5.1 Development of the Periodic Table Explain the role of Mendeleev in the development of the periodic table Explain how the periodic law can be used to predict properties Describe the modern periodic table.

J.W. Dobereiner 1780 - 1849 Triad – groups of three elements. Noticed that some elements in groups of THREE demonstrated similar chemical properties. Also noted that the middle element was approximately the average mass of the other two elements.

J.A.R. Newlands 1837 - 1898 Law of Octaves - presented another way of organizing the elements by octaves. He noticed that every 8 elements shared similar properties. The 1st was like that of the 9th element. The 2nd was like that of the 10th element, etc.

Dmitri Mendeleev 1834-1907 Noticed that when the elements were placed in order of atomic mass Similar properties appeared at regular intervals Created a table with columns of elements with similar properties The First Periodic Table Many holes were filled in over time

Henry Moseley 1887-1915 Recognized a pattern based upon Atomic Number Periodic Table is ordered by Atomic Number, not Atomic Masses Ar (At# 18, AtW 39.95) – K (At# 19, AtW 39.01) Periodic Law – properties of elements are periodic functions of their atomic number. Cycles are observed as the Atomic Number increases

5.2 Reading the Periodic Table Arrangement of elements in order of their atomic number so that elements with similar properties fall into the same column. Creating Groups or Families. Periodicity repeated patterns w/ the increasing of atomic number Like a spiral growing in diameter outwards Atomic Number Differences between groups 2,8,8,18,18,32,32

Electron Configuration & The Periodic Table A fully occupied energy level is stable Low Energy Noble Gases Every element wants to have a fully occupied outer energy level Np6 Every element wants to have a Noble Gas Config. Electron configurations can be used to determine the reactivity of elements

Electron Configuration 7 6 5 4 3 2 1 Lr Lu Ac La s - block d – block (s-1) p - block f – block (s-2) Electron Configuration

Noble Gas Configuration or Short-Hand Notation Since every element tries to complete its valence shell, like noble gases, the short-hand notation takes this into account. Write the last noble gas to occur prior to the element that will be configured. Count down the periods to the element being configured and start the s orbital at that energy level. Move from the left to the right of the period until the element being configured is reached. Every s and p will have the same energy level, d will be 1 less than the s and f will be 2 less.

Writing in Short-hand notation Al – Write the short hand notation for the following: V Rb I Hg U W

Classwork Use Short-hand notation: Yb Au At Cs Rn

Chapter 6 - Groups, Periods, and Blocks Groups – Vertical Columns Periods – Horizontal Rows Blocks – Sections of P.T. (s,p,d,f) Main Block Elements – s & p together Representative group s-block – Alkali & Alkaline Earth Metals d-block – Transition Metals p-block – Metalloids, Halogens, Noble Gases, And other Non-Metals & Metals f-block – Lanthanide & Actinide

The Alkali & Alkaline Metals Alkali Metals Group 1 or 1A Li, Na, K, Rb, Cs, Fr Soft (cut w/ Knife Extremely Reactive Not found as Free Elements Less Dense than H2O React w/ Non-Metals to form Salts E-Config. ns1 Alkaline Earth Metals Group 2 or 2A Harder, Denser, & Less reactive than Gr1 Too Reactive to be found as Free Elements Electron Config. ns2

Hydrogen & Helium Both Non-metals Colorless & Odorless Hydrogen Placed In Group 1 Properties do not resemble elements in Group 1 Helium Group 18 even though it has a ns2 config. Placed in Gr. 18 because of its non-reactive nature (similar to the noble gases)

The Transition Metals (Groups 3-12) Electron Config ns2(n-1)d1-10 Doesn’t always follow a normal pattern Groups 3-12 Properties Metallic Good Conductors Heat Electricity Compared to s-block Metals Harder, Denser Hg is an exception (liquid) Higher Melting Point Less Reactive Common Transition metals Gold, Silver, Copper, Iron, Tungsten, Cobalt, Nickel, Platinum

Metals vs. Nonmetals Metals Good conductors of heat and electricity. Luster Malleable Non-metals Poor conductors of heat and electricity. Dull Brittle

