Metals  Lose e-s  positive ion Octet Rule & Ions Atoms may or may not be stable due to their nuclear composition They become more stable by undergoing radioactive decay Atoms may or may not be stable due to their electronic composition They can become more stable by undergoing ion formation Metals  Lose e-s  positive ion called cations Na+ Mg2+ Fe3+ Ions are electrically charged particles Nonmetals  Gain e-s  negative ion called anions Cl - O2- P3-

The most stable atoms are the Noble gases Atoms lose or gain e-s to acquire a noble gas electron configuration (8 valence electrons) is called the octet rule. Exception [He] 2 valence e-s.

Na Na+ + e- Metals lose electrons due to their low ionization energy. IE = energy required to remove an electron M + energy → M+ + e- Na Na+ + e- [Ne] = Stable State 8 valence e-s

Nonmetals gain electrons due to their high electron affinity. EA = energy change when an electron is added. X + e− → X− + energy [Ar] = stable state 8 valence e-s Cl + e- Cl-

Most common ionic charges Transition metals Charge is always given by means of: The name Iron (III) chloride The formula FeCl3

Ionic Compounds http://youtu.be/uCwHzTsx5yY

Na2S From charges to chemical formulas Mg202 Mg0 When ionic particles react, they do so in a ratio that cancels their charge. This is what is referred to as “charge balance”. Na2S “cross-the-charges”. Al3+ 02- 2 3 Al203 Always keep the simplest ratio Mg2+ 02- 2 2 Mg202 Mg0

Practice 1. How many e-s will the following elements need to lose or gain to fulfill the octet rule? N Cl Na Al O Mg 2. What is the ionic charge developed by the following elements? N Cl Na Al O Mg 3. The symbol for the ion formed by : N Cl Na Al O Mg 4. Write the formula of the ionic compound formed by: a. Calcium and chlorine b. Magnesium and nitrogen c. Barium and oxygen

Naming & Writing Ionic Formulas NaCl Sodium chloride FeCl3 Iron (III) chloride

Polyatomic Ions Note: A polyatomic ion taken more than once Compounds containing polyatomic ions are named as ionic compounds. Na+ SO42- 2 1 Na2S04 Sodium sulfate Al3+ SO42- 2 3 Al2 (S04)3 Aluminum sulfate Figure 05-T07 Title: Names and Formulas of Some Common Polyatomic Ions Caption: Note: A polyatomic ion taken more than once use parenthesis

Practice 5. Provide the name for the following ionic compounds: Mg3N2 Fe2O3 Cu2S AuCl3 CaSO4 Cu(NO2)2 6. Provide the formula for the following ionic compounds: iron(III) chloride chromium(III) oxide aluminum bicarbonate calcium hydroxide 7. Provide the formula for a compound containing ammonium ion(s) and phosphate ion(s). 8. Draw the electron-dot structure for SO2

Between two nonmetal atoms. Covalent Compounds Covalent bonds form: Between two nonmetal atoms. When atoms share valence electrons to complete octets (duet for hydrogen). Denoted by a line Gilbert Newton Lewis Figure 05-03-05UN Title: Formation of a Hydrogen Molecule Caption: In the covalent bond in H2, the shared electrons give the noble gas configuration of He to each of the H atoms. Thus the atoms bonded in H2 are more stable than two individual H atoms. Has a noble gas electron count Becomes more stable Has lower energy

H H An electron pair being shared is denoted by a line. In his honor, we refer to electron-dot structures as Lewis structures. H H Lewis further proposed that by counting valence electrons, it would be possible to predict how many bonds an atom can form: H O N C 1 2 3 4

Electron-dot structure = Lewis structure All valence electrons are shown as dots or bonding electrons are shown as lines. F2 Fluorine In a single bond One pair of electrons is shared. Carbon dioxide C : O .. C : O .. In a double bond, Two pairs of electrons are shared. Hydrogen cyanide In a triple bond. Three pairs of electrons are shared. : N C H : N C H

Steps Guidelines to Writing Lewis Structures NH3 N 5 H  1 x 3 = 3 1. Use the group number to determine the total number of valence e-s N 5 H  1 x 3 = 3 Total = 8 2. Draw a “skeleton” by connecting the atoms with lines. First element of the formula is at the center, the others are surrounding. H | H—N—H 3. Provide octets or duets to surrounding atoms & find the remaining e-s. H | H—N—H 8 -6 =2 4. Use the remaining electrons to provide an octet to the central atom. 5. If not enough e-s remain, pull surrounding e-s to form double or triple bonds. H | H—N—H 6. Any e-s remaining after all octets are provided, are given to the central atom.

