Chemical Bonding Chapters 8-9 (Ionic, Covalent)

1. Chemical BondingChapters 8-9(Ionic, Covalent) Chemistry

2. Answers to Explain (Analysis), Part 1: 1. Examine the data collected for melting point. What conclusions can you draw about the melting point of these chemicals? - The chemicals that took longer to melt have a higher melting point than those that melted more quickly. 2. Which substances have higher melting points? Which have lower melting points? What does this indicate about the bonds in the substances?   -Substances with ionic bonds had higher melting points and those with covalent bonds had lower melting points. A lower melting point indicates weaker bonds that will be more easily broken. A higher melting point indicates stronger bonds. 3. Summarize the solubility of the substances in the Explore Activity. -All substances are soluble in water, except benzoic acid.

3. 4. How is solubility associated with the type of bond present? -All our ionic substances are soluble in water. However, not all ionic compounds are soluble in water. One of our covalent molecules are soluble (dextrose) and one is not (benzoic acid).  5. What does the solubility of the different substances indicate about the type of bond present?   -Substances with weaker intermolecular forces are more soluble in water. Those with stronger intermolecular forces are less soluble in water. 6. How is conductivity related to the type of bond present? -Ionic substances, when dissolved in water, conduct electricity. Covalent substances do not conduct electricity when dissolved in water.

4. 7. Why do substances with certain types of bonds conduct electricity well, while some substances are not good conductors? -Ionic compounds conduct electricity well because they possess ions, which allows electrons to flow from atom to atom. This occurs only when they are melted or dissolved in water. 8. Is there a significant difference in appearance between the substances with covalent bonds and those with ionic bonds? What properties did you notice you could not see with the naked eye? -Both ionic and covalent compounds appear to be white solids. Under the hand lens, however, the covalent substances are smaller in particle size than the particles in ionic compounds. Ionic particles also have a more geometric, crystalline shape while covalent particles vary in shape.

5. 9. Imagine looking at the substances under a microscope. What do you think the substances might look like on a microscopic level? -(Answers will vary.) Lewis Structures are used to show bonding in molecules and ionic compounds. -dots represent valence electrons -in ionic bonding, the charge of each ion must be shown -in covalent bonding, bonded electrons is shown by lines

6. Lewis Structures: Ionic Bonds Arrows represent transfer of electrons from the metal to the nonmetal. -the charge of each atom must be shown Example: CaS

7. Lewis Structures: Single Covalent Bonds When we show bonding, shared electron pairs can be shown by either a pair of dots or a single line. -Lewis Structures are used to show how bonding electrons are arranged in molecules -example: NH3 -sigma bond (s): single covalent bond formed when an electron pair is shared by the direct overlap of orbitals ♦can occur between s & s, s & p , or p & p orbitals

8. Explain(Analysis), Part 2: Draw Lewis structures for the following ionic compounds. 1. NaCl 2. MgO 3. LiF Draw Lewis structures for the following covalent compounds. 1. H2O 2. CO2 3. NH4

9. Lewis Structure: Multiple Bonds A multiple bond forms when two atoms share more than 2 electrons. -double bond: 4 electrons shared ( 2 pairs) ♦ O2 -triple bond: 6 electrons shared (3 pairs) ♦ N2 Some molecules have both single and multiple bonds. ♦HCN pi bond (p): forms when parallel orbitals overlap to share electrons -only occurs with multiple bonds because the first overlap is always a sigma bond

10. Lewis Structures Practice 1 Show the formation of the ionic compound for the following pairs of elements 1. strontium and fluorine 2. aluminum and oxygen 3. cesium and phosphorus 4. lead and chlorine 5. potassium and iodine 6. magnesium and chloride 7. aluminum and bromide 8. cesium and nitride 9. barium and sulfide

11. Lewis Structure Practice 2 1. PH3 2. H2S 3. HCl 4. SCl2 5. SiH4 6. CO2 7. CH2O 8. C2H2

12. Forming Chemical Bonds chemical bond: force that holds two atoms together -creates stability in the atom Two types of bonds: 1. Attraction between a positive nucleus and negative electrons (covalent bonding) 2. Attraction between a positive ion and a negative ion (ionic bonding) Remember: It is the valence electrons that are involved in this bonding.

