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Chapter 6

Chapter 6. Chemical Bonding. Introduction to Chemical Bonding. Atoms seldom exist as independent particles in nature--instead they are combinations of atoms that are held together by chemical bonds.

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Chapter 6

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  1. Chapter 6 Chemical Bonding

  2. Introduction to Chemical Bonding • Atoms seldom exist as independent particles in nature--instead they are combinations of atoms that are held together by chemical bonds.

  3. A chemical bond is a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together. By bonding, the atoms decrease in potential energy, thereby creating more stable arrangements of matter.

  4. Types of chemical bonding: • Ionic bonding: chemical bonding that results from the electrical attraction between large numbers of cations and anions (metal and nonmetal) • Atoms completely give up electrons to other atoms

  5. Covalent bonding- chemical bonding that results from the sharing of electron pairs between two atoms (nonmetal and nonmetal)

  6. Ionic or Covalent??? • Bonding between atoms of different elements is rarely purely ionic or covalent. It usually falls somewhere between, depending on how strongly the atoms of each element attract electrons. • Electronegativity- measure of an atom’s ability to attract electrons

  7. Remember… • Electronegativity increases across a period and decreases down a group. • Which atom has the highest electronegativity? • Flourine • Which atom has the lowest? • Francium

  8. Ionic vs. Covalent

  9. Practice

  10. Nonpolar covalent bond- a covalent bond in which the bonding electrons are shared equally by the bonding atoms, resulting in a balanced distribution of electrical charge

  11. Polar bond- bonds that have an uneven distribution of charge(one end of the molecule is more electronegative than the others) Polar covalent bond- a covalent bond in which the bonded atoms have an unequal attraction for the shared electrons

  12. Examples: • Nonpolar H-H • Polar H-Cl • What about HBr, CCl4, CO2, NH3?

  13. The composition of a compound is given by its chemical formula. • Chemical formula- indicates the relative numbers of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts

  14. Covalent Bonding Molecular Compounds • Many chemical compounds including most of the chemicals that are in living things, are composed of molecules • Molecule- a neutral group of atoms that are held together by covalent bonds • Molecular compound- a chemical compound whose simplest units are molecules

  15. Molecular formula- shows the types and numbers of atoms combined in a single molecule of a molecular compound • Diatomic molecule- a molecule containing only two atoms • examples: H2 O2

  16. Covalent Bond • Bond Length- the distance between two bonded atoms at their minimum potential energy, that is, the average distance between two bonded atoms • Bond Energy- the energy required to break a chemical bond and form neutral isolated atoms (in kj/mol) • Bond length and bond energies vary with the types of atoms that have combined

  17. Types of covalent bonds • Single • Hydrogen (H2) • Double • Ethene (C2H4) • Triple • Nitrogen (N2)

  18. Octet Rule • Octet rule- chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an octet of electrons in its highest occupied energy level • They want to be like the noble gases, with their outer s and p orbitals completely filled with 8 electrons.

  19. Exceptions to the Octet Rule • Hydrogen-only needs 2 electrons to be happy • Be can have only 4 electrons • B and Al can have only 6 electrons • 3rd row or lower can sometimes have more than 8 electrons (SF6, PCl5)

  20. Lewis Structures Draw the following: • IBr • CH3Br • C2HCl • SiCl4 • F2O

  21. Ionic Bonding and Ionic Compounds • Ionic compound- composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal • The chemical formula of an ionic compound represents the simplest ratio of the compounds combined ions

  22. Formula unit- the simplest collection of atoms from which an ionic compound’s formula can be established • The ratio depends on the charges of the ions combined

  23. Characteristics of Ionic Bonding • In order for ions to minimize their potential energy, they combine in an orderly arrangement known as a crystal lattice. • Lattice energy- the energy released when one mole of an ionic crystalline compound is formed from gaseous ions • Energy is released when crystals are formed

  24. Ionic vs. Molecular Compounds

  25. Ionic Compounds • Remember, valence electrons are transferred from one atom to another • Must have an element with low electronegativity and an element with high electronegativity • 2 ions are formed-one positive and one negative (cation, anion) • They are attracted to each other by their opposite charges

