Download Atoms, Molecules, and Ions Chapter 2

Atoms, Molecules, and Ions
Chapter 2
The Discovery of Atomic Structure
The ancient Greeks were the first to postulate that
matter consists of indivisible constituents.
Later scientists realized that the atom consisted of
charged entities.
The Atomic Theory of Matter
 John Dalton:
 Each element is composed of atoms
 All atoms of an element are identical.
 In chemical reactions, the atoms are not changed.
 Compounds are formed when atoms of more than one element
combine.
 Dalton’s law of multiple proportions: When two elements form
different compounds, the mass ratio of the elements in one
compound is related to the mass ratio in the other by a small whole
number.
The Discovery of Atomic Structure
Cathode Rays and Electrons
A cathode ray tube (CRT) is a hollow vessel with an electrode at either end.
A high voltage is applied across the electrodes.
Cathode Rays and Electrons
The voltage causes negative particles to move from the
negative electrode to the positive electrode.
The path of the electrons can be altered by the presence
of a magnetic field.
The Discovery of Atomic Structure
Cathode Rays and Electrons
The Discovery of Atomic Structure
Cathode Rays and Electrons
Consider cathode rays leaving the positive electrode
through a small hole.
•If they interact with a magnetic field perpendicular to
an applied electric field, the cathode rays can be
deflected by different amounts.
•The amount of deflection of the cathode rays depends
on the applied magnetic and electric fields.
•In turn, the amount of deflection also depends on the
charge to mass ratio of the electron.
The Discovery of Atomic Structure
Cathode Rays and Electrons
In 1897, Thomson determined the charge to mass ratio of
an electron to be 1.76  108 C/g.
Goal: find the charge on the electron to determine its
mass.
The Discovery of Atomic Structure
Millikan Oil Drop Experiment
The Discovery of Atomic Structure
Cathode Rays and Electrons
Consider the following experiment:
•Oil drops are sprayed above a positively charged plate
containing a small hole.
•As the oil drops fall through the hole, they are given a
negative charge.
•Gravity forces the drops downward. The applied
electric field forces the drops upward.
•When a drop is perfectly balanced, the weight of the
drop is equal to the electrostatic force of attraction
between the drop and the positive plate.
The Discovery of Atomic Structure
Cathode Rays and Electrons
Using this experiment, Millikan determined the charge
on the electron to be 1.60  10-19 C.
Knowing the charge to mass ratio, 1.76  108 C/g,
Millikan calculated the mass of the electron:
9.10  10-28 g.
With more accurate numbers, we get the mass of the
electron to be 9.10939  10-28 g.
The Discovery of Atomic Structure
Radioactivity
Consider the following experiment:
•A radioactive substance is placed in a shield containing
a small hole so that a beam of radiation is emitted from
the hole.
•The radiation is passed between two electrically
charged plates and detected.
•Three spots are noted on the detector:
•a spot in the direction of the positive plate,
•a spot which is not affected by the electric field,
•a spot in the direction of the negative plate.
The Discovery of Atomic Structure
Radioactivity
The Discovery of Atomic Structure
Radioactivity
A high deflection towards the positive plate corresponds
to radiation which is negatively charged and of low
mass. This is called b-radiation (consists of electrons).
No deflection corresponds to neutral radiation. This is
called g-radiation.
Small deflection towards the negatively charged plate
corresponds to high mass, positively charged radiation.
This is called a-radiation (He atom).
The Discovery of Atomic Structure
The Nuclear Atom
From the separation of radiation
we conclude that the atom
consists of neutral, positively, and
negatively charged entities.
Thomson assumed all these
charged species were found in a
sphere.
The Discovery of Atomic Structure
The Nuclear Atom
Rutherford’s a-particle experiment:
The Discovery of Atomic Structure
The Nuclear Atom
Rutherford carried out the following experiment:
A source of a-particles was placed at the mouth of a
circular detector.
