dr kresimir rupnik lsu chemistry 1001 part 9 electrons and photons in chemistry-electronic chemical...

DR Kresimir RupnikLSU

CHEMISTRY 1001Part 9

Electrons and PhotonsIn Chemistry-Electronic Chemical Properties (the

Periodic Table)-PT

2005, Intelliome LLC , Prentice Hall

Periodic Trends in the Properties of the Elements

so,the periodic table is actually thetable of electronic structure of atoms of different elements !!!

Chemistry depends on the electronic structure!!!!

http://wps.prenhall.com/wps/media/objects/477/488764/flash/periodic_table/PeriodicTable.html

http://wps.prenhall.com/wps/media/objects/477/488764/flash/periodic_table/PeriodicTable.html

Prentice - Intelliome A. S. LLC 2005 3

Electron Configurations and the Periodic Electron Configurations and the Periodic TableTableThe periodic table can be used as a guide for electron configurations.The period number is the value of n.Groups 1A and 2A have the s-orbital filled.Groups 3A - 8A have the p-orbital filled.Groups 3B - 2B have the d-orbital filled.The lanthanides and actinides have the f-orbital filled.Note that the 3d orbital fills after the 4s orbital. Similarly, the 4f orbital fills after the 5d orbital.


Electron Configurations and the Periodic Electron Configurations and the Periodic TableTable


One of the most obvious atomic property-One of the most obvious atomic property-electron structure dependence: size of the electron structure dependence: size of the atomatom

Start from the Electron Shells in Atoms!!Start from the Electron Shells in Atoms!!

As the principal quantum number increases, the size of the orbital increases.We can plot the 2 (probability of finding an electron) versus distance, r, is called a radial plot of electron density (see electronic structure Ch 3.).

Atomic size varies consistently through the periodic table.

Periodic Trends in atomic PropertiesPeriodic Trends in atomic Properties


Electron Shells and the Sizes of AtomsElectron Shells and the Sizes of AtomsElectron Shells in AtomsElectron Shells in AtomsThe ns orbitals all have the same shape, but have different sizes and different numbers of nodes.Consider: He: 1s2, Ne: 1s2 2s22p6, and Ar: 1s2 2s22p6 3s23p6.The radial electron density is the probability of finding an electron at a given distance.

For He there is only one maximum (for the two 1s electrons).For Ne there are two maxima: one largely for the 1s electrons (close to the nucleus) and one largely for the n = 2 electrons (further from the nucleus).For Ar there are three maxima: one each largely for n = 1, 2, and 3.


Electron Shells and the Sizes of AtomsElectron Shells and the Sizes of Atoms

Electron Shells in AtomsElectron Shells in AtomsConsider a simple diatomic molecule.The distance between the two nuclei is called the bond distance.If the two atoms which make up the molecule are the same, then half the bond distance is called the covalent radius of the atom.


Electron Shells and the Sizes of AtomsElectron Shells and the Sizes of Atoms

Atomic SizesAtomic SizesAs the principle quantum number increases (i.e., we move down a group), the distance of the outermost electron from the nucleus becomes larger. Hence, the atomic radius increases.As we move across the periodic table, the number of core electrons remains constant. However, the nuclear charge increases. Therefore, there is an increased attraction between the nucleus and the outermost electrons. This attraction causes the atomic radius to decrease.


Electron Shells and the Sizes of AtomsElectron Shells and the Sizes of AtomsAtomic Sizes –Atomic radii (Table 4.5 textbook)Atomic Sizes –Atomic radii (Table 4.5 textbook)


Trends in Atomic Size• either volume or radius

treat atom as a hard marble

• Increases down a groupvalence shell farther from nucleuseffective nuclear charge fairly close

• Decreases across a period (left to right)adding electrons to same valence shelleffective nuclear charge increasesvalence shell held closer


Trends in Atomic Size

http://wps.prenhall.com/wps/media/objects/166/170446/PeriodicTrendsAtomicRadii.html


4 p+

2e-

2e-

12 p+

2e-

8e-

2e-

Be (4p+ & 4e-)

Group IIA

Mg (12p+ & 12e-)

Ca (20p+ & 20e-)

16 p+

2e-

8e-

2e-

8e-


Li (3p+ & 3e-)

Period 2

Be (4p+ & 4e-) B (5p+ & 5e-)

6 p+

2e-

4e-

C (6p+ & 6e-)

8 p+

2e-

6e-

O (8p+ & 8e-)

10 p+

2e-

8e-

Ne (10p+ & 10e-)

2e-

1e-

3 p+

2e-

2e-

4 p+

2e-

3e-

5 p+


Covalent Radius, elements 1 - 58

0

50

100

150

200

250

Atomic Number

Rad

ius,

pm


Example 9.6 – Choose the Larger Atom in Each Pair

• C or O

• Li or K

• C or Al

• Se or I


Example 9.6 – Choose the Larger Atom in Each Pair

• C or O

• Li or K

• C or Al

• Se or I?


