1. Chemical Bonds: Metallic & Ionic chapter 7, 9, 13 Vocabulary: leave enough space for definition AND an example 1.Kinetic energy 2.Phase change 3.Evaporation 4.Condensation 5.Metallic bond 6.Alloy 7.Ionic bond 8.Cation 9.Anion 10.Crystal

2. Solids: Particles are strongly attracted to each other Much energy is needed to overcome attractions High melting point Particles vibrate in place fixed shape fixed volume Intermolecular forces Liquids: Particles are weakly attracted to each other Energy is needed to overcome attractions Low melting point Particles move around each other freely no fixed shape fixed volume

3. Gases: Particles are not attracted to each other Gases have enough energy to overcome attractions between particles Extremely low melting point & low boiling point Particles move rapidly and randomly no fixed shape, no fixed volume

4. study question 1 If a substance has no fixed shape it could be: _________ or _________ If a substance has no fixed shape and no fixed volume it would be: ___________ _____ energy is required to melt a solid because particles are ________ attracted to each other A gas is formed when particles gain enough _________ to overcome any _________ between each other

5. Physical state: The state of a material depends on the balance between: 1.the kinetic energy of the particles. 2.the attractions between particles Kinetic energy > attractions = gas Kinetic energy < attractions = liquid Kinetic energy << attractions = solid

6. study question 2 In terms of kinetic energy and intermolecular attractions: explain why iron is a solid at room temperature but oxygen is a gas explain why ice melts at room temperature

7. Phase changes. Melting a solid or evaporating a liquid requires energy to overcome the forces holding the particles together. Freezing a liquid or condensing a gas is caused by removing energy so attractive forces between particles dominate

8. study question 3 Explain why a gas condenses when energy is removed from the particles

9. gas liquid solid melting evaporating Condensing Energy released/removed intermolecular attractions take over freezing Physical state Energy added/absorbed intermolecular attractions are overcome

10. Study question 4 Explain why melting requires energy and freezing releases energy

11. Types of Bonds (intra-molecular) Type of atoms Electrons Solid/ liquid/gas Metallic bond Metals Shared between atoms Solid Ionic bond Metals & non-metals Transferred from one atom to another Solid Covalent bond Non-metals Shared between atoms Gas, liquid or solid

12. study question 5 Provide an example of: A metallic compound An ionic compound A covalent solid A covalent liquid A covalent gas

13. All metals have extra valence electrons. Instead of losing electrons to form ions, metal atoms can also share valence electrons with other metals.  Atoms are very close together ஃ Metals are solid compounds  Electrons move around (sea of electrons) ஃ Metals conduct electricity ஃ Metals are malleable.  ex: lead Bonding in Metals

14. study question 6 Describe the unique characteristic that allows metals to conduct electricity when non-metals can not. Based on the observation that mercury is a liquid, what can you say about the attractions between mercury atoms? Do you think mercury conducts electricity? why/why not?

15. An Ionic bond is formed between metal and non- metal atoms Each atom gains or loses electrons in order to form an octet An ion is a charged particle A cation = positive ion = a Metal. oe.x. Na +, Ca 2+, Al 3+ An anion = negative ion = a Non-metal oe.x. Cl -, S 2-, P 3- Ionic bonds.

16. study question 7 For the compound K 2 O: Which atom is the cation? Which atom is an anion?

17. Crystal Lattice Structure All ionic compounds form a crystal lattice structure oformed by a very large network of electrostatic attractions (positive and negative ions attracted to each other). oLattice energy: The energy needed to break the electrostatic attractions holding the lattice structure. Ionic compounds are always solid because of the strong electrostatic attractions between ions

18. Crystal Lattice Structure. + + ++ +

19. study question 8 Explain why crystal lattice structures are always solid study question 9 Which of the following are solid at room temperature? CaO CO NO 2 Na 2 O

20. e.x. Na 2 O naming ionic compounds 1.Use cation name 2.modify anion name (ide) 3.Do NOT use prefixes Name = sodium oxide More examples: CaF 2 = calcium fluoride K 2 O = potassium oxide

21. study question 10 Name the following 1.MgCl 2 2.Al 2 O 3 3.NaBr

22. Polyatomic ions: act as one charged particle in an ionic bond Anion names are not modified for polyatomic ions  Example: NH 4 OH = Ammonium hydroxide  Example: Al(NO 3 ) 3 = aluminum nitrate Weird things

23. study question 11 Name the following 1.NaOH 2.K 2 SO 4 3.Mg(NO 3 ) 2

24. example:  Sodium hydroxide (Na + and OH - )  Charges must cancel out = NaOH  example:  Magnesium hydroxide (Mg 2+ and OH - ) For each Mg 2+ there must be 2 x OH - = Mg(OH) 2 note: use parenthesis only when showing more than one polyatomic ion Writing ionic formulas from names

25. study question 12 Write formulas for the following compounds: 1.Aluminum hydroxide 2.Potassium oxide 3.Magnesium nitrate