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Lecture 4
Periodicity, Ionization Energy and the
proposed “Shell” structure of the atom;
Ch 3, 4.1-4.3
Dr. Harris
Suggested HW: (Ch 3) 4, 28
(Ch 4) 1, 4, 5, 12
Chemical Reactions
• When elements undergo a chemical reaction, the products may be
quite different from the reactants
• The simplest reactions are those between metals and nonmetals.
The product of such a reaction is an ionic compound
• Lets consider the reaction between sodium metal and chlorine gas
Stark Differences Between Reactants
and Products
𝑁𝑎 𝑠 + 𝐶𝑙2 𝑔 → 𝑁𝑎𝐶𝑙 (𝑠)
• As you can see from the chemical equation
above, products can exhibit physical
characteristics that are vastly different from
those of the reactants
• Recall the law of conservation of mass. Based
on this law, can you find a problem with the
equation written above?
Balancing Reactions
• Mass can not be created or destroyed. This means that every element
involved in a reaction must be accounted for in a chemical equation.
𝑁𝑎 𝑠 + 𝐶𝑙2 𝑔 → 𝑁𝑎𝐶𝑙 (𝑠)
• As you can see, there are two chlorine atoms on the reactant side, and
only one chlorine atom one the product side. To balance the chlorine
atoms, we add a coefficient of 2 to the NaCl(s)
𝑁𝑎 𝑠 + 𝐶𝑙2 𝑔 → 𝟐 𝑁𝑎𝐶𝑙 (𝑠)
• We have balanced the chlorine atoms, but the sodium atoms are now
unbalanced. We add a coefficient of 2 to the Na (s). The reaction is now
balanced.
𝟐 𝑁𝑎 𝑠 + 𝐶𝑙2 𝑔 → 𝟐 𝑁𝑎𝐶𝑙 (𝑠)
Coefficients vs. Subscripts
𝟐 𝑁𝑎 𝑠 + 𝐶𝑙2 𝑔 → 𝟐 𝑁𝑎𝐶𝑙 (𝑠)
• The balanced equation above says that two Na atoms react with one
chlorine gas molecule to produce two molecules of NaCl
• The coefficient of 2 means that there are two separate Na atoms
• The subscript of 2 indicates two Cl atoms bonded together in a single
molecule
Na (s)
Cl
Cl
Na (s)
• Do not confuse coefficients and subscripts
NaCl (s)
NaCl (s)
Balancing Equations
• Before carrying out any calculations, it is imperative that you first
confirm that a given chemical equation is balanced.
• The rules for balancing a chemical equation are provided below.
1.
2.
3.
First, balance those elements that appear only once on each
side of the equation
Balance the other elements as needed. Pay attention to
subscripts.
Include phases
Balancing Equations
• Let’s balance the equation below using the rules from the previous
slide.
C3H8 (s) + O2 (g)
CO2 (g) + H2O (L)
• We’ll balance C first.
C3H8 (s) + O2 (g)
3 CO2 (g) + H2O (L)
• Now balance H.
C3H8 (s) + O2 (g)
3 CO2 (g) + 4 H2O (L)
• Now balance O.
C3H8 (s) + 5 O2 (g)
3 CO2 (g) + 4 H2O (L)
Group Examples
• Balance the following
Chemical Groups
• As more and more elements were discovered, chemists began to
notice patterns in the chemical properties of certain elements.
• Consider the three metals Li, Na, and K
• All 3 metals are soft
• All 3 metals are less dense than water
• All 3 metals have similar appearance and low melting points
• The most interesting feature is that all 3 metals react with the
same elements in a nearly identical manner
• As you see in the periodic table, these elements are all
listed in the same group.
• Elements in a group behave similarly. Recognizing patterns allows
us to predict reactions without memorizing every characteristic of
every element
http://www.youtube.com/watch?v=qRmNPKVEGeQ&feature=related
http://www.youtube.com/watch?v=MTcgo46nxNE
Periodicity
• Dmitri Mendeleev created the periodic table in in 1869 by arranging
the elements in order of increasing atomic mass.
• In doing so, he observed repetitive patterns in chemical behavior
across periods
• This periodicity is described in the next slide.
Periodicity
2
He
1
H
3
Li
Highly reactive,
highly conductive
metal
Totally
unreactive gas
4
Be
Decreasing metallic
character
Less reactive,
less conductive
metal
6
C
Nonconductive,
nonmetallic
solid
9
F
Highly reactive,
diatomic,
nonmetallic gas
Decreasing metallic
character
11
Na
Highly reactive,
highly conductive
metal
19
K
Highly reactive,
highly conductive
metal
14
Si
12
Mg
Less reactive,
less conductive
metal
20
Be
Less reactive,
less conductive
metal
Slightly
conductive
semi-metal
Decreasing metallic
character
22
Ge
Slightly
conductive
semi-metal
17
Cl
Highly reactive,
diatomic,
nonmetallic gas
25
F
Highly reactive,
diatomic,
nonmetallic liq.
10
Ne
Totally
unreactive
gas
18
Ar
Totally
unreactive
gas
26
Kr
Totally
unreactive
gas
Mendeleev’s Genius
• At the time in which the periodic table was being constructed, not
all of the elements had been discovered.
