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1. IONS UNIT 1B.10Ionic Bondinghttps://www.youtube.com/watch?v=g_hwgHGeRhAhttps://www.youtube.com/watch?v=WWc3k2723IM https://www.youtube.com/watch?v=hiyTfhjeF_U

2. Bohr’s Model • Electrons move like planets around the sun. • In circular orbits at different levels. • Amounts of energy separate one level from another.

3. III. The Bohr Model of the Atom • Electrons of hydrogen circle the nucleus in orbits 1. orbits have a fixed amount of energy in the ground state 2. orbits are a fixed distance from the nucleus 3.orbits furthest from the nucleus have the greatest energy

4. Niels Bohr (1885 – 1962) Bohr Model of the Atom

5. Bohr’s Model Nucleus Electron Orbit Energy Levels

6. Bohr’s Model } • Further away from the nucleus means more energy. • There is no “in between” energy • Energy Levels Fifth Fourth Third Increasing energy Second First Nucleus

7. IONS • An atom usually has a neutral charge. That means it has the same number of protons as electrons • Remember, a proton has a positive charge and an electron has a negative charge • ION – an atom that has lost or gained one or more electrons and has become charged either positively or negatively

8. Positive Ions • When an atom LOSES electrons, it becomes more POSITIVE • Why? • If you are getting rid of negative particles (electrons) but your number of positive particles (protons) are staying the same. • In other words, you are subtracting negative numbers

9. Examples • What would the charge be if: • The neutral form of Gold (Au) lost 4 of its 79 electrons. It now has 79 protons and 75 electrons • The neutral form of Mg lost 2 of its 12 electrons. It now has 12 protons and 10 electrons.

10. Negative Ions • When an atom GAINS electrons it becomes more NEGATIVE • Why? • Electrons have a negative charge, so the more you have, the more negative you become

11. Representing Ions • Ions are represented by placing a “superscript” charge number next to the atomic symbol. • Ex. • O-2 = oxygen with a negative 2 charge • K+ = potassium with a positive 1 charge • N-3 = nitrogen with a negative 3 charge • And so on

13. A. Chemical Bond and valence electron 1.The electrons responsible for the chemical properties of atoms are those in the outer energy level: VALENCE ELECTRONS. a.Valence electrons - The electrons in the outer energy level. b.Inner electrons -those in the energy levels below.

14. Keeping Track of Electrons 2. Atoms in the same column a. Have the same outer electron configuration. b. Have the same valence electrons. c. Easily found by looking up the group number on the periodic table. d. Group 2A - Be, Mg, Ca, etc.- 2 valence electrons

16. List the number of valence shell electrons are in each of the elements in groups • 1-2 13-18

17. B. Electron Dot Diagrams: Lewis Structures 1. A way of keeping track of valence electrons. 2. How to write them 3. Write the symbol. 4. Put one dot for each valence electron 5. Don’t pair up until they have to X

18. The Electron Dot diagram for Nitrogen • Nitrogen has 5 valence electrons. • First we write the symbol. N • Then add 1 electron at a time to each side. • Until they are forced to pair up.

19. Write the electron dot diagram for Na Mg C O F Ne He

20. C. Ion Formation in Representative Elements Group Gain or Lose Charge of Ion 1 (1A) lose 1 +1 2 (2A) lose 2 +2 13 (3A) lose 3 +3 14 (4A) lose or gain 4* +4,-4* 15 (5A) gain 3 -3 16 (6A) gain 2 -2 17 (7A) gain 1 -1

21. C. Electron Configurations for Cations 1. Metals lose electrons to attain noble gas configuration. 2. They make positive ions. 3. If we look at electron configuration it makes sense. • Na 1level- 2e • 2 level -8e • 3 level -1e: 1 valence electron

22. Electron Dots For Cations • Metals will have few valence electrons Ca

23. Electron Dots For Cations • Metals will have few valence electrons • These will come off Ca

24. Electron Dots For Cations • Metals will have few valence electrons • These will come off • Forming positive ions Ca+2

25. D. Electron Configurations for Anions 1. Nonmetals gain electrons to attain noble gas configuration. 2. They make negative ions. 3. If we look at electron configuration it makes sense. • S 1s22s22p63s23p4: 6 valence electrons • S-2 1s22s22p63s23p6:noble gas configuration.

26. Electron Dots For Anions • Nonmetals will have many valence .electrons. • They will gain electrons to fill outer shell. P P-3

27. E. Stable Electron Configuration 1.All atoms react to achieve noble gas configuration. 2.Noble gases 8 electrons on last energy level. 3. 8 valence electrons . 4. Also called the octet rule. Ar

28. Write the electron configuration diagram label as anion or cation • Na • Mg • P • O • F • Cl • K

29. I. Chemical Bonds • A. The force that holds two atoms together. 1. Why do atoms form bonds? a. to acquire 8 electrons in the valence shell (like noble gases 2. How do atoms form bonds? a. atoms may lose, gain or share electrons to get 8 in the valence shell

30. I. Properties of Ionic Compounds View Ionic video a. Crystalline structure. b. A regular repeating arrangement of ions in the solid. c. Structure is rigid.

31. II. Ionic Bonding A. Anions and cations are held together by opposite charges. B. Ionic compounds are called salts. C. Simplest ratio is called the formula unit. D.The bond is formed through the transfer of electrons. E. Electrons are transferred to achieve noble gas configuration.

32. Formation and Nature of Ionic Bonds F. 1. atom “M” loses electron(s)  cation 2. atom “N” gains electron(s)  anion 3. cation and anion attract each other a. electrostatic attraction 4. the electrostatic force that holds the oppositely charged ions together is the ionic bond

34. Sodium loses an electron forming a (+) ion. Chlorine gains an electron forming a ( - ) ion. Electrostatic attraction between the (-) and (+) ion forms the ionic bond between sodium and chlorine

35. B. Ionic Compounds 1. compounds containing ionic bonds 2. types of ionic compounds a. oxides– metal + oxygen Na2O, CaO, Al2O3, Fe2O3 b. salt – metal + nonmetal NaCl CaF AgCl KI c. binary compounds – two elements 1)all of the compounds in a and b are binary compounds

36. D. monoatomicion - one atom 1) ex. K+ Fe3+ O2- N3- E. polyatomic ion – ion with more than one atom that acts as a single ion NO3- OH- SO42- NH4+

37. III. Names and Formulas • Formulas for Ionic Compounds 1. vocabulary a. formula unit – simplest ratio of ions in a compound 1) ex. NaCl MgBr2 AlCl3 b.monoatomicion - one atom 1) ex. K+ Fe3+ O2- N3-

38. B. Formulas For Ionic Compounds 1. write formula for the cation first, then the anion 2. use subscripts to indicate number of ions (criss-cross the charges) a. sum of charges should equal 0 b. never change subscripts in polyatomic ions

39. Sodium Chloride Crystal

40. Ionic Bonding Na Cl

41. Ionic Bonding: Lewis Structure Na+ Cl -

42. Ionic Bonding • All the electrons must be accounted for! Ca P

43. Ionic Bonding Ca P

44. Ionic Bonding Ca+2 P

45. Ionic Bonding Ca+2 P Ca

46. Ionic Bonding Ca+2 P-3 Ca

47. Ionic Bonding Ca+2 P-3 Ca P

48. Ionic Bonding Ca+2 P-3 Ca+2 P

49. Ionic Bonding Ca Ca+2 P-3 Ca+2 P

50. Ionic Bonding Ca Ca+2 P-3 Ca+2 P