The a -particles were shot through a piece of gold foil.
Most of the a-particles went straight through the foil
without deflection.
Some a-particles were deflected at high angles.
If the Thomson model of the atom was correct, then
Rutherford’s result was impossible.
The Discovery of Atomic Structure
The Nuclear Atom
In order to get the majority of a-particles through a
piece of foil to be undeflected, the majority of the atom
must consist of a low mass, diffuse negative charge - the
electron.
To account for the small number of high deflections of
the a-particles, the center or nucleus of the atom must
consist of a dense positive charge.
Atoms are mostly empty space!!
The Discovery of Atomic Structure
The Nuclear Atom
Rutherford modified Thomson’s
model as follows:
assume the atom is spherical but
the positive charge must be
located at the center, with a
diffuse negative charge
surrounding it.
The ModernView of Atomic Structure
The atom consists of positive, negative, and neutral
entities (protons, electrons, and neutrons).
Protons and neutrons are located in the nucleus of the
atom, which is small. Most of the mass of the atom is
due to the nucleus.
There can be a variable number of neutrons for the same number of protons.
Isotopes have the same number of protons but different numbers of neutrons.
Electrons are located outside of the nucleus. Most of the
volume of the atom is due to electrons.
The ModernView of Atomic Structure
The ModernView of Atomic Structure
Isotopes, Atomic Numbers, and Mass Numbers
Atomic number (Z) = number of protons in the nucleus.
Mass number (A) = total number of nucleons in the
nucleus (i.e., protons and neutrons).
A
By convention, for element X, we write Z X
Isotopes have the same Z but different A.
The ModernView of Atomic Structure
Isotopes, Atomic Numbers, and Mass Numbers
Periodic Table
The Periodic Table
Columns in the periodic table are called groups
(numbered from 1A to 8A or 1 to 18).
Rows in the periodic table are called periods.
Metals are located on the left hand side of the periodic
table (most of the elements are metals).
Non-metals are located in the top right hand side of the
periodic table.
Elements with properties similar to both metals and nonmetals are called metalloids and are located at the
interface between the metals and non-metals.
The Periodic Table
The Periodic Table
Some of the groups in the periodic table are given special
names.
These names indicate the similarities between group
members:
Group 1A: Alkali metals.
Group 2A: Alkaline earth metals.
Group 6A: Chalcogens.
Group 7A: Halogens.
Group 8A: Noble gases.
Molecules and Molecular Compounds
Molecules and Chemical Formulas
•Molecules are assemblies of two or more atoms bonded
together.
•The chemical formula indicates
•which atoms are found in the molecule, and in what proportion they are found.
•Compounds formed from molecules are called
molecular compounds.
Molecules and Molecular Compounds
Molecular and Empirical Formulas
Molecular formulas
give the actual numbers and types of atoms in a
molecule.
Examples: H2O, CO2, CO, CH4, H2O2, O2, O3, and
C2H4.
Empirical formulas
give the relative numbers and types of atoms in a
molecule.
That is, they give the lowest whole number ratio of
atoms in a molecule.
Examples: H2O, CO2, CO, CH4, HO, CH2.
Molecules and Molecular
Compounds
Picturing Molecules
If the structural formula does show the
shape of the molecule, then either a
perspective drawing, ball-and-stick
model, or space-filling model is used.
Molecular Naming
Two non-metals
Same side of Table
Naming 2 nonmetals
Binary Molecular Compounds
The most metallic element is usually written first
(i.e., the one to the farthest left on the periodic
table). Exception: NH3.
If both elements are in the same group, the lower
one is written first.
Greek prefixes are used to indicate the number
of atoms.
Naming Inorganic Compounds
Names and Formulas of Binary Molecular
Compounds
Ionic Compounds
Opposite sides of Periodic Table
MetalsNonmetals
Ions and Ionic Compounds
Ionic Compounds
The majority of chemistry involves the transfer of
electrons between species.
