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CHEMICAL BONDS

CHEMICAL BONDS. The forces that hold two or more atoms together and make them function as a unit . A chemical bond is based on the electronegativity difference between the atoms involved.

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CHEMICAL BONDS

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  1. CHEMICAL BONDS The forces that hold two or more atoms together and make them function as a unit. A chemical bond is based on the electronegativity difference between the atoms involved.

  2. Chemical bonds form so that each atom has an “octet” of electrons in its highest occupied energy level. In Ionic bonds – atoms transfer electrons In Covalent bonds – atoms share electrons

  3. IONIC BONDS AND Ions • Ionic bonds form between a metal and a non-metal • Metals lose electrons to form positively-charged ions called cations • Nonmetals gain electrons to form negatively-charged ions called anions

  4. Ionic Bonding

  5. The Formation of Sodium Chloride • Sodium has 1 valence electron • Chlorine has 7 valence electrons • An electron transferred from sodium to chlorine, gives each an octet Na:1s22s22p63s1 Cl: 1s22s22p63s23p5

  6. The Formation of Sodium Chloride This transfer forms ions, each with an octet in the outermost shell: sodium ion,Na+ : 1s22s22p6 chloride ion, Cl- : 1s22s22p63s23p6

  7. The Formation of Sodium Chloride The resulting ions come together due to electrostatic attraction (opposites attract): Na+ Cl- The net charge on the compound must equal zero.

  8. Predicting Ionic Charges Group 1: Lose 1 electron to form 1+ ions

  9. Predicting Ionic Charges Group 2: Lose 2 electrons to form 2+ ions

  10. Predicting Ionic Charges Lose 3 electrons to form 3+ ions Group 13:

  11. Predicting Ionic Charges Group 14: Different possibilities depending on the element

  12. Predicting Ionic Charges Group 15: Gain 3 electrons to form 3- ions

  13. Predicting Ionic Charges Gain 2 electrons to form 2- ions Group 16:

  14. Predicting Ionic Charges Group 17: Gain 1 electron to form 1- ions

  15. Predicting Ionic Charges Stable Noble gases do not form ions! Group 18:

  16. Predicting Ionic Charges Many transition metals have more than one possible valence number. (“multivalent” elements) Groups 3 - 12: Fe3+ Fe2+

  17. Sodium Chloride Crystal Lattice Ionic compounds form solid crystalsat ordinary temperatures. All saltsare ionic compoundsand formcrystals.

  18. Ionic Compounds are composed of Formula Units, not Molecules! It is impossible to identify which pairs of adjacent ions participated in transferring electrons Therefore, we call the smallest particle of an ionic compound a “formula unit”, instead of a “molecule” ?

  19. Properties of Ionic Compounds

  20. Lewis Dot Diagrams Unlike Bohr-Rutherford drawings, Lewis Dot diagrams only show the valence electrons

  21. [Na]1+ Lewis dot diagrams for ions ▪Mg▪ [Mg]2+ An ion is indicated by using square brackets around the symbol and placing the charge outside of the brackets Since metal ions have lost their valence electrons, no dots are required Since non-metal ions gain electrons to complete their octet, they will generally have 8 dots ▪N▪ [:N:]3- ▪▪ ▪▪ ▪ ▪▪

  22. Covalent bonds form between two or more non-metals Electrons are shared between atoms As many as 3 covalent bonds can form between 2 atoms, so that both atoms have a stable octet The term “molecule” is used exclusively for covalent bonding Covalent Bonds

  23. The HONCRule • Hydrogen (and Halogens) form one covalent bond • Oxygen (and sulfur) form two covalent bonds • One double bond, or two single bonds • Nitrogen (and phosphorus) form three covalent bonds • One triple bond, or three single bonds, or one double bond and a single bond • Carbon (and silicon) form four covalent bonds. • Two double bonds, or four single bonds, or a triple and a single, or a double and two singles The HONC Rule DOES NOT work for polyatomic ions!

  24. RULES FOR COMPLETING LEWIS DOT DIAGRAMS FOR COVALENT MOLECULES AND POLYATOMIC IONS • Arrange atoms symmetrically around the central atom (the atom that occurs the least and usually the atom with the lowest electronegativity). • The central atom is rarely oxygen and NEVER hydrogen. • Add the number of valence electrons for all atoms in the compound. • Add additional electrons for a negative ion, or subtract electrons for a positive ion.

  25. Place a pair of electrons (i.e. a bonding pair) between central atom and each of the surrounding atoms. • Complete the octets of the surrounding atoms using lone pairs of electrons. (Remember hydrogen only needs 2 electrons to complete its outer shell). • Any remaining electrons go on the central atom (as lone pairs). • If the central atom does not have an octet, move lone pairs from the surrounding atoms to form double or triple bonds as required. • Put square brackets around polyatomic ions showing the charge outside the brackets.

  26. PROPERTIES OF COVALENT COMPOUNDS

  27. Polar covalent bonds • The electrons in a covalent bond are not always shared equally • How electrons are shared depends on differences in electronegativity and molecular structure • Unequal sharing of bonding pairs results in a polar covalent bond and can result in a polar covalent molecule, where one end is slightly negative (δ-) and the other end is slightly positive (δ+) • The direction of polarity is indicated by a “dipole” vector, which points towards the negative end δ+δ-

  28. Polar covalent bonds • EXAMPLE 1: Hydrogen gas ENH = 2.1 ENH2 = 2.1 – 2.1 = 0 H – H • Non-polar covalent bond • Non-polar covalent molecule • EXAMPLE 2: Hydrogen chloride ENH = 2.1 ENCl = 3.0 ENHCl = 3.0 – 2.1 = 0.9 δ + δ- H – Cl: • Polar covalent bond • Polar covalent molecule : :

  29. POLAR COVALENT BONDS • EXAMPLE 3: Carbon monoxide ENC = 2.5 ENO = 3.5 ENCO = 3.5 – 2.5 = 1.0 δ+ δ- :C  O: • Polar covalent bond • Polar covalent molecule • EXAMPLE 4: Carbon dioxide ENC = 2.5 ENO = 3.5 ENCO = 3.5 – 2.5 = 1.0 δ-δ+ δ- . . . . O = C = O ' ' ' ' • Polar covalent bonds • Equal and opposite forces cancel • Non-polar molecule

  30. Polar covalent bonds EXAMPLE 5: Water ENH = 2.1 ENO = 3.5 ENH2O = 3.5 – 2.1 = 1.4 O H H • Polar covalent bonds • Non-symmetrical molecule • Polar covalent molecule EXAMPLE 6: Ammonia ENN = 3.0 ENH = 2.1 ENNH3 = 3.0 – 2.1 = 0.9 . . N H H H • Polar covalent bonds • Non-symmetrical molecule • Polar covalent molecule δ- δ- •• •• δ+ δ+

  31. The bonding continuum The type of bond that forms between two atoms is determined by their difference in ELECTRONEGATIVITY (EN) TYPE OF BOND FORMED non-polar slightly polar polar highly polar covalent covalentcovalentcovalentionic 0 0.5 1.0 1.5 1.7 2.0 ∆EN

  32. Intermolecular forces • By far the strongest intermolecular forces are the electrostatic attractions within ionic compounds • Covalent compounds have much weaker intermolecular forces (IMF’s) • The strongest IMF’s in covalent compounds occur in polar covalent compounds (dipole-dipole attractions)

  33. Dipole-dipole attractions in water • Many of the unique characteristics of water are due to the fact that it is a very polar molecule • Water is known as the “universal solvent • Ice is less dense than liquid water • Capillary action • These powerful IMF’s are known as “hydrogen bonds”

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