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Chemical Bonds

Chemical Bonds. Forces that hold atoms together Ionic bonds : the forces of attraction between ions; involves a transfer of electrons. Covalent bonds : the forces of attraction between two atoms that are sharing electrons. Ionic Bonds.

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Chemical Bonds

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  1. Chemical Bonds • Forces that hold atoms together • Ionic bonds: the forces of attraction between ions; involves a transfer of electrons. • Covalent bonds: the forces of attraction between two atoms that are sharing electrons

  2. Ionic Bonds • Result from reaction between metal and nonmetal to form a cation and an anion • An ionic bond is the attraction between a positive ion and negative ion. • The ions are arranged in a pattern called a crystal lattice.

  3. Crystal Structures sodium chloride

  4. Covalent Bonds • Sharing pairs of electrons • Molecules attracted to each other weakly • Often found between nonmetal atoms N2 Cl2 HF

  5. • •  H F  Bond Polarity • Covalent bonding between unlike atoms results in unequal sharing of the electrons. • The result is bond polarity; HF

  6. Electron Sharing in HF Actual molecule If electrons were equally shared

  7. Electronegativity Measure of the inherent ability of an atom to attract shared electrons A larger electronegativity means that an atom attracts shared electrons more strongly; fluorine in HF The difference in electronegativity between two atom in a bond is a measure of bond polarity. A larger difference in electro- negativity means a more polar bond. If the difference is ≥2, the bond is ionic If the difference is >0 and <2, the bond is polar covalent If the difference is 0, the bond is covalent

  8. Electronegativity (cont.)

  9. Choose the bond in each pair that will be more polar. • H-P or H-C: • O-F or O-I: • N-O or S-O: • N-H or Si-H:

  10. Classify the following bonds as ionic, covalent, or polar covalent • HCl • NaF • Cl2 • Cl-F • NH3

  11. Dipole Moment • Any molecule that has a center of positive charge and a center of negative charge at different points in space is said to have a dipole moment. H-F

  12. Dipole Moment (cont.) • In polyatomic molecules the dipole moment depends upon the atoms involved and the three-dimensional structure.

  13. Dipole Moment (cont.) The dipole moment affects the attractive forces between molecules, and therefore the physical properties of the substance. For example, water readily dissolves NaCl.

  14. Electron Arrangements and Ionic Bonding • Metals lose their valence electrons to form cations. • Nonmetals gain electrons to form anions. • Both try to acquire the electron configuration of a noble gas; for example, NaCl Na: [Ne]3s1→ [Ne] + 1e- Cl: [Ne]3s23p5 + 1e- → [Ar]

  15. What are the electron configurations of the ions in the following compounds? • Al2S3 • MgO • SrF2 • LiCl

  16. Electron Arrangements and Ionic Bonding (cont.) • Formulas are predicted by achieving electrical neutrality using the component cations and anions. What is the formula of the ionic compound formed from Al and S? • In polyatomic ions, the atoms in the ion are connected with covalent bonds. The ions are attracted to oppositely charged ions to form an ionic compound.

  17. Properties of Ionic Compounds • All are solids at room temperature. • Melting points are greater than 300°C. • Many are soluble in water. When dissolved the solution becomes an electrical conductor.

  18. Bonding & Structure of Ionic Compounds • Crystal lattice: geometric pattern determined by the size and charge of the ions. • Anions are almost always larger than cations. • Anions are generally considered “hard” spheres packed as efficiently as possible, with the cations occupying the “holes” in the packing. • Compounds with polyatomic anions contain covalent bonds within the anion structure, which are ionically bonded to the cation; CuSO4

  19. Copper(II) sulfate SO42- Cu2+

  20. Lithium fluoride

  21. Lewis Structures • Use the symbol of the element to represent the nucleus and inner electrons. • Use dots around the symbol to represent valence electrons. • Elements try to achieve a noble gas configuration. sodium chloride

  22. Lewis Structures (molecules with covalent bonds) • Hydrogen shares two electrons (duet rule). • Helium already has two, so it does not form bonds. • The second row nonmetals require eight electrons to • fill the 2s and 2p orbitals (octet rule). • example: Cl2 • Neon has eight electrons, so it does not form bonds.

  23. Writing Lewis Structures of Molecules • Count the total number of valence electrons from all the atoms. • Attach the atoms together with one pair of electrons. • Arrange the remaining electrons in pairs so that all hydrogen atoms have two electrons (one bond) and other atoms have eight electrons (combination of bonding and nonbonding). Nonbonding pairs of electrons are also know as lone pairs.

  24. Draw Lewis structures for the following molecules: • NH3 • CCl4 • LiBr • CH3OH

  25. Multiple Covalent Bonds • Single covalent bond: atoms share two electrons (one pair); HBr • Double covalent bond: atoms share four electrons (two pairs); C2H4 • Triple covalent bond: atoms share six electrons (three pairs); HCN

  26. Lewis Structures of Molecules with Multiple Bonds • Determine the total number of valence electrons. • Form a single bond around each atom. • Distribute the remaining electrons. • Make double or triple bonds as needed until each atom (except H) acquires an octet of electrons.

  27. Draw Lewis structures for the following molecules: • CO2 • NO3-1 • HCO3-1 • SO3

  28. Resonance • When there are multiple Lewis structures for a molecule that differ only in the positions of the electrons, they are called resonance structures.(Lone pairs and multiple bonds are in different positions.) The real structure is a hybrid. NO3-1

  29. Draw resonance structures for HCO3-1

  30. Exceptions to the Octet Rule Some atoms may violate the octet rule. beryllium: forms two bonds; BeCl2 boron and aluminum: normally form three bonds; BH3, AlCl3 phosphorus: can form up to five bonds; PCl5 sulfur: can form up to six bonds; SF6

  31. Molecular Structure • Refers to the three-dimensional geometry of a molecule. • Can specify a bond angle. • We will consider five different structures: - linear - trigonal planar - tetrahedral - trigonal pyramid - bent 120°

  32. Molecular Structure (cont.) 180° • Linear: BeCl2 • Trigonal planar: BH3 • Tetrahedral: CCl4 120° 109.5°

  33. Trigonal pyramid: NH3 • Bent: H2O

  34. Predicting Molecular Structure • VSEPR Theory • Valence Shell Electron Pair Repulsion • The shape around the central atom can be predicted by assuming that the electrons surrounding the central atom will position themselves as far apart as possible. BeCl2

  35. Steps for Predicting Molecular Structure • Draw the correct Lewis structure. • Count the electron pairs and place them as far apart as possible. • Determine the positions of the atoms. • Determine the name of the molecular structure based on the positions of the attached atoms.

  36. Determine the molecular structures of the following molecules: • NH3 • CCl4 • AlCl3 • BeF2

  37. Some molecules contain a combination of more than one shape. CH3OH

  38. Molecules with Multiple Bonds • When using the VSEPR model to predict the molecular geometry, multiple bonds are counted the same as a single electron pair. CO2 HCN

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