look like?
There are three main shapes of orbitals relevant to most of chemistry, and these are referred to as s, p, and d orbitals. The designations s, p, and d are abbreviations for sharp, principle, diffuse, and fundamental, which have historical significance describing results of early experiments to probe the electronic structure of atoms. You can see what these orbitals look like in the graphic on the preceding page.
The shapes of these orbitals are determined by their orbital angular momentum, which is a property that describes the motion of the electron around the nucleus.
What is the valence shell of electrons?
Electrons fill up orbitals in “shells.” The innermost shell consists of just one s-type orbital and can hold just two electrons. The next shell consists of one s-type and three p-type orbitals, and can hold eight electrons. Higher shells consist of more and more orbitals and can thus hold more and more electrons. The valence shell of electrons is the highest occupied, or partially occupied, set of orbitals.
What is the atomic radius of an atom?
The atomic radius of an atom is defined as half of the distance between two atoms of the same element held together in a chemical bond. Not surprisingly, these are very small distances! For hydrogen, the smallest atom, the atomic radius is 0.37 Ångströms, or 3.7 × 10−11 meters.
How do the atomic radii of atoms change across the table?
The atomic radii of atoms generally decrease going from left to right across a period, and increase going top to bottom down a group (see graphic on next page).
The increase in atomic radius going down a group is fairly straightforward to understand: additional shells of electrons are added and they must surround the inner shells, resulting in an increased atomic radius. Though the number of protons in the nucleus increases going down a group, the inner shells of electrons serve to shield the valence shell from the attractive force of the nucleus, resulting in an overall increase in atomic radius.
Moving to the right across a period, the number of protons increases, increasing the attractive force on electrons in the valence shell. Within a period, additional electrons go into the same valence shell, and an increasing attractive pull from the nucleus results in a more contracted valence shell, resulting in a smaller atomic radius. The situation is complicated by the rightmost group (known as the Noble gases), but the atomic radius of these elements is typically not important as they are rarely involved in chemical bonds to other atoms.
The atomic radii of atoms generally decrease going from left to right across a period, and increase going top to bottom down a group. (Atomic sizes are not to precise ratios and are for illustrative purposes only.)
What is the ionization energy of an atom?
The ionization energy of an atom is the amount of energy it takes to remove an electron from the atom. The process of removing an electron leaves the atom with an extra proton, relative to the number of electrons, and thus creates a positively charged ion, known as a cation. The ionization energy can be thought of as a measure of how strongly an atom holds on to its electrons. In general, ionization energies increase from left to right across a period (though there are exceptions) due to an increasing number of protons to attract electrons in the valence shell. Ionization energies decrease going down a group in the periodic table, due to the valence electrons being farther from the nucleus, and thus more shielded from its positive charge. Note that the trends in atomic radii and ionization energy go in the same direction—larger atoms tend to have lower ionization energies.
What keeps an electron from crashing into the nucleus?
Opposites attract, so electrons and protons are attracted to each other, making it somewhat difficult to understand why an electron wouldn’t just get as close as possible to the nucleus and crash into it. The key to answering this question has to do with the fact that electrons are very, very small particles, so they are governed by rules that don’t apply to larger objects. As we’ve talked about a little already, electrons are best thought of as clouds of negative charge surrounding the nucleus. Their properties are governed by rules that describe the cloud as a whole, rather than as a single particle. It turns out that there is something favorable about the electron being spread out, or delocalized, around the nucleus. For reasons we won’t go into in detail, when the electron’s cloud gets packed closer to the nucleus, the energy associated with its motion (its kinetic energy) begins to rise, which makes the situation unstable. There’s a balance between the stability associated with placing the electron close to the nucleus (the favorable positive-negative charge attraction) and that associated with spreading out the electron’s cloud (to keep its kinetic energy low). This prevents the electron’s cloud from getting too close to the nucleus or the electron just crashing into the nucleus.
MOLECULES AND CHEMICAL BONDS
What is a molecule?
A molecule is a set of atoms held together by chemical bonds. Molecules are the smallest units of a substance that behave as that substance, and separating the atoms of a molecule will change its properties.
What is a substituent?
A substituent is an atom, or group of atoms, attached to a specific position in a molecule. For example, in the molecule 3-bromopentane (see drawing below), we could refer to the bromine as a substituent on the third carbon atom.
What is a chemical bond?
A chemical bond is an attractive interaction that binds atoms together through a sharing of electron density. The simplest bonding arrangement involves just two electrons shared between nuclei such that each effectively has a stable octet of eight valence electrons (or just two in the case of H–H). When two atoms are sharing a total of two electrons between them, the atoms are referred to as singly bonded to each other.
Bonds are what hold atoms together in molecules, and they are usually not easily broken. The arrangement of atoms in a molecule determines the identity of a chemical compound. The making or breaking of bonds is a chemical reaction, which converts one chemical compound into another.
Can I think of chemical bonds as springs between atoms?
Chemical bonds can be thought of as springs holding together the atoms in a bond. When atoms in a bond are stretched or compressed from their equilibrium separation, the bond provides a force to pull the atoms back together or to keep them from getting too close. For relatively small displacements, the bond actually provides a force that is physically very similar to that of a spring connecting two objects. This model of a spring as a chemical bond can be very useful for getting an intuitive idea of how a chemical bond connects atoms in a molecule.
What is a Lewis structure?
Lewis structures are a simple way of depicting the electronic structure of atoms and molecules. They show us which atoms are bonded to each other in a molecule and also show how many nonbonded electrons are present in the valence electron shell of each atom. The easiest way to understand them is probably to just take a look at a few.
The simplest
