Science: An Elementary Teacher’s Guide/Chemical Reactions
Any objects you see around you are made of molecules that are held together by chemical bonds. Those bonds were formed during chemical reactions. Some reactions make bonds between atoms or molecules, while other reactions break those bonds. For example, plants use sunlight as the energy source that powers a series of chemical reactions to build sugar molecules out of carbon dioxide and water. In taking apart water molecules, the Hydrogen atoms become part of the sugar, but the Oxygen atoms do not get used--they combine with each other and produce O2, which exits the plant leaf, enters the atmosphere, and is the oxygen you are breathing in right now. The plant can combine those sugar molecules into larger molecules such as starch, or even wood. In turn, when you eat a leaf of spinach or a potato, or you burn wood, you break the bonds between the sugar molecules. An interesting surprise is that in order to do this you need to use oxygen (which you get through breathing) and in the process of taking apart the food you produce carbon dioxide and water! The energy that was stored in those chemical bonds helps warm your body and also acts as an energy source for new chemical reactions, allowing you to build new molecules within your cells.
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When it comes to forming chemical bonds, it is important to understand the idea of valence electrons, which refers to the electrons in the outer "shell" or "cloud" or "orbital" around the atom. The electrons in the outermost occupied shell determine the chemical properties of the atom. Basically, as more electrons are added, they automatically arrange themselves in a manner that keeps them as far apart from each other as possible (since their negative charges repel each other), while still being attracted to the nucleus. For our purposes, just know that the first shell can hold only 2 electrons. When that first shell is full, electrons start to fill up the next shell which holds 8 electrons. Additional shells can hold even more, but let's just emphasize that having a full outer shell is very stable. To help understand, we are going to assign some "personalities" to different elements depending on where they are in the Periodic Table.
Some atoms will not combine with any other atoms--they are completely non-reactive. They can be found in the far right column of the periodic table and are called noble gases. Their personality is "prideful" and "stuck up," and they refuse to play with anyone and will never share. The reason they are non-reactive is that their outer (valence) shell of electrons is completely full. Helium is the first in that column and has just 2 electrons, which perfectly fills the first electron shell. Next down is Neon--it has 10 electrons, which is 2 plus 8--again, the outer shell is naturally all the way full. All other atoms "wish" they could be like their "noble" cousins. They will react in some way that will bring them closer to having their own outer shells full. For some, the quickest way to complete the outer shell is by getting rid of an electron or two (if your outer shell just has 1 electron and you donate that to another atom, then that shell just disappears and the full shell beneath is now the outer shell), and for others they need to add an electron or two. If your first shell has 2 electrons and your second shell has 8 electrons, and your next shell has 7 electrons (Chlorine, Cl), then you desperately want one more electron so you can be like your noble next-door neighbor Argon. On the other hand, if your first shell has 2 electrons, your second shell has 8 electrons, and your third shell has 1 electron (Sodium, Na), then you desperately want to get rid of 1 electron (which would leave you looking like the noble gas Neon). In their pure state both Chlorine and Sodium are dangerously reactive, but once they have found each other and sodium has given its electron away to chlorine then both are incredibly stable and nonreactive. This exchanging of an electron is discussed further under Ionic bonding. If, in order to get to a full outer shell, you need 2 or 3 or 4 electrons then it is easier to share electrons with others, which will be discussed under covalent bonding.
Four types of bonding
editIonic bonds
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Atoms can give or take electrons to get to a state where they resemble their noble cousins, but they can never change their number of protons. An atom in its native state has the same number of electrons as it does protons, and the negative charges cancel out the positive charges, so the atom is electrically neutral. If it adds an electron, it will end up with a -1 charge because it now has 1 electron more than it has protons. If a different atom subtracts an electron, it will end up with a +1 charge (because it has one more proton than it does electrons). Some atoms will give one electron (1st column of the Periodic Table), others are willing to give 2 electrons (2nd column).
Ions are atoms that have lost or gained electrons, which make them cations (they have a positive charge, because they lost one or more electrons) or anions (they have a negative charge by having gained one more electrons). An electron can only be given away if another atom is willing to take it, so a cation and an anion are produced together when the electron transfer occurs. Now you have two atoms next to each other that are both happy and non-reactive because their outer shells are full, but they now have an electric charge and are electrically attracted to each other! Positively charged cations and are attracted to the negatively-charged anions. Positively charged and negatively charged ions stay stuck to one another via an electrostatic attraction.
