Transcript Periodic Table & Bonding Chapter 6

UNIT 3 – Periodic Table & Bonding
Chapter 6 – The Periodic Table
Chapter 7 – Ionic Bonding
Chapter 8 – Covalent Bonding
Chapter 6
The Periodic Table
Anything in black letters = write it in
your notes (‘knowts’)
Objectives for Chapter 6
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
6.1 – Organizing the Elements
Dmitri Mendeleev (1869) – created 1st modern
periodic table; arranged elements based on
atomic mass and chemical properties.
Ga & Ge
Discovered later
Mendeleev arranged
elements with similar
properties in the
same row.
He also left gaps
where proposed
elements should be.
Similar
properties
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.
Chapter 7
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.
Electron configurations for Group 1
1s1
(valence e- underlined)
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
[Xe]6s1
[Rn]7s1
1s22s22p5
1s22s22p53s23p5
Get the idea?...
Why is it called the Periodic Table
of the Elements?
The properties of the elements repeat
going across each row.
Three broad classes of elements;
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
6.2 – Classifying the Elements
Columns are called groups or families.
Horizontal rows are called periods.
Chapter 6 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.
How is the modern periodic table arranged?
4.
State 4 properties of metals.
5.
What is the explanation for the reason elements
in the same column have similar chemical
properties?
6.
How can you tell if an elements is a metal,
nonmetal or metalloid from the periodic table?
7.
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
8.
A horizontal row in the periodic table is called a
_____.
9.
Write the electron configuration for
a) Nitrogen
9.
Write the electron configuration for
a)
Nitrogen
b)
Chlorine
c)
Rubidium
10. How many valence electrons are in each element
from question 1?
Chapter 6 ASSIGNMENT (p. 166-173)
#1-17
6.3 – Periodic Trends
Atomic size
Ionic size
Ionization Energy
Electronegativity
Atomic radius (pm)
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, more energy
levels are in 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 + ions (cations)
Nonmetals tend to form - ions (anions)
Atom
Ion
Atom
Ion
Atom
Ion
Li
Li+
Be
Be2+
O
O2-
Na Na+
Ca
Ca2+
S
S2-
Sr
Sr2+
Se
Se2-
Ba
Ba2+
Te
Te2-
K
K+
Rb Rb+
Ionic Size
Cations are smaller than the atoms they formed from
Anions are larger than the atoms they formed from.
Ionic Size
Ionization Energy – energy required to remove an
electron from an atom.
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)
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
Chapter 6 ASSIGNMENT (p. 182)
#18-25
Chapter 6 ASSIGNMENT (p. 188)
#26, 27, 30, 34,
39-42, 45-47, 49-51, 53, 62, 64,
75
A little more for Chapter 6…
Energy levels can also be called electron shells
Each shell corresponds to a period on the table.
2
8
8
18
18
32
32
Electrons in the s and p orbitals of the outer shell are
the valence electrons.
8 is the maximum number of valence electrons
The noble gases are chemically stable because they
have a full outer shell (valence).
Atoms tend to gain or lose electrons to have a full shell
Sodium:
1s22s22p63s1
Magnesium: 1s22s22p63s2
Fluorine:
1s22s22p5
Nitrogen:
1s22s22p3
Chapter 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)
Possible Short Answer Questions:
1. Why was Mendeleev’s table an accepted success?
2. Why is the periodic table called the “periodic” table?
3. What causes elements in the same column to have similar
chemical properties?
4. What is an ion and how are ions formed?
5. Why is the 2nd ionization energy of Na so much larger than
the 1st ionization energy?