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

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Presentation on theme: "Chapter 7– Ionic Compounds & Metals 7.0Bonding Overview 7.1 Ion Formation 7.2 Ionic Bonds and Ionic Compounds 7.3Names & Formulas for Ionic Compounds 7.4."— Presentation transcript:

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

2 Section 7.0 Bonding Overview 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 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.

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 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. Section 7.0 Bonding Overview

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) CO 2 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 – – – – – 4.0

12 Fig (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 Electronegativity Difference % Ionic Character Ionic Bonds 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  EN > 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 7.0Bonding Overview 7.1 Ion Formation 7.2 Ionic Bonds and Ionic Compounds 7.3Names & Formulas for Ionic Compounds 7.4 Metallic Bonds & Properties of Metals

19 Section 7.1 Ion Formation 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.. Ions are formed when atoms gain or lose valence electrons to achieve a stable octet electron configuration.

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 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. Section 7.1 Ion Formation

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 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 1 [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 ns 2 configuration, so generally can form +2 cation Also have (n-1)d x 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: Sc 3+ (many sources don’t list) Zn 2+ Ag 1+ Cd 2+ (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]3d 2 4s 2 (filling order 4s then 3d) To get electron configuration of ion, remove valence electrons first Ti 2+ [Ar]3d 2

33 Special Case Transition Metals Pseudo-noble gas configuration Groups , periods Full sublevels have extra stability s 2, p 6, d 10 – 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]3s 2 3p 6 3d 10 4s 2 Zn +2 = [Ne]3s 2 3p 6 3d 10 Compare to Ar = [Ne]3s 2 3p 6 Consequence: Zn +2 only Zn ion formed

35 Zn +2 – Pseudo-Noble Gas

36 Special Case Transition Metals Pseudo-noble gas configuration Ag [Kr]4d 10 5s 1 Same situation as Cu (not 5s 2 ) Ag 1+ [Kr]4d 10 = [Ar] 3d 10 4s 2 4p 6 4d 10 Compare to Kr = [Ar] 3d 10 4s 2 4p 6 Consequence: Ag 1+ only Ag ion formed Reasoning for Ag 1+ applies to Cd 2+ Cd [Kr]4d 10 5s 2 = [Ar]3d 10 4s 2 4p 6 4d 10 5s 2 Cd 2+ = [Ar]3d 10 4s 2 4p 6 4d 10

37 Special Case Transition Metals Noble gas configuration Sc [Ar]3d 1 4s 2 Sc 3+ = [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 (O 2- ) Sulfur (S)  Sulfide (S 2- ) Nitrogen  Nitride (N 3- ) Phosphorus  Phosphide (P 3- ) Carbon  Carbide (C 4- )

41 Atom vs Anion Radius Formation of chlorine ion Cl + e -  Cl - Chlorine atom Chloride anion [Ne]3s 2 3p 5 [Ne]3s 2 3p 6 or [Ar]

42 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 N Add 1 electron at a time to each side Until they are forced to pair up.

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) P P 3- Gain electrons to fill outer energy level

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

47 Section 7.2 Ionic Bonds and 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. Oppositely charged ions attract each other, forming electrically neutral ionic compounds.

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 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. Section 7.2 Ionic Bonds and Ionic Compounds

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 AlCl 3 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) Al 2 O 3 (2:3 ratio) Be 4 Al 2 Si 6 O 18 formula unit for beryl Simple binary compounds = 2 different elements MgO, KCl, FeBr 2, Al 2 O 3

53 NaCl Formation Electron Configuration Picture Transferred Electron

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

55 NaCl Formation Electron Dot Picture NaCl

56 Na + Cl - 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) Neutral Na Atom Neutral Cl Atom Electron loss Chloride Ion Electron gain Sodium Ion 10 e - 18 e - 17 e - 11 e -