Property: Ionization Energy of AtomsProperty: Ionization Energy of Atoms

The first ionization energy, I1, is the amount of energy required to remove an electron from a free atom:

Na(g) Na+(g) + e-.The second ionization energy, I2, is the energy required to remove an electron from a gaseous ion:

Na+(g) Na2+(g) + e-.The larger ionization energy, the more difficult it is to remove the electron.There is a sharp increase in ionization energy when a core electron is removed.


Ionization Energy

• minimum energy needed to remove an electron from an atom gas stateendothermic processvalence electron easiest to removeM(g) + 1st IE M1+(g) + 1 e- M+1(g) + 2nd IE M2+(g) + 1 e-

first ionization energy = energy to remove electron from neutral atom; 2nd IE = energy to remove from +1 ion; etc.

http://wps.prenhall.com/wps/media/objects/166/170446/IonizationEnergy.html


Ionization EnergyIonization Energy


Ionization Energy of Elements 1-56

0

500

1000

1500

2000

2500

Elements by Atomic Number

Ion

izati

on

En

erg

y,

kJ/m

ol


Ionization Energy of Group IA

H

LiNa

K Rb Cs

0

200

400

600

800

1000

1200

1400

H Li Na K Rb Cs

Elements by Period Number

Ion

izati

on

En

erg

y,

kJ/m

ol


Covalent Radii of Group IA

1 Hydrogen

2 Lithium

3 Sodium4 Potassium

5 Rubidium

6 Cesium

0

50

100

150

200

250

Hydrogen Lithium Sodium Potassium Rubidium Cesium

1 2 3 4 5 6

Group Number

Ra

dii

, p

m

Ionization Energy, Group IA

H

LiNa

K Rb Cs

0

200

400

600

800

1000

1200

1400

H Li Na K Rb Cs

Group Number

En

erg

y, k

J/m

ol


Ionization Energy of Periods 2 & 3

Li

BeB

C

NO

F

Na

Mg

Al

Si

PS

Cl

Ar

Ne

0

500

1000

1500

2000

2500

Li Be B C N O F Ne Na Mg Al Si P S Cl Ar

Elements by Group Number

Ion

izat

ion

En

erg

y, k

J/m

ol


Ionization EnergyIonization Energy

Periodic Trends in Ionization EnergyPeriodic Trends in Ionization EnergyIonization energy decreases down a group. This means that the outermost electron is more readily removed as we go down a group.As the atom gets bigger, it becomes easier to remove an electron from the most spatially extended orbital.

Ionization energy generally increases across a period.As we move across a period, Zeff increases. Therefore, it becomes more difficult to remove an electron.

Two exceptions: removing the first p electron and removing the fourth p electron.


Ionization EnergyIonization EnergyPeriodic Trends in Ionization EnergyPeriodic Trends in Ionization Energy


Trends in Ionization Energy


Trends in Ionization Energy• as atomic radius increases, the IE generally

decreasesbecause the electron is closer to the nucleus

• 1st IE < 2nd IE < 3rd IE …• 1st IE decreases down the group

valence electron farther from nucleus

• 1st IE generally increases across the periodeffective nuclear charge increases

http://wps.prenhall.com/wps/media/objects/166/170446/PeriodicTrendIonization.html


Example 9.7 – Choose the Atom with the Highest Ionization Energy in Each Pair

• Mg or P

• As or Sb

• N or Si

• O or Cl


Example 9.7 – Choose the Atom with the Highest Ionization Energy in Each Pair

• Mg or P

• As or Sb

• N or Si

• O or Cl?


Metallic Character• how well an element’s properties match the general

properties of a metal• Metals

malleable & ductileshiny, lusterous, reflect lightconduct heat and electricitymost oxides basic and ionic form cations in solution lose electrons in reactions - oxidized

• Nonmetalsbrittle in solid statedullelectrical and thermal insulatorsmost oxides are acidic and molecular form anions and polyatomic anionsgain electrons in reactions - reduced


Trends in Metallic Character


Example 9.8 – Choose the More Metallic Element in Each Pair

• Sn or Te

• Si or Sn

• Br or Te

• Se or I


Example 9.8 – Choose the More Metallic Element in Each Pair

• Sn or Te

• Si or Sn

• Br or Te

• Se or I?


Metals, Nonmetals, and MetalloidsMetals, Nonmetals, and Metalloids

MetalsMetals


Property: Electron AffinitiesProperty: Electron AffinitiesElectron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion:

Cl(g) + e- Cl-(g)

Look at electron configurations to determine whether electron affinity is positive or negative.

