Transcript Slide 1
Chapter 6: Chemical Bonding Cartoon courtesy of NearingZero.net 6-1 Chemical Bonds • Forces that hold groups of atoms together and make them function as a unit. • Ionic bonds – transfer of electrons Metal + Nonmetal Ex) NaCl Li2O • Covalent bonds – sharing of electrons. 2 nonmetals Ex) H2O CO2 • Metallic bonds- electrons are free to move throughout the material. Metals Covalent Bonds 1. Nonpolar-Covalent bonds (H2) • Electrons are equally shared • Electronegativity difference of 0 to 0.3 2. Polar-Covalent bonds (HCl) • Electrons are unequally shared • Electronegativity difference between .3 and 1.7 Using Electronegativity differences • SCs • S=2.5 • Cs= 0.7 2.5-0.7= 1.8 look on table pg 162 1.7 thru 3.3= Ionic Bond 6-2 Covalent Bonding • Molecule- is the smallest unit quanitity of matter which can exist by itself and retain all the properties of the original substance. • Examples H2O & O2 • Diatomic molecule- is a molecule containing 2 identical atoms. (H2 N2 O2 F2 Cl2 Br2 I2) • H NO F • Chemical Formula- represents the relative # of atoms of each kind in a chemical compound by using atomic symbols and numerical subscripts. Example: H2O H=2 O=1 • Molecular compound (Covalent Compounds) - simplest formula unit are molecules. Have low melting & boiling pts. • Molecular formula- shows the types and numbers of atoms combined in a single molecule. • Bond Length- is the average distance between 2 bonded atoms. • Bond Energy- is the energy required to break a bond. • It gives us information about the strength of a bonding interaction. • Resonance-Occurs when more than one valid Lewis structure can be written for a particular molecule. The Octet Rule • Chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons, has an (8) octet of electrons in its valence shell. H He 8 is Great! Lewis Dot • Shows how valence electrons are arranged among atoms in a molecule. • Reflects central idea that stability of a compound relates to noble gas electron configuration. I started it for you CH4 Try one yourself NH3 H2O Structural Formula • Shows shared pair of electrons by a dashed line. • • • • • • Single bond- 1 pair of electrons Double bond- 2 pair of electrons Triple bond- 3 pair of electrons Try a couple: O2 N2 6-3 Ionic Bonding and Compounds • Ionic Compound- is composed of positive and negative ions combined so that the positive & negative charges are equal. (Metal + nonmetal) • Formula Unit- is the simplest collection of atoms from which a compounds formula can be established. 1. 2. 3. Writing Formulas Write the symbols of each element Put their charge in their upper right corner Crisscross the numbers down (Not the charges). Example: Write the formula for Magnesium Chloride Mg Cl Mg+2 Cl-1 Mg+2 Cl-1 MgCl MgCl22 Writing Formulas Practice +1 • 1. 2. 3. 4. +2 Write the formula for: Aluminum Bromide Calcium Oxide Calcium Nitride Sodium Chloride +3 4 -3 -2 -1 • Lattice Energy- is the energy released when 1 mole of an ionic crystalline compound is formed from gaseous ions. • Ionic Compounds have high melting points & boiling points, are hard and brittle, have crystalline structure. Polyatomic Ions • Many atoms with a charge. • Example SO4-2 6-4 Metallic Bonding • Metals- conduct heat, have low ionization energy & electronegativity, give up e• Metallic Bond- is a chemical bond resulting from the attraction between positive ions and surrounding mobile e-. • Malleability and ductility 6-5 Molecular Geometry • VSEPR Theory- “Valence- shell, electron-pair repulsion” • states that repulsion between the sets of valence-level electrons surrounding an atom cause these sets to be oriented as far apart as possible. Determining VSEPR • 1. 2. 3. 4. 5. Determine the VSEPR for H2O Draw the Lewis Dot Draw the Structural Formula Label the central atom as A Label any atoms attached to the center atom as B Label any paired electrons on the central atom that are not used in the bond as E H-O-H B E2 A B VSEPR AB2E2 Shape Bent (look on chart) VSEPR Chart VSEPR AB or AB2 AB2E AB3 AB4 AB3E AB2E2 AB5 AB6 SHAPE Linear Bent Trigonal-Planar Tetrahedral Trigonal-Pyramidal Bent Trigonal-Bipyramidal Octahedral • Hybridization-The Blending of Orbitals. • Dipole- is created by equal but opposite charges that are separated by a short distance. • Dipole-Dipole Attractions-Attraction between oppositely charged regions of neighboring molecules. • Hydrogen Bonding- Bonding between hydrogen and more electronegative neighboring atoms such as oxygen and nitrogen. Hydrogen bonding in Kevlar, a strong polymer used in bullet-proof vests. • London Dispersion Forces- The temporary separations of charge that lead to the London force attractions are what attract one nonpolar molecule to its neighbors. London forces increase with the size of the molecules. • • • • • • • “Electronegativity chart”. Table. Aug. 9, 2006. http://www.chemistry210.com/notes/u01s06f.htm “Lewis Structures”. Drawings. Aug. 9, 2006. http://www.avonchemistry.com/chem_bond_explain.html “Oscar”. Photo. Aug. 9, 2006. http://www.musicmerchant.com/22061.htm “Water Structural Formula”. Drawing. Aug. 10, 2006. http://www.accs.net/users/kriel/ch4notes/water_structura l_formula.gif “Periodic Table of Elements”. Chart. Aug. 9, 2006. http://users.erols.com/kdennis/periodictable.jpg “Information”. Aug 11, 2006. http://www.sciencegeek.net/Chemistry/Powerpoin t/Unit3/Unit3_files/frame.htm Holt, Rinehart and Winston. Modern Chemistry. Harcourt Brace & Company. 1999.