Chemical Bonding

1. Chemical Bonding

2. Questions you already know the answers to: • What does the atomic number tell us about the electrons in an atom? • Where are electrons located in an atom? • What does a row in the Periodic Table represent? • What and where are valence electrons? • How do we show valence electrons with Lewis Dots?

3. Lewis Dot Review An atom’s valence electrons can be represented by Lewis dot notations.

4. Lewis Dot Review – Period 2 Lewis dot notations for the valence electrons of

5. Lewis Dot Notations

6. Define: • Cation– a positively charged ion • Anion – a negatively charged ion • Oxidation number – the charge on an ion • Electronegativity – How strongly an atom attracts other electrons. • Lone Pair - Those two valence electrons in a Lewis Dot structure that are NOT bonded. • Bonding Pair - Two valence electrons shared between atoms that ARE bonded.

7. An atom that gains one or more electrons will have a ____________________charge. An atom that loses one or more electrons will have a ____________________ charge. An atom that gains or loses one or more electrons is called an _________. A positive ion is called a ______________ and a negative ion is called an _______________. NEGATIVE POSITIVE ION___ CATION ANION

8. Ionic Bonds Atoms will transfer one or more ________________to another to form the bond. Each atom is left with a ________________valence shell… (usually 8, sometimes 2) An ionic bond forms between a ___________cationwith a positive charge and a ________________anion with a negative charge. ELECTRONS FULL METAL NONMETAL

9. Because opposites attract! - anion + cation Na Cl

10. Ionic Bonds Positive cations and the negative anions are attracted to one another… Get a few trillion anions bonded to a few trillion cations and you’ll have a lattice structure, e.g. crystal

11. Properties of Ionic Compounds

12. List of Common Metal & Nonmetal Ions

13. Lewis Dots for Ioniccompounds [ ]1+ [ ]1- Li F [ ]2+[ ]2- Ca O [ ]2+ 2[ ]1- Ca F

14. Form an ionic bond between… • Al + Cl • Mg + P • Ba + O • Ca + Cl • H + P • K + O

15. Examples of Ionic compounds Mg2+ Cl- 2( ) Magnesium chloride: Magnesium loses two electrons and eachchlorine gains one electron 2( ) Na+O2- Sodium oxide: Eachsodium loses one electron and the oxygen gains two electrons 3( ) Aluminum sulfide: Eachaluminum loses three electrons (six total) and eachsulfur gains two electrons (six total) 2( ) Al 3+ S 2-

16. Covalent Bonds Atoms ___________ one or more electrons with each other to form the bond. Each atom is left with a ________________ outer shell. A covalent bond forms between two or more _________________. SHARE COMPLETE NONMETALS

17. Covalent Bonds • Electrons are shared between nonmetals & other nonmetals • Can share 1, 2, or 3 electrons • Neither atom is “strong” enough to swipe It… so they share it.

18. Properties of Covalent Compounds

19. Sharing is caring! O O

20. Ionic compound OR Covalent molecule? • Ionic compound • Covalent molecule • Ionic compound • Covalent molecule • Ionic compound • Covalent molecule • CaO • H2O • AlF3 • N2O • Fe2O3 • C2H6

21. Covalent Bonds 1 line = 2 electrons Single Bond! Cl Cl Cl Cl

22. 2 lines = 4 electrons Covalent Bonds Double Bond! O O O O

23. 3 lines = 6 electrons Covalent Bonds Triple Bond! N N N N

24. Step 1:Write the Lewis dot symbols for each atom. If there are more than 2 atoms, put the least electronegative one in the middle. C O O The least electronegative is usually the one with the least valence electrons

25. Step 2:Bond each atom to its neighbor with a single bond. O O C

26. Step 3:Count the valence electrons around EACH atom… If it has an “octet”, you’re done. If it has less than 8, try another bond. O O C

27. Step 4:Repeat step 3… Once all atoms have full valence shells, pair any unpaired electrons. C O O

28. Lewis Dots for Covalent Bonds O O Cl Cl

29. Dihydrogen monoxide H2O H H O

30. Carbon has lots of ways to shareCH4 is Carbon tetrahydride is methane H H H H C

31. Ammonia NH3 is Nitrogen trihydride is ammonia H H H N

32. Your turn… • I2 • S2 • P2 • OF2 • NF3 • SiH4

33. Non-Polar Covalent Bonds • electronegativity difference between 0.0 - 0.3 • …usually it’s the same diatomic element bonding to itself. • Or if the middle atom is pulled symmetrically in all directions • Ex. O2, N2, CO2, CH4

34. Non-Polar Covalent Bonds Each H is pulling just as strongly on Carbon as all the others are, so this bond is… Non-Polar

35. Polar Covalent Bonds • electronegativity difference between 0.4 – 1.7 • One atom here is tugging harder on the electron than the other one... Which one and why?

36. We can usually tell the type of bondby finding the difference in electronegativity of the two atoms that are bonded.

37. ElectronegativityDifferences Difference Bond Type • 1.7 to 4.0: Ionic • 0.3 to 1.7: Polar Covalent • 0.0 to 0.3: Non-Polar Covalent • Example: NaCl • Na= 0.9Cl= 3.2 3.2 - 0.9 = 2.3

38. Exceptions to the Octet Rule 3 types • Molecules with an odd number of electrons; • Molecules in which one atom has less than an octet; • Molecules in which one atom has more than an octet.

39. Odd Number of Electrons • Few examples. Generally molecules such as ClO2, NO, and NO2 have an odd number of electrons.

40. Less than an Octet • Relatively rare. • Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A. • Most typical example is BF3. • Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds.

41. More than an Octet • This is the largest class of exceptions. • Atoms from the 3rd period onwards can accommodate more than an octet. • Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density.