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Chemical Bonding

Chemical Bonding . Chapter 7. The Octet Rule. Atoms tend to gain, lose, or share electrons in order to acquire a full set of valence electrons. “octet” – most atoms need 8 valence electrons for a full set Gaining or losing g ions = ionic bonding Sharing = covalent bonding. Ionic Bonding.

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Chemical Bonding

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  1. Chemical Bonding Chapter 7

  2. The Octet Rule • Atoms tend to gain, lose, or share electrons in order to acquire a full set of valence electrons. • “octet” – most atoms need 8 valence electrons for a full set • Gaining or losing g ions = ionic bonding • Sharing = covalent bonding

  3. Ionic Bonding

  4. Properties of Ionic Compounds • High melting points • Brittle • Usually salts • Many dissolve in water • Can conduct electricity because ions separate and are charged in the solution

  5. Ionic Bonds • Positively charged ion attracted to negatively charged ion • Positive ions = cations • Negative ions = anions • Metal + nonmetal • Metals form cations • Nonmetals form anions

  6. Types of Ions • Monatomic = “one-atom” • H+, Ca2+, Br-, N3- • Polyatomic = “many-atoms” • NH4+, OH-, SO42-,

  7. Lewis Dot Structures • Developed by American chemist Gilbert Lewis (1875-1946) • Valence electrons represented by dots around the element symbol • No more than two dots per side • Can be used to show rearrangement of electrons during chemical reactions

  8. Binary Ionic Compounds • Contain ions of only two elements • Formula: Cation written first, then anion • Charges of ions written as superscripts, # of atoms in a compound written as subscripts • Ratio written in lowest terms = empirical formula

  9. Binary Ionic Compounds • Draw the Lewis Dot Structures for sodium and chlorine • Using an arrow, identify how the transfer of 1 electron can create 2 new ions

  10. Sodium transfers an electron to chlorine. • Sodium becomes a positive ion with a +1 charge. • Chlorine becomes a negative ion with a -1 charge.

  11. Binary Ionic Compounds Na+ + Cl-g NaCl • The total (net) charge on the compound should be zero. • You must determine how many of each ion will need to be in the compound to balance out the charges.

  12. Compound Formula Practice magnesium ion + oxide ion Mg2+ + O2-g Mg2+ + O2-g MgO calcium ion and bromide ion Ca2+ + Br-g strontium ion and nitride ion Sr2+ + N3-g Mg2O2 CaBr2 Sr3N2

  13. The Crisscross Method for Writing Compound Formulas • Write the ion symbols (with their charges as superscripts) for the cation and anion • Criss-cross the two charges, moving them diagonally from one ion’s superscript to the other ion’s subscript • Drop the sign!

  14. Crisscross Method Practice magnesium ion and chloride ion Mg2+ Cl-1 Mg Cl = MgCl2

  15. Naming Ionic Compounds • Name the cation using its element name. • Name the anion by dropping the ending of the element name and adding –ide. Ca3P2 calcium phosphide • If the anion is polyatomic, simply name it using the ion’s name Mg3(PO4)2 magnesium phosphate

  16. Naming Ionic Compounds • If the cation has more than one valence (it can have different charges), indicate the charge using roman numerals in parenthesis after the cation name. FeO = iron (II) oxide Fe2O3 = iron (III) oxide

  17. Covalent Bonding

  18. Covalent Bonds • Formed by a shared pair of electrons between two atoms • Make up molecules (which make up molecular substances) • Between nonmetals

  19. Formulas • Empirical formula gives the lowest ratio of types of atoms in a compound • Molecular formula gives the exact number of atoms of each element in a single molecule of a compound • Structural formula shows how atoms are bonded together

  20. Formula Example: Glucose molecular formula C6H12O6 empirical formula CH2O structural formula

  21. Lewis Dot Structures • For molecules: • Show pairs of electrons that are shared between atoms using 2 dots or 1 dash. • Leave electrons not involved in bonds as dots.

  22. Lewis Dot Structures Draw the Lewis dot structures for: F2 NH3 H2O H2CO C2H2

  23. Exceptions to the Octet Rule • Less than an octet • BF3 • More than an octet • SF4 • Odd number of electrons • NO

  24. Properties of Covalent Bonds • Polar covalent bonds • Due to electronegativity difference • More electronegative atom gets slightly negative charge (higher electron density) • Less electronegative atom gets slightly positive charge (lower electron density) • Nonpolar covalent bonds • No electronegativity difference • Share electrons equally

  25. Practice • Look at page 184. • Which compound has the more polar bond, HCl or F2? How do you know? • Which atom in HCl has the higher electronegativity? • Draw the Lewis Dot Structure for HCl and indicate the partial charges on the atoms.

  26. Naming Covalent Compounds (Molecules) • Similar to naming ionic compounds, but prefixes must be added to tell the ratio of atoms in the compound.

  27. Naming Covalent Compounds (Molecules) • Most electronegative element written last in formula and name. • Drop ending of this element’s name and add –ide. Si2Br6 disilicon hexabromide • Don’t include mono- prefix for 1st element listed. CF4 carbon tetrafluoride

  28. Naming Covalent Compounds (Molecules) • Shorten prefixes to make names easier to say. H2O dihydrogen monoxide not dihydrogen monooxide • Sometimes common names are used. O2 = oxygen NH3 = ammonia

  29. Hydrates and Acids

  30. Naming Hydrates • Hydrates are ionic compounds that absorb water into their solid structures. • Anhydrous substances are water-free • Naming: • Name the ionic compound • Using the prefixes that you have learned, identify the degree of hydration MgSO4s 7 H2O magnesium sulfate heptahydrate

  31. Naming Acids • Acids are molecular substances that dissolve in water to produce hydrogen ions (H+). • Can separate into ions even though they are molecular compounds • Hydrogen is the cation in acids.

  32. Naming Acids • If the anion ends in –ide • Begin the name with hydro- • Add the root name of the anion, but change the ending from –ide to –ic • Add the word acid HBr HCl H2S Hydrobromic acid Hydrochloric acid Hydrosulfuric acid

  33. Naming Acids • If the anion ends in –ate • Do NOT begin with hydro- • Keep the root of the anion, but change the ending from –ate to –ic • Add the word acid HNO3 H3PO4 HC2H3O2 nitric acid phosphoric acid acetic acid

  34. Naming Acids • If the anion ends in –ite • Do NOT begin with hydro- • Keep the root of the anion, but change the ending from –ite to –ous • Add the word acid HNO2 H2SO3 HClO2 nitrous acid sulfurous acid chlorous acid

  35. anion _______ide _______ate ________ite add hydrogen ions (H+) acid hydro____ic acid ______ic acid ______ous acid

  36. Metallic Bond, A Sea of Electrons

  37. Ionic Bonds: One Big Greedy Thief Dog!

  38. Polar Covalent Bonds: Unevenly matched, but willing to share.

  39. Metallic Bonds: Mellow dogs with plenty of bones to go around.

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