Unit 3: The Atom
I. History and Development of the Atom
• Believed atoms were:
o The smallest particle of matter
o ___________________________ - could not be
divided or broken down any further
A. Democritus (around 400 B.C.)
indivisible and indestructible
• Based on his observations of the natural world around
him, Democritus was the first to suggest that _________
was _______________________—called them _______all matter
made up of small particles atoms
• Studied the ratios in which elements
combined in chemical reactions.
• Based on his experiments, he formulated
the first real theory about atoms:
Dalton’s Atomic Theory:
1. All _____ is ________ of indivisible particles called _____
B. John Dalton (1803-1805)
2. ___________________________________
(same mass and properties). Atoms of
different elements have different
masses/properties
All atoms of a given element are identical
3. Atoms of elements _______________________ to
_______________; compounds are formed when 2 or more
different atoms bond together
atomsmatter made up
combine in definite ratios
form compounds
4. Atoms _______________________________________
________ - they are just rearranged
• Based on his theories, Dalton viewed the atom as a _______
____________
Dalton’s Atomic Model: ______________
*Note – Dalton didn’t have one specific experiment regarding the atom
Billiard Ball Model
cannot be created or destroyed in a chemical
reaction
hard,
solid sphere
• Performed experiments using a _______________
o Involved shooting a cathode ray (a stream of electricity)
through a tube that had a magnetic field
• Observed two main things:
1. The rays were actually streams of unknown particles that
were so light, they were lighter than the mass of the smallest
known atom (hydrogen)
2. The rays were attracted to the positive plate
• Concluded two main things:
1. The atom really is divisible and it is made up of even smaller
particles
2. One of the particles is negatively charged
C. JJ Thomson (1906)
VIDEO CLIP
cathode ray tube
• He called these __________________________________
(particles “beneath” the atom) ____________
• Based on his experiments, Thomson pictured the atom as a
sphere of positively charged matter with electrons
mixed/embedded in it
Thomson Atomic Model: _________________
VIDEO CLIP
Plum Pudding Model
negatively charged subatomic particles
electrons (e-)
*Plum pudding is a British
dessert. If it helps, think of a
chocolate chip cookie
instead—the chocolate
chips are the electrons and
the dough is the positively
charged sphere
• Performed an experiment called the __________________
o He bombarded (fired) alpha particles (42He)—which are
positively charged—at a thin piece of gold foil
D. Ernest Rutherford ( 1911 )
o If the Thomson model was correct, all the alpha
particles would pass through the foil undisturbed due to
the charge of the positive sphere cancelling out the
negative, free-floating electrons. However, some
particles were slightly deflected
Gold Foil Experiment
• Based on the observation that some alpha particles
were deflected, he concluded that:
1. ______ are made up of ________________
2. There was a ______________________________
_______ = the discovery of the __________!
Rutherford Atomic Model: _____________Nuclear Model
• Based on the observation that most alpha particles
passed through un-deflected, he concluded that:
*Note - Provided no information about ________ other than the fact that
they were located _________________.
VIDEO CLIPhttp://www.mhhe.com/physsci/chemi
stry/essentialchemistry/flash/ruther1
4.swf
Atoms mostly empty space
small, dense, positively charged
center nucleus
electronsoutside the nucleus
E. Neils Bohr (1913)
• Expanded the atomic model by analyzing _________
_____________________
o Emission spectra = a chart of lines of light given off
when an electric current is run through an atom
• Concluded that the __________________was
_______ by the ___________________
o ______________on the spectrum made him
conclude that __________ must be moving from
____________________
the emission
spectra of hydrogen
light emitted/given off
caused movement of electrons
Different colors
electrons
different energy levels
• Based on his experiments, Bohr’s model of the atom had
__________________________________ in well-defined
_________ called _________
o ____________ in ________________had ___________
__________________
• It looked like a solar system – the nucleus was like the sun and
the electrons orbited around the nucleus like planets
Bohr Atomic Model: _____________Planetary Model
electrons traveling around the nucleus
paths orbits
Electrons different orbits different
amounts of energy
F. Werner Heisenberg (1926)
• Bohr’s model only explained the hydrogen atom
with one electron. Did not explain multi-electron
atoms
• Based on his research with multi-electron atoms,
Heisenberg proposed that electrons do not travel
around in circles around the nucleus, instead, they
____________________________
o Orbital: A region in which an electron is most
likely (high probability) located
randomly move in regions (orbitals)
Heisenberg Atomic Model: ___________________
• The nucleus was still the dense, positive center but now it was believed that one cannot know the exact position of an electron; there were only areas where the electron is most likely found (orbitals)
• Heisenberg viewed the atom more like a bee and a hive whereas the Bohr model is like orbiting planets around the sun
Wave-Mechanical Model
* Note - also called the Modern Atomic Model, Quantum-Mechanical Model, or Electron Cloud Model
Practice
1. Which of the following did Rutherford’s Gold Foil experiment prove?
a) That the atom was a uniformly dense sphere.
