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Chemical Bonding Chapters 8-9 (Ionic, Covalent)

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1 Chemical Bonding Chapters 8-9 (Ionic, Covalent)

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

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

6 Lithium reacts with nitrogen to form lithium nitride.
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 Lattice Energyies of Some Ionic Compounds Compound Lattice Energy Name (kJ/mol) KI -632 KF -808 KBr -671 AgCl -910 RbF -774 NaF NaI -682 LiF -1030 NaBr -732 SrCl2 -2142 NaCl -769 MgO -3795

11 Lattice Energyies of Some Ionic Compounds
Name (kJ/mol) KI -632 KF -808 KBr -671 AgCl -910 RbF -774 NaF NaI -682 LiF -1030 NaBr -732 SrCl2 -2142 NaCl -769 MgO -3795 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
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?


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
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

51 Covalent Bonding Review 2
1. Describe single, double, and triple bonds. 2. How is a pi bond different from a sigma bond? 3. Can a molecule with single bonds have a pi bond? Why or why not? 4. What affects bond strength? 5. What two things determines bond length? Describe them. 6. What is bond dissociation energy and what does it indicate? 7. What occurs when a bond forms or breaks? 8. List the properties of molecules.

52 Naming Molecules (9.2) Molecules are represented by both names and formulas. Rules for Naming Binary Molecular Compounds 1. The first element in the formula is named first, using the entire element name. 2. The second element in the formula is named using the root of the element and adding the suffix –ide. 3. Prefixes are used to indicate the number of atoms of each type that are present in the compound. -exception: 1st element never uses the prefix mono- -drop the final letter of the prefix if element name begins with a vowel.

53 Prefixes you need to know: # atoms prefix 1 mono- 2 di- 3 tri- 4 tetra- 5 penta- 6 hexa- 7 hepta- 8 octa- 9 nona- 10 deca-

54 Naming Binary Molecules-Example
Name the compound P2O5, which is used as a drying and dehydrating agent. 1st atom: P = phosphorus 2nd atom: O = oxygen = oxide There are 2 phosphorus = diphosphorus There are 5 oxygens = pentoxide (drop the –a of penta-) Put it together: diphosphorus pentoxide

55 Naming Binary Molecules Practice
Name the following molecules: 1. CCl4 2. As2O3 3. CO 4. SO2 5. NF3

56 Naming Acids (We will talk more about acids in Ch 19) There are two types of acids: 1. binary acid: contains hydrogen and one other element -when naming use the prefix hydro- plus the root of the second element with the suffix –ic, followed by the word acid. -ex: HCl H = hydro- Cl = chloride = chloric hydrochloric acid

57 Some acids are not binary, but are named according to the binary acid rules when oxygen is not present, as in HCN. H = hydro CN = cyanide = cyanic hydrocyanic acid 2. oxyacid: an acid that contains an oxyanion (oxygen containing polyatomic ion) -the name depends on the oxyanion present -the name consists of the root of the anion, a suffix, and the word acid ♦if the anion suffix is –ate, it is replaced with -ic ♦if the anion suffix is –ite, it is replaced with -ous

58 -examples: ~HNO3 NO3 = nitrate = nitric nitric acid ~HNO2 NO2 = nitrite = nitrous nitrous acid

59 Naming Acids Practice Name the following acids: 1. HBr 2. H3PO4 3. H2SO4 4. H2SO3 5. H2CO3

60 Writing Formulas Use the prefixes in the molecule’s name to determine the subscript for each atom in the compound. - phosphorus tribromide P Br 1 (no prefix) 3 (tri) PBr3 - the formula for an acid can be derived from the name as well ♦charge of the oxyanion or anion gives the number of hydrogens hydrofluoric acid = HF (1 H because fluorine has a -1 charge)

61 Writing Formulas Practice
1. oxygen difluoride 2. dinitrogen tetrasulfide 3. phosphorus pentachloride 4. iodic acid 5. phosphoric acid

62 Molecular Structure Objectives
1. Draw structural formulas. 2. Explain why resonance occurs and identify resonance structures. 3. Discuss exceptions to the octet rule.

63 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)

64 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

65 central atom and each terminal atom. H C O H
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

66 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

67 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 NBr5

68 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

69 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

70 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

71 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)

72 Structural Formulas-Practice 1
1. SO3 2. N2O 3. SF6 4. ClF3 5. SiF4 6. PO BF3 8. SO3-2

73 Molecular Structure Review 1
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. PO Write formulas for the following: a. sulfur trioxide c. chlorous acid b. hydrosulfuric acid 6. Draw structural formulas a. SO2 b. H2O c. BrCl5


75 Molecular Shape & VSEPR Objectives
1. Discuss the VSEPR bonding theory. 2. Predict the shape of and the bond angles in a molecule. 3. Define hybridization. 4. Describe how electronegativity is used to determine bond type. 5. Compare and contrast polar and nonpolar bonds.

76 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.

77 VSEPR model The repulsion between electron pairs result in fixed angles between atoms -bond angle: angle formed by any two terminal atoms and the central atom ♦lone pairs take up slightly more space than bonded pairs ♦multiple bonds have no affect on the geometry because they exist in the same region as single bonds -example: H2O See page 260 for the Molecular Geometries (Shapes)

78 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)

79 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

80 Shapes and Polarity Review 1
1. What determines many of the physical and chemical properties of molecules? 2. Describe the VSEPR model. 3. What does the repulsion between electron pairs result in? 4. Why do multiple bonds have no affect on geometry of a molecule? 5. Why do molecules with lone pairs have shorter bond angles? 6. What is electronegativity and what does it predict? 7. What is the difference between a nonpolar covalent bond and a polar covalent bond?

81 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 = 0.84 polar covalent C-Cl = 0.61 S-Se = 0.03 nonpolar covalent

82 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.

83 Let’s look at H2O and CCl4. O—H C—Cl d- d+ d+ d- 1. 24 0
Let’s look at H2O and CCl4. O—H C—Cl d- d+ d+ d 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.

84 -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

85 If water is polar, why will oil not dissolve in it
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”

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

87 7. True or False. Explain your answer if false. a
7. True or False? Explain your answer if false. a. “Orbital hybridization theory can describe both the shape and bonding of molecules.” b. “Covalent bonds differ in the way electrons are shared by the bonded atoms, depending on the kind and number of atoms.” 8. Draw Lewis structures for the following and determine if they are polar or nonpolar? Why? a. CO2 b. NH3 c. HCl


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