Halogens, Metalloids, & Noble Gases (p-block elements) semiconductors mostly brittle solids metal & non-metal properties Noble Gases (Group 18 or 8A) Inert (non-reactive) Stable (low energy) Monatomic Gases P- Block Elements Electron Config. np(1-6) Halogens (Group 17 or 7A) np5 Configuration Most reactive of non-metals React w/ metals to form salts

Lanthanides & Actinides ‘Inner Transition Metals’ 4f (1-14) Configuration “Rare-Earth Metals” Which are not actually rare Shiny Luster Reactive Metals All have practical uses Phosphors in TV’s contain many Lanthanide metals. Actinides 5f (1-14) Configuration Unstable & Radioactive 1st four are found in nature The rest are synthetic Not very Practical Uranium is an exception Nuclear Power ess05_vid_fusion_300.mov

5.3 Periodic Trends Elements are grouped according to their Physical and Chemical Properties. Period Trends – how a characteristic increases/decreases across a period Group Trends – how a characteristic increases/decreases up & down a group

Atomic & Ionic Radii Atomic Radius – ½ the distance between the nuclei of identical atoms joined in a molecule Period Trend Decreases L to R Greater # of p+ pulls electrons closer Group Trend Increases T to B Increase in energy level – greater size Ionic Radii 2 Types of Ions Cations ( + Charge) Loss of Electrons Anions ( - Charge) Gaining of Electrons Cations – smaller than neutral counterpart Less electrons Anions – larger than neutral counterpart Gained Electrons

Valence Electrons Valence Electrons Valence Electrons – electrons available to be lost, gained, or shared in the formation of compounds. Electrons in the outer energy shell 1s22s22p63s23p64s23d104p4 - 6 V.E. Electrons in the outer s & p sublevels Representative Group Valence Electrons Group Number 1 2 13 14 15 16 17 18 Valence Electrons 3 4 5 6 7 8

Ions –atoms that have gained or lost a # of electrons. Cations – Atoms that need to lose electrons to reach a complete valence shell. Na will lose 1e- to become Na 1+ Na = 1s2 2s2 2p6 3s1 Na 1+ = 1s2 2s2 2p6 Anions – Atoms that need to gain electrons to reach a complete valence shell. Cl will gain 1e- to become Cl 1- Cl = 1s2 2s2 2p6 3s2 3p5 Cl 1- = 1s2 2s2 2p6 3s2 3p6

Noble Gas Configuration When atoms form ions their electron configurations resemble that of a noble gas element. Pseuodo noble gas- Ions that resemble noble gases, but contain a d or f block configuration.

Ionization Energy Ionization – any process that results in the formation of an ion. Ionization energy – the energy required to remove an electron from an atom A + energy  A+ + e- Measured w/ atoms in gas phase to eliminate influence from nearby atoms. Period Trend – Increase from L to R Group Trend – Decreases from T to B

Multiple Ionization Energies More electrons more energy needed. Usually there is a large jump in ionization energy when an atom loses an entire energy level. Example: Sodium has 1 electron in valence shell. 1st = 496 2nd = 4562 3rd = 6912 1st ionization energy, Ability to remove one electron from an atom’s valence shell. 2nd ionization energy, Ability to remove the first and second electron from an atom’s valence shell. 3rd ionization energy, Ability to remove three electrons from an atom’s valence shell.

Electron Affinity Electron Affinity – is the energy change that occurs when an electron is acquired by a neutral atom. Some atoms want electrons A + e -  A- + Energy ( Exothermic) Some atoms do not want electrons A + e - + Energy  A- (Endothermic) Will be unstable & lose electron spontaneously Period Trend – More easily acquires electrons from L to R ( More Negative/ More Exothermic) Group Trend – More easily acquires electrons from B to T

Electronegativity Electronegativity – the ability of an atom in a compound to attract electrons Fluorine – has the strongest electronegativity arbitrarily assigned a value of 4 all other atoms were then compared to F Period Trends – increases from L to R Group Trend – decreases from T to B Table 5-23, page 184, summarizes all trends

Trend summary Electronegativity Increases Ionization Increases Electron Affinity Increases Atomic Radii / Ionic Radii Decreases