• Based on the observed periodicity, Mendeleev realized that gaps in
the initial periodic table belonged to undiscovered elements
• For example, in 1869, the element following Zn on the periodic
table was As. Yet, he knew to put As in group 15 rather than 13
because As behaved like P, and he knew that two undiscovered
elements (Ga and Ge) would fill the gaps.
Transition Metals
transitions metals
• Transition metals
span the region
where the
transition from
metal to nonmetal
occurs.
• Transition metals
are very dense and
have very high
melting points.
Semiconductors
Intro to Ch 4
• In ch. 4, we begin to answer many questions about chemical reactivity
• Why is it that some atoms join together and form molecules, while
others don’t?
• Why is there such wide variation in the reactivity and physical
properties of elements?
• Why is there periodic repetition (periodicity) of the chemical/physical
properties of elements as we move across the periodic table?
How to Interpret the Findings of Mendeleev
• As previously discussed, Mendeleev noticed that chemical behavior
was repeated periodically when elements were sorted by increasing
atomic number
• The existence of periodicity proves a very important point:
Atomic number, and therefore, atomic mass, has no effect on
chemical behavior. Otherwise, chemical behaviors would never
repeat.
Therefore, the chemical behavior of an element must be due to the
configuration of electrons around the nucleus.
Ionization Energy
• A direct indication of the arrangement of electrons about a nucleus is
given by the ionization energies of the atom
• Ionization energy (IE) is the minimum energy needed to remove an
electron (form a cation) completely from a gaseous atom
• Ionizations are successive.
• As you remove one electron, it becomes increasingly difficult to remove
the next because of the increasing attraction between the remaining
electrons and the protons in the nucleus
𝑀 → 𝑀+ + 𝑒 −
1st Ionization Energy
𝑀+ → 𝑀2+ + 𝑒 −
2nd Ionization Energy
IE1 < IE2 < IE3 …….IEn
What Can Ionization Energy Tell Us
About Chemical Behavior?
• By measuring the energy required to remove electrons from an
element, you can gain an idea of how “willing” an atom is to lose an
electron, and relate this to its reactivity
• In the next slide, you will see data from an experiment in which the
1st ionization energies of elements are plotted against atomic
number.
1st Ionization Energies Show A Periodic Trend
Trends in 1st Ionization Energies
• It is relatively easy to remove
electrons from group 1 metals.
• It becomes increasingly
difficult as you move right
across the periodic table, and
up a group.
• It takes a very large amount of
energy to ionize a noble gas.
• Like chemical properties,
ionization energies are also
periodic.
The lower the ionization energy of
an element, the more METALLIC
and REACTIVE it is.
Electron Arrangement
• The closer an electron is to the nucleus, the harder it would be to
pull the electron away.
• By carrying out multiple ionizations, we can gain insight into the
arrangement of electrons around the nucleus of the element.
Example
• Using the table of ionization energies in the previous slide, calculate
the energy required to ionize Be to Be3+
• In order to go from Be to Be3+, you must LOSE 3 electrons.
This will require 3 ionization steps (see pg 107 in book).
𝐵𝑒 𝑔 → 𝐵𝑒 + 𝑔 + 𝑒 −
𝐼𝐸1 = 1.49 𝑎𝐽
𝐵𝑒 + 𝑔 → 𝐵𝑒 2+ 𝑔 + 𝑒 −
𝐼𝐸2 = 2.92 𝑎𝐽
𝐵𝑒 2+ 𝑔 → 𝐵𝑒 3+ 𝑔 + 𝑒 −
𝐼𝐸3 = 24.7 𝑎𝐽
29.1 aJ
Remember, energy is always in Joules (J). atto (a) = 10-18
Graph of Successive Ionization Energies
• Let’s take a look at the electron configurations of Lithium (atomic # = 3)
and Beryllium (atomic # = 4)
Li
3 electrons
Single electron
that is easily
removed
Be
4 electrons
Pair of tightly
bound
electrons
Pair of electrons
that are more
easily removed
Larger Atomic Numbers
Ne
10 electrons
Na
11 electrons
Eight electrons of
similar attraction to the
nucleus
Same two tightly
bound electrons
11th electron
enters different
“shell”
Electrons reside in “shells” of different
distances from the nucleus
• From these plots, Niels Bohr derived the Bohr model of the atom. In it,
electrons reside in shells that orbit at different distances from the nucleus.
• Each shell has a finite number of electrons that it can hold
• The two electrons closest to the nucleus are the hardest to remove
Na
atom
• Each shell holds 2n2
electrons, where the
n=1 shell is the closest
to the nucleus.
Same Outer Electron Configuration Along A Group Leads
to Similarities in Reactivity
Na
Li
K
All group 1 metals have a lone
electron in the outermost shell.
Chemical properties of an
element are determined by the
outer electron configuration.
Periodicity is Due To Repeating Valence Electron
Configurations
Li
Be
B
C
N
O
F
Ne
Na
Mg
Al
Si
P
S
Cl
Ar
Noble Gas Configurations
• The inner-most electrons of an element comprise what is known as
a noble gas core.
• At the close of each shell, you have a noble gas configuration.
Noble gases are chemically inactive because they have
completely filled shells.
• Lithium, for example, has a two electron core, which we call a
Helium core, and one outer, or valence electron. Sodium has a 10electron, Neon core, and one valence electron; and so on.
• The electron configuration of an element can be represented with a
Lewis dot formula
Full Lewis dot
configuration
Valence Lewis dot
configuration
We use these
representations to
describe the electron
configurations of an
element.