Atomic Structure and Electron Configuration - mrgssciencepage

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Atomic Structure and Electron Configuration Rapid Learning Core Tutorial Series

Wayne Huang, PhD Kelly Deters, MA Russell Dahl, PhD

Rapid Learning Center www.RapidLearningCenter.com/ © Rapid Learning Inc. All rights reserved.

© Rapid Learning Inc. All rights reserved.

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Objectives By studying this tutorial you will learn… Basic structure of atoms How to determine the number of electrons How to place electrons in energy levels, subshells and orbitals How to show electron configurations using three methods How to write and understand Quantum Numbers 3/56

Electron Configuration Concept Map Previous Previouscontent content Chemistry Chemistry

New Newcontent content

Studies

Quantum QuantumNumbers Numbers Matter Matter Location described by Made of

Electrons Electrons

Chemical properties determined by

Atoms Atoms

3 ways to show configurations

Boxes Boxesand andArrows Arrows

Spectroscopic Spectroscopic Notation Notation

Noble NobleGas Gas Notation Notation

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Atomic Structure

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Definition: Atom

Atom - n. smallest piece of matter that has the chemical properties of the element.

Often called the “Building Block of Matter”

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

Location

Mass

Charge

Proton

Nucleus

1 amu = 1.67×10-27 kg

+1

Neutron

Nucleus

1 amu = 1.67×10-27 kg

0

Electron

Outside the nucleus

0.00055 amu 9.10×10-31 kg

-1

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

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The Atom Electron cloud

Nucleus

Mass = # of protons + # of neutrons

Charge = # of protons

Charge = - (# of electrons)

Very small relative mass

Overall Charge = # of protons - (# of electrons)

Overall Mass = # of protons + # of neutrons 8/56


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Protons Versus Electrons Protons

Electrons

+ Charge

- Charge

Contributes to mass of atom

Does not contribute significantly to mass of atom

Found in nucleus

Found outside nucleus

# determines the “identity” of the atom

# and configuration determine how the atom will react

Cannot be lost or gained without changing which element it is (nuclear reaction)

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

The ratio of protons to electrons determines the charge on the atom. 9/56

Electron Locations

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Definition: Electron Cloud Electron cloud – It is the area outside of the nucleus where the electrons reside.

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Electron Clouds Electron cloud

Principle energy levels

The electron cloud is made of energy levels.

Subshells

Energy levels are composed of subshells.

Orbitals

Subshells have orbitals.

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Definition: Subshell and Orbital

Subshell – A set of orbitals with equal energy.

Orbital – Area of probability of the electron being located. Each orbital can hold 2 electrons.

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Types of Subshells

Energy increases

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

Begins in energy level

Number of equal energy orbitals

Total number of electrons possible

s

1

1

2

p

2

3

6

d

3

5

10

f

4

7

14

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Pictures of Orbitals

s orbital

3 p orbitals

5 d orbitals 15/56

Electron Configuration

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Definition: Electron Configurations

Electron Configurations – Shows the grouping and position 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.

Electron configurations use boxes for orbitals and arrow for electrons. 17/56

Aufbau Principle The first of 3 rules that govern electron configurations

1

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

What does this mean? Electrons must fill the lowest available subshells and orbitals before moving on to the next higher energy subshell/orbital.

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Energy and Subshells The energy diagram below shows the relative energy levels. 6p 6s

5p

5d

4f

4d

5s 4p 3d

4s 3p 3s 2p

Energy

2s

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Subshells are filled from the lowest energy level to increasing energy levels. Not that this does not always go in numerical order.

1s

Hund’s Rule The second of 3 rules that govern electron configurations.

2

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

How does this work? If you need to add 3 electrons to a p subshell, add 1 to each before beginning to double up.

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Pauli Exclusion Principle The last of 3 rules that govern electron configurations. Pauli Exclusion Principle: Two electrons that occupy the same orbital must have different spins.

3

“Spin” describes the angular momentum of the electron “Spin” is designated with an up or down arrow.

How does this work? If you need to add 4 electrons to a p subshell, you’ll need to double up. When you double up, make them opposite spins.

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Determining the Number of Electrons In order to properly construct an electron configuration, you must be able to determine how many electrons to use. Charge = # of protons – # of electrons Atomic number = # of protons

Example: Br1-

How many electrons does the following have? Charge = -1 Atomic number for Br = 35 = # of protons -1 = 35 - electrons Electrons = 36

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Another Example In order to properly construct an electron configuration, you must be able to determine how many electrons to use. Charge = # of protons – # of electrons Atomic number = # of protons

Example:

How many electrons does the following have? No charge written Æ Charge is 0

Cl

Atomic number for Cl = 17 = # of protons 0 = 17 - electrons Electrons = 17 23/56

Applying the Rules Use the 3 rules of electron configurations.

