Unit 2 Chemistry Atomic Theory and Periodicity

Atomic Structure

The Atom and Periodic Table

I. Scientific Notation is another way to express a very large or
very small number

Give an example:

6.02 x 1023

Practice Problems:

Multiple Choice

Three protons have a total charge of _______.

• +3

• 0

• -3

• 2

Multiple Choice

Three neutrons have a total charge of ________.

• +3

• 0

• 2

• -3

Multiple Choice

A nucleus with three protons and three neutrons has a total charge of ________.

• +3

• 0

• -3

• +6

Multiple Choice

How many protons are in a Helium atom? (Hint: same as atomic number!)

• 2

• 8

• 5

• 6

Multiple Choice

How many protons are in an Oxygen atom? (Hint: same as atomic number!)

• 2

• 8

• 5

• 6

Multiple Choice

This is a neutrally-charged lithium ion. It has 3 protons and 3 neutrons. How many electrons does it have?

• 9

• 0

• 6

• 3

Dimensional Analysis

•A method for converting units using
proportions.

•Example:
•Convert 4.6 hours into seconds

The Elements

A. An element is: the most fundamental form of matter

B. Elements can exist as pure substances or as parts of
compounds.

•

Examples: Fe, H2, H2O

•

Symbols for the Elements

Chemical symbols are: abbreviations for the elements

Be forewarned, the one or two letter chemical symbol for
many elements is not the same as the first one or two letters
in the element name.

•Examples: Cu - Copper

Dalton’s Atomic Theory

John Dalton (1766-1844) – English Scientist

1.Elements are made of tiny particles called atoms.

2.All atoms of a given element are identical

3.The atoms of a given element are different from those

of any other elements.

4.Atoms of one element can combine with atoms of other

elements to form compounds.

5.Atoms are indivisible in chemical processes.

The Structure of the Atom

A. J. J. Thomson’s Plum Pudding Model (1890’s) described atoms

as being:

Summary of Contributions

​John Dalton

​​J.J. Thomson​

​Ernest Rutherford

​Neils Bohr

​​first atomic theory

​

Inferred all atoms have the same four characteristics​

discovered the electron

​

"Chocolate Chip Cookie" Model

discovered the positive charge of the nucleus + the proton

​

Gold Foil Experiment

theorized about electron orbitals around the nucleus

​

Orbital Model​

​(1600 - present time)

Reorder

Reorder the following

• Democritus

• Dalton

• Thomson

• Rutherford

• Bohr

1

2

3

4

5

• 1
• 2
• 3
• 4
• 5

B. Ernest Rutherford’s “Nuclear” Model (1911) of the atom described the
structure of atoms as follows:

1. Gold Foil Experiment (Rutherford)

Introduction to the Modern Concept of Atomic Structure
A. From Rutherford’s gold foil experiment we learned that.

1. Atoms have a central positive nuclear charge
2. Atoms are 99.99% empty space.

B. Subatomic Particles

Particle

Mass

Charge

Location in the Atom

Electron

1

1-

Outside nucleus

Proton

1836

1+

Nucleus

Neutron

1839

None

Nucleus

Drag and Drop

Which subatomic particle has:

mass of 1 and located in nucleus ​​ ​

Placeholder

​

Placeholder

​ ​

mass of 0 and located outside of nucleus ​

Placeholder

Drag these tiles and drop them in the correct blank above

proton

neutron

electron

Match

Match the following

• proton

• neutron

• electron

positive charge

neutral charge

negative charge

• positive charge

• neutral charge

• negative charge

Drag and Drop

Element 11 is called ​

Placeholder

has symbol ​

Placeholder

. It has a ​

Placeholder

of 23, so that means it has ​12 ​

Placeholder

.

Drag these tiles and drop them in the correct blank above

Sodium

Na

mass number

neutrons

Multiple Choice

What are the correct responses for Carbon's symbol, atomic number, protons, neutrons, and electrons?

