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Molecular Nomenclature and Geometry

Molecular Nomenclature and Geometry. Chemistry Text Ch 6.1,6.2,6.5,16.1-16.3. Covalent Bonding. Covalent bonding entails a sharing of electrons. Covalent bonding usually occurs between nonmetals. Form individual molecules. Shapes of molecules determines properties.

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Molecular Nomenclature and Geometry

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  1. Molecular Nomenclature and Geometry Chemistry Text Ch 6.1,6.2,6.5,16.1-16.3

  2. Covalent Bonding • Covalent bonding entails a sharing of electrons. • Covalent bonding usually occurs between nonmetals. • Form individual molecules. • Shapes of molecules determines properties.

  3. Properties of Covalent Molecules • Gases, liquids, or solids (made of molecules) • Low melting and boiling points • Poor electrical conductors in all phases • Many soluble in nonpolar liquids but not in water

  4. Covalent Bonds – attain the octet or full valence by sharing pairs of valence electrons.

  5. Covalent Bonds • For every pair of electrons shared between two atoms, a single covalent bond is formed.  • Some atoms can share multiple pairs of electrons, forming multiple covalent bonds. 

  6. Building a Dot Structure Hydrogen - H2 1. How many valence electrons? (1) 2. Add up the number of valence electrons that can be used. H = 1 and H = 1 Total = (1 + 1) = 2 = 2 electrons 3. Put dots together With dots- H:H Electrons

  7. Building a Dot Structure Hydrogen - H2 With dots- H:H 4. When two dots are together, replace with a line. H:H H-H Electrons

  8. Building a Dot Structure Ammonia, NH3 1. Number of Valence electrons H= 1 N=5 Decide on the central atom; never H. 2. Add up the number of valence electrons that can be used. H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons / 4 pairs

  9. H H N H H H N H Building a Dot Structure 3. Form a single bond between the central atom and each surrounding atom (each bond takes 2 electrons!) .. 4. Remaining electrons form LONE PAIRS to complete the octet as needed (or duet in the case of H). 3 BOND PAIRS and 1 LONE PAIR.

  10. H H N H Building a Dot Structure .. 5. Check to make sure there are 8 electrons around each atom except H. H should only have 2 electrons. This includes SHARED pairs. 6. Also, check the number of electrons in your drawing with the number of electrons from step 2. If you have more electrons in the drawing than in step 2, you must make double or triple bonds. If you have less electrons in the drawing than in step 2, you made a mistake!

  11. Simplest Organic molecule The H’s can be replaced with any group 17 Halogen(s) to make a multitude of molecular compounds

  12. Diatomic Elements - There are seven elements that exist as diatomic molecules in which two atoms of the same element bond together. • They arebromine, iodine, nitrogen, chlorine, hydrogen, oxygen, andfluorine. If the symbols are written for these elements in the order given, they spell outBrINCl H OF. • Whenever these elements appear as free elements (by themselves) in a chemical equation, they MUST have a subscript "2" • Ex. Br2 I2 N2 Cl2 H2 O2 F2

  13. Double Covalent Bond • Oxygen (which has six valence electrons) needs two electrons to complete its valence shell. Two oxygen atoms form the compound O2, they share two pairs of electrons, forming two covalent bonds (double bond).  

  14. Triple Covalent Bonds • Nitrogen (which has five valence electrons) needs three electrons to complete its valence shell. Two nitrogen atoms form the compound N2, they share three pairs of electrons, forming three covalent bonds (triple bond). 

  15. Covalent Bonds • – Draw the Lewis structures for each atom, draw circles to show the electrons that are shared, and then write the bond structure and chemical formula. • Fluorine + Fluorine • (B) 3 Hydrogen + 1 Phosphorus • (C) 2 Hydrogen + 1 Sulfur • (D) 1 Carbon + 2 Oxygen

  16. lowest energy state- allows four bonds by electron promotion Excited state 4 valence e-/sp3 Hybridization

  17. VSEPR • VSEPR stands for Valence Shell Electron Pair Repulsion.   • Basically, the idea is that covalent bonds and lone pair electrons like to stay as far apart from each other as possible under all conditions.  • This is because covalent bonds consist of electrons, and electrons don't like to hang around next to each other much because they have the same charge (like charges repel).

