Download The Development of the Periodic Table - wths

The Development of the Periodic
Table
Chapter 7 Section 1
Timeline of Development…

1790’s



Antoine Lavoisier: compiled a list of
elements (about 23)
Mid-1800’s

Scientists developed a way to
determine atomic mass
1870

Ah Ha!
My life has
purpose again
About 70 known elements
Organization

Meyer, Mendeleev & Moseley
Mendeleev gets most of the credit
 Organized by atomic mass (just as
Newlands) but changed columns
 Organized into columns with similar
properties
 Left blank spaces for places where he
thought elements should be, but weren’t
discovered yet
 Table 7.1?

Mendeleev’s Predictions
Why not atomic #?
It was found that some of Mendeleev’s
elements were incorrectly placed
 Why didn’t he use atomic number
instead of atomic mass?
 Answer: atomic #’s weren’t discovered
until the early 1900’s

Moseley’s Adaptation
After Henry Moseley discovered protons
(and atomic number) he changed the
organization and fixed Mendeleev’s
problems
 Periodic Law:


Periodic repetition of chemical and physical
properties of the elements when arranged
by increasing atomic number
Parts of the Periodic Table
Columns = Groups (or families)
Rows = Periods
Sections of the PT
Transition Elements
Inner Transition Elements
Other periodic tables…
Why?

Why do things behave the way they do?

The best predictor/explanation of why
elements react are found in:
 Their
# of electrons
 The way their electrons are organized
 The size of the atoms
 How much they want electrons or how much they
want to get rid of electrons
Valence Electrons
Electrons in the outermost energy level
of an atom
 Core Electrons: all electrons that are
not in the valence shell

Na
1s22s22p63s1
Electron Shielding
Positives & Negatives are attracted to
each other
 Effective Nuclear Charge: describes
the pull on the electrons by the nucleus
Zeff = Z - S

Effective Nuclear Charge
Shielding constant
(# of non-valence electrons)
Nuclear Charge
(# of protons)
Atomic Size
50ml + 50ml = ?
 Atomic Size

Atoms of different elements have different
sizes
 What happens to Zeff
as we go down a group?
As we go across a period?

Atomic Radius

What is it?
Atomic Radius Trend
Increases
Increases
Increases
Atomic Radius Trend
Increases
Why?
 1) As you go down a group, principle
energy levels are added



(n=1, n=2, n=3)
This increases the radius
Increases
Atomic Radius Trend
Why?
 2) As you go across a period:

No energy levels are added
 Protons are added

Increases
Ionic Radius

Ions:
An atom that has an overall positive or
negative charge
 Examples:

 Cl-1 (Chlorine
with 17 protons and 18 electrons)
 Mg2+ (Magnesium with 12 protons and 10
electrons)

What happens to size when atoms do this?
Ionic Radius Trend

Positive Atoms
To become positive, atoms lose electrons
 What happens to size if you lose
electrons?

 Hint:
You now have more positives pulling in
less negatives
Positive
Nucleus
Ionic Radius Trend

Negative Atoms
To become negative, atoms gain electrons
 What happens to size if you gain
electrons?

 Hint:
out
You now have more negatives pulling
Positive
Nucleus
Chapter 7 Test



Monday – January 7th
Development of the Periodic Table
Periodic Trends (what & why)







Atomic radius
Ionic radius
Ionization energy
Electron Affinity
Isoelectronic
Ions
Groups of the Periodic
Comparing Atomic Size

Remember isoelectronic


When atoms have the same electron
configuration, which one is bigger?
Example:
a) Na+
b) F-1 c) O-2
Radius decreases with increasing nuclear
charge (# of protons)
O-2 > F-1 > Na+
Na = +11
F = +9
O = +8
Sample 7.6
Ionization Energy
The energy required to remove an
electron from an atom
 1st IE: Energy to remove the first
electron
Na  Na+ + e 2nd IE: Energy to remove the 2nd
Na+  Na2+ + e 3rd IE, 4th IE etc…

Hard to
steal
electrons
Easy to
steal
electrons
Increases
Increases
Trend in 1st Ionization Energy
IE Equations & Energies
We show the change through an
equation:
Na  Na+ + e- E=+495
Na+  Na+2 + e- E= +4562
 Why is the 2nd IE so much bigger?

Spikes in IE
Sample 7.7
Electron Affinity
The measure of how much an atom
wants to gain an electron
 For most atoms, energy is released
when this happens


Delta E = negative
Affinity vs Ionization

Ionization energy


Cl  Cl+ + e-
DE = 1251 kJ/mol
Electron Affinity

Cl + e-  Cl-
DE = -349 kJ/mol
More negative = more
energy given off = more
favorable
Fluorine has the
most electron
affinity
Increases
Increases
Electron Affinity