Chemical Bonding Chapters 8-9 (Ionic, Covalent)

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

2. Forming Chemical Bonds chemical bond: force that holds two atoms together -creates stability in the atom Bonds may form in two ways: 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.

3. 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

4. Example: Sodium reacts with chlorine to form sodium chloride. Electron Configuration Notation: Orbital Notation: Lewis Dot Notation:

5. Try this # 1: Magnesium reacts with oxygen to form magnesium oxide. Electron Configuration Notation: Orbital Notation: Lewis Dot Notation:

6. Try this # 2: Lithium reacts with nitrogen to form lithium nitride. Electron Configuration Notation: Orbital Notation: Lewis Dot Notation:

7. Ionic Bonding Review 1 1. Define chemical bond. 2. What is an ionic bond? How does it form? 3. What are two ways bonding can occur? Describe each. 4. Draw the orbital notation and Lewis dot notation showing the bonding between sodium and sulfur. (you may use noble gas notation).

8. 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.

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

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

11. 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

12. Ionic Bonding Review 1-2 (finish for HW) 1. How do positive ions form? How do negative ions form? What are each called? 2. Why do atoms bond? . 3. What determines the properties of an element? 4. What is a crystal lattice? 5. List 5 characteristics of ionic compounds. 6. What is the difference between endothermic and exothermic? Which occurs in ionic reactions? 7. What is lattice energy? 8. What does lattice energy depend on? 9. Which substance has a stronger bond: NaCl or MgO? Why?

13. Names and Formulas-Ionic Compounds A universal set of rules must be used so chemists around the world can communicate. formula unit: simplest ratio of ions represented in an ionic compound -remember that ionic compounds form a crystal lattice, consisting of many cations and anions. -the overall charge for the compound is 0 Most ionic compounds are binary, consisting of two monatomic ions. -monatomic ion: one atom ion, either positively or negatively charged

14. Remember that we determine the charge of each ion by its oxidation number. Formula Rules for Ionic Compounds 1. write the cation first, followed by the anion 2. state the charges of both ions 3. cross the number for the charge of one ion to become the subscript for the other ion. -subscripts are used to state the number of each atom in the compound

15. Example: Determine the formula for the ionic compound formed when potassium reacts with oxygen. 1. Cation = potassium = K Anion = oxygen = O 2. K+1 O-2 3. K+1 O-2 K2O1 K2O You try: Determine the formula for the ionic compound formed when aluminum reacts with chlorine.

16. Ionic Bonding Practice 2 Write the correct formula for the following pairs of atoms: 1. potassium and iodine 2. magnesium and chloride 3. aluminum and bromide 4. cesium and nitride 5. barium and sulfide

17. Ionic Bonding Review 3 1. Why do we need a universal set of rules for naming and writing formulas? 2. Define monatomic and binary. 3. What is meant by a formula unit? 4. Briefly describe the steps to writing ionic formulas. 5. Explain how we determine the charge of the cation and anion. 6. What is the purpose of subscripts. 7. Determine the formula for the ionic compound formed when lithium reacts with nitrogen.

18. Ionic Compounds with Polyatomic Ions We write formulas for ionic compounds containing polyatomic ions the same way as in binary compounds. -the cation comes first, followed by the anion -state the charges -cross over the number for the charges However: -if you have more than one polyatomic ion, place parenthesis around the polyatomic ion, with the subscript outside the parenthesis.

19. Example: Determine the formula for the ionic compound formed when beryllium reacts with cyanide. 1. Cation = beryllium = Be Anion = cyanide = CN- 2. Be+2 CN-1 3. Be+2 CN-1 Be1(CN)2 Be(CN)2 You try: Determine the formula for the ionic compound formed when ammonium reacts with iodine.

20. Ionic Bonding Practice 3 Write the correct formula for the following pairs of atoms: 1. ammonium and oxygen 2. lithium and nitrate 3. aluminum and hydroxide 4. ammonium and phosphate 5. strontium and acetate

21. Ionic Bonding Practice 4 Write the correct formula for the following pairs of atoms: 1. aluminum and carbon 2. ammonium and carbonate 3. calcium and oxygen 4. aluminum and chromate 5. sodium and phosphate 6. potassium and hydrogen sulfate 7. magnesium and phosphorus

22. Ionic Bonding Review 4 1. What is meant by a formula unit? 2. Explain how we determine the charge of the cation and anion. 3. What is the purpose of subscripts. 4. Describe what a polyatomic ion is? 5. When do we use parenthesis for writing ionic compounds with polyatomic ions? 6. Determine the formula for the ionic compound formed when lead reacts with sulfur. 7. Determine the formula for the ionic compound formed when magnesium reacts with phosphate.

