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Transcript
```Lesson 6
The Periodic Table
Anything in black letters = write it in
Objectives for Lesson 6
Section 1 – Organizing the Elements
Section 2 – Periodic Trends
1. Describe ways in which the modern periodic table is organized
2. Understand electron configuration patterns in the periodic table
3. Describe and explain trends in the periodic table
Section 1 – Organizing the Elements
Dmitri Mendeleev (1869) – created 1st modern
periodic table.
Mendeleev arranged
elements with similar
properties.
Ga & Ge
Discovered later
Similar
properties
He also left gaps
where proposed
elements should be.
These gaps were
later filled in as more
elements were
discovered.
Mendeleev’s table was an accepted success
because it predicted the properties of elements
that had not yet been discovered.
Woo Hoo!
Today’s periodic table is arranged in order of
increasing atomic number (not mass).
Also, elements with similar chemical properties
are placed in the same vertical column.
Columns are called groups or families.
Horizontal rows are called periods.
Valence Electrons – Electrons in the highest
occupied energy level; maximum of 8.
Elements in the same column have similar
properties because they have the same
number of valence electrons.
Electrons in the s and p orbitals of the outer shell are
the valence electrons.
8 is the maximum number of valence electrons
The Octet Rule –
Atoms tend to gain or lose electrons to have 8 e-
Sodium:
1s22s22p63s1
Magnesium: 1s22s22p63s2
Fluorine:
1s22s22p5
Nitrogen:
1s22s22p3
The noble gases are chemically stable because they
have a full outer energy level (valence).
Atoms tend to gain or lose electrons to have 8 e-
Sodium:
1s22s22p63s1
Magnesium: 1s22s22p63s2
Fluorine:
1s22s22p5
Nitrogen:
1s22s22p3
Electron configurations for Group 1
1s1
(valence e- underlined)
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
[Xe]6s1
[Rn]7s1
1s22s22p5
1s22s22p63s23p5
Get the idea?...
Why is it called the Periodic Table
of the Elements?
The properties of the elements repeat
going across each row.
metals, metalloids, nonmetals
Metals – good conductors of heat and
electricity, shiny, most are solid at
room temp (except Hg), malleable,
ductile
Nonmetals – not metals!, most are
gases at room temp
Metalloids – can show properties of
both metals and nonmetals
Practice
1.
Explain why Mendeleev’s table was an accepted
success.
2.
Why is the table of elements called the “periodic”
table of elements?
3.
State 4 properties of metals.
4.
Explain the reason that elements in the same
column have similar chemical properties?
5.
How can you tell if an elements is a metal,
nonmetal or metalloid from the periodic table?
6. Name an element that is part of the
a)
Halogen family
b)
Alkali metal family
c)
Alkaline earth metal family
d)
Transition metals
e)
Inner transition metals
f)
Noble gas family
7. A horizontal row in the periodic table is called a
_____.
8. Write the electron configuration for
a)
Nitrogen
b)
Chlorine
c)
Rubidium
9. How many valence electrons are in each element
from question 8?
Section 2 – Periodic Trends
Atomic size
Ionic size (skip!)
Ionization Energy
Electronegativity
Atomic Size
Atomic number
Atomic size generally decreases from left to right
across a period.
As Z increases across a row, the +/- electrical
attraction increases, making the atom smaller.
As Z increases down a group, another energy
level is added to the atom which ‘shield’ the
outer electrons from this nuclear attraction.
Ion – atom or group of atoms that has a positive or
negative charge.
Ions are formed when electrons are transferred
between atoms.
Cation – ion with a positive charge.
Anion – ion with a negative charge.
Metals tend to form cations
Nonmetals tend to form anions
Ionization Energy – energy required to remove an
electron from an atom.
lithium +1 ion
lithium atom
1st ionization energy = 520 kJ/mol
lithium ion
Lithium +2 ion
2nd ionization energy = 7297 kJ/mol
Ionization Energy – energy required to remove an
electron from an atom.
lithium +1 ion
lithium atom
1st ionization energy = 520 kJ/mol
lithium ion
Lithium +2 ion
2nd ionization energy = 7297 kJ/mol
Ionization Energies of Some Common Elements
Symbol
First
Second
Third
H
1312
He (noble gas)
2372
5247
Li
520
7297
11,810
Be
899
1757
14,840
C
1086
2352
4619
O
1314
3391
5301
F
1681
3375
6045
Ne (noble gas)
2080
3963
6276
Na
496
4565
6912
Mg
738
1450
7732
S
999
2260
3380
1520
2665
3947
K
419
3096
4600
Ca
590
1146
4941
Ar (noble gas
Atomic number
First ionization energy (kJ/mol)
1. What does 1st ionization energy mean?
2. Explain why the 2nd ionization energy of Li and Na
is so much higher than the 1st ionization energy.
3. Explain why the 1st ionization energy of Na is
smaller than Li.
4. Why are the ionization energies of the noble gases
so large?
Electronegativity – tendency of an atom to attract
electrons of another atom.
Metals have low e-neg values,
Nonmetals have high e-neg values
B<H<C
Noble gases do not have e-neg values
Electronegativity Values for Selected Elements
H
2.1
Li
1.0
Be
1.5
B
2.0
C
2.5
N
3.0
O
3.5
F
4.0
Na
0.9
Mg
1.2
Al
1.5
Si
1.8
P
2.1
S
2.5
Cl
3.0
K
0.8
Ca
1.0
Ga
1.6
Ge
1.8
As
2.0
Se
2.4
Br
2.8
Rb
0.8
Sr
1.0
In
1.7
Sn
1.8
Sb
1.9
Te
2.1
I
2.5
Cs
0.7
Ba
0.9
Tl
1.8
Pb
1.9
Bi
1.9
Lesson 6 Practice
1.
How does the size of an atom change from left to
right a) across a period? b) down a column?
2.
Give the explanation for question 1.
3.
What is an ion and how are they formed?
4.
Metals tend to form _____ ions and nonmetals
tend to form _____ ions.
5.
Define Ionization Energy.
6.
Describe the trend in ionization energy in the
periodic table.
7.
Define electronegativity.
8.
Which atom is the a) most electronegative, b)
least electronegative?
9.
Which atom has the highest ionization energy?
10. How do electronegativity values differ between
metals and nonmetals?
A little more for Chapter 6…
Lesson 6 Quiz Review
Terms to know:
valence electron,
cation,
anion,
electronegativity,
ionization energy (1st & 2nd)
Things to know:
Metal, nonmetals, metalloids locations
4 properties of metals
metals form cations, nonmetals form anions
family names (alkali, alkaline earth, noble,
halogens, transition and inner transition)
electronegativity and ionization energy trends
electron configurations (w/out aufbau diagram)