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Chapter 7– Ionic Compounds & Metals

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1 Chapter 7– Ionic Compounds & Metals
7.0 Bonding Overview 7.1 Ion Formation 7.2 Ionic Bonds and Ionic Compounds 7.3 Names & Formulas for Ionic Compounds 7.4 Metallic Bonds & Properties of Metals

2 Section 7.0 Bonding Overview
There are 3 basic types of chemical bonding which vary in how the valence electrons are used. Electronegativity difference and the type of element (metal or nonmetal) are the key parameters in determining the type of bonding that will occur. Name the three basic types of chemical bonding. Distinguish between metals and nonmetals using a periodic table Explain the “octet rule”, the role it plays in chemical bonding and its relationship to noble gas electron configurations

3 Section 7.0 Bonding Overview
Describe the role of electronegativity difference in determining the type of bonding between two elements Distinguish between an ionic and a covalent (molecular) compound and describe the basic difference in bond formation between these two types of compounds

4 Section 7.0 Bonding Overview
Key Concepts A chemical bond is the force that holds two atoms together. There are three main types of chemical bonds – ionic, covalent and metallic. All involve valence electrons in some way. Both the electronegativity difference and the category of an element (metal or nonmetal) determine the type of bond that will form. Ionic bonds are formed between a metal and nonmetal with a large electronegativity difference (> 1.7) Covalent bonds are generally formed between nonmetals or between a metal and a nonmetal that have a small (< 1.7) electronegativity difference. Metallic bonds are formed between metals and other metals.

5 Electron Categories The electrons responsible for the chemical properties of atoms are those in the outer energy level Valence electrons - The s and p electrons in the outer energy level highest occupied energy level Core electrons - those in energy levels below the valence electrons

6 Chemical Bonds Bond is force holding two atoms together
When describing bond formation, focus is on valence electrons Elements tend to lose or gain electrons to achieve an octet of electrons Extra stability associated with noble gas configuration (filled outmost energy level) For low AN elements, have [He], not octet

7 Chemical Bonds & Valence Electrons (VEs)
Ionic Bonding Transfer of VEs from one atom to another Results in charged ions with opposite sign that attract each other (e.g. Na+ Cl- ) Covalent Bonding Sharing of VEs between atoms Metallic Bonding VEs become part of “sea” of electrons

8 Types of Chemical Bonds
Ionic Formed between metal and nonmetal Na (metal) Cl (nonmetal) NaCl sodium chloride Covalent Generally formed between nonmetals C (nonmetal) O (nonmetal) CO2 carbon dioxide Metallic Formed between metals (same or similar) Fe iron

9 Ionic vs Covalent Bonding
Key parameter in distinguishing types is the electronegativity

10 Electronegativity Relative ability to attract electrons in a chemical bond Max value F Min value Fr Elements with high EN tend to form anions Noble gases not tabulated Very few compounds to get info from

11 Electronegativity Ranges (values slightly different than in book)
Below 1.0 2.0 – 2.4 1.0 – 1.4 2.5 – 2.9 1.5 – 1.9 3.0 – 4.0

12 Fig. 8.20 (p. 265) Electronegativity

13 Electronegativity & Bonding
[see p. 266 in section 8.1] Difference in electronegativity (EN) between the atoms involved in bond formation determines type of bond = EN(atom 1) – EN (atom 2) If difference large, electron transferred  ionic bond If difference small or zero, electron shared  covalent bond

14 EN Difference & Bond Character
1.0 2.0 3.0 Electronegativity Difference Ionic Bonds % Ionic Character Covalent Bonds

15 Electronegativity (EN) & Bonding Type
Large EN difference (> 1.7) usually occurs between 2 elements when one is a metal and the other a nonmetal Group 1 and 2 metals and highly electronegative nonmetals form compounds with high ionic character (>50% or DEN > 1.7) Polar covalent bonds involve unequally shared valence electrons (but not so unequal that ions can form) Most bonds have some mix of ionic & covalent character

16 Electronegativity Difference and Bond Character (Table 8.7, p. 266)

17 Ionic Bonding Subjects for Sections 7.1 and 7.2 are formation of ions and formation of ionic bonds

18 Chapter 7– Ionic Compounds & Metals
7.0 Bonding Overview 7.1 Ion Formation 7.2 Ionic Bonds and Ionic Compounds 7.3 Names & Formulas for Ionic Compounds 7.4 Metallic Bonds & Properties of Metals

19 Section 7.1 Ion Formation Ions are formed when atoms gain or lose valence electrons to achieve a stable octet electron configuration. Define a chemical bond. Describe the formation of positive and negative ions from the elements. Describe the size change that occurs when an atom becomes an anion or a cation. Relate ion formation to electron configuration. Determine the oxidation state of metals and nonmetals based on their position in the periodic table..

