Slide 1: Draw atomic models of lithium, magnesium, and fluorine.
From the models, predict whether the ions of these elements will be larger or smaller than the atoms. Be sure to justify your predictions. Slide 2: Objectives Describe periodic trends in ionization energy, and relate them to the atomic structures of the elements.
Describe periodic trends in atomic radium, and relate them to the atomic structures of the elements.
Describe periodic trends in electronegativity, and relate them to the atomic structures of the elements.
Describe periodic trends in ionic size, electron affinity, and melting and boiling points, relate themto the atomic structures of the elements. Slide 3: Periodic Trends The arrangement of the periodic table reveals trends in the properties of the elements.
A trend is a predictable change in a particular direction.
Understanding a trend among the elements enables you to make predictions about the chemical behavior of the elements.
These trends in properties of the elements in a group or period can be explained in terms of electron configurations. Slide 4: Ionization Energy The ionization energy is the energy required to remove an electron from an atom or ion. Slide 5: Ionization Energy Slide 6: Ionization Energy Decreases as You Move Down a Group Each element has more occupied energy levels than the one above it has.
The outermost electrons are farthest from the nucleus in elements near the bottom of a group.
As you move down a group, each successive element contains more electrons in the energy levels between the nucleus and the outermost electrons.
Electron shielding is the reduction of the attractive force between a positively charged nucleus and its outermost electrons due to the cancellation of some of the positive charge by the negative charges of the inner electrons. Slide 7: Ionization Energy Decreases as You Move Down a Group Slide 8: Ionization Energy Increases as You Move Across a Period Ionization energy tends to increase as you move from left to right across a period.
From one element to the next in a period, the number of protons and the number of electrons increase by one each.
The additional proton increases the nuclear charge.
A higher nuclear charge more strongly attracts the outer electrons in the same energy level, but the electron-shielding effect from inner-level electrons remains the same. Slide 9: Ionization Energy Increases as You Move Across a Period Ionization : Ionization Slide 11: Atomic Radius The exact size of an atom is hard to determine.
The volume the electrons occupy is thought of as an electron cloud, with no clear-cut edge.
In addition, the physical and chemical state of an atom can change the size of an electron cloud.
One method for calculating the size of an atom involves calculating the bond radius, which is half the distance from center to center of two like atoms that are bonded together.
The bond radius can change slightly depending onwhat atoms are involved. Slide 12: Atomic Radius Bond Length : Bond Length Slide 14: Atomic Radius Increases as You Move Down a Group As you proceed from one element down to the next in a group, another principal energy level is filled.
The addition of another level of electrons increases the size, or atomic radius, of an atom.
Because of electron shielding, the effective nuclear charge acting on the outer electrons is almost constant as you move down a group, regardless of the energy level in which the outer electrons are located. Slide 15: Atomic Radius Decreases as You Move Across a Period As you move from left to right across a period, each atom has one more proton and one more electron than the atom before it has.
All additional electrons go into the same principal energy level—no electrons are being added to the inner levels.
Electron shielding does not play a role as you move across a period.
As the nuclear charge increases across a period, theeffective nuclear charge acting on the outer electronsalso increases. Check out the Atomic Radius Interactive Tutorial
http://college.cengage.com/chemistry/intro/zumdahl/intro_chemistry/5e/students/protected/periodictables/pt/pt/pt_ar5.html Slide 16: Atomic Radius Slide 17: Atomic Radius Slide 18: Periodic Trends of Radii Slide 19: Periodic Trends of Radii Atomic Radius : Atomic Radius Slide 21: Electronegativity Not all atoms in a compound share electrons equally.
Knowing how strongly each atom attracts bonding electrons can help explain the physical and chemical properties of the compound.
Linus Pauling, an American chemists, made a scale of numerical values that reflect how much an atom in a molecule attracts electrons, called electronegativity values.
Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons. Slide 22: Electronegativity The atom with the higher electronegativity will pull on the electrons more strongly than the other atom will.
Fluorine is the element whose atoms most strongly attract shared electrons in a compound. Pauling arbitrarily gave fluorine an electronegativity value of 4.0.
Values for the other elements were calculated in relation to this value. Slide 23: Electronegativity Decreases as You Move Down a Group Electronegativity values generally decrease as you move down a group.
The more protons an atom has, the more strongly it should attract an electron.
However, electron shielding plays a role again. Slide 24: Electronegativity Increases as You Move Across a Period Electronegativity usually increases as you move left to right across a period.
As you proceed across a period, each atom has one more proton and one more electron—in the same principal energy level—than the atom before it has.
Electron shielding does not change as you move across a period because no electrons are beingadded to the inner levels. Slide 25: Electronegativity Increases as You Move Across a Period, continued The effective nuclear charge increases across a period.
As this increases, electrons are attracted much more strongly, resulting in an increase in electronegativity.
The increase in electronegativity across a period is much more dramatic than the decrease in electronegativity down a group. Slide 26: , Slide 27: Electronegativity Slide 28: Other Periodic Trends The effective nuclear charge and electron shielding are often used in explaining the reasons for periodic trends.
Effective nuclear charge and electron shielding also account for electron affinity. Slide 29: Periodic Trends in Electron Affinity The energy change that occurs when a neutral atom gains an electron is called the atom’s electron affinity.
This property of an atom is different from electronegativity.
The electron affinity tends to decrease as you move down a group because of the increasing effect of electron shielding.
Electron affinity tends to increase as you move across a period because of the increasing nuclear charge. Slide 30: Periodic Trends Electron Affinity Electron Affinity : Visual Concepts Electron Affinity Chapter 4