Bonding An Introduction to Chemical Reactions

1. BondingAn Introduction to Chemical Reactions Pg. 156-215

2. Chemical Bonds • Properties of many materials can be understood in terms of their microscopic properties: • connectivity between atoms, • three dimensional shape of the molecule. • When atoms are strongly attracted to one another = chemical bond • What causes this attraction between atoms?

3. Electrostatics • Electrostatics- attraction and repulsion determines bonding between atoms and forces of attraction that can exist between molecules.

4. Coulomb’s Law q1 & q2 are charges on particles 1 & 2 D is the distance between the particles ke = constant • As distance between charges increases, the electrostatic force __________. • As the charge on the particles increases, the electrostatic force __________.

5. Questions to Consider • Where do these charges exist in an atom? • How does the organization of the atom’s electrons affect this electrostatic force? • Make a connection between reactivity of atoms in the periodic table and the organization of electrons using this concept of electrostatic force.

6. Bond Types • 1. Ionic Bonds- ________________________________________ These oppositely charged ions are attracted to each other through electrostatic forces. • 2. Covalent Bonds- • 3. Metallic Bonds-

7. Metallic Bonds “Positive ions in a sea of mobile electrons.” Delocalized Valence Electrons

8. Metallic Bonds • Form between two or more metals • Atoms of metals achieve stability by sharing their valence electrons. Delocalized valance electrons. • Metallic bonds are the attractive forces between fixed positive ions and the moving valence electrons of the metal.

9. Composition of Selected Alloys

10. Ionic Bonds • Static electricity and the clothes dryer • Static electricity is the basis for ionic bonds. • Octet Rule dictates that some substances gain electrons- __________, while others lose electrons- ___________. • Positive and negative ions are attracted to one another.

11. Ionic Bonds

12. Characteristics of Substances with Ionic Bonds • Composed of _______ • Have ________ melting points • Solids at room temperature, many soluble in water • ________________________________________________ • Tend to be ______________

13. Covalent Bonds • Formed by a shared pair of electrons between two atoms. • Molecule = Glycine- AA

14. Types of Formulas • Molecular formula- indicates the number of atoms that are in a single molecule of a compound. C6H12O6 • Empirical formula- indicates the lowest whole number ratio of atoms in a molecule. CH2O • Structural Formulas- specifies which atoms are bonded to each other in a molecule

16. Structural Formulas- Lewis Structures • Valence electrons are indicated around the symbol for the element Oxygen has 6 valence electrons Nitrogen has 5 valence electrons

17. Drawing Lewis Structures • Imagine each side (top, bottom, left, right) of the symbol of the element can hold 2 electrons for a total of 8 electrons. • Each side will hold one electron first, then will double up. • In covalent bonding the number of single electron sides (unpaired electrons) indicates the number of covalent bonds the atom must have to satisfy its octet.

18. Oxygen has 6 valence electrons. • Two unpaired electrons means that oxygen must form two bonds to satisfy its octet. • Draw the Lewis structure for the following: • Chlorine • Phosphorus • Carbon

19. Lewis Structures • Atoms share electrons to fill their octets. • A solid line indicates a shared pair of electrons. • Dots are used to indicate unshared pairs of electrons. Formation of a single covalent bond

20. Double and Triple Bonds • A unique characteristic of covalent compounds is their ability to form multiple bonds between two atoms. • Refer back to the Lewis Structures for nitrogen and oxygen. • Nitrogen needs to share three electrons • Oxygen needs to share two electrons.

21. Technique for Drawing Lewis Structures • Determine the number of valence electrons in each atom making up the molecule • Add the valence electrons and divide by two • Draw the “skeleton.” If carbon is present, place it at the center of the molecule. • Distribute the pairs of electrons around the skeleton to satisfy each atoms octet. (Remember: Hydrogen only needs two electrons to fill its octet.)

22. Practice • Draw Lewis Structures for the following compounds: • Ammonia • Ethyne- C2H2 • Carbon Dioxide • HCN

23. Exceptions to the Octet Rule • Atoms with more than an octet • SF4 • Molecules with an odd number of electrons • NO • Generally short lived, unstable molecules

24. Properties of Molecular Compounds • Composed of 2 or more ___________ • ___________ electrons in bond formation • Can be solids, liquids, or gases at room temperature. • Some are soluble in water, others are not. • Tend to be ___________________ conductive. • Generally have _________ melting points.

25. Questions to Consider for Lewis Structures • What does it mean to “share” electrons in the formation of a bond. • In your experience, is “sharing” always equal? • Pick a bond in your Lewis structure and decide if the sharing of electrons is equal or unequal. Why is it so? • How might this “sharing” affect the physical and chemical characteristics of the molecule?

