Chapter 7
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Chapter 7. The Structure of Atoms and Periodic Trends. Arrangement of Electrons in Atoms. Electrons in atoms are arranged as: Shells (n) Subshells ( l ) Subshell orientation (m l ). Pauli’s Exclusion Principle. discovered in 1925 by Wolfgang Pauli

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Chapter 7

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Chapter 7

Chapter 7

The Structure of Atoms and Periodic Trends


Arrangement of electrons in atoms

Arrangement of Electrons in Atoms

Electrons in atoms are arranged as:

Shells (n)

Subshells (l)

Subshell orientation (ml)


Pauli s exclusion principle

Pauli’s Exclusion Principle

  • discovered in 1925 by Wolfgang Pauli

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

Practice:What are the 4 quantum numbers for

each electron in He?


Aufbau 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

Electron Configuration

To remember the correct filling order for electrons in atoms:


Electron configuration1

Electron Configuration


Writing electron configurations

Example: H atomic number = 1

1

no. of

s

1

electrons

value of l

value of n

Writing Electron Configurations

Two ways to express electron configuration:

1. spdf notation


Writing electron configurations1

Writing Electron Configurations

2. Orbital box notation

spdf notation


Electron configurations

Electron Configurations

Using the Aufbau Principle to determine the electronic configurations of the elements

1st row elements:


Electron configurations1

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 configurations2

Electron Configurations


Electron configurations and quantum numbers

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


Electron configurations and quantum numbers1

Electron Configurations and Quantum Numbers

l

l

The ml and ms are interchangeable


Electron configurations and quantum numbers2

Electron Configurations and Quantum Numbers

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-

Skip the first 18 electrons


Electron configurations and quantum numbers3

Electron Configurations and Quantum Numbers

l

l


Electron configurations and quantum numbers4

Electron Configurations and Quantum Numbers

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


Electron configurations of ions

Electron Configurations of Ions

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

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


Electron configurations of ions1

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

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


More about the periodic table

More About the Periodic Table

Representative Elements

Groups 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


More about the periodic table1

More About the Periodic Table

d-Transition Elements

All have d electrons

-With n s-orbitals

-With n-1 d–orbitals

Have small property variations from row-to-row


More about the periodic table2

More About the Periodic Table

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


More about the periodic table3

More About the Periodic Table

Noble Gases

-Have filled electron shells

-have similar chemical reactivities

-similar electronic structures

He1s2

Ne[He] 2s2 2p6

Ar[Ne] 3s2 3p6

Kr [Ar] 4s2 4p6

Xe[Kr] 5s2 5p6

Rn[Xe] 6s2 6p6


Periodic properties

Periodic Properties

  • Atomic radii describes the relative sizes of atoms

  • Atomic radii increase within

    a column

  • Atomic radii decrease within

    a row


Periodic properties1

Periodic Properties

Example: Arrange these elements based on their atomic radii:

Se, S, O, Te

O < S < Se < Te


Periodic properties2

Periodic Properties

Example: Arrange these elements based on their atomic radii:

P, Cl, S, Si

Cl < S < P < Si


Periodic properties3

Periodic Properties

Electronegativity: 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 properties4

Periodic Properties

Example: Arrange these elements based on their electronegativity:

Se, Ge, Br, As

Ge < As < Se < Br


Periodic properties5

Periodic Properties

Example: Arrange these elements based on their electronegativity:

Be, Mg, Ca, Ba

Ba < Ca < Mg < Be


Periodic properties6

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+ ion

    Atom(g) + energy  Atom+(g) + e-

Example:

Mg(g) + 738kJ/mol  Mg+ + e-


Periodic properties7

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 properties8

Periodic Properties

Ionization Energy:

  • IE2 > IE1

    always takes moreenergy to remove a second electron from an ion

  • IE1 increases to the right

    Important exceptions are Be & Mg, N & P, etc. due to filled and half-filled subshells

  • IE1 decrease down


First ionization energies

First Ionization Energies

He

Ne

F

Ar

N

Cl

C

P

H

Be

O

Mg

S

Ca

B

Si

Li

Al

Na

K


Periodic properties9

Periodic Properties

Example: Arrange these elements based on their first ionization energies:

Sr, Be, Ca, Mg

Sr < Ca < Mg < Be


Periodic properties10

Periodic Properties

Example: Arrange these elements based on their first ionization energies:

Al, Cl, Na, P

Na < Al < P < Cl


Periodic properties11

Periodic Properties

Electron Affinity: Energy absorbed when an electron is added to an atom to form a negative ion

Sign 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 properties12

Periodic Properties

Examples of electron affinity values:

Mg(g) + e- + 231 kJ/mol Mg-(g)

EA = +231 kJ/mol

  • Br(g) + e- Br-(g) + 323 kJ/mol

    • EA = -323 kJ/mol

Increasing ability to

add electrons

decreasing ability

to add electrons


Electron affinity

Electron Affinity

He

Be

B

N

Ne

Mg

Al

Ar

Ca

P

Na

K

H

Li

O

C

Si

S

F

Cl


Periodic properties13

Periodic Properties

Example: Arrange these elements based on their electron affinities:

Al, Mg, Si, Na

Si < Al < Na < Mg


Periodic properties14

Periodic Properties

Ionic Radius: diameter of an atom in its ionized form

-Cations are always smaller


Periodic properties15

Periodic Properties

Anions are always larger


Periodic properties16

Periodic Properties

Cation radii decrease from left to right across a period

  • Increasing nuclear charge attracts the electrons and decreases the radius.


Periodic properties17

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


Ionic radii

Ionic Radii

Active Figure 8.15


Periodic properties18

Periodic Properties

Example: Arrange these elements based on their ionic radii:

Ca2+, K+, Ga3+

K1+ > Ca2+ > Ga3+


Periodic properties19

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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