Unit III. Chemical Periodicity

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1 Unit III. Chemical Periodicity History of Periodic Table Modern Periodic Table Periodic Trends Links Return to Notes PDF Format READING ASSIGNMENT 1: Read Ch. 14.1, p Answer questions 1-5 p. 396 A. History of the Periodic Table -Sb, C, Cu, Au, Fe, Pb, Hg, Ag, S, Sn were only elements known before As comes slightly later -Prior to late 1700 s only 24 elements were known 1. Johann Wolfgang Döbereiner Germany Triads- groups of 3 elements that share similar properties Li, Na, K Cl, Br, I 2. John Newlands- c England Law of Octaves- when the elements are placed in order of atomic weights, a cycle of properties is repeated with every eight elements 3. Dmitri Mendeleev Russia -devised 1 st periodic table- coupled properties of elements and organization of increasing atomic mass -Law- Elements arranged according to their atomic masses present a clear periodicity of properties. Groupings- organized elements into an eight-column format. Mendeleev's Periodic Table a. Some spaces had to be skipped for not-yet discovered elements ex. Eka-Silicon- (Germanium) similar properties to that of Silicon & Tin (others- Gallium, Scandium) Predicted properties for Mendeleev s Eka-Silicon and properties of Germanium: Chloride formula Element Atomic Weight Density Oxide formula Eka-Silicon (predicted 1871) Germanium (discovered 1886) g/cm 3 EsO 2 EsCl g/cm 3 GeO 2 GeCl 4 b. Problems with Mendeleev s Periodic Table 1. Following increasing atomic mass, sometimes elements had to be reversed. ex. Ni & Co, I & Te, Ar & K 2. Newly discovered elements had no spaces available ex.holmium and Samarium 3. Elements in the same group were sometimes quite different in their reactivity ex. Group I- alkali metals & Coinage metals 4. Lothar Meyer- c Germany -Worked on periodic relationship similar to that of Mendeleev. Acknowledged that Mendeleev had original idea. Meyer's Periodic Table 5. Henry Moseley England Atomic number- the number of protons found in the nucleus of a specific element -current periodic table utilizes increasing atomic numbers instead of atomic masses High Frequency Spectra of the Elements, Original work by Henry Moseley REVIEW: Contributors to the Periodic Table. PRACTICE: Periodic Table of Elements Quizzes 1 through 4.

2 B. Modern Periodic Table Periodic Law- When the elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties. 1. Elementsa. Identified by: 1. name 2. symbol 3. atomic number. 4. atomic mass b. Characterized by 1. Physical Properties- Boiling/Melting Points, Density, Color, Crystalline Structure, State of Matter, etc. 2. Chemical Properties- Oxidation States, Acid/Basic Properties, Ionization Energy, Electron Affinity, etc. 2. Periods- (Rows)- represents the number of energy levels 3. Columns- (Families/Groups)- representative of elements with similar properties- -analagous to the number of electrons in outer energy level(s) a. Representative Elements (A groups) IA- VIIIA- fills s and p orbitals -IA. Alkali Metals- -IIA. Alkaline Earth Metals -IIIA. The Aluminum Family -IVA. The Carbon Family -VA. The Pnicogens -VIA. The Chalcogens -VIIA. The Halogens -VIIIA- Noble Gases- completely filled outer s and p orbitals. These demonstrate chemical stability. b.transition Metals (B groups) BI-VIIIB- fills outermost s (ns) and prior d orbitals (n-1d) c. Inner-Transition Metals- fills outermost s orbital (ns) and twice prior f orbitals (n-2f). a.k.a. Rare Earth Metals Lanthanide Series. The elements in the 4f orbitals (top row of the inner transition metals) Actinide Series. The elements in the 5f orbitals (botton row of the inner transition metals) EXPLANATION: -Group Numbering. Explanation of IUPAC and CAS group numbering systems 4. Line of Demarkation: Staggered line used to separate metals from nonmetals a. metals- elements to the left of the line b. nonmetals- elements to the right of the line c. metalloids/weak-metals- elements along the line. PRACTICE: Periodic Table Quizzes: (there are 5 quizzes here): Groups on the Periodic Table, Atomic Number - Quiz 1, Periodic Table of Elements - Quizzes 1-3. ASSIGNMENT 1: Section 14.1 Worksheet READING ASSIGNMENT 2: Read section 14.2 pp , answer questions 6-9, and construct a concept map from p.408 C. Periodic Properties of the Elements -The properties are based on the electrons and their positions Factors that affect the properties: a. the number of valence electrons b. the magnitude of the nuclear charge (Z) and the total number of electrons surrounding the nucleus c. the number of filled shells lying between the nucleus and the valence shell d. the distances of the electrons in the various shells from each other and from the nucleus