20 Section 7.1 Ion Formation Determine the electron configuration of an ion Explain why transition metals can have multiple oxidation states. Name the transition metals that do not have multiple oxidation states, list the oxidation number associated with each, and explain using an argument involving the ion’s electron configuration why this single oxidation state is highly preferred.

21 Section 7.1 Ion Formation Key Concepts
A chemical bond is the force that holds two atoms together. Some atoms form ions to gain stability. This stable configuration involves a complete outer energy level, usually consisting of eight valence electrons. Ions are formed by the loss or gain of valence electrons. The number of protons remains unchanged during ion formation. Transition metals can use electrons in d orbitals as valence electrons to attain multiple oxidation states. The electron configuration of most ions is obtained by removing electrons in reverse order of highest n from the atom’s electrons configuration.

22 Bonding Atom will try to form octet by gaining or losing valence electrons

23 Bonding Metals are reactive because they lose valence electrons easily

24 Na + ionization energy  Na+ + electron
Formation of Cations Positive ions – called cations Ca+ions – write “t” as a plus sign Energy equal to ionization energy (IE) must be supplied to remove electron Expressed in kJ/mol (kilojoules per mole) Na + ionization energy  Na+ + electron IE = 498 kJ/mol

25 Atom vs Cation Radius Ionization of sodium Na  Na+ + e-
Sodium atom Sodium cation [Ne]3s [Ne]

26 Formation of Cations In most cases, lose enough electrons to achieve noble gas configuration

27 Formation of Cations For group 1 and 2 metals (and H), cation charge = group number Aluminum always has +3 charge (Above metals lose all valence electrons) Na+1, Mg+2, Al+3 First two rows of table 7.7, page 218 lists group 1 and 2 cations

28 Formation of Cations For transition metals
Generally have ns2 configuration, so generally can form +2 cation Also have (n-1)dx configuration, so some d electrons also lost Books says that “rule of thumb” – generally can form +2 or +3 Not very good rule – wide range from +1 to +7 with +8 possible (osmium, ruthenium)

29 Transition Metal Oxidation States (uses old group designations)
Common Less Common

30 Transition Metal Oxidation States

31 Special Case Transition Metals
Several transition metals only commonly form a single type of ion: Sc3+ (many sources don’t list) Zn2+ Ag1+ Cd2+ (some sources don’t list) Roman numerals are not used in compound names involving these ions

32 Forming Transition Metal Cations
For atom, following standard Aufbau order for transition metal means ns sublevel fills before (n - 1)d sublevel Example: Ti [Ar]3d24s2 (filling order 4s then 3d) To get electron configuration of ion, remove valence electrons first Ti2+ [Ar]3d2

33 Special Case Transition Metals
Pseudo-noble gas configuration Groups , periods 4 - 6 Full sublevels have extra stability s2, p6, d10 – all with same When an ion can be formed with this configuration it is especially stable and become the preferred (or only) ion formed by this metal

34 Special Case Transition Metals
Zn = [Ne]3s23p63d104s2 Zn+2 = [Ne]3s23p63d10 Compare to Ar = [Ne]3s23p6 Consequence: Zn+2 only Zn ion formed

35 Zn+2 – Pseudo-Noble Gas

36 Special Case Transition Metals
Pseudo-noble gas configuration Ag [Kr]4d105s1 Same situation as Cu (not 5s2) Ag1+ [Kr]4d10 = [Ar] 3d104s24p64d10 Compare to Kr = [Ar] 3d104s24p6 Consequence: Ag1+ only Ag ion formed Reasoning for Ag1+ applies to Cd2+ Cd [Kr]4d105s2 = [Ar]3d104s24p64d105s2 Cd2+ = [Ar]3d104s24p64d10

37 Special Case Transition Metals
Noble gas configuration Sc [Ar]3d14s2 Sc3+ = [Ar] = Only common oxidation state for Sc (+3)

38 Formation of Anions Applies to nonmetals on upper RHS of periodic table Negative ions called anions aNions – N for negative Some nonmetals can gain or lose electrons to complete an octet

39 Formation of Anions Gain enough electrons to form octet
Halogen anions: –1 charge Oxygen (O), Sulfur (S): -2 charge Nitrogen (N), Phosphorus (P): -3 charge Carbon (C): - 4 charge Charge = group # - 18 (groups 14 to 18)

