Because there’s a lot of chemistry vocabulary out there, I’ve decided to break it down into several pages so you don’t have to scroll so much. This page has vocabulary words that start with the letters I-L, as I’m sure you probably already figured out.
Incidentally, if you see an asterisk (*) after the definition, there will be more information about the topic at the bottom of the page when you scroll down. I’ve included what I think is most important in the main definition, but figured that some of you might need a bit of supplemental information.
ideal gas: A gas that follows all of the postulates of the kinetic molecular theory (infinitely small particles, kinetic energy proportional to temperature, random travel, no intermolecular forces). Though no gas behaves as an ideal gas in the real world, real world gases are usually within 5% or so of what would be expected of an ideal gas.
ideal gas law: PV = nRT, where P = gas pressure, V = gas volume (L), n = moles of gas, R = ideal gas constant, and T = temperature (in K).*
ideal solution: A solution in which the vapor pressure is directly proportional to the mole fraction of solvent present.
immiscible: When two liquids do not form a solution with one another. Oil and water are immiscible, while water and rubbing alcohol are not.
indicator: A compound that turns different colors at different pH values. The most useful indicators for many purposes change color at a pH near 7 because that’s the equivalence point in acid-base titrations. The most commonly-used indicator in acid-base titrations is phenolphthalein (colorless in acid, pink in base) and one of the most commonly-recognized is litmus (red in acid, blue in base).
inhibitor: A substance that slows down a chemical reaction. Think of it as a catalyst in reverse, because instead of decreasing the activation energy of a reaction it increases it.
inorganic compound: Any non-carbon containing compound, with the exception of the oxides of carbon (carbon monoxide, carbon dioxide, carbonate-containing compounds).
insoluble: If something is insoluble, it does not dissolve to any appreciable extent.*
intermediate: A chemical species that is formed in one step of a reaction and consumed in another.
intermolecular force: A force (not a bond!) that exists between covalent molecules. Intermolecular forces all depend, in one way or another, on the attraction of partial positive charges on one molecule to partial negative charges on another (a.k.a. electrostatic forces). The three intermolecular forces include dipole-dipole forces (partial charges in adjacent molecules attract), hydrogen bonds (strong dipole-dipole forces in molecules with H-F, O-F, and N-F bonds that include the partial connection of lone pairs between the hydrogen atom and other atom on adjacent molecules, and London dispersion forces (in which the temporary polarization of nonpolar molecules causes this attraction.) London dispersion forces are the weakest intermolecular force, followed by dipole-dipole forces, followed by hydrogen bonding.*
ion: An atom, or group of atoms, that contain positive charge (cations) or negative charge (anions). Ions that contain only one atom are called “monatomic ions”, while ions with more than one atom are called “polyatomic ions.” (Examples of common ions)
ionic bond: A bond formed when ions stick to one another electrostatically.*
ionization energy: The amount of energy needed to remove an electron from an atom in the gas phase. This trend generally follows those for electron affinity and electronegativity.
irreversible reaction: A reaction that can proceed only in the forward reaction, with no observed back reaction. Though all reactions are technically reversible, in practice many reactions behave as if they were irreversible.*
isotonic solutions: Solutions that contain the same osmotic pressure as one another. This is mainly useful in understanding cell biology, as an isotonic cell will remain the same size and a cell with a different tonicity than its surroundings will tend to either expand or collapse.
isotopes: The different forms of an element in which the atoms have the same number of protons but different numbers of neutrons (and hence, different masses). Many isotopes are stable, while others are radioactive.*
Kelvin (K): The SI unit of temperature. The Kelvin scale is an absolute temperature scale, meaning that 0 K is defined as absolute zero. One Kelvin is the same size as one degree Celsius – the only difference is that the Celsius scale is offset by the Kelvin scale by 273 degrees Celsius (i.e. 0° C = 273 K).*
ketone: An organic molecule containing R-CO-R’ functionality. Acetone (dimethyl ketone, 2-propanone) is a very common one.*
kinetic energy: The energy that an object has due to its motion. The kinetic energy of an object is equal to 1/2 mv², which indicates that mass plays a role in kinetic energy, as does the object’s velocity.*
lanthanide: Elements in the 4f section of the periodic table (lanthanum → ytterbium).*
lanthanide contraction: The tendency of the elements in the lanthanide series to get smaller as you move across (left → right) the periodic table.*
lattice energy: The energy that’s released when one mole of a crystal is formed from gaseous ions. Though this is the official definition, the useful lesson to get out of lattice energies is that some crystals (i.e. ones with more negative lattice energies) are more stable than others (i.e. ones that are less negative).*
lattice: The three-dimensional arrangement of atoms in a crystal. Lattices are made up of repeating units called “unit cells”, which can be stacked like a bunch of identical Legos to generate a structure for the whole crystal (link).
