An atom is defined as the smallest constituent unit of an element that still retains all of the original properties of the element, and all matter is composed of atoms. Atoms may be broken into further components: protons, neutrons, and electrons. All atomic nuclei are comprised of positively charged protons and neutrally charged neutrons, meaning nuclei have an overall positive charge.
Negatively charged electrons orbit the nucleus in orbitals, with the orbitals closer to the nucleus having less energy than those farther away. Thus, overall atomic charge is determined by the number of positively charged protons and negatively charged electrons in an atom.
Every atom of an element has the same number of protons, which is that element’s atomic number. Elements are arranged on the Periodic Table of the Elements by their atomic number which increases from top to bottom and left to right on the table. Hydrogen, the first element on the periodic table, has one proton while helium, the second element, has two, and so on.
Along with atomic charge, atoms have measurable mass. An element’s mass number is the number of protons and neutrons it contains. (The mass of electrons is very small compared to protons and neutrons, so it is not included.) The number of neutrons in an atom can be found by subtracting the atomic number from the mass number.
While atoms of the same element have the same number of protons, their number of neutrons may vary. Atoms which differ in their number of neutrons but have equal numbers of protons are isotopes of the same element.
When writing the atomic symbol of an element, isotopes are differentiated by writing the mass number in the upper left-hand corner of the symbol. The atomic symbol for ordinary hydrogen is written as 1H, to signify that it has no neutrons and 1 proton, while deuterium, which is a hydrogen isotope with 1 neutron, is written as 2H.
The atomic mass of an atom, which is different from the mass number, is the average mass of all known isotopes of an element. For each element on the Periodic Table, the atomic number is listed above the symbol of the element and the atomic mass (measured in atomic mass units, or AMU) is listed underneath the symbol.
Atoms may lose or gain electrons, creating charged particles called ions. Ions are called cations if they are positively charged (due to the loss of electrons) or anions if they are negatively charged (due to the gaining of electrons). Ionic charges are denoted by adding a plus or minus sign onto the elemental symbol; for example, a sodium ion with a charge of +1 would be written as Na+.
Ions may be composed of two or more atoms known as molecular ions or polyatomic ions. The overall charge of a polyatomic ion is equal to the sum of the charges of all constituent atoms.
Common Polyatomic Ions
There are many useful physical and chemical patterns represented in the Periodic Table of the Elements. The periodic table is organized into rows called periods and columns called groups. The position of an element’s symbol on the periodic table indicates its electron configuration. The elements in each group on the table all contain the same number of electrons in their valence shell, which results in all elements in a group having similar chemical properties.
The majority of the elements in the periodic table are metals. Metals have the following properties:
Solid metals usually consist of tightly packed atoms, resulting in fairly high densities. Metals begin on the left side of the periodic table and span across the middle of the table, almost all the way to the right side. Examples of metals include gold (Au), tin (Sn), and lead (Pb).
Nonmetals are elements that do not conduct electricity and tend to be more volatile than metals. They can be solids, liquids, or gases. The nonmetals are located on the right side of the periodic table. Examples of nonmetals include sulfur (S), hydrogen (H), and oxygen (O).
Metalloids, or semimetals, are elements that possess both metal and nonmetal characteristics. For example, some metalloids are shiny but do not conduct electricity well. Metalloids are located between the metals and nonmetals on the periodic table. Some examples of metalloids are boron (B), silicon (Si), and arsenic (As).
Specific names are given to certain groups on the periodic table. Group 1 elements (belonging to the leftmost column) are known as the alkali metals and are characterized by the fact that they are very unstable and react violently with water. Other notably reactive elements are in Group 17, the halogens. In contrast to both of these groups, Group 18 contains the noble gases, which are inert because they have a full outer shell of electrons.
There are two periods below and separated from the main periodic table. These are called lanthanides and actinides. They are set apart from the other elements for two reasons: first, to consolidate the periodic table, and second, because they are more complicated chemically than the rest of the elements—which means that they do not follow any of the trends outlined below.
The Periodic Table is organized so that element show trends across periods and groups. Some of these trends are summarized below.
Electronegativity and ionization energy follow the same periodic trends. These two properties are simply different ways of describing the same basic property: the strength with which an atom holds electrons.
An atom’s electron configuration—the location of its electrons—influences its physical and chemical properties, including boiling point, conductivity, and its tendency to engage in chemical reactions (also called the atom’s stability). The chemical reactivity of an atom is determined by the electrons in the outermost (valence) shell, as they are first to interact with neighboring atoms.
Conventionally, electrons are depicted as orbiting a nucleus in defined pathways, much like a planet orbits the sun. In reality, electrons move in clouds surrounding the nucleus known as orbitals. Each orbital in an atom holds two electrons.
Orbitals are grouped into four types of subshells labeled with the letters s, p, d, and f. Each subshell has a specific number of orbitals:
The orbitals in each type of subshell have a particular shape. For example, the s subshell is spherical, while the p subshell is shaped like a bow tie.
Subshells are further grouped into shells, which are labeled with integers (1, 2, 3, …). The shell numbered 1 is closest to the nucleus, and the energy of the electrons in shells increases the further the shell is from the nucleus.
The location of an electron is described by its shell number and subshell letter, with the number of electrons in that orbital given as a superscript. The one electron in hydrogen, for example, is written as 1s1.
The orbitals for the first four shells are described in the table below.
Electron Configuration Notation
No. of Orbitals
No. of Electrons in Subshell
Notation for Full Subshell
Electrons fill orbitals in order of increasing energy, meaning they fill orbitals close to the nucleus before filling in outer orbitals. The order in which orbitals are filled is shown in the figure below.
The electrons in an atom’s outermost shell are its valence electrons. Most elements require eight electrons to fill their outermost shell (2 in s and 6 in p). So, elements with six or seven valence electrons are likely to gain electrons (and become cations). Conversely, elements with one or two electrons are very likely to lose electrons (and become anions). Elements with exactly eight electrons (the noble gases), are almost completely unreactive.
The electron configuration of each element correlates to its position on the periodic table: Group 1 and Group 2 (defined as the s-block) have valence electrons in s-orbitals. Elements in Groups 13 to 18 (defined as the p-block) have valence electrons in their p-orbitals. These groups (with the exception of the noble gases) are very reactive.
Group 3 through Group 12 elements (defined as the d-block) have valence electrons in d-orbitals. The lanthanides and actinides have valence electrons in their f-orbitals (and are called the f-block). The properties of these elements are less predictable because atoms’ d and f orbitals do not fill in a straightforward order.
The periodic table can be used to remember the order in which orbitals are filled: start at the upper left corner and move from left to right, then move down to the next row.