Electron configuration

High School Chemistry Rapid Learning Series - 13

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Atomic Structure andAtomic Structure and Electron Configuration

HS Ch i t R id L i S i

Rapid Learning Centerwww.RapidLearningCenter.com/© Rapid Learning Inc. All rights reserved.

HS Chemistry Rapid Learning Series

Wayne Huang, PhDKelly Deters, PhDRussell Dahl, PhD

Elizabeth James, PhD



Learning Objectives

Basic structure of atoms.How to determine the

By studying this tutorial you will learn…

How to determine the number of electrons.How to place electrons in energy levels, subshells and orbitals.How to show electron configurations using three

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configurations using three methods.How to write and understand Quantum Numbers.

Concept Map

Chemistry

Studies

Previous content

New content

Matter

Studies

Atoms

Made of

Electrons

Quantum Numbers

Chemical properties determined by

Location described by

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Boxes and Arrows SpectroscopicNotation

Noble GasNotation

3 ways to show configurations



Atomic Structure

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Definition: Atom

Atom - Smallest piece p(basic unit) of matter that has the chemical properties of the element.

Often called the Graphical Rendering of an Atom

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Often called the“Building Block of Matter”.

p g

Protons

Neutrons

Electrons



What’s in an Atom?An atom is made of three sub-atomic particles.

Particle Location Mass

1

Charge

Nucleus

Nucleus

Outside the nucleus

1 amu = 1.67×10-27 kg

1 amu = 1.67×10-27 kg

0.00055 amu9.10×10-31 kg

+1

0

-1

Proton

Neutron

Electron

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1 amu (“atomic mass unit”) = 1.66 × 10-27 kg

The AtomNucleus

Ch

Electron Cloud

M Very small relative mass

Charge = - (# of

electrons)

Charge = # of protons

Mass = # of protons

+ # of neutrons

Overall Charge = # of protons

-# f l t

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# of electrons

Overall Mass = # of protons

+ # of neutrons



Protons Versus Electrons

Protons Electrons

+ Charge - Charge

Found in nucleus.

# determines the “identity” of the atom (atomic number).

Found outside nucleus.

# and configuration determine how the atom will react.

Contributes to mass of atom.

Not contribute significantly to mass of atom.

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Cannot be lost or gained without changing which element it is (nuclear reaction).The ratio of protons to electrons determines the charge on the atom (since neutrons are “neutral”).

Can be lost or gained—results in an atom with a charge (ion).

Electron Locations

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Definition: Electron Cloud

Electron cloud – It is the area outside of thethe area outside of the nucleus where the electrons reside (i.e. the probability of finding electrons).

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Electron Clouds

Electron cloud

Principal energy levels

Subshells

The electron cloud is made of energy levels (n).

Energy levels are

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Subshellscomposed of subshells (l).

Subshells have orbitals (ml).



Definition: Subshell and Orbital

Subshell – A set of orbitals with equal energy.gy

Orbital – Area of probability of an electron being located.

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Each orbital can hold 2 electrons (spin up and down).

Types of Subshells

Begins inNumber of Total number

There are 4 types of subshells that electrons reside in under ordinary circumstances.

Subshell Begins in energy level

equal energy orbitals

of electrons possible

s

p 2

1

3

1

6

2

gy In

crea

ses

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d

f

3

4

5

7

10

14

Ener

g

Subshell Mnemonic: spdf = Smart People Don’t Fail.



Pictures of Orbitals

1 s orbital1 s orbital

3 p orbitals

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5 d orbitals

Electron Configuration

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Definition: Electron Configurations

Electron Configurations –Shows the grouping and g p gposition of electrons in an atom.

Since the number of electrons and their configuration determines the chemical properties of the atom, it is important to understand them.

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Box (and Arrow) Notation: Electron configurations use boxes for orbitals and arrows for electrons.

Aufbau Principle

Aufbau (building-up) Principle: Electrons must fill subshells (and orbitals) so that the total energy of 1

The first of 3 rules that govern electron configurations:

( ) gyatom is at a minimum.

1

What does this mean?

Electrons must fill the lowest available subshells and orbitals

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before moving on to the next higher energy subshell/orbital.



Energy and SubshellsThe energy diagram below shows the relative energy levels.

6p5d 4f

3s

4s

5s

3p

4p

5p

3d

4d6s

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2s

2p

Ener

gy

Subshells are filled from the lowest energy level (1s) to increasing energy levels (follow the arrows).

Not that this does not always go in numerical order.

Hund’s Rule

Hund’s Rule: Place electrons in unoccupied

The second of 3 rules that govern electron configurations.

Hund s Rule: Place electrons in unoccupied orbitals of the same energy level (spin up) before doubling up.