40 Naming Monatomic Anions
For monatomic anions, name becomes element root + “ide” ending Chlorine (Cl)  Chloride (Cl-) Oxygen (O)  Oxide (O2-) Sulfur (S)  Sulfide (S2-) Nitrogen  Nitride (N3-) Phosphorus  Phosphide (P3-) Carbon  Carbide (C4-)

41 Atom vs Anion Radius Formation of chlorine ion Cl + e-  Cl-
Chlorine atom Chloride anion [Ne]3s23p [Ne]3s23p6 or [Ar]

42 X Electron Dot diagrams Way of keeping track of valence electrons
How to write them? Write the symbol Put one dot for each valence electron Don’t pair up until you have to (Hund’s rule) X

43 Electron Dot diagram for Nitrogen
5 valence electrons First write chemical symbol Add 1 electron at a time to each side Until they are forced to pair up. N

44 Electron Dots For Cations
Metals generally have few valence electrons These will come off Forming positive ions Ca 2+

45 Electron Dots For Anions
Nonmetals have many valence electrons (usually 5 or more) Gain electrons to fill outer energy level P P3-

46 Chapter 7– Ionic Compounds & Metals
7.0 Bonding Overview 7.1 Ion Formation 7.2 Ionic Bonds and Ionic Compounds 7.3 Names & Formulas for Ionic Compounds 7.4 Metallic Bonds & Properties of Metals

47 Section 7.2 Ionic Bonds and Ionic Compounds
Oppositely charged ions attract each other, forming electrically neutral ionic compounds. Describe the formation of ionic bonds using multiple methods including the use of electron configurations, orbital diagrams and dot diagrams. Describe the structure of ionic compounds. Generalize about the strength of ionic bonds based on the physical properties of ionic compounds. Explain why ionic compounds are brittle.

48 Section 7.2 Ionic Bonds and Ionic Compounds
Identify the conditions under which ionic compounds do or don’t conduct electricity and explain why. Categorize ionic bond formation (starting with the component ions) as exothermic or endothermic. Describe lattice energy and the sign convention used to report it. Relate the strength of ionic bonds and the lattice energy to the size of the ions and the charge on the ions using reasoning based on Coulomb’s law.

49 Section 7.2 Ionic Bonds and Ionic Compounds
Key Concepts Ionic compounds contain ionic bonds formed by the attraction of oppositely charged ions. Ions in an ionic compound are arranged in a repeating pattern known as a crystal lattice. Ionic compound properties are related to ionic bond strength. Ionic compounds are electrolytes; they conduct an electric current in the liquid phase and in aqueous solution. Lattice energy is the energy needed to remove 1 mol of ions from its crystal lattice.

50 Ionic Bonding Bond formed through transfer of electrons to form anion and cation In most cases, electrons transferred to achieve noble gas configuration in each ion No net loss or gain of electrons Total number lost = total number gained NaCl Sodium loses 1, chlorine gains 1 AlCl3 Al loses 3, three Cl each gain 1

51 Ionic Bonding Elements start out electrically neutral (no charge) and ionic compound must also be neutral

52 Ionic Compounds Ionic compounds called salts
Salts involving oxygen called oxides CuO Copper (II) Oxide Simplest ratio called formula unit (for ionic compounds, there are no molecules) NaCl (1:1 ratio) Al2O3 (2:3 ratio) Be4Al2Si6O18 formula unit for beryl Simple binary compounds = 2 different elements MgO, KCl, FeBr2, Al2O3

53 NaCl Formation Electron Configuration Picture
Transferred Electron

54 NaCl Formation Orbital Notation Picture
Transferred Electron Na Cl Na+ Cl-

55 NaCl Formation Electron Dot Picture

56 NaCl Formation Electron Dot Picture
Note that after electron transfer, the atoms are charged

57 Sodium Chloride Formation
No single isolated unit of + and – charge (no molecule as such) 11 e- 10 e- Electron loss Sodium Ion Neutral Na Atom 18 e- 17 e- Electron gain Neutral Cl Atom Chloride Ion

58 NaCl Crystalline structure

59 Ionic Bonding All the electrons must be accounted for! Ca P

60 Ionic Bonding Ca P

61 Ionic Bonding Ca2+ P 2-

62 Ionic Bonding P 2- Ca2+ Ca

63 Ionic Bonding Ca2+ P 3- Ca1+

64 Ionic Bonding Ca2+ P 3- Ca1+ P

65 Ionic Bonding Ca2+ P 3- Ca2+ P 1-

66 Ionic Bonding Ca Ca2+ P 3- Ca2+ P 1-

67 Ionic Bonding Ca Ca2+ P 3- Ca2+ P 1-

68 Ionic Bonding Ca2+ Ca2+ P 3- Ca2+ P 3-

69 Ionic Bonding – Formula Unit
Formula Unit for Calcium Phosphide = Ca3P2 Simplest ratio of ions in compound is called the formula unit

