2. Chapter 6.1 Introduction to Chemical Bonding

3. Why do elements bond?  They want to become more stable elements, which they achieve by having 8 valence electrons, causing a decrease in the atom’s potential energy.

4. Types of Chemical Bonding  When atoms bond, their valence electrons are moved around in a way to make the atom more stable.  Ionic Bonding: the electrical attraction between cations and anions.  Covalent Bonding: the sharing of electrons between non-metals

5. Ionic or Covalent?  Looking at the difference in electronegativity of two elements will tell us which bonding it will favor.  Electronegativity difference of 1.7 or less will favor covalent and a difference of 1.8 or more will favor ionic.

6. Types of Covalent  Nonpolar-covalent: electrons are shared equally. Difference of 0.0 – 0.3  Polar-Covalent: electrons are not shared equally, one atom will have a stronger pull on the electrons. Difference of 0.4 to 1.7

7. Example  Use electronegativity differences to classify bonding between sulfur, and hydrogen, cesium, and chlorine. In each pair, which atom will be more negative.  Look at figure 20 on page 161 to find electronegativities.

8.  Sulfur and Hydrogen ▪ 2.5-2.1 = 0.4 ▪ polar-covalent ▪ sulfur  Sulfur and Cesium  2.5-0.7 = 1.8  ionic  sulfur  Sulfur and Chlorine  3.0-2.5 = 0.5  polar-covalent  chlorine

9. Chapter 6.2 Covalent Bonding and Molecular Compounds Fructose Carbon Dioxide Water Ammonia

10. Why Do Atoms Bond?  To get eight valence electrons  To become more stable  In ionic bonds, metals lose electrons and non-metals gain electrons.  What happens when both elements need electrons?

11. Molecules and Molecular Compounds  Compounds that are NOT held together by an electrical attraction, but instead by a sharing of electrons.  Atoms held together by sharing joined by a covalent bond.  Atoms held together by sharing electrons and filling the outer energy levels are joined by a covalent bond.  NONMETALS ONLY!! - No metals

12. Molecules and Molecular Compounds  A molecule is a neutral group of atoms joined together covalent bonds. A compound composed of molecules is called a molecular compound.  The chemical formula for a molecule is called the molecular formula.  A chemical formula tells you how many atoms of each element one molecule of a compound contains.

13. Learning Check Indicate whether a bond between the following would be 1) Ionic2) covalent ____A. sodium and oxygen ____B. nitrogen and oxygen ____C. phosphorus and chlorine ____D. calcium and sulfur ____E. chlorine and bromine Ionic Covalent

14. Monatomic (One Atom)  Noble gases are monatomic.  They exist as single atoms and do not combine with any other elements.  Ex: He, Ne, Ar, Kr, Xe, Rn

15. 7 Diatomic Molecules  Some elements will covalently bond to themselves to form a molecule composed of TWO atoms.  Some elements occur as “diatomic” molecules in nature because they are more stable than individual atoms  The 7 diatomic elements are all gases: H 2, O 2, N 2, Cl 2, Br 2, I 2, F 2

16. Strength of Covalent Bonds  Distance between two bonding nuclei at the position of maximum attracting is bond length  Bond length is determined by the size of the atoms and how many electron pairs are shared  Bond energy is the energy required to break a chemical bond and form neutral isolated atoms.

17. Homework  Page 177 #1-3  Page 209 #6, 10-11

18. 6.2 Notes Continued

19. Octet Rule in Covalent Bonds  Remember that all compounds want to attain the electron configuration of noble gases.  Hydrogen only needs 2, the rest need 8.  Regarding covalent bonds, electrons are shared between the atoms so that they attain the electron configuration of noble gases.