law of conservation of energy: Energy is neither created nor destroyed in any process. Or, put another way, the total amount of energy in the universe is constant. Or put another way, it’s the first law of thermodynamics. No matter how you phrase it, however, this law basically states that energy can be converted from kinetic to potential (and back again), but no matter what happens it’s around somewhere.*
law of conservation of mass: Mass is neither created nor destroyed in any process. A simple example: If you make a pizza out of 1 kg of crust, 1 kg of cheese, and 1 kg of sauce (not a good pizza, I know), the final pizza will weigh 3 kg, even though it doesn’t look the same.*
Le Châtelier’s principle: If you disturb an equilibrium by manipulating one of its components, the equilibrium will respond in such a way as to minimize whatever it is you did. For example, if you have an equilibrium where A ⇌ B, adding more of compound A will disrupt the equilibrium. As a result, the forward reaction rate of A → B will increase to get rid of the added A (this is referred to as “shifting the equilibrium to the right” or “shifting the equilibrium toward the products.”) Eventually the equilibrium will be reestablished, though not in the exact same way as it was before the disruption was made.*
Lewis acid/base: A Lewis acid is an electron-pair acceptor and a Lewis base is an electron-pair donor. Though this sounds a lot different than the Arrhenius and Brønsted-Lowry definitions, it’s really just an extension of them that was designed to cover a wider range of reactions. General chemistry students won’t run into this definition all that much in practice, but organic, organometallic, and coordination chemists use it a lot.*
Lewis structure: A structural formula that shows all of the atoms and valence electrons in a molecule. You know those diagrams you see with the atoms and the lines between them? That’s a Lewis structure.*
ligand: A molecule or ion that sticks to the central atom in a coordination complex. Examples include carbon monoxide, water, and ammonia (link).
limiting reagent: If you do a reaction with two reagents, it’s the reagent that runs out first. It’s called the “limiting reagent” because the amount of product that can be formed depends on the quantity of the limiting reagent present. This is in contrast to the “excess reagent”, of which there is more than enough.*
line spectrum: A spectrum that shows only certain wavelengths of light. Each of these lines in the spectrum corresponds to a particular energy difference between a ground state and excited state orbital.*
London dispersion force: A very weak intermolecular force that arises when nonpolar molecules, through the random motion of electrons, become very slightly polar and stick to one another. This force tends to be stronger in large molecules because there are more electrons to become unbalanced in the first place. The term “London dispersion forces” and “Van der Waals forces” are often used interchangeably, which isn’t technically correct but not a big deal.*
lone pair (a.k.a. unshared pair): Two electrons on an atom that aren’t involved in chemical bonding. These lone pairs can act as nucleophiles/Lewis bases when reacting with other molecules.
ideal gas law: The value of P in the ideal gas law is typically given in atmospheres or kilopascals, and the value of the ideal gas constant is either 8.314 J/K·mol or 0.08206 L·atm/mol·K, depending on the pressure units used for P. Incidentally, the Van der Waals equation is sometimes used in place of the ideal gas law when additional accuracy is needed, because it has terms that account for the nonideality of various gases.
indicator: There are other types of indicator as well, but you’re unlikely to run into them in an introductory chemistry class. What all indicators have in common, however, is that a color change indicates the endpoint of the process.
insoluble: It should be noted, however, that everything dissolves to some degree or another in any solvent. From this table, you can see that even compounds normally thought of as insoluble (such as calcium carbonate) have very slight solubility in water.
