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Ionic Compounds and Metals

Ionic Compounds and Metals. Chapter 7. 7.1 – Ion Formation. What is a chemical bond? The force that holds two atoms together. Bond formation attraction between the positive nucleus of one atom and the negative electrons of another atom attraction between positive and negative ions. Ions.

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Ionic Compounds and Metals

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  1. Ionic Compounds and Metals Chapter 7

  2. 7.1 – Ion Formation • What is a chemical bond? • The force that holds two atoms together. • Bond formation • attraction between the positive nucleus of one atom and the negative electrons of another atom • attraction between positive and negative ions

  3. Ions • Atoms or groups of atoms with a positive or negative charge • Cation • Positively charged ion • Forms when an atom loses one or more valence electrons • Metal ions • Metals atoms are reactive because they lose valence electrons easily. • Examples: Magnesium to Mg2+, Potassium to K1+

  4. Ions • Anion • Negatively charged ion • Forms when an atom gains one or more electrons • Anions designated by the ending –ide • Nonmetal ions • Gain the number of electrons that, when added to their valence electrons, equals 8 • Examples: Phosphorus to Phosphide (P3-) , Fluorine to Fluoride (F1-), Oxygen to Oxide (O2-)

  5. 7.2 – Ionic Bonds and Ionic Compounds • Ionic bond • Bond formed by electrostatic force holding oppositely charged particles together • Ionic Compounds • Compounds that contain ionic bonds • Binary Ionic Compounds • Contain a metallic cation and a nonmetallic anion • Examples: sodium chloride (NaCl), magnesium oxide (MgO) • Overall charge of Compound must equal zero

  6. Ionic Compounds • Examples • Sodium and Nitrogen • Lithium and Oxygen • Strontium and Fluorine • Aluminum and Sulfur

  7. Properties of Ionic Compounds • Physical Structure • Cations and Anions exist together in a ratio determined by the # of electrons transferred • Crystal lattice • 3-dimensional geometric arrangement of particles • Each cation is surrounded by anions and each anion is surrounded by cations • Physical Properties • Dependent upon how strongly the particles that make up the matter are attracted to one another • Melting point, boiling point, hardness, conductivity

  8. 7.3 – Names and Formulas for Ionic Compounds • Formula unit – the chemical formula for an ionic compound which represents the simplest ratio of the ions involved. • Monatomic ions = one atom ions • Examples: Li1+, Mg2+, N3-, Te2-, F1- • The superscripted number next to an ion is referred to as the oxidation number.

  9. Binary Ionic Compounds • Bond between a cation and an anion. • Cation written first followed by anion • The ratio of ions must balance the oxidation numbers (no overall charge) • Examples: K1+ and O2- create K2O • Na1+ and Cl1- • Cs1+ and N3- • Al3+ and Br1-

  10. Polyatomic Ions • Ions made up of more than one atom. • Pg. 221 Table 9 • Bond in the same manner as binary ionic compounds • Examples • Na1+ and NO31- • Ca2+ and ClO31- • Al3+ and CO32-

  11. Naming Ionic Compounds • Oxyanions – polyatomic ion composed of an element, usually a nonmetal, bonded to one or more oxygen atoms. • Chemical nomenclature • Name the cation followed by the anion. (cation is always written first) • For monatomic cations, use the element name. • For monatomic anions, use the root of the element name plus the suffix –ide. Example: CsBr • To distinguish between multiple oxidation numbers of the same element, the name of the chemical formula must indicate the oxidation number of the cation. The oxidation number is written as a Roman numeral in parentheses after the name of the cation. Example: Fe2+ and O2- ions form FeO, iron(II)oxide Fe3+ and O2- ions form Fe2O3, iron(III)oxide • When the compound contains a polyatomic ion, simply name the cation followed by the name of the polyatomic ion. Example: NaOH (NH4)2S CaCl2

  12. 7.4 Metallic Bonds and the Properties of Metals • Metallic Bonds are formed by the attraction of a metallic cation for delocalized electrons. • Electron sea model – proposes that all the metal atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons. • Delocalized electrons – electrons present in the outer energy levels of metallic bonds are not held by any specific atom and can move easily from one atom to the next.

  13. Properties of Metals • Melting and boiling points • In general, metals have moderately high melting points and high boiling points. • Malleability, ductility, and durability • Metals are malleable (hammered into sheets) • Metals are ductile (drawn into wire) • Metals are durable (strong attractions between cations and electrons) • Thermal conductivity and electrical conductivity • The movement of mobile electrons around positive metallic cations make metals good conductors. • Hardness and strength • D block electrons increase the hardness and strength

  14. Metal Alloys • Alloy – a mixture of elements that has metallic properties. • Due to the unique blend of properties, alloy have a wide range of commercial applications. • List of common commercial alloys pg. 228 Table 13 • Substitutional Alloy • Some of the atoms in the original metallic solid are replaced by other metals of similar atomic size. • Sterling silver • Interstitial Alloy • Formed when the small holes (interstices) in a metallic crystal are filled with smaller atoms. • Carbon steel

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