Bonding Honors Chemistry Unit 6
Bond Types Ionic: transfer of electrons Covalent: sharing electron pair(s) Metallic: delocalized electrons
Predicting Bonds Based on electronegativity difference (look at chart on p. 194 in online book) Examples of calculation: H-F 4.0 2.1 = 1.9 H-Br 2.8 2.1 =.7 H-I 2.5 2.1 =.4 The greater the difference, the stronger the bond and the more ionic character it has
Bond Character Large difference: ionic bond Small difference: covalent bond Dividing line is 1.67 > 1.67 is ionic, < 1.67 is covalent, = 1.67 is 50% ionic and 50% covalent Unless bonded to the same type atom, the bond has both ionic and covalent character (use chart)
Finding % Bond Character 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 100-59 = 41% covalent H-Br, H-I
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
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)
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
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
Molecules Covalently bonded compounds Diatomic molecules: always as 2 atoms when in element form (like O 2 ) 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
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: N 2 O 5 Dinitrogen pentaoxide
Lewis Dot Structures Find the total number of valence electrons using group numbers for each element. (Add together) Arrange atoms to form skeleton structure with lines connecting the atoms. Carbon is always central. Otherwise, least electronegative element is central. H is NEVER central-only one bond Each line counts as 2 electrons. Subtract from total valence number. Compare the electrons left to what is needed. If the same #, add unshared dot pairs to give each nonmetal or metalloid a full octet (except H). (Be, Mg, B, Al just double their electrons don t get a full octet, Gr. 2 shares to get 4, Gr. 13 shares to get 6)
Lewis Structures (continued) Count electrons to verify the same number 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. Example: CH 2 O Ions - Add electrons if negative, take away if positive, put brackets around ion and charge in upper right corner. Example: CO 3 2-
Example 1 H 2 O 2(1) + 6 = 8 valence electrons Skeleton:.. H-O-H.. Subtract 2 for each line 8-4 = 4 e - left Put dots to complete octet for oxygen
Example 2 CH 3 I 4+3(1)+7 = 14 valence electrons Skeleton: H H C H : I :.. Subtract 2 for each line 14-8 = 6 e - left Add dots to I to complete the octet. Other Examples: NH 3, AlI 3, SeO 2, CO 2, SO 3
Resonance Using more than one Lewis structure to explain when bonds are in between drawn structures (from lab measurements)
Molecular Shape Based on VSEPR theory: valence-shell electronpair repulsion theory Electrons want to be as far apart as possible (like charges repel) Electron areas around central atom give angles 2 areas: linear 180 o angle 3 areas: trigonal planar 120 o angle 4 areas: tetrahedral 109.5 o angle Repulsion is greater between lone pairs: they push harder on bonded atoms, decreasing the expected bond angle 2 lone>1 shared with one unshared>2 shared
VSEPR phet Link Possible Shapes: Linear: 2 bonded atoms180 o angle Trigonal planar: 3 bonded atoms 120 o angle Bent: 2 bonded, 1 lone pair<120 o angle Tetrahedral: 4 bonded atoms 109.5 o angle Trigonal pyramidal: 3 bonded, 1 lone pair<109.5 o angle (107 o angle) Bent: 2 bonded, 2 lone pairs<109.5 o angle (104.5 o angle)
Determining Shape Draw Lewis structure Count bonded atoms and lone pairs on central atom (ONLY!) to determine shape Example: H 2 O Lewis structure:.. H-O-H.. 2 bonded atoms (lines), 2 lone pairs (dot pairs) Shape: bent, Angles: <109.5 o angle (104.5 o angle) Other Examples: NH 3, AlI 3, CH 4, HF, SO 3
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)
Hybrid Orbitals Count bonded atoms and lone pairs to see how many orbitals are needed. sp hybrids Starting with s, add p orbitals to make enough. Name the sp 3 hybrids hybrids. sp 2 hybrids
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
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: NH 3, H 2 O)
Induced Dipole Peer pressure model Electrons of nonpolar molecule are disturbed by presence of charged particle (ion or dipole)
Dipole-Induced Dipole
Temporary Dipole Movement of electrons may cause electron distribution to become asymmetrical for an instant
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)
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.
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
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
Paper Chromatography Separation on paper into spots or lines on the strip Has limitations
Gas Chromatography Used to analyze volatile liquids and gas or vapor mixtures. Mixed with inert gas (like He) in mobile phase Interpreted by computer
Gas Chromatogram
Chromatography Applications Drug testing uses column and gas chromatography Car emissions are done with gas chromatography