Last updated 11-1-10
<-CHM130

Chemical Bonds: The Atom Connection

(PowerPoint about bonds presented to Quinn Thacker's class:  BondsForQuinnsClass.pptx) 
(PowerPoint about solubility presented to Quinn Thacker's class:   SolubilityForQuinnsClass.pptx)

In the Building Block Tutorial Series, one of the tutorials covered the basics of chemical bonding. We will review some of those concepts and also update it using concepts from the Modern Atom tutorial.

Atoms are made from protons, neutrons, and electrons. Since bonding involves electrical attraction, it's the protons and electrons that are more important.

Here are two hydrogen atoms. Notice that the electron on the left atom is attracted to both its own proton and to the right proton of the right hydrogen. Likewise, the electron in the right hydrogen is attracted to both protons. So there's a balance of repulsion and attraction. With no attraction, atoms would never make a bond. With no repulsion, the atoms would just merge together as one atom. So balancing attraction and repulsion allows for atoms to form compounds.
Two hydrogen atoms will stay together because the electrons have room to stay away from each other as they orbit both hydrogen atoms. So even though there is repulsion between the electrons and between the protons, this arrangement minimizes the repulsions allowing for the opposite charge attractions to keep them together.To see animation move cursor over the image. When two atoms share electrons, we say they are bonded together.
To the left are two fluorine atoms. The nucleus of the fluorine atom has nine protons. This is possible because the neutrons help hold them together. The nine fluorine protons have a tremendous pull (attraction) to electrons in the vicinity. Two electrons orbit close to the nucleus. The outer seven spread out to maximize space between them. However there is still space for one more electron. Notice how the two electrons are being attracted by protons from both atoms. This pulls the atoms together, and because there is room, the atoms bond.

Here the fluorine atoms have pulled together and are sharing one electron each. They are now bonded. This is called a single bond (even though two electrons are involved) Because they pull on these shared electrons equally, they call this a covalent bond. "co" meaning "shared" and "valent" meaning "strength". So the strength of the bond is shared equally.

Atoms naturally pull on each other because the protons in them pull on the electrons of the other atoms. However, there has to be room for the electrons to merge and be shared by both atoms, otherwise electron repulsion keeps them apart.

Note: Most non-metal atoms have room for eight outer electrons.

Here is carbon. It has six protons, six neutrons, and six electrons. You can see that there are four outer electrons and spaces for four more electrons. From this picture you can see that carbon could accommodate four electrons and share four of its electrons. Can carbon connect to fluorine?

Fluