Download 4.3 Ionization Energy

The Periodic Table
Chapter 4
4.1 Mendeleev’s Periodic
Table (1869)
Reading Check
 Why did Mendelev
place certain
elements together in
the same group?
 Arranged elements
with similar
properties into
groups.
4.1 Mendeleev’s Periodic
Table (1869)
 Why did Mendeleev
leave gaps in his
periodic table?
 ANS = He predicted
that elements would
later be discovered
to occupy these gaps
4.1 Periodic Table
 Why does the PT have the shape it does?
4.1 Periodic Table
 Elements are organized in order of
increasing atomic number
 Atomic mass follows the same order (in
most cases)
4.1 Periodic Table
 Mendeleev did not know why certain
elements have similar properties
because he did not know about atomic
theory.
 But we do now! Why do properties
repeat themselves?
 The properties of an element are
determined by it’s valence electrons.
3.3 Valence Electron
Anelectronthatisfoundontheoutermostshello
fanatomandthatdeterminestheatom’schemi
calproperties.
3.3 Valence Electron
An electron that is found on the outermost
shell of an atom and that determines the
atom’s chemical properties.
The electrons that are on the inside are
called core electrons
Which element is this? 
Think about it…..
How many valence electrons are in there in
a sulfur atom? (first, write the electron
configuration)
Think about it…..
1s2 2s2 2p6 3s2 3p4
2-8-6
How many valence electrons are in there in
a sulfur atom?
6 valence electrons
The rest are core electrons.
How many core electrons?
10
4.1 Valence electrons
 How many
valence electrons
are there in a
fluorine atom?
 (first, write the
electron
configuration for
fluorine)
4.1 Valence electrons
 How many
valence electrons
are there in a
fluorine atom?
 1s22s22p5
 [He]2s22p5
 2-7
4.1 Periodic Table
 A period (group/period) is a horizontal
row in the periodic table 
 numbered 1-7
 Corresponds to the energy level of the
valence electrons
4.1 Periodic Table
A group is a vertical column in the periodic
table ↓
Groups 1-18
See fig 3 p 119
s-block, p-block, d-block, f-block
4.3 Trends in the periodic
table
Standard
Use the periodic table as a model to predict
the relative properties of elements based
on the patterns of electrons in the outer
energy level of atoms.
4.3 Trends in the periodic
table
Learning Target
Students learn about periodic trends in
atomic radius by plotting/interpreting a
graph
Think about it….
Atoms are small……but are they all the
same size?
Is a carbon atom the same size as an iron
atom?
4.3 Trends in the PT
Worksheet: Family Characteristics
Part 1: Atomic Radius
How is atomic radius defined?
See fig 19 p. 135
4.3 Atomic Radius
fig 21 p. 136
Inference: What is this diagram telling us?
4.3 Trend in Atomic
Radius
Draw in the lines for groups 1,2,13,14,
17,18 on your worksheet.
Inference: What is your graph telling you?
What is the trend in atomic radius as you
read down a group in the PT?
What is the trend in atomic radius as you
read across a period in the PT?
4.3 Trends in the PT
 How does atomic radius change as you
read down a group in the PT?
 ANS = Atomic radius increases
(increases/decreases).
4.3 Atomic radius
Notes: Explain why (1 sentence)
4.3 Trends in the PT
 How does atomic radius change as you
read across a period in the PT?
 ANS = Atomic radius decreases
(increases/decreases).
4.3 Trends in the PT
COUNTERINTUITIVE 
Factors:
• Effective nuclear charge
• Shielding
Read p. 136 ‘Atomic Radius Decreases as
You Move Across a Group’
4.3 Electron Shielding
Calculate effective nuclear charge for Na and Ar
Factors:
• Effective nuclear charge
• Shielding
4.3 Atomic Radius
Summary
fig 20 p. 135
Inference: where are the largest elements
found on the PT? The smallest?
4.3 Atomic Radius
Summary
fig 21 p. 136
Inference: what is this diagram telling us?
How does this diagram differ from the one
we plotted?
Check For Understanding:
Atomic Radius
1. Which of the following elements has the
smallest atomic radius?
a. Magnesium
b. Calcium
c. Strontium
d. Barium
2. Which of these elements has the largest atomic
radius?
a. Gallium
b. Germanium
c. Selenium
d. Bromine
4.3 Trends in the periodic
table
Standard
Use the periodic table as a model to predict
the relative properties of elements based
on the patterns of electrons in the outer
energy level of atoms.