26. Covalent Bonds- Are the Atoms Really “Sharing” Electrons? Chlorine Hydrogen

27. Covalent Bond Types • Polar Covalent Bonds- electrons in bond are ________________. • Nonpolar Covalent Bonds- electrons in bond ___________________________________.

28. Polar Covalent Bonds   When a bond is classified as polar covalent (H-O), the atom with the higher electronegativity has the greater attraction for the shared electrons ·As a result, a charge unbalance is produced in the molecule + by H and – by O   Dipole = charge unbalance d+ H –O d- • The “positive” and “negative” ends of the dipole are not real charges (such as positive and negative ions) because no electrons have actually been transferred between the atoms. The dipole represents only an unbalanced charge distribution along the bond.

29. Nonpolar Covalent MoleculesBrINCl HOF Elements • Diatomics- elements that can combine with themselves in a nonpolar covalent molecule to form a stable compound. • Memorize!

31. Electronegativity

32. Bond Type by Electronegativity

33. Water’s polarity allows it to pull at the ions in an ionic crystal.

36. Metallic Vs Ionic Bonding • Much easier to deform materials with metallic than with ionic bonding. Why? Ag (s) NaCl (s) • Sliding atom planes over each other (deformation) very unfavorable energetically in ionic solids! •  metals are ductile & ceramics (ionic) are brittle

37. Generally, intermolecular forces are much weaker than intramolecular forces. Intermolecular Vs Intramolecular Forces Intermolecular forces are forces ___________ molecules. Arises from interaction between dipoles. Bond Polarity Intramolecular forces _________________________________ • Intermolecular vs Intramolecular • 41 kJ to vaporize 1 mole of water (inter) • 930 kJ to break all O-H bonds in 1 mole of water (intra)

38. Types of Intermolecular Forces • Dipole-Dipole Forces • Hydrogen Bonding Forces • London Dispersion Forces Tend to be less than 15% as strong as covalent or ionic bonds. “Measure” of intermolecular force boiling point DHvap melting point DHfus DHsub

39. liquid Intermolecular Forces 1. Dipole-Dipole Forces: solid

40. or … … H H B A A A Intermolecular Forces 2. Hydrogen Bond: a special dipole-dipole interaction between the hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. A & B are N, O, or F

41. Intermolecular Forces 3. London Dispersion Forces:

42. separated Cl2 molecules London Dispersion Forces among nonpolar molecules instantaneous dipoles

43. Chemical Reactions • A process in which one or more substances are converted into new substances with different physical and chemical properties. • Reactant- a substance that enters into a chemical reaction. • Product- a substance that is produced by a chemical reaction.

44. The Reason for Reactions • During a chemical reaction, new substances are produced as existing bonds are broken, atoms are rearranged, and new bonds are formed. • Substances undergo chemical reactions with other substances _____________________

45. Chemical Equations • Describes what happens in a chemical reaction- similar to mathematic equations. • Word Equations- give the names of the reactants and the products. Calcium + oxygen yields calcium oxide • Formula Equations-chemical symbols replace the names of the reactants and products. Ca + O2 CaO

46. Law of Conservation of Mass and Balancing Chemical Equations • Matter is neither created nor destroyed during a chemical reaction. Therefore, all the atoms that were present at the start of the reaction must be present at the end of the reaction.

47. Balanced?Ca + O2 CaO • Coefficients are used in chemical equations to balance an equation. • Subscripts cannot be changed once the compound is written. Changing the subscript would change the compound! Ca + O2 CaO A coefficient of 2 is placed in front of calcium and calcium oxide to balance the equation. 2Ca + O2 2CaO

48. Steps to Balance Chemical Equations • Write the formula equation with the correct symbols and formulas. Na + Cl2 NaCl • Count the number of atoms of each element on each side of the arrow. • Balance atoms by using coefficients. 2Na + Cl2 2NaCl • Check your work by counting atoms of each element.

49. Edible Equations • 1. Gather several thin pretzel sticks and a package of M&Ms. • 2. Use the pretzels and M&Ms to make models of the following chemical reactions: 2KClO3 2KCl + 3O2 U+ 3F2  UF6 Cd + HCl  CdCl2 + H2 Cs2 + O2  CO2 + SO2 • 3. How do your models illustrate the Law of Conservation of Matter?

50. Practice • Sodium phosphate is used to cut grease. Write a balanced equation for the reaction in which iron (II) chloride reacts with sodium phosphate to produce sodium chloride and iron (II) phosphate. • Chlorine reacts with lithium bromide to produce lithium chloride and bromine.