58 NaCl Crystalline structure

59 Ionic Bonding All the electrons must be accounted for! CaP

60 Ionic Bonding CaP

61 Ionic Bonding Ca 2+ P 2-

62 Ionic Bonding Ca 2+ Ca

63 Ionic Bonding Ca 2+ P 3- Ca 1+

64 Ionic Bonding Ca 2+ P 3- Ca 1+ P

65 Ionic Bonding Ca 2+ P 3- Ca 2+ P 1-

66 Ionic Bonding Ca 2+ P 3- Ca 2+ Ca P 1-

67 Ionic Bonding Ca 2+ P 3- Ca 2+ Ca P 1-

68 Ionic Bonding Ca 2+ P 3- Ca 2+ P 3- Ca 2+

69 Ionic Bonding – Formula Unit = Ca 3 P 2 Formula Unit for Calcium Phosphide 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 - /r 2  q + = magnitude of positive (cation) charge  q - = magnitude of negative (anion) charge  r = distance between charge centers

73 Force vs distance for 1/r 2

74 Coulomb’s Law & Ionic Bonds F  q + q - /r 2 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 KBr NaBr CaCl 2 782>1600 CaI NaCl MgO

77 Ionic solids are brittle

78 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 H 2 O 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, -2 72, 140 pm) 4 X more negative LE than NaF (+1, , 133 pm)

86 Lattice Energy (LE) Comp ound LE (kJ/mol) Cation r (pm) Anion r (pm) KI RbF KF AgCl SrCl MgO See next 2 slides for explanation of arrows

87 Lattice Energy (LE) – Size Effects Comp ound LE (kJ/mol) Cation r (pm) Anion r (pm) KI RbF KF I vs F – anion size decrease Rb vs K – cation size decrease

88 Lattice Energy (LE) – Charge Effect Comp ound LE (kJ/mol) Cation r (pm) Anion r (pm) AgCl SrCl Ag vs Sr – cation charge increase from +1 to +2

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

90 Section 7.3 Names and Formulas for Ionic Compounds 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 In written names and formulas for ionic compounds, the cation appears first, followed by the anion.

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 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. Section 7.3 Names and Formulas for Ionic Compounds

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

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 2.Monatomic cations use element name sodium magnesium lead 3.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 (O 2- ) Sulfur (S)  Sulfide (S 2- ) Nitrogen  Nitride (N 3- ) Phosphorus  Phosphide (P 3- ) Carbon  Carbide (C 4- )

97 Ionic Compound Names Examples for Monatomic Ions Cations Either Al or Groups 1A, 2A NaCl– Sodium chloride MgO– Magnesium oxide Al 2 O 3 – Aluminum oxide Ca 3 P 2 – Calcium phosphide Li 3 N– 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) SO 4 2- (sulfate) NH 4 + (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 Polyatomic Ions Table 7.9, page 221 lists common ones You must know: AmmoniumNH 4 + Nitrite, NitrateNO 2 - NO 3 - HydroxideOH - Bicarbonate*, Carbonate HCO 3 - CO 3 2- PhosphatePO 4 3- PeroxideO 2 2- SulfateSO 4 2- ChlorateClO 3 - * 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 3 - NO 2 - Nonmetal N 3 vs 2 O SO 4 2- SO 3 2- 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 NO 3 - nitrate NO 2 - nitrite SO 4 2- sulfate SO 3 2- sulfite

107 Naming Ions Oxyanions with Halogen or P Rules: More complicated because can have three or four different anions ClO 4 - ClO 3 - ClO 2 - ClO - perchlorate chlorate chlorite hypochlorite PO 4 3- PO 3 3- PO 2 3- 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 3 - BrO 3 - IO 3 - IO 4 - ? periodate

110 Polyatomic Ions In series with varying oxygen (only), charge fixed Nitrate, NitriteNO 3 - NO 2 - Sulfate, SulfiteSO 4 2-, SO 3 2- Phosphate, Phosphite, Hypophosphite PO 4 3-, PO 3 3-, PO 2 3- Perchlorate, Chlorate, Chlorite, Hypochlorite ClO 4 - ClO 3 - ClO 2 - ClO -

111 Polyatomic Ions For series differing by an H, charge increases by +1 for each added H Carbonate, Hydrogen carbonate CO 3 2- HCO 3 - Sulfate, Hydrogen Sulfate SO 4 2- HSO 4 - Phosphate, Hydrogen Phosphate Dihydrogen Phosphate PO 4 3- HPO 4 2- H 2 PO 4 -

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

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 Fe 2 O 3 – Iron(III) oxide CuCl 2 – 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 +, Zn 2+, Cd 2+, Al 3+ (silver, zinc, cadmium, aluminum) You must know the four ions above ZnCl 2 zinc chloride, not zinc(II) chloride Sometimes Sc 3+ included in this list