2

How does this work?

If you need to add 3 electrons to a p subshell

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If you need to add 3 electrons to a p subshell, add 1 to each (in parallel spins) before beginning to double up.



Pauli Exclusion Principle

Pauli Exclusion Principle: Two electrons that th bit l t h diff t i

3

The last of 3 rules that govern electron configurations.

occupy the same orbital must have different spins.

“Spin” describes the angular momentum of the electron.

“Spin” is designated with an up or down arrow.

How does this work?SpinUp Spin

Down

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How does this work?

If you need to add 4 electrons to a psubshell, you’ll need to double up. When you double up, make them opposite spins.

Down

Determining the Number of ElectronsIn order to properly construct an electron configuration, you must be able to determine how many electrons to use.

Br1- Charge = -1

Charge = # of protons – # of electrons

Atomic number = # of protons

Example: How many electrons does the following have?

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-1 = 35 - Electrons

Atomic number for Br = 35 = # of protons

Electrons = 36 35Br1-



Another ExampleIn order to properly construct an electron configuration, you must be able to determine how many electrons to use.

No charge written Charge is 0Cl

Charge = # of protons – # of electrons

Atomic number = # of protons

Example: How many electrons does the following have?

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Electrons = 17

0 = 17 - Electrons

Atomic number for Cl = 17 = # of protons

17Cl

Applying the Rules

Aufbau Principle: Electrons must fill subshells (and orbitals) so that the total energy of atom is at a minimum.1

Use the 3 rules of electron configurations.

Hund’s Rule: Place electrons in unoccupied orbitals of the

Example: Give the electron configuration for a Cl atom. No charge written Charge is 0

17Cl Atomic number for Cl = 17 = # of protons

Pauli Exclusion Principle: Two electrons that occupy the same orbital must have different spins.3

Hund s Rule: Place electrons in unoccupied orbitals of the same energy level before doubling up.2

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0 = 17 - Electrons Electrons = 17Place 17 electrons

1s 2s 2p 3s 3p

4231567910111213141516178

Electron Configuration Rules Mnemonic: Aufbau (stays low); Hund (does not double up); Pauli (spins up and down) = “Alligator stays low; Hippo does not pair up and Penguin jumps up and down.”



Spectroscopic Notation

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Definition: Spectroscopic Notation

Spectroscopic Notation – Shorthand way of showing electron configurations.g g

The number of electrons in a subshell are shown as a superscript after the subshell designation.

Box

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1s 2s 2p 3s 3p

1s2 2s2 2p6 3s2 3p5 Spectroscopic Notation

BoxNotation

1S2

Principal Energy Level (n)Subshell (l)

Number of Electrons



Writing Spectroscopic NotationDetermine the number of electrons to place.1

Follow Aufbau’s Principle for filling order.2Fill in subshells until they reach their max (s = 2, p = 6, d = 10, f = 14).3The total of all the superscripts is equal to the number of electrons.4

Example: Give the spectroscopic notation for S.

No charge written Charge is 0S

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0 = 16 - Electrons

No charge written Charge is 0

16S Atomic number for S = 16 = # of protonsElectrons = 16

Place 16 electrons

1s 2s 2p 3s 3p2 2 6 2 4

2 2 6 2 4+ + + + = 16

Electron Configurations and the Periodic Table

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1 2 2 2 2 5

Configurations Within a GroupLook at the electron configurations for the Halogens (Group 7).

F 1s2 2s2 2p5F

Cl

Br

1s2 2s2 2p6 3s2 3p5

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5

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I 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5

All of the elements in Group 7 end with 5 electrons in a p subshell.

Configurations and Periodic TableIn fact, every Group ends with the same number of electrons in the highest energy subshell.Each area of the periodic table is referred to by the hi h t b h ll th t t i l t

s1 s2

p1 p2 p3 p4 p5 p6

highest energy subshell that contains electrons.

d-block

p-blocks-block

IA

IIA

IIIB IVB VB VIB VIIB VIIIB VIIIB VIIIB IB IIB

IIIA IVA VA VIA VIIA

VIIIAGroup

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d1 d2 d3 d4 d5 d6 d7 d8 d9 d10

p1 p2 p3 p4 p5 p6

f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14f-block



Wondering how to remember the order of filling of the subshells? Just use the periodic table as a mnemonic device.

Periodic Table as a Road-Map - 1

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In order to do this, the “f” block needs to be placed in atomic order.(It’s usually written below to fit it on the paper).

To see the filling order of subshells, read from left to right, top to bottom!

Periodic Table as a Road-Map - 2

1s 1s

This tool shows that the 3d energy level is filled after the 4s energy level!