Example:
To form NaCl, the neutral sodium atom, Na, must lose
an electron to become a cation: Na+.
The electron cannot be lost entirely, so it is
transferred to a chlorine atom, Cl, which then
becomes an anion: Cl-.
The Na+ and Cl- ions are attracted to form an ionic
NaCl lattice which crystallizes.
Ions and Ionic Compounds
Important: note that there are no easily identified NaCl
molecules in the ionic lattice. Therefore, we cannot use
molecular formulas to describe ionic substances.
Ions and Ionic Compounds
When an atom or molecule loses electrons, it
becomes positively charged.
For example, when Na loses an electron it
becomes Na+.
Positively charged ions are called cations.
Ions and Ionic Compounds
When an atom or molecule gains electrons, it
becomes negatively charged.
For example when Cl gains an electron it
becomes Cl-.
Negatively charged ions are called anions.
An atom / molecule can lose more than 1 electron.
Naming Inorganic Compounds
Names and Formulas of Ionic Compounds
Name the cation then anion for the ionic compound.
Example: BaBr2 = barium bromide.
Ions and Ionic Compounds
Predicting Ionic Charge
The number of electrons an atom loses is related
to its position on the periodic table.
Metals tend to form cations whereas non-metals
tend to form anions.
Naming Inorganic Compounds
Names and Formulas of Ionic Compounds
Cations formed from a metal have the same name
as the metal.
Example: Na+ = sodium ion.
If the metal can form more than one cation, then
the charge is indicated in parentheses in the name
(usually Transition Metals).
Examples: Cu+ = copper(I); Cu2+ = copper(II).
Naming Inorganic Compounds
Names and Formulas of Ionic Compounds
Monatomic anions (with only one atom) are
called -ide.
Example: Cl- is chloride ion.
Exceptions: hydroxide (OH-), cyanide (CN-),
peroxide (O22-).
Polyatomic anions (with many atoms) containing
oxygen end in -ate or -ite. (The one with more
oxygen is called -ate.)
Examples: NO3- is nitrate, NO2- is nitrite.
Naming Inorganic Compounds
Names and Formulas of Ionic Compounds
Polyatomic anions containing oxygen with more
than two members in the series are named as
follows (in order of decreasing oxygen):
per-…..-ate
-ate
-ite
hypo-….-ite
ClO4ClO3ClO2ClO -
perchlorate
chlorate
chlorite
hypochlorite
Naming Inorganic Compounds
Names and Formulas of Ionic Compounds
Polyatomic anions containing oxygen with
additional hydrogens are named by adding
hydrogen or bi- (one H), dihydrogen (two H),
etc., to the name as follows:
CO32- is the carbonate anion
HCO3- is the hydrogen carbonate (or
bicarbonate) anion.
H2PO4- is the dihydrogen phosphate anion.
Naming Inorganic Compounds
Names and Formulas of Acids
The names of acids are related to the names of
anions:
-ide becomes hydro-….-ic acid;
-ate becomes -ic acid;
-ite becomes -ous acid.
PERIODIC
TRENDS
General Periodic Trends
 Atomic and ionic size
 Ionization energy
 Electron affinity
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.
Effective Nuclear Charge, Z*
(kernel Charge)
 Z* is the nuclear charge experienced
by the outermost electrons.
 Explains why Energy(2s) < Energy(2p)
 Z* increases across periodic table
Orbitals in Many Electron Atoms
Effective Nuclear Charge
- Electrons are attracted to the nucleus, but
repelled by the electrons that screen it from the
nuclear charge.
- The nuclear charge experienced by an electron
depends on its distance from the nucleus and the
number of core electrons.