117 Figure top of p 224

118 Rules for Naming Ionic Compounds 5.If compound contains a polyatomic ion, use the ion name NH 4 Cl ammonium chloride NaOH sodium hydroxide (NH 4 ) 2 SO 4 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 ? A.calcium oxide B.calcium (II) hydroxide C.calcium hydroxide D.calcium oxyhydride ??? Section 7.3 Assessment

121 Ionic Compounds Practice SrSO 3 barium nitrate magnesium phosphate calcium sulfate strontium sulfite Name the following: Ba(NO 3 ) 2 CaSO 4 Mg 3 (PO 4 ) 2

122 Name the following: BaO CaF 2 Mg 3 N 2 SrS Mn(H 2 PO 4 ) 3 barium oxide calcium fluoride magnesium nitride strontium sulfide manganese (III) dihydrogen phosphate Ionic Compounds Practice

123 Name the following: FeBr 2 FeBr 3 SnCl 2 SnCl 4 iron(II) bromide iron(III) bromide tin(II) chloride tin(IV) chloride Ionic Compounds Practice

124 Name the following: AgCl Na 2 Se Fe 2 O 3 CrI 3 Ionic Compounds Practice silver chloride sodium selenide iron(III) oxide chromium(III) iodide

125 Name the following: BaBr 2 K 2 S AlN SnF 4 Cd(OH) 2 Ionic Compounds Practice barium bromide potassium sulfide aluminum nitride tin(IV) fluoride cadmium hydroxide

126 Name the following: Fe(OH) 3 NH 4 I Na 2 O 2 Ca(ClO) 2 Ionic Compounds Practice iron(III) hydroxide ammonium iodide sodium peroxide calcium hypochlorite

127 Ionic Compounds Practice Name following compounds: Fe(NO 3 ) 2 Fe(NO 2 ) 3 Sn(ClO) 2 Sn(ClO 2 ) 4 iron(II) nitrate iron(III) nitrite tin(II) hypochlorite tin(IV) chlorite

128 Ionic Compounds Practice Name following compounds: AgHSO 4 (NH 4 ) 2 CO 3 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 7.0Bonding Overview 7.1 Ion Formation 7.2 Ionic Bonds and Ionic Compounds 7.3Names & Formulas for Ionic Compounds 7.4 Metallic Bonds & Properties of Metals

132 Section 7.4 Metallic Bonds and the Properties of Metals 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. Metals form crystal lattices and can be modeled as cations surrounded by a “sea” of freely moving (delocalized) valence electrons.

133 Section 7.4 Metallic Bonds and the Properties of Metals 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. Metals form crystal lattices and can be modeled as cations surrounded by a “sea” of freely moving (delocalized) valence electrons.

134 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. Section 7.4 Metallic Bonds and the Properties of Metals

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 (+) Free electron (-) Metal lattice structure

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

138 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 Electrons allow atoms to slide by

143 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 Atomic Number Melting Point (K) Peak occurs at W (4 unpaired d electrons)

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, Fe 3 C (iron carbide) Fe 3 C 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 SubstitutionalInterstitial

152 Alloys

153 Substitutional 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 Ferrous MetalsNon- Ferrous Metals IronAluminum Low Carbon SteelCopper Medium Carbon SteelBrass High Carbon SteelBronze Cast IronZinc Alloy SteelLead Stainless SteelTin Others Classification of Commerically Important Metals

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 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 Steel and Carbon

158 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 Steel - Effect of Increasing Carbon

159 Steel Composition - % by Weight Balance is Fe Type SAE CMnPSSiRemarks Common Tools Tempered Parts NonmetalsMetalloid

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

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 Grain Boundary

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 TreatmentProcessEffect on metal properties Effect on metal structure AnnealingA metal is heated to a moderate temperature and allowed to cool slowly The metal is softer with improved ductility Larger metal crystals form QuenchingA 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. TemperingA 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 Arvind Thekdi - E3M, Inc. Sales 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 Mechanical & Heat Treatment of Metals

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

174 Arvind Thekdi - E3M, Inc. Sales 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 Mechanical Treatment of Metals

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


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