1 / 74

Bonding Between Atoms

Bonding Between Atoms. Why do Atoms Form Bonds?. To get a stable octet of valence electrons. Called a “noble gas configuration”. Changes in Energy. Energy is released when bonds form. (exothermic) Results in lower energy, more stability Know example…. Bonding Type #1. Ionic Bonding.

richeyj
Télécharger la présentation

Bonding Between Atoms

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


  1. Bonding Between Atoms

  2. Why do Atoms Form Bonds? • To get a stable octet of valence electrons. • Called a “noble gas configuration”

  3. Changes in Energy • Energy is released when bonds form. (exothermic) • Results in lower energy, more stability • Know example…

  4. Bonding Type #1 Ionic Bonding

  5. Atoms Become Ions • Lose electrons, become a (+) ion (cation) • Ex: Mg is 2-8-2 • Loses 2 electrons to become Mg+2 2-8

  6. Gain electrons, become a (–) ion (anion) • Ex: Cl is 2-8-7 • Gains 1 electron to become Cl-1 2-8-8

  7. Metals lose electrons (form + ions) • Nonmetals gain them (form – ions)

  8. Electron Transfer • As ions form, an exchange or transfer of electrons happens. Lithium (metal) 2-1 Fluorine (nonmetal) 2-7

  9. Electronegativity • Atoms in ionic compounds have a large difference in their EN values > 1.7

  10. Ex: EN Values • Metal Na = 0.9 • Nonmetal Cl = 3.2 • Nonmetal with higher EN “takes” electron(s) from the metal

  11. A positively charged ion and a negatively charged ion attract each other. • This attraction forms an IONIC BOND - ion + ion

  12. Ionic compounds have a uniform crystalline lattice structure.

  13. Classic Example of Ionic Bond Formation • http://youtu.be/xTx_DWboEVs

  14. Properties of Ionic Compounds

  15. High Melting Point Attraction between the ions is very strong. Requires a lot of heat energy to separate ions and make solid crystal melt.

  16. Solubility in Water • Most ionics will dissolve in water, or be “soluble”. • When dissolved in water they are “aqueous” Ex: NaCl (aq)

  17. Water is a “polar molecule”. • Acts like a magnet to pull ions apart and into solution. • Ions are now “dissociated” or “hydrated” ions.

  18. http://youtu.be/gN9euz9jzwc • http://youtu.be/EBfGcTAJF4o

  19. Conductivity • Ionic compounds conduct when ions are “mobile” or free to move about.

  20. Conductivity • Ionic compounds conduct when: • Molten (melted/liquid) (l) • Aqueous (aq) Note: They DO NOT conduct when solid as the ions are locked in place.

  21. Types of Ionic Compounds • Binary: Contain 2 elements Ex: MgCl2, Al2O3, NaCl

  22. Ternary: Contain 3 elements • Polyatomic ion present. • Contain both ionic and covalent bonds!! Ex: NaNO3, Ca3(PO4)2, NH4Cl Elements inside the polyatomic ion are covalently bonded (all nonmetals).

  23. Ionic Compounds are Neutral • Criss-Cross charges if necessary to balance the formula

  24. Roman Numeral (Stock System) • Roman numeral is used in the name of ionic compounds in which the metal can have more than one possible charge. Ex: NiBr2 Nickel II Bromide NiBr3 Nickel III Bromide

  25. Bonding Type #2

  26. Covalent Bonding • Share valence electrons between atoms • Electron clouds overlap

  27. Happens between nonmetals Ex: H2O CH4 C6H12O6 NH3 CO2

  28. Electronegativity • Difference in EN is smaller than in ionics and is usually < 1.7 • Ex: HCl H = 2.1 Cl = 3.2 Difference = 1.1

  29. Molecular Formulas • All covalent compounds are called molecules. • Molecular formulas: show actual number of atoms of each element present in compound Ex: H2O 2 hydrogen atoms and 1 oxygen

  30. Empirical Formulas • Empirical formulas: • Show simplest whole number ratio of atoms or ions in the compound. • All ionic compounds have empirical formulas • Ex: MgCl2 1 : 2 ion ratio

  31. You can simplify some molecular formulas to make them empirical ratios Ex: C6H12O6 Simplest ratio of atoms CH2O

  32. Structural Formulas • Show how the atoms are bonded together in a molecule. • Use “lines” to show covalent bonds

  33. Molecule vs. Ionic Crystal CH4 = 5 atoms in molecule NaCl = 1:1 ion ratio

  34. Ionic Character • Note: The greater the EN difference is between atoms the more “Ionic Character” the bond has.

  35. Single, Double, Triple Bonds • Atoms can share single double or triple bonds between them. • Each bond represents a shared pair of electrons. • http://youtu.be/1wpDicW_MQQ

  36. Bond Polarity • Polar Bond: when there is a difference in EN values. (unequal sharing) • Ex: H Cl EN=2.1 EN=3.2

  37. NonPolar Bond: no difference in EN values. (equal sharing) Ex: O2, N2, Cl2, H2 (all the diatomics!)

  38. Shapes of Molecules

  39. VSPER • Valence Shell Electron Pair Repulsion • Valence electrons will orient themselves around the “central” atom to be as far apart from each other as possible. • This influences the “shape” of the molecule.

  40. Polarity of a Molecule • Polar Molecules: • Have polar bonds and are not symmetrical • Nonpolar Molecules • Have nonpolar bonds OR • Have polar bonds and are symmetrical

  41. Tetrahedral • Has 4 atoms bonded (no free pairs)

  42. Symmetry? Depends on what atoms are attached. Can be polar (asymmetrical) or nonpolar (symmetrical)

  43. Pyramidal • Three atoms bonded (one free pair)

  44. Symmetry? • All pyramids are asymmetrical. • These molecules are always POLAR!

  45. Bent • Two atoms attached (2 free pair) The 2 free pair make it bent and not linear. These are always asymmetrical so are always polar. H2O

  46. Hey, Water is Polar!!!!! Never forget this!!!

More Related