- As the average number of screening electrons (S)
increases, the effective nuclear charge (zeff)
decreases.
zeff = z - S
48
Effective Nuclear Charge, Z*
 Atom
 Li
 B
 C
 N
 O
 F
Z* Experienced by 2s Electrons in
Valence Orbitals
+1.28
+2.58
Increase in Z*
+3.22
across a
+3.85
period
+4.49
+5.13
[Values calculated using Slater’s Rules]
General Periodic Trends
 Atomic and ionic size
 Ionization energy
 Electron affinity
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.
Atomic Size
 Size Increases going down a group.
 Because electrons are added further from the
nucleus, there is less attraction.
 Size Decreases going across a period. Electrons
are held closer.
Atomic Size
Size decreases across a period due to an
increase in Z*. Each added electron feels a greater
and greater (+) charge.
Large
Small
Increase in Z*
Atomic Radii
Questions
 Of the elements magnesium, Mg, Chlorine, Cl,
sodium Na, and phosphorus P which has the largest
atomic radius? Explain.
 Of the elements Ca, Be, Ba, and Sr, which has the
largest atomic radius? Explain.
 Of the elements Br, At, F, I and Cl which one has the
smallest atomic radius and which has the largest
atomic radius?
General Periodic Trends
 Atomic and ionic size
 Ionization energy
 Electron affinity
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.
Trends in Ionization Energy
1st Ionization energy (kJ/mol)
2500
He
Ne
2000
Ar
1500
Kr
1000
500
0
1
H
3
Li
5
7
9
11
Na
13
15
17
19
K
21
23
25
27
29
31
Atomic Number
33
35
Ionization Energy
IE = energy required to remove an electron from an
atom in the gas phase.
Mg (g) + 738 kJ ---> Mg+ (g) + e-
Ionization Energy
IE = energy required to remove an electron from
an atom in the gas phase.
Mg (g) + 738 kJ ---> Mg+ (g) + eCalled 1st IE
Mg+ (g) + 1451 kJ ---> Mg2+ (g) + eCalled 2nd IE
Ionization Energy
See Screen 8.12
Mg (g) + 735 kJ ---> Mg+ (g) + eMg+ (g) + 1451 kJ ---> Mg2+ (g) + e-
Mg2+ (g) + 7733 kJ ---> Mg3+ (g) + eEnergy cost is very high to dip into a shell of lower
n. Does not happen.
Trends in Ionization Energy
Trends in Ionization Energy
 IE increases across a period
because Z* increases.
 Metals lose electrons more
easily than nonmetals.
 Metals are good reducing
agents.
 Nonmetals lose electrons with
difficulty.
Trends in Ionization Energy
 IE decreases down a group
 Because size increases.
 Electrons are held less
closely
Questions
Consider the four hypothetical main group elements Q, R ,T, X
with the outer electron configurations indicated below. Then
answer the questions that follow.
Q = 3s23p5 R=3s1 T=4d105s25p5 X=4d105s25p1

a.
b.
c.
d.
Identify the block location of each hypothetical element.
Which of these elements are in the same period? Which are in
the same group?
Which element would you expect to have the highest first IE?
Which would have the lowest first IE?
Which element is most likely to form a 1+ ion?
General Periodic Trends
 Atomic and ionic size
 Ionization energy
 Electron affinity
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.
Electron Affinity
 New terms required
 Anion= negative charged atom
 Cation= positive charged atoms
 Easy to remember!!!
Electron Affinity
A few elements GAIN electrons to form
anions. Halogens are the easiest.
Electron affinity is the energy involved
when an atom gains an electron to form
an anion.
A + e- ---> AE.A. = ∆E
Electron Affinity Trends
 - numbers indicate an ease in adding electrons.
 Indicates how much energy is given off in the process.
 It is like a bank money
 (energy in, uses) = positive, spending money
 (energy out, releases) = negative numbers.
 Ex. EA of F =-339.9 kJ/mol (easy to do)
 EA of Mg = 0 kJ/mol (more difficult to do)
Electron affinity Trends Cont
 In general electrons are more difficult to
add going down a group. Because
 Nuclear charge increases
 Atomic radius increases down a group
decreasing electron affinity
Adding more electrons
 Is it possible to add more electrons to
something that is already negatively charged.