• C, 6, 7, 7, 7

• Ca, 6, 7, 7, 7

• C, 6, 6, 6, 6

• Ca, 6, 6, 6, 6

Multiple Choice

This is the Bohr model of which element?

• Oxygen

• Sulfur

• Silicon

• Carbon

Multiple Choice

This is the Bohr model of which element?

• hydrogen

• beryllium

• lithium

• fluorine

Draw

Draw the Bohr model of nitrogen. Just place the electrons on the orbits, making sure they are correctly paired.

Draw

Draw the Bohr model of fluorine. Just place the electrons on the orbits, making sure they are correctly paired.

A “Box” in the Periodic Table:

Mass Number = Number of Protons + Number of Neutrons

Also, for a neutrally charged atom,

Atomic Number = Number of Protons = Number of Electrons

*the atomic number determines the identity of an element

Ions

A. A neutrally charged atom has a(n) __same

number_________ of protons and electrons.

B. A neutrally charged atom can become an ion if _____ it loses

or gains electrons__.

C. An ion is ___an atom with a different number of

electrons___________
1. A positively charged ion is known as a __cation________.
2. A negatively charged ion is an ____anion_____________.

Isotopes
A. Atoms of the same element always have the same number of

__protons_____.

*the number of protons an atom of a certain element contains is

given by its ____atomic number__________________.

B. Isotopes are: atoms of the same element with different numbers of

nuetrons

C. The sum of the number of protons and the number of neutrons

contained in an atom is know as the atom’s __atomic mass_____.

D. Isotope Notation:

1. Isotopes are often are symbolized by an element symbol with

superscripts and subscripts denoting the isotope’s mass number and
atomic number.

Examples: C

C

C

2. Isotopes are can also be written as the element’s name, a dash,
and the mass number of the isotope.

Examples:

C

C

C

Weighted Average Atomic Mass

What is a weighted average?
Your grades are weighted averages. (Major – 60%,

Minor – 40%)

Why do we need one?
It is the best estimate we can use.

There are two different types (isotopes) of copper atoms.

Cu-63 (62.93) accounts for 69.09% of all copper atoms
Cu-65 (64.94) accounts for 30.91% of all copper

What is the mass of copper on the Periodic Table?

63.546

The atomic masses on the Periodic Table are averages. Specifically,
they are weighted averages.

B. How to calculate a weighted average atomic mass

Take the sum of the products of each isotope’s mass and its
corresponding relative abundance as a decimal (take percent and
move decimal 2 places left).

Example 1

Isotope name
Isotopemass

(amu)

Percent

Abundance

Silicon-28

27.98

92.21

Silicon-29

28.98

4.70

Silicon-30

29.97

3.09

Example 2

Isotope name

Isotope

abundance

Isotopemass

(amu)

Iron-54

5.90%

53.94

Iron-56

91.72%

55.93

Iron-57

2.10%

56.94

Iron-58

0.280%

57.93

Math Response

Element J has two isotopes. J-35 has an abundance of 75.78%; and J-37 has an abundance of 24.22%. What is the average atomic mass of this element?

Type answer here

Correct Answer: 35.48

Deg^°

Rad

Rutherford’s Atom

A. His model leaves many questions about electrons unanswered

1. How are the electrons arranged?

2. How do they move?

3. Since the nucleus and the electrons are oppositely charged,

why doesn’t the atom collapse?

The Bohr Model of the Atom
A. Bohr’s model of the atom included the following main points.

1. Central nucleus made up of ___protons____ and

___neutrons_____.

2. Electrons were restricted to circular orbits.

Electromagnetic Radiation – energy transmitted as a wave

A. Parts of a wave

1. Wavelength is – distance between trough or crests

2. Frequency is – # of wave per unit time

B. Types of EM Radiation and wavelength.

Emission of Energy by Atoms

A.

