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Unit 5: Periodic Table Notes
Name _____________________________________
Period ________ Date _______________________
6.1: Development of the Periodic Table
1) History of the Periodic Table – By the end of the 1700’ scientists had identified only 30 elements
By the mid 1800’s with the assistance of electricity and spectroscopy
additional elements were identified by using their line spectra and about
65 elements had been identified.
A. J.W. Dobereiner: Organized the elements into ______________ with similar _____________.
He called these groups ________________.
The middle element is often the
____________________ of the other two.
Ex)
B. J.A.R. Newlands: Law of _______________________.
He said properties repeated every
_________________ element. This worked for
________________ elements but did not work well for
_______________ elements.
C. Dmitri Mendeleev:
1. Organized the 1st periodic table according to increasing ________________________
and put elements with similar properties in the same ___________________.
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2. He rearranged some elements out of atomic mass _______________ to keep them
together with other elements with similar _____________________. He also left
three ________________ in his table and correctly __________________ the properties
of these 3 unidentified elements that were later identified and matched his predictions.
D. Henry Moseley: Each element has a certain amount of positive charge in the nucleus which
are called _________________.
1. Mosely reorganized the periodic table by _________________ __________________.
2. The Periodic Law:
When elements are arranged in order of increasing _______________
___________________, their physical and chemical ________________________
show a ________________________ pattern.
6.2: Reading the Periodic Table
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A. Information in each square:
Al
Aluminum
26.9815
B. Organizing the Squares

Vertical Columns – called ___________ or _____________. Elements in the same
group have similar _________________ ___________________.

Horizontal Rows – called __________________.

The modern periodic table has ________ groups and _________ periods (just like the
school day!)
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C. Electron Configuration & Families
1. ______________ electrons are the electrons involved in chemical bonds.
2. Elements in the same ____________ have the same number of valence electrons which
results in them having similar chemical properties.
Group number
1
2
13
14
15
16
17
18
Number of valence eLocation of valence e-
D. Family Names – Label the following regions on a periodic table:
alkali metals, alkaline earth metals, transition metals, coinage
metals, halogens, noble gases, lanthanides, actinides, main group elements.
E. Metals vs. Nonmetals
F. Metalloids (def)
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ChemQuest 5A Notes
A - Coulombic Attraction – The force of attraction between oppositely charged particles such as
__________ and _____________.
B) Two main trends for coulombic attraction between the nucleus and the valence electrons
1) If two elements are in the same vertical group (family), ________________ is the dominant
factor. The more shells of electrons that atoms have, the farther away the valence electron(s)
are from the protons in the nucleus, the ________________ the attraction will be.
2) If two elements are not in the same family, but are in the same horizontal period, then the
number of __________________ in the nucleus is the dominant factor.
The more protons in
the nucleus, the greater the force of attraction (to be explained more in CQ5B)
ChemQuest 5B Notes – Atomic Size
A. Effective Nuclear Charge – the net attractive force that pulls the valence electrons towards the
nucleus.
Group number
1
2
13
14
15
16
17
18
Effective nuclear charge
B. Shielding Effect – when atoms have more than one energy level of electrons, the core electrons
(electrons in the inner shells) shield/block the valence electrons from feeling the full positive
charge of the nucleus. Each core electron shields the valence electrons from 1+ charge of
the nucleus.
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C. Bohr Model – We simplify the electron configuration of atoms to better understand atomic
trends and chemical bonding. In Bohr models we draw electrons in shells (energy levels).
Energy Level
Maximum e-
Sublevels
1
2
1s2
2
8
2s22p6
3
18
3s23p63d10
4
32
4s24p64d104f14
Application: In the space below, draw a Bohr Model of calcium (Ca) on the left and chlorine
(Cl) on the right. Identify the core electrons and valence electrons. Determine the effective
nuclear charge in each atom.
6.3: Trends in the Periodic Table
A. Atomic Radius
1. The distance from ________________________________________________
____________________________________________________________
2. Atoms get ______________________ going down a group and
______________________ going across a period.
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Why do atoms get bigger as you go down a group?
Why do atoms get smaller as you go across a period from left to right?
B. Ionic Size
1. When atoms _________ electrons, they become _______ and get _______________.
2. When atoms _________ electrons, they become _______ and get _______________.
3. Ions get _____________________ as you go down a group.
Summary: When an atom adds additional electrons, it makes it harder for the protons to
hold on to the extra electrons which results in a bigger radius. Removing electrons from an
atom, makes it easier for the protons to hold on to the remaining electrons so the atom gets
smaller.
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C. Ionization Energy:
1. The energy need to __________________ one electron from an atom.
2. Elements that do not want to lose their electrons have ______________ ionization
energies.
3. Elements that easily lose electrons have ___________ ionization energies.
4. I.E. _____________________ down a group. (valence e- get farther from nucleus)
5. I.E. _____________________across a period. (ENC increases)
IONIZATION ENERGY VALUES
Questions –
1) Why do metals tend to lose electrons and form positive ions (cations)?
2) Noble gases do not lose electrons to other atoms. Why?
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D. Successive Ionization Energies:
1. Energy required to ___________________ electrons beyond the 1st electron.
2. Ionization energies will _______________ for every electron removed. In other
words, it will get ______________ to steal each additional electron.
3. Na [Ne]3s1
Na•
1st = ______ kJ
2nd = ______ kJ
Why is it so much harder to remove the 2nd electron from sodium?
4. Mg [Ne]3s2
Mg:
1st = ______ kJ 2nd = ______ kJ 3rd = ______kJ
Why is it so much harder to remove the 3rd electron from magnesium?
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E. Electronegativity (EN):
1. Reflects an atom’s ability to _______________ electrons involved in a chemical bond.
2. The higher the EN value, the ____________ the ability of the atom to pull electrons in a
chemical bond towards its nucleus.
3. The lower the EN value, the ____________ the ability of the atom to pull electrons in a
chemical bond towards its nucleus.
4. Metals tend to have ______ EN values and nonmetals tend to have ______ EN values.
This helps explain why metals tend to ________ electrons and form __________ ions
and nonmetals tend to ________ electrons and form __________________ ions.
ELECTRONEGATIVITY CHART
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G. Octet Rule:
1. An atom will tend to ______________, ______________ or _______________
electrons in order to acquire a full set of valence electrons.
2. “Octet” = “happiness” = “stability” = 8 valence electrons = s2p6 configuration
H. Oxidation Number:
Group
1
2
13
14
15
16
17
18
Oxidation
Number
What will it do to
become like a
noble gas?
FINAL QUESTIONS – Draw Bohr model to support your answers.
 Why is oxygen chemically similar to sulfur?
 Why is oxygen considered to be more reactive than sulfur?
 Why is potassium more reactive than calcium?
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