1. Chapter 5 Ions and Ionic Compounds

2. What are the characteristics of ionic compounds? Unit Essential Question:

3. Why do atoms form ions? Lesson Essential Question:

4. Section 1: Simple ions How reactive atoms are depends upon the number of valence electrons. Certain number makes atoms stable. Octet rule – elements with full outer energy levels tend to be stable. Full energy levels have eight electrons. All noble gases have a satisfied octet rule except He (2 valence electrons). Don’t want to gain or lose an electron.

5. Stability This will be an important theme in this chapter and in the next. Being stable is the goal of atoms and is the reason why bonding takes place. By bonding, atoms change their electron configurations to become more stable. There are several terms/phrases that you will see that indicate an atom is stable: #1: Octet rule is satisfied/achieved. #2: The atom has a noble gas configuration. #3: The atom has a full valence/outer shell.

6. Alkali Metals & Halogens Most reactive elements. Alkali metals-1 valence e - ; halogens- 7 valence e - Full valence shell is achieved if alkali metals lose one e - and halogens gain one e -. Often bond together- the alkali metal gives an e - to the halogen. Example: K: [Ar]4s 1 and Cl: [Ne]3s 2 3p 5 K will lose one e - and Cl will gain one e -. Each now has the same configuration: [Ne]3s 2 3p 6 Both have the same configuration as the noble gas argon- they’re now stable, have an octet!

7. Ions Ion – atom or molecule that has a charge because it has gained or lost electrons. Identity does not change! Cation – ion with a positive charge; e - are lost. The ‘t’ in cation resembles a + sign. This can help you remember that cations are positive. Anion – ion with a negative charge; e - are gained. Ions form as atoms change their valence e - to try and achieve a noble gas configuration. As we saw before: K = [Ar]4s 1  loses 1 e - ; K is now a K + ion. Cl = [Ne]3s 2 3p 5  gains 1 e - ; Cl is now a Cl - ion.

8. Patterns of Charges on the Periodic Table Since elements in the same group have the same # of valence e -, elements in the same group often form the same ions (in the s and p blocks). +1 +2 *Metals tend to lose e - (become cations) and nonmetals tend to gain e - (become anions). +3 -3-2

9. No patterns in the d-block Transition metals do not follow a pattern according to group number. Almost all transition metals form more than one cation. Pb 4+

10. Ions- what changes and what doesn’t Ions and Ionic compounds Properties change! Compounds formed from ions have very different properties from the original atoms. Ex: sodium and chlorine vs. sodium chloride. Remember: Identities don’t change! Ions do not BECOME noble gases! Have electron configurations of noble gases and therefore behave like noble gases. They become more stable.

11. What properties result from forming ionic bonds? Lesson Essential Question:

12. Section 2: Ionic Bonding and Salts Ionic bonds form when oppositely charged ions attract. Salt – name given to many ionic compounds. Ex: NaF, KCl, CaSO 4, etc. Thousands of ionic compounds are called salts. Ionic compounds as a whole are neutral in charge. Not just one cation and one anion come together- many are attracted. Results in a tightly packed crystal lattice structure.

13. Energy and Ionic Bonding There are numerous steps to forming ionic compounds from elements (ie: forming NaCl from Na and Cl). Energy is involved in every step. Overall, more energy is released than is absorbed. This makes forming ionic compounds spontaneous (favorable). Lattice energy: energy released when ionic bonds form.

14. Properties of Ionic Compounds Attractive forces (+ and -) make bonds very strong. High melting and boiling points. Usually solids at room temperature. Conduct electricity when melted (molten) or in solution (dissolved in water). Most dissolve well in water. Hard and brittle as crystals. Due to repeating pattern of ions. Crystal lattice – regular pattern in which a salt crystal is arranged. Made up of many repeating units (smallest called a unit cell).

15. How are ionic bonds represented using formulas and names? Lesson Essential Question:

16. Section 3: Names and Formulas of Ionic Compounds Naming Monatomic Ions Monatomic ions – one atom with a charge. Mono = one. Ex: Al +3, O -2 Cation names do not change. Add ‘ion’ after the name. Na + = sodium ion; Ba +2 = barium ion Transition metals can form more than one ion. In order to tell them apart Roman Numerals are used. Cu + = copper( I ); Cu +2 = copper( II ) Sn +2 = tin( II ); Sn +4 = tin( IV ) Anion names do change. “-ide” ending is added at the end of the element name. S -2 = sulfide; N -3 = nitride; F - = fluoride Value of Roman Numeral tells you the size of the charge.

17. Naming Polyatomic Ions Polyatomic ions – 2 or more atoms with a charge. Poly = many. Ex: NH 4 +, SO 4 -2 (ammonium, sulfate) These are names that you must memorize. Ion Chart You should know all cations and all starred* anions on the ion chart. Remember- monatomic ions in the s and p blocks can be determined just by looking at the periodic table! These are not included on the ion chart but you are expected to know them! There will be a quiz on ion names and charges!

18. Naming Ionic Compounds ALWAYS name and write the cation before the anion. Binary – compounds with 2 elements. Use rules for naming monatomic ions. Ex: NaCl: sodium chloride (Na + and Cl - ) Recall use of Roman numerals for transition metal cations with multiple possible charges! Ex: FeCl 3 : iron (III) chloride (Fe +3 and Cl - ) To help you figure out the charge: For right now looking at the subscript for the anion will help you determine the charge on the transition metal cation. CaBr 2 calcium bromide (Ca +2 and Br - ) Cu 2 O: copper (I) oxide (Cu + and O -2 )

19. Naming ionic compounds cont. Naming compounds with 3 or more elements: Use rules for naming monatomic ions AND polyatomic ion names. Ex: K 3 PO 4 : potassium phosphate (K + and PO 4 -3 ) Notice that parentheses are placed around polyatomic ions whenever there are more than one used in the chemical formula. oOtherwise looks like: NH 42 S Another ex: NH 4 NO 3 : ammonium nitrate (NH 4 + and NO 3 - ) (NH 4 ) 2 S : ammonium sulfide (NH 4 + and S -2 )

20. Writing Chemical Formulas for Ionic Compounds Compounds must be neutral! Positive charge = negative charge. Use the “criss-cross” method: Step 1: determine ions from the name. Step 2: criss-cross numbers only. Step 3: be sure to give the lowest ratio possible. Reduce if necessary! (ex: Sn 2 O 4  SnO 2 ) Note that the positive ion is written first! Use parentheses if there is more than one polyatomic ion! Ex: calcium nitrate: Ca(NO 3 ) 2 (NOT: CaNO 32 ) Don’t write 1’s in formulas. We assume there’s 1 if there’s no number written after the charge. Step 2: Al +3 O -2 Step 3: Al 2 O 3 aluminum oxide Step 1: Al +3 O -2