Atoms, Molecules, & Bonding
In this lesson, we will explore the structure of atoms and how they bond to form various molecules.
Atoms
A particle of matter that defines an element or material.
Atoms are extremely small, for example, a carbon atom weighs less than 2 × 10‾²³ grams. Atoms are composed of three subatomic particles: protons, neutrons, and electrons. The development of modern atomic theory revealed much about the inner structure of atoms. It was learned that an atom contains a very small nucleus composed of positively charged protons and uncharged neutrons, surrounded by a much larger volume of space containing negatively charged electrons. The nucleus contains the majority of an atom’s mass because protons and neutrons are much heavier than electrons, whereas electrons occupy almost all of an atom’s volume. For a perspective on their relative sizes, consider this: If the nucleus were the size of a blueberry, the atom would be about the size of a football stadium (Figure 2.11).
Figure 2.11 If an atom could be expanded to the size of a football stadium, the nucleus would be the size of a single blueberry. (credit middle: modification of work by “babyknight”/Wikimedia Commons; credit right: modification of work by Paxson Woelber)
A proton has a mass of 1.0073 amu (atomic mass unit; 1 amu = 1.6605 × 10-24 grams and a charge of 1+. A neutron is a slightly heavier particle with a mass 1.0087 amu and a charge of zero; as its name suggests, it is neutral. The electron has a charge of 1− and is a much lighter particle with a mass of about 0.00055 amu (it would take about 1800 electrons to equal the mass of one proton). The properties of these fundamental particles are summarized below.
Properties of Subatomic Particles
| Name | Location | Charge | Mass (amu) |
|---|---|---|---|
| electron | outside nucleus | 1− | 0.00055 |
| proton | nucleus | 1+ | 1.00727 |
| neutron | nucleus | 0 | 1.00866 |
The number of protons in the nucleus of an atom is its atomic number (Z). This is the defining trait of an element: Its value determines the identity of the atom. For example, any atom that contains six protons is the element carbon and has the atomic number 6, regardless of how many neutrons or electrons it may have. A neutral atom must contain the same number of positive and negative charges, so the number of protons equals the number of electrons. Therefore, the atomic number also indicates the number of electrons in an atom. The total number of protons and neutrons in an atom is called its mass number (A). The number of neutrons is therefore the difference between the mass number and the atomic number: A – Z = number of neutrons.
Atoms are electrically neutral if they contain the same number of positively charged protons and negatively charged electrons. When the numbers of protons and electrons are not equal, the atom is electrically charged and is called an ion. The charge of an atom is defined as follows:
Atomic charge = number of protons − number of electrons
The chemical symbol of an element contains information about the subatomic particles that make up the atom.
Remember that the periodic table identifies each element by its chemical symbol.
Some Common Elements and Their Symbols
Element | Symbol | Element | Symbol |
|---|---|---|---|
aluminum | Al | iron | Fe (from ferrum) |
bromine | Br | lead | Pb (from plumbum) |
calcium | Ca | magnesium | Mg |
carbon | C | mercury | Hg (from hydrargyrum) |
chlorine | Cl | nitrogen | N |
chromium | Cr | oxygen | O |
cobalt | Co | potassium | K (from kalium) |
copper | Cu (from cuprum) | silicon | Si |
fluorine | F | silver | Ag (from argentum) |
gold | Au (from aurum) | sodium | Na (from natrium) |
helium | He | sulfur | S |
hydrogen | H | tin | Sn (from stannum) |
iodine | I | zinc | Zn |
Chemical Bonding
A particle of matter that defines an element or material.
In ordinary chemical reactions, the nucleus of each atom (and thus the identity of the element) remains unchanged. Electrons, however, can be added to atoms by transfer from other atoms, lost by transfer to other atoms, or shared with other atoms. The transfer and sharing of electrons among atoms govern the chemistry of the elements. During the formation of some compounds, atoms gain or lose electrons, and form electrically charged particles called ions (see the figure below).
(a) A sodium atom (Na) has equal numbers of protons and electrons (11) and is uncharged. (b) A sodium cation (Na+) has lost an electron, so it has one more proton (11) than electrons (10), giving it an overall positive charge, signified by a superscripted plus sign.