2p3p4p5p6p

3d4d5d6d

4f5f

1s2s3s4s5s6s7s

1s

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p subshells begin in level 2, so begin the p-block with “2p”.

s subshells begin in level 1, so begin the s-block with “1s”.

d subshells begin in level 3, so begin the d-block with “3d”.f subshells begin in level 4, so begin the f-block with “4f”.



Another Tool for Filling Order

1s

There is another tool (mnemonic device) commonly used to remember orbital filling order.

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

Building-Up Principle:To read the chart, start with 1s and follow the arrows. Move down one diagonal as far as possible, then jump to the top of the next di l d k

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5s 5p 5d 5f

6s 6p 6d

7s 7p

8s

diagonal and keep going.

ElectronElectron Configurations of Ions

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Definition: Ion

Ion – an atom (or group of atoms) that has gained or ) glost electrons resulting in a net charge.

Atoms gain and lose electrons to be in a more stable state.

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Usually, the “more stable state” is a full valence shell.

Outermost shell of electrons

Look at the electron configurations for the following (#p = # of protons and #e = # of electrons):

Full Valence Shell Ions

1s 2s 2p2 2 6

Br-

O2-

1s 2s 2p 3s 3p2 2 6 2 6 4s 2 3d 10 4p 6

#p = 35 -1 = 35 - #e #e = 36

#p = 8 -2 = 8 - #e #e = 10

Charge = p-e

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Na+

Ca2+

1s 2s 2p 3s 3p2 2 6 2 6

#p = 11 +1 = 11 - #e #e = 10

#p = 20 +2 = 20 - e #e = 18

1s 2s 2p2 2 6



What do you notice about each of these configurations?

Full Valence Shell Ions

They all end with full p subshells.

Notice that O2- and Na+ have the same number and configuration of

1s 2s 2p2 2 6

Br -

O2-

1s 2s 2p 3s 3p2 2 6 2 6 4s2 3d 10 4p 6

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electrons.

Na+

Ca2+ 1s 2s 2p 3s 3p2 2 6 2 6

1s 2s 2p2 2 6 This makes them isoelectric.

Noble Gas Configuration

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Definition: Noble Gas Notation

Noble Gas – Group 8 of the Periodic Table. They contain full valence shells.

Noble Gas Notation – Noble gas is used to represent the core (inner) electrons and only the valence shell is shown.

35Br

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1s 2s 2p 3s 3p2 2 6 2 6 4s 2 3d 10 4p 5

4s 2 3d 10 4p 5[Ar]

35BrSpectroscopic Notation:

Noble Gas Notation:

The “[Ar]” represents the core electrons and only the valence electrons are shown.

How do you know which noble gas to use to symbolize the core electrons?

Which Noble Gas Do You Choose?

Think: Price is Right.

H d i th P i i Ri ht?How do you win on the Price is Right?

By getting as close as possible without going over.

Choose the noble gas that’s closest without going over!

Noble Gas # of electrons

He 2

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Ne

Ar

Kr

Xe

10

18

36

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How do you know where to start off after using a noble gas?Use the periodic table!

Where Does the Noble Gas Leave Off?

2p3p4p5p6p

3d4d5d6d

4f5f

1s2s3s4s5s6s7s

HeNeArKrXeRn

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6d5f7s

The noble gas fills the subshell that it’s at the end of.

Begin filling with the “s” subshell in the next row to show valence electrons.

Noble Gas Notation Example

Determine the number of electrons to place.1

Determine which noble gas to use.2Start where the noble gas left off and write spectroscopic notation for the valence electrons.

3

Example: Give the noble gas notation for As.

No charge written Charge = 0A

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0 = 33 - electrons33As Atomic number for As = 33 = # of protons.

Electrons = 33 Place 33 electrons.

[Ar] 4s 3d 4p2 10 3 18 2 10 3+ + = 33

Closest noble gas: Ar (18)Ar (1s22s22p63s23p6) is full up through 3p.



Comparing the Different Notations

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Pros and Cons of Each NotationEach notation has it’s advantages and disadvantages.

Pro Con

Shows if electrons are paired or

unpaired.

Quicker than “Boxes and arrows”.

Longest method.

Does not show pairing of electrons.

“Boxes and arrows”

Spectroscopic Notation

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Allows focus on the valence electrons

(that control bonding).

Quickest method.

Does not show core electrons.

Does not show pairing of electrons.

Noble Gas Notation



Exceptions to the Aufbau Rule

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Stability of d Subshells with 5 or 10d subshells have 5 orbitals…They can hold 10 electrons.

