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Unit 2 - Bonding

Unit 2 - Bonding. Group. e - configuration. # of valence e -. ns 1. 1. I. II. ns 2. 2. III. ns 2 np 1. 3. IV. ns 2 np 2. 4. V. ns 2 np 3. 5. VI. ns 2 np 4. 6. VII. ns 2 np 5. 7. Valence electrons are the outer shell electrons of an

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Unit 2 - Bonding

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  1. Unit 2 - Bonding

  2. Group e- configuration # of valence e- ns1 1 I II ns2 2 III ns2np1 3 IV ns2np2 4 V ns2np3 5 VI ns2np4 6 VII ns2np5 7 Valence electrons are the outer shell electrons of an atom. The valence electrons are the electrons that particpate in chemical bonding. 9.1

  3. Lewis Dot Symbols for the Representative Elements & Noble Gases 9.1

  4. Trends in the Periodic Table • The elements have a regular (periodic) recurrence of physical and chemical properties • The horizontal rows are called periods • The vertical columns are called groups

  5. Atomic Size • Covalent Atomic Radius – half of the distance between the nuclei of two atoms in a homonuclear diatomic molecule • Group trend – atomic size increases as we move down a group • The farther away the e- from the nucleus, the less strongly they are held • Periodic trend – atomic size decreases as we move from left to right • Nuclear charge increases as we move left to right, pulling the e- closer to the nucleus

  6. Ionic Size • Cations are always smaller than neutral atoms because they have less e- • Loss of outer shell e- results in increased attraction by the nucleus for the remaining e- • Anions are always larger than neutral atoms because they have more e- • Addition of outer shell e- results in less attraction to the protons of the nucleus

  7. Ionization Energy • The energy required to remove an electron from a gaseous atom Na (g) + Energy Na+(g) + e- • Group trend – Ionization energy decreases as we move down a group because the size of the atom is increasing, with more orbitals of e-, allowing the outermost e- to be easily removed • Periodic trend – Ionization energy increases from left to right. Nuclear charge is increasing so more energy is required to remove e-

  8. Electron Affinity • The energy change (release) that accompanies the addition of an electron to a gaseous atom F (g) + e- F-(g) + Energy • Group trend – Electron affinity generally decreases with increasing atomic size (releases less energy as you go down) • Periodic trend – Generally increases from left to right because atoms become smaller due to increased nuclear charge (releases more energy as you go left to right)

  9. Electronegativity • The tendency for an atom to attract electrons to itself. It is affected by the distance from the atom’s valence electrons to the nucleus. • Group trend – Electronegativityusually decreases as you move down a group • Periodic trend – Electronegativityusually increases as you move across a period • Do the Noble Gases have electronegativity?

  10. Why do substances bond? • More stability • Atoms want to achieve a lower energy state

  11. Chemical bonds: an attempt to fill electron shells • Ionic bonds – • Covalent bonds – • Metallic bonds

  12. Ionic Bonding • Between a metal and a non-metal with very different electronegativity. • Metals lose electrons becoming a cations, while non-metals gain electrons becoming anions. • An ionic bond is an electrostatic attraction between the oppositely charged ions.

  13. Ionic Bonds: One Big Greedy Thief Dog!

  14. . Ionic bond– electron from Na is transferred to Cl, this causes a charge imbalance in each atom. The Na becomes(Na+)and the Cl becomes(Cl-),charged particles or ions.

  15. - - - - + Li+ Li Li Li+ + e- e- + Li+ Li+ + F F F F F F The Ionic Bond [He] [Ne] 1s22s1 1s22s22p5 1s2 1s22s22p6 9.2

  16. - - - - + Li+ Li Li Li+ + e- e- + Li+ Li+ + F F F F F F The Ionic Bond [He] [Ne] 1s22s1 1s22s22p5 1s2 1s22s22p6 9.2

  17. Ionic Structures • In an ionic compound (solid), the ions are packed together into a repeating array called a crystal lattice. • The simplest arrangement is one in which the spheres in the base are packed side by side. Opposite charges are attracted to each other. • Its called simple cubic packing (NaCl is an example) • Ionic formulas are always Empirical Formulas (simplest)

  18. Properties of Ionic Compounds • All ionic compounds form crystals. • Ionic compounds tend to have high melting and boiling points. • To break the positive and negative charges apart, it takes a huge amount of energy. • Ionic compounds are very hard and very brittle. • Again, this is because of the way that they're held together – strong attraction of oppositely charged ions. These ions simply don't move around - so they don't bend at all.  This also explains the brittleness of ionic compounds. If we give a big crystal a strong enough whack with a hammer, we usually end up using so much energy to break the crystal that the crystal doesn't break in just one spot, but in a whole bunch of places. Instead of a clean break, it shatters. 