3 RESOURCE: Periodic Properties of the Elements. The elements can be sorted by properties 1. Atomic Radius a. Ways to measure the radius of an atom 1. covalent radius- ½ distance from nuclei of 2 identical atoms joined by a single covalent bond 2. van der Waals radius- ½ distance from nuclei of 2 atoms of neighboring molecules 3. metallic radius- ½ distance from nuclei of 2 atoms in a solid metal 4. atomic radius- based on the quantum model. Theoretical/Mathematical approach b. Trends in the periodic table 1. Period- radius decreases from left to right- increase in (Z) with same number of energy levels- 2. Group- radius increases from top to bottom- increase in the number of energy levels (Principal Q.N. increases) The atomic radii for the elements in the first 3 energy levels Effective Nuclear Charge (Z eff )- dependent upon (Z) and the shielding effect of other electrons ** Shielding- interior orbitals that contain electrons shield the attraction of the nucleus on the valence electrons. What effects are seen from shielding? Slater s Rule: (Z eff = Z σ) where σ is the shielding factor for valence electrons calculating σ for : a. for valence electrons in s and p type orbitals 1. (ns & np) electrons shield at 35% 2. for (n-1) orbitals, these shield at 85% 3. for (n-2) orbitals, these shield at 100% b. for valence electrons in d and f type orbitals 1. (nd, nf) shield at 35% 2. higher orbitals (n+1) shield at 0% (ns, np) 3. s and p electrons in the same energy level (ns & np)and lower energy levels (n-1& et.al.) shield at 100% 2. Ionization Energy: The amount of energy needed to remove an electron from a gaseous atom. Production of a positive ion a. Ion- an atom which has gained or lost electrons- dependent upon the Effective Nuclear Charge 1. cation- positive ion- due to a loss of electrons 2. anion- negative ion- due to a gain of electrons b. Trends 1. Groups- First Ionization energies decrease from top to bottom 2. Periods- General- First Ionization energies increase from left to right

4 g g some exceptions occuri. Due to shielding of ns on 1 st electron in a np orbital (B, Al, Ga) ii. Losing 1 paired electron is less than losing a parallel electron (O, S, Se) ** these exceptions fail at higher energy levels** 1st Ionization energy for the elements in the first 3 energy levels c. Successive Ionization energies 1. First Ionization- removing the first valence electron 2. Second Ionization- removing the second valence electron 3. Third Ionization- removing the third valence electron -removing successive electrons reduces the shielding factor but maintains Z. This increases the Z eff on the remaining electrons. RESOURCE: Table of Successive Ionization Energies 3. Electron Affinity: The energy change that accompanies the addition of an electron to a gaseous atom. For most atoms, energy is released when an electron is gained. This is seen as a negative energy change. General Trend: Atoms that have a high ionization energy, will typically have a larger negative electron affinity. Noble Gases have a positive electron affinity because it would take energy to add electrons to these atoms. TUTORIAL & PRACTICE: Ionization Energies & Electron Affinities 4. Ion Size a. Cations are always smaller than their neutral atom counterpart -losing electrons may: lose valence shells and/or increase z eff values b. Anions are always bigger than their neutral atom counterpart. -gaining electrons decrease z eff values c. Trends- 1. Period- Ion size decreases from left to right, (cation/anion specific) 2. Groups- Ion size increases from top to bottom. Isoelectric species- atoms and ions that have the same electron configurations (ex. N -3, O -2, F -1, Ne, Na +1, Mg +2, Al +3 ) Octet Rule- All atoms strive to have full valence shells (typically 8 electrons in the ns & np orbitals) exception is 1 st energy level (no p orbital: 2 electrons) Oxidation Number- the charge of the stable ion after gaining/losing electrons. Net difference between protons & electrons 5 El t ti it Th tt ti th t t h l t h it i h i ll bi d ith th t

5 5. Electronegativity: The attraction that an atom has on an electron when it is chemically combined with another atom. a. Trends 1.Periods- electronegativities increase from left to right 2. Groups- decrease from top to bottom. APPLET: Periodic Table & properties PRACTICE: Periodic Table Trends quizzes 1 through 5 PRACTICE Periodicity Quiz. ASSIGNMENT 2: Ch 14 Review Questions: p.409; #s 12, 14, 17, 18, 19, 21, 25, 26, 29, 30, 34, & 35 D. Links Review of the Periodic Table: Frostburg State WebElements Periodic Table Chemical elements.com Cool Periodic Table Periodic Tables in the English Classroom The Periodic Table and Poetry

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