The most famous ionic compound is NaCl or Sodium Chloride aka table salt. It's what makes our steaks, corn, eggs, popcorn and coffee taste good. . . Wait. You guy don't put salt in your coffee? Ignore I said that. Any "salts" (such as magnesium chloride or potassium chloride) are the result of ionic bonds. They can form crystals because the charged atoms arrange themselves in patterns, but the crystals dissolve easily in water (the polar water molecules are able to surround the ions and pull them away from each other).
The Sodium or Na in salt is actually Na+ and the Chloride or Cl is actually Cl-; the Sodium in its native state has just one electron in its outer shell--by giving up this electron it then has a complete outer shell. Chlorine has 7 electrons in its outer shell and by gaining one more it completes its outer shell. Before reacting, both sodium and chlorine are dangerous elements, but once they have completed the electron transfer from sodium to chlorine they become chemically inert (nonreactive).
Just to emphasize one more time, the periodic table has atoms in columns with similar properties. The next-to-last column on the right of the periodic table has the halogens (Fluorine, Chlorine, etc.). These are the most electronegative atoms, meaning they want an electron so strongly they are ready to steal one from anything! (this is why chlorine gas is dangerous--it will quickly react with your eyes and skin, burning you as it takes electrons). These atoms will always gain an electron during a reaction, so they will be negatively charged as ions. The column to the left of that has atoms that are capable of gaining two electrons. To the very left of the periodic table (1st column), you will get the alkali metals and alkali earth metals (Lithium and Beryllium respectively and the elements below them) and they always lose electrons and have a positive charge.
Ionic bonds form when elements interact and completely move an electron from one atom to another. It is considered a weak bond since the ions are only attracted to each other electrically, but not held together chemically (you can break ionic bonds just by putting the crystals in water). Compare this to covalent bonds, where electrons are shared between atoms.
Covalent bonds
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Atoms have electrons. Bonds are made of electrons and atoms either share them or take them. Covalent bonds are the "sharing" of electrons. The result is that atoms stick tightly together and make molecules! Covalent bonds are very strong, and almost all objects around you are held together with covalent bonds (an exception would be if you are looking at salt crystals--they are held together with ionic bonds!). One atom, Carbon, is like a happy, generous kid who wants to share with everyone. Carbon has 4 electrons in its outer shell--it is too much work to take 4 electrons from other atoms, and it is too much work to give 4 electrons away. Instead, Carbon says, "I'll share with you if you share with me." Carbon gets along very well with many different types of atoms, including Hydrogen. Hydrogen is the smallest atom, with just one proton and one electron. Hydrogen would love to have 2 electrons, but is such a little guy that it cannot take an electron from anyone. But if someone wants to share, Hydrogen is always ready to do that. Carbon will share an electron with Hydrogen, and another electron with another Hydrogen, and third electron with another Hydrogen, and fourth electron with a 4th Hydrogen! This would be CH4, which is methane gas. Or, two Carbon atoms can share one electron with each other, then each of them has 3 more electrons available to share, so you can get C2H6, which is butane gas. Carbon is super important for life, because it can form 4 different covalent bonds--it is like a Lego or a Tinker Toy that can join all sorts of things together. Even plastics are made with long molecules built mainly out of Carbon and Hydrogen!
The usual convention is that the closer 2 elements exist on the periodic table the more likely they form a covalent bond with each other because they have similar electron affinities (electron desires). Two Oxygen atoms are both very electronegative and will share equally with each other in an O2 molecule. Remember covalent means to co-exist with electrons, or to share the same electrons between the valence shells of two different atoms. Now I just said that covalent bonds form more likely between atoms right next to each other or close to each other. . . Water has covalent bonds but oxygen and hydrogen are so far from each other!! But they are both non-metals. Also atom type (i.e. non-metal, metal, alkali metal, halogens, chalcogens, etc.) can dictate bond type and generally the same types are grouped together on the Periodic Table.
Covalent bonds can be polar or non-polar. . . Non-polar covalent are things like H-H (hydrogen gas H2) or Cl-Cl (poisonous chlorine gas, Cl2) or (C-C) carbon-carbon bond. These are 2 atoms of the exact same affinity or electronegativity. It is analogous to saying we have identical twins of same strength pulling on either side of a rope. . . The rope is not moving. Hence electrons are equally shared between participants in the bond, and the two ends of the atom are not any different from each other--there is no "pole," so they are "non-polar." Oils and grease and wax and gasoline and most gases are examples of non-polar compounds.