intermolecular force: The term “London dispersion force” is sometimes used interchangeably with “Van der Waals force.” Strictly speaking, Van der Waals forces include several other intermolecular forces, none of which are typically mentioned in an introductory chemistry class. Additionally, it should be noted that hydrogen bonding does not actually involve covalent bonding, but as an unusually strong intermolecular force has sometimes been mistaken for them.
ionic bond: Some people (including me!) would argue that ionic bonding isn’t really bonding at all, as true bonding involves the sharing of electrons. There are arguments for and against this. Supporting this belief are theories of bonding that involve the overlap of orbitals – this is something that doesn’t occur in ionic bonds. However, one might correctly make the argument that, no matter what the electronegativity difference between the two atoms, there remains some sharing of electrons. You’ll have to make the call for yourself, but whatever you decide, I would recommend that you remain aware that the distinction between ionic and covalent bonding is fairly arbitrary, as both types of bonds involve some degree of electron sharing and some degree of electron transfer.
irreversible reaction: The principle of microscopic reversibility states that all reactions that proceed in the forward reaction can proceed, under the right conditions, in exactly the opposite fashion to regenerate the reagents. This is the main idea that explains how chemical equilibria are established, though it should be noted that there are many factors that keep reactions from being significantly reversible. For example, if a process has a very high activation energy in the reverse direction, one would expect the reverse reaction to proceed at a very slow rate. Another example would be one of combustion, where the products of the reaction are blown away via smoke, making it nearly impossible to collect them and cause them to revert into reagents.
isotope: A weighted average of the isotopic masses of an element make it possible to find the average atomic mass of an element. Interestingly, there is no one “average atomic mass”, as different sorting processes cause the concentrations of each isotope to be slightly different in different geographic locations. Also, when the TV says that there’s an “isotope” present, it often implies that isotopes are inherently dangerous. That’s not true (all atoms are isotopes of some element), and you shouldn’t believe what your stupid TV says.
Kelvin (K): A common and somewhat annoying mistake is the tendency of some people to refer to a temperature as being “350 degrees Kelvin” or “350 Kelvins.” The Kelvin scale is not described as being in degrees and is not pluralized, so the temperature above would be described simply as “350 Kelvin.” One other point: The actual conversion between the Celsius and Kelvin scales is 273.15 degrees Celsius, but people usually drop the 0.15 degrees because it’s usually pretty insignificant.
ketone: The naming of ketones is very rarely done via the IUPAC method, but rather by other methods that people have devised. As a result, you’ll have naming differences such as those present in methyl ethyl ketone (MEK), which is called butanone by IUPAC.
kinetic energy: To give a more technically-correct definition, the kinetic energy of an object is the amount of energy that’s needed to accelerate it from rest to its current velocity. Kinetic energy can be converted to potential energy and vice-versa.
lanthanide: There is some debate about which elements are actually lanthanides. Some people say that it includes the elements from lanthanum to lutetium, while others (including me) believe that it includes the elements from lanthanum to ytterbium, with lutetium in group 3 of the transition metals. Muddying the waters is the fact that lanthanum is sometimes treated as a transition metal, with the elements cerium through lutetium treated as lanthanides. Which is true? Who knows? IUPAC, and more importantly, Wikipedia say that lanthanum through lutetium are lanthanides. I’d recommend just going with whatever periodic table your teacher gives you.
lanthanide contraction: Why this happens is kind of interesting. It seems that the 4f electrons in the lanthanides, while good at doing chemistry-type stuff, aren’t so good at shielding the 6s electrons from the positive charge in the nucleus. As a result, these elements get smaller as the number of protons increases. There is also an “actinide contraction”, but people talk about it a lot less (J. Am. Chem. Soc., 1995, 117 (24), pp 6597–6598. Link)
lattice energy: In real life, this value usually can’t be measured experimentally because it’s hard to get ions such as Na+ and Cl- into the gas phase. As a result, there are a wide variety of equations you can use to figure out what this should be, and they seem to work pretty well, as far as people can tell (link).
law of conservation of energy and/or mass: Though it may seem obvious to all of us that these laws are true, looking at it a little more closely shows why it took people quite a while to figure it out.
