3.03 Periodic Trends Chemistry

Ion

An atom or group of atoms that has a positive or negative charge due to a loss or gain of electrons.

ionization energy: the energy required to remove 1 electron from an element, resulting in a positive ion.

-An electron can be removed from an atom if enough energy is provided. When one or more electrons are removed, the atom is left with a positive charge becuase it has more protons than electrons. -Na+ is the symbol for an ion of sodium with a single positive charge because it now has one more proton than electron. -The measurement of ionization energy is made on atoms in the gaseous state to reduce the influence of nearby atoms. Elements with a lower effective nuclear charge felt by their electrons will give up an electron easier than other elements, so ionization energy follows a pattern, or trend, on the periodic table. -Successive Ionization Energies: if you have enough energy, electrons can be removed from positive ions just like they can be removed from neutral atoms. The energy required to remove a second electron from an atom is called the "second ionization energy," to remove a third electron would be called the "third ionization energy," and so on. -Across a period, ionization energy has a general increase from left to right. The stronger the effective nuclear charge felt by an atom's valence electrons, the more energy is required to remove one of these electrons from the atom. As effective nuclear charge increases across a period, ionization energy also increases. -Down a group, the ionization energy of elements decreases going down a group becuase the atomic radius of the atoms increases. The farther an atom's valence electrons are from its nucleus, the lower the amount of energy required to remove one of those electrons.

effective nuclear charge: the charge (from the nucleus) felt by the valence electrons after you have taken into account the number of shielding electrons that surround the nucleus.

-In order to compare many of the periodic trends on the periodic table, it is important to understand a measure that chemists use to compare the attraction of valence electrons to the nuclei of different elements. -Subtracting the number of shielding (core) electrons from the total nuclear charge (number of protons) gives the effective nuclear charge felt by an atom's valence electrons. The values of the effective nuclear charge can be calculated or measured more accurately, but this generalization will be sufficient to help us understand the rest of the periodic trends. -across a period, increase from left to right. This is because each element has one additional proton in its nucleus, adding to the overall nuclear charge. Each atom also has an additional electron in the valence energy level, but the number of core electrons in the lower energy levels remains the same across the entire period. This means that the nuclear charge increases, but the shielding stays constant, so the effective nuclear charge increases by a factor of one for each element across a period. Down a group- stays constant for each element. This is because the number of additional protons in each element's nucleus is equal to the number of additional shielding electrons in the inner energy levels of the atom.

electronegativity: a measure of the attraction of an atom for the electrons in a chemical bond.

-Valence electrons, the electrons that get involved in chemical bonding, are responsible for holding atoms together in molecules. The higher an element's electronegativity, the greater its attraction for all of the electrons in a chemical bonds. Electronegativity is related to effective nuclear charge. Electrons that have a high effective nuclear charge also have high electronegativity values due to the weak pull exerted on the electrons by the nucleus. -scientists determined an arbitrary rating for electronegativity, a scale from 0.0 to 4.0. Most of the noble gases, from helium to argon, were not given electronegativity values because they do not form chemical bonds. This means that Fluorine (F) is the most electronegative element on the entire periodic table. Fluorine was given the highest electronegativity value, of 4.0, and all other elements were given values by comparing their electronegativity to that of Fluorine. The lowest electronegativity values on the periodic table are 0.7 for both cesium (Cs) and francium (Fr). -Across a period, electronegativity increases from left to right becuase of the increase in effective nuclear charge. The greater an atom's attraction for its own valence electrons, the greater it is able to attract another atom's electrons in a chemical bond. -Down a group, electronegativity decreases becuase of increased distance between the valence electrons and the nucleus (greater atomic radius).

ionic radii: 1/2 the diameter of an ion

-We use ionization to examine the energy required to remove electrons and form positive ions, but in nature, not all ions have a positive charge. A positive ion is called a cation, and a negative ion is called an anion. -Metals naturally form cations but losing one or more electrons. Most metals lose enough electrons to empty out their valence, giving the ion an electron configuration of the noble gas from the previous period. When the outermost energy level of electrons is emptied, the remaining electron cloud is smaller. This is why all cations are smaller than they were as neutral atoms. -Nonmetals naturally form anions by gaining one or more electrons. Most nonmetals gain enough electrons, if they can, to fill their valence energy level. This would give the ion an electron configuration like the noble gas at the end of their row on the periodic table. When electrons are added to an atom without a change in the number of protons, the electrons experience additional repulsion among them. This causes the electron cloud to spread out, becuase each electron experiences a lower effective nuclear charge. This is why all anions are larger than they were as neutral atoms. -Across a period, the metals at the left form cations and the nonmetals at the right form anions. There is a decrease in the ionic radii of the cations from left to right across the metals, and also in the ionic radii of the anions from left to right across the nonmetals. However, all of the anions (nonmetal negative ions) are larger in radius than all of the cations (metal positive ions) in a period. -Down a group, ionic radii increase, following the same trend as atomic radii, because there is still an increase in the number of occupied energy levels in the electron cloud.

atomic radius: 1/2 the diameter of an atom, measured by taking 1/2 of the distance between the centers of 2 atoms of that element that are bonded together.

-You have already seen that the electron cloud makes up the majority of the size of an atom. Because an atom's electron cloud does not have a definite edge that can be easily detected, it is difficult to measure the size of an individual atom. -Across a period, there is a gradual decrease in atomic radii from left to right. The decrease in radius is due to the additional number of protons in each nucleus. -Down a group, there is a general increase in atomic radii going down each group of elements on the periodic table. Going down in the group, there is an increase in the number of occupied energy levels in the electron cloud, as well as an increase in protons.

The periodic table is arranged based on the properties of the elements

Reading across each period and down each group, you will see repeated trends in some of these properties as atomic number increases.