CHEMISTRY LAB C

• Team of 2 students

• 50 minutes

• Safety RequirementsWear apron or lab coatOSHA approved goggles with indirect vents

• Do not bring reference material or calculators

CHEMISTRY LAB C

• A series of stations with various activitiesCould Include:

Hands-on Activities: ExperimentsInterpretation of Experimental Data

– (graphs, charts, diagrams, data tables, etc.)

Observation of Running Demonstration

Redox Reactions & Periodicity

Periodicity

• Understand the periodic nature of the elements

Demonstrated ConceptuallyPredicting and explaining trends

Demonstrated ExperimentallyCollecting and/or accounting for data

Topics Covered

Physical Properties

Electronic Structure and Bonding

Chemical Properties

Physical Properties

Atomic and Ionic Radii

Ionization Energy

Melting Point

Electronegativity

Electronic Structure

Electron Configuration

Ionic and Covalent Bonding

Charges on Ions

Metallic Properties

Chemical Properties

Precipitation Formation

(Solubility)

Reaction with Acids

Acidity of Oxides

Dmitri Mendeleev

• Periodic Properties

• Arrange Elements According to Properties

• Families have similar properties– All alkali metals react with water– But to different degrees or reactivity

• Predict Ekasilicon between Si and Sn

• Later arranged according to atomic number not mass

Electron Configuration - I

• H 1s 1

• He 1s 2 [He]• Li 1s2 2s 1 [He] 2s 1

• Be 1s2 2s 2 [He] 2s 2

• B 1s2 2s 2 2p 1 [He] 2s 2 2p 1

• C 1s 2 2s 2 2p 2 [He] 2s 2 2p 2

• N 1s 2 2s 2 2p 3 [He] 2s 2 2p 3

• O 1s 2 2s 2 2p 4 [He] 2s 2 2p 4

• F 1s 2 2s 2 2p 5 [He] 2s 2 2p 5

• Ne 1s 2 2s 2 2p 6 [He] 2s 2 2p6 = [Ne]

Electron Configuration - II

• Na [Ne] 3s 1

• Mg [Ne] 3s 2

• Al [Ne] 3s 2 3p 1

• Si [Ne] 3s 2 3p 2

• P [Ne] 3s 2 3p 3

• S [Ne] 3s 2 3p 4

• Cl [Ne] 3s 2 3p 5

• Ar [Ne] 3s 2 3p6 == [Ar]

Order of Electron Filling

7s 7p

6s 6p 6d

5s 5p 5d 5f

4s 4p 4d 4f

3s 3p 3d

2s 2p

1s

Electron Configuration - III• K [Ar] 4s 1

• Ca [Ar] 4s 2 Or this order is OK !• Sc [Ar] 4s 2 3d 1 [Ar] 3d 1 4s 2

• Ti [Ar] 4s 2 3d 2 [Ar] 3d 2 4s 2

• V [Ar] 4s 2 3d 3 [Ar] 3d 3 4s 2

• Cr [Ar] 4s 1 3d 5 • Mn [Ar] 4s 2 3d 5

• Fe [Ar] 4s 2 3d 6 Either order will be OK !• Co [Ar] 4s 2 3d 7 But it’s normally best to• Ni [Ar] 4s 2 3d 8 put the one filling last!!!• Cu [Ar] 4s 1 3d 10 • Zn [Ar] 4s 2 3d 10

Anomalies to Filling

Anomalies to Filling

Orbital Box Diagrams - III Na ArAtomic Number Orbital Box Condensed Electron Element Diagrams(3s&3p) Configuration