70 Properties of Ionic Compounds
Crystalline structure, usually solids A regular repeating 3D arrangement of ions in the solid – crystal lattice Crystals vary in shape due to variation in relative number and sizes of ions No single isolated unit of + and – charge (no molecule as such)

71 Crystalline structure

72 Ionic Bonding Anions and cations are held together by force between opposite charges Strong electrostatic force Magnitude of force (F) given by Coulomb’s Law F proportional to q+q- /r2 q+ = magnitude of positive (cation) charge q- = magnitude of negative (anion) charge r = distance between charge centers

73 Force vs distance for 1/r2

74 Coulomb’s Law & Ionic Bonds
F  q+q-/r2 The larger the ionic charges (q+, q-), the stronger the force between them The closer together the ions are, the stronger the force Smaller ions can get closer together than larger ions

75 Properties of Ionic Compounds
Ions are strongly bonded together Structure is hard, rigid, brittle High melting and boiling points Strength of bond depends on relative size and charge of ions

76 Ionic Solids – Melting, Boiling Points
Compound MP ( C) BP ( C) NaI 660 1304 KBr 734 1435 NaBr 747 1390 CaCl2 782 >1600 CaI2 784 1100 NaCl 801 1413 MgO 2852 3600

77 Ionic solids are brittle
+ -

78 Ionic solids are brittle
Strong Repulsion breaks crystal apart. + - + - + - + -

79 Conductivity of Ionic Solids
Substance that conducts electricity is allowing charges to move In a solid, ions locked in place Ionic solids are insulators When melted, ions can move around Melted ionic compounds conduct NaCl: must get to about 800 ºC

80 Conductivity of Ionic Solids
Dissolved in water (aqueous) they conduct Solution called electrolyte Each individual ion, surrounded by H2O molecules, free to move about and carry charge from one place to another

81 Conductivity of Ionic Solids
Conduct when dissolved in water (electrolyte)

82 Energy and the Ionic Bond
Energy absorbed or released during a chemical reaction Released – Exothermic Absorbed – Endothermic Formation of ionic compounds from cations & anions always exothermic

83 Energy and the Ionic Bond
Ions sit in crystal lattice Takes energy to separate them into individual ions Energy required is called lattice energy Like ionization energy, lattice energy expressed in kJ/mol (kilojoules/mole) The more negative the value, the stronger the forces of attraction Can correlate LE with melting point

84 Lattice Energy (LE) Related to size and charge of ions
Smaller ions generally have more negative values (stronger forces) Electrons approach closer to + nucleus Li (period 2) compound more negative LE than K (period 4) compound Follows ionic radius trend – Li ion < K ion

85 Lattice Energy (LE) Related to size and charge of ions
Bonds formed from ions with larger charges have more negative LE MgO (+2, , 140 pm) 4 X more negative LE than NaF (+1, , 133 pm)

86 See next 2 slides for explanation of arrows
Lattice Energy (LE) See next 2 slides for explanation of arrows Compound LE (kJ/mol) Cation r (pm) Anion KI -632 138 220 RbF -774 152 133 KF -808 AgCl -910 126 181 SrCl2 -2142 118 MgO -3795 72 140

87 Lattice Energy (LE) – Size Effects
Compound LE (kJ/mol) Cation r (pm) Anion KI -632 138 220 RbF -774 152 133 KF -808 I vs F – anion size decrease Rb vs K – cation size decrease

88 Lattice Energy (LE) – Charge Effect
Compound LE (kJ/mol) Cation r (pm) Anion AgCl -910 126 181 SrCl2 -2142 118 Ag vs Sr – cation charge increase from +1 to +2

89 Chapter 7– Ionic Compounds & Metals
7.0 Bonding Overview 7.1 Ion Formation 7.2 Ionic Bonds and Ionic Compounds 7.3 Names & Formulas for Ionic Compounds 7.4 Metallic Bonds & Properties of Metals

90 Section 7.3 Names and Formulas for Ionic Compounds
In written names and formulas for ionic compounds, the cation appears first, followed by the anion. Determine the oxidation state of metals and nonmetals based on their position in the periodic table Relate a formula unit of an ionic compound to its composition. Know the formulas and charges of common polyatomic ions Name ionic compounds and produce the formula of a compound given its name

