1. Trends on the Periodic Table
The Nuclear Charge
The Shielding Effect
Trends on The Periodic Table
Atomic Radius
Ionic Radius
Ionization Energy
Electronegativity

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The Nuclear Charge
The Nuclear Charge is defined as the Positive charge in the nucleus
Increases Significantly across a period
Not as much going down a group…
WHY???

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The Shielding Effect
The Shielding Effect is the dampening affect or reduction of the interactions between the positive charge in the nucleus (the Nuclear Charge) and the outer shell electrons due to the inner shell electrons.
Increases with addition of Principle Energy Levels
Does not increase across a period…
WHY??

4. Atomic Size Atomic Radius- The radius of the atom,
distance from the Nucleus to the outer electron,
half the distance between atoms of the same element when in diatomic molecule
The distances between atoms in a molecule are extremely small.
The atomic radius is often measured in picometers (pm).
Recall that there are one trillion, or 1012, picometers in a meter.

5. Distance between the nuclei
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Trends in Atomic Size
Because the atomic radius is one-half the distance between the nuclei, a value of 140 pm (280/2) is assigned to the radius of the iodine atom.
Distance between the nuclei
280 pm
140 pm
Atomic radius

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Trends in Atomic Size
In general, atomic size increases from top to bottom within a group, and decreases from left to right across a period.

7. Group Trends in Atomic Size
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Interpret Graphs
Group Trends in Atomic Size
The atomic radius within these groups increases as the atomic number increases.
Atomic radius (pm)
Atomic number
This increase is an example of a trend.

8. Trends in Atomic Size Periodic Trends in Atomic Size
Across a period, the electrons are added to the same principal energy level.
The shielding effect is constant for all elements in a period.
The increasing nuclear charge pulls the electrons in the highest occupied energy level closer to the nucleus, and the atomic size decreases.
Atomic radius (pm)
Atomic number

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Trends in Atomic Size
The figure below summarizes the group and period trends in atomic size.

10. Ions What is an Ion? How do ions form?
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Ions
What is an Ion?
How do ions form?

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Ions
An ion is an atom or group of atoms that has a positive or negative charge.
The charge comes from the addition or subtraction of ELECTRONS
Positively Charged Ions are called Cations
Negatively charged ions are called Anions

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Ions
An atom is electronically neutral because it has equal numbers of protons and electrons.
For example, an atom of sodium (Na) has 11 positively charged protons and 11 negatively charged electrons.
The net charge on a sodium atom is zero [(+11) + (-11) = 0].
An Ion is NOT electrically Neutral because it has either more or less electrons than Protons
The Cl- anion (Atomic Number 17) has 18e- and 17 protons and a charge of -1
The Ca+2 cation has 20 protons and 18e- and there for has a charge of +2

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Ions
Atoms of metals, such as sodium, tend to form ions by losing one or more electrons from their highest occupied energy levels.

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Ions
Atoms of metals, such as sodium, tend to form cations by losing one or more electrons from their highest occupied energy levels.
In the sodium+ cation, the number of electrons (10) is not equal to the number of protons (11).
Because there are more positively charged protons than negatively charged electrons, the sodium ion has a net positive charge.

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Ions
Atoms of nonmetals, such as chlorine, tend to form anions by gaining one or more electrons.
In a chloride ion, the number of electrons (18) is not equal to the number of protons (17).
Because there are more negatively charged electrons than positively charged protons, the chloride ion has a net negative charge.

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What type of element tends to form anions? What type tends to form cations?

17. Nonmetals tend to form anions. Metals tend to form cations.
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What type of element tends to form anions? What type tends to form cations?
Nonmetals tend to form anions. Metals tend to form cations.

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Ionization Energy
Electrons can move to higher energy levels when atoms absorb energy. Sometimes the electron has enough energy to overcome the attraction of the protons in the nucleus.
The energy required to remove an electron from an atom is called ionization energy.
This energy is measured when an element is in its gaseous state.
The energy required to remove the first electron from an atom is called the first ionization energy.

19. Trends in Ionization Energy
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Trends in Ionization Energy
First ionization energy tends to decrease from top to bottom within a group and increase from left to right across a period.

20. Trends in Ionization Energy
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Trends in Ionization Energy
Ionization energies can help you predict what ions an element will form.
It is relatively easy to remove one electron from a Group 1A metal atom, but it is difficult to remove a second electron.
This difference indicates that Group 1A metals tend to form ions with a 1+ charge.

21. 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
Ar (noble gas
1520
2665
3947 K
419
3096
4600 Ca
590
1146
4941

22. Group Trends in Ionization Energy
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Interpret Graphs
Group Trends in Ionization Energy
Look at the data for noble gases and alkali metals.
First ionization energy (kJ/mol)
Atomic number
In general, first ionization energy generally decreases from top to bottom within a group.

23. Trends in Ionization Energy
Group Trends in Ionization Energy
The atomic size increases as the atomic number increases within a group.
As the size of the atom increases, nuclear charge has a smaller effect on the electrons in the highest occupied energy level.
Less energy is required to remove an electron from this energy level, and the first ionization energy is lower.

24. Trends in Ionization Energy
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Trends in Ionization Energy
Periodic Trends in Ionization Energy
In general, the first ionization energy of representative elements increases from left to right across a period.
This trend can be explained by the nuclear charge and the shielding effect.
The nuclear charge increases across the period, but the shielding effect remains constant.
As a result, there is an increase in the attraction of the nucleus for an electron.
Thus, it takes more energy to remove an electron from an atom.

25. Trends in Ionization Energy
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Trends in Ionization Energy
The figure below summarizes the group and period trends for the first ionization energy.
Energy generally increases
Energy generally decreases

26. Trends in Ionic Size
During reactions between metals and nonmetals, metal atoms tend to lose electrons and nonmetal atoms tend to gain electrons.
This transfer of electrons has a predictable effect on the size of the ions that form.
Cations are always smaller than the atoms from which they form.
Anions are always larger than the atoms from which they form.

27. Trends in Ionic Size Group Trends in Ionic Size
For each of these elements, the ion is much smaller than the atom.
The radius of a sodium ion (95 pm) is about half the radius of a sodium atom (191 pm).
When a sodium atom loses an electron, the attraction between the remaining electrons and the nucleus is increased. As a result, the electrons are drawn closer to the nucleus.

28. Trends in Ionic Size Group Trends in Ionic Size
The trend is the opposite for nonmetals, like the halogens in Group 7A.
For each of these elements, the ion is much larger than the atom.
For example, the radius of a fluoride ion (133 pm) is more than twice the radius of a fluorine atom (62 pm).
As the number of electrons increases, the attraction of the nucleus for any one electron decreases.

29. Trends in Ionic Size Group Trends in Ionic Size
Metals that are representative elements (NOT D BLOCK) tend to lose all their outermost electrons during ionization.
Therefore, the ion has one fewer occupied energy level.

30. Trends in Ionic Size Period Trends in Ionic Size
From left to right across a period, two trends are visible—a gradual
decrease in the
size of the positive
ions (cations),
followed by a
gradual decrease
in the size of the
negative ions
(anions).

31. Trends in Ionic Size
The figure below summarizes the group and period trends in ionic size.
Size of cations decreases
Size of anions decreases
Size generally increases

32. What are the trends for ionic size?
Ionic size tends to increase from top to bottom within a group. Generally, the size of cations and anions decreases from left to right across a period.

33. Trends in Electronegativity
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Trends in Electronegativity
What is Electronegativity?
Electronegativity is the ability of an atom to attract electrons or electron density within a molecule.
Phenol and the phenolate ion

34. Trends in Electronegativity
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Trends in Electronegativity
There is a property that can be used to predict the type of bond that will form during a reaction.
This property is called electronegativity.
Electronegativity is the ability of an atom of an element to attract electrons when the atom is in a compound.
Scientists use factors such as ionization energy to calculate values for electronegativity.

35. Electronegativity Values for Selected Elements
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Interpret Data
This table lists electronegativity values for representative elements in Groups 1A through 7A.
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 Si
1.8 P S Cl K
0.8 Ca Ga
1.6 Ge As Se
2.4 Br
2.8 Rb Sr In
1.7 Sn Sb
1.9 Te I Cs
0.7 Ba Tl Pb Bi
The data in this table is expressed in Pauling units.

36. Trends in Electronegativity
In general, electronegativity values decrease from top to bottom within a group. For representative elements, the values tend to increase from left to right across a period.
Metals at the far left of the periodic table have low values.
By contrast, nonmetals at the far right (excluding noble gases) have high values.
Values among transition metals are not as regular.

37. Trends in Electronegativity
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Trends in Electronegativity
The least electronegative element in the table is cesium, with an electronegativity of 0.7.

38. The Periodic Table and Electronegativity
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The Periodic Table and Electronegativity
The data in this table is expressed in Pauling units.

39. Trends in Electronegativity
This figure summarizes several trends that exist among the elements.

40. What are the trends for electronegativity values?
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What are the trends for electronegativity values?

41. What are the trends for electronegativity values?
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What are the trends for electronegativity values?
In general, electronegativity values decrease from top to bottom within a group. For representative elements, the values tend to increase from left to right across a period.

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Key Concepts
In general, ATOMIC SIZE increases from top to bottom within a group and decreases from left to right across a period.
Positive and negative ions form when electrons are transferred between atoms.
First ionization energy tends to decrease from top to bottom within a group and decrease from left to right across a period.
Ionic size tends to increase from top to bottom within a group. Generally, the size of cations and anions decreases from left to right across a period.
In general, electronegativity values decrease from top to bottom within a group. For representative elements, the values tend to increase from left to right across a period.

43. Electrons and the Structure of Atoms
These trends can be explained by variations in atomic structure.
The increase in nuclear charge within groups and across periods explains many trends.
Within groups, an increase in electron shielding has a significant effect on these trends.

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Glossary Terms
atomic radius: one-half the distance between the nuclei of two atoms of the same element when the atoms are joined
ion: an atom or group of atoms that has a positive or negative charge
cation: any atom or group of atoms with a positive charge
anion: any atom or group of atoms with a negative charge

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Glossary Terms
ionization energy: the energy required to remove an electron from an atom in its gaseous state
electronegativity: the ability of an atom to attract electrons when the atom is in a compound