Chapter 7

Chapter 7

The Structure of Atoms and Periodic Trends

Electrons in atoms are arranged as:

Shells (n)

Subshells (l)

Subshell orientation (ml)

Arrangement of Electrons in Atoms

Practice: What are the 4 quantum numbers for each electron in He?

discovered in 1925 by Wolfgang Pauli

-No two electrons in an atom can have the same set of 4 quantum numbers

Pauli’s Exclusion Principle

Aufbau Principle

Describes the electron filling order in atoms-electrons are placed in the lowest available energy orbital-the periodic table is a

function of electron configurations for the elements

Electron Configuration

To remember the correct filling order for electrons in atoms:

Electron Configuration

11 s

value of nvalue of l

no. of

electrons

Example: H atomic number = 1

Two ways to express electron configuration:

1. spdf notation

Writing Electron Configurations

Arrowsdepictelectronspin

ORBITAL BOX NOTATIONfor He, atomic number = 2

1s

21 s

Arrowsdepictelectronspin

ORBITAL BOX NOTATIONfor He, atomic number = 2

1s

21 s

2. Orbital box notation

spdf notation

Writing Electron Configurations

Electron ConfigurationsUsing the Aufbau Principle to determine the

electronic configurations of the elements

1st row elements:

22

11

1s He

1s H

ionConfigurat 1s

Electron Configurations

Hund’s rule: electrons fill suborbitals by placing electrons in each suborbital unpaired first with the same spin direction, then the electrons pair

Electron Configurations

NeAr

Ne Cl

Ne S

Ne P

Ne Si

3p s3 Ne Ne Al

s3 Ne Ne Mg

s3 Ne Ne Na

ionConfigurat 3p 3s

18

17

16

15

14

1213

212

111

Electron Configurations and Quantum Numbers

We can write a complete set of quantum numbers for all of the electrons in every element:

– Na– Ca– Fe

111 s3 Ne NeNa

ionConfigurat 3p 3s


electron s 31/2 0 0 3 e 11

electrons p 2

1/2 1 1 2 e 10

1/2 0 1 2 e 9

1/2 1 1 2 e 8

1/2 1 1 2 e 7

1/2 0 1 2 e 6

1/2 1- 1 2 e 5

electrons s 21/2 0 0 2 e 4

1/2 0 0 2 e 3


1/2 0 0 1 e 1

m m n

-th

-th

-th

-th

-th

-th

-th

-th

-rd

-nd

-st

s

l l

The ml and ms are interchangeable


Noble Gas Notation (or short hand notation):

The first 18 electrons in Ca are represented with the preceding noble gas ([Ar])

- we only concern ourselves with the outermost e-

220 4s Ar [Ar]Ca

ionConfigurat 4p 4s 3d

Skip the first 18 electrons


1/2 0 0 4 e 19]Ar[

m m n

-th

-th

s

l l


There is only one set of 4 quantum numbers for each of the 26 electrons in Fe:

– To save space, we use the symbol [Ar] to represent the first 18 electrons in Fe

6226 3d 4s Ar Ar Fe

ionConfigurat 4p 4s 3d


Electrons are removed from subshell of highest energy level (n-level)

P0 [Ne] 3s2 3p3 -3e- ---> P3+ [Ne] 3s2 3p0

1s

2s

3s3p

2p

1s

2s

3s3p

2p

Electron Configurations of Ions

For transition metals, remove the highest s-orbital electrons first:

Fe [Ar] 4s2 3d6

-2 electrons Fe2+ [Ar] 3d6

-3 electrons

Fe3+ [Ar] 3d5

Electron Configurations of Ions

To form cations, always remove electrons of highest n value first!

More About the Periodic Table

Representative ElementsGroups IA, IIA, IIIA-VIIIA– These elements will have

their “outermost” electron in an outer s or p orbital

– Variations in their properties are similar from top-to-bottom


d-Transition ElementsAll have d electrons

-With n s-orbitals -With n-1 d–orbitals

Have small property variations from row-to-row


f - transition metals -Sometimes called

inner transition metals

-Electrons are being added to f orbitals

Extremely small variations in properties from one element to another


Noble Gases-Have filled electron shells-have similar chemical reactivities-similar electronic structures

He 1s2

Ne [He] 2s2 2p6

Ar [Ne] 3s2 3p6

Kr [Ar] 4s2 4p6

Xe [Kr] 5s2 5p6

Rn [Xe] 6s2 6p6

Periodic PropertiesPeriodic Properties

• Atomic radii describes the relative sizes of atoms

• Atomic radii increase within a column

• Atomic radii decrease within a row

Periodic Properties

Example: Arrange these elements based on their atomic radii:Se, S, O, Te

O < S < Se < Te

Periodic Properties

Example: Arrange these elements based on their atomic radii:P, Cl, S, Si

Cl < S < P < Si

Periodic PropertiesElectronegativity: measure of the tendency of

an atom to attract electrons to itself-Fluorine is the most electronegative element-Cesium is the least electronegative element

