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Chemical Bonds Section 1 Introduction to Chemical Bonding

Chemical Bonds Section 1 Introduction to Chemical Bonding. Objectives. Define chemical bond. Explain why most atoms form chemical bonds. Describe ionic and covalent bonding. Explain why most chemical bonding is neither purely ionic or covalent.

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Chemical Bonds Section 1 Introduction to Chemical Bonding

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  1. Chemical BondsSection 1 Introduction to Chemical Bonding

  2. Objectives • Define chemical bond. • Explain why most atoms form chemical bonds. • Describe ionic and covalent bonding. • Explain why most chemical bonding is neither purely ionic or covalent. • Classify bonding type according to electronegativity differences.

  3. Atomic Stability Q. Why do atoms form compounds? • Search for outer shell stability by losing or gaining electrons! Nobel Gases  8 electrons in os = stable Ex: Na loses 1 e- to Cl Chemical Bond  attractive force that holds atoms together in a compound; when atoms gain, lose, or share e-, an attractive force pulls them together to form a compound

  4. Oxidation Number  the number that tells you how many electrons an atom has gained, lost or shared to become stable

  5. Ion • Def. Charged particle that has either gained or loses electrons • RESULT: Now has either more or fewer number of electrons than protons • Ex: sodium fluoride NaF active ingredient in toothpaste  Na loses 1e- to F; Na is +1; F is -1 • Ex: potassium iodide KI – ingredient in iodized salt  K loses 1 e -to I; K is +1; I is -1 • Cation = positive ion #P > #E • Anion = negative ion #P < #E

  6. Ionic Bond • Def. Force of attraction between the opposite charges of the ions in an ionic compound (cations and anions) • Transfer of electrons btwn metals & nonmetals • Ex: magnesium chloride MgCl2 Mg loses 2 electrons to each Cl; Mg is +2; each Cl is -1 • RESULT neutral compound = sum of the charges of the ions equals 0 2(Mg) + -1(Cl) + -1(Cl) = 0

  7. Covalent Bond • Def. Attraction that forms between atoms when they share electrons • Atoms will be more stable by sharing e- rather than losing or gaining e-

  8. Unequal Sharing • Electrons don’t always share equally between atoms in a covalent bond Strength of attraction of atom to electrons due to: • 1. size of atom • 2. size of the positive charge in the nucleus (a strong magnet will hold a metal better than a weak magnet) • 3. total # of electrons • 4. how far are the electrons from the nucleus being shared (a magnet has a stronger pull to a metal when it is next to it rather than a couple inches away)

  9. Ex: HCl hydrochloric acid used to clean metal and found in your stomach to digest food • Cl – atoms have a stronger attraction for electrons than H atoms  electrons shared will spend most time near the chlorine atom • RESULT Cl atom has a partial negative charge (Greek delta)  H atom has a partial positive charge • VISUAL: Tug-of-war  stronger team pulls the rope towards them

  10. Polar or Nonpolar • Polar molecule  molecule that has a slightly positive end and a slightly negative end although the overall molecule is neutral ex: water • Nonpolar molecule  molecule in which electrons are shared equal in bonds; doesn’t have oppositely charged ends; found in 2 identical atoms or molecules that are symmetric ex: CCl4

  11. Ionic or Covalent??? • Bonding between atoms of different elements is rarely purely ionic or purely covalent. • Falls somewhere between 2 extremes depending on electronegativity measure of atom’s ability to attract e-

  12. Electronegativity Values

  13. less than 0.3 0.3-1.7 greater than 1.7

  14. If you still need help understand polarity and electronegativity, watch this video: http://www.bing.com/videos/search?q=Examples+Polar+and+Nonpolar+Covalent+Bonds+With&Form=VQFRVP#view=detail&mid=295CAC087126B74D1565295CAC087126B74D1565

  15. HOMEWORK Section Review pg 177 #1-5

  16. Objectives • Define molecules and molecular formula. • Explain relationships among potential energy, bond length, and bond energy. • Learn the basic steps used in writing Lewis structures. • Explain how resonance structures are used to represent molecules.

  17. Section 2 Covalent Bonding Molecular Compounds • molecule neutral group of atoms that are held together by covalent bonds • Red = O, White = H, Black = C • Molecular compound a chemical compound whose simplest units are molecules

  18. chemical formula indicates what elements atoms and numbers of atoms in a chemical compound by using atomic symbols and numerical subscripts • molecular formula shows types and number of atoms combined in a single molecule of a molecular compound

  19. Formation of a Covalent Bond • Nature favors chemical bonding most atoms have lower potential energy when bonded to other atoms than as independent atoms. • separated H atoms do not affect each other • PE decreases as atoms are drawn together by attractive forces • PE minimum when attractive forces are balanced by repulsion forces = ideal distance • PE increases when repulsion btwn like charges outweighs attraction between opposite charges

  20. Characteristic of the Covalent Bond • Bond length average distance between 2 bonded atoms H-H 75 pm • Form CB= H atoms release energy= amt of energy equals drop in PE • Bond energy energy required to break a chemical bond and form neutral isolated atoms (kJ/mol)

  21. Octet Rule • Chemical compounds tend to form so that each atom has an octet of e- in outer energy level by gaining, losing or sharing e- • Draw Fluorine electron configurations:

  22. Objectives • Learn the basic steps used in writing Lewis structures. • Explain how resonance structures are used to represent molecules.