91 Section 7.3 Names and Formulas for Ionic Compounds
Know how to transform an oxyanion name to adjust for increased or decreased oxygen, addition of hydrogen, or change in halogen Manipulate subscripts (including use of parentheses for polyatomic ions) in the chemical formula of an ionic compound to produce a neutral compound Know when to use roman numerals in the names for ionic compounds

92 Section 7.3 Names and Formulas for Ionic Compounds
Key Concepts A formula unit gives the ratio of cations to anions in the ionic compound. A monatomic ion is formed from one atom. The charge of a monatomic ion is its oxidation number. Roman numerals indicate the oxidation number of cations having multiple possible oxidation states. Polyatomic ions consist of more than one atom and act as a single unit. To indicate more than one polyatomic ion in a chemical formula, place parentheses around the polyatomic ion and use a subscript.

93 Formulas formula unit = simplest ratio of ions in compound
Overall charge is zero KBr (K+ Br-) MgCl2 (Mg2+ 2 x Cl-) Monatomic ion = one atom ion Na+1, O2-

94 Oxidation Number Charge on monatomic ion = oxidation number or oxidation state Na+1 has oxidation number of +1 Equals number of electrons transferred from atom to form the ion + sign transferred from - sign transferred to Oxidation numbers of elements in ionic compound must sum to zero

95 Rules for Naming Ionic Compounds (see page 223)
1. Name cation first, anion second Cation is always written first in formula Monatomic cations use element name sodium magnesium lead Monatomic anions use element root name plus suffix –ide chloride oxide sulfide nitride

96 Naming Monatomic Anions
For monatomic anions, name becomes element root + “ide” ending Chlorine (Cl)  Chloride (Cl-) Oxygen (O)  Oxide (O2-) Sulfur (S)  Sulfide (S2-) Nitrogen  Nitride (N3-) Phosphorus  Phosphide (P3-) Carbon  Carbide (C4-)

97 Ionic Compound Names Examples for Monatomic Ions Cations Either Al or Groups 1A, 2A
NaCl – Sodium chloride MgO – Magnesium oxide Al2O3 – Aluminum oxide – Calcium phosphide Ca3P2 Li3N – Lithium nitride

98 Naming Ionic Compounds Binary Compounds
Write formulas from names Practice problems 19-23, page 221 Write names from formulas Practice problems 28-30, page 223

99 Polyatomic Ions Ions made up of more than one atom
OH- (hydroxide) SO42- (sulfate) NH4+ (ammonium) Note: atoms within the polyatomic ion are covalently bonded to each other Never change the subscripts of polyatomic ions

100 Polyatomic Ions

101 p 221

102 * AKA hydrogen carbonate
Polyatomic Ions Table 7.9, page 221 lists common ones You must know: Ammonium NH4+ Nitrite, Nitrate NO2- NO3- Hydroxide OH- Bicarbonate*, Carbonate HCO3- CO32- Phosphate PO43- Peroxide O22- Sulfate SO42- Chlorate ClO3- * AKA hydrogen carbonate

103 Polyatomic Ions Table R-5, page 970 has more comprehensive list sorted by net charge See also Wikipedia

104 Naming Ionic Compounds With Polyatomic Ions
Write formulas from names Practice problems 24-27, page 222 Write names from formulas Practice problems 30-33, page 223

105 Naming Ions - Oxyanions
Oxyanion = Polyatomic anions Same nonmetal element Differing number of oxygen atoms NO NO Nonmetal N 3 vs 2 O SO SO Nonmetal S 4 vs 3 O

106 Naming Ions - Oxyanions
Rules: Greater O atoms, use nonmetal root + suffix –ate Fewer O atoms, use nonmetal root + suffix –ite NO3- nitrate NO2- nitrite SO42- sulfate SO32- sulfite

107 Naming Ions Oxyanions with Halogen or P
Rules: More complicated because can have three or four different anions ClO ClO3- ClO ClO- perchlorate chlorate chlorite hypochlorite PO PO PO23- Phosphate Phosphite Hypophosphite

108 Oxyanion Naming Conventions for Chlorine (table 7.11, p 223)

109 Naming Oxyanions with Br and I
All rules followed by chlorine are followed by bromine and iodine as well Change root name Chlorate Bromate Iodate ClO BrO IO3- IO4- ? periodate

110 Polyatomic Ions In series with varying oxygen (only), charge fixed
Nitrate, Nitrite NO3- NO2- Sulfate, Sulfite SO42-, SO32- Phosphate, Phosphite, Hypophosphite PO43-, PO33- , PO23- Perchlorate, Chlorate, Chlorite, Hypochlorite ClO4- ClO3- ClO2- ClO-

