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Section 7.1 Ion formation

Chapter 7 Ionic Compounds and Metals. Section 7.1 Ion formation. Chemical Bonds. A chemical bond is the force that holds two atoms together. Can form by the attraction between the positive nucleus of one atom and the negative electrons of another

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Section 7.1 Ion formation

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  1. Chapter 7Ionic Compounds and Metals Section 7.1Ion formation

  2. Chemical Bonds A chemical bond is the force that holds two atoms together. • Can form by the attraction between the positive nucleus of one atom and the negative electrons of another • Can form between positive and negative ions

  3. Valence Electrons • Electrons in the outermost principal energy level • Shown in the electron dot structures • Octet rule – atoms will gain, lose or share electrons to obtain 8 valence electrons • The valence electrons determine the bonding properties of the atom

  4. Positive Ion Formation • A positively charged ion is called a cation. • Positive ions are formed when an atom loses one or more valence electrons • Metals make positive ions

  5. Negative Ion Formation • A negatively charged ion is called an anion. • Negative ions are formed when an atom gains one or more electrons in its valence shell. • Nonmetals make negative ions.

  6. 7.2: Ionic bonds and ionic compoundsFormation of an Ionic Bond • An ionic bond is the electrostatic force that holds oppositely charged particles together in an ionic compound • Compounds that contain ionic bonds are called ionic compounds. • Ionic compounds are formed between metals (+ charge) and nonmetals (- charge).

  7. Binary Ionic Compounds • Contain a metallic cation and a nonmetallic anion. • Formation of Binary Ionic Compounds • Electron(s) is/are transferred from metal to nonmetal • Metal becomes positive, nonmetal becomes negative • Opposite charges attract

  8. Properties of Ionic Compounds • Take the structure of a crystal lattice • Many units of positive and negative ions stick together in a three-dimensional geometric arrangement • Can conduct electricity when dissolved in water (they are electrolytes and break into ions when dissolved in water), but not in solid form • Melting point, boiling point and hardness depend upon how strongly the ions are attracted to each other

  9. Formulas for Ionic Compounds • Monatomic ions are one-atom ions • Examples: Mg2+ , Br-1 • Oxidation numbers are the charges on ions • Note: some elements have multiple oxidation states – you will have a periodic table to tell this • Binary ionic compounds are made of two monatomic ions (one positive, one negative)

  10. Formulas for Binary Ionic Compounds • Symbol for cation is written first, anion second • Subscripts tell the number of atoms of each element • What are the following compounds made of? • CaF2 1 calcium, 2 fluorine • Na2S 2 sodium, 1 sulfur • NaCl 1 sodium, 1 chlorine

  11. Naming Binary Ionic Compounds • Name the cation first • Name the anion second with –ide at the end • Examples • CaF2 calcium fluoride • Na2S  sodium sulfide • NaCl  sodium chloride

  12. Try Naming a few moreBinary Ionic Compounds • K2O potassium oxide • Al2S3 aluminum sulfide • Na3N sodium nitride

  13. What if the cation has more than one oxidation state? • You tell which ion was used by putting a Roman Numeral after the name of the cation • Example: • CuS • We know S was -2 (that’s the only one it makes) • If there is only one atom of each element, the Cu must have been +2 • So, the name is written as Copper (II) sulfide [the “II” indicates the charge] • Make sure, especially with transition elements, that you are checking the oxidation states

  14. Writing Formulas for Binary Ionic Compounds • Look up the charges for each element • For a compound to form, the total charge must balance out to zero (positive charges must equal negative charges) • Example: • Sodium bromide • Na is +1, Br is -1 • Only need one of each to balance • Formula is NaBr

  15. Try writing some more formulasBinary Ionic Compounds • Potassium Iodide KI • Aluminum bromide AlBr3 • Magnesium chloride MgCl2 • Cesium nitride Cs3N

  16. Formulas for Polyatomic Ionic Compounds • Polyatomic ions are ions that are made up of more than one atom • You will have a chart for these and do not have to memorize them. • Examples: • SO42- = sulfate • CN- = cyanide • NH4+ = ammonium

  17. Naming Polyatomic Ionic Compounds • Name the cation first, anion second • Name the polyatomic as is – don’t change its name at all • Examples: • Ca3(PO4)2 calcium phosphate • Mg(CN)2 magnesium cyanide • NH4Cl ammonium chloride

  18. Now you try naming Polyatomic Ionic Compounds • NaNO3 sodium nitrate • Ca(ClO3)2 calcium chlorate • Al2(CO3)3 aluminum carbonate

  19. Writing formulas for Polyatomic Ionic Compounds • Same as binary ionic compounds EXCEPT you may not change anything in the polyatomic ion formula • Put them in a (parenthesis) and put subscripts outside that parenthesis • Example: • Calcium Nitrate • Ions are Ca2+ and NO3- • Formula will be Ca(NO3)2

  20. Now you try writing formulas forPolyatomic Ionic Compounds • Sodium hydroxide NaOH • Copper (II) nitrate Cu(NO3)2 • Silver chromate Ag2CrO4

  21. 7.3: Metallic bonds and theproperties of metals • The electron sea model proposes that all the metal atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons • Since the electrons are free to move, they are called delocalized electrons • A metallic bond is the attraction of a metallic cation for delocalized electrons

  22. Properties of Metals(revisited) • Moderately high melting points • High boiling points • Malleable, ductile, durable • Conduct heat and electricity well • Transition metals are harder/stronger than alkali metals because the transition metals have more delocalized electrons

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