  18. VSEPR explains why molecules have their shapes.  • If carbon has four atoms stuck to it (as in CH4), these four atoms want to get as far away from each other as they can.  This isn't because the atoms necessarily hate each other, it's because the electrons in the bonds “hate” each other.  That's the idea behind VSEPR.

  19. Methane is Tetrahedral Sp3 hybridized carbon 4 equivalent C-H bonds (s-bonds) All purely single bonds are called s-bonds

  20. Molecular Geometry

  21. Polarity • Depending on the percent covalent vs. ionic characteristic of the bond, molecular compounds can have polar covalent or nonpolar covalent bonds • The higher the percentage of ionic characteristic the more polar the bond will be

  22. Naming Covalent Compounds Covalent compounds are named by adding prefixes to the element names. A prefix is added to the name of the first element in the formula if more than one atom of it is present. (The less electronegative element is typically written first.) A prefix is always added to the name of the second element in the formula. The second element will use the form of its name ending in ‘ide’.

  23. Naming Covalent Compounds Prefixes Note: When a prefix ending in ‘o’ or ‘a’ is added to ‘oxide’, the final vowel in the prefix is dropped.

  24. Naming Binary Covalent Compounds:Examples N2S4 dinitrogen tetrasulfide NI3 nitrogen triodide XeF6 xenon hexafluoride CCl4 carbon tetrachloride P2O5 diphosphorus pentoxide SO3 sulfur trioxide

  25. Writing Formulas for Covalent Compounds The names of covalent compounds contain prefixes that indicate the number of atoms of each element present. If no prefix is present on the name of the first element, there is only one atom of that element in the formula (its subscript will be 1). A prefix will always be present on the name of the second element. The second element will use the form of its name ending in “ide”

  26. Writing Formulas for Binary Covalent Compounds:Examples nitrogen dioxide NO2 diphosphorus pentoxide P2O5 xenon tetrafluoride XeF4 sulfur hexafluoride SF6

  27. Molecular Solid • Within molecules, covalent bonding holds molecule together • Between molecules, intermolecular forces hold molecules together to make them liquids or solids. • Molecules are held in place by intermolecular forces • Properties • Low to moderate melting point and boiling point • Soft • Non conductive

  28. Intermolecular Forces Intermolecular Forces are electrostatic forces of attraction that exist between an area of negative charge on one molecule and an area of positive charge on a second molecule.

  29. TYPES OF INTERMOLECULAR FORCES • (only attractions not bonds) • Hydrogen Bonding ( generally the STRONGEST) • Van Der Waals • a. Dipole-dipole (polar molecules) • b. London dispersion forces (nonpolar • molecules

  30. Two of the intermolecular forces are associated with POLAR structures. • Hydrogen Bonding • Dipole-dipole Forces • One of the intermolecular forces is associated with NONPOLAR structures. • London Dispersion Forces

  31. Hydrogen Bonding : • These occur between polar covalent molecules that possess a hydrogen bonded to an extremely electronegative element, specifically - N, O, and F.

  32. Hydrogen Bonding : d - d + • Hydrogen’s single electron is pulled toward the very electronegative oxygen resulting in: • Large partial charges • The unshielded nucleus of hydrogen attraction to the unshared electron pairs O H d + d - H d + O Hydrogen bond H Covalent bond d + H

  33. Hydrogen Bonding :

  34. Properities due to Hydrogen Bonding • Higher than expected melting/boiling points • More viscous substances (liquids are “thicker” to pour) • Surface tension – an inward pull that minimizes the surface area of a liquid. • Capillary Action

  35. Hydrogen Bonding :

  36. Hydrogen Bonding : Surface Tension

  37. Only polar covalent molecules have the ability to form dipole-dipole attractions between molecules. Polar covalent molecules act as little magnets, they have positive ends and negative ends which attract each other. • Dipole-Dipole :

  38. Dipole-Dipole :

  39. London Dispersion Force : • Because of the constant motion of the electrons, an atom or molecule can develop a temporary (instantaneous) dipole when its electrons are distributed asymmetrically about the nucleus. • The attractive forces are responsible for Bromine being a liquid and Iodine a solid at room temperature.

  40. London Dispersion Force : • The electron asymmetry about the nucleus induces a temporary attraction between the non-polar molecules causing the London Dispersion Force.

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