111 Polyatomic Ions For series differing by an H, charge increases by +1 for each added H Carbonate, Hydrogen carbonate CO32- HCO3- Sulfate, Hydrogen Sulfate SO42- HSO4- Phosphate, Hydrogen Phosphate Dihydrogen Phosphate PO43- HPO42- H2PO4-

112 Rules for Naming Ionic Compounds
4. A. Group 1, 2 metals or Al – no additional work necessary B. Some group 13 to 15 metals and most transition metals (see following slide) have multiple oxidation states - use Roman numeral in parentheses to indicate which one

113 f-block transition metals
d-block transition metals f-block transition metals Group 13 to 15 metals with multiple oxidation states Zn Ga Ag Cd In Sn Tl Pb Bi Lanthanides Actinides

114 Rules for Naming Ionic Compounds
4. Group 13 to 15 metals + transition metals Iron(II) Iron(III) Copper(I) Copper(II) FeO – Iron(II) oxide CuCl – Copper(I) chloride Fe2O3 – Iron(III) oxide CuCl2 – Copper(II) chloride Outdated but still commonly used naming system uses –ous, -ic suffixes (not responsible for these) Ferrous Ferric Cuprous Cupric

115 Table 7.8, page 219: values for some transition and group 3A / 4A (13/14) metals

116 Cation Oxidation Numbers
There are certain common cations (beyond group 2) with fixed oxidation numbers that do not get roman numerals in their compound names For table 7.8, these are: Ag+, Zn2+, Cd2+, Al3+ (silver, zinc, cadmium, aluminum) You must know the four ions above ZnCl2 zinc chloride, not zinc(II) chloride Sometimes Sc3+ included in this list

117 Figure top of p 224

118 Rules for Naming Ionic Compounds
If compound contains a polyatomic ion, use the ion name NH4Cl ammonium chloride NaOH sodium hydroxide (NH4)2SO ammonium sulfate [note use of parentheses and subscript in this compound to obtain neutral compound]

119 Section 7.3 Assessment Which subscripts would you most likely use for an ionic compound containing a group 1 metal and a group 17 nonmetal? (Remember, 1 = no written subscript) A. 1 and 2 B. 2 and 1 C. 2 and 3 D. 1 and 1 ???

120 Name of the compound Ca(OH)2?
Section 7.3 Assessment Name of the compound Ca(OH)2? A. calcium oxide B. calcium (II) hydroxide C. calcium hydroxide D. calcium oxyhydride ???

121 Ionic Compounds Practice
Name the following: Ba(NO3)2 barium nitrate calcium sulfate CaSO4 magnesium phosphate Mg3(PO4)2 SrSO3 strontium sulfite

122 Ionic Compounds Practice
Name the following: BaO CaF2 Mg3N2 SrS Mn(H2PO4)3 barium oxide calcium fluoride magnesium nitride strontium sulfide manganese (III) dihydrogen phosphate

123 Ionic Compounds Practice
Name the following: FeBr2 FeBr3 SnCl2 SnCl4 iron(II) bromide iron(III) bromide tin(II) chloride tin(IV) chloride

124 Ionic Compounds Practice
Name the following: AgCl Na2Se Fe2O3 CrI3 silver chloride sodium selenide iron(III) oxide chromium(III) iodide

125 Ionic Compounds Practice
Name the following: BaBr2 K2S AlN SnF4 Cd(OH)2 barium bromide potassium sulfide aluminum nitride tin(IV) fluoride cadmium hydroxide

126 Ionic Compounds Practice
Name the following: Fe(OH)3 NH4I Na2O2 Ca(ClO)2 iron(III) hydroxide ammonium iodide sodium peroxide calcium hypochlorite

127 Ionic Compounds Practice
Name following compounds: Fe(NO3)2 Fe(NO2)3 Sn(ClO)2 Sn(ClO2)4 iron(II) nitrate iron(III) nitrite tin(II) hypochlorite tin(IV) chlorite

128 Ionic Compounds Practice
Name following compounds: AgHSO4 (NH4)2CO3 Silver hydrogen sulfate Ammonium carbonate

129 Naming Ionic Compounds
Problem 81 page 233 (get formula) Problem 82 page 233 (get name)

130 Systematic vs Common Names
Elaborate rules exist for assigning names to chemical substances on basis of their structures Called systematic names; they uniquely identify given substance Rules for these names are defined by international body (IUPAC)

131 Chapter 7– Ionic Compounds & Metals
7.0 Bonding Overview 7.1 Ion Formation 7.2 Ionic Bonds and Ionic Compounds 7.3 Names & Formulas for Ionic Compounds 7.4 Metallic Bonds & Properties of Metals

132 Section 7.4 Metallic Bonds and the Properties of Metals
Metals form crystal lattices and can be modeled as cations surrounded by a “sea” of freely moving (delocalized) valence electrons. Describe a metallic bond. Describe the meaning of the words/terms “delocalized electron”, malleable, and ductile. Describe how the properties of conductivity, reflectivity malleability and ductility are related to the presence of delocalized electrons (electron sea model). Describe the similarities and differences between ionic and metallic bonding.

