charge of 2+. • Mg Mg2+ + 2e– • loses all its valence electrons Magnesium atom (electrically neutral, charge = 0) Magnesium ion (2+ indicates two units of positive charge) (2 in front of e– indicates two units of negative charge)

Formation of Cations The figure at right lists the symbols of the cations formed by metals in Groups 1A and 2A. Cations of Group 1A elements always have a charge of 1+. Cations of Group 2A elements always have a charge of 2+.

Formation of Anions Atoms of chlorine have seven valence electrons. Atoms of nonmetallic elements attain noble-gas electron configurations more easily by gaining electrons than by losing them because these atoms have relatively full valence shells. Atoms of chlorine have seven valence electrons. A gain of one electron gives a chlorine atom an octet and converts a chlorine atom into a chloride atom. Cl 1s22s22p63s23p5 Cl– 1s22s22p63s23p6 +e– octet

Formation of Anions Chlorine atoms need one more valence electron to achieve the electron configuration of the nearest noble gas.