Electronegativity increase from left-to-right and decrease from top-to-bottom

increase

decrease

Periodic Properties

Example: Arrange these elements based on their electronegativity:Se, Ge, Br, As

Ge < As < Se < Br

Periodic Properties

Example: Arrange these elements based on their electronegativity:Be, Mg, Ca, Ba

Ba < Ca < Mg < Be

Periodic Properties

Ionization Energy: energy required to remove an electron from an atom in the gas state

First ionization energy (IE1) – Energy required to remove the first electron

from an atom in the gas state to form a 1+ ionAtom(g) + energy Atom+

(g) + e-

Example: Mg(g) + 738kJ/mol Mg+ + e-

Periodic Properties

Second ionization energy (IE2)– The amount of energy required to

remove the second electron from a gaseous 1+ ion

Atom+ + energy Atom2+ + e-

Mg+ + 1451 kJ/mol Mg2+ + e-

- Atoms can have 3rd (IE3), 4th (IE4), etc. - Each IE is significantly higher than the

previous IE

Periodic PropertiesIonization Energy:

1. IE2 > IE1

always takes more energy to remove a second electron from an ion

2. IE1 increases to the rightImportant exceptions are Be & Mg, N & P, etc. due to filled and half-filled subshells

3. IE1 decrease down

First Ionization Energies

0

500

1000

1500

2000

2500

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

Atomic Number

Ionization Energy (kJ/mol)

H

He

Li

Be

B

C

N

O

F

Ne

Na

Mg

AlSi

P

S

Cl

Ar

K

Ca

Periodic Properties

Example: Arrange these elements based on their first ionization energies:Sr, Be, Ca, Mg

Sr < Ca < Mg < Be

Periodic Properties

Example: Arrange these elements based on their first ionization energies:Al, Cl, Na, P

Na < Al < P < Cl

Periodic Properties

Electron Affinity: Energy absorbed when an electron is added to an atom to form a negative ionSign conventions for electron affinity:– If electron affinity > 0 energy is absorbed– If electron affinity < 0 energy is released

Electron affinity is the measure of an atom’s ability to form negative ions

atom(g) + e- + EA atom-

(g)

Periodic Properties

Mg(g) + e- + 231 kJ/mol Mg-(g)

EA = +231 kJ/mol

Examples of electron affinity values:

Increasing ability to add electrons

decre

asin

g a

bilit

y

to a

dd

ele

ctr

on

s

Br(g) + e- Br-(g) + 323 kJ/mol

EA = -323 kJ/mol

Electron Affinities of Some Elements

-400-350-300-250-200-150-100-50

0

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20

Atomic Number

Ele

ctro

n A

ffin

ity

(kJ/

mo

l)

Electron Affinity

H

He

Li

Be B

C

N

O

F

NeNa

Mg Al

Si

P

S

Cl

ArK

Ca

Periodic Properties

Example: Arrange these elements based on their electron affinities:Al, Mg, Si, Na

Si < Al < Na < Mg

Periodic PropertiesIonic Radius: diameter of an atom in its ionized form

-Cations are always smaller

Element Li Be

Atomic Radius (Å)

1.52 1.12

Ion Li+ Be2+

Ionic Radius (Å)

0.90 0.59

Periodic Properties

Anions are always larger

Element N O F

AtomicRadius(Å)

0.75 0.73 0.72

Ion N3- O2- F1-

IonicRadius(Å)

1.71 1.26 1.19

Periodic Properties

Cation radii decrease from left to right across a period– Increasing nuclear charge attracts the electrons and decreases

the radius.

Ion Rb+ Sr2+ In3+

IonicRadii(Å)

1.66 1.32 0.94

Periodic Properties

Anion radii decrease from left to right across a period– Increasing electron numbers in highly charged ions cause the

electrons to repel and increase the ionic radius

Ion N3- O2- F1-

IonicRadii(Å)

1.71 1.26 1.19

Active Figure 8.15Active Figure 8.15

Ionic Radii

Periodic Properties

Example: Arrange these elements based on their ionic radii:

Ca2+, K+, Ga3+

K1+ > Ca2+ > Ga3+

Periodic Properties

Example: Arrange these elements based on their ionic radii:Cl-1, Se-2, Br-1, S-2

Cl1- < S2- < Br1- < Se2-

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Chapter 7