The O-H bond of water is very polar due to difference in electronegativity. Covalent bonds are sharing, sharing, sharing. . . So the more one atom in the bond relationship is more selfish the less covalent in nature it is BUT we don't call it "non-covalent covalent". . . That would not make sense . Instead, it is called polar covalent bond to indicate the one-sided nature of the electron sharing relationship (i.e. in the O-H bond the electrons spend most of their time with the oxygen because the pull is so strong). Think of Oxygen as a bit of a bully. It may take your electrons (and become an ion), but it is more likely to "share" your electrons in the same way that a bully might "share" a pencil with you as long as you share a pen with him, and both the pencil and the pen spend most of their time on the bully's desk! So when Oxygen shares an electron with one Hydrogen, the Hydrogen gets to feel like its outer shell is full (because it went from having just 1 electron of its own to now having 2), but both electrons spend more time with Oxygen than they do with Hydrogen. Oxygen gets to feel like its outer shell is full by forming 2 covalent bonds, because it needed 2 to complete the orbital. This is why a water molecule is H2O. But since Oxygen is more electronegative, ALL the electrons spend more of their time with Oxygen than they do with the two Hydrogens. This results in the Oxygen side of the molecule being slightly negative and the Hydrogen side of the molecule being slightly positive--in other words, the two sides of the molecule are not the same, so it has polarity. (Don't feel bad for Hydrogen--it gets a special bond named after it!).
Hydrogen bonds
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A hydrogen bond is a partial electrostatic attraction between different molecules, or between different parts of a single, larger molecule. In other words, it is similar to ionic bonding, but weaker because instead of two ions (which have full +1 or -1 charges) the attraction is between molecules that have polarity and only partial positive or negative charges. It is kind of like two really weak magnets--they are attracted to each other, but the bond is also easy to break. Hydrogen bonding is seen when a hydrogen (H) is bound to a more electronegative atom such as nitrogen (N), oxygen (O), or fluorine (F), with oxygen being the most common example. We discussed polarity in water molecules, and the resulting hydrogen bonding between water molecules gives water some unique properties, which are discussed separately. Even though hydrogen bonding is the weakest of the bond types we are discussing, it is biologically significant--the three-dimensional structure of proteins relies on hydrogen bonds, and the two strands of DNA are held together by hydrogen bonds (because they are weak, this also allows the strands to be unzipped for copying), and hydrogen bonding is important in cellulose, cotton, and other fibers.
Metallic bonds
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Metallic bonds are very important--most elements on the periodic table are metals, in fact, and they form strong bonds with each other. However, we will not be discussing the nature of these bonds in depth because metallic bonding can get very confusing (it has confused chemists and physicists for a long time!). The simple way to say it is that metals are kind of like super-sharers when they are with other metals. They don't exactly share electrons to fill their outer shell, and in fact they don't hold on very well to their own electrons--instead, the metal nuclei are "floating" in a sea or cloud of delocalized electrons. Each nucleus is strongly attracted to the cloud of electrons around it, giving metal strength. Every atom nucleus has protons and neutrons (big, charged, dense center) and metal atoms are big. For example, atomic element Copper (Cu) has atomic number 29, so every Copper atom in a lattice (arrangement) has 29 protons in the nucleus and 29 electrons to donate to the Electron Sea which is a part of the Copper mass . . . given that in an old copper penny there are around 3 X 1022 atoms, that is a whopping lot of electrons!! So each time you look at a hunk of metal don't be shocked that humans have created their electricity from metals and not plastics, wood, or some other solid. The other atoms just don't have enough electrons floating around. And that is what electricity is! The movement of electrons. The electrons of metals move very freely between the atoms, and this arrangement of electrons is responsible for many of the properties of metal--they are shiny, malleable (can be bent and shaped without breaking), they are strong, they can be combined well with other metals (to form alloys like brass or stainless steel), and they are excellent conductors of electricity and heat. Metals also bind with non-metals (especially oxygen), and often that is how they are found in nature, as ore bound tightly in rocks. Gold, silver, and platinum are "precious metals" because they are non-reactive and won't corrode and have been assigned special value by humans, with a long use as art, jewelry, and currency.
Combining different ratios of metals as alloys modifies the properties of pure metals to produce desirable characteristics. The aim of making alloys is generally to make them less brittle, harder, resistant to corrosion, or have a more desirable color and luster. Of all the metallic alloys in use today, the alloys of iron (including steel, stainless steel, cast iron, and more) are the most common. Being able to extract metals from ore, and being able to combine metals into alloys have been important steps in human history (the Bronze Age and the Iron Age).
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Making and breaking molecules
editExothermic and endothermic
editExothermic and Endothermic reactions can be taken apart by their word roots. Exo- "out", Endo- "in", thermic- "heat". An exothermic reaction releases energy meanwhile and endothermic reaction needs energy. So, sometimes in a reaction scheme you'll see . . A + B ---> C + D + Heat or A + B ---> C + D + # of Joules. In both examples you see that the heat or joules are on the products side of a reaction which means the heat was released. In a similar example you'll see that heat was on the reactants side meaning the heat or energy inputted was needed for reaction to happen.