1

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

2

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

3

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

Example:

Give the electron configuration for a Cl atom No charge written Æ Charge is 0

Cl

Atomic number for Cl = 17 = # of protons 0 = 17 - electrons Place 17 electrons 1s

2s

Electrons = 17

4 9 70 6 5 3 2 1 8 17 16 15 14 13 12 11 2p

3s

3p

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Spectroscopic Notation

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Definition: Spectroscopic Notation Spectroscopic Notation – Shorthand way of showing electron configurations. The number of electrons in a subshell are shown as a superscript after the subshell designation.

1s

2s

2p

3s

3p

1s2 2s2 2p6 3s2 3p5 26/56


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Writing Spectroscopic Notation 1

Determine the number of electrons to place.

2

Follow Aufbau’s Principle for filling order.

3

Fill in subshells until they reach their max (s = 2, p = 6, d = 10, f = 14).

4

The total of all the superscripts is equal to the number of electrons.

Example: S

Give the spectroscopic notation for S. No charge written Æ Charge is 0 Atomic number for S = 16 = # of protons Electrons = 16 0 = 16 - electrons

Place 16 electrons

2 + 2 + 6 + 2 + 4 = 16 1s 2 2s 2 2p 6 3s 2 3p 4

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Electron Configurations and the Periodic Table

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Configurations Within a Group Look at the electron configurations for the Halogens (Group 7). 1s2 2s2 2p5

F Cl

1s2 2s2 2p6 3s2 3p5

Br

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

I

All of the elements in Group 7 end with 5 electrons in a p subshell. 29/56

Configurations and the Periodic Table In 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 highest energy subshell that contains electrons. p-block s-block d-block s1 s2

p1 p2 p3 p4 p5 p6 d1 d2 d3 d4 d5 d6 d7 d8 d9 d10

f-block

f1

f2

f3

f4

f5

f6

f7

f8

f9 f10 f11 f12 f13 f14

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Periodic Table as a Road-Map Wondering how to remember the order of filling of the subshells? Just use the periodic table.

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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)

Periodic Table as a Road-Map To see the filling order of subshells, read from left to right, top to bottom! This tool shows that the 3d energy level is filled after the 4s energy level! 1s

1s 2p

2s

3p

3s 4s

3d

4p

5s

4d

5p 6p

6s

4f

5d

7s

5f

6d

s subshells begin in level 1, so begin the s-block with “1s” p subshells begin in level 2, so begin the p-block with “2p” 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” 32/56


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Another Tool for Filling Order There is another tool commonly used to remember orbital filling order. 1s

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2s

2p

3s

3p

3d

4s

4p

4d

4f

5s

5p

5d

5f

6s

6p

6d

7s

7p

To read the charge, move down one diagonal as far as possible, then jump to the top of the next diagonal and keep going.

8s

Electron Configurations of Ions

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Definition: Ion

Ion – an atom that has gained or lost electrons resulting in a net charge. Atoms gain and lose electrons to be in a more stable state. Usually, the “more stable state” is a full valence shell. Outermost shell of electrons

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Full Valence Shell Ions Look at the electron configurations for the following: Br-1

p = 35

-1 = 35 - e

e = 36

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 O2-

p=8

-2 = 8 - e

e = 10

+1 = 11 - e

e = 10

+2 = 20 - e

e = 18

1s 2 2s 2 2p 6 Na+

p = 11 1s 2 2s 2 2p 6

Ca2+

p = 20

1s 2 2s 2 2p 6 3s 2 3p 6 36/56


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Full Valence Shell Ions What do you notice about each of these configurations? They all end with full p subshells. Br-1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 O21s 2 2s 2 2p 6 Na+ 1s 2 2s 2 2p 6

Notice that O2- and Na+ have the same number and configuration of electrons. This makes them isoelectric.

Ca2+ 1s 2 2s 2 2p 6 3s 2 3p 6 37/56

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. Br Spectroscopic

1s 2

2s 2

2p 6

3s 2

3p 6 4s 2 3d 10 4p 5 [Ar] 4s 2 3d 10 4p 5

Noble gas

The “[Ar]” represents the core electrons and only the valence electrons are shown. 39/56

Which Noble Gas Do You Choose? How do you know which noble gas to use to symbolize the core electrons? Think: Price is Right. 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!