You can use the periodic table to predict whether an atom will form an anion (negative charge) or a cation (positive charge), and you can often predict the charge of the resulting ionTo illustrate, an atom of an alkali metal (group 1) loses one electron and forms a cation with a 1+ charge; an alkaline earth metal (group 2) loses two electrons and forms a cation with a 2+ charge, and so on. For example, a neutral calcium atom, with 20 protons and 20 electrons, readily loses two electrons. This results in a cation with 20 protons, 18 electrons, and a 2+ charge. It has the same number of electrons as atoms of the preceding noble gas, argon, and symbolized Ca2+. The name of a metal ion is the same as the name of the metal atom from which it forms, so Ca2+ is called a calcium ion.
When atoms of nonmetal elements form ions, they generally gain. Atoms of group 17 gain one electron and form anions with a 1− charge; atoms of group 16 gain two electrons and form ions with a 2− charge, and so on. For example, the neutral bromine atom, with 35 protons and 35 electrons, can gain one electron to provide it with 36 electrons. This results in an anion with 35 protons, 36 electrons, and a 1− charge. It has the same number of electrons as atoms of the next noble gas, krypton, and is symbolized Br−.
Note the usefulness of the periodic table in predicting likely ion formation and charge (see the figure below). Notice that transition metals and some other metals often exhibit variable charges that are not predictable by their location in the table. For example, copper can form ions with a 1+ or 2+ charge, and iron can form ions with a 2+ or 3+ charge.
Some elements exhibit a regular pattern of ionic charge when they form ions.
Ionic Bonding
The chemical bonding that transfers electrons.
| NaCl, sodium chloride | Na2O, sodium oxide |
|---|---|
| KBr, potassium bromide | CdS, cadmium sulfide |
| CaI2, calcium iodide | Mg3N2, magnesium nitride |
| CsF, cesium fluoride | Ca3P2, calcium phosphide |
| LiCl, lithium chloride | Al4C3, aluminum carbide |
Covalent Bonding
The chemical bonding that shares electrons.
Many compounds do not contain ions but instead consist solely of discrete, electrically neutral molecules. These molecular compounds contain covalent bonds (and are covalent compounds). Covalent bonds result when atoms share, rather than transfer (gain or lose), electrons. Covalent bonding is an important and extensive concept in chemistry, and we will not need to discuss it in detail in this course. Whereas ionic compounds are usually formed when a metal and a nonmetal combine, covalent compounds are usually formed by a combination of nonmetals. Thus, the periodic table can help us recognize many of the compounds that are covalent. While we can use the positions of a compound’s elements in the periodic table to predict whether it is ionic or covalent at this point in our study, you should be aware that this is a very simplistic approach.
We can often identify molecular compounds on the basis of their physical properties. Under normal conditions, molecular compounds often exist as gases, low-boiling liquids, and low-melting solids, although many important exceptions exist.
When two nonmetallic elements form a molecular compound, several combination ratios are often possible. For example, carbon and oxygen can form the compounds CO and CO2. Since these are different substances with different properties, they cannot both have the same name (they cannot both be called carbon oxide). To deal with this situation, we use a naming method that is somewhat similar to that used for ionic compounds, but with added prefixes to specify the numbers of atoms of each element. The name of the more metallic element (the one farther to the left and/or bottom of the periodic table) is first, followed by the name of the more nonmetallic element (the one farther to the right and/or top) with its ending changed to the suffix –ide. The numbers of atoms of each element are designated by the Greek prefixes shown in the table below.
Nomenclature Prefixes
Number | Prefix | Number | Prefix | |
|---|---|---|---|---|
1 (sometimes omitted) | mono- | 6 | hexa- | |
2 | di- | 7 | hepta- | |
3 | tri- | 8 | octa- | |
4 | tetra- | 9 | nona- | |
5 | penta- | 10 | deca- | |
When only one atom of the first element is present, the prefix mono– is usually deleted from that part. Thus, CO is named carbon monoxide, and CO2 is called carbon dioxide. When two vowels are adjacent, the a in the Greek prefix is usually dropped. Some other examples are shown in the table below.
| Compound | Name | Compound | Name |
|---|---|---|---|
| SO2 | sulfur dioxide | BCl3 | boron trichloride |
| SO3 | sulfur trioxide | SF6 | sulfur hexafluoride |
| NO2 | nitrogen dioxide | PF5 | phosphorus pentafluoride |
| N2O4 | dinitrogen tetroxide | P4O10 | tetraphosphorus decaoxide |
| N2O5 | dinitrogen pentoxide | IF7 | iodine heptafluoride |
Polarity
The unequal distribution of charge.
If the atoms that form a covalent bond are identical, as in H2, Cl2, and other diatomic molecules, then the electrons in the bond must be shared equally. We refer to this as a pure covalent bond. Electrons shared in pure covalent bonds have an equal probability of being near each nucleus.
