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Understanding Electron Dot Structures and Ionic Compounds in Chemistry

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fiona
14 diapositives · 282 Views · Oct 05, 2014

This chapter outlines the fundamentals of valence electrons, their role in chemical reactions, and how to predict valence numbers using the periodic table. It introduces electron dot structures as a method to visualize valence electrons, specifically in ionic compounds where the overall charge must equal zero. The document explains how to write chemical formulas using cations, anions, and polyatomic ions, including examples such as lithium fluoride and ammonium chloride. Also covered are the proper naming conventions for chemical compounds.

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  1. 20 October 2008 Chapter 8 Notes Mr. Herbst Room 2705

  2. Electron Dot Structures • Valence electrons are the furthest away from the nucleus and are involved in chemical reactions • Valence number can be predicted from the periodic table (Same as Column #) • Electron dot structures are a simple way to show the number of valence electrons

  3. Lewis Structures for Ionic Compounds • The overall charge on the compound must equal zero. (that is, the number of electrons lost by one atom must equal the number of electrons gained by the other atom.) • The Lewis Structure (electron dot diagram) of each ion is used to construct the Lewis Structure (electron dot diagram) for the ionic compound.

  4. Dot Structure to show Ionic Bonds • Examples • Lithium fluoride, LiF Li. . F Li+ [ F ]-

  5. Naming and Formula Writing • Cation – ion with a positive charge • Anion – ion with a negative charge • Polyatomic Ion – tightly bound group of atoms that behave as a unit and carry a charge

  6. Chemical Formulas • ‘shorthand’ for chemical names • subscripts represent numbers of atoms; superscripts represent charge • molecular formula – chemical formula that shows that actual number and kinds of atoms present in a molecule of a compound • empirical formula – shows the lowest whole number ratio of elements in a compound; used to get molecular formula

  7. Writing ionic Formulas (metal/nonmetal) • Write the symbol for the cation first. • Write the symbol for the anion second. • find the oxidation number of each ion. • Figure out how many of each ion is required to make the charge equal to zero • Write the number of each required as subscripts. ** Remember: Charges are never included in the formula; they are only a tool.

  8. +1 -2 • Example: Lithium sulfide • Li+ S-2 • Li+ • formula is Li2S • b/c it takes 2 lithium ions to equal the charge of the sulfide ion.

  9. Formulas with Polyatomic Ions • Write the formulas the same way as binary formulas using a polyatomic ion in place of one of the simple ions. • Ex. Ammonium chloride • NH4+ & Cl- • formula is NH4Cl • b/c charges already equal each other – no more ions needed

  10. When more than one of a polyatomic ion is needed, put parentheses ( ) around the polyatomic ion in the formula. • Ex. Magnesium hydroxide • Mg+2 & OH- • formula is Mg(OH)2 • b/c two hydroxides are required to equal the magnesium

  11. Writing Chemical Names • Write the name of the cation first. • Use a roman numeral with a cation that can have a variable oxidation number. • Write the name of the anion second with the ending changed to ‘ide’. (ex. chlorine becomes ‘chloride’) • When the anion is a polyatomic ion, do NOT change the ending.