Atomic structures

AP Chemistry Rapid Learning Series - 12

Rapid Learning Center
Chemistry :: Biology :: Physics :: Math

Rapid Learning Center Presents …
p
g

Teach Yourself
AP Chemistry Visually in 24 Hours

1/60

http://www.RapidLearningCenter.com

Atomic Structure and
Electron Configuration

AP Ch i t R id Learning Series
Chemistry Rapid L
i
S i
Wayne Huang, PhD
Kelly Deters, PhD
Russell Dahl, PhD
Elizabeth James, PhD
Debbie Bilyen, M.A.

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

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

1


Learning 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
3/60

Concept Map
Previous content
Chemistry

New content
Studies

Quantum Numbers
Matter
Location described by
Made of

Electrons

Chemical properties
determined by

Atoms

3 ways to show configurations

Boxes and Arrows

Spectroscopic
Notation

Noble Gas
Notation

4/60


2


Atomic Structure

5/60

Definition: Atom

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

Often called the
“Building Block of Matter”

6/60


3


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

7/60

The Atom
Electron
cloud

Nucleus

Mass =
M
# 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
8/60


4


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.
9/60

Electron
Locations

10/60


5


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

11/60

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.

12/60


6


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.

13/60

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

Energ increases
gy

Subshell

4

7

14

14/60


7


Pictures of Orbitals

s orbital

3 p orbitals

5 d orbitals
15/60

Electron
Configuration

16/60


8


Definition: Electron Configurations
Electron Configurations –
g
p g
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.
atom,
them

Electron configurations use boxes for orbitals
and arrow for electrons.
17/60

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
El t
t
th l
t
il bl
b h ll
and orbitals before moving on to the next higher
energy subshell/orbital.

18/60


9


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

19/60

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
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.

20/60


10


Pauli Exclusion Principle
The last of 3 rules that govern electron configurations.

3

Pauli Exclusion Principle: Two electrons that
occupy th same orbital must have different spins.
the
bit l
th
diff
t i
“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.

21/60

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

22/60


11


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:
Cl

How many electrons does the following have?
No charge written

Charge is 0

Atomic number for Cl = 17 = # of protons
0 = 17 - electrons
Electrons = 17
23/60

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 R l Place electrons in unoccupied orbitals of the
H d’ Rule: Pl
l t
i
i d bit l f th
same energy level before doubling up.

3

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

Example:
Cl

Give the electron configuration for a Cl atom.
No charge written

Charge is 0

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

2s

Electrons = 17

40
9
8
17
16
15
14
13
12
11
7
6
5
1
3
2
2p

3s

3p

24/60


12


Spectroscopic
Notation

25/60

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
26/60


13


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
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

27/60

Electron Configurations
and the Periodic Table

28/60


14


Configurations Within a Group
Look at the electron configurations for the Halogens
(Group 7).
1s 2s 2p
1 22 22 5

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.
29/60

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
hi h t energy subshell that contains electrons.
b h ll th t
t i
l t
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

30/60


15


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

31/60

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

2s

2p

3s

3p

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”
32/60


16


Another Tool for Filling Order
There is another tool commonly used to remember
orbital filling order.
1s
2s

2p

3s

3p

3d

4s

4p

4d

4f

5s

5p

5d

5f

6s

6p

6d

7s

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

7p

8s
33/60

Electron
Configurations
of Ions

34/60


17


Definition: Ion
Ion – an atom that has
g
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

35/60

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
36/60


18


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
37/60

Noble Gas
Configuration

38/56


19


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

2

2

6

2

1s 2s 2p 3s 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.
39/60

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 d you win on the Price is Right?
H
do
i
th P i i Ri ht?
By getting as close as possible without going over.
Choose the noble gas that’s closest without going over!

Noble Gas
He

2

Ne

10

Ar

18

Kr
40/56

# of electrons

36

Xe

54


20


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
2p

2s

Ne

3p

Ar

4s

3d

4p

Kr

5s

4d

5p

Xe

6p

Rn

3s

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.
41/60

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
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

42/60


21


Comparing the
Different Notations

43/60

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

44/60


22


Exceptions to the
Aufbau Rule

45/60

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

46/60


23


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
g

5s2 4d9

5s1 4d10

Au

6s2 5d9

6s1 5d10

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

47/60

Quantum
Numbers

48/60


24


Definition: Quantum Numbers

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

49/60

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
S b h ll
type

Whole # < n

Magnetic

ml

Orbital

Spin

ms

Spin

-l

+l

+ ½ or – ½

50/60


25


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

51/60

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

2s

2p

3s

3p
0

It s
It’s in level “3”
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

52/60


26


Identifying Incorrect Quantum Numbers
Example: What’s wrong with the following sets of quantum numbers?
1, 1, 0, + ½

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
l

1, 0, 0, -1

53/60

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

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

Atomic Structure &
The AP Exam

54/60


27


Atomic Structure in the Exam
Common atomic structure problems:
How many paired or unpaired electrons are in an
atom of an element?
t
f
l
t?
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?
This topic isn’t used in the Free Response section very
often…usually just the multiple choice!
55/60

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

A.
B.
C.
D.
E.

Se
Mg
Al
N
F

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4
1s2 2s2 2p6 3s2
1s2 2s2 2p6 3s2 3p1
1s2 2s2 2p3
1s2 2s2 2p5

1. The atom that contains exactly 2 unpaired electrons
A. Se (4 e-1 in the p subshell means 2 are unpaired)
2. The atom that contains exactly 2 electrons in the highest
occupied subshell
B. Mg
3. The atom that must obtain 3 electrons to obtain a full subshell
56/60

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


28


Another Multiple Choice Question
Example:

Which of the following atoms is in an excited state?
A.
B.
B
C.
D.
E.

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4
1s2 2s2 2p6 3s2 3p2
1s2 2s2 2p6 3s2 3p2 4s1
1s2 2s2 2p6 3s2 3p6 4s2 3d8
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6

An atom in an excited state is one in which an electron has been
p
promoted to a higher subshell.
g
Option “C” has an electrons in the 4s subshell when the 3p isn’t
full—this shows a promoted electron.

Answer: C
57/60

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

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

Electrons are
organized in levels,
subshells and orbitals.

Quantum numbers
describe the
ocat o o a
location of an
electron in an atom
and are a series of
4 numbers.

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

58/60


29


Congratulations
You have successfully completed
the core tutorial

Atomic Structure and
Electron Configuration

59/60

Chemistry :: Biology :: Physics :: Math

What’s N t
Wh t’ Next …

Step 1: Concepts – Core Tutorial (Just Completed)
Step 2: Practice – Interactive Problem Drill
Step 3: Recap – Super Review Cheat Sheet

Go for it!

60/60

http://www.RapidLearningCenter.com


30

Atomic structures

Recomendados

Recomendados

Mais conteúdo relacionado

Mais procurados

Mais procurados (20)

Destaque

Destaque (6)

Semelhante a Atomic structures

Semelhante a Atomic structures (20)

Mais de Timothy Welsh

Mais de Timothy Welsh (20)

Último

Último (20)

Atomic structures