13. Formation of Ionic Bonds ionic bond: electrostatic force that holds oppositely charged particles together -called ionic compounds -forms between metals and nonmetals ◊metals lose electrons, forms a cation ~cation: positive ion from loss of electrons ◊nonmetals gain electrons, forms an anion ~anion: negative ion formed from gain of electrons -most are binary, which means they contain 2 different elements, such as MgO, Al2O3

14. Properties of Ionic Compounds It is the chemical bonds between atoms that determines many of the physical properties of the compound. -alternating positive and negative ions form an ionic crystal -the ratio of positive to negative ions is determined by the number of electrons transferred -strong attraction results in a crystal lattice, a 3-D arrangement of atoms.

15. Other characteristics include: -high melting and boiling points -very hard and rigid -brittle -electrolyte when dissolved in water (aqueous solution) During chemical reactions, energy is either absorbed (endothermic) or released (exothermic) -the formation of ionic bonds is always exothermic

16. lattice energy: energy required to separate one mole of ions of an ionic compound -the more negative the lattice energy, the stronger the bond

17. Depends on: 1. smaller ions -more negative value because the attraction is stronger between the nucleus and valence electrons 2. larger the positive/negative charge, the more negative the lattice energy because the attraction is stronger when more electrons are lost/gained

18. Ionic Bonding Review 1 (finish for HW) 1. Draw the Lewis dot notation showing the bonding between beryllium and chlorine. 2. What determines the properties of an element? 3. What is a crystal lattice? 4. List 5 characteristics of ionic compounds. 5. What is the difference between endothermic and exothermic? Which occurs in ionic reactions? 6. What is lattice energy? 7. What does lattice energy depend on? 8. Which substance has a stronger bond: NaCl or NaBr? Why?

19. Covalent Bonds (9.1) Remember that atoms bond to increase stability, which occurs when an atom gets a full outer shell of electrons. -in ionic bonding, one atom loses electrons (metal) and another gains electrons (nonmetal) to form oppositely charged ions with a full outer shell However, sometimes there is not a transfer of electrons, but a sharing of electrons. -covalent bond: attractive force between atoms due to the sharing of valence electrons

20. Covalent bonds can form between: -2 or more nonmetal atoms -metalloids (especially the ones to the right of the metalloid line) and nonmetals molecule: when two or more atoms bond covalently Covalent bonds can have either single bonds or multiple bonds. -single bonds: 2 shared electrons (1 pair) -multiple bonds: 4 or 6 electrons shared (2 pair= double or 3 pair = triple)

21. Properties of Molecules (Covalent Compounds) 1. low melting and boiling points. 2. many vaporize readily at room temperature 3. relatively soft solids (but not all, some are gases/liq.) 4. can form weak crystal lattices 5. do not conduct electricity when dissolved in water

22. Properties of Molecules These properties are due as a result of differences in attractive forces -attraction between atoms within a molecules is strong -attraction between different molecules is weak ~called intermolecular forces or van der Walls forces Types of Intermolecular Forces (van der Walls forces) • dispersion force (induced dipole) • dipole-dipole force • hydrogen bonding

23. Properties of Molecules dispersion force (induced dipole) -occurs between nonpolar molecules -very weak dipole-dipole force -occurs between polar molecules -the more polar the molecule, the stronger the force hydrogen bonding -strong intermolecular force between the hydrogen end of one dipole and a fluorine, oxygen or nitrogen atom on another molecule’s dipole

24. Strength of Covalent Bonds All bonds can be broken, though some more easily than others. -due to the strength of the bond What affects bond strength? bond length: distance that separates the bonded nuclei -determined by the size of the atoms and how many electron pairs are shared ♦larger the atom, the longer the bond length, the weaker the bond ♦more shared electrons gives a shorter, stronger bond