Electrons surrounding an atom can absorb a discrete packet of energy
called a ______photon_____ to become “excited”.

B. When excited electrons lose that extra energy they fall back into their

__ground state_______. The release of energy by the electron results
in emission of a photon of a certain wavelength. Each element has its
own unique spectrum of wavelengths that are released.

Examples:

Electronic
Transitions
in the Bohr
Model for
the
Hydrogen
Atom

The Energy Levels of Hydrogen

A. Electrons can only absorb quantized amounts of energy…this means
___that they exist in measureable areas around the nucleus___________

B. With only one electron, a hydrogen atom is the simplest way to view
what can happen when electrons get excited.

C. How can Hydrogen produce photons with 4 different energies (colors)
when it only has 1 electron to excite?

The Wave Mechanical Model of
the Atom (Quantum Mechanics

Model)

This model of the atom is our most current and up to

date model.

The biggest difference between this model and Bohr’s

model is “Orbits vs. Orbitals”.

1. An Orbit is – an elliptical path around a central

object.

2. An Orbital is – is a “best guess” area for an

electron’s location

(a) The Probability
Distribution for the
Hydrogen 1s Orbital
in
Three-Dimensional
Space (b) The
Probability of Find
the Electron at
Points Along a Line
Drawn From the
Nucleus Outward in
Any Direction for the
Hydrogen 1s Orbital

The Hydrogen Orbitals

A. Within each principal energy level there can be one or more

orbitals.

1. “s” orbitals:
2. “p” orbitals:
3. “d” orbitals:

B. Each principle energy level is a little larger and further away from

the nucleus than the last and contains more orbitals than the last.

Two
representations of
the Hydrogen 1s,
2s, and 3s Orbitals
(a) The Electron
Probability
Distribution (b)
The Surface
Contains 90% of
the Total Electron
Probability (the
Size of the Oribital,
by Definition)

Electron Configuration

Mrs. Nebzydoski's Lesson

Multiple Choice

What is the shape of a p orbital?

• sphere
• it's just too complex to thing about it
• dumbbell
• I don't know this stuff.

Electron Configuration tells you the location that an electron "lives" at in an atom

• There are four types of orbital shapes

• S, p, d, & f

Multiple Choice

1.  What is the shape of an s orbital?

• sphere
• dumbbell
• double dumbbell
• TOO COMPLEX TO KNOW IT.

Multiple Choice

There are _____ types of atomic orbitals.

• 1

• 2

• 3

• 4

Multiple Choice

How many d orbitals are there in a given sublevel?

• 1
• 3
• 5
• 7

Multiple Choice

How many p orbitals are there in a sublevel?

• 2
• 1
• 4
• 3

The orbital shapes can be described as the path that the electron takes around the nucleus.

• There is:

• 1 s orbital

• 3 p orbitals

• 5 d orbitals

• 7 f orbitals

Multiple Choice

How likely is it that an electron occupying a p or a d orbital would be found very near an atom’s nucleus? 

• likely...that's where you find it
• not likely
• not likely, but there is a f orbital near the nucleus
• what's near the orbital

Orbital shapes

As you can see, the s orbital is a circular path around the nucleus and that is why there is only one option. The p orbital is a figure-8 path that can be on the x-axis, the y-axis, or the z-axis in a 3-D atom. This is why there are 3 different p orbitals.

Multiple Choice

How many electrons can the d sublevel hold?

• 8
• 10
• 2
• 4

Multiple Select

Which of the following are TRUE of atomic orbitals?

• There is 1 s orbital

• There are 3 p orbitals

• There are 5 d orbitals

• There are 7 f orbitals

Multiple Choice

Which is associated with more energy? 

• 2p
• 2s
• 3p
• 1s

Each orbital can contain up to 2 electrons

• s orbitals can hold a total of 2 electrons (1 orbital)

• p orbitals can hold a total of 6 electrons (3 orbitals)

• d orbitals can hold a total of 10 electrons (5 orbitals)

• f orbitals can hold a total of 14 electrons (7 orbitals)

Multiple Choice

How many total electrons can the f orbitals in a sublevel hold?

