The Periodic Table

1. The Periodic Table Unit 3

2. Origins • In 1817, Johann Dobereiner found that the properties of the metals calcium, barium and strontium were very similar. He grouped these elements in a triad.

3. Origins (Cont.) • In 1863, John Newlands arranged the elements in order of their increasing atomic masses. He noted that there appeared to be a repetition of similar properties every eighth element. • The Law of Octaves shows the same properties repeated every eight elements.

4. Dimitri Mendeleev • In 1860’s, Dimitri Mendeleev studied elements and also believed that the properties of the elements were a function of their atomic masses. • However, he believed that similar properties occurred after periods (horizontal rows) that could vary in length.

5. Mendeleev (Cont.) • In 1869 he published his first periodic table, leaving some blank spots in order to group all the elements with similar properties in the same column. • He explained the blank spots were for elements not yet discovered, but in 1871 he predicted their atomic mass and properties.

6. The Periodic Law • His predictions were very nearly correct. Mendeleev stated that the properties of the elements are a periodic function of their atomic masses. This is called the periodic law. • For his work, he is known as the Father of the Periodic Table.

7. The Modern Periodic Law • There was a problem with this arrangement. Henry Moseley found the reason for these exceptions to Mendeleev’s periodic law. • Moseley’s X-ray experiments showed that elements were positioned better if done by their atomic number. • The modern statement of the periodic law is the properties of the elements are a periodic function of their atomic numbers.

8. The Modern Periodic Table • A horizontal row is referred to as a period. • A vertical column is referred to as a group or family. • There are 18 groups in the periodic table.

9. Stability of Atoms • One of the primary rules in chemistry is that atoms with a full outer level are particularly stable (less reactive). • For all such elements, except hydrogen & helium, the outer level contains 8 electrons, these outer electrons are called an octet. • The fact that eight electrons in the outer level render an atom unreactive is called the octet rule.

11. Characteristics of Metals: • They tend to lose electrons to become positive ions. • Hard and shiny • Conduct heat and electricity • Malleable & ductile • They are to the left of the stair step. • Have only a few electrons in their outer level Gold & Copper

12. Characteristics of Non-metals: • Gases or brittle solids at room temperature • Dull and they are insulators • Are to the right of the stair step • Have more electrons in their outer level • Can share their electrons or gain electrons to become negative ions. Carbon & Sulfur

13. Metalloids • They are elements right along the stair steps on the chart. (Except Al) • Metalloids share characteristics of both metals and non-metals. • For example, Silicon is shiny like a metal, but is brittle and is not a good conductor of electricity.

14. Column IA: Alkali Metals • Alkali metals have a silvery appearance and are soft enough to be cut with a knife. • They are extremely reactive with moisture and are never found as free elements in nature. Potassium in water

15. Column IIA: Alkaline Earth Metals • Alkaline earth metals are harder, denser and stronger than alkali metals. • They are less reactive but are never found as free elements in nature. • Pure Calcium is too reactive to be free in nature but can be found in Marble, Calcium Carbonate (CaCO3).

16. Transition Metals • They occupy columns IIIB – IIB. • They are typically less reactive than alkali or alkaline earth metals. • Some are so unreactive they do not easily form compounds and exist in nature as free elements. • Palladium, Platinum and Gold are transition metals.

17. The halogens are the most reactive nonmetals. They react vigorously with most metals to form “salts”. Fluorine and chlorine are gases at room temperature. Bromine is a reddish liquid. Iodine is a dark purple solid. Column VIIA: Halogens

18. Physical Properties of Halogens

19. Column VIIIA: Noble Gases • First discovered in 1894 when Argon was discovered. • Noble gases were undetected due to their total lack of chemical reactivity. • They have full valence shells, 8 valence electrons, with the exception of Helium which has 2.

20. The Lanthanides and Actinides • Lanthanides contain elements from #58, Cerium, to 71, Lutetium in the top pullout row. • Actinides contain elements from #90, Thorium, to 103, Lawrencium in the bottom pullout row. • Both series are off of group IIIB and are in periods 6 & 7 respectively and are metals.

21. Practice • WS #1 – Blank Ptable • WS #1 Make your Own Periodic Table

22. Periodicity • Periodicity of properties can be observed in any group on the periodic table. • The reason for periodicity is explained in the arrangement of electrons around the nucleus.

23. Periodicity • There are trends in these properties as you go down in a group or across in a period. • NOTE: When discussing trends in properties, the Noble gases are never included.

24. Atomic Radius • It measures the size of the atom • Decreases as you go across a period(left to right). • Increases as you go down a group because you add another level of electrons.

26. Ionic Radius

27. First Ionization Energy • The energy needed to remove the most loosely held electron from an atom. • It increases as you go across a period(left to right). • It decreases down a group.

29. Electron Affinity • Electron Affinity is the attraction of an atom for an electron. • The electron affinity increases as you go across a period (left to right) and decreases as you go down a group. • In general, as electron affinity increases, an increase in ionization energy can be expected.

31. Electronegativity • The relative tendency of an atom to attract electrons to itself when it is bonded to another atom. • Electronegativities of elements have the same trend as ionization energy and electron affinity.

33. Conceptual Short Cut! • As the atomic radius gets smaller, the: • valence electrons are closer to the nucleus, • nucleus pulls harder on the electrons, • Ionization energy, electron affinity, and electronegativity will increase.

34. Periodic Trends Summary All arrows show increases!!

35. Zeff: Solution for Difficult Problems • Effective Nuclear Charge: Zeff • Zeff = Z – S • Where: • Z = atomic number • S = # of shielded (non-valence) electrons • The higher the Zeff, the greater pull of the positive nucleus on the valence electrons, which means a smaller atomic radius

36. ZeffContinued • Example: Mg vs. Mg2+ • Zeff for Mg = 12 -10 n= +2 • Zefffor Mg2+ = 12 – 2 = +10 • Therefore: the effective positive charge on Mg2+ is WAY bigger, which means the nucleus pulls on the outer electrons WAY harder, and the atom is subsequently WAY smaller. • Higher Zeff = smaller atomic radius

37. Electronegativity Cont. • The most reactive metals (lower left) have the lowest electronegativities. • Fluorine, a nonmetal, is the most electronegative element with a value of 4.0. • The difference in values for two elements tells you how the elements are bonded together when they combine with each other.

38. Electronegativity Cont. • If the difference in electronegativity values between 2 elements is below 1.7, it indicates a covalent bond. • If it is greater than 1.7, an ionic bond is indicated.

39. Reactivity Trend • Some metals and nonmetals are more reactive then others. • The easier it is for a metal to lose electrons, and the easier it is for a nonmetal to gain electrons, the more reactive they are.

40. Reactivity of Metals • Lose electrons and become positive ions. • The farther the valence electrons are from the nucleus, the less the protons pull and the easier (takes less energy) it is to remove an electron . • Reactivity increases down a group and decreases across a period. • What is the most reactive metal?

41. Reactivity of Non-Metals • Gain electrons and become negative ions. • The closer the valence electrons are to the nucleus, the more the protons pull and the easier it is to pull in an electron. • Reactivity decreases down a group and increases across a period. • What is the most reactive non-metal?

42. Metal/Nonmetal Reactivity Most reactive nonmetals – top, right Most reactive metals – bottom, left

43. Practice • WS #2 Periodic Trends • WS #2 Periodic Trends