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Bonding Honors Chemistry Unit 6.

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Presentation on theme: "Bonding Honors Chemistry Unit 6."— Presentation transcript:

1 Bonding Honors Chemistry Unit 6

2 Bond Types Ionic: transfer of electrons
Covalent: sharing electron pair(s) Metallic: delocalized electrons

3 Covalent Bonds Characteristics
Low melting points Don’t conduct electricity Most are brittle if solid, but usually gas or liquid Particle of a covalent compound is called a molecule (most are between nonmetals)

4 Covalent Bonds Two types:
Polar covalent – one atom attracts shared pair of electrons more strongly (most) sides of bond appear to be partially charged Nonpolar covalent – electrons are being shared equally, no charge difference (no electronegativity difference) usually between two atoms of same element

5 Terms to Know Bond axis: line joining nuclei
Bond angle: angle between 2 axes Bond length: distance between nuclei Bond energy: energy to break bond Bonds are not fixed More like a stiff spring Average position is given as bond length or bond angle

6 Molecules Covalently bonded compounds
Diatomic molecules: always as 2 atoms when in element form (like O2) 7 elements, make a 7 in the periodic table (begin with N) and most are in group 17 Elements: Br, I, N, Cl, H, O, F

7 Naming a compound with two non-metals
Use the prefixes: (1) mono-, (2) di-, (3) tri-, (4) tetra-, (5) penta, (6) hexa-, (7) hepta-, (8) octa-, (9) nona-, (10) deca If the first element listed has a quantity of just one then you don’t use mono- as a prefix. Put the appropriate prefix in front of the name of each element change the ending to –ide. Example: N2O5 Dinitrogen pentaoxide

8 Lewis Dot Structures Find the total number of valence electrons using group numbers for each element Arrange atoms to form skeleton structure with lines connecting the atoms. If carbon is present, it is central. Otherwise, the least electronegative element is central. H is NEVER central.

9 Lewis Structures (continued)
Each line counts as 2 electrons. Subtract these from total valence electrons. Compare the electrons left to what each needs to be full. If they are the same, add unshared pairs to give each nonmetal or metalloid a full octet (except H). Add all electrons to see if they equal the valence electrons. (Gr2 and 13 just double their electrons don’t get a full octet, Gr. 2 gets 4, Gr. 13 gets 6) If there are not enough electrons to give each its own dots, one more line needs to be drawn for each 2 electrons you are short (2 atoms share). Recalculate from the valence electrons and dots can be given.

10 Example 1 H2O 2(1) + 6 = 8 valence electrons Skeleton: .. H-O-H
. . Subtract 2 for each line = 4 e- left Put dots to complete octet for oxygen

11 Example 2 CH3I 4+3(1)+7 = 14 valence electrons Skeleton: H H – C – H
| H – C – H : I : . . Subtract 2 for each line = 6 e- left Add dots to I to complete the octet. Other Examples: NH3, AlI3, SeO2, CO2, SO3 :

12 Resonance Using more than one Lewis structure to explain when bonds are in between drawn structures (from lab measurements)

13 Molecular Shape Based on VSEPR theory: valence-shell electron-pair repulsion theory Electrons want to be as far apart as possible (like charges repel) Pairs around central atom will give angles 2 pairs: linear 180o angle 3 pairs: trigonal planar 120o angle 4 pairs: tetrahedral 109.5o angle Repulsion is greater for unshared pairs: they push harder on shared pairs, decreasing the expected bond angle 2 unshared>1 shared with one unshared>2 shared

14 VSEPR Oklahoma State Link Possible Shapes: p. 186 of book
Linear: 2 shared (bonded pairs) 180o angle Trigonal planar: 3 shared pairs 120o angle Tetrahedral: 4 shared pairs 109.5o angle Trigonal pyramidal: 3 shared, 1 unshared (lone) <109.5o angle (107o angle) Bent: 2 shared, 2 unshared (lone) <109.5o angle (104.5o angle)

15 Determining Shape Draw Lewis structure
Count shared and lone pairs on central atom (ONLY!) to determine shape Example: H2O Lewis structure: .. H-O-H . . 2 shared (lines), 2 unshared (dot pairs) Shape: bent, Angles: <109.5o angle (104.5o angle) Other Examples: NH3, AlI3, CH4, HF, SO3