• 2
• 14
• 6
• 10

Multiple Choice

How many electrons can fit in one orbital?

• 1

• 2

• 3

• 4

Multiple Choice

Which orbital has the least amount of energy? 

• p orbital
• d orbital
• s orbital 
• What is an orbital?

Periodic Table and Orbitals

The periodic table gets its shape from the type of atomic orbitals that are filled in the atoms of the collumns.

Representation of the 2p Orbitals (a) The

Electron Probability Distribution for a 2p Orbital

(b) The Boundary Surface Representations of

all Three 2p Orbitals

The blocks of the periodic table

As you can see here the "s" block is shown on the left in pink, the "p" block is on the right in orange, the "d" block is in the middle shown in blue, and the "f" block is shown below in yellow.

Representation of the 3d Orbitals (a) Electron
Density Plots of Selected 3d Orbitals (b) The
Boundary Surfaces of All of the 3d Orbitals

Multiple Choice

Look at a periodic table. Count how many columns on the periodic table are in the "p" block. How many columns did you find?

• 2

• 6

• 10

• 14

Representation of the 4f Orbitals in Terms of

Their Boundary Surfaces

Columns on the periodic table and the orbitals...

• The s block contains 2 columns

• The p block contains 6 columns

• The d block contains 10 columns

• The f block contains 14 columns

Organization of Principle Energy
Levels, Orbitals, and Sublevels

A. Principle energy levels: Where orbitals are
located.

B. Orbitals: The only place an electron can exist.

C. Sublevel: The combination of a Principle
Energy Level and an Orbital.

D. Pauli exclusion principle: An atomic orbital
contains a maximum of two electrons

E. Aufbau Principle: Electrons enter orbitals of
lowest energy first.

Think about it....

The s block contains 2 columns and the "s" orbitals can contain 2 electrons. The p block contains 6 columns and the "p" orbitals can contain up to 6 electrons. The d block contains 10 columns and the the "d" orbitals can contain up to 10 electrons! (This is NOT a coincidence!)

House Diagram

Multiple Choice

If the f block contains 14 columns, then it would make sense that an f orbital can hold up to _____ electrons.

• 2

• 5

• 7

• 14

The Orders of the Energies of the Orbitals in the First

Three Levels of Polyelectronic Atoms

Orbital filling

Electrons fill these orbitals in a way that allows them to take the easiest root possible. The LOWEST energy levels are filled first.

So what does all of

this mean?

Principle

Energy Level

Type of
Orbitals
Present

Number of

Orbitals

Maximum
Electrons in

Orbital

1

s

1

2

2

s

1

2

2

p

3

6

3

s

1

2

3

p

3

6

3

d

5

10

4

s

1

2

4

p

3

6

4

d

5

10

4

f

7

14

Orbital filling cheat sheet

You copied this orbital filling diagram cheat sheet in your notes. It is important that you always use it to fill electrons into orbitals when completing electron configurations.

So how do we write an

“electron configuration”?

• The electron configuration for an element

shows where all of the electrons are within
the electron cloud.

• All electrons in the atom must be accounted

for.

• You cannot overfill an orbital.

• Must is in the correct format:

2p4

Principle
Energy
level

orbital

Electrons

Multiple Choice

Using your orbital filling cheat sheet, tell me which orbital fills after the 2 p orbital is full.

• 3p

• 3s

• 2s

• 1p

There is an order that
you need to follow….

3d10

1s22s22p6

3p64s2

4p65s24d105p66s2

4f145d106p67s25f147p6

3s2

Why?

ENERGY SAYS SO!!!!!!!!!

Multiple Choice

Using your orbital filling cheat sheet, which orbital fills after the 4s orbital?