 Yes. But it is more difficult.
 Only certain cases when it creates a noble gas
configuration do electrons continue to add like
in O-2 or N-3
Ion Radii
Li,152 pm
3e and 3p
Does+the size of the
atoms
and ions go
+
Li , 60 pm
up
or down
when
2e and
3p
losing an electron
to form a cation?
Ion Sizes
+
Li,152 pm
3e and 3p
Li + , 78 pm
2e and 3 p
Forming a
cation.
 CATIONS are SMALLER than the atoms from
which they come.
 The electron/proton attraction has gone UP and
so size DECREASES.
 Atoms also decrease in size.
Ion Sizes
Does the size of the atoms and ions go up or down
when gaining an electron to form an anion?
Ion Sizes
F, 71 pm
9e and 9p
F- , 133 pm
10 e and 9 p
Forming
an anion.
 ANIONS are LARGER than the atoms from
which they come.
 The electron/proton attraction has gone
DOWN and so size INCREASES.
 Atom size also increases.
 Trends in ion sizes are the same as atom sizes.
Trends in Ion Sizes
Active Figure 8.15
Atomic Radii Size Trends
 Metals at the left end tend to form
cations
 Nonmetals at the upper right tend to
form anions.
Electronegativity
 A measure of the ability of an atom in a chemical
compound to attract electrons.
 Most electronegative is Fluorine.
 Least electronegative are alkali and alkaline
metals
 Most electronegative are the halogens nitrogen
and oxygen.
Trends
Electronegativty is strongest at the
upper right of periodic table
Weakest in the lower right.
Decrease or stay the same down a
group.
Electronic Structure of
Atoms
Chapter
7
78
The Wave Nature of Light
- Visible light is a small portion of the
electromagnetic spectrum
79
The Wave Nature of Light
Frequency (v, nu) – The number of times per second that one complete
wavelength passes a given point.
Wavelength (l, lambda) – The distance between identical points on successive
waves.
80
lv=c
c = speed of light, 2.997 x 108 m/s
V=frequency (1/s = hertz= Hz)
l =wavelenth (m)
The Wave Nature of Light
- We can also say that light energy is quantized
- This is used to explain the light given-off by hot
objects.
- Max Plank theorized that energy released or
absorbed by an atom is in the form of “chunks” of
light (quanta).
E=hv
h = plank’s constant, 6.63 x 10-34J/s
E= Energy in system (Joules= kg . m2/s2)
V= frequency (Hz=1/s)
- Energy must be in packets of (hv), 2(hv), 3(hv), etc.
81
Quantized Energy and Photons
The Photoelectric Effect
82
Quantized Energy and Photons
Einstein’s The Photoelectric Effect
- The photoelectric effect provides evidence for the
particle nature of light.
- It also provides evidence for quantization.
- If light shines on the surface of a metal, there is
a point at which electrons are ejected from the
metal.
- Below the threshold frequency, no electrons are
ejected.
- Above the threshold frequency, the number of
electrons ejected depend on the intensity of the
light.
83
Quantized Energy and Photons
The Photoelectric Effect
- Einstein assumed that light traveled in energy
packets called photons.
- The energy of one photon, E = hn.= hc/ l
- This equation means that the energy of the
photon is proportional to its frequency.
84
Chapter 6
De Broglie’s Contribution
 Thought:
 if waves show particle properties, can particles show wave
properties?
 Answer is yes
 Mass=h / l x v (Einstein) so,
 Rearrrange
 l = h / mass x velocity
Example
 Find the wavelength for a ball with a mass of .10 kg thrown at 35 m/s.
 l = h / mass x velocity
 Plug in numbers,
 6.626 x
.m/s
10-34 kg.m
(.10 kg) ( 35 m/s)
= 1.9 x 10-34 m
Conclusions
 All matter exhibits both particulate and wave properties.