b) That the atom is mostly empty space with a dense, positive core.
c) That most the atom consists of a uniform positive “pudding” with
small negative particles called electrons embedded throughout.
d) That electrons travel around the nucleus in well-defined paths
called orbits.
2. J.J. Thomson’s Cathode Ray Tube experiment led to the discovery of
a) the positively charged subatomic particle called the electron
b) the positively charged subatomic particle called the proton
c) the negatively charged subatomic particle called the proton
d) the negatively charged subatomic particle called the electron
3. According to the Bohr Model, a) electrons are found in areas of high probability called orbitals
b) electrons travel around the nucleus in circular paths called orbits
c) electrons are found in areas of high probability called orbits
d) electrons travel around the nucleus in random paths called orbitals
4. According to the Wave-Mechanical Model, a) electrons are found in areas of high probability called orbitals
b) electrons travel around the nucleus in circular paths called orbits
c) electrons are found in areas of high probability called orbits
d) electrons travel around the nucleus in random paths called orbitals
II. Atomic Structure
p + 1 1 amu nucleus
n 0 1 amu nucleus
Particles in
nucleus are
called
________
e- - 1 0 amu Outside
nucleus
* Amu = _____________
A.Subatomic Particles
• Subatomic Particles =_____________________
• There are ___of them
particles inside the atom
3
Atomic mass unit
nucleons
Practice1. Which subatomic particle is neutral?
Neutron
2. Where is most of the mass of an atom
located?
the Nucleus
3. What is the charge of the nucleus of any
atom?
Positive
• Atomic Number = __________________________
o Found _______________(the bolded number)
o ________ the number of ________ in an atom
Example: Iron : ________ = ___= _________Atomic # 26
• Nuclear Charge = the ________ of the ________
o The particles in the nucleus are protons and
neutrons. Protons have a charge of +1 and neutrons
have a charge of 0
o Therefore, the nuclear charge is _____________and
_________________________
B. Vocabulary and Notation
26 protons
Example: Carbon – Atomic # =____= ___ protons =
nuclear charge of ____
6 6+6
Vocabulary
identifies the type of element it is
Equals protons
charge nucleus
always positive equal to the number of protons
on periodic table
Example: Cobalt (Co) : _____p and _____ e-27
• Mass Number = The mass of a specific
isotope(sample) of an element
o Mass # = _________________
Why does it make sense that electrons
aren’t included?
o Always a whole number
27
• Atomic Charge = The total charge of an atom
o An atom is ALWAYS _____________
___________ = __________
Example: If an isotope of nitrogen has 7 protons
and 7 neutrons, its mass number is ___
neutral (zero)
# protons # electrons
# protons + # neutrons
So light they barely contribute to mass of element
14
*C – 14 - ____ p: ____ n; ____ e-6 8 6
Mass #
Examples: 94 Be ______ p; ____ n; _____ e-4 5 4
Mass #
Atomic #
• Isotopic Notation: Shows the mass number of an atom
along with element symbol
*C-14 14C Carbon – 14
They all mean the element carbon with a mass number of 14
Notation
Practice
Element Atomic # Mass # Number
of
Protons
Number
of
Neutrons
Number
of
Electrons
Nuclear
Charge
Na - 23
35Cl
K-40
Silver 108
Use your Periodic Table and your knowledge of the atom
to fill in the following chart
11 23 11 1112 + 11
17 35 17 18 17 +17
19 40 19 21 19 +19
47 47 61 47 +47
C. Electrons
• Electrons and how they behave are responsible for many parts of
chemistry
• Even though it is not technically correct, Bohr’s model of the atom is
often used when discussing electrons and the structure of the atom.
It is easiest to visualize and it is “good enough”
• According to Bohr’s model, electrons are located outside of the
nucleus in energy levels. Each energy level can hold a certain
amount of electrons.
