Periodic Table Unit 4

 Discovery of Modern Elements –Antoine Lavoisier suggested that burning was actually a chemical combination with oxygen. –Lavoisier realized that there needed to be a new concept of elements, compounds, and chemical change. –We now know that there are 89 naturally-occurring elements and at least 23 short-lived and artificially prepared.

 Names of Elements –The first 103 elements have internationally accepted names, which are derived from:  The compound or substance in which the element was discovered  An unusual or identifying property of the element  Places, cities, and countries  Famous scientists  Greek mythology  Astronomical objects.

 The Periodic Law

 Dmitri Medeleev gave us a functional scheme with which to classify elements. –Mendeleev’s scheme was based on chemical properties of the elements. –It was noticed that the chemical properties of elements increased in a periodic manner. –The periodicity of the elements was demonstrated by Medeleev when he used the table to predict to occurrence and chemical properties of elements which had not yet been discovered.

 Mendeleev left blank spaces in his table when the properties of the elements above and below did not seem to match. The existence of unknown elements was predicted by Mendeleev on the basis of the blank spaces. When the unknown elements were discovered, it was found that Mendeleev had closely predicted the properties of the elements as well as their discovery.

 The Periodic Law –Similar physical and chemical properties recur periodically when the elements are listed in order of increasing atomic number.

 Introduction –The periodic table is made up of rows of elements and columns. –An element is identified by its chemical symbol. –The number above the symbol is the atomic number –The number below the symbol is the rounded atomic weight of the element. –A row is called a period –A column is called a group

 (A) Periods of the periodic table, and (B) groups of the periodic table.

 Periodic Patterns –The chemical behavior of elements is determined by its electron configuration –Energy levels are quantized so roughly correspond to layers of electrons around the nucleus. –A shell is all the electrons with the same value of n.  n is a row in the periodic table. –Each period begins with a new outer electron shell

–Each period ends with a completely filled outer shell that has the maximum number of electrons for that shell. –The number identifying the A families identifies the number of electrons in the outer shell, except helium –The outer shell electrons are responsible for chemical reactions. –Group A elements are called representative elements –Group B elements are called transition elements.

 Chemical “Families” –IA are called alkali metals because the react with water to from an alkaline solution –Group IIA are called the alkali earth metals because they are reactive, but not as reactive as Group IA.  They are also soft metals like Earth. –Group VIIA are the halogens  These need only one electron to fill their outer shell  They are very reactive. –Group VIIIA are the noble gases as they have completely filled outer shells  They are almost non reactive.

 Four chemical families of the periodic table: the alkali metals (IA), the alkaline earth metals (IIA), halogens (VII), and the noble gases (VIIIA).

Metal: Elements that are usually solids at room temperature. Most elements are metals. Non-Metal: Elements in the upper right corner of the periodic Table. Their chemical and physical properties are different from metals. Metalloid: Elements that lie on a diagonal line between the Metals and non-metals. Their chemical and physical properties are intermediate between the two.

Predicting Valence Electrons For the main group elements, the Group number indicates the number of valence electrons. Ion Electron Configurations When we write the electron configuration of a cation, we remove one electron for each positive charge: When we write the electron configuration of an anion, we add one electron for each negative charge: Na → Na+ 1s2 2s2 2p6 3s1 → 1s2 2s2 2p6 O → O2- 1s2 2s2 2p4 → 1s2 2s2 2p6

Isoelectronic Ions Isoelectronic atoms and ions are those which have the same number of electrons For example: – A Sodium ion has 10 electrons: – A Fluorine ion also has 10 electrons: Na → Na+ 1s2 2s2 2p6 3s1 → 1s2 2s2 2p6 F → F- 1s2 2s2 2p5 → 1s2 2s2 2p6

Atomic Radii  atomic radius is the distance from the nucleus to the outer electron shell  From left to right in a period, there is a repeating pattern of decreasing atomic radii.  In each period, metals have a larger radii than nonmetals.  The valence electrons of all elements in a period are basically in the same shell, but the # of protons increases, attracting the electrons to them stronger, therefore causing the radii to decrease.

Down a group  As you move down a group, atomic radii increases  As you move down a group each preceding element has more inner-level electrons, shielding the outside electrons (valence electrons) from the nucleus, allowing less attraction to the nucleus therefore making the radii increase.

Ionic Radius  Is the distance from the nucleus to the outer energy level of the ion.  As an atom loses (metal) it’s valence electron, it’s losing an energy level therefore it’s radius decreases in size.  When a nonmetal gains an electron, it becomes an octet resulting in an increased radius.

Atomic radius vs. atomic number

Trends in Ionization Energy  First Ionization Energy (IE) - The energy required to remove the first valence electron from an atom

Trends in Ionization Energy 6.3

Transition Metals  Groups 3-12  Form colored ions (have empty energy levels, when hit with light electrons get excited and can………………????  Multiple oxidation states (valence states)  All have similar properties. Typically hard solids with high melting points except Hg

Ionization Energy (IE) Periodic Trend  Periodic Trend for Ionization Energy (IE)  Down a Group:  Trend:  Reason:  Across a Period:  Trend:  Reason:

Trends in Electronegativity  Trends in Electronegativity  Electronegativity is the ability of an atom, in a chemical bond, to attract the shared valence electrons to itself. (i.e. the shared valence electrons are physically closer to the higher EN value atom than the other atom in the chemical bond).  What EN really means?  Fluorine? 6.3

Trends in Electronegativity  Representative Elements in Groups 1A through 7A 6.3

Electronegativity (EN) Periodic Trend  Periodic Trend for Electronegativity (EN)  Down a Group:  Trend:  Reason:  Across a Period:  Trend:  Reason:

Using ionic charge to determine the formula of an ionic compound  Ionic charges can be used to determine the formula of an ionic compound.  All compounds are neutrally charged. The total positive and total negative charges must balance out, they must add up to 0.  Remember, for the non-metals, they only form ions with charges equal to the first negative value.

 1. K and F  2. Ca and Br  3. Fe (III) and I