23. Naming Ionic Compounds The names of ionic compounds include the ions of which they are composed. 1. The element whose symbol appears first in the formula also appears first in the name. -this is always the positively charged ion, or metal 2. The name of the second ion follows, with its ending changed to –ide for single atom ions. Ex: What is the name of MgCl2? magnesium chloride

24. Ionic Compounds Practice 5 Write the formula and the name. 1. Na2S 2. Ga2S3 3. CaSe 4. LiF

25. Naming with Polyatomic Ions You follow the same rules when naming polyatomic ions as when you have binary ionic compounds, however: -you do not change the ending of the polyatomic ions, even when they are the second atom. Example: Al2(SO4)3 aluminum (III) sulfate Rule: You must state the charge of all metals not included in groups 1 and 2 because many have multiple charges.

26. Rules for Transition Metals *According to the previous rules, FeO and Fe2O3 would both be named iron oxide,even though they are not the same compound* Since many transition metals can have more than one charge, the name must show this. This is done using roman numerals. -FeO is named iron (II) oxide because Fe has a +2 charge -Fe2O3 is named iron (III) oxide because Fe has a +3 charge *The roman numeral states the charge of the metal*

27. Q: How do I know the iron in FeO has a +2 charge? A: The oxide ion has a –2 charge, so the Fe must have a +2 charge so the compound is overall neutral. Q: How do I know the iron in Fe2O3 has a +3 charge? A: There are three oxide ions with a –2 charge: (3 ions)(-2 charge/ion) = a total of –6 charge Since the overall charge must be neutral, the iron must have a total charge of +6. Therefore: (2 ions)(x charge/ion) = +6 x = +3

28. Ionic Compounds Practice 6 Write the formula given & the name of each compound. 1. FeCl3 2. Zn3P2 3. CuS 4. AuF 5. CuC2H3O2 6. AgHCO3 7. ZnSO4 8. Pb(CO3)2

29. Ionic Compounds Practice 7 Name the following compounds: 1. NaBr 2. CaCl2 3. KOH 4. Cu(NO3)2 5. Ag2CrO4 6. PbNO2 7. AlCl3

30. Ionic Bonding Review 5 1. Describe what a polyatomic ion is? 2. What is the relationship between lattice energy and the strength of ionic bonds? 3. What is the ending of the second atom changed to when naming ionic compounds? 4. Write the name for (NH4)3P 5. Write the name for AlS. 6. Determine the formula for the ionic compound formed when magnesium reacts with phosphate.

31. Metallic Bonds Metallic bonds are similar to ionic bonds because they often form lattices in the solid state. -eight to twelve metal atoms surround another, central metal atom Instead of sharing electrons or losing electrons, the outer orbitals overlap. -electron sea model: all metal atoms in a metallic solid contribute their valence electrons to form a ‘sea’ of electrons around the metal atoms. -valence electrons are free to move from atom to atom (delocalized electrons), forming metallic cations

33. metallic bond: attraction of a metallic cation for the delocalized electrons that surround it This bonding contributes to the unique properties of metals: 1. generally have high melting and boiling points, with especially high boiling points -due to the amount of energy needed to separate the electrons from the group of cations 2. malleable (hammered into sheets)and 3. ductile (drawn into wire) -mobile electrons can easily be pulled and pushed past each other

34. 4. durable -though electrons move freely, they are strongly attracted to the metal cations and are not easily removed from the metal

35. 5. good conductors -free movement of the delocalized electrons, allowing heat and electricity to move from one place to another very quickly 6. luster -interaction between light and delocalized electrons

36. As the number of delocalized electrons increases, as in transition metals (d electrons), the hardness and strength also increases. -alkali and alkaline earth metals are soft (s valence electrons only) It is easy to combine 2 or more different metals to make a metallic crystal -alloy: mixture of elements with metallic properties -the properties of alloys differ from those of the individual elements that make it up

37. Metallic Bonding Bellringer • What is a metallic bond? • What is an alloy? • Describe the electron sea model. • What occurs with orbitals in metals? • How is metallic bonding similar to ionic bonding? • What are delocalized electrons? • What contributes to a metal’s high boiling point, malleability, ductility and conductivity? • List the other 2 properties of metals. • What happens to strength and hardness as you decrease the number of delocalized electrons?

38. TEST

39. 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

40. 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)

41. 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

42. Single Bond Practice 1. PH3 2. H2S 3. HCl 4. SCl2 5. SiH4

43. Covalent Bonding Review 1 • Describe a covalent bond. • What types of atoms do covalent bonds form between? • Describe single and double bonds. • What do we mean by sigma bonds? • What do we call covalent compounds?

44. 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

45. Multiple Bonds Practice 1. CO2 2. CH2O 3. C2H2

46. 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

47. 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)

48. 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

49. 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

50. 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