- Why the law of conservation of energy looks false: If you have a piece of wood, it just sits there at room temperature. When you light it on fire, it gets very hot, indicating that energy was created. When it goes out, it gets cold again, indicating that the energy vanished. From the viewpoint of people who don’t know much science, the obvious conclusion is that energy can be forced to appear and vanish under the right conditions. To us, it’s clear that the wood has energy stored in chemical bonds, which then goes into heating up various molecules when it is released, and then the heat dissipates into the environment, making it seem as if it vanished.
- Why the law of conservation of mass looks false: If you burn wood, you may start out with 1 kg of wood and end up with 0.1 kg of ash. The conclusion that ancient people would naturally draw from this is that, as the wood vanished, it turned into heat somehow. When the heat was used up, the wood was gone. Of course, we know that the mass didn’t vanish – it just got converted to gaseous carbon dioxide and water vapor and drifted away.
Le Châtelier’s principle: I gave the example of addition of A to this reaction, but there are other ways in which this equilibrium could also be disturbed:
- If B is added, the reaction will shift to the left to get rid of some of it (forming more A in the process).
- If B is removed, the reaction will shift to the right to replace the B that was removed (this is a trick that’s used to force equilibria to make more product than you’d normally expect to see).
- If A is removed, the reaction will shift to the left to make more. Why you’d do this, nobody knows.
- For exothermic reactions (i.e. where heat is a product), cooling the reaction will shift the reaction to the right (to reform some of the lost heat) and heating it will shift it to the left (to get rid of the added heat).
- For endothermic reactions (i.e. where heat is a reagent), heating the reaction will shift it to the right (heat acts as a reagent in this case, so the reaction tries to get rid of it) and cooling it will shift it to the left (to replace the lost heat).
- For gaseous equilibria, changing the pressure of the compounds might make a difference in the position of the equilibrium. If you increase the partial pressure of one of the reagents or products, that’s the same thing as adding them in the bullet points above – similarly, decreasing partial pressures serves to remove them. However, if you add another gas that’s not in the equilibrium, there will be no shift to the equilibrium because the partial pressures (i.e. concentrations) of the reagents and products will be constant.
- If a catalyst is added, the position of the equilibrium will not be affected. However, one would expect the equilibrium to be established more quickly, as both the forward and reverse reactions would proceed at a faster rate.
Lewis acid/base: Like the Brønsted-Lowry definition, compounds are not inherently Lewis acids or bases, but behave as Lewis acids or bases in a particular reaction. That said, it’s unusual to find some classes of compounds behaving as other than Lewis acids (the carbon in organic carbonyl groups comes to mind) or Lewis bases (amines do this).
Lewis structure: There are actually all kinds of structures that can be drawn in addition to Lewis structures, including Fisher projections (common in biochemistry), Newman projections (common in organic chemistry), and skeletal formulas (simplified Lewis structures used in organic chemistry). Incidentally, the Lewis structures you see with two dots showing each bond (rather than the lines) are used only to teach the basic ideas of Lewis structures and are not used by chemists.
limiting reagent: You’d think that it’s a good idea to always have the perfect amount of each reagent present, but that’s not always the case. For example, if you have one reagent that’s considerably more expensive or rare than the other, you’d want to make sure there was plenty of the other reagent to make sure that all of it reacts. Additionally, equilibria with a large amount of excess reagent will lie far to the right/products, which is a good thing.
line spectrum: To say more about this, a line spectrum occurs when the following process takes place: 1) An atom has energy added to it through electricity, heat, or something else; 2) This causes a ground state electron to be raised in energy into an excited state orbital; 3) When the electron returns to the ground state, the energy it absorbed is given off as light that has the same energy (and corresponding color) as the difference in energy between the ground and excited state. This idea is the basis behind spectroscopy.
London dispersion forces: Though I’ve described this force as taking place only in nonpolar molecules, it’s actually present in all compounds. However, because London dispersion forces are extremely weak when compared to other forces, we normally notice them only in nonpolar molecules.
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