Salts & Metals

download report

Transcript Salts & Metals

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

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

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

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.

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

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

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

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

Ionic vs Covalent Bonding

Key parameter in distinguishing types is the electronegativity

Electronegativity

Relative ability to attract electrons in a chemical bond • Max value 3.98 F • Min value 0.7 Fr Elements with high EN tend to form anions Noble gases not tabulated • Very few compounds to get info from

Electronegativity Ranges (values slightly different than in book) Below 1.0

1.0 – 1.4

1.5 – 1.9

2.0 – 2.4

2.5 – 2.9

3.0 – 4.0

Fig. 8.20 (p. 265) Electronegativity

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 shared  covalent or zero, electron bond

EN Difference & Bond Character

Ionic Bonds Covalent Bonds 0 1.0

2.0

3.0

Electronegativity Difference

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 D EN > 1.7) Polar covalent bonds involve unequally shared valence electrons (but not so unequal that ions can form) Most bonds have some mix character of ionic & covalent

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

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

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

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..

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

each, and

explain

the oxidation number associated with using an argument involving the ion’s electron configuration why this single oxidation state is highly preferred.

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.

Bonding

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

Bonding

Metals are reactive because they lose valence electrons easily

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

Atom vs Cation Radius

Ionization of sodium Na  Sodium atom [Ne]3s 1 Na + + e Sodium cation [Ne] -

Formation of Cations

In most cases, lose enough electrons to achieve noble gas configuration

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

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)

Transition Metal Oxidation States

(uses old group designations) Common Less Common

Transition Metal Oxidation States

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

Forming Transition Metal Cations

For atom, following standard Aufbau order for transition metal means n s 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

Special Case Transition Metals

Pseudo-noble gas configuration • Groups 3 - 14, periods 4 - 6 • 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

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

Zn

+2

– Pseudo-Noble Gas

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

Special Case Transition Metals

Noble gas configuration • Sc [Ar]3d 1 4s 2 • Sc 3+ = [Ar] = Only common oxidation state for Sc (+3)

Formation of Anions

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

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)

Naming Monatomic Anions

For monatomic anions, name becomes element root + “ ide ” ending • Chlor ine (Cl)  Chlor ide (Cl ) • Ox ygen (O)  Ox ide (O 2 ) • Sulf ur (S)  Sulf ide (S 2 ) • Nitr ogen  Nitr ide (N 3 ) • Phosph orus  Phosph ide (P 3 ) • Carb on  Carb ide (C 4 )

Atom vs Anion Radius

Formation of chlorine ion Cl + e  Cl Chlorine atom [Ne]3s 2 3p 5 Chlor ide anion [Ne]3s 2 3p 6 or [Ar]

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

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

Electron Dots For Cations

Metals generally have few valence electrons These will come off Forming positive ions Ca 2+

Electron Dots For Anions

Nonmetals have many valence electrons (usually 5 or more) Gain electrons to fill outer energy level P P 3-

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

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.

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.

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.

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

Ionic Bonding

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

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 elements compounds = • MgO, KCl, FeBr 2 , Al 2 O 3 2 different

NaCl Formation Electron Configuration Picture Transferred Electron

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

NaCl Formation Electron Dot Picture

Na Cl

NaCl Formation Electron Dot Picture

Na

+

Cl

Note that after electron transfer, the atoms are charged

Sodium Chloride Formation

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

NaCl Crystalline structure

Ionic Bonding

All the electrons must be accounted for!

Ca P

Ca

Ionic Bonding

P

Ionic Bonding

Ca

2+

P

2-

Ionic Bonding

Ca

2+

Ca P

2-

Ca

Ionic Bonding

2+

Ca

1+

P

3-

Ionic Bonding

Ca

2+

Ca

1+

P

3-

P

Ionic Bonding

Ca

2+

Ca

2+

P

3-

P

1-

Ionic Bonding

Ca Ca

2+

Ca

2+

P

3-

P

1-

Ionic Bonding

Ca Ca

2+

Ca

2+

P

3-

P

1-

Ionic Bonding

Ca

2+

Ca

2+

Ca

2+

P

3-

P

3-

Ionic Bonding – Formula Unit

= Ca 3 P 2 Formula Unit for Calcium Phosphide Simplest ratio of ions in compound is called the formula unit

Properties of Ionic Compounds

Crystalline structure, usually solids A regular

repeating

ions in the solid – 3D arrangement of 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)