 Large mater exhibit mostly particulate properties
 ex. Baseballs, footballs, cannon balls, Rocks
 Small matter exhibit mostly wave properties
 ex. Photons,
 Intermediate matter exhibit like electrons demonstrate both.
Quantum Model of the Atom
Bohr’s Model of the Hydrogen Atom
Bohr’s Model
- Assumed that a single
electron moves around
the nucleus in a circular
orbit.
- The energy of a given
electron is assumed to be
restricted to a certain
value which corresponds
to a given orbit.
89
Bohr’s Model of the Hydrogen Atom
Line Spectra for light is just like hydrogen
90
Chapter 6
Bohr’s Model of the Hydrogen Atom
Bohr’s Model
- Since the energy states are quantized, the light emitted
from excited atoms must be quantized and appear as
line spectra.
91
Chapter 6
Bohr’s Model of the Hydrogen Atom
Line Spectra
Line spectra can be “explained” by the following
equation:

-18  Z
E  - 2.179 x10 J 2
n

n=energy level
Z= number of protons
92




Bohr’s Model of the Hydrogen Atom
Bohr’s Model – Important Features
- The first orbit in the Bohr model has n = 1
and is closest to the nucleus.
- The furthest orbit in the Bohr model has n
close to infinity and corresponds to zero
energy.
- Electrons in the Bohr model can only
move between orbits by absorbing and
emitting energy in quanta (hn).
93
Bohr’s Model of the Hydrogen Atom
Bohr’s Model – Line Spectra
Ground State – When an electron is in
its lowest energy orbit.
Excited State – When an electron
gains energy from an
outside source and moves
to a higher energy orbit.
94
Chapter 6
More Equations
 ΔE= energy of final – energy of initial
 Will tell us if an atom has gained or lost energy.
 Negative sign means more stable and a loss of
energy.
 Plus sign means less stable and a gain of energy.
 From that we can then determine the wavelength
of that emitted photon using the equation:
 l hc/ ΔE
Example
 Calculate the energy required to excite the hydrogen electron from level
1 to level 2.
 Use equation
Z
E  - 2.179 x10 -18 J  2
n





-2.179 x 10-18 (12/12) = -2.179 x 10-18 J
 For n=2 -2.179 x 10-18 (12/22) = -5.445 x 10-19 J
 For n=1
 So, ΔE=
energy of final – energy of initial
 ΔE= (-5.445 x 10-19 J) - (-2.179 x 10-18 J)
 1.633 x 10-18 J (+ sign means energy is absorbed)
2 Things to Remember About the
Bohr Model
 Only works with Hydrogen
 As you get closer to the ground state, energy is being
released.
 This was a good
start, and helps us
To understand the
atom, it was
fundamentally
Incorrect.
The Wave Behavior of Matter
- Remember DeBroglie proposed that there is a
wave/particle duality.
- Knowing that light has a particle nature, it seems
reasonable to assume that matter has a wave nature.
- DeBroglie proposed the following equation to describe
the relationship:
h
l
mv
98
The Wave Behavior of Matter
The Uncertainty Principle
Heisenberg’s Uncertainty Principle - on the
mass scale of atomic particles, we cannot
determine exactly the position, speed, and
direction of motion simultaneously.
- For electrons, we cannot determine their
momentum and position simultaneously.
99
Quantum Mechanics
- These theories (wave/particle duality
and the uncertainty principle) mean
that the Bohr model needs to be
refined.
 Quantum Mechanics 
100
Quantum Mechanics
- The path of an electron can no longer be
described exactly, now we use the
wavefunction(y).
Wavefunction (y) – A mathematical expression to
describe the shape and energy of an electron in
an orbit.
