# General Chemistry/Chemistries of Various Elements/Inner Transition Metals

## Inner Transition Metals

The inner transition metals are found in the f-block, usually put at the bottom of the Periodic Table. These elements were sometimes called rare earth metals due to their extremely low natural occurrence. Except for extremely-unstable promethium which quickly decays to another lanthanoid metal, these elements are not rare. Indeed cerium is abundant in Earth's crust.) Many of them do not occur naturally, but are instead created in labs artificially. Furthermore, these elements all have nearly identical properties, both chemically and physically, making them very difficult to identify and separate. They are almost as reactive as the alkali metals, and all actinoids are radioactive, so they have little commercial significance. However, the radioactive elements can be used in nuclear power plants or as weapons.

Most of the inner transition metals form ions with a +3 charge. Some of the lighter actinoids can use their f-electrons for bonding, giving them a wider range of oxidation states, but the rest do not use f-electrons and have only a +3 oxidation state. Cerium is a notable exception: it has a somewhat common +4 oxidation state, seen in curium(IV) oxide CeO2.

These elements tarnish quickly in oxygen. Some will ignite in oxygen. They react with water to release hydrogen:

${\displaystyle 2{\hbox{M}}+3{\hbox{H}}_{2}{\hbox{O}}\to {\hbox{M}}_{2}{\hbox{O}}_{3}+3{\hbox{H}}_{2}}$

### Lanthanoids

Lanthanoids burn in oxygen easily and react violently with non-metals. They are used in lasers and sometimes steels depending on the element.

Neodymium magnets (Nd2Fe14B) are the strongest known permanent magnets. Gadolinium exhibits ferromagnetism below room temperature.

The terbium(III) cation is very fluorescent—it glows in the dark.

Lanthanoid contraction is a phenomenon that causes the lanthanoids (and all elements after them) to have much smaller atomic radii than expected. The f-electrons do not shield the nuclear charge as much as expected, so the outermost electrons are attracted to the nucleus more.

### Actinoids

Only thorium and uranium occur naturally in Earth's crust (along with neptunium and plutonium in trace amounts).

The actinoids are radioactive and decay into more stable elements. The actinoids that do not occur naturally have been created in labs for experiments and research.

## Nuclear Chemistry

A nuclear reaction

Throughout your study of General Chemistry, you have undoubtedly heard of "radioactive elements" and "unstable isotopes". These elements are the study of nuclear chemistry. Normal chemical reactions occur between atoms and electrons. Atoms gain, lose, and share electrons to form different substances. Chemical reactions are essentially interactions of electrons. Nuclear reactions, on the other hand, occur within the nucleus of an atom. They involve the gaining, losing, and transformation of protons, neutrons, and sometimes other particles (electrons and photons). Nuclear chemistry is something that you can study only within your textbook—radioactive substances are deadly to living things, can cause explosions, and are difficult to procure.

You should already know what isotopes are: elements with the same number of protons, but different numbers of neutrons (and a different total mass). Some isotopes are stable and do not decay. They last indefinitely. Other isotopes are unstable, meaning that they are radioactive. They will undergo nuclear reactions to become a more stable isotope. Some elements are always unstable, regardless of how many neutrons, so all of their isotopes are unstable.

For example, carbon-12 (6 protons, 6 neutrons) is stable. Carbon-14 (6 protons, 8 neutrons) is unstable and decays into nitrogen-14. This is unusual from a chemical point of view—there is no way for an atom to change into a different element. This is nuclear chemistry, though, and elements do change frequently in their quest to become more stable.

### Stability

There is no formula or exact rule to determine which isotopes are stable and which are unstable. That must be determined experimentally. Patterns have emerged throughout the study of the elements, and there are some general guidelines you can use to guess if an isotope will be stable or radioactive:

• Lighter elements are stable when they have roughly equal numbers of protons and neutrons.
• Heavier elements are stable when they have more neutrons than protons in about a 3:2 ratio.
• Elements that have a "magic number" of protons or neutrons are especially stable: 2, 8, 20, 28, 50, 82, 126.

In regard to the magic numbers, notice how helium-4 (2 p, 2 n) is the most abundant isotope in the universe. Lead-208 is the heaviest stable isotope known (82 p, 126 n). The air we breathe is filled with oxygen-16 (8 p, 8 n). The stability of these isotopes is no coincidence.

### Fusion and Fission

The fusion reaction that powers the Sun

Fusion reactions take two small nuclei and "fuses" them together into one large nucleus. Fission reactions split a large nucleus into smaller nuclei. Fission releases tremendous amounts of energy, which is why fission reactions are used in both nuclear power plants (to provide electricity to an entire city) and nuclear bombs (to destroy an entire city). Fusion reactions release even greater amounts of energy, but they only occur at unfathomably high temperatures. Fusion reactions occur in stars in outer space. Our sun is basically one giant fusion reactor. Hydrogen nuclei fuse together into helium nuclei, releasing the light and heat that warms our planet. Here are some example nuclear reactions:

 ${\displaystyle _{1}^{2}{\hbox{H}}+_{3}^{6}{\hbox{Li}}\to 2_{2}^{4}{\hbox{He}}}$ Fusion ${\displaystyle _{92}^{235}{\hbox{U}}+_{0}^{1}n\to _{52}^{141}{\hbox{Te}}+_{40}^{91}{\hbox{Zr}}+3_{0}^{1}n}$ Fission

Notice that the Law of Conservation of Matter is bent but not broken. If you add up the mass numbers, they will be equal on both sides of the reaction. The total charge numbers will also be equal.

### Decay Modes

An unstable isotope will decay to become more stable. There are many decay modes, but a few are common:

Common Decay Modes
• Alpha decay releases an alpha particle (helium-4, 2 p + 2 n). Occurs when the isotope is too big to be stable.
• Beta+ decay converts a neutron into a proton, releasing a beta particle (electron). Occurs when there are too many neutrons to be stable.
• Beta decay converts a proton into a neutron, releasing a beta particle (positron). Occurs when there are too many protons to be stable.
• Gamma decay releases a gamma particle (photon). Occurs when the nucleus has too much energy.
• Electron capture converts a proton into a neutron by absorbing an electron. Occurs when there are too many protons to be stable.

As far as health concerns, alpha particles are the most dangerous. They can be inhaled, causing bodily damage. They are heavy and have a double positive charge, but they are easily stopped by a piece of paper or skin. Beta particles are simply electrons (or positrons, an antielectron). They are somewhat dangerous, and they are stopped by a piece of wood or aluminum foil. Gamma rays are only stopped by thick slabs of lead. They are essentially x-rays that have extreme amounts of energy. Although they have the most energy, they only cause damage to things directly exposed to a radioactive substance. The other particles are worse because they can travel through the atmosphere.

A particular isotope always uses the same decay mode. These reactions will summarize the decay modes (notice the law of conservation):

 ${\displaystyle _{92}^{238}{\hbox{U}}\to _{90}^{234}{\hbox{Th}}+_{2}^{4}{\hbox{He}}^{2+}}$ Alpha decay ${\displaystyle _{92}^{238}{\hbox{U}}\to _{90}^{234}{\hbox{Th}}+\alpha }$ The same reaction, written with the more common notation. ${\displaystyle _{55}^{137}{\hbox{Cs}}\to _{56}^{137}{\hbox{Ba}}+_{-1}^{0}e^{-}}$ Beta— decay ${\displaystyle _{11}^{22}{\hbox{Na}}\to _{10}^{22}{\hbox{Ne}}+_{1}^{0}e^{+}}$ Beta+ decay