25. When a bond forms or breaks, an energy change occurs. -bond formation: energy released (exergonic) -bond breaking: energy absorbed (endergonic) bond dissociation energy: amount of energy required to break a specific covalent bond -always a positive number -indicates the strength of a covalent bond larger the bond dissociation energy, stronger the bond (see p 246 for examples)

26. Covalent Bonding Review • Describe a covalent bond. • What types of atoms do covalent bonds form between? • Describe single, double and triple bonds. • What do we mean by sigma and pi bonds? • What do we call covalent compounds? • What affects bond strength? • Describe the two things that determine bond length. • What does bond dissociation energy indicate? • What occurs when a bond forms or breaks?

27. Electronegativity and Polarity Remember that atoms have different attractions for electrons (electronegativity). -electronegativity increases left to right and decreases down a period The character and type of bond can be predicted using the difference in electronegativities between bonded atoms. -pure covalent bond: electronegativity difference = 0 (usually occurs between identical atoms, H2)

28. Most atoms do not have equal sharing of electrons, producing a purely covalent bond. -polar covalent bond: unequal sharing of electrons ♦the larger the electronegativity difference, the more ionic the bond character -ionic bonds form when the electronegativity difference is > 1.7 and nonpolar covalent bonds form when the difference is < 0.5 -the cutoff between polar covalent, nonpolar, and ionic is sometimes inconsistent with experimental data

29. Electronegativity Practice Remember: bonding is not clearly ionic or covalent, with ionic character increasing as the difference in electronegativity increases. Decide if the following pairs of atoms are polar covalent, nonpolar covalent or ionic. • N-H 3.04-2.20 = 0.84 polar covalent • C-Cl 2.55-3.16 = 0.61 polar covalent • S-Se 2.58-2.55 = 0.03 nonpolar covalent

30. When a polar bond forms the shared electrons are pulled more strongly toward one atom. -this creates partial charges at opposite ends of the molecule, which is called a dipole ♦ d- indicates a partial negative d+ indicates a partial positive Polar molecule or not? A molecule can have individual polar bonds, but make a nonpolar molecule. How? We look at the shape of the molecule.

31. Let’s look at H2O and CCl4. O—H C—Cl d- d+ d+ d- 1.24 0.61 both O-H and C-Cl have polar covalent bonds One molecule is polar and the other is nonpolar? How do we know? We look at the shape of the molecule and the terminal atoms.

32. -symmetric molecules like CCl4 are nonpolar because the polar bonds cancel each other out. CCl4 -asymmetric molecules like H2O are polar because the polar bonds do not cancel each other out. H2O

33. If water is polar, why will oil not dissolve in it? Oil must be nonpolar because A substance is only soluble (dissolvable) when combined with a like molecule. “Like Dissolves Like” hydrophobic- “fear of water” hydrophilic- “likes water”

34. Polarity Review 1. What is electronegativity and what does it predict? 2. What is the difference between a nonpolar covalent bond and a polar covalent bond? 3. What is a dipole and what indicates them? 4. Describe the electronegativity trend both across a period and down a group. 5. Are the following bonds polar or nonpolar covalent? a. H-Br b. C-O c. S-C 6. Describe the relationship between polarity and solubility. 7. What do we mean by symmetric and asymmetric?

35. Final Bonding Questions: • Draw a table comparing the properties of ionic and covalent bonds. -leave room to add more properties (we will discuss the table and add more to it) • What is a general definition of a bond? • What are the two types of bonds? Describe each. • What is the octet rule? • What do we mean by polar or nonpolar? • What is electronegativity? How do we use this in bonding? • What are intermolecular forces?

36. 1. 2. A bond is a force holding two atoms together to create stability in an atom 3. An ionic bond is an attraction of oppositely charged ions due to a transfer of electrons from a metal atom to a nonmetal atom. A covalent bond is the sharing of electrons between nonmetals or nonmetals and some metalloids.

