Atomic structures

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Atomic Structure andAtomic Structure and Electron Configuration

AP Ch i t R id L i S i

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

AP Chemistry Rapid Learning Series

Wayne Huang, PhDKelly Deters, PhDRussell Dahl, PhD

Elizabeth James, PhDDebbie Bilyen, M.A.



Learning Objectives

Basic structure of atomsHow to determine the

By studying this tutorial you will learn…

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

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configurations using three methodsHow 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 of matter thatAtom - Smallest piece of matter that has the chemical properties of the element.

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Often called the“Building Block of Matter”



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 Electron

cloud

M Very small relative mass

Charge = - (# of electrons)

Charge = # of protons

Mass = # of protons

+ # of neutrons

Overall Charge = # of protons

- (# of electrons)

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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

Found outside nucleus

# and configuration determine how the atom will react

Contributes to mass of atom

Does 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.

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 ofthe area outside of the nucleus where the electrons reside.

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Electron Clouds

Electron cloud

Principle energy levels

Subshells

The electron cloud is made of energy levels.

Energy levels are

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Subshells

Orbitals

composed of subshells.

Subshells have orbitals.



Definition: Subshell and Orbital

Subshell – A set of orbitals with equal energy.with equal energy.

Orbital – Area of probability of the electron being located.

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Each orbital can hold 2 electrons.

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



Pictures of Orbitals

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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of the atom, it is important to understand them.

Electron configurations use boxes for orbitals and arrow for electrons.

Aufbau Principle

Aufbau Principle: Electrons must fill subshells (and orbitals) so that the 1

The first of 3 rules that govern electron configurations:

( )total energy of atom is at a minimum.

1

What does this mean?

El t t fill th l t il bl b h ll

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Electrons must fill the lowest available subshells and orbitals 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

4d

6s

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2s

2p

Ener

gy

Subshells are filled from the lowest energy level to increasing energy levels.

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 before doubling up.

2

How does this work?

If you need to add 3 electrons to a p

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If you need to add 3 electrons to a p subshell, add 1 to each 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?

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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.

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



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

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.

H d’ R l Pl l t i i d bit l f th

Example: Give the electron configuration for a Cl atom.

No charge written Charge is 0Cl

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

No charge written Charge is 0ClAtomic number for Cl = 17 = # of protons

Electrons = 17

Place 17 electrons

1s 2s 2p 3s 3p

4231567910111213141516178



Spectroscopic Notation

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Definition: Spectroscopic Notation

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

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

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1s 2s 2p 3s 3p

1s2 2s2 2p6 3s2 3p5



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 0SAtomic number for S = 16 = # of protons

Electrons = 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 2p5 F

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 the 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

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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.

Periodic Table as a Road-Map

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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

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 To read the charge,

There is another tool commonly used to remember orbital filling order.

2s 2p

3s 3p 3d

4s 4p 4d 4f

5s 5p 5d 5f

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

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5s 5p 5d 5f

6s 6p 6d

7s 7p

8s

ElectronElectron Configurations of Ions

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Definition: Ion

Ion – an atom that has gained or lost electrons gresulting 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:

Full Valence Shell Ions

1s 2s 2p2 2 6

Br-1

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

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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-1

O2-

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

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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]

Br

Spectroscopic

Noble gas

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 ExampleDetermine 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 is 0As

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0 = 33 - electrons

As 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 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 than the above configuration, so it is:

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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 that describes the electron’s placement in the atom.

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4 Quantum Numbers

2, 1, -1, + ½

n ml

Quantum Number

Symbol

n

Describes

Shell number

S b h ll

Possible Numbers

Whole #s ≥ 1Principal

l ms

50/60

l

ml

ms

Subshell type Whole # < n

- l + l

+ ½ or – ½

Azimuthal

Magnetic

Spin

Orbital

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

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 ↓ = - ½

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”

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”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

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1, 0, 0, -1

ml must be between l and l

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

Atomic Structure & The AP Exam

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Atomic Structure in the Exam

How many paired or unpaired electrons are in an t f l t?

Common atomic structure problems:

atom of an element?

Given a set of quantum numbers, what’s the next higher level or sublevel?

When shown a set of boxes & arrows, which element is it?

Is a set of quantum numbers possible?

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Is a set of quantum numbers possible?

This topic isn’t used in the Free Response section very often…usually just the multiple choice!

Multiple Choice QuestionsThe first few questions on the AP MC exam give a list of choices and then ask several questions about that same list.

A. SeB. MgC. AlD. NE. F

1. The atom that contains exactly 2 unpaired electrons

Example: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4

1s2 2s2 2p6 3s2

1s2 2s2 2p6 3s2 3p1

1s2 2s2 2p3

1s2 2s2 2p5

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2. The atom that contains exactly 2 electrons in the highest occupied subshell

3. The atom that must obtain 3 electrons to obtain a full subshell

A. Se (4 e-1 in the p subshell means 2 are unpaired)

B. Mg

D. N (if a p has 3, it needs 3 more to be full)



Another Multiple Choice Question

Example: Which of the following atoms is in an excited state?

A. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4

B 1s2 2s2 2p6 3s2 3p2B. 1s2 2s2 2p6 3s2 3p2

C. 1s2 2s2 2p6 3s2 3p2 4s1

D. 1s2 2s2 2p6 3s2 3p6 4s2 3d8

E. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

An atom in an excited state is one in which an electron has been promoted to a higher subshell.

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p g

Option “C” has an electrons in the 4s subshell when the 3p isn’t full—this shows a promoted electron.

Answer: C

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

Electron configurations Electron configurations

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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gare written following the Aufbau principle, Hund’s Rule and the

Pauli Exclusion Principle.

gare 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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Atomic structures