Probability density = y2
101
Quantum Mechanics
Quantum Numbers
- The use of wavefunctions generates four
quantum numbers.
- Principal Quantum Number (n)
- This is the same as Bohr’s n
- Allowed values: 1, 2, 3, 4, … (integers)
102
Quantum Mechanics
Quantum Numbers
Secondary (Azimuthal) Quantum Number (l)
- Allowed values: 0, 1, 2, 3, 4, . , (n – 1)
(integers)
- Each l represents an orbital type
103
l
orbital
0
s
1
p
2
d
3
f
Quantum Mechanics
Quantum Numbers
Magnetic Quantum Number (ml ).
- This quantum number depends on l.
- Allowed values: -l  +l by integers.
- Magnetic quantum number describes the
orientation of the orbital in space.
104
l
Orbital
ml
0
s
0
1
p
-1,
2
d
-2, -1,
0, +1
0, +1, +2
Quantum Mechanics
Quantum Numbers
Spin Quantum Number (s)
- Allowed values: -½  +½.
- Electrons behave as if they are spinning
about their own axis.
- This spin can be either clockwise or counter
clockwise.
105
Chapter 6
Quantum Mechanics
Orbitals and Quantum Numbers
106
Representation of Orbitals
The s Orbitals
- All s-orbitals are spherical.
- As n increases, the s-orbitals get larger.
- As n increases, the number of nodes
increase.
- A node is a region in space where the
probability of finding an electron is zero.
107
Representation of Orbitals
The s Orbitals
108
Representation of Orbitals
The p Orbitals
- There are three p-orbitals px, py, and pz. (The
letters correspond to allowed values of ml of -1,
0, and +1.)
- The orbitals are dumbbell shaped.
109
Representation of Orbitals
The p Orbitals
110
Representation of Orbitals
The d and f Orbitals
- There are 5 d- and 7 f-orbitals.
- Four of the d-orbitals have four lobes each.
- One d-orbital has two lobes and a collar.
111
Representation of Orbitals
The d and f Orbitals
112
Electron configuration
 Electron configuration is the arrangement of electron in an
atom.
 All atoms have different arrangements
 Like everything in nature, they want to have the lowest
possible energy called ground states.
3 Rules Determining Ground State
Electron Configurations
 Aufbau principle
 Pauli excusion principle
 Hund’s rule
Aufbau Principle
 Electrons will occupy the lowest energy orbital that can
receive it.
 Notice the order changes after 3p. It does not always go in
order.
Help is on the way!!!
 If you can count to seven
then you can determine
which energies are required.
Or
Use your periodic table
Electron Configurations and the
Periodic Table
117
Chapter 6
Pauli Exclusion Principle
 No two electrons can have the same 4 quantum numbers.
 This means that the numbers may look the same, but the spin
will be different for electrons that occupy different orbitals.
 Ex.
1s2 2s2 2p3
1s2 2s2 2p6 3s2 3p6
Hund’s Rule
 Orbitals of equal energy are each occupied by one electron
before any orbital is occupied by a second electron, and all
electrons in singly occupied orbitals must all have the same
spin.
 Ex.
OR
3 Methods to representing electron
configurations
 Orbital notation
 Electron configuration notation
 Noble Gas notation
Orbital Notation
 An unoccupied orbital is represented by a line with the
orbital name under it. Electron spins are then entered.
 Ex.
 H
or He
1s1
1s2
Electron configuration notation
 Eliminates the lines and arrows of orbital notation.
 The number of electrons in a sublevel is shown by adding a superscirpt
to the sublevel designation.
 Ex. H 1s1
He 1s2
Be
1s2 2s2
Noble gas notation
 Same as electron configurations, but it uses noble gases to
shorten the work.
 Noble gases are placed in brackets to indicate full electron
shells and then the proper amount of energy levels is added
after it.
 Ex.
K [Ar] 4s1
Or Zn [Ar]3d10 4s2