Download The Structure of the Atom- Chapter 4, 3

Unit 4A Glow in the Dark
4.1 History of Atomic Theory
**The ______________ is defined as the smallest particle of an element that retains the properties of
that element.**
History
400 BC Democritus, a Greek philosopher expressed his idea that matter was made of very small,
indivisible particles that he named “ATOMOS”
1803 John Dalton
Assumptions of the theory:
1
All matter is made up of _________________ particles called atoms.
2
Atoms of the ___________ element are identical in properties. Atoms
of different elements are different in properties.
3
Atoms of different elements combine in whole number ratios to form
_____________________.
4
Chemical _______________ involve the rearrangement of atoms. Atoms
cannot be created, divided, destroyed or changed into other types of
atoms
Diagram of Dalton’s Atomic Model
tiny solid ball that could not be broken
up into parts.
What was wrong with Dalton’s Theory?
o The Discovery of Subatomic Particles led to the idea that the atom is not “INDIVISIBLE.”
o Atoms of the same element can have different masses(isotopes)
Discovery of Subatomic Particles
1897 Thompson discovers the ______________ through the use of a cathode ray tube. He
knows it is ________________ charged and has an extremely small mass.
Goldstein determines there are positive particles called protons and have a mass 1837 times heavier
than an electron.
Plum Pudding Model (ball of positive charge with
negatively charged particles evenly distributed)
1910 Rutherford discovers the nucleus through the gold foil experiment and that atoms are mostly
_____________________ .
Rutherford’s Atomic Model (sphere with dense
middle center called the _________________ with
electrons dispersed around it.
1932 Chadwick confirms the ______________________ which has a mass similar to the
_____________________and no charge. They are located in the nucleus.
1913 Bohr performed experiments with hydrogen and light.
 Electrons are in levels according to how much
energy they have and that only certain energy
amounts are allowed.
 Think of the energy levels as rungs of a
_________________.
 The farther away an energy level is from the nucleus,
the more _____________ it attains!
 The first level can hold 2 electrons, then the next two levels can each hold 8 and then levels
farther out can hold 18.
 Electrons can move from one energy level to the next by gaining or losing energy(quanta).
 Ground State: An electron is as close to the nucleus as it can get.
 Excited State: An electron in a higher energy level than it should be.
 Drawing of simple Bohr Models
1920 Schrodinger developed the Quantum Mechanical Model (Modern view of the atom)
 Modern atomic theory uses calculus to show how electrons act as both particle &
_____________
 These equations show the most probable location of electrons in the atom (known as atomic
______________________________)
4.2 Atomic Structure
Subatomic Particles and their Properties
Particle
Symbol
Location
Electrical
Charge
Actual mass in
grams
-1
Relative
Mass in
a.m.u*
1 amu
+1
1 amu
1.6710-27 kg
0
.00055 amu
9.1010-31 kg
1.6710-27 kg
ep+
n0
*a.m.u : atomic mass unit
1 amu (“atomic mass unit”) = 1.66  10-27 kg
What are the 2 regions of the atom as of now?
 Nucleus: dense center containing ___________________ and _____________________.
 Electron Cloud: region surrounding nucleus containing electrons and mostly
________________________.
Counting Subatomic Particles
Mass Number
 The number of protons and neutrons in a nucleus is called the mass number.
 Round atomic mass to a whole number to get an element’s mass number.
 mass # = _________________ + _________________
Nitrogen’s mass number is 14.
Atomic Number
 Every atom has a different number of protons.
 The number of protons determines the identity of the atom
 The atomic number shows the number of protons. ___________________= protons
Nitrogen’s atomic number is
7!
Calculating Neutrons
 To calculate the number of neutrons, subtract the atomic number from the mass number
 # _______________________ = mass # - atomic #
14 – 7 = 7 neutrons
Charges
 A neutral atom has the ________________number of protons and electrons.
 An ion has a charge.
 Overall Charge = protons - electrons
Example:
How many electrons does Br-1 have?
How mane electrons does Al+3 have?
Element Information
Nuclear Symbol
____ p ____ e
____ n
Hyphen Notation:
copper – 65 ___________
Example
Symbol
Name
Atomic #
Mass #
Charge
Magnesium-25
Proton
Neutron
Electron
126
82
+2
82
How do Atoms Differ: Isotopes
 Isotopes are atoms of the same element with a different number of ____________________.
 Most elements contain a mixture of 2 or more isotopes Each one having its own
mass and abundance.
Isotope
Atomic
number
Protons
Neutrons
Electrons
Mass (a.m.u)
Lithium-6
Lithium-7
Lithium-8
You Try ? ( Be Careful: Might not be most common isotope. Never use the given atomic mass on
the periodic table unless absolutely necessary!!)
What the number of protons, neutrons and electrons in each atom?
19
57
F
9
204
Fe
Hg
26
80
How to calculate Average Atomic Mass?
35
37
Cl
17
35.45 amu
WHY?
Cl
17
 Average atomic mass is the weighted _____________________ of the masses of all naturally
occurring isotopes.
Equation:
Average atomic mass = (% abundance of isotope x mass of 1st isotope) +
(% abundance of isotope x mass of 2nd isotope) + ………
Example:
Element x has 2 natural isotopes. Calculate the average atomic mass. 1st isotope has a mass of
10.012 a.m.u with 19.91% abundance. 80.09% of the 2 nd element has a mass of 11.009 a.m.u.
You Try!
Calculate the average atomic mass if copper if it has 2 isotopes. 69.11% has a mass of 62.93
a.m.u and the rest has a mass of 64.93 a.m.u.
4.3 Electron Structure
Where are the electrons?
1. Within the Electron Cloud are “Energy Levels”
 There are ______ on the periodic table.
 The period number on the periodic table corresponds to the energy level.
 Energy levels are also called ___________________.
 Ca is in energy level _______
Cl is in energy level _____________
2. Within each energy level are “Subshells”
 Subshells are a set of orbitals with equal _______________
“s” subshell
 spherical shaped
 there is only ______ orientation (position) = orbital
 represented in the periodic table as
groups 1A and 2A + helium
 first seen in the 1st energy level
 maximum ______ electrons
“p” subshell
 dumbbell shaped
 there are ______ orientations(positions) = orbitals
 represented in the periodic table
as groups 3A- 8A
 first seen in the 2nd energy level
 maximum ____ electrons
“d” subshell
 four lobed shaped
 there are _______ orientations(positions)= orbitals
 represented in the periodic table as the transition metals, group 3B – 2B
 first seen in the 3rd energy level
 maximum ____ electrons
“f” subshell
 too complex of a shape to name
 there are _______ orientations(positions)= orbitals
 represented in the periodic table as
the inner transition metals, lower block
 first seen in the 4th energy level
 maximum ____ electrons
3. Within each sublevel are “Orbitals”
 An orbital is defined as an area of high ______________ of the electron being located.
 Each orbital can hold ___________ electrons
 To calculate the total number of electrons in an energy level, use 2(n2)
Summary
n2
2(n2)
Energy Level # of subshells
subshell
Total # of orbitals Total # of electrons
1
s
1
2
1
2
s&p
1&3=4
8
3
s&p&d
1&3&5=9
18
2
3
4
s & p & d & f 1 & 3 & 5 & 7 = 16
32
4
Electron Configuration

