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Unformatted text preview: CHAPTER 3 PERIODICITY (IB TOPICS 3 AND 13) SUMMARY Basic concepts • Periodic Table is arranged according to increasing atomic number and consists of horizontal rows called periods, and vertical columns called groups (or families). The arrangement is such that elements with similar chemical properties fall directly beneath each other in the same group. • A period is a series of elements arranged according to increasing atomic number, which begins with the first element having one electron in a new main energy level. • A group is a vertical column consisting of elements with the same electron arrangement in their outer energy levels, which gives the group similar chemical properties. • Modern Periodic Law states that chemical and physical properties of elements vary periodically if elements are arranged in order of increasing atomic numbers. • Periodicity is the regular repeating of properties according to arrangement of elements in the periodic table (i.e., after regular intervals) such as atomic radius, ionisation energy, etc. arising from the systematic filling of successive energy level. • Valence electrons are electrons in the outermost energy level (the highest energy level) of an atom and are usually the electrons that take part in a chemical reaction. • Properties of elements are chiefly due to the number and arrangement of electrons in the outer energy level of atoms. • The number of valence electrons is the same for a group, but increases across a period. • The transition metals are the large d-block elements in the middle of the Periodic Table from Sc to Zn etc. (includes three transition series). • Ionization Energy (IE) is the minimum energy required to form a mole of 1+ ions (by removing an electron from each atom) in the gaseous state: M (g) M+ (g) + e– (units: kJ mol–1) • Electronegativity is a measure of how strongly the atom attracts the electrons in a covalent bond. © IBID Press 2007 1 CHAPTER 3 PERIODICITY (IB TOPICS 3 AND 13) SUMMARY Alkali metals are extremely reactive, electropositive metals. They react vigorously with O2 in the air and oxidize fast to form ionic oxides. contain one valence electron that is very easily lost due to low ionization energy; thus they are very good reducing agents: Na (s) → Na+ (aq) + e–. In compounds, these exist as +1 cations with noble gas electron configuration. Since a valence e– further away from the nucleus down a group is more easily removed, reactivity increases from Li to Cs. Reaction with Water: Reactivity increases down the group: 2 Li (s) + 2 H2O (l) • 2 LiOH (aq) + H2 (g) + heat Reaction with Halogens: form ionic salts (e.g. Na+Cl–); the reaction is highly exothermic. The alkali halides are ionic, neutral, water-soluble, white crystalline compounds. 2 Na (s) + Cl2 (g) 2 NaCl (s) + heat Halogens Are non-metallic; metallic character increases down the group (I2 is a shiny solid). Higher ionization energies indicate little tendency to lose electrons. Have 7 valence electrons and achieve noble gas configuration by gaining an electron to form an anion (e.g. Cl–), or by sharing an e– pair with another atom (e.g., Cl2). Are diatomic (F2, Cl2, Br2, I2), non-polar molecules. They are simple molecular substances with only weak van der Waal’s between molecules (Bonding, Ch. 4). Higher electronegativities mean that halogens have a tendency to accept an electron, and act as oxidizing agents: X2 +2 e– 2 X–. Oxidizing strength decreases down the group, since the atom gets larger and attraction for the electrons decreases. Reactions of halogens X2 with halide ions X–: Halogen Displacement Reactions: Weaker halogens are displaced from their salts by more powerful oxidizing agents. Thus Cl2 can displace Br2 from Br– and I2 from I– (as it is a stronger oxidizing agent than both Br and I), but it cannot displace F2 from F–. Identification of Halide Ions – Reaction with silver ions, Ag+: Silver halides, with different colored precipiates can be used to identify a halide ion: 1. AgNO3 (aq) + NaF (aq) → No precipitate formed. 2. AgNO3 (aq) + NaCl (aq) → AgCl (s) + NaNO3 (aq), a white precipitate. 3. Ag+ (aq) + Br– (aq) → AgBr (s), a cream precipitate. 4. Ag+ (aq) + I– (aq) → AgI (s), a yellow precipitate. © IBID Press 2007 2 CHAPTER 3 PERIODICITY (IB TOPICS 3 AND 13) SUMMARY Period 3 elements Ionization energy and electronegativity increase across the period; thus elements do not lose electrons as easily and metallic character decreases across the period. Bonding changes from metallic to covalent. • Na, Mg and Al are metallic, good conductors of heat and electricity, with low to medium melting points, and lower electronegativities. • Si is a semi-conductor / metalloid. It forms a network covalent solid of very high melting point. • P, S and Cl are elements of higher electronegativity and are non-metallic. These form simple molecular substances (see bonding) with lower boiling points; the bonding between atoms is covalent and the bonding between molecules is weak van der Waal’s forces. Period 3 oxides and chlorides Period 3 oxides and chlorides s block elements: Na and Mg p block elements Al, Si; P, S, Cl © IBID Press 2007 Bonding in oxides and chlorides Acid/base properties of oxides Acid/base properties of chlorides Ionic solids with high melting points; bonding between active metals of low IE and active nonmetals of high electronegativity Ions present, e.g. Oxides form basic solutions Na+ + Cl− Mg2+ + O2−; conduct electricity in the molten state (and in aqueous sol) Na2O + H2O 2NaOH; Chlorides (NaCl, MgCl2) form neutral solutions. MgCl2 is actually very slightly acidic. Al2O3 is ionic; SiO2 is a giant covalent structure (with very high melting point). No mobile ions present; nonconductors Oxides are acidic except Al2O3 which is amphoteric (can act as an acid or a base) Others: covalent bonding; form simple molecular substances Electrical conductivityof molten oxides and chlorides MgO + H2O Mg(OH)2 Chlorides are acidic; undergo hydrolysis reaction to produce HCl vapor pH of chlorides in water NaCl : pH =7 MgCl2: pH ∼ 7 Al2Cl6: pH ∼ 3 SiCl4, PCl5, S2Cl2 : pH ∼ 2 3 CHAPTER 3 PERIODICITY (IB TOPICS 3 AND 13) SUMMARY d-block elements • d-block elements with characteristic properties form at least one ion with a partially filled d sub-level (with 1 to 9 electrons); Sc and Zn are not typical of d-block elements. • Characteristic properties are: presence of variable oxidation states, formation of complex ions, colored complexes, and catalytic activity. • 4s is lower in energy (but further from the nucleus) than 3d sub-level; 4s electrons are lost before 3d; both 4s and 3d can behave as valence electrons because these are close in energy. • Multiple oxidation states are due to 4s and 3d being close in energy. All d-block elements (except Sc) show an oxidation state of +2. All d-block elements except Zn show an oxidation state of +3. • Complexes are formed when a central metal ion is bonded to ligands; ligands contain (at least one) lone e− pair that form coordinate bonds with the central transition metal ion. • Color is due to d-d electron transitions between split d orbitals. • Transition metals behave as surface catalysts and their compounds more often as intermediate catalysts, which are important for industrial and biological reactions. (N.B. Shading indicates Topic 13 (AHL) material.) © IBID Press 2007 4 ...
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This note was uploaded on 09/23/2010 for the course CS 001 taught by Professor Jix during the Spring '10 term at Riverside Community College.

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