  19. Ionic Compounds are poor conductors of electricity in the solid state. • Ions are held tightly together and cannot move. Therefore ions cannot conduct electricity • Ionic compounds conduct electricity when molten (melted). • Ions are mobile and can therefore conduct electricity. • Ionic compounds conduct electricity when aqueous (dissolved in water). • If we take a salt and dissolve it in water, the water molecules pull the positive and negative ions apart from each other. Instead of the ions being right next to each other, they are able to move around in the water and conduct electricity. • View here

  20. Covalent Bonding

  21. COVALENT BONDbond formed by the sharingof electrons

  22. Covalent Bond • Between nonmetallic elements of similar electronegativity. • Formed by sharing electron pairs • Stable non-ionizing particles, they are not conductors at any state • Examples; O2, CO2, C2H6, H2O, SiC

  23. Covalent Bonds

  24. Covalent bonds- Two atoms share one or more pairs of outer-shell electrons. Fluorine Atom Fluorine Atom Fluorine Molecule (F2)

  25. Why should two atoms share electrons? + 8e- 7e- 7e- 8e- F F F F F F F F lonepairs lonepairs single covalent bond single covalent bond lonepairs lonepairs A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Lewis structure of F2 Dot diagram of F2 9.4

  26. single covalent bonds H H H H or H H O 2e- 2e- O 8e- O C O C O O double bonds 8e- 8e- 8e- double bonds O N N triple bond N N triple bond 8e- 8e- Lewis structure of water + + Double bond – two atoms share two pairs of electrons or Triple bond – two atoms share three pairs of electrons or 9.4

  27. Properties of Covalent Compounds • Covalent compounds tend to have low melting and boiling points. • Many simple molecular substances are gases or liquids at room temperature. • Most molecular compounds are poor conductors of electricity in all states. • There are no free electrons available to move and conduct an electric current. • Covalent compounds are usually much softer than ionic material. • Covalent compounds tend to be flammable than ionic compounds. • Most molecular compounds are insoluble in water, but will dissolve in nonpolar organic solvents.

  28. 9.4

  29. Drawing Lewis Structures Write the dot diagram for each atom present in the compound. Take the TWO atoms with the MOST unpaired dots and join (bond) them together. Add the remaining atoms where they are needed the most. (starting with remaining atoms with most unpaired dots) Try the following: HOF N2H4 CH2O H4CO HCN

  30. Bonds in all the polyatomic ions and diatomics are all covalent bonds

  31. In covalent bonding, one or more pair of electrons are shared. However, all ‘sharing’ is NOT the same. Therefore, we can have Nonpolar (Pure) Covalent Bonds andPolar Covalent Bonds.

  32. NONPOLAR COVALENT BONDS when electrons are shared equally H2 or Cl2

  33. NonpolarCovalent bonds- Two atoms equally share one or more pairs of outer-shell electrons. Fluorine Atom Fluorine Atom Fluorine Molecule (F2)

  34. POLAR COVALENT BONDS when electrons are shared but shared unequally. H2O

  35. Polar Covalent Bonds: Unevenly matched, but willing to share.

  36. - water is a polarmolecule because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

  37. F H F H Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms (electrons are shared unequally) electron rich region electron poor region e- poor e- rich d+ d-

  38. To determine the type of Bonding present in a compound, we compare the electonegativities of the bonding atoms. The difference in electronegativity between these atoms predicts the type of bond present

  39. The Electronegativities of Common Elements

  40. F H Electronegativityis the ability of an atom to attract toward itself the electrons in a chemical bond. Electronegativity - relative, F is highest electron rich region electron poor region

  41. Increasing difference in electronegativity Nonpolar(Pure)Covalent Polar Covalent Ionic partial transfer of e- Unequal sharing transfer e- share e- equally Classification of bonds by difference in electronegativity Difference Bond Type 0  0.4 Nonpolar (Pure) Covalent > 1.7 Ionic 0.4 < and ≤1.7 Polar Covalent Creates a “Bonding Continuum”

  42. Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H2S; and the NN bond in H2NNH2. Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 NonpolarCovalent N – 3.0 N – 3.0 3.0 – 3.0 = 0 NonPolar Covalent 9.5

  43. Predicting Molecular Geometry • Draw Lewis structure for molecule. • Count number of lone pairs on the central atom and number of atoms bonded to the central atom. • Use VSEPR to predict the geometry of the molecule.

  44. # of atoms bonded tocentral atom # lone pairs on central atom Arrangement ofelectron pairs Molecular Geometry Class linear linear B B Valence shell electron pair repulsion (VSEPR) model: Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs. AB2 2 0

  45. Cl 0 lone pairs on central atom Be Cl 2 atoms bonded to central atom 10.1

  46. # of atoms bonded tocentral atom # lone pairs on central atom trigonal planar trigonal planar Arrangement ofelectron pairs Molecular Geometry Class VSEPR AB2 2 0 linear linear AB3 3 0

  47. # of atoms bonded tocentral atom # lone pairs on central atom trigonal planar trigonal planar AB3 3 0 Arrangement ofelectron pairs Molecular Geometry Class tetrahedral tetrahedral VSEPR AB2 2 0 linear linear AB4 4 0

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