16 Hybridization Hybridized orbitals merge s and p orbitals by borrowing empty p orbitals to put one electron in each. This allows them to share that orbital with an electron from another atom in a covalent bond. The new hybrids have an energy that is in between that of s and p Examples: Be, Al & B, C & Si (& others)

17 Hybrid Orbitals Count bonds to see how many orbitals are sp hybrids
needed. Start with s, then add p orbitals to make enough. This names the hybrids. sp2 hybrids sp3 hybrids

18 Predicting Bonds Based on electronegativity difference
Examples of calculation: (use table on p. 151) H-F – 2.1 = 1.9 H-Br – 2.1 = .7 H-I – 2.1 = .4 The greater the difference, the stronger the bond

19 Bond Character Large difference: ionic bond
Small difference: covalent bond Dividing line is 1.7 > 1.7 is ionic, < 1.7 is covalent, = 1.7 is 50% ionic and 50% covalent Unless bonded to the same type atom, the bond has both ionic and covalent character (use chart on back of per. table)

20 Find the electronegativity difference in black in the chart
The percent ionic character is underneath in red Subtract from 100 to find the covalent character Example: H-F difference was 1.9 Ionic character is listed as 59%, so covalent character is = 41% covalent H-Br, H-I

21 Bonding Demo Record color and intensity (brightness) as each bond forms in your journal, then calculate the % character for each bond S-O 3.5 – 2.5 = 1.0 difference 22% ionic, 78% covalent Mg-O 3.5 – 1.2 = 2.3 difference 74% ionic,26% covalent

22 Polarity If charge of polar bonds is distributed
equally in all directions, the molecule is nonpolar If charge of polar bonds is not equal in all directions, the molecule is polar Look for something that makes the charge asymmetrical (either of these makes it polar) Bonded atoms are not all the same element attached to the central atom Unshared pairs of electrons on the central atom A polar molecule is called a dipole (has + and – poles) Polarity is measured as dipole moment

23 van der Waals Forces Intermolecular: Weak forces between molecules (van der Waals forces) Intramolecular: strong forces inside a molecule holding atoms together (bonds) Types of van der Waals forces Dipole-dipole: between polar molecules Dipole-induced dipole: between dipole and nonpolar (peer pressure model) London Dispersion Forces: temporary dipoles that happen because of electron movement Induced by concentrations of electrons in nonpolar molecules Only attractive force operating in nonpolar substances 85% of force in most polar molecules (exceptions: NH3, H2O)

24 Induced Dipole Peer pressure model
Electrons of nonpolar molecule are disturbed by presence of charged particle (ion or dipole)

25 Dipole-Induced Dipole

26 Temporary Dipole Movement of electrons may cause electron distribution to become asymmetrical for an instant

27 Effects of IM Forces Properties are affected by IM forces
Boiling and melting points give an indication of how strong the IM forces are Nonpolar substances have the weakest IM forces: gases or lowboiling liquids (lower melting and boiling points) Polar substances have dipole forces that are stronger: liquid or solid at room temp (higher melting and boiling points)

28 Soaps and Detergents There are polar and nonpolar sides to a soap molecule The nonpolar side embeds or dissolves in greasy dirt The polar side is attracted to water molecules (polar) Agitation breaks globule up into small pieces which are then pulled away into the water and washed away. Detergents have an additive to keep soap scum from forming.

29 Chromatography Fractionation (separation) based on polarity
Two phases: Mobile phase: mixture to be separated dissolved in liquid or gas Stationary phase: solid or liquid adhering to a solid Types: column, paper, gas

30 Column Chromatography
Stationary phase is in a column. Used for delicate separations such as vitamins, hormones, and proteins. HPLC and ion are special kinds of column chromatography

31 Paper Chromatography Separation on paper into spots or lines on the strip Has limitations

32 Gas Chromatography Used to analyze volatile liquids and gas or vapor mixtures. Mixed with inert gas (like He) in mobile phase Interpreted by computer

33 Gas Chromatogram

34 Chromatography Applications
Drug testing uses column and gas chromatography Car emissions are done with gas chromatography

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