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Covalent Compounds and Intermolecular Forces

Covalent Compounds and Intermolecular Forces. Bonds. Chemical bonds are __________ forces They act between atoms ________ a molecule. Why does bonding occur?. Bonding occurs to maximize stability of the atoms involved. More stable = LOWER potential energy. Bond types.

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Covalent Compounds and Intermolecular Forces

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  1. Covalent Compounds and Intermolecular Forces

  2. Bonds • Chemical bonds are __________ forces • They act between atoms ________ a molecule

  3. Why does bonding occur? • Bonding occurs to maximize stability of the atoms involved. • More stable = LOWER potential energy

  4. Bond types • Dependent on the difference between the electronegativities of the elements involved in bond • Electronegativity • ______________________________ • Highest found in small non-metals • ______________ are not ranked

  5. To Determine Type Subtract the electronegativities If difference is Zero – bond is ________________ 0.1-1.67- bond is ________________ Greater than 1.67- __________ Bond Types (cont)

  6. Bonding is all about the electrons! • Bond type tells us what will happen to the electron(s) • Octet rule will give us an idea of how many electrons will be involved • Have to look at valance electrons • Remember most atoms are stable with 8 • Common exceptions • H, He, Li, and Be can be stable with _____ • B is stable with ______ • Elements with d orbitals (can have more than 8)

  7. 1 2 3 4 5 6 7 8 Valence electrons

  8. Covalent bonding • Covalent bonding involves the ___________ of valence electrons. • Atoms can share electrons in order to have full valence shell around them. (usually 8 electrons)

  9. Two Types of Covalent Bonds • Non-polar • _________ sharing of electrons • Occurs between the same element bonded together • For example: ______ • Polar • __________ sharing of electrons • Occurs between two different elements bonded together • For example: _______

  10. Polarity of Bonds • Polarity indicates how much electrons are pulled to one end of the molecule or the other. • Creates two ends (or poles) • One end is slightly negative and the other is slightly positive • Creates a ______________

  11. This atom has greater electronegativity. δ+ δ-

  12. Labeling a Polar Bond • Need to draw a vector to show polarity • Vectors • A vector is an arrow that represents the strength of something (like a charge), and the direction in which it is acting.

  13. Vectors for Dipoles • The vector points in the direction of the partial ______________ • It has a + on the partial positive end.

  14. Vector of dipole δ+ δ-

  15. Polarity • Ionic compounds are very polar. • They have one atom that’s so strong it can pull one electron (or more) away from the other atom. So the electrons are pulled to one end of the molecule. • This is a matter of degrees and there really isn’t a distinct line between polar covalent and ionic.

  16. Covalent Compounds • _________ Covalent • Form molecules • Example: ________ • ________ Covalent • Form large interconnected networks • Example: _________

  17. Naming Binary Covalent Compounds • Compound made from two nonmetals (covalent bond) • For example: H2O, CO, NH3, CH4 • No _________ Involved!! • Form ___________

  18. Creating Binary Covalent Compound Formulas • Can only be created from the name or a description of how many atoms of each element there are

  19. Creating Binary Covalent Compound Formulas (cont) • For example, • What is the formula for a compound made of 2 boron atoms and three oxygen atoms? • _____________

  20. Naming Binary Covalent Compounds • Base name with _____________ • First element has element name • Second element’s ending is changed to –ide • Prefix is put before each element to designate how many atoms there are • Mono- is never put in front of the first element

  21. Binary Covalent Compound Prefixes • Mono 1 • Di 2 • Tri 3 • Tetra 4 • Penta 5 • Hexa 6 • Hepta 7 • Octa 8 • Nona 9 • Deca 10

  22. Naming Binary Covalent Compounds (cont) • For example: • CO2 is _______________ • BF3 is _______________ • B2O3 is _______________ • P2O5 is _________________

  23. Representing bonds • There are several ways we represent bonding so that we can have a visual picture of how the molecules look. • The most common is called • Lewis structure • Lewis dot diagram • Electron dot diagram • (they all mean the same thing.)

  24. Drawing Lewis Dots • Write the symbol. • Find out how many valence electrons. • Draw that many dots • 2 Dots max / side • Max of 8 dots!