11 Na [He] 3s1

12 Mg [He] 3s2

13 Al [He] 3s23p1

14 Si [He] 3s23p2

15 P [He] 3s23p3

16 S [He] 3s23p4

17 Cl [He] 3s23p5

18 Ar [He] 3s23p6

3s

3s

3s

3s

3s

3s

3s

3px

3px

3px

3py

3py

3py

3py

3py

3py

3py

3px

3px

3px

3px

3pz

3pz

3pz

3pz

3pz

3pz

3pz

Orbital Box Diagram - IV : Sc Zn

4s 3d

Z = 21 Sc [Ar] 4s2 3d1

Z = 22 Ti [Ar] 4s 2 3d 2

Z = 23 V [Ar] 4s 2 3d 3

Z = 24 Cr [Ar] 4s1 3d 5

Z = 25 Mn [Ar] 4s 2 3d 5

Z = 26 Fe [Ar] 4s 2 3d 6

Z = 27 Co [Ar] 4s 2 3d 7

Z = 28 Ni [Ar] 4s 2 3d 8

Z = 29 Cu [Ar] 4s 1 3d 10

Z = 30 Zn [Ar] 4s 2 3d 10

Electronic Configuration Ions

• Na 1s 2 2s 2 2p 6 3s 1 Na+ 1s 2 2s 2 2p 6

• Mg 1s 2 2s 2 2p 6 3s 2 Mg+2 1s 2 2s 2 2p6

• Al 1s 2 2s 2 2p 6 3s 2 3p 1 Al+3 1s 2 2s 2 2p 6

• O 1s 2 2s 2 2p 4 O- 2 1s 2 2s 2 2p 6

• F 1s 2 2s 2 2p 5 F- 1 1s 2 2s 2 2p 6

• N 1s 2 2s 2 2p 3 N- 3 1s 2 2s 2 2p 6

Fig. 8.23

Atomic Size

Atomic Size

• Across a rowDiameter Decreases

Electrons added to the same shellMore protons pull in electrons closer

• Down a columnDiameter Increases

Electrons fill into further out shells

Fig. 8.13

Transition Metals

• Across the transition series (d block) the atomic radii initially decrease, then increase.

• Initially, the increase in the nuclear charge decreases the size when d electrons are added into a shell closer than the valence shell.

• Later the increased electron - electron repulsion from many electrons in the d orbitals cause the atomic radii to increase.

Law of Dulong and Petit

• Heat Capacity is the amount of energy needed to raise the temperature of an amount of a substance

• 1819 Pierre Dulong and Alexis PetitProduct of molar mass and heat capacity is a constant for metals

• Heat capacity decreases with molar mass

Ionization Energy

• The energy required to remove an electron from a neutral atom

A + energy A+ + e-

Second Ionization Energy

• The energy required to remove an electron from a +1 cation

A+ + energy A2+ + e-

• Successive ionization energies are greater than earlier ionization energies

Periodicity of First IonizationEnergy (IE1)

Fig. 8.14

Fig. 8.27

Size of Ions

• Size of anions are larger than atomsAdding electrons to an atom increases the size: Higher -/+ ratio

• Size of cations are smaller than atomsRemoving electrons from an atom decreases the size: Lower -/+ ratio and often lose electrons in furthest shell

Crystal Structures

• Ionic Crystals are lattice of large anions with smaller cations inbetween the anions

• (r+ / r-) > 0.732 cations in cubic hole

• 0.732 > (r+ / r-) > 0.414 cations in octahedral holes

• 0.414 > (r+ / r-) cations in tetrahedral holes

Crystal Structures

• CsCl (r+ / r-) = 0.169 nm/0.181nm > 0.732 cations in cubic hole BCC

• NaCl (r+ / r-) = 0.095 nm/0.181nm 0.732 > (r+ / r-) > 0.414

cations in octahedral holes FCC

• ZnS (r+ / r-) = 0.074 nm/0.184nm 0.414 > (r+ / r-)

cations in tetrahedral holes FCC

Electron Affinity

• Energy released when an electron is added to a neutral atom

A + e- A- + energy

(Sometimes defined as energy needed to remove an electron from an anion)

More Negative

Trends in Three Atomic Properties

Fig 8.18

Fig. 9.2

Metals and Nonmetals

• MetalsShiny luster, various colors - mostly silver

Malleable and ductile

Good conductors of heat and electricity

Most metal oxides are basic Na2O(s) + H2O(l) ==> 2 NaOH(aq)

Generally form cations

Metals and Nonmetals

• NonmetalsNo luster, various colors

Usually brittle - some hard, some soft

Poor conductors of heat and electricity

Most nonmetallic compounds are acidicCO2(g) + H2O(l) ==> H2CO3(aq)