  23. Lewis Dot Diagrams Electron Dot Structure or Lewis Dot Diagram (Gilbert Lewis) • Def. A notation showing the valence electrons (electrons in outer energy level) surrounding the atomic symbol.

  24. Lewis Structures • 1)Write the element symbol. • 2)Carbon is in the 4th group, so it has 4 valence electrons. • 3)Starting at the right, draw 4 electrons, or dots, counter-clockwise around the element symbol. On your sheet, try these elements on your own: • a)H b)P c)Ca d)Ar e)Cl f)Al

  25. unshared pair- (lone pair) e- not involved in bonding and belong to one atom • Structural formula- shows bonds but not unshared pairs of e- in molecule

  26. Single Covalent Bond • Made of 2 shared electrons • 1 comes from one atom in the bond and 1 comes from the other atom in the bond • Ex: water – O now is stable with 8 e in outer shell and H is stable with 2

  27. Multiple Bonds • Bonds with multiple pairs of shared electrons • Ex: N2 Nitrogen has 5 e in os and needs 3 to be stable  shares 3 e (triple bond)

  28. Lewis Structures for Molecules Draw the LS for CH3I 1. Write the LS for each atom in the molecule. 2. Determine the total number of valence e- available. 3. Arrange atoms to form skeleton structure of molecule. When present C is central atom. Otherwise, least EN atom is central (Except H). 4. Add unshared pairs of e- to each nonmetal (except H) such that each is surrounded by 8. 5. Count e- to be sure that # of VE=number available. Check to see that atoms have octet.

  29. Practice • Draw and build the Lewis Structure of NH3 • Draw and build the Lewis Structure of H2S • Draw and build the Lewis Structure of SiH4 • H-white, N-orange (should only have 3 holes) S-red, Si-black

  30. Resonance Structures • def. bonding in molecules or ions that cannot be correctly represented by a single Lewis structure • resonance aka hybrids constantly alternating from one form to another O3 has a single structure that is avg of 2 structures • use double arrow to indicate resonance

  31. Draw the 3 Resonance Structures for SO3

  32. Homework • Electron Dot Diagrams and Lewis Structures Worksheet

  33. Objectives • Compare and contrast a chemical formula of molecular and ionic compounds • Compare and contrast properties of ionic and molecular compounds • Write Lewis structures for polyatomic ions

  34. Section 3 Ionic Bonding and Ionic Compounds • Ionic compound def. composed of + and – ions combined so that # of + and – charges are equal • most exist as crystalline solids= 3D structure of +/- ions attracted to each other • Formula unit def. simplest unit of atoms from which ionic compound can be established ex: NaCl • http://fikus.omska.cz/~bojkovsm/termodynamika/Obrazky/bindingstyper.swf

  35. Formation of Ionic Compounds

  36. Show formation of …KF, potassium fluorideNa20, sodium oxide

  37. Characteristic of Ionic Bonding *Remember: Nature favors arrangements where potential energy is minimized. In an ionic crystal, ions minimize PE by combining in orderly arrangement crystal lattice

  38. Lattice Energy def. energy released when one mole of an ionic crystalline compound is formed from gaseous ions negative values mean energy released when crystals are formed

  39. Polyatomic Ions def. a charged group of covalently bonded atoms • combine w/ ions of opposite charge to form ionic compounds • https://www.youtube.com/watch?v=A9qzV1j3NuE

  40. Metallic Bonding • Occurs bc e- move freely among a metal’s + charged ions e- form a cloud around the metal ion • RESULT: ductility & malleable • Ex: metal hammered into sheets doesn‘t break bc ions are in layers that slide past each other w/out losing their attraction to the electron cloud • Ex: good conductor of electricity bc outer-level electrons are held weakly • http://cd1.edb.hkedcity.net/cd/science/chemistry/resource/animations/metallic_bond/metallic.html • https://www.ck12.org/physical-science/Metallic-Bonding-in-Physical-Science/enrichment/Metallic-Bond-Animation/?referrer=concept_details

  41. Metallic Bonding • Ductility- ability of a substance to be drawn, pulled, or extruded through a small opening to produce a wire • Malleability- ability of a substance to be hammered into thin sheets

  42. Homework • Chemical Speed Dating Profile • Chapter Review

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