Energy level # of electrons
n=1 2
n=2 8
n=3 18
n=4 32
The __________________is from the nucleus, the ___________it
has; therefore, it is ___________and _____________
Closest to nucleus
Furthest away
Energy Levels
farther the electron more energy
less stable easier to move
Electron Configurations
• Electron Configurations = a dashed chain of numbers that shows _______________________________________
o found in the lower left corner of an element box (see below)
• Tells us the number of energy levels as well as the number of electrons in each level
Example: Carbon’s electron configuration is ____
This means it has __ electrons in the ________energy shell and __ in the ________ energy shell (so a total of __ electrons in the atom)
*All electron configurations on the Periodic Table are for atoms when they are most stable (notice #p= #e-)
how electrons are arranged around nucleus
2-4
2 first4 second 6
SUBSTANCE ELECTRON CONFIGURATION
Magnesium
Bromine
*Lead
(see the * at
bottom of
periodic table)
*shortcut allows you to cut out the first two energy levels to shorten
the configuration so it can fit in the box
Practice
Use your Periodic Table to fill in the electron configurations
for the atoms of the following elements
2-8-2
2-8-18-7
On PT = -18-32-18-4
Actually = 2-8-18-32-18-4
Types of electrons
• There are two types of electrons: valence electrons and kernel
electrons
Valence Electrons:
• electrons found in the ___________ shell or energy level
• the ___________in the electron configuration
• the electrons that get lost or gained because they are the furthest away
from the nucleus so they are the easiest to remove
• An element is ___________when its ___________________________
(valence shell)
o __________
o *Hydrogen and Helium are exceptions-stable with 2* Why?
Kernel Electrons:
• Inner electrons (all the other, non-valence electrons)
Example: Calcium’s configuration is 2-8-8-2; therefore it has __
valence electrons and ____ kernel electrons.
outermost
last number
most stable last occupied energy level is full
8 is great!
The first shell IS full with only 2 e-
218
Practice
Electron
configuration
# valence e- # kernel e-
Chlorine
Nitrogen
Sodium
Use your PT to fill in the electron configurations for the atoms of the
following elements. Then identify the # of valence and kernel e-
2-8-7 7 10
2-5 5 2
2-8-1 1 10
D. Atom Diagrams
• There are two common diagrams used to represent the
structure of the atom: Bohr Diagrams and Lewis Dot
Diagrams
Bohr Diagrams
• As previously mentioned, Bohr’s model is often used when
visualizing an atom
• Bohr Diagrams are models of the atom that have the
electrons in rings (orbits) around the nucleus
Steps for drawing Bohr Diagrams:
1. Draw a circle representing the nucleus
2. Find the number of protons and neutrons and write them inside the
nucleus
• To find # of protons-find the element’s atomic number using the
Periodic table
• To find # of neutrons – subtract the atomic number (or number
of protons) from the mass number
3. Look up the element’s electron configuration on the Periodic Table
4. Use the electron configuration to determine how many rings will be
around the nucleus (# of energy levels = # of rings)
Example: Magnesium’s configuration is 2-8-2 so there will be
______________around the nucleus
5. Using dots to represent electrons, fill in the number of electrons in
each ring
3 rings/circles
Example: Draw the Bohr Diagram for C-14
6 p
8 n...
.
.
.
Electron configuration: 2-4
Lewis Dot Diagrams
• Whereas Bohr Diagrams illustrate all the electrons of an
element, lewis dot diagrams or ___________________,
only illustrate the valence electrons
o Valence electrons are often seen as the most important
ones because they are the electrons that are gained or
lost when elements bond to form compounds
• Electron dot diagrams consist of the _______________
surrounded by dots that represent its _______________
electron dot diagrams
element symbol
valence electrons
Examples:
Ca N F
2-8-8-2
Valence e-
Ca
2-5
N
2-7
F
Notice – Put one electron on each side then double up!
Steps to Drawing Lewis Dot Diagrams:1. Write the element’s symbol
2. Find the electron configuration from Periodic Table. The last number
in the configuration is the number of valence electrons
3. Using dots to represent the electrons, place the electrons around
the element symbol, one at a time, starting first at the 12 spot on a
clock. Then add any remaining valence electrons one at a time to
the 3, 6, and 9 spots and then double up if there are more valence
electrons.
• Note: you must add only one electron at a time
because of bonding
o Bonding site = Where there is only a
single electron or an ________________
(lone electrons are open to attach to other
e- and/or easily lost)
2-4
C
Bonding Sites
unpaired electron
Practice1.What is the maximum number of electrons an atom or an ion can have in its
valence shell?
a. 2
b. 4
c. 6
d. 8
*this means that the most dots you can have in a Lewis dot diagram is 8!
