2. Bonding between atoms

3. Bonds Forces that hold groups of atoms  Forces that hold groups of atoms together and make them function together and make them function as a unit. as a unit. Ionic bonds – transfer of electrons  Ionic bonds – transfer of electrons  Covalent bonds – sharing of electrons

4. Molecule: collection of atoms that are bound together What holds a molecule together? Covalent bond Force of attraction that results from valence electrons being attracted to two nuclei. Valence electrons being shared between two nuclei Core electrons are not involved in bonding

5. Electronegativity The ability of an atom in a molecule to attract shared electrons to itself. The ability of an atom in a molecule to attract shared electrons to itself. Linus Pauling 1901 - 1994

6. Table of Electronegativities

7. Polar-Covalent bonds Nonpolar-Covalent bonds Covalent Bonds  Electrons are unequally shared  Electronegativity difference between.3 and 1.7  Electrons are equally shared  Electronegativity difference of 0 to 0.3

8. Energy is released when forming a covalent bond and absorbed when breaking it

9. Octet Rule and electron sharing

10. Lewis Dot diagrams Show an atom’s valence electrons The number of dots is equal to the group number Pairs of dots(electrons) are electron pairs Single electrons are unpaired electrons

11. Examples of Lewis Dot diagrams All elements in group IA will have one dot All elements in group IIA will have two dots All elements in group IIIA will have three dots Group IV will have four dots, V five, VI six, VII seven, and VIII eight dots

12. Metal Metalloids Nonmetals IIIIIIIVVVIVIIVIII Transition metals

13. C H H.. Completing a Lewis Structure - CH 3 Cl Add up available valence electrons: C = 4, H = (3)(1), Cl = 7 Total = 14 Join peripheral atoms to the central atom with electron pairs. Complete octets on atoms other than hydrogen with remaining electrons Make carbon the central atom H H Cl..

14. Lewis Structures 1) Count up total number of valence electrons 2) Connect all atoms with single bonds - “multiple” atoms usually on outside - “single” atoms usually in center; C always in center, H always on outside. 3) Complete octets on exterior atoms (not H, though) 4) Check - valence electrons math with Step 1 - all atoms (except H) have an octet;if not, try multiple bonds - any extra electrons?Put on central atom

15. Comments About the Octet Rule 2nd row elements C, N, O, F observe the octet rule. 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals. When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

16. Exceptions to the Octet Rule Atoms with less than an octet: BF 3 Atoms with more than an octet: most notably phosphorus and sulfur (SF 4 ) additional electrons fill the 3d orbitals Molecules with odd number of electrons: NO (nitrogen monoxide)

17. Multiple bonds Determine the total number of valence electrons that will be in the final diagram Connect atoms with single bonds Put in remaining dots two at a time as lone pairs Send lone pairs to atoms that do not have an octet(lone pairs must come from adjacent atoms)

18. Multiple Covalent Bonds: Double bonds Two pairs of shared electrons

19. Multiple Covalent Bonds: Triple bonds Three pairs of shared electrons

20. Ionic Bonding  Forms between metals and nonmetals  There is a strong attraction between oppositely charged ions  Form ionic lattices (network of positively and negatively charged ions)  Examples: sodium oxide, magnesium chloride, calcium phosphate

21. Properties of Ionic Compounds Structure:Crystalline solids Melting point:Generally high Boiling Point:Generally high Electrical Conductivity: Excellent conductors, molten and aqueous Solubility in water: Generally soluble

22. Ionic Bonds  Electrons are transferred  Electronegativity differences are generally greater than 1.7  The formation of ionic bonds is always exothermic!

23. Ionic Bonding: The Formation of Sodium Chloride  Sodium has 1 valence electron Cl 1s 2 2s 2 2p 6 3s 2 3p 5 Na 1s 2 2s 2 2p 6 3s 1  Chlorine has 7 valence electrons  An electron transferred gives each an octet

24. Ionic Bonding: The Formation of Sodium Chloride Cl - 1s 2 2s 2 2p 6 3s 2 3p 6 Na + 1s 2 2s 2 2p 6 This transfer forms ions, each with an octet:

25. Ionic Bonding: The Formation of Sodium Chloride Cl - Na + The resulting ions come together due to electrostatic attraction (opposites attract): The net charge on the compound must equal zero

26. Sodium Chloride Crystal Lattice Ionic compounds form solids at ordinary temperatures. Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions.

27. Monatomic Cations Name H+H+ Hydrogen Li + Lithium Na + Sodium K+K+ Potassium Mg 2+ Magnesium Ca 2+ Calcium Ba 2+ Barium Al 3+ Aluminum

28. Writing Ionic Compound Formulas Example: Barium nitrate 1. Write the formulas for the cation and anion, including CHARGES! Ba 2+ NO 3 - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! ( ) 2

29. Writing Ionic Compound Formulas Example: Ammonium sulfate 1. Write the formulas for the cation and anion, including CHARGES! NH 4 + SO 4 2- 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! ( ) 2

30. Writing Ionic Compound Formulas Example: Iron(III) chloride 1. Write the formulas for the cation and anion, including CHARGES! Fe 3+ Cl - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! 3

31. Writing Ionic Compound Formulas Example: Aluminum sulfide 1. Write the formulas for the cation and anion, including CHARGES! Al 3+ S 2- 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! 23

32. Writing Ionic Compound Formulas Example: Magnesium carbonate 1. Write the formulas for the cation and anion, including CHARGES! Mg 2+ CO 3 2- 2. Check to see if charges are balanced. They are balanced!

33. Writing Ionic Compound Formulas Example: Zinc hydroxide 1. Write the formulas for the cation and anion, including CHARGES! Zn 2+ OH - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! ( ) 2

34. Writing Ionic Compound Formulas Example: Aluminum phosphate 1. Write the formulas for the cation and anion, including CHARGES! Al 3+ PO 4 3- 2. Check to see if charges are balanced. They ARE balanced!

35. Naming Ionic Compounds 1. Cation first, then anion1. Cation first, then anion 2. Monatomic cation = name of the element2. Monatomic cation = name of the element Ca 2+ = calcium ionCa 2+ = calcium ion 3. Monatomic anion = root + -ide3. Monatomic anion = root + -ide Cl  = chlorideCl  = chloride CaCl 2 = calcium chlorideCaCl 2 = calcium chloride

36. Naming Ionic Compounds 1. Cation first, then anion1. Cation first, then anion 2. Monatomic cation = name of the element2. Monatomic cation = name of the element Ca 2+ = calcium ionCa 2+ = calcium ion 3. Monatomic anion = root + -ide3. Monatomic anion = root + -ide Cl  = chlorideCl  = chloride CaCl 2 = calcium chlorideCaCl 2 = calcium chloride

37. Naming binary covalent compounds Name less electronegative element first The more electronegative element takes a suffix-ide For each element, use a prefix to indicate number of atoms present

38. Naming Binary Covalent Compounds  Compounds between two nonmetals  First element in the formula is named first.  Second element is named as if it were an anion.  Use prefixes  Only use mono on second element - P 2 O 5 = CO 2 = CO = N 2 O = diphosphorus pentoxide carbon dioxide carbon monoxide dinitrogen monoxide

39. Covalent compounds Phosphorous trichloride Carbon dioxide Sulfur hexafluoride

40. Naming chemical compounds Compounds are named according to atoms or ions that compose them Binary molecular are named after component atoms, with prefixes An acid is named after its characteristic anion

41. Bond type by electronegativity Electronegativity difference less or equal to 0.4  bond type nonpolar covalent Between 0.4 and 2.0 polar covalent Greater than 2.0 ionic