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Chapter 6.2 and 6.5

Chapter 6.2 and 6.5. Covalent Compounds. Covalent Bonds. Sharing Electrons Covalent bonds form when atoms share one or more pairs of electrons nucleus of each atom is attracted to electron cloud of other atom neither atom removes an electron from the other

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Chapter 6.2 and 6.5

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  1. Chapter 6.2 and 6.5 Covalent Compounds

  2. Covalent Bonds • Sharing Electrons • Covalent bonds form when atoms share one or more pairs of electrons • nucleus of each atom is attracted to electron cloud of other atom • neither atom removes an electron from the other • Usually formed by a nonmetal + a nonmetal

  3. Covalent Bonding

  4. Covalent Bonds • Sharing Electrons • Covalent bonds • Molecules are formed • space where electrons move is called molecular orbital • made when atomic orbitals overlap

  5. Molecules

  6. Covalent Bonds • Energy and Stability • Noble gases are stable (full octet) (low potential energy) • Other elements are not stable (high potential energy) • covalent bonding decreases potential energy because each atom achieves an electron configuration like a noble gas

  7. Covalent Bonds • Energy and Stability • because potential decreases when atoms bond, energy is released • i.e., atoms lose potential energy when they bond • loss of potential energy implies higher stability

  8. Covalent Bonds • Energy and Stability • Potential energy determines bond length • At minimum potential energy, the distance between two bonded atoms is called bond length • The bonded atoms vibrate about the bond, therefore, bond length is an average length

  9. Covalent Bonds • Energy and Stability • bonds vary in strength • bond energy: the amount of energy required to break the bonds in 1 mol of a chemical compound • bond energy predicts reactivity • bond energy is equal to loss of potential energy during formation

  10. Bond Energies and Lengths

  11. Bonds and Energy • Single bonds are the longest bonds with the least bond energy • Double bonds are shorter, stronger and have intermediate bond energy • Triple bonds are the shortest, strongest, and have the highest bond energy

  12. Covalent Bonds • Electronegativity • Atoms share electrons equally or unequally • nonpolar covalent bond: bonding electrons shared equally • polar covalent bond: shared electrons more likely to be found around more electronegative atom (shared unequally)

  13. Covalent Bonds • Electronegativity • The difference in electronegativity can be used to predict the type of bond (but boundaries are arbitrary) • Electronegativities are listed on your periodic table

  14. Bond Types 0.3 1.7

  15. Covalent Bonds • Polar molecules have partially positive and partially negative ends • such molecules are called “dipoles” • “δ” means “partial” in math and science • The positive end is designated as δ+ and the negative end as δ- • example: Hδ+Fδ-

  16. N and H F and F Ca and Cl C and O Polar covalent Non-polar covalent Ionic Polar covalent Practice: Calculate the bond type

  17. Covalent Bonds • Polarity is related to bond strength • greater electronegativity difference means greater polarity, which means greater bond strength

  18. Electronegativity Difference for Hydrogen Halides

  19. Covalent Bonds • Bond type determines properties of substances • metallic bonds: electrons can move from one atom to another—good conductors • ionic bonds: hard and difficult to break apart • covalent bonds: low melting/boiling points

  20. Properties of Substances with Different Types of Bonds

  21. Lewis Structures • G. N. Lewis was an American chemistry who developed the system of showing valence electrons and the resulting bonds • Lewis structure: a structure in which atomic symbols represent nuclei and inner-shell electrons, and dots are used to represent valence electrons

  22. Drawing Lewis Structures for Atoms • Lewis structures represent valence electrons with dots • position of electrons is symbolic (not literal) • the symbol represents the nucleus and core electrons • dots around the atomic symbol represent electrons

  23. Lewis Structures of Second-Period Elements

  24. Vocabulary for Lewis Structures • Unshared pairs: AKA lone pairs, a pair of valence electrons not involved in bonding to another atom • Single bond: a covalent bond in which one pair of electrons is shared between two atoms • Multiple bond: double bond shares 2 • pairs of electrons, triple bond shares 3 pairs of electrons