Is an _________________ of an electron

Electrons must be placed in the ____________
possible energy levels first (ground state)
4p1
3 rules that govern electron configuration
1.
Aufbau Principle: electrons must be fill the lowest available subshells
and orbitals before moving to the next higher energy subshell/orbital.
 Filling Order: 1s 2s 2p 3s 3p 4s 3d 4p 5s
4d 5p 6s 4f

Can use the periodic table as a guide or memorize filling order

An orbital can ____________ within different energy levels

__________is lower in energy than 3d.
2.
Hund’s rule: place electrons in unoccupied energy orbitals of the same energy level before
doubling up.
Example: You need to add 3 electrons to a p subshell.
3. Pauli Exclusion Principle: two electrons in the same orbital must have opposite ___________.
Example: you need to add 4 electrons to a p subshell
Electron Configuration
1. Determine the number of electrons to place
2. Follow Aufbau Principle for filling order
3. Fill in subshells until they reach their max (s = 2, p = 6, d = 10, f = 14) or use
periodic table as a guide.
4. The total of all the superscripts is equal to the number of electrons
Example:
Write electron configuration for S
Write electron configuration for K
Write electron configuration for Ti
Orbital Notation
 use boxes or lines for orbitals and arrows for electrons.
Examples
Write the boxes & arrows (orbital notation) for Cl
Write the orbital notation for N
Write the orbital notation for Fe
Shorthand Notation
 Noble gas is used to represent the core (inner) electrons and only the valence
shell is shown.
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 electron configuration
for the valence electrons
Example:
Write the shorthand notation for Br.
Write the shorthand notation for As.
Exceptional (special) Configurations
 Elements up to vanadium (V) follow Aufbau principle
 Half-filled or completely filled d & f sublevels have lower energies and are more stable than
partially filled d’s and f’s.
 This means that an atom can “borrow” one of its “s” electrons from the previous orbital to
become more stable.
Example: Ag
___
___ ___ ___ ___ ___
5s
4d
becomes
___
___ ___ ___ ___ ___
5s
4d
 Because the 4d sublevel is now full, the atom is at a lower energy state and therefore more
stable.
Electron Configuration for Ions
o Determine the number of electrons to place. Positive ions lose electrons;
negative ions gain electrons.
o Follow Aufbau Principle for filling order
o Fill in subshells until they reach their max (s = 2, p = 6, d = 10, f = 14) or use
periodic table as a guide.
Example: Write the electron Configuration for the sulfide ion, S-2