133 Section 7.4 Metallic Bonds and the Properties of Metals
Metals form crystal lattices and can be modeled as cations surrounded by a “sea” of freely moving (delocalized) valence electrons. Define alloys, categorize them into two basic types, list the two types of solution alloys, and give examples of each. List possible advantages of using an alloy over using a pure metal Explain the role that carbon plays in steel alloys. Describe the roles that imperfections play in the properties of metals and list various physical methods that are used to alter these imperfections.

134 Section 7.4 Metallic Bonds and the Properties of Metals
Key Concepts A metallic bond forms when metal cations attract freely moving, delocalized valence electrons. In the electron sea model, electrons move through the metallic crystal and are not held by any particular atom. The electron sea model explains the physical properties of metallic solids. Metal alloys are formed when a metal is mixed with one or more other elements.

135 Metallic Bonds Metals don’t form ionic bonds
Do form solid state lattices Lattice similar to ionic crystal lattice Have valance electrons, but these are free to roam in a “sea” of other electrons Electrons are “delocalized” – not confined to any particular location

136 Metallic Bonds – “Electron Sea Model”
Metal ion (+) Metal lattice structure Free electron (-)

137 Metallic Bonds – “Electron Sea Model”
Metal ion (+) Free electron “sea”

138 Not very directional, so metal atoms can be rearranged without problem
Metallic Bonds Metallic bond is the attraction of a metallic cation for the delocalized electrons Not very directional, so metal atoms can be rearranged without problem Gives ductility and malleability

139 Metallic Bonds – MP & BP Indicate strength of metallic bond
BP more extreme than MP – large energy required to separate atoms from soup of cations and electrons

140 Metals – Malleable & Ductile

141 Malleable

142 Malleable Electrons allow atoms to slide by

143 Impart good electrical conductivity
Mobile Electrons Impart good electrical conductivity Interact with light, absorbing & releasing photons Redirected light gives luster

144 Delocalized Electrons & Properties
As number of delocalized electrons increases, so does hardness and strength Alkali metals soft (1 valence electron) In transition metals, unpaired d electrons are delocalized, so transition metals in the middle of the d block tend to be harder and stronger and also to have higher MPs

145 Melting Points (C)

146

147 Period 6 – s block & TM Melting Points
Peak occurs at W (4 unpaired d electrons) Atomic Number Melting Point (K)

148 Alloys Alloys have more than one element (one a metal)
Alloy has metal characteristics Pure metals and alloys have different physical and chemical properties Strength, hardness, corrosion resistance In jewelry, alloy of gold & copper used alloy harder (& cheaper) than pure gold

149 Alloys - Types Solution alloys are homogeneous
Heterogeneous alloys: components are not dispersed uniformly Steel with >1.4% C has 2 phases: almost pure Fe and cementite, Fe3C (iron carbide) Fe3C is white, hard, brittle – makes steel less ductile but much stronger

150 Alloys Two types of solution alloy
Substitutional alloys - some atoms in the original metallic solid are replaced by other metals of similar atomic structure Interstitial alloys - formed when small holes in a metallic crystal are filled with smaller atoms (solute occupies interstitial sites in metallic lattice)

151 Alloys Substitutional Interstitial

152 Alloys

153 Alloys Substitutional alloys Interstitial alloys
atoms must have similar atomic radii elements must have similar bonding characteristics Sterling silver – Ag 92.5% Cu 7.5% Interstitial alloys one element must have a significantly smaller radius than the other (must fit into interstitial site) e.g. a nonmetal – Carbon Steel

154 Metal Properties The chemical composition (alloying elements) of a metal is only one factor that determines metal properties Properties such as hardness and strength also depend on any mechanical and heat treatments that may be applied These treatments effect how the alloying elements are distributed within the alloy, the crystal size, and the number and type of crystal defects within the material

155 Classification of Commerically Important Metals
Ferrous Metals Non- Ferrous Metals Iron Aluminum Low Carbon Steel Copper Medium Carbon Steel Brass High Carbon Steel Bronze Cast Iron Zinc Alloy Steel Lead Stainless Steel Tin Others