  26. Electron-Dot Notation • Electron-dot notation- an electron configuration in which only the valence electrons of an atom of a particular element are shown, indicated by dots placed around the elements symbol • Examples: F H N

  27. Dot Diagrams for Ionic Compounds • Metals lose valence electrons and have positive charges • Nonmetals gain enough electrons to have 8 valence electrons and have a negative charge • Alternate positive and negative ions in the compound formula (if possible)

  28. Practice • NaCl • MgF2 • Ga2O3 • CaO

  29. Lewis Structures • Electron-dot notation can also be used to represent molecules by combining the notations of two individual atoms • H : H F F • Shared pair- a pair of electrons involved in bonding • Unshared pair (lone pair)- a pair of electrons that is not involved in bonding • You can change the shared pair to a dash • H - H F - F

  30. Lewis Structures- formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs or dashes between two atomic symbols represent electron pairs in covalent bonds, and dots adjacent to only one atomic symbol represent unshared pairs • Structural formula- indicates what kind, number, arrangement and bonds, but not the unshared pairs of the atoms in a molecule

  31. Drawing Lewis Structures • Determine the type and number of atoms in the molecule • Write the electron dot notation for each type of atom in the molecule • Determine the total number of valence electrons in the atoms to be combined • Arrange the atoms to form a skeleton structure for the molecule (hint: if carbon is present, it is always in the middle…why?)

  32. 5.Add unshared pairs of electrons so that each hydrogen atom shares a pair of electrons and each other nonmetal is surrounded by 8 electrons. 6a. Count the electrons in the structure to be sure that the number of valence electrons used equals the number available 6b. If too many electrons have been used, subtract one or more lone pairs until the total number of valence electrons is correct. Then move lone pairs until all have 8. (this will require a double or triple bond)

  33. Practice • NH3 • H2S • CH4 • H2O

  34. Multiple Covalent Bonds • Some atoms can share more than one electron pair • Double bond- a covalent bond produced by the sharing of 2 pairs of electrons between two atoms • Triple bond- a covalent bond produced by the sharing of 3 pairs of electrons between two atoms • Examples: C2H4, C2H2, N2

  35. More examples… • CH2O • CO2 • HCN

  36. Resonance Structures • Some molecules and ions cannot be represented by a single Lewis structure • Resonance- refers to bonding in molecules or ions that cannot be correctly represented by a single Lewis structure • Example: O3 (ozone), SO3-

  37. Polyatomic Ions • Polyatomic ions- a charged group of covalently bonded atoms • Polyatomic ions combine with ions of opposite charge to form ionic compounds • The charge of the ion results from an excess of electrons (- charge) or a shortage of electrons (+ charge) • Examples: NH4+, NO3-, SO4-2

  38. Molecular Geometry • The properties of molecules depend not only on the bonding of the atoms but also on molecular geometry • Molecular polarity- the uneven distribution of molecular charge

  39. VSEPR Theory • VSEPR (valence shell, electron-pair repulsion) states that repulsion between the sets of valence-level electrons surrounding an atom causes these sets to be oriented as far as possible away from each other • Examples: BeF2, BF3, CH4

  40. Linear Trigonal planar Tetrahedral

  41. Practice (use page 184 & 186) • Give the molecular geometry and bond angles of the following: • HI • CBr4 • AlBr3 • CH2Cl2

  42. VSEPR and Unshared Pairs • The lone pairs occupy space around the central atom, but the actual shape of the molecule is determined by the positions of the atoms only • Examples • NH3 • H2O • ClO3-

  43. Hybridization • To explain how the orbitals of an atom become rearranged when the atom forms covalent bonds, a different model is used • Hybridization- the mixing of 2 or more atomic orbitals of similar energies of the same atom to produce new orbitals of equal energies

  44. Hybrid orbitals- orbitals of equal energy produced by the combination of 2 or more orbitals on the same atom • Look at carbon, 2 of carbon’s valence electrons occupy the 2s and 2 occupy the 2p. To achieve four equivalent bonds, carbon’s 2s and 2p orbitals hybridize • They form 4 new identical orbitals called sp3.

  45. Metallic Bonding • Metals have unique properties • Excellent electrical properties • Shiny appearance • Malleability- the ability of a substance to be hammered into thin sheets • Ductility- the ability of a substance to be pulled into wire

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