  25. Examples of Lewis Dots • Carbon – ______________ C • Nitrogen – ____________ N

  26. Lewis Structure for Compounds • Ionic Compounds • Show the Lewis Structure of each ion (including charge) • Sit them next to each other

  27. Lewis Structure for Compounds • Covalent Compounds • The central atom is usually the one with the lowest electronegativity (but never ______) • Determine total valence electrons • Move electrons so that each terminal atom has an octet (but not H who gets 2) • Any extra electrons go on the _____________

  28. Covalent bonding • Let’s make carbon tetrachloride, CCl4 • Start with the atoms C Cl Cl Cl Cl

  29. Covalent bonding

  30. Double and Triple Bonds • Atoms can form single, double or triple bonds with other atoms. • One example – carbon dioxide

  31. Structural formula • Electron dot diagrams show all valence electrons as “dots”. • We frequently represent BONDED electrons (shared electrons) as lines. One line is a single bond, two lines is a double bond, three lines is a triple bond. • This is called the structural formula. • _______________ are still dots.

  32. Examples Cl Cl Cl - C - Cl Cl C Cl Cl Cl

  33. Limitations of Lewis diagrams • Work great on paper, but they are drawn in two dimensions. • Molecules exist in ________ dimensions.

  34. Electrons repel each other • In order to understand how molecules behave in three dimensions, we need to realize that electrons (or groups of electrons) repel each other. • When do we have groups of electrons on a Lewis diagram? • _______________ • ________________

  35. VSEPR • VSEPR (Valence Shell Electron Pair Repulsion) Theory states that pairs of electrons repel each other. • This allows us to predict the shapes of molecules in three dimensions.

  36. Rules for VSEPR • Used for covalently bonded molecules only (all non-metals) • “Electron groups” can be bonds or lone pairs. • Double and triple bonds behave like single bonds (so they’re really 2 or 3 pairs of electrons, but they act like 1 group and we count them as 1 group).

  37. Steps for using VSEPR to predict molecular shape (geometry) • Draw the Lewis structure. • Count the number of electron groups (on the central atom). • Look at the chart to determine electron group geometry (first row). • Count the number of lone pairs on the central atom (if none, you’re done) and move down the chart to name the molecular geometry.

  38. VSEPR Chart

  39. Linear

  40. Trigonal Planer

  41. Tetrahedral

  42. Let’s try some examples: • Draw the Lewis structure for methane, CH4. • How many electron groups are on the central atom? (electron groups = bonds + lone pairs) • So the electron group geometry is TETRAHEDRAL. • Since there are no lone pairs on the central atom, the molecular geometry is the same as the electron group geometry: TETRAHEDRAL. • Now make a model of methane. • In the Lewis structure, the terminal atoms are 90 degrees apart. • In the 3-D model, they’re more like 109.5°.

  43. Another example: • Draw the Lewis structure for water. • How many electron groups on the central atom? • Electron group geometry is _______. • How many lone pairs on the central atom? • Find the 3D model that matches the electron group geometry and pull of an atom for each lone pair. • Molecular geometry is ____________.

  44. One more example: • Draw the Lewis structure of COH2 • How many electron groups are on the central atom? • Electron group geometry is _______. • How many lone pairs are on the central atom? • Molecular geometry is ________. • Choose the best model to represent this. • How does it compare to MY model of this? • Does the double bonded atom repel the other terminal atoms differently than a singly bonded atom does?

  45. Polar Bond vs Polar Molecule • Polar Bond • Between two different non-metals • Electronegativity difference creates a pull in the electrons to one side of bond • Slightly negative and positive end of bond • Polar Molecule • Must contain at least one polar bond • Polar bonds can cancel each other out • Creates slightly negative and slightly positive end of a molecule

  46. Deciding if a Molecule is Polar • Need to draw a vector to show polarity • Vectors of equal strength and in opposite directions cancel out!

  47. Molecules with 2 Atoms • Any 2-atom molecule with a polar bond has a dipole moment. • The molecule will be polar

  48. Molecules with 3+ atoms. • Vectors are very important for helping us determine the polarity of a molecule that has more than one bond. • In some molecules, the polar bonds will cancel each other out making the molecule non-polar (even though it contains polar bonds) • This is bases on the molecules geometry

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