In the case of Cl2, each atom starts off with seven valence electrons, and each Cl shares one electron with the other, forming one covalent bond:
Cl+Cl⟶Cl2
The total number of electrons around each individual atom consists of six nonbonding electrons and two shared (i.e., bonding) electrons for eight total electrons, matching the number of valence electrons in the noble gas argon. Since the bonding atoms are identical, Cl2 also features a pure covalent bond.
When the atoms linked by a covalent bond are different, the bonding electrons are shared, but no longer equally. Instead, the bonding electrons are more attracted to one atom than the other, giving rise to a shift of electron density toward that atom. This unequal distribution of electrons is known as a polar covalent bond, characterized by a partial positive charge on one atom and a partial negative charge on the other. The atom that attracts the electrons more strongly acquires the partial negative charge and vice versa.
For example, the electrons in the H–Cl bond of a hydrogen chloride molecule spend more time near the chlorine atom than near the hydrogen atom. Thus, in an HCl molecule, the chlorine atom carries a partial negative charge and the hydrogen atom has a partial positive charge.
The distribution of electron density in the HCl molecule is uneven. The electron density is greater around the chlorine nucleus. The small, black dots indicate the location of the hydrogen and chlorine nuclei in the molecule.
Whether a bond is nonpolar or polar covalent is determined by a property of the bonding atoms called electronegativity. Electronegativity is a measure of the tendency of an atom to attract electrons (or electron density) towards itself. It determines how the shared electrons are distributed between the two atoms in a bond. The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity. Electrons in a polar covalent bond are shifted toward the more electronegative atom; thus, the more electronegative atom is the one with the partial negative charge. The greater the difference in electronegativity, the more polarized the electron distribution and the larger the partial charges of the atoms.
Intermolecular Forces
The forces that hold atoms together.
The differences in the properties of a solid, liquid, or gas reflect the strengths of the attractive forces between the atoms, molecules, or ions that make up each phase. The phase in which a substance exists depends on the relative extents of its intermolecular forces (IMFs). molecules. IMFs are the various forces of attraction that may exist between the atoms and molecules of a substance due to electrostatic phenomena. Under appropriate conditions, the attractions between all gas molecules will cause them to form liquids or solids. This is due to intermolecular forces, not intramolecular forces. Intramolecular forces are those within the molecule that keep the molecule together, for example, the bonds between the atoms. Intermolecular forces are the attractions between molecules, which determine many of the physical properties of a substance.
Intramolecular forces keep a molecule intact. Intermolecular forces hold multiple molecules together and determine many of a substance’s properties.
The strengths of these attractive forces vary widely, though usually, the IMFs between small molecules are weak compared to the intramolecular forces that bond atoms together within a molecule.
All of the attractive forces between neutral atoms and molecules are known as van der Waals forces.
Van der Walls Forces
Dispersion Forces
Because the electrons of an atom or molecule are in constant motion (or, alternatively, the electron’s location is subject to quantum-mechanical variability), at any moment in time, an atom or molecule can develop a temporary, instantaneous dipole if its electrons are distributed asymmetrically. The presence of this dipole can, in turn, distort the electrons of a neighboring atom or molecule, producing an induced dipole. These two rapidly fluctuating, temporary dipoles thus result in a relatively weak electrostatic attraction between the species—a so-called dispersion force. Dispersion forces that develop between atoms in different molecules can attract the two molecules to each other. The forces are relatively weak, however, and become significant only when the molecules are very close.
Dispersion forces result from the formation of temporary dipoles, as illustrated here for two nonpolar diatomic molecules.
Dipole-Dipole Attractions
Polar molecules have a partial positive charge on one side and a partial negative charge on the other side of the molecule—a separation of charge called a dipole.
A dipole-dipole attraction is an electrostatic force between the partially positive end of one polar molecule and the partially negative end of another.
This image shows two arrangements of polar molecules, such as HCl, that allow an attraction between the partial negative end of one molecule and the partial positive end of another.
Hydrogen Bonding
Hydrogen bonding occurs when exceptionally strong dipoles attract. It is a bond that exists when hydrogen is bonded to one of the three most electronegative elements: F, O, or N.
Despite use of the word “bond,” keep in mind that hydrogen bonds are intermolecular attractive forces, not intramolecular attractive forces (covalent bonds). Hydrogen bonds are much weaker than covalent bonds, only about 5 to 10% as strong, but are generally much stronger than other dipole-dipole attractions and dispersion forces.
Water molecules participate in multiple hydrogen-bonding interactions with nearby water molecules.