Crystalline structure

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

Force vs distance for 1/r

2 0 0 0 0 0 0 0 0 0 5 15 25 35

r (distance)

45

Coulomb’s Law & Ionic Bonds

F  q + q /r 2 The larger the ionic charges stronger (q + , q the force between them ), the The closer together the ions are, the stronger the force • Smaller ions can get closer together than larger ions

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

Ionic Solids – Melting, Boiling Points

Compound MP (  C) BP (  C) NaI KBr NaBr CaCl 2 CaI 2 NaCl MgO 660 734 747 782 784 801 2852 1304 1435 1390 >1600 1100 1413 3600

Ionic solids are brittle

+ + + + + + + +

Ionic solids are brittle

Strong Repulsion breaks crystal apart.

+ + + + + +

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

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

Conductivity of Ionic Solids

Conduct when dissolved in water ( electrolyte )

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

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

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

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 , -1 102, 133 pm)

Comp ound

Lattice Energy (LE)

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

Lattice Energy (LE) – Size Effects

Comp ound LE (kJ/mol) KI RbF KF -632 -774 -808 Cation r (pm) 138 152 138 Anion r (pm) 220 133 133 I vs F – anion size decrease Rb vs K – cation size decrease

Lattice Energy (LE) – Charge Effect

Comp ound LE (kJ/mol) AgCl SrCl 2 -910 -2142 Cation r (pm) 126 118 Anion r (pm) 181 181 Ag vs Sr – cation charge increase from +1 to +2

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

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 compound given its name

produce

the formula of a

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

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.

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-

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

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 chlor ide ox ide sulf ide nitr ide

Naming Monatomic Anions

For monatomic anions, name becomes element root + “ ide ” ending • Chlor ine (Cl)  Chlor ide (Cl ) • Ox ygen (O)  Ox ide (O 2 ) • Sulf ur (S)  Sulf ide (S 2 ) • Nitr ogen  Nitr ide (N 3 ) • Phosph orus  Phosph ide (P 3 ) • Carb on  Carb ide (C 4 )

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

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

Polyatomic Ions

Ions made up of more than one atom • OH (hydroxide) • SO 4 2 • NH 4 + (sulfate) (ammonium) Note: atoms within the polyatomic ion are covalently bonded to each other Never change the subscripts of polyatomic ions

Polyatomic Ions

p 221

Polyatomic Ions

Table 7.9, page 221 lists common ones You must know : • Ammonium • Nitrite, Nitrate • Hydroxide NH 4 + NO 2 OH NO 3 • Bicarbonate * , Carbonate HCO 3 • Phosphate PO 4 3 • Peroxide O 2 2 * • Sulfate • Chlorate SO 4 2 ClO 3 CO AKA 3 hydrogen carbonate 2-

Polyatomic Ions

Table R-5, page 970 has more comprehensive list sorted by net charge See also Wikipedia http://en.wikipedia.org/wiki/Polyatomic_ion

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

Naming Ions - Oxyanions

Oxyanion = Polyatomic anions • Same nonmetal element • Differing number of oxygen atoms NO 3 SO 4 2 NO 2 SO 3 2 Nonmetal N 3 vs 2 O Nonmetal S 4 vs 3 O

Naming Ions - Oxyanions

Rules: • Greater O atoms, use nonmetal root suffix – ate + • Fewer O atoms, use nonmetal root suffix – ite + NO 3 SO 4 2 nitr ate sulf ate NO 2 SO 3 2 nitr ite sulf ite

Naming Ions Oxyanions with Halogen or P

Rules: More complicated because can have three or four different anions ClO 4 per chlor ate ClO 3 chlor ate ClO 2 ClO chlor ite hypo chlor ite PO 4 3 Phosph ate PO 3 3 PO 2 3 Phosph ite Hypo phosph ite

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

Naming Oxyanions with Br and I

All rules followed by chlorine are followed by bromine and iodine as well Change root name Chlor ate Brom ate ClO IO 4 ?

