1.  UNIT 4 Bonding and Stereochemistry

2. Stable Electron Configurations  All elements on the periodic table (except for Noble Gases) have incomplete outer energy levels  Valence electrons- electrons in outer energy level of atom  Elements will gain, lose, or share electrons to get full outer levels (octet rule)  Eight electrons = STABLE!!!  Electron dot diagrams help to visualize valence electrons  Symbol represents nucleus and inner electrons  Dots represent valence electrons  Group # = valence electrons

4. Drawing Electron Dot Diagrams  Determine number of valence electrons from periodic table  Draw the symbol for the element  Place dots around the symbol, one per side, until all valence electrons are accounted for  Example- Aluminum with 3 valence electrons Al

6. Ions  Charged atoms where the number of protons and electrons is not equal  Charge indicates how many electrons are added or subtracted  Negative charge- ADD electrons  Positive charge- SUBTRACT electrons  Example Sodium ion  Na atomic number 11 = 11 electrons  Na + subtract one electron = 10 electrons

8. Chemical Bonds  Forces that hold groups of atoms together and make them function as a unit.  A bond will form if the energy of the pairing is lower than that of the separate atoms.  Some elements have stronger attractions to e- when bonded  ELECTRONEGATIVITY (EN)  Relative attraction an atom has for shared electrons in a covalent bond  Unit- paulings  Arbitrary number used for comparison purposes  F is 4.0, Cs/Fr 0.7 Copyright © Cengage Learning. All rights reserved 8

9.  Increase from left to right in a period- nonmetals higher than metals  Decrease from top to bottom in a group  Metals on left side and nonmetals on right side most reactive (alkali metals and halogens)  Electrons attracted to the higher EN element  Using EN to predict bonds  Ionic Bonds- Metal + nonmetal  Covalent Bonds (nonmetal + nonmetal)  Polar- EN is different  Nonpolar- EN is same value

11. Types of Chemical Bonds  Ionic Bonds  Some elements achieve stable configurations by transferring electrons  Example- sodium and chlorine  Sodium 1 valence electron Chlorine 7 valence electrons  Both want to be stable  Sodium will lose the one electron, and chlorine will gain that electron, forming IONS (atoms that have gained or lost electrons) Copyright © Cengage Learning. All rights reserved 11 Na + Cl -

12.  Charge on ion represented by + or – sign  Positive ion- cation  Negative ion- anion (use suffix –ide)  Na + Cl - is sodium chloride (NaCl)  Groups 1, 2, and 3 will lose electrons  Groups 5, 6, and 7 will gain electrons  Group 4 will go either way- usually share though

13.  Ionic compounds- compounds that contain ionic bonds  Can be made with single elements or polyatomic ions  Empirical formula- shows ratios of ions contained in the bond  Na + Cl - one to one NaCl  Mg 2+ Cl - one to two MgCl 2

14.  Crystal Lattices  Each ionic compound makes specific shape based on arrangement  Crystal- solid whose particles are arranged in a lattice structure (NaCl- cubes, ruby- hexagonal)

15.  Properties of ionic compounds  High melting point, boiling point  Poor conductor when solid, good when molten/dissolved  Crystal structure- shatters when hit

17.  Covalent Bonds  Nonmetals have high ionization energy  Don’t usually form ions-share electrons to get to stable energy level  Covalent bond-chemical bond in which two atoms share a pair of valence electrons  May share one (single bond), two (double bond), or three pairs (triple bond) 17

18.  Form molecules  Neutral group of atoms that are joined together by one or more covalent bonds  May exist as diatomic molecules  Made of 2 atoms of same element

19.  May form single or multiple bonds  Molecular Formula- expression of the number and type of atoms that are present in a single molecule of a substance.  Subscript tells how many of each element are present  N 2 O- 2 atoms of N, 1 atom of O

20.  When atoms share electrons, they rarely share equally  One element will “attract” electrons more than the others  Polar Covalent Bond- a covalent bond in which electrons are not shared equally  Atom with greater attraction gets a partial negative charge (δ-), lesser attraction partial positive charge (δ+)

21. Metallic “bonding”  Attraction of metal atoms and the sea of electrons surrounding them  Gives metals their properties  Malleability  Good conductors

22. Writing Lewis Diagrams for Molecules  Steps  Count all valence e-  Draw skeleton structure  Put a pair of e- between all atoms to show a covalent bond  All should have 8 (exc H which has 2)  Distribute lone pairs around atoms (exc H)  If an atom needs more e- then move pairs between atoms to get 8  Most multiple bonds are in C, N, and O  Double bond (share two), triple bond (share three)

24.  A double covalent bond, or simply a double bond, is a covalent bond in which two pairs of electrons are shared between two atoms.  Double bonds are often found in molecules containing carbon, nitrogen, and oxygen.  A double bond is shown either by two side-by-side pairs of dots or by two parallel dashes

25.  A triple covalent bond, or simply a triple bond, is a covalent bond in which three pairs of electrons are shared between two atoms. example 1—diatomic nitrogen: example 2—ethyne, C 2 H 2 :

26. Stereochemistry- VSEPR Theory  All molecules have 3D shape  Stereochemistry- study of shapes of molecules  VSEPR theory  Valence Shell Electron Pair Repulsion  Helps to understand and predict molecuar geometry (from Lewis Dot diagrams)  Developed by Gillespie and Nyholm in 1956-57  Rules based on the idea that the arrangement in space of the covalent bonds formed by an atom depends on the arrangement of valence e-  e- try to push each other far away while still bonding to central atom

27.  Restricted VSEPR rules  Valence e- pairs (both shared and lone) arrange themselves around the central atom in a molecule in such a way as to minimize repulsion (as far away from each other as possible)  When predicting molecular geometry, double and triple bonds act like single bonds  Lone pairs of e- occupy more space than bonding e-

28.  Steps to draw VSEPR molecules  Draw Lewis Diagram  Determine the central atom (lowest EN)  Count the number of bonding and lone pairs surrounding the central atom  Multiple bonds count as one pair  Shape molecule in order to minimize repulsion  Find on VSEPR chart

30. EXAMPLES  Water, H 2 O  2 bond pairs  2 lone pairs  The molecular geometry is BENT.