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Ch. 12 Chemical Bonding

Explore the concepts of ionic, covalent, and polar covalent bonds, electronegativity, bond types, and Lewis structures in chemistry.

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Ch. 12 Chemical Bonding

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  1. Ch. 12 Chemical Bonding

  2. 12.1 Ionic, covalent and polar covalent bonds. • A bond is a force that holds atoms together. • Ionic Bonding a. When a metal reacts with a nonmetal, electrons are transferred from a metal to a nonmetal and an ionic compound is made.

  3. Ionic Bonding b. In ionic bonding, electrostatic attraction holds atoms together.

  4. Covalent Bonding a. Atoms make covalent bonds by sharing electrons. b. Electrons are attracted to the nucleus of both atoms in the bond. c. Nonmetals make covalent bonds. Let’s share! H H

  5. Electronegativity • Electronegativity is a measure of how strongly an atom attracts electrons. Look at the electronegativity chart on p403. • Fluorine has an electronegativity value of 4.0. • Hydrogen has an electronegativity value of 2.1. • Difference in electronegativity = 4.0 – 2.1 = ________

  6. Determined by difference in electronegativity values—absolute value of DEN Pauling Electronegativity Values Bond Types

  7. Electronegativity d. At left is a picture of hydrofluoric acid (HF). At right is a picture of “The Blob.” e. F has a higher electronegativity than H, so electrons are closer to F.

  8. In a covalent bond, atoms have a difference of electronegativity of 0-0.2. Electrons are shared equally. • In a polar covalent bond, atoms have a difference in electronegativity of 0.3-1.7. Electrons are not shared equally.

  9. Electronegativity • Exampe: Si—C • Electronegativity of Si is ________. Electronegativity of C is _______. • Difference in electronegativity is ____________. • Si – C bond is ( covalent / polar covalent ) [circle one]. • Which attracts more electrons? ( Si / C ) 1.8 2.5 0.7

  10. In the picture below, • Label each molecule or compound as ionic bonding, covalent bonding or polar covalent bonding. • For polar and ionic bonds, label the more electronegative atom. Covalent Polar covalent Ionic

  11. Identify each of the following bonds as ionic, covalent or polar covalent.

  12. Dipoles • In a polar molecule, one side has a partial positive charge, and the other has a partial negative charge. • A dipole moment is represented with an arrow pointing to the negative side and the Greek letter “delta” δ to show the partial positive and negative charges:

  13. Dipoles • Write the partial charges and draw the dipole moment on Cl –I

  14. 12.2 Ionic Bonding • Ions • Metals ( lose / gain ) electrons. • Nonmetals ( lose / gain ) electrons. • Group 1 elements form ions with a charge of ____. • Group 2 elements form ions with a charge of ____. • Group 6 elements form ions with a charge of ____. • Group 7elements form ions with a charge of ____.

  15. Ions • That’s interesting…but WHY??? • Atoms gain or lose electrons to get the electron configuration of a noble gas. • Noble gases have completely filled energy levels, so they are very stable. • He has a completely filled 2s sublevel. • Other noble gases have filled s and p sublevels.

  16. Ions • Example: Li • The electron configuration of Li is ______ • Li loses one 2s electron and becomes Li+. • The electron configuration of Li+ is _____ • Li + has the same electron configuration as He.

  17. Ions • The electron configuration of F- • The electron configuration of F is • F gains one 2p electron and becomes F- • The electron configuration of F- is • F- has the same electron configuration as ______

  18. Ions • What would happen if Li reacted with F? • Li gives an electron to F • Li + F  Li+ + F- • And they form _______ (write the formula of the compound) • Ionic bonding and structures • LiF is packed together in a group in order to maximize attractions of the cations and anions. • It makes a hard, tight crystal.

  19. The size of ions • Which is larger, Na or Na+? Why? • Na. Na loses a 3s electron, and then only has electrons in the n=2 level. n=2 orbitals are smaller than n=3 orbitals. Na Na+ F- F

  20. The size of ions • Which is larger, F or F-? Why? • F- because it gains electrons. • Which is larger F- or Na+? Why? • F-. They both have the electron configuration of Ne, but Na+ has more protons (a stronger + charge), which pulls electrons closer. • Which is larger, Ca or Ca2+? Why?

  21. Ionic SizeTaken from: http://www.chem.umass.edu/people/botch/Chem121F06/Chapters/Ch15/IonicRadii.jpg

  22. 12.3 Lewis Structures • The octet rule: Sharing of electrons usually occurs so that atoms acquire the electron configurations of noble gases (1s2 or ns2np6)

  23. Lewis dot structures • The element symbol represents the core electrons. • Dots to show the valenceelectrons.

  24. Lewis dot structures • 1. First, write the symbol for the element. • 2. Imagine the molecule has four sides (but don’t draw the “x”) • 3. Draw one dot at a time in each empty section. • 4. You should only have pairs of e- if there are no empty sections. Cl

  25. Lewis dot structures • The “paired” electrons cannot usually make bonds. • The three “unpaired” electrons can make bonds.

  26. Dot Diagram? In order to get the configuration of a noble gas, how many bonds will arsenic form? Dot Diagram for Hydrogen? In order to get the configuration of a noble gas, how many bonds will H form? H Arsenic 3 1

  27. Arsenic Trihydride, AsH3 Lewis dot structure Structural formula As H H As H H H H As

  28. Drawing Lewis Dot Structures • Use the molecular formula to find the total number of valence electrons • 1 + 1 + 1 + 5 = 8 H H H As

  29. Drawing Lewis Dot Structures 2. Draw the symbols of each element. Draw the backbone/skeleton structure • Terminal atoms go around the central atoms. • Least Electronegative is usually central • Hydrogen is always terminal. • Carbon is usually central • Oxygen and Halogens are usually terminal

  30. Drawing Lewis Dot Structures 3. Draw the valence electrons 4. Make single bonds with pairs

  31. Drawing Lewis Dot Structures 5. Each pair represents a bond. 6. Count electrons again and make sure you get the same number (8). Single bond

  32. O O N N Structural Formula • The structural formula is drawn with a “–“ line to represent the bond. • Double and triple bonds • Double bonds can be formed by sharing two pairs of electrons • Triple bonds involve sharing three pairs of electrons.

  33. O O O H H Cl Cl B Cl I I Cl I O O Si I I I Si Cl Cl B I I O H H Try the Dot Diagrams for each of the following: O2 H2O BCl3 SiI4 Each pair of shared electrons can be represented by a line.

  34. Structural formulas for polyatomic ions • Example: NO3- • The negative charge means there is one more electron • How many total electrons are there? • 5 + 3(6) + 1 = 24 electrons. • NO3- • Draw:

  35. O O O O O O Resonance • Resonance • Example: ozone • Ozone’s structure can be drawn like this: • The actual bonding in ozone is not like either of these structures. The actual structure lies somewhere in between these two. • These drawings are resonance structures of ozone

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