3 periodate BrO 3 Iod ate IO 3 -

Polyatomic Ions

In series with varying oxygen (only), charge fixed • Nitr ate , Nitr ite NO 3 NO 2 • Sulf ate , Sulf ite SO 4 2 , SO 3 2 • Phosph ate , Phosph ite , Hypo phosph ite PO 4 3 , PO 3 3- , PO 2 3 • Perchlorate, Chlorate, Chlorite, Hypochlorite ClO 4 ClO 3 ClO 2 ClO -

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 -

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

d-block

13 14 15

transition metals Zn

Ga

Ag Cd

In Sn Tl Pb Bi

Group 13 to 15 metals with multiple oxidation states Lanthanides Actinides f-block transition metals

Rules for Naming Ionic Compounds

4. Group 13 to 15 metals metals + transition Iron(II) Iron(III) Copper(I) Copper(II) FeO Fe 2 O – Iron(II) oxide CuCl – Copper(I) chloride 3 – Iron(III) oxide CuCl 2 – Copper(II) chloride Outdated but still commonly used naming system uses – ous , ic suffixes (not responsible for these) Ferr ous Ferr ic Cupr ous Cupr ic

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

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+ cadmium, aluminum) (silver, zinc, You must know the four ions above ZnCl 2 zinc chloride, not zinc( II ) chloride Sometimes Sc 3+ included in this list

Figure top of p 224

Rules for Naming Ionic Compounds

5. If compound contains a polyatomic ion, use the ion name NH 4 Cl ammonium chloride Na OH sodium hydroxide ( NH 4 ) 2 SO 4 ammonium sulfate [note use of parentheses and subscript in this compound to obtain neutral compound]

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.

B.

1 and 2 2 and 1 C.

2 and 3 D.

1 and 1 ???

Section 7.3 Assessment

Name of the compound Ca(OH) 2 ? A.

calcium oxide B.

calcium (II) hydroxide C.

D.

calcium hydroxide calcium oxyhydride ???

Ionic Compounds Practice

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

Ionic Compounds Practice

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

Ionic Compounds Practice

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

Ionic Compounds Practice

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

Ionic Compounds Practice

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

Ionic Compounds Practice

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

Ionic Compounds Practice

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

Ionic Compounds Practice

Name following compounds: AgHSO 4 (NH 4 ) 2 CO 3 Silver hydrogen sulfate Ammonium carbonate

Naming Ionic Compounds

Problem 81 page 233 (get formula) Problem 82 page 233 (get name)

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) http://www.chem.qmul.ac.uk/iupac/ http://www.acdlabs.com/iupac/nomenclature/

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

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.

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 Describe

the role that carbon plays in steel alloys.

the roles that imperfections play in the properties of metals and

list

various physical methods that are used to alter these imperfections.

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.

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

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

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

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

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

Metals – Malleable & Ductile

Malleable

Malleable

Electrons allow atoms to slide by

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

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

Melting Points (

C)

Trend of Unpaired d Electrons and 6 5 4 3 2 1 0 -1 20 Melting Point vs AN - Period 4 2500 25 Atomic Number 30 2250 2000 1750 1500 1250 1000 750 500 250 Unpaired d MP (K)

Period 6 – s block & TM Melting Points Peak occurs at W (4 unpaired d electrons)

Atomic Number

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

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

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)

Substitutional

Alloys

Interstitial

Alloys

Substitutional

Alloys

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

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

Classification of Commerically Important Metals

Ferrous Metals

Iron Low Carbon Steel Medium Carbon Steel High Carbon Steel Cast Iron Alloy Steel Stainless Steel Others

Non- Ferrous Metals

Aluminum Copper Brass Bronze Zinc Lead Tin Others

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

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

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

Steel Composition - % by Weight Balance is Fe Type C SAE 1010 0.08 0.13

1040 0.37 0.44

1552 0.47 0.55

Mn P S Si Remarks 0.30 0.60

0.60 0.90

1.20 1.55

0.030 0.035

- 0.030 0.035 0.35

Common Tools 0.030 0.035 0.35 Tempered Parts Nonmetals Metalloid

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

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

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 Grain Grain Boundary Fe crystal structure

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 liquid crystals growing grain structure

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

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

Area Defects: Grain Boundaries

Grain boundaries • boundaries between crystals • produced by solidification process • have change in crystal orientation across them • impede dislocation motion

grain boundaries

Imperfections in Solids

Schematic drawing of poly-crystal with many defects

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

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

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

Treatment Process Effect on metal properties Effect on metal structure

Annealing Quenching A metal is heated to a moderate temperature and allowed to cool slowly The metal is softer with improved ductility A metal is heated to a moderate temperature and cooled quickly (sometimes by plunging into water The metal is harder and brittle.

Larger metal crystals form 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.

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

Arvind Thekdi - E3M, Inc.

Also uses heat treatment

Cold Working Increases strength at the expense of ductility

Arvind Thekdi - E3M, Inc.

Sales

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

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