According to the Aufbau principle, Cr should have the following valence electron configuration:

4s2 3d4

But a half-full or completely full d subshell is more stable

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than the above configuration, so it is:4s1 3d5



Elements with ExceptionsThe following elements are excepts to the Aufbau Principle:

Element Should be Actually is

4s2 3d4

5s2 4d4

6s2 5d4

4s2 3d9

5s2 4d9

4s1 3d5

5s1 4d5

6s1 5d5

4s1 3d10

5s1 4d10

Cr

Mo

W

Cu

Ag

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g

6s2 5d9 6s1 5d10Au

They are the two groups on the periodic table that begin with Cr and Cu.

Quantum Numbers

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Definition: Quantum Numbers

Quantum Numbers – A set of 4 numbers (n, l, ml, & ms) that describes the electron’s placement in the atom.

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4 Quantum Numbers

2, 1, -1, +½

n ml

2p1 0 +1

n = 2

or

ms = ½(up)

Quantum Number

Symbol

n

Describes

Shell Number (Size)

S b h ll

Possible Numbers

Whole # ≥ 1Principal

A i th l

l ms

-1 0 +1

l = 1ml = -1

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l

ml

ms

Subshell Type (Shape)

Whole # < n(0 n-1)

-l +l

+½ or –½

Azimuthal (Angular)

Magnetic

Spin

Orbital (Orientation)Spin (up or down spin)



Determining Quantum Numbersn: principal energy level

l: subshell s = 0

Give the number of the shell.

4p 3

l: subshell s = 0p = 1d = 2f = 3

ml: orbital 0s -1 0 1p -2 -1 0 1 2d

-3 -2 -1 0 1 2 3f

Coding system: 0,1… n-1.

Number-line system of identifying orbitals.0 is always in the middle.Number line from l to + l

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Number line from –l to + l.

ms: spin

Coding system↑ = + ½ (spin up)↓ = - ½ (spin down)

Quantum Number ExamplesGive the quantum numbers for the red arrow.Example:

1s 2s 2p 3s 3p

It’s in level “3” 0It s in level 3 .

___, ___, ___, ___3

It’s in subshell “s” - the “code” for “s” is “0”.

0It’s in orbital “0”.

0It’s a down arrow. -½

Give the quantum numbers for the red arrow.Example:

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1s 2s 2p 3s 3p

It’s in level “2”.

___, ___, ___, ___2

It’s in subshell “p”—the “code” for “p” is “1”.

1It’s in orbital “-1”.

-1It’s an up arrow. +½

-1 0 +1



Identifying Incorrect Quantum Numbers

Example: What’s wrong with the following sets of quantum numbers?

1, 1, 0, +½ n = 1…OK as n (energy level) can be any whole # > 0l = 1…subshell is “p”, but if n = 1 so l must be 0 (i.e. s subshell).

There is no p subshell in energy level 1

2, 1, -2, -½

There is no p subshell in energy level 1.

n = 2…OK as n can be any whole # >0l = 1…subshell is “p”.

OK as level 2 has “p”, i.e. “2p”.ml = -2…on the “-2” orbital

“p” subshell has 3 orbitals: ___ ___ ___-1 0 +1

No “-2” orbital in a “p” subshell. ml must be between –l and l (i.e. -1, 0, +1), not -2.

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1, 0, 0, -1

ml must be between l and l (i.e. 1, 0, 1), not 2.

n = 1…OK as n can be any whole # >0l = 0…subshell is “s”.

OK as level 1 has an “s”.ml = 0…on the “0” orbital

OK as “s” has 1 orbital and it’s “0”.ms = -1

ms must be either +½ or -½, not -1.

Electron configurations can

be shown with

Electron configurations can

be shown with

Atoms are made of protons, neutrons

and electrons. The configuration of the

Atoms are made of protons, neutrons

and electrons. The configuration of the

Quantum numbersdescribe the

location of an

Quantum numbersdescribe the

location of an

Learning Summary

ff

boxes and arrows, in spectroscopic notation, or noble

gas notation.

boxes and arrows, in spectroscopic notation, or noble

gas notation.

gelectrons

determines the chemical properties

of the atom.

gelectrons

determines the chemical properties

of the atom.

ocat o o aelectron in an atom and are a series of

4 numbers.

ocat o o aelectron in an atom and are a series of

4 numbers.

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Electron configurations are written following the Aufbau principle, Hund’s

Rule and the Pauli Exclusion Principle.

Electron configurations are written following the Aufbau principle, Hund’s

Rule and the Pauli Exclusion Principle.

Electrons are organized in levels, subshells and

orbitals.

Electrons are organized in levels, subshells and

orbitals.



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Atomic Structure and Electron ConfigurationElectron Configuration

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Electron configuration