2.The number of bonds an atom of an element can form is the same as the
number of
a. electrons in its valence shell.
b. paired electrons in its valence shell.
c. unpaired electrons in its valence shell.
3. Looking back at your Lewis Dot Diagrams, which element can form the
most bonds?
a. Calcium b. Nitrogen c. Fluorine
III. Ions
• Ion = ____________________________(# of protons
DOES NOT EQUAL # of electrons)
• Ions ____________whereas atoms do not!
Example: 2311 Na +1
Atomic # =
Mass # =
Ion Charge =
23
11
+1
# of p =
# of n =
# of e- =
1123 – 11 = 12
(# protons - Ion Charge)
11- (+1) = 10
A.What is an ion?
atom that lost or gained electrons
have a charge
1. Anion =
Negatively charged ion (atom GAINED e-)
2. Cation =
Positively charged ion (atom LOST e-)
Remember:
a CATion is
PAWsitive
• There are two types of ions
Examples:
7 Li + 1 ______ p ________ n ___________e ________
31 P – 3 ______ p ________ n ________ e ___________
79 Se – 2 ______ p ________ n ________ e ___________
19 F – 1 ______ p ________ n ________ e ___________
3
Atomic #Mass # -
Atomic #
7 – 3 = 4
Atomic #-
Charge
3 – (+1) = 2 Lost e-
(cation)
15 16 18 Gained e-
(anion)
34 45 36 gained e-
(anion)
9 10 10 gained e-
(anion)**Think of weight loss – losing weight/electrons is a positive
thing, gaining weight/electrons is a negative thing. It’s
opposite!(when you gain something, it’s negative)
Cation/anion?
Practice
1. When a neutral atom gains an electron, it becomes a
a) negative cation
b) positive cation
c) negative anion
d) positive anion
2. When a neutral atom loses an electron, it becomes
a) negative cation
b) positive cation
c) negative anion
d) positive anion
3. What is the charge on a magnesium ion that has lost two
electrons? _______
4. What is the charge on a fluoride ion that has gained one
electron? _______
5. The chemical symbol Fe+3 represents
a) cation formed as a result of a iron atom losing 3 electrons
b) cation formed as a result of a iron atom gaining 3
electrons
c) anion formed as a result of a iron ion losing 3 electrons
d) anion formed as a result of a iron ion gaining 3 electrons
6. Give the correct chemical symbol for the ion formed when
oxygen gains 2 electrons: ______
+2
-1
O-2
B. Ion Diagrams
• Bohr Diagrams and Lewis Dot Diagrams can also be
used to represent ions
• The steps are the same as the atom except you must
add or subtract electrons from the last number in the
electron configuration
o The last number represents the electrons in the
shell/energy level furthest from the nucleus so they
are the least stable and the easiest to access.
Remember: if it is a __________, you _________
electrons. If it is a ____________, you ____electrons
(opposite!)
positive ion subtract
negative ion add
Examples:
Draw the Bohr Diagram for 40Ca and 40Ca+2
Draw the Bohr Diagram for 19F and 19F-1
• For a lewis dot diagram you must also change the valence
electrons (add or subtract electrons) in the configuration
before doing the diagram. Also, your final diagram must
include _____________________________
Examples:Ex 1: S vs S-2
ADD 2 e- to the 6 that S normally
has in its valence shell.
Ex 2: K vs K+1
REMOVE 1 e- from the valence
shell of K.
*negative ions always end up
with 8 valence e- (8 dots)
*positive ions always end up
with 0 valence e- (0 dots)
brackets and the charge of the ion.
IV. Electron Transitions
A. Ground State vs. Excited State
What do you notice in the diagrams?
• Ground State = Electrons in lowest energy configuration/energy levels possible (____________________________________)
o Stable
• Excited State = Electrons are found in a higher energy configuration (_______________________________)
o Unstable
o excited state electron configuration for Li could be 1-2, 1-1-1 vs. 2-1 ground state
the configuration found on periodic table
any configuration not found on PT
Distinguish between ground state and excited state electron
configurations below:
2-5
2-8-8-1
2-7-1
1-6
Examples:
Ground
Ground
Excited
Excited
• Hint: When atoms are in the excited state, they are still atoms, meaning
protons=electrons. Instead of searching aimlessly for the configuration on
the table, do the following:
1. add up the total number of electrons in the configuration
2. Because it’s an atom, p=e so now that you have the e- you can find the
protons/atomic #/what element it is.
3. Compare the configuration you are given to the one on the table. If it’s
the same=ground state; if it’s different=excited state
• Warning: Both the formation of ions and the excited vs. ground state
involve electrons “doing things” but there are important differences
between the two.
