Ionic Compounds and Metals

1. Ionic Compounds and Metals

2. Types of bonds • A chemical bond is the force that holds atoms together in a molecule. • Ionic bonds form when electrons are transferred between atoms. • Covalent bonds form when electrons are shared between atoms.

3. Electronegativity • Electronegativity is a measure of the attraction a bonded atom has for electrons. • Electronegativity increases from bottom to top and from left to right.

4. Formation of Ionic Bonds • If the electronegativity difference is large (a metal with a nonmetal) an ionic bond will form. • Ionic bonds are the strongest type of bond. • Substances resulting from ionic bonds are classified as crystalline solids.

5. Let’s Practice • Identify the type of bond (ionic or covalent) that will form between each pair of atoms. • Na and Cl • C and O • N and H • Ca and O • K and N • Mg and F

6. Example

7. Forming Ions • When neutral atoms gain or lose valence e- an ion is formed. • Ions are atoms with a charge • Cations have a positive charge (K+, Li+) • Anions have a negative charge (F-, I-)

8. Ions forming bonds • Cations are attracted to anions because of opposite charges. • “Opposites attract”

9. Why do ions form? • Ions form to attain a more stable electron configuration. • Remember that a Noble gas configuration is the most stable Ne - ___ ___ ___ ___ ___ 1s 2s 2p • All Noble Gases have a full octet (8 valence e-)

10. Noble Gases

11. Na __ __ __ __ __ __ 1s 2s 2p 3s Na+ __ __ __ __ __ __ 1s 2s 2p 3s Now Na+ looks like neon Ne __ __ __ __ __ 1s 2s 2p All atoms are trying to have a Noble Gas electron configuration so, they will lose or gain electrons to get there Why do ions form?

12. Why do ions form? • F __ __ __ __ __ 1s 2s 2p F- __ __ __ __ __ 1s 2s 2p Ne __ __ __ __ __ 1s 2s 2p

13. Review Draw the dot notation for each element and the charge of its ion • N • S • Ba • Li

14. Ionic Bonds • Ionic bonds- the force that holds two oppositely (+ and -) charged particles together. • Ionic compounds contain ionic bonds and are neutral • Usually forms between a metal and a nonmetal

15. Forming Ionic Bonds

16. Forming Ionic Compounds Write an equation showing how ionic compounds are formed for the following elements. • Na and I • Mg and O • Al and Cl

17. Properties of Ionic CompoundsShape • The Ionic bond produces unique characteristics for ionic compounds • Crystals are created by repeating patterns of cations and anions. This balances the forces between the ions. • This is called a crystal lattice

18. Properties of Ionic CompoundsPhysical Properties • Electrolytes (aqueous ions- Ions dissolved in water) conduct electricity • Solid ionic compounds DO NOT conduct electricity because the charges are locked in place

19. Properties of Ionic CompoundsPhysical Properties • Ionic solids are strong and hard to break apart • High Melting points • High Boiling points • Transition metals in the crystal give the ionic solid color.

20. Properties of Ionic CompoundsPhysical Properties The solid ionic crystals are hard, rigid, and brittle The strong forces between ions produce this property

21. Ionic Compounds • Ionic compounds are repeating patterns of the cation and anion. • The simplest ratio of ions involved is called a formula unit. • NaCl is the formula unit for table salt

22. Ions in Ionic Compounds • Monatomic ions- ions made of only one type of atom (Cl-, F-, Li+) • Polyatomic ions- ions made of at least two different types of ions (OH-, NH4+) • Cation + Anion  Ionic Compound

23. Writing formulas for ionic compounds Binary ionic compounds contain a single cation and anion. Remember that the final compound must be neutral so the charges must balance. Ca+2 Cl-  CaCl2

24. Cris-Cross method • Record the ion symbol and charge • The # of the charge of the cation becomes the subscript for the anion and the # of the charge for the anion becomes the subscript for the cation. The #1 does not have to be recorded it is understood. Al+3 O-2 Al2O3

25. Cris-Cross method Al2O3 is the formula unit for aluminum oxide. The subscripts indicate how many ions are needed of each in order to balance the charges. The formula unit consists of 2Al3+ ions and 3O2- ions. (6 positive charges cancel out 6 negative charges resulting in a neutral compound)

26. Try a few • Potassium and Iodide • K1+ I1- KI • Magnesium and chloride • Mg2+ Cl1- MgCl2 • Aluminum and bromide • Al3+ Br1- AlBr3 • Cesium and nitride • Cs1+ N3- Cs3N

27. Naming Ionic Compounds NaBr • Name the cation then the anion. Cations are always written first • Monatomic cations always keep their name. If the cation is an ion that may carry multiple charges then a roman numeral is used to indicate which charge the cation is carrying. [Fe+3 (III) and Fe+2 (II)]. • Note: Cations that are not located in groups 1A, 2A, or 3A, require a roman numeral.

28. Naming Ionic Compounds The monatomic anion’s ending is changed to -ide. Fluorine is changed to Fluoride. NaBr- sodium bromide Fe3N2 – Iron (II) nitride Note: Cris cross the subscripts back up to determine the charges: Fe3N2 Fe2+N3-

29. Let’s practice • CaCl2 • Calcium chloride • K2O • Potassium oxide • CuCl2 • Copper (II) chloride • Fe2O3 • Iron (III) oxide

30. What about Polyatomic ions? • Al+3 and PO4-3  AlPO4 • The same rules apply, however NEVER CHANGE A SUBSCRIPT THAT IS ALREADY THERE. • Polyatomic ions may also require () around them if there is a need for more than one to balance the charge

31. Examples • Write the formula for copper (II) sulfate. • Cu2+ • Look up sulfate on the table and record the formula • SO42- • Apply the cris cross method as usual. • Cu2+SO42- CuSO4

32. Another Example • Iron (II) nitrate • Fe2+ • Look up nitrate on the table • NO31- • Apply the cris cross method as usual • Fe2+NO31- Fe(NO3)2

33. Naming compounds containing polyatomic ions • Write the name for Al2(SO4)3 • name the cation • Look up the polyatomic ion on the table and write its name (Do not change the ending) • aluminum sulfate • Try this one: Na3PO4 • Sodium phosphate

34. Metallic Compounds • Compounds containing only metals • Form lattice crystals as solids

35. Metallic Bonds • Metals do not give away or take on valence electrons. • They do not share electrons to make bonds either. • Electrons from metals overlap and create a sea of electrons that bond metal atoms

36. Metallic Bonds • The electrons are free to move between atoms (delocalized electrons) • As an electron moves from one atom to the next it creates a cation which “bonds” the atoms together creating a metallic bond.

37. Properties of Metals • Metals Usually have high melting points. • Metals usually have high boiling points • Why?

38. Properties of Metals • Malleable- ability to be hammered into a thin sheet • Ductile- ability to be stretched into a thin wire • Durable-long lasting • Why?

39. Properties of Metals • Good conductors of heat • Good conductors of electricity • Why?

40. Properties of Metals • Metals are strong and hard • Why?

41. Alloys • Alloys are mixtures of metals • Ex: stainless steel, brass, cast iron • Properties of alloys are similar to those of its elements. Alloys tend to be stronger than the element