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Noble Gas

# of electrons

He

2

Ne

10

Ar

18

Kr

36

Xe

54

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Where Does the Noble Gas Leave Off? How do you know where to start off after using a noble gas? Use the periodic table! 1s

He

2s 3s

2p

Ne

3p

Ar

4s

3d

4p

Kr

5s

4d

5p

Xe

6p

Rn

6s

4f

5d

7s

5f

6d

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. 41/56

Noble Gas Notation Example 1

Determine the number of electrons to place.

2

Determine which noble gas to use.

3

Start where the noble gas left off and write spectroscopic notation for the valence electrons.

Example: As

Give the noble gas notation for As. No charge written Æ Charge is 0 Atomic number for As = 33 = # of protons 0 = 33 - electrons Electrons = 33 Place 33 electrons

Closest noble gas: Ar (18) [Ar] 4s 2 3d 10 4p 3

Ar is full up through 3p 18 + 2 + 10 + 3 = 33

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Comparing the Different Notations

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

Con

“Boxes and arrows”

Shows if electrons are paired or unpaired

Longest method

Spectroscopic notation

Quicker than “Boxes and arrows”

Does not show pairing of electrons Does not show core electrons

Noble Gas notation

Allows focus on the valence electrons (that control bonding) Quickest method

Does not show pairing of electrons

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Exceptions to the Aufbau Rule

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Stability of d Subshells with 5 or 10 d 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 than the above configuration, so it is: 4s1 3d5

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Elements with Exceptions The following elements are excepts to the Aufbau Principle: Element

Should be

Actually is

Cr

4s2 3d4

4s1 3d5

Mo

5s2 4d4

5s1 4d5

W

6s2 5d4

6s1 5d5

Cu

4s2 3d9

4s1 3d10

Ag

5s2 4d9

5s1 4d10

Au

6s2 5d9

6s1 5d10

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

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Quantum Numbers

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Definition: Quantum Numbers

Quantum Numbers – A set of 4 numbers that describes the electron’s placement in the atom.

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4 Quantum Numbers ml

n

2, 1, -1, + ½ l Quantum Number

ms

Symbol

Describes

Possible Numbers

Principal

n

Shell number

Whole #s ≥ 1

Azimuthal

l

Subshell type

Whole # < n

Magnetic

ml

Orbital

-lÆ+l

Spin

ms

Spin

+ ½ or – ½

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Determining Quantum Numbers 4p 3

n: principal energy level Give the number of the shell

l: subshell

s=0 p=1 d=2 f=3

coding system

ml: orbital

s

0

p

-1

0

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

1

f

d -3

-2

-2

-1

0

1

2

-1

0

1

2

3

↑= + ½ ↓=-½

ms: spin Coding system

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Quantum Number Examples Example: Give the quantum numbers for the red arrow. 1s

2s

2p

3s

3p 0

It’s in level “3”

It’s in subshell “s”—the “code” for “s” is “0” It’s in orbital “0” ___, 3 ___, 0 ___, 0 ___ -½

It’s a down arrow

Example: Give the quantum numbers for the red arrow. 1s

2s

It’s in level “2”

2p

3s -1

3p

0 +1

It’s in subshell “p”—the “code” for “p” is “1” It’s in orbital “-1” It’s an up arrow

___, 2 ___, 1 ___, -1 ___ +½

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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 # > 0 l = 1…subshell is “p” There is no p subshell in energy level 1

2, 1, ; -2, - ½

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

1, 0, 0, ; -1

n = 1…OK as n can be any whole # >0 l = 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

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Summary Electron Electron configurations configurations can can be be shown shown with with boxes boxes and and arrows, arrows, in in spectroscopic spectroscopic notation, notation, or or noble noble gas gas notation. notation.

Atoms Atomsare are made made of of protons, protons, neutrons neutrons and andelectrons. electrons. The The configuration configuration of of the the electrons electronsdetermines determines the the chemical chemical properties properties of of the the atom. atom.

Electrons Electrons are are organized organized in in levels, levels, subshells subshells and and orbitals. orbitals.

Quantum Quantum numbers numbers describe describe the the location location of of an an electron electron in in an an atom atom and and are are aa series series of of 44 numbers. numbers.

Electron Electron configurations configurations are are written written following following the the Aufbau Aufbau principle, principle, Hund’s Hund’s Rule Rule and and the the Pauli Pauli Exclusion Exclusion Principle. Principle.

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Atomic Structure and Electron Configuration Rapid Learning Center

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Atomic Structure and Electron Configuration - mrgssciencepage

Atomic Structure and Electron Configuration - mrgssciencepage

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