SO3 S 6 O  6 x 3 = 18 Total = 24 O | O—S—O 24 O -24 | =0 O—S—O O || 1. Use the group number to determine the total number of valence e-s S 6 O  6 x 3 = 18 Total = 24 2. Draw a “skeleton” by connecting the atoms with lines. First element of the formula is at the center, the others are surrounding. O | O—S—O 3. Provide octets or duets to surrounding atoms & find the remaining e-s. 4. Use the remaining electrons to provide an octet to the central atom. O | O—S—O 24 -24 =0 5. If not enough e-s remain, pull surrounding e-s to form double or triple bonds. O || O—S—O 6. Any e-s remaining after all octets are provided, are given to the central atom.

Naming & Writing Covalent Formulas STEP 1 Name the first nonmetal by its element name. STEP 2 Name the second nonmetal with an --ide ending. STEP 3 Add prefixes to indicate the number (from subscripts) of atoms of each nonmetal. Mono is always omitted for the first element and usually omitted for the second one. Exception is CO carbon monoxide Examples of compounds of N and O: NO nitrogen oxide nitric oxide nitrogen (I) oxide NO2 nitrogen dioxide N2O dinitrogen oxide N2O3 dinitrogen trioxide N2O4 dinitrogen tetroxide nitrous oxide nitrogen (II) oxide N2O5 dinitrogen pentoxide

Chapter 6, Figure 6.4, A flowchart illustrates naming for ionic and molecular compounds.

Practice 9. Write the electron-dot formula for PCl3, phosphorus trichloride. 10. Write the electron-dot formula for CO2, carbon dioxide. 11. What is the name for P4S3 12. Write the correct formula for each of the following: A. phosphorus pentachloride B. dinitrogen trioxide C. sulfur hexafluoride 13. Identify each compound as ionic or covalent and give its correct name. A. SO3 B. BaCl2 C. (NH4)3PO4 D. Cu2CO3 E. N2O4

Electronegativity & Bond Polarity Electrons can be Equally shared Electrons can be Unequally shared +1 +17 Yes, separation of charge = polar +1 -1 = 0 -1 +1 = 0 No separation of charge = nonpolar

Electronegativity indicates the attraction of an atom for shared electrons. Linus Pauling When a separation of charges exists in a bond, it is a polar bond. Always on the highest EN value

Chapter 5, Table 5.14

Practice 14. Which element is most electronegative? a. N b. F c. O d. C 15. Use the electronegativity difference to identify the type of bond between the following: nonpolar covalent (NP), polar covalent (P), or ionic (I). A. K-N B. N-O C. Cl-Cl D. H-Cl 16. Which compound contains a polar covalent bond? a. HBr b. NaF c. F2 d. H2 17. Which of the following would have the most polar bond? a. Cl – Cl b. F – Cl c. F – I d. I - Br

Shapes & Polarity of Molecules On paper, molecules appear to be flat, two dimensional objects. Instead, they have a 3-D structure The valence-shell electron-pair repulsion theory (VSEPR) predicts the shape based on bonding and lone pairs (“electron groups”) around the central atom. 4 Electron Groups  Four bonding 4 Electron groups  Two Bonding and Two Lone Pairs VSPER pronounced V (like in Victor) – and (spr)

Chapter 5, Table 5.16

Polar bonds can produce  a polar molecule  or a nonpolar molecule Based on EN, a bond can be polar. Based on shape, a bond can be polar, but a molecule nonpolar

For multiple dipoles, it helps to use “center of (+)” and “center of (-)”

Example Determine the polarity of the H2O molecule. Solution: Write the electron-dot structure Determine the shape & bond polarity. Determine the molecular polarity. 29 29

Practice 18. What is the shape of the SO3 molecule. Provide a drawing of the molecule. a. Trigonal planar b. Bent c. Tetrahedral d. Trigonal Pyramidal 19. What is the shape of the NH3 molecule. Provide a drawing of the molecule. a. Linear b. Bent c. Tetrahedral d. Trigonal pyramidal 20. Determine the shape of each of the following molecules and whether they are polar or nonpolar. If polar show the direction of the dipole. PBr3 B. HBr C. Br2 D. SiBr4