Basically Endothermic and Exothermic are opposites. BUT we must remember that heat needed or expelled in a reaction does not indicate the speed of reaction or thermodynamic availability. If a reaction is exothermic it is generally spontaneous (it will happen on its own) but that doesn't mean it will be fast or slow. So, there are a mix of categories that could go with one another but for now we are defining exothermic (heat released) and endothermic (heat needed).
Exothermic processes Endothermic processes making ice cubes melting ice cubes formatino of snow in clouds conversion of frost to water vapor condesation of rain from water vapor evaporation of water a candle flame forming a cation from an atom in the gas phase mixing sodium sulfite and bleach baking bread rusting iron cooking an egg burning sugar producing sugar by photosynthesis forming ion pairs separating ion pairs combining atoms to make a molecule in the gas phase splitting a gas molecule apart mixing water and strong acids mixing water and ammonium nitrate mixing water with an anhydrous salt making an anhydrous salt from a hydrate crystallizing liquid salts (as in sodium acetate in chemical handwarmers) melting solid salts nuclear fission reaction of barium hydroxide octahydrate crystals with dry ammonium chloride mixing water with calcium chloride reaction of thionyl chloride (SOCl2) with cobalt(II) sulfate heptahydrate
Oxidation and Reduction
editOxidation-reduction reactions are electron transfers between two objects. For simplicity this type of reaction should have simply been called "electron transfer," but it was named oxidation-reduction because 1) oxidation is what happens when a metal is exposed to oxygen--for instance, iron turning to rust. Chemists of old realized this and said the iron had been oxidized. Later, when studying this type of reaction in depth it was learned that oxygen does not have to be involved. It was also realized that there was a loss, or reduction, in charge from the other agent involved in the reaction. Hence, reduction and oxidation happen at the same time, and sometimes this reaction is just called "Redox." In redox reactions the oxidation number of an ion, molecule, or atom changes by either losing or gaining an electron. One way to remember the process of Redox reaction is through the mnemonic OIL RIG. OIL = Oxidation Is Loss (of an electron) and RIG = Reduction Is Gain (of an electron).
Reduction
Oxidant + e– ⟶ Product
(Electrons gained; oxidation number decreases)
Oxidation
Reductant ⟶ Product + e–
(Electrons lost; oxidation number increases)
In a Redox reaction there is always a receiver and a donor, since an electron is being transferred. As an analogy, in an exchange money if an individual goes to a store the individual will lose money and the store will gain money through the transaction. Oxidation and reduction go hand in hand. Without one, the other ceases to exist. There is a Crash Course video about redox reactions if you want to learn more.
Through the process of Redox reactions, molecules are built and taken apart.
Acid-base reactions
editAn acid is a molecule that is capable of donating a proton or hydrogen ion (H+). Remember, Hydrogen has one proton, one electron, and no neutrons. If you remove the electron from hydrogen, what you are left with is called a hydrogen ion (because it has a charge), but you could also say what you had left was just a single proton. In some instances, people talk about protons being donated, and in other cases, they talk about hydrogen ions being donated--it is the same thing! A strong acid dissociates (comes apart) very strongly. For example, when hydrochloric acid, HCl, is mixed in water it breaks apart into H+ and Cl- almost completely (very few of the HCl molecules remain intact). Vinegar is a much weaker acid. It begins as CH3COOH and breaks down into CH3COO- and H+. However, many of the molecules remain intact. By donating fewer H+ it is considered a weaker acid. Remember, pH is a measure of H+ concentration.
A base is a molecule that accepts H+, thus removing them from the solution. Hydroxy ions (OH-) are the most common base that accepts the H+ (because if you join H+ and OH- you get H2O, which is water), so sometimes a base is defined as a molecule that donates hydroxy ions. Like with acids, a strong base dissociates more completely than a weak base. Lye, or sodium hydroxide (NaOH) is a very strong base. An example of a weaker base would be baking soda, sodium bicarbonate (NaHCO3).
In an acid-base reaction, the acid and base neutralize each other and produce a salt. A simple example is mixing our strong acid, Hydrochloric Acid, and our strong base, Sodium Hydroxide. HCl + NaOH --> H2O + NaCl
The acidic properties of hydrochloric acid were neutralized, and the basic properties of sodium hydroxide were neutralized--we combined two dangerous chemicals and made salt water! When you combine vinegar and baking soda you get a release of carbon dioxide gas (CO2), but you also neutralize the acid and base.
This is an oversimplified discussion of acid-base reactions but should be sufficient for our purposes.