Ions Excited State
• Definition: When an atom gains
or loses electrons and becomes
charged
• The amount of total electrons
changes
Example:
Na (atom) 2-8-1 Total e = 11
Na+1 (ion) 2-8 Total e = 10
Definition: When electrons absorb
energy and are found in a higher
energy configuration
• The amount of total electrons stay
the same, they just move shells!
(therefore still a neutral atom)
Example:
Na (ground state) 2-8-1 Total e = 11
Na (excited state) 2-7-2 Total e = 11
B. Bright Line Spectra
• When _________________________________, they __________________________ or an excited state.
o This is a very unstable/temporary condition
• ________________ rapidly ____ back down or drop ____________________(because they are unstable in the excited state)
• When excited electrons fall from an excited state to lower energy level (to the ground state), they _______________in the _____________(photons).
• One way this light is commonly analyzed is through a bright-line spectrum
o Recall, bright-line spectrum = a chart of lines of light given off when an electric current is run through an atom
ground state electrons absorb energy
jump to a higher energy level
Excited electrons fall
to a lower energy level
release energy form of light
• ________________________________________________
________________
o Fireworks are an example of this
• Spectra are unique for each element (like fingerprints are
unique for each person) so we can use
________________________________________
http://www.mhhe.com/physsci/chemistry/essentialchemis
try/flash/linesp16.swf
Different elements produce different colors of light or
different spectra
spectral lines to identify different elements
What elements are present in the mixture based on the
bright-line spectra?
Example:
Strontium and lithium
V. Isotopes• Isotopes = Atoms of the ______________but
______________________
o ________________ but __________________
Example:
1
1
1
0 1 11H
1 2 21H
2 3 31H
same # protons different # neutrons
same element different mass number
Practice
1. Determine the amount of each subatomic particle for
the following isotopes of Carbon (C-12, C-13, & C-14)
p = p = p =
n = n = n =
e = e = e =
6
6
6
6
7
6
6
8
6
*Notice-___________ have a ____________________
whereas _______ have a ______________________isotopes different # of neutrons ions different # of electrons
2. Two different isotopes of the same element must contain the
same number of
a. protons b. neutrons c. electrons
3. Two different isotopes of the same element must contain a
different number of
a. protons b. neutrons c. electrons
4. Isotopes of a given element have
a. the same mass number and a different atomic number
b. the same atomic number and a different mass number
c. the same atomic number and the same mass number
• We have learned that mass number is defined as the # of protons + the # of
neutrons.
• We have seen that mass numbers are all whole numbers. So what’s up with
the atomic mass given in the periodic table?
• Atomic mass and mass number are not the same; though they are similar
(the mass number should always be somewhat close to the atomic mass).
• So what is the atomic mass?
Atomic Mass = the ________________of
___________________________of an
element.
A weighted average takes in account
relative abundance, or
percentages/amount of each isotope
VI. Atomic MassA. Atomic Mass vs. Mass Number
weighted average all naturally occurring isotopes
• Yes, the atomic mass for each element is in the upper left-hand corner
on the periodic table. But how is it calculated?
Atomic Mass = the weighted average of an element’s naturally occurring
isotopes
(% abundance of isotope in decimal form) x (mass of isotope 1)
(% abundance of isotope in decimal form) x (mass of isotope 2)
+ (% abundance of isotope in decimal form) x (mass of isotope 3)
B. Calculating Atomic Mass
Mass Number Atomic Mass
The _____ of _________
of a given element.
The _____________of
___________of a given
element
mass one isotope average mass
all isotopes
Average Atomic Mass of the Element
Examples:
1. Carbon has two naturally occurring stable isotopes . 98.89% of
carbon atoms are C-12, while the remaining 1.108% are C-13.
What is the atomic mass of carbon?
Step 1: Convert % to decimal (by dividing by 100)
(0.9889)
Step 3: Add up the masses of isotopes
11.87 amu + 0.1440 amu =
(12 amu) = 11.87 amu
(0.01108) (13 amu) = 0.1440 amu
12.01 amu
Step 2: Multiply the decimal by its mass number
98.89%/100 = (0.9889) 1.108%/100 = (0.01108)
2. 92.21% of Si is found to be 27.98 amu, 4.70% is found to be
28.98 amu, and the remaining 3.09% is found to be 29.97.
Calculated the atomic mass of silicon
Step 1: Convert % to decimals (by dividing by 100)
Step 2: Multiply the decimals by the mass number
Step 3: Add up the masses of isotopes