Let’s begin with lithium. Chemists have known for more than a century about lithium’s strong tendency to ionize as well as its very small atomic mass. In fact, being from row 2 as well as a group 1 metal of the periodic table makes lithium able to carry large amounts of charge in a relatively small mass of material. Does this sound like something that might be useful?
Lithium in Daily Life
Lithium-ion batteries are some of the most powerful and compact energy sources available. Lithium’s low mass, but strong tendency to become a charged ion, has been employed to provide electrical power whenever weight matters, as it does in portable electronic devices and vehicles.
Lithium’s light weight and metallic properties have made it valuable for other applications. Pure lithium metal is far too reactive to be used as a manufacturing material, but when alloyed with other light metals like aluminum, lithium’s low density can be exploited without its vigorous reactivity getting in the way.
In alloys between lithium and aluminum, the metallic bonding is much more stable than in pure lithium, while being even lighter than aluminum, without sacrificing strength. Such alloys have been particularly useful in aviation, and have been responsible for an initial surge in global lithium production that began in the 1950s, long before the development of the lithium-ion battery.
This article comes directly from content in the video series Understanding the Periodic Table. Watch it now, on Wondrium.
Sodium is much more abundant in our environment, but its atomic mass is not particularly low like lithium. Made of atoms more than three times the mass, and forming a metal nearly twice the density of lithium, sodium’s usefulness is less linked to the size of its atoms. It is, however, remarkably useful for some of its other properties, such as the way it emits light.
Most of us are familiar with the warm, yellow glow of certain street lamps and lighting in other industrial settings. But did you know that sodium is the element responsible for this type of outdoor lighting?
While most other elements produce line emission spectra with multiple intense lines of varying color, the excited sodium atom produces two unusually intense lines at about 589 nm in its emission spectrum. This means that when sodium vapor is stimulated by electricity, it emits intense, yellow light. Sounds like the perfect design for a light bulb, doesn’t it?
By enclosing sodium in an evacuated bulb, heating it until that sodium vaporizes, and passing an electric current through it, a handy light bulb can be made. Sodium lamps were developed in the 1920s and were valued at the time by scientists who needed monochromatic light for their experiments. Decades before the laser, sodium lamps offered a simple, affordable source of nearly monochromatic light that often fit the bill.
Sodium Lighting Has Problems
But the impact of sodium in lighting goes beyond just the laboratory. Because of their economy, availability, and easiness on the eye, sodium lights became a very popular option for urban lighting in the mid-1900s. By the 1960s many urban streets were bathed in the warm, yellow glow of sodium lamps each night.
Astronomers used to love urban sodium lighting. Its single intense yellow wavelength minimized its impacts on their astronomical observations, and the impacts it did have on those observations were very easy to subtract because of its nearly monochromatic nature.
But the problem these lamps posed was that monochromatic light can make certain objects difficult to see on the ground. For example, bathing a red traffic sign in pure yellow sodium light doesn’t make it very visible, since there is hardly any red light to reflect and send back to our eyes.
So, in the interest of public safety, many communities have recently abandoned the use of sodium lights in favor of newer technologies that produce a broad spectrum of wavelengths. This improves the visibility of road signs and obstacles, protecting motorists and pedestrians.
Healthy Group 1 Metals
Unlike sodium, potassium does not produce such a simple atomic emission spectrum, nor does it vaporize quite as easily. So potassium never became a staple material in the lighting industry. But sodium and potassium do share remarkably important roles in human health.
Both sodium and potassium are critical nutrients in organisms, helping to manage such critical functions as muscle contraction and hydration. Because both are so abundant and ionize easily, the body has evolved to use sodium and potassium transport to help maintain and establish the charge separations that create such essential functions as muscle contraction.
In 2003, Roderick MacKinnon won the Nobel Prize in Chemistry for his characterization of potassium ion channels from cells. Ion channels are remarkable proteins that span the cell membranes in a variety of tissues, shuttling potassium back and forth as the body needs.
But these ion channels make us vulnerable to certain venoms produced by predatory animals. For example, the venom of scorpions contains chemical compounds that stick to the opening of such ion channels, halting the transport of potassium, causing paralysis and even death in any animal unfortunate enough to receive a sting or a bite. Simply by halting the flow of this single element through the body, these venoms can paralyze or even kill.
Common Questions about the Group 1 Metals that Influence Our Everyday Lives
Lithium is a Group 1 metal that can carry large amounts of charge in small amounts of mass. Lithium’s low mass and strong tendency to become ionized make it a great candidate for use in batteries.
Though sodium has excellent properties to be used as a light source, there is a downside to its monochromatic nature. Since the light is almost entirely monochromatic it makes things, such as traffic signs, less visible. Abandoning sodium lights is in the interest of pedestrians and motorists.
Both of these group 1 metals are crucial nutrients for organisms because they help to manage essential functions such as muscle contraction and hydration.