20. Exceptions to the Octet Rule 1) A small group of molecules has an odd number of valence electrons and cannot form an octet around each atom -Ex: NO 2 2) Fewer than eight electrons: BORON is stable with 6! - Ex: BH 3 H - B - H H O N O

21. Exceptions to the Octet Rule 3) Some central atoms have more than eight valence electrons - Referred to as an “expanded” octet - Explained by d-orbitals PCl 5 (10 e - ) SF 6 (12 e - )

22.  “Electron-dot notation”: Electrons are represented as dots located around the symbol of the element. You must put one electron on each side before you double up. Examples: Nitrogen = Hydrogen = Carbon = Drawing Valence Electrons X NH C

23. In-Class Examples  Chlorine  Neon  Magnesium  Sulfur  Silicon

24. To draw Lewis structures for covalent bonds, use the NASB method: N (Needed): Find the number of electrons needed to form full octets for all elements. For most nonmetals, they need 8. Hydrogen needs only 2. N (Needed): Find the number of electrons needed to form full octets for all elements. For most nonmetals, they need 8. Hydrogen needs only 2. A (Available): Find the number of electrons available by adding up all of the valence electrons for all elements involved. A (Available): Find the number of electrons available by adding up all of the valence electrons for all elements involved. S (Shared): Subtract the two numbers. S= N-A S (Shared): Subtract the two numbers. S= N-A B (Bond): A bond is formed with 2 electrons, so divide by 2 to find how many bonds to draw between the elements. B (Bond): A bond is formed with 2 electrons, so divide by 2 to find how many bonds to draw between the elements. Draw the molecule. Put first atom in the center. H’s are always outside. Draw in the bonds, then fill in the rest of the electrons. Draw the molecule. Put first atom in the center. H’s are always outside. Draw in the bonds, then fill in the rest of the electrons. Check to ensure all atoms have a full octet. Check to ensure all atoms have a full octet.

25. Draw the Lewis-dot-structure for the following molecules 1. HF 2. CCl 2 H 2

26. Draw the Lewis-dot-structure for the following molecules 1. H 2 O 2. CO 2

27. Types of Bonds Each bond involves the sharing of _____ _________ of electrons. Single Bonds= __ e - ’s Double Bonds= __ e - ’s Triple Bonds=__ e - ’s one pair 2 4 6

28. Resonance Structures  Occurs when more than one valid Lewis Structure can be written for a molecule or ion  Differ only in the position of electron pairs, never the atom’s positions  Actual molecule behaves as if it has one structure  Example: O 3

29. Homework  6.2 page 209 #15-19, 21, 23

31.  Bond formed between two or more ions to form an electrically neutral compound by the transfer of electrons.  Formula Unit: the simplest collection of atoms from which an ionic compound’s formula can be established. Ch 6.3 IONIC BONDS

32. Ionic Bonding How Ionic Bonding Works The negative and positively-charged ions are attracted to each other (like a magnet). ****Ionic bonding – only 2 types 1 Metal ion + 1 Nonmetal ion or 1 Metal ion + 1 Polyatomic ion

34. Ionic Bonds: One Big Greedy Thief Dog!

35. Ionic bond – electron from Na is transferred to Cl.

36. Characteristics of Ionic and Covalent Compounds CharacteristicIonic CompoundCovalent Compound Representative UnitFormula UnitMolecule Bond FormationTransfer of electronsSharing of electron pairs Type of ElementsMetals and nonmetalsNonmetals Physical StateSolidSolid, Liquid, or Gas Melting PointHighLow Solubility in WaterUsually HighHigh to Low Electrical Conductivity of Aqueous Solution Good ConductorPoor to nonconducting

37. Polyatomic Ions  A charged group of covalently bonded atoms.  They behave as one group.  If more than one is needed, written with parentheses around the ion.  NaCl 2 vs. Mg(OH) 2  They are held together by covalent bonds, but form ionic bonds with other ions.

38. Metallic bonds -Bonds found in metals -Holds metal atoms together very strongly. Ch 6.4 Bonding in Metals

39.  Good conductors at all states, shiny, very high melting points  The valence electrons of metal atoms can be modeled as a sea of electrons and the electrons can move freely.  Malleability and Ductile Metallic Bond

40. Crystalline Structure of Metals  Metal atoms are arranged in very compact and orderly patterns.  Resembles how apples and oranges are stacked in a grocery store.

41. Homework  6.3 and 6.4 pg 210 #25-26, 28, 30-31