156 Steel 0.001% to 1.5% carbon Wide range of properties due to
Variation in carbon content Cold working (work hardening) Heat treatment Addition of alloying elements

157 Steel and Carbon Carbon even at relatively low levels has an impact on steel properties Because iron and carbon form an interstitial alloy, carbon acts as a “stiffener” to prevent the layers of iron ions from moving freely relative to each other Result is a harder, stronger but more brittle alloy as the carbon content increases *Note, some texts including CES state med carbon as .3% to .5%

158 Steel - Effect of Increasing Carbon
Decreases ductility Decreases machinability Lowers melting point Increases tensile strength Increases hardness Makes steel easier to harden with heat treatments Lowers temperature required to heat treat steel Increases difficulty of welding *Note, some texts including CES state med carbon as .3% to .5%

159 Steel Composition - % by Weight Balance is Fe
Nonmetals Metalloid Type SAE C Mn P S Si Remarks 1010 0.030 0.035 -- Common 1040 0.35 Tools 1552 Tempered Parts

160 Metal Properties – Other Factors
Although chemical composition (% Fe, % C, etc) plays important role, other factors strongly influence metal’s properties (hardness, toughness, etc) Mechanical treatment (working) Heat treatment (tempering, quenching) Distribution of elements within metal (often not homogeneous) All of above can interact – study is field of metallurgy

161 Cooling Rate and Crystal Size
The way metal prepared can have large impact on how it behaves Many metals prepared in liquid state & cooled; rate of cooling can have significant effect on properties of solid because it controls crystal size/grain structure

162 Grain Structure & Imperfections (NIB)
Structure not continuous throughout As metal cools, have l  s phase change, atoms come together to form grains Crystal structure not continuous Steel paper clip Fe crystal structure Grain Grain Boundary

163 Formation of Grain Structure
Solidification of molten material Two steps starting with molten material (all liquid) 1) Nuclei form 2) Nuclei grow to form crystals Crystals grow until they meet each other to form grain structure nuclei crystals growing grain structure liquid

164 Metal Crystal Size Small crystals make metal harder because ions less able to move; also means there is more disruption between crystals making them brittle (easy to break) Larger crystals make metal soft

165 Imperfections and Alloys
Many imperfections within each crystal Flaws produce weak points in bonds between atoms Adding other elements to produce an alloy can counteract effects of imperfections and make metal harder and stronger Heat and mechanical treatment also effect these imperfections

166 Area Defects: Grain Boundaries
• boundaries between crystals • produced by solidification process • have change in crystal orientation across them • impede dislocation motion

167 Imperfections in Solids
Schematic drawing of poly-crystal with many defects

168 Grain Structure & Imperfections (NIB)
Micrograph of metal that has undergone intergranular corrosion Grain Boundary Grain

169 Heat Treatment of Metals
3 ways of treating a metal with heat: Annealing Quenching Tempering Steel is alloy most commonly treated Used to: Soften part that is too hard Harden part that is not hard enough Put hard skin on part that is soft Make good magnets out of ordinary material Make selective property changes within parts

170 Heat Treatment Metal on striking face of hammer heat-treated differently than that on rest of head Hardness on front traded for toughness at back

171 A metal is heated to a moderate temperature and allowed to cool slowly
Treatment Process Effect on metal properties Effect on metal structure Annealing A metal is heated to a moderate temperature and allowed to cool slowly The metal is softer with improved ductility Larger metal crystals form Quenching A metal is heated to a moderate temperature and cooled quickly (sometimes by plunging into water The metal is harder and brittle. Tiny metals crystals form. Tempering A quenched metal is heated (to a lower temperature than is used for quenching and allowed to cool The metal is harder but less brittle. Crystals of intermediate size form.

172 Mechanical & Heat Treatment of Metals
Blacksmith creates objects from wrought iron or steel by forging the metal (using tools to hammer, bend, and cut) and in the process can also change the characteristics of the metal Also uses heat treatment Arvind Thekdi - E3M, Inc. Sales

173 Cold Working Increases strength at the expense of ductility
Arvind Thekdi - E3M, Inc. Sales

174 Mechanical Treatment of Metals
Work hardening (aka strain hardening or cold working) is strengthening of metal by plastic deformation. Strengthening occurs because of dislocation movements and dislocation generation within crystal structure of the material Most non-brittle metals with a reasonably high melting point as well as several polymers can be strengthened in this fashion Alloys not amenable to heat treatment, including low-carbon steel, are often work-hardened Arvind Thekdi - E3M, Inc. Sales

175 Mechanical Treatment of Metals
Cold rolling increases strength via strain hardening – metal grains become elongated Cold rolled steel


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