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  1. Bonding Forces of attraction that hold atoms together making compounds

  2. Chemical symbols • Symbols are used to represent elements • Either one capital letter, or a capital letter with a lower case letter • Know names and symbols of elements: • 1 – 30, plus • Rb, Cs, Sr, Ba, Ag, Au, Cd, Hg, Pt, Ga, Ge, As, Sn, Pb, Se, Br, I, and U

  3. Basic idea... • All chemical bonds form because they impart stability to the atoms involved • lower energy = greater stability

  4. Quick review • All types of chemical bonds involve electrons • Valence electrons, the electrons in the outermost occupied energy level of an atom, are usually the electrons involved in bonding

  5. The representative elements have the same number of valence electrons as their family number in the American system • Example: Mg, column IIA, 2 valence electrons • The transition metals all have two valence electrons • ns2(n-1)dx

  6. Lewis dot structures are used to represent the valence electrons • each dot represents a valence electron • no more than 8 dots total • no more than 2 dots on a side • example = Mg: Na. .

  7. Lewis dot structures of representative elements

  8. The Octet Rule • Atoms will gain, lose, or share electrons in order to achieve an ns2np6 valence configuration

  9. Sizes of atoms • Periodic trend: atomic radii increase moving down a group • Increasing energy level • Periodic trend: atomic radii decrease moving left to right in a period • The charge felt by the valence electrons becomes larger

  10. Sizes of atoms • There is a general decrease in atomic radius from left to right, caused by increasing positive charge in the nucleus. • Valence electrons are not shielded from the increasing nuclear charge because no additional electrons come between the nucleus and the valence electrons.

  11. For metals, atomic radius is half the distance between adjacent nuclei in a crystal of the element. • For elements that occur as molecules, the atomic radius is half the distance between nuclei of identical atoms.

  12. Atomic Radius

  13. Atomic Radius • Atomic radius generally increases as you move down a group. • The outermost orbital size increases down a group, making the atom larger.

  14. Sizes of ions • Periodic trend: anions are always larger than the atom they were formed from • Electrons repel each other • Periodic trend: cations are always smaller than the atom they were formed from • Fewer electrons to share same positive nuclear charge

  15. Ionic Radius • When atoms loseelectrons and form positively charged ions, they always become smaller for two reasons: The loss of a valence electron can leave an empty outer orbital resulting in a small radius. Electrostatic repulsion decreases allowing the electrons to be pulled closer to the radius.

  16. Ionic Radius • When atoms gain electrons, they can become larger, because the addition of an electron increases electrostatic repulsion.

  17. Ionic Radius • Both positive and negative ions increase in size moving down a group.

  18. Ionic Radius • The ionic radii of positive ions generally decrease from left to right. • The ionic radii of negative ions generally decrease from left to right, beginning with group 15 or 16.

  19. Bonding Forces of attraction that hold atoms together making compounds

  20. Ionization energy • The energy needed to remove a valence electron from an atom • A measure of how tightly the electrons are being held • periodic trend • increases from the bottom up • increases left to right

  21. In general, metals have lower IE than nonmetals • alkali metals are the lowest IE family • noble gases are highest IE family

  22. Ionization energy • The energy required to remove the first electron is called the first ionization energy. • First ionization energy increases from left to right across a period. • First ionization energy decreases down a group because atomic size increases and less energy is required to remove an electron farther from the nucleus.

  23. Ionization energy

  24. Ionization energy

  25. Ionization energy • Removing the second electron requires more energy, and is called the second ionization energy. • Each successive ionization requires more energy, but it is not a steady increase. • The ionization at which the large increase in energy occurs is related to the number of valence electrons.

  26. Ionization energy

  27. Electron affinity • A measure of how strongly an element would like to gain an electron • periodic trend • increases from the bottom up • increases left to right • ignore the noble gases

  28. Atoms that lose electrons easily have little attraction for additional electrons (and vice versa) • metals have low IE, low EA • Nonmetals have high IE, high EA • Octet rule: when atoms react, they tend to strive to achieve a configuration having 8 valence electrons • This results in some form of bond formation

  29. Periodic trends… • As you move from left to right along a period… • Atoms get …. Smaller • Ionization energy goes …. Up • Electron affinity goes …. Up

  30. Periodic trends… • As you move down a group/family • Atoms get …. Larger • Ionization energy goes …. Down • Electron affinity goes …. Down

  31. Check your understanding The lowest ionization energy is the ____. A.first B.second C.third D.fourth

  32. Check your understanding The ionic radius of a negative ion becomes larger when: A.moving up a group B.moving right to left across period C.moving down a group D.the ion loses electrons

  33. Electron Configuration of Ions • Na 1s22s22p63s1 • will lose one e- to gain ns2np6 configuration • Na+ 1s22s22p6 • S 1s22s22p63s23p4 • will gain 2 e- to gain ns2np6 configuration • S2-  1s22s22p63s23p6

  34. Ionic Bonding • Metals lose electrons easily, nonmetals have a strong attraction for more electrons • metal atoms will lose electrons to nonmetal atoms, causing both to become ions

  35. Metals, having lost one or more electrons, become cations (+) • Nonmetals, having gained one or more electrons, become anions (-) • Opposites attract: the cations and anions are held together electrostaticly • called “ionic bonds”

  36. In summary... • Ionic bonds are electrostatic attractions between cations and anions formed when electron(s) are transferred from the low IE, EA metal to the high IE, EA nonmetal

  37. Ionic compound = crystalline solid Cation (+)

  38. Ionic Compounds • High melting points • brittle solids • nonconducting as solids • conduct electricity as liquids or aqueous

  39. Ionic Compounds • As solids, exist in a 3-D repeating pattern called a crystal “lattice” • the lattice energy is the energy lowering (stability) accomplished by the formation from “free” ions • Also a measure of the energy required to break apart the ionic compound once formed • The greater the lattice energy, the stronger the force of attraction

  40. Bonding Forces of attraction that hold atoms together making compounds

  41. Ion dissociation • Many ionic compounds will dissolve in water if it results in more stability (lower E) than in the solid ionic compound • the ions “dissociate” from each other • Ex: CaCl2(s) + H2O  Ca2+(aq) + 2Cl-(aq)

  42. Ionic Bond Strength • A measure of the attractive force between the ions • smaller atoms = stronger ionic bonds • fewer atom ratio = stronger bond • evidence: melting points

  43. Compare the melting points: • KCl : 776oC • KI : 723oC • smaller atoms result in stronger ionic bonds

  44. Compare the melting points: • CaCl2 : 772oC • NaCl : 800oC • fewer atoms result in stronger ionic bonds

  45. Bonding Forces of attraction that hold atoms together making compounds

  46. Covalent Bonding • Covalent bonding involves the sharing of electron pairs • usually between two high EA, high IE nonmetals • both want more e-’s, neither is willing to lose the e-’s they have

  47. A nonmetal will form as many covalent bonds as necessary to fulfill the octet rule • example: C, with 4 valence e-’s, will form 4 covalent bonds • results in 8 valence e-’s around the carbon atom at least part of the time • double and triple covalent bonding is a possibility

  48. When does the octet rule fail?

  49. H, He and Li • Helium strives for 2 valence electrons • 1s2 configuration • Hydrogen will sometimes will share its one electron with another atom, forming a single covalent bond • Lithium will lose its lone valence electron, gaining the 1s2 configuration of He

  50. Be • Be will sometimes lose its 2 valence electrons, gaining the Is2 configuration of He • Be will sometimes form 2 covalent bonds, giving it 4 valence electrons • nuclear charge of +4 cannot handle 8 valence electrons