Fundamentals of Elements and the Periodic Table

Periodic Table - The periodic table is a compressed collection of specific information about the elements.

• Period - Each row within the periodic table is called a period.

• Group - Each column within the periodic table is called a group.

  • Elements within the same groups share chemical and physical similarities and the closer they are in the group, the more similar two elements are.

• As we go up and down various groups and periods, we can identify trends in the atomic properties.

• Periods are numbered going down and groups are numbered going right.

• Each element is represented in the periodic table by a box that contains its atomic symbol, number, mass and possibly other information, depending on how detailed a given copy of the periodic table is.

• Atomic Mass - The atomic mass is the average mass of the atoms of an element

  • It is measured in atomic mass units, amus for short.

  • Atomic mass is experimentally calculated by taking the average weight of various isotopes of the element, multiplying each value by the abundance in nature of that particular isotope, and adding together all the resulting values.

• Atomic Symbol - Atomic symbols are the abbreviations used in chemistry for chemical elements, that usually consist of one or two identifying letters, where the first letter is capitalized.

  • Most atomic symbols are derived from the English name for the element, but some are derived from the name of the element in other languages, with the most common being Latin.

  • Chemists will use subscripts and superscripts along with the atomic symbol to represent certain information about the element, as seen below.

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  • Most people will typically only note the bottom-right subscript, used to indicate how many atoms of the element are present in a formula unit of a compound.

• Atomic Number - The atomic number of an element is the charge number of its nucleus.

  • Within an ordinary, unbonded atom, the atomic number represents the number of protons and electrons the atom contains.

  • The number of electrons in an atom can change naturally due to various reactions and such but the number of protons cannot be changed as easily.

Protons - For any given element, the number of protons is equal to the atomic number.

Electrons - For any given element, the number of electrons is equal to the atomic number.

• For an ion of an element, the number of electrons will change because after a chemical reaction, the atom has either lost or gained electrons.

• A positive ion has lost electrons and a negative ion has gained electrons.

Neutrons - The neutron is another subatomic particle present within the nucleus of the atom.

• Neutrons have about the same mass as a proton but without an electric charge.

• The number of neutrons in an element depends on the specific isotope of the element in question.

Isotopes - Atoms with the same number of protons but different numbers of neutrons are called isotopes.

• Different isotopes have different atomic masses as a result of the differences in the number of neutrons.

Radioactive Decay - Radioactivity is a property of matter whereby an unstable nucleus spontaneously emits small particles and/or energy in order to attain a more stable nuclear state.

• Radioactive isotope - An isotope that contains an unstable nucleus is termed a radioactive isotope.

• One radioactive nucleus may decay into another radioactive nucleus and so on, until the radioactive decay results in an isotope with a stable nucleus.

• Radioactive isotopes can be found in nature but they can also be created in nuclear experiments.

• Both natural and artificial radioactive isotopes emit subatomic particles as they disintegrate and when these particles are emitted, the nuclear mass and/or charge changes and one isotope is converted into another.

• Radioactive decay can also result in energy, released in the form of x-rays or gamma rays.

Periodic Properties and Trends

The periodic table can be broken up into groups of related elements such as the alkali metals, alkaline metals, earth metals, transition elements, halogens, and noble gasses.

Each column or group has the same number and type of outermost electrons resulting in the chemical similarities of these  electrons.

• Chemical reactions occur when atoms collide with one another and since in these collisions, the outermost electrons make the first contact, elements with similar electronic structures have similar chemical properties.

DIfferentiating Electron - A differentiating electron is the electron in a neutral element that makes it different from the previous element

Atoms and molecules will undergo chemical reactions to gain or lose electrons such that they have a full valence shell and are isoelectronic with a noble gas.

Isoelectronic - Isoelectronic refers to atoms and ions that have identical electron configurations.

Of all the known elements on the periodic table, two elements, Mercury and Bromine, are solid at room temperature. The noble gasses are gaseous at room temperature and the remaining elements are solids.

Most elements in the periodic table may be considered as individual atoms, however, some will exist naturally as diatomic molecules for the sake of electronic stability.

Diatomic - A diatomic molecule contains two atoms of the same element.

Metals -

Non-Metals -

Metalloids - Elements bordering the line between metals and non-metals are often termed metalloids since they exhibit properties of both metals and nonmetals.

Allotropes - Allotropes are different physical forms with different sets of chemical and physical properties  in which an element can exist.

• For example graphite, charcoal, and diamond are all allotropes of carbon.

The melting and boiling points of metals tend to decrease from the top to the bottom of a group, bu non-metals show an increase in their melting and boiling points.

• Similar trends can be noted in each group for electrical properties, densities, and specific heats.

The atomic radius increases from the top to the bottom of a group because eac increase in period number involves another, larger, energy level. But when moving left to right across a period, there is a general decrease in atomic radius because of an increase in the effective nuclear charge.

Effective Nuclear Charge - core electrons shield valence electrons from n equal amount of positive electrons, but as you move from left to right across a period, the number of core electrons stays the same while nuclear charge increases, creating an overall increase in effective nuclear charge and thus a decrease in atomic radius.

Ionization Energy - the ionization energy is the energy needed to remove an electron from an atomic completely.

• Ionization energy is always endothermic  because energy must be added to remove an electron.

• First ionization energy for most elements decreases from the top to the bottom of a group.

• First ionization energy for most elements increases from the left to the tight end of a period.

• valence electrons have lower ionization energy when compared to the notation of energies of the same atom’s core electrons. This is because the core electrons are closer to the nucleus.

Photoelectron Spectroscopy - experimental technique that measures the relative energies of electrons in atoms and molecules

Electron affinity - energy change that accompanies the addition of an electron to an atom.

• Most atoms do not accept additional electrons steadily and the electron affinity is a positive valve, indicating that energy must be used to add the electron.

• Some atoms readily attract electrons and electron affinity has a negative value,  meaning that energy is released.

• Electron affinity varies diagonally across the table, with fluorine having the highest and francium having the lowest.

Electronegativity - Electronegativity is a combination of ionization energy, electron affinity and other factors, that in total is used to describe the attraction of electrons by individual atoms.

• Electronegativity shows the same diagonal trend as ionization energy and electron affinity.

Ionic Radii - the radius of an atom changes after it becomes an ion due to a change in the number of electrons.

• Cations - Cations are atoms that have lost one or more electrons and carry a positive charge.

• Anions - Anions are atoms that have gained one or more electrons and carry a negative charge.

• Cations are smaller than the neutral atom of the same element, while anions are always larger than the neutral atoms.

• With cations, it's easy to identify the cause of the change in size: by losing its valence electrons, an atom loses its entire outer shell, thus making it smaller.

• In the case of anions, the cause of the difference is harder to identify. The added electron(s) still go in the already present valence shell so why is the radius getting bigger? Chemists reason that the change in size is due to extra repulsive forces due to the added electrons, which push the electrons apart and make the atomic radius larger.