Generally form anions or oxyanions

Metalloids (Semimetals)

• Intermediate properties between metals and nonmetals

Some metallic characteristics and some nonmetal characteristics

Some, most notably Si, are electrical semiconductors

Lattice Energy

Li+ (g) + F- (g) ==> LiF (s)

HoLattice of LiF = -1050 kJ

Periodic Trends in Lattice Energy

Electrostatic Force = (C+) (A-) / Distance

• Ionic Size

• Ionic Charge

Melting and Boiling Points of Some Ionic Compounds

Compound mp( oC) bp( oC)

CsBr 636 1300NaI 661 1304MgCl2 714 1412KBr 734 1435CaCl2 782 >1600NaCl 801 1413LiF 845 1676KF 858 1505MgO 2852 3600

Table 9.1 (p. 340)

Electronegativity

• A scale to show the relative attraction of an atom for electrons shared in a bond

• Linus Pauling Scale

Lowest Fr = 0.7

Highest F = 4.0

The Periodic Table of the Elements2.1

0.9 1.5

0.9 1.2

0.8 1.0 1.3

0.8

0.7

0.7

1.0

0.9

1.5 1.6 1.61.5 1.8

1.2

1.1

1.8 1.8 1.9 1.6

1.4 1.6

1.5

1.8

1.7

1.9

1.9

2.2 2.2

2.2

2.2

2.2

1.9

2.4

1.7

1.9

2.0 2.5 3.0 3.54.0

He

Ne

Ar1.5 1.8 2.1 2.5 3.0

1.6 1.8 2.0 2.4 2.8 Kr

Xe

Rn

2.52.1

2.2

1.9

2.01.9

1.81.7

1.81.8

1.1 1.1 1.1 1.1

1.3

1.2 1.2 1.2 1.2 1.2 1.2 1.2 1.2 1.21.3

1.5 1.7 1.3 1.3 1.3 1.3 1.3 1.3 1.31.3 1.5

0.9

1.3 2.2

Electronegativity

1.1

Th Pa U Np No Lr

1.3

Ce Pr Nd Pm Yb Lu

Bond Polarity

Nonpolar Covalent Bonds:Electronegativity Difference is ideally 0 Very small differences are still considered to be mostly covalent bonds, up to about 0.4

Polar Covalent Bonds:Electronegativity Difference measurableHas polar covalent characteristics up to 2.0

Mostly Ionic Bonds: High Electronegativity Differences

Bond Polarity

Cl2 is a nonpolar covalent bond E = (3.0 - 3.0) = 0

HCl is a polar covalent bond E = (3.0 - 2.1) = 0.9

NaCl is a very polar bond - ionic E = (3.0 - 0.9) = 2.1

The Redox Process in Compound Formation

Fig. 4.13

Oxidation-Reduction Reactions

• How can we predict if a oxidation-reduction reaction will occur

• Experimental trials give reactivity relationships

Metal Activity

• Higher activityMore the metal wants to be oxidizedMore the metal wants to gain electronsBetter reducing agent

• Compare to other metals

• Compare to H+ in water and acids

Activity Series

Mg ==> Mg2+ has a higher activity than

Zn ==> Zn2+

Therefore:

Mg + Zn2+ ==> Mg2+ + Zn

and

Zn + Mg2+ ==> No Reaction

Activity Series

Cr ==> Cr3+ has a higher activity than

Ni ==> Ni2+

Therefore:

2Cr + 3Ni2+ ==> 2Cr3+ + 3Ni

and

Ni + Cr3+ ==> No Reaction

Basic and Acidic Oxides

• More ionic oxides formed on left side of periodic table

• If dissolve in water form basic solutions

MO(s) + H2O(l) M+2(aq) + 2 OH-(aq)

Basic and Acidic Oxides

• More covalent oxides formed on right side of periodic table

• If dissolve in water form acidic solutions

MO(g) + H2O (l) H2MO2(aq)

H2MO2(aq) +H2O H3O+(aq) + HMO3-(aq)