The extra electron in Ar needs to be placed in the 4s orbital which is significantly higher in energy than the 3p orbital.


Electron AffinitiesElectron AffinitiesThe added electron in Cl is placed in the 3p orbital to form the stable 3p6 electron configuration.


Properties of Main-Group ElementsProperties of Main-Group Elements


Metals, Nonmetals, and MetalloidsMetals, Nonmetals, and Metalloids

MetalsMetalsMetal character refers to the properties of metals (shiny luster, malleable and ductile, oxides form basic ionic solids, and tend to form cations in aqueous solution).Metal character increases down a group.Metal character decreases across a period.Metals have low ionization energies.Most neutral metals are oxidized rather than reduced.


Active MetalsActive Metals

Group 1A: The Alkali MetalsGroup 1A: The Alkali MetalsAlkali metals are all soft.Chemistry dominated by the loss of their single s electron:

M M+ + e-

Reactivity increases as we move down the group.Alkali metals react with water to form MOH and hydrogen gas:

2M(s) + 2H2O(l) 2MOH(aq) + H2(g)



Group 1A: The Alkali MetalsGroup 1A: The Alkali MetalsAlkali metal produce different oxides when reacting with O2:

4Li(s) + O2(g) 2Li2O(s) (oxide)2Na(s) + O2(g) Na2O2(s) (peroxide)

K(s) + O2(g) KO2(s) (superoxide)Alkali metals emit characteristic colors when placed in a high temperature flame.The s electron is excited by the flame and emits energy when it returns to the ground state.



Group 1A: The Alkali MetalsGroup 1A: The Alkali Metals

Na line (589 nm): 3p 3s transition

Li line: 2p 2s transition

K line: 4p 4s transition



Group 2A: The Alkaline Earth MetalsGroup 2A: The Alkaline Earth MetalsAlkaline earth metals are harder and more dense than the alkali metals.The chemistry is dominated by the loss of two s electrons:

M M2+ + 2e-.Mg(s) + Cl2(g) MgCl2(s)2Mg(s) + O2(g) 2MgO(s)

Be does not react with water. Mg will only react with steam. Ca onwards:

Ca(s) + 2H2O(l) Ca(OH)2(aq) + H2(g)


Important NonmetalsImportant Nonmetals

HydrogenHydrogenHydrogen is a unique element.Most often occurs as a colorless diatomic gas, H2.It can either gain another electron to form the hydride ion, H, or lose its electron to become H+:

2Na(s) + H2(g) 2NaH(s)2H2(g) + O2(g) 2H2O(g)

H+ is a proton.The aqueous chemistry of hydrogen is dominated by H+

(aq).



Group 6A: The Oxygen GroupGroup 6A: The Oxygen GroupAs we move down the group the metallic character increases (O2 is a gas, Te is a metalloid, Po is a metal).There are two important forms of oxygen: O2 and ozone, O3. Ozone can be prepared from oxygen:

3O2(g) 2O3(g) H = +284.6 kJ.Ozone is pungent and toxic.



Group 6A: The Oxygen GroupGroup 6A: The Oxygen GroupOxygen (or dioxygen, O2) is a potent oxidizing agent since the O2- ion has a noble gas configuration.There are two oxidation states for oxygen: 2- (e.g. H2O) and 1- (e.g. H2O2).Sulfur is another important member of this group.Most common form of sulfur is yellow S8.Sulfur tends to form S2- in compounds (sulfides).



Group 7A: The HalogensGroup 7A: The HalogensThe chemistry of the halogens is dominated by gaining an electron to form an anion:

X2 + 2e- 2X-.Fluorine is one of the most reactive substances known:

2F2(g) + 2H2O(l) 4HF(aq) + O2(g) H = -758.7 kJ.

All halogens consists of diatomic molecules, X2.



Group 7A: The HalogensGroup 7A: The HalogensChlorine is the most industrially useful halogen. It is produced by the electrolysis of brine (NaCl):

2NaCl(aq) + 2H2O(l) 2NaOH(aq) + H2(g) + Cl2(g).The reaction between chorine and water produces hypochlorous acid (HOCl) which disinfects pool water:

Cl2(g) + H2O(l) HCl(aq) + HOCl(aq).Hydrogen compounds of the halogens are all strong acids with the exception of HF.


Group 8A: The Noble GasesGroup 8A: The Noble GasesThese are all nonmetals and monatomic.They are notoriously unreactive because they have completely filled s and p sub-shells.In 1962 the first compound of the noble gases was prepared: XeF2, XeF4, and XeF6.To date the only other noble gas compound known is KrF2.

dr kresimir rupnik lsu chemistry 1001 part 9 electrons and photons in chemistry-electronic chemical...

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