Chemistry: The Periodic Table Flashcards
Periodic Law
The chemical properties of the elements are dependent, in a systematic way, upon their atomic numbers.
Periods (Rows)
There are seven, representing the principal quantum numbers 1 to 7. Each period is filled subsequently.
Groups (Columns)
Represent elements that have the same electronic configuration in their valence, and share similar chemical properties.
Representative (A) Elements
One set of two groups. Have either s or p sublevels as their outermost orbitals. These elements are those in Groups IA through VIIA, all of which have incompletely filled s or p subshells of the highest principal quantum number.
Note: With the exception of noble gases (Group VIII), elements all have a completely filled p subshell.
Nonrepresentative (B) Elements
Includes transition elements, which have partly filled d sublevels, and the lathanide and actinide series, which have partly filled f sublevels.
Atomic Radius
The atomic radius of an element is equal to 1/2 the distance between the centers of two of that element that are just touching each other. In general, the atomic radius decreases across a period from left to right and increases down a given group; the atoms with the largest radius will be located at bottom of groups, in Group IA.
As one moves from left to right across period, electrons are added one at a time to the valance shell. Electrons within a shell cannot shield one another from the attractive pull of protons. So, since the number of protons is also increasing, producing a greater positive charge attracting the valance electrons, the effective nuclear charge increases steadily across a period. This causes the atomic radius to decrease.
As one moves down a group of the periodic table, the number of electrons and filled electron shells will increase, but the number of valance electrons remain the same. So, the outermost electrons in a given group will feel the same amount of effective nuclear charge, but electrons will be found farther from the nucleus as the number of filled energy shells increases. Thus, the atomic radii will increase.
Ionic Radius
Radius of a cation or an anion. It will affect the physical and chemical properties of an ionic compound. For example, the 3D infrastructure of a given ionic compound will depend largely upon the sizes of the cations and anions. If a neutral atom is converted into an ion, one would expect the radius to increase secondary to the repulsion of the additional electrons that will increase the size of the electron cloud.
Ionization Energy
Or ionization potential. it’s the energy required to remove an electron completely from a gaseous atom or ion. Removing an electron from an atom always requires an input of energy. The closer and more tightly bound an electron is to the nucleus, the more difficult it will be to remove and the higher the ionization energy will be.
The first ionization energy is the energy required to remove one valance electron from the parent atom. The second ionization energy is the energy needed to remove a second valance electron from the univalent ion to form a divalent ion, and so on. Successive ionization energies grow increasingly large.
Ionization energy increases from left to right across a period as the atomic radius decreases. Moving down a group, the ionization energy decreases as the atomic radius increases. Group IA elements have low ionization energies because the loss of an electron results in the a stable octet.
Electron Affinity
The energy change that occur when an electron is added to a gaseous atom, and it represents the ease with which the atom can accept an electron. The stronger the attractive pull of the nucleus for electrons (effective nuclear charge of Zeff), the greater the electron affinity will be. In discussing electron affinities, two sign conventions are used. The more common one states that a positive electron affinity value represents energy release when an electron is added to an atom; the other states that a negative electron affinity represents a release of energy. In this discussion, the first convention is used.
Electron affinity can be best represented in this equation, where X is an atom of a given element in the gaseous state (g):
X(g) + e- –> X-(g)
Generalizations can be made about the electron affinities of particular groups in the periodic table. Group IIA elements (alkaline earths) have low electron affinity values because their s subshell is filled. Group VIIA (halogens) elements have high electron affinities because the addition of an electron to the atom results in a completely filled shell, which represents a stable electron configuration.
Achieving the stable octet involves a release of energy, and the strong attraction of the nucleus for the electron leads to a high energy change. The group VIIIA elements (noble gases) have electron affinities on the order of zero because they already possess a stable octet and cannot readily accept an electron. Elements of other groups generally have low values of electron affinity.
Electronegativity
A measure of the attraction an atom has for electrons in a chemical bond. Greater the electronegativity, the greater its attraction for bonding electrons.
Electronegativity values not determined directly. Most common scale is the Pauling electronegativity scale where the values range from 0.7 for the most electropositive elements, like cesium, to 4.0 for the most electronegative element, flourine.
Electronegatives are related to ionization energies: Elements with low ionization energies will have low electronegatives because their nuclei don’t attract electrons strongly, while elements with high ionization energies will have high electronegativities because of the strong pull the nucleus has on electrons.
Therefore, electronegativity increases from left to right across periods. In any group, electronegativity decreases as the atomic number increases, as a result of the increased distance between the valance electrons and the nucleus.
Metals
Shiny solids (except for mercury) at room temperature and generally have high melting points and densities. Metals have the characteristic ability to be deformed without breaking. The ability of a metal to be hammered into shapes is called malleability, and the ability to be drawn into wires is called ductility. Many of the characteristic properties of metals, such as large atomic radius, low ionization energy, and low electronegativity, are due to the fact that the few electrons in the valence shell of a metal atom can easily be removed. Because the balance electrons can move freely, metals are good conductors of heat and electricity.
Groups IA and IIA represent the most reactive metals. Metals are located on the left side and in the middle of the periodic table.
Nonmetals
Generally brittle in the solid state and show little or no metallic luster. They have high ionization energies and electronegatives and are usually poor conductors or heat and electricity. Most nonmetals share the ability to gain electrons easily, but otherwise they display a wide range of chemical behaviors and reactivities.
The nonmetals are located on the upper right side of the periodic table; they are separated from the metals by a line cutting diagonally through the region of the periodic table containing elements with partially filled p orbitals.
Metalloids
Or semimetals. Found along the ling between the metals and nonmetals int he periodic table, and their properties vary considerably. Their densities, boiling points, and melting points fluctuate widely. The electronegativities and ionization energies of metalloids lie between those of metals and nonmetals; their, these elements possess characteristics of both those classes.
The reactivity of metalloids is dependent upon the element with which they are reacting.
Hydrogen
There’s no suitable place for hydrogen in the periodic table. Hydrogen does resemble alkali metals because it has a single s valence electron and forms the H+ ion, which is hydrated in solution. However, it can also form the hydride ion, which is far too reactive to exist in water. In this respect, hydrogen resembles the halogens.
Alkali Metals
Elements of Group IA. They possess most of the physical properties common to metals, yet their densities are lower than those of other metals. The alkali metals have only one loosely bound electron in their outermost shell, giving them the largest atomic radii of all the elements in their respective periods. Their metallic properties and high reactivity are determined by the fact that they have low ionization energies; they easily lose their valence electron to form univalent cations. Alkali metals have low electronegativities and react very readily with nonmetals, especially halogens.