Learning Target
Students learn about periodic trends in
ionization energy by plotting/interpreting a
graph
4.3 Ions
Notes:
Whenatomshaveequalnumbersof
protonsandelectronstheyareelect
ricallyneutral
4.3 Ions
When atoms have equal
numbers of protons and
electrons they are electrically
neutral.
Butwhenenoughenergyisadded,t
heattractiveforcecanbeovercome.
4.3 Ions
But when enough energy is
added, the attractive force can
be overcome.
Whenthishappens,anelectronisre
movedfromtheatom,anditbecome
sapositivelychargedion
4.3 Ions
When this happens, an electron
is removed from the atom, and it
becomes a positively charged
ion. (p. 133)
Positively charged ion = ‘cation’
IONS
4.3 Ionization Energy
Ionization Energy = The energy that is needed
to remove an electron from an atom.
Na atom  Na+ ion + electron
Ionization Energy
Unit = (kJ/mole)
‘First Ionization energy’ =
energy needed to remove the first electron from an atom
4.3 Ionization energy
If ionization energy is high, it is difficult
(difficult/easy) to remove an electron from
the atom.
4.3 Ionization energy
If ionization energy is low, it is easy
(difficult/easy) to remove an electron from
the atom.
4.3 Trend in ionization
energy
Read Family Characteristics worksheet
Part 2
Draw in the lines for groups 1,13,15,17,18
on your worksheet.
Inference: What is your graph telling you?
4.3 Ionization Energy
 Ionization Energy decreases (increases
or decreases) as you read down a group
in the PT.
 Ionization Energy increases
(increases/decreases) as you read
across a period in the PT
Think about it…..
 In the Brainiac alkali metals video clip,
which group 1 element was the most
reactive when put into water?
ANS =
Those at the bottom
Rb & Cs
4.3 Ionization Energy
‘Ionization Energy decreases as you read
down a group in the PT.’
 Which elements in group 1 have the
lowest ionization energy?
ANS=
Those at the bottom
Rb & Cs
4.3 Electron Shielding
Notes: Explain why (1 sentence)
Read p. 133 ‘Ionization
Energy Decreases As
You Move Down a
Group’
4.3 Effective Nuclear
Charge
Why is the ionization energy for chlorine greater
than that for magnesium? Read Ionization
Energy Increases as you Move Across a Period
p. 134
4.3 Ionization Energy
See fig 17 p. 134
Inference: which elements have the highest
ionization energy? The lowest?
Periodic Trends
What is the relationship between atomic
radius and ionization energy?
An atom with a small atomic radius has a
high (high/low) ionization energy.
e.g. Helium
Periodic Trends
What is the relationship between atomic
radius and ionization energy?
An atom with a large atomic radius has a
low (high/low) ionization energy.
e.g. Cesium
Check For Understanding:
Ionization Energy
1. Which of the following elements has the lowest
ionization energy?
a. Magnesium
b. Calcium
c. Strontium
d. Barium
2. Which of these elements has the highest
ionization energy?
a. Gallium
b. Germanium
c. Selenium
d. Bromine
Electronegativity
Definition:
Ameasureoftheabilityofanatominachemical
compoundtoattractelectronstoitself.
A measure of the ability of an atom in a
chemical compound to attract electrons
to itself. (p. 137)
Facts:
Read/Notes ‘Electronegativity’ p. 137 (first
3 paragraphs)
think about it……..
Periodic Trends
What is this diagram telling us?
How does ‘electronegativity’ vary as we move around the PT?
Scale: 0-4
Electronegativity
See p. 137 fig 22
Where are the most electronegative elements
found in the PT?
Where are the least electronegative elements
found in the PT?
5.1 Electronegativity
 Fluorine has the strongest demand for
electrons
 Cesium has the lowest demand for electrons
Periodic Trends
What is the relationship between ionization
energy and electronegativity?
An atom with a high ionization energy has
a high (high/low) electronegativity
e.g. Fluorine
Directly proportional
5.1 Electronegativity
 Electronegativity difference between two atoms
tells us what kind of bond will form between
them.
5.1 Electronegativity
 Calculate the EN difference between
sodium and chlorine
= Ionic Bond
5.1 Electronegativity
 Electronegativity difference between two atoms
tells us what kind of bond will form between
them.
5.1 Electronegativity
 Calculate the EN difference between
sulfur and oxygen
= covalent bond
Other Periodic Trends
p. 139
 Electron Affinity
 Melting and Boiling Points