• 4p

• 5s

• 3d

• 4f

Another Way to Remember

the ORDER

7s 7p 7d 7f

6s 6p 6d 6f

5s 5p 5d 5f

4s 4p 4d 4f

3s 3p 3d

2s 2p

1s

Up next....Electron Configuration!

The end.

Complete Electron

Configurations

• You will write the complete electron configurations
for all of the elements with atomic numbers 1 through
54.
•Use a separate piece of paper.

Example:
H 1s1

He1s2

Li
Be
B

Hund’s Rule

•When filling orbitals in a sublevel, each orbital must
have one electron before any orbital can have two
electrons.

•Examples:
•N 1s2 2s2 2p3

• Co 1s2 2s2 2p6 3s2 3p6 4s2 3d6

Types of Electrons

•Valence
•Electrons in the outermost Principle Energy
Level

•Core
•All other electrons

Exceptions to the

Order

•Atoms are most stable when they have full or ½
filled sublevels

•Copper
1s2 2s2 2p6 3s2 3p6 4s2 3d9

•Silver
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d9

Ion Configuration

• How many electrons would sodium have if it lost

one electron?
10

• What would the electron configuration look like?

1s2 2s2 2p6

• How many protons does it have?

11

• +11 + (-10) = +1

Therefore…

•If atom becomes a POSITIVE ion it has lost an
electron(s)
•Lost 1 = +1

•Lost 2 = +2

•If an atom becomes a NEGATIVE ion it has
gained an electron(s)
•Gain 1 = -1

•Gain 2 = -2

Excited State

•Look at Sulfur’s configuration

1s2 2s2 2p6 3s2 3p4

•What does it look like in the excited state?

1s2 2s2 2p6 3s1 3p5

Or

1s2 2s2 2p6 3s2 3p3 3d1

Video Summary of the Atom

http://www.youtube.com/watch?v=R79SGQ02C2Q

The orbitals being filled for elements in various parts of the

periodic table.

Electron Configurations and the Periodic Table

Modern Periodic Law – Elements are in order according to their atomic
number

Things to notice:

Transition Metals (one energy level below)

Lanthanide and Actinide Series (two energy levels below)

Orbital Filling

S, p, d, and f blocks

Groups 1, 2, 13 - 18 indicate the total number of valence electrons.

Main-group elements

Areas of the Periodic Table

Metals

Location

-to the left of the “staircase”

Properties

-luster, conduct electricity/heat, malleable, ductile, lose

electrons

Non-metals

Location

-to the right of the “staircase”

Properties

-brittle, exist in all states, gain electrons

Metalloids

Location

-on, or near, the “staircase”

Properties

-can have metallic or non-metallic characteristics

Families

atoms that have similar chemical characteristics (groups, columns)

Alkali Metals – Column 1

Alkaline Earth Metals - Column 2

Halogens – Column 17

Noble Gases – Column 18

Transition Metals – Columns 3-12

Lanthanides – La – Lu
Actinides – Ac - Lr

Natural States of the Elements

Noble Gases

Why are they Noble?

They have a full valence shell

Diatomic Elements

Definition: Elements that exist as two atom molecules in their natural
state

What are the seven diatomic elements?

H2, N2, O2, F2, Cl2, Br2, I2

What are the two elements that exist in the liquid state at room temperature?

Br and Hg

Allotrope:

Different forms of the same element.

What are carbon’s allotropic forms?

Trends of the
Periodic Table

•Atomic Radii

•Electronegativity

•Metallic Character

•Non-Metallic Character

•Ionization Energy

Metallic vs.
Non-Metallic

Metallic

Non-Metallic

Atomic Radii

• Distance between 2 nucleii of the same element,

divided in half.

•Trend

Electronegativity

•Atom’s ability to attract electrons to itself while
in a bond

Electronegativity

Trend

Ionization Energy

The amount of energy required to remove an
electron from the outermost energy level of an
atom

Ionization Energy

Trend