37. The octet rule states that atoms are stable if they have a full valence shell of electrons. For most atoms, the number is 8, but the period 1 elements are stable with 2. • A covalent bond is polar if there is an unequal sharing of electrons due to the electronegativity difference between the atoms. It is nonpolar if there is an equal sharing of electrons. • Electronegativity is the attraction an atom has for electrons. The more electronegative the atom, the stronger the attraction. We use electronegativity to determine the polarity of molecules. • Intermolecular forces are the force that holds atoms together. They can be weak, allowing atoms to be pulled apart easily, or strong.

38. TEST #1

39. Molecular Structures (9.3) structural formula: uses letter symbols and bonds to show relative positions of atoms -one of the most useful -can be predicted for many molecules by drawing Lewis structures -H is always an end (terminal) atom, never a central atom -less electronegative atom is the central atom (nm or metalloid closest to the left of the PT-usually)

40. Structural Formulas-Example CH2O 1. Predict the location of the atoms C is least electronegative & farthest to left on PT, therefore it is the central atom 2. Find the total number of electrons available for bonding. 1 C-4, 2 H-2, 1 O-6 for a total of 12 valence e- 3. Determine the number of bonding pairs 12 valence e- / 2 = 6 electron pairs

41. 4. Place one bonding pair (single bond) between the central atom and each terminal atom. H C O H 5. Subtract the number of pairs you used in step 4 from the number of bonding pairs determined in step 3. 6 – 3 used = 3 e- pairs left

42. 5. Subtract the number of pairs you used in step 4 from the number of bonding pairs determined in step 3. -take the remaining electron pairs and place electron pairs around the terminal atoms to satisfy the octet rule H C O H

43. 6. If the central atom is not surrounded by 4 electron pairs, it does not have an octet -convert one or two of the lone pairs on a terminal atom to a double or triple bond between that terminal atom and the central atom H C O H Practice: 1. CH3Cl 2. NBr5

44. Structural Formulas-Polyatomic Ions Writing structural formulas for polyatomic ions is the same with one exception: -the total number of electrons may differ due to the negative and positive charge. ♦negative charge, more electrons are present SO4-2 add two electrons ♦positive charge, less electrons are present NH4+1 subtract one electron

45. Resonance Structures Let’s look at CO3-2. -when one or more valid Lewis structure can be written for a molecule, resonance occurs -let’s look at another resonance molecule/ion: NO3-1 -each molecule/ion that undergoes resonance behaves as if it only has one Lewis structure

46. Exceptions to the Octet Rule Some molecules do not obey the octet rule. Three reasons exist: 1. when a small group of molecules have an odd number of valence electrons: -NO2 for a total of 17 valance electrons-one unpaired electron on N

47. 2. Some form with fewer than eight, though this is relatively rare: -B in BH3 is stable with six because it only has 3 valence electrons. 3. When the central atom has more than 8 electrons, which is referred to as an expanded octet. -can occur in elements that are found in period three or higher elements (because of the d orbitals). -P in PCl5 (1 s orbital, 3 p orbitals, and 1 d orbital)

48. Structural Formulas Practice 1. SO3 2. N2O 3. SF6 4. ClF3 5. SiF4 6. PO4-3 7. BF3 8. SO3-2

49. Molecular Structure Review 1. What is a structural formula? 2. Describe resonance. 3. List three reasons for exceptions to the octet rule. 4. Name the following: a. BH3 b. SO2 c. PO4-3 5. Write formulas for the following: a. sulfur trioxide c. chlorous acid b. hydrosulfuric acid 6. Draw structural formulas a. SO2 b. H2O c. BrCl5

50. Molecular Shape Many of the physical and chemical properties of molecules is determined by the shape of the molecule. -the shape of molecules determines if two or more molecules can get close enough for a reaction to occur. VSEPR (Valence Shell Electron Pair Repulsion) model: atoms in a molecule are arranged so that the pairs of electrons (bonded and lone) minimize repulsion.