  25. Drawing Lewis Structures for Molecules • Lewis Electron-Dot Structures • draw the Lewis structure for each atom in the compound; place one electron on each side before pairing • Count and write down the total number of valence electrons for all atoms involved • Count and write down all the electrons in lone pairs

  26. Drawing Lewis Structures for Molecules • Subtract the number of electrons in lone pairs from the total number of valence electrons • Divide this number by 2 and this is usually the number of bonds you have • Draw sticks for bonds (2 electrons each), draw in the lone pairs, check for octets

  27. Drawing Lewis Structures for Molecules • The atom with the lowest electronegativity is usually the central atom • If carbon is present, it is the central atom

  28. Lewis Structure Practice • Draw Lewis structures for the following molecules: • SiF4 HBr • PBr3 CO2 • SiS2 H2O • CO ammonia (NH3)

  29. Drawing Ions • You must consider the absence of an electron or the extra electrons • Subtract an absent electron from the total • Add any extra electrons to the total (add on an atom that is not the central atom)

  30. Drawing Ions • Place the entire structure in brackets and put the charge in the upper right corner outside the brackets • Coordinate-covalent bond: resulting bond when a shared pair of electrons originates from one atom

  31. Drawing Ions • Draw the Lewis structure for the sulfate ion, SO42-

  32. Resonance Structures • Resonance Structures: bonding in molecules or ions that cannot be correctly represented by a single Lewis structure • Draw each possible structure and the hybrid structure

  33. Practice Resonance Structure • Draw the Lewis structure for SO3 • Possible structures:

  34. Naming Covalent Compounds • First name: name of first element in formula • usually least electronegative • requires a prefix if more than one of them • Second name: ends in –ide • requires a prefix if more than one of them • oxygen requires the prefix “mono” if only one of them

  35. Naming Covalent Compounds

  36. Practice Naming

  37. More Naming Practice

  38. Molecular Shapes • Three-dimensional shape helps determine physical and chemical properties • valence shell electron pair repulsion (VSEPR) theory predicts molecular shapes • based on the idea that electrons repel one another

  39. Molecular Shapes • Lewis structures show which atoms are connected where, and by how many bonds, but they don’t properly show the 3-D shapes of molecules • To find the actual shape, first draw the Lewis structure, and then use VSEPR theory

  40. MOLECULAR GEOMETRY VSEPR • Valence Shell Electron Pair Repulsion theory. • Most important factor in determining geometry is the relative repulsion between electron pairs.

  41. MOLECULAR GEOMETRY Molecule adopts the shape that minimizes the electron pair repulsions.

  42. VSEPR Rules • To apply VSEPR theory: • Draw the Lewis structure of the molecule and identify the central atom • Count the number of electron charge clouds (lone and bonding pairs) surrounding the central atom. • Predict molecular shape by assuming that clouds orient so they are as far away from one another as possible

  43. Bond Angles • Lone-pairs of electrons behave as if they are slightly bigger than bonded electron pairs and act to distort the geometry about the atomic center so that bond angles are slightly smaller than expected

  44. Polarity of Molecules • Two atoms: bond polarity is the molecular polarity • More than 2 atoms: the geometry of the molecule must be considered • If the bonds are non-polar, the molecular is non-polar • Some molecules with polar bonds • can be non-polar

  45. More • Sometimes the partial charges cancel each other out because they are directly opposite each other • Consider CO2 and CCl4 • The symmetrical distribution of the bonds leads to cancellation of the charges

  46. Polar and Non-polar Molecules

  47. Linear with Only 2 Atoms • Molecules made of only two atoms are always linear • Bond angle is 180° • Examples: HCl, F2 • Bond polarity is the molecular polarity

  48. Linear with More than 2 Atoms • Sometimes molecules made of more than 2 atoms are shaped linearly • Example: CO2 • The oxygens pull the electrons in equal, but opposite, directions • The molecule is sometimes nonpolar and sometimes polar

  49. Bent • There are two shapes that are called bent • The first one has one lone pair on the central atom and two bonded atoms • Example: NOCl (nitrosyl chloride) • Bond angle: 117° • The molecule is polar

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