Lesson 1 Intro to Chemical Bonding

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Covalent Bonding Unit …

Covalent Bonding Unit
Lesson 1 Intro to Bonding

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  • 1. Lesson 1: Introduction to Chemical Bonding Covalent Bonding Unit
  • 2. TargetsI can define chemical bond.I can describe covalent bonding.I can classify bonding type according to electronegativity differences.
  • 3. DefinitionsChemical Bond – mutual electrical attraction between the nuclei and valence electrons of different atoms that bind the atoms togetherValence electrons – outermost electrons that are available to be lost, gained, or shared to form a chemical bond
  • 4. Chemical Bond A force that holds groups of 2 or more atoms together and makes them function as a unit Atom – smallest unit of an element Molecule – Group of covalently bonded atoms Atoms Molecule
  • 5. Types of Chemical BondsIonic Bonding – (covered in next chapter) a type of bond in which a metal and a nonmetal transfer electronsCovalent Bonding – type of bond in which 2 or more nonmetal atoms share electrons
  • 6. IONIC – Metal + nonmetalPeriodic Table COVALENT – 2 nonmetals
  • 7. Types of Covalent BondsNonpolar covalent bond – electrons are shared equally
  • 8. Types of Chemical BondsPolar covalent – electrons are not shared equally because one atom attracts the shared electrons more than the other atom
  • 9. Bond TypesVideo
  • 10. ElectronegativityElectronegativity - measure of an atom’s ability to attract electrons.Electronegativities tend to increase across a period and decrease down a group
  • 11. Classifying Chemical BondsThe polarity of a bond depends on the difference between the electronegativity values of the atoms forming the bonds.Nonpolar covalent – 0 to 0.3Polar covalent – 0.4 to 1.7Ionic – greater than 1.8
  • 12. Electronegativity ValuesIncreases from left to right across a periodDecreases down a group of representative elements
  • 13. PracticeUse electronegativity values to classify the following bonds: a. Sulfur and Hydrogen b. Lithium and Fluorine c. Potassium and Chlorine d. Iodine and Bromine e. Carbon and Hydrogen
  • 14. PracticeUse electronegativity values to classify the following bonds: a. Sulfur and Hydrogen 2.5 – 2.1 = 0.4; polar covalent b. Lithium and Fluorine 4.0 – 1.0 = 3.0; Ionic c. Potassium and Chlorine 3.0 – 0.8 = 2.2; Ionic d. Iodine and Bromine 2.8 – 2.5 = 0.3; Nonpolar covalent e. Carbon and Hydrogen 2.5 – 2.1 = 0.4 ; polar covalent
  • 15. Covalent BondingCovalent Bonding and Molecular Compounds
  • 16. TargetsI can explain why most atoms form chemical bonds.I can explain the relationships among potential energy, distance between approaching atoms, bond length and bond energy.I can state the octet rule.I can determine the number of valence electrons for a given atom.
  • 17. Formation of a Covalent BondNature favors chemical bonding because most atoms have lower potential energy when they are bonded to other atoms.
  • 18. Formation of a Covalent BondEach atom has a positive nucleus in the center and negative electrons surrounding the nucleus in a spherical pattern.The positively charged nuclei are attracted to the negatively charged electrons.
  • 19. Formation of a Covalent BondAs the atoms approach each other, the charged particles interact: nucleus on one atom attracts electrons on the other atom.
  • 20. Formation of a Covalent BondAs the atoms approach one another, the potential energy decreases.A bond forms when the potential energy is at a minimum.
  • 21. Formation of a Covalent BondIf the atoms continue to approach one another once the bond forms, the nuclei will begin to repel one another and the potential energy will start to increase.
  • 22. Characteristics of the Covalent Bond Bond length – distance between two bonded atoms at their minimum potential energy or the average distance between two bonded atoms Bond energy – energy required to break a chemical bond and form neutral isolated atoms - kilojoules per mole (kJ/mol) Bond lengths and bond energies vary with the types of atoms that have combined
  • 23. The Octet RuleThe octet rule states that atoms tend to lose, gain or share electrons until they are surrounded by 8 electrons in their valence shell.The number of valence electrons is equal to the group number. (Groups 13-18; Group # -10)LABEL YOUR PERIODIC TABLE 1A 8A 2A 3A 4A 5A 6A 7A
  • 24. PracticeWhat is the relationship between bond energy and bond length?
  • 25. PracticeWhat is the relationship between bond energy and bond length? The bond length decreases as the strength of the bond increases.
  • 26. PracticeArrange the following in order of increasing bond strength: C–Cl, C–I, H–F, and I–ISKIP
  • 27. PracticeArrange the following in order of increasing bond strength: C–Cl, C–I, H–F, and I–I I-I, C-I, C-Cl, H-F
  • 28. Practice ProblemsWhich pair of bonded atoms has the strongest bond?
  • 29. Practice ProblemsWhich pair of bonded atoms has the strongest bond? H – F
  • 30. Practice ProblemsWhich pair of bonded atoms has the weakest bond?
  • 31. Practice ProblemsWhich pair of bonded atoms has the weakest bond? I – I
  • 32. Practice ProblemsArrange the following bond lengths in order of increasing bond strength: 72 pm, 149 pm, 53 pm, and 398 pm SKIP
  • 33. Practice ProblemsArrange the following bond lengths in order of increasing bond strength: 72 pm, 149 pm, 53 pm, and 398 pm 398 pm, 149 pm, 72 pm, 53 pm
  • 34. Practice ProblemsDetermine the number of valence electrons in each of the following atoms. Lithium Sulfur Carbon Neon
  • 35. Practice ProblemsDetermine the number of valence electrons in each of the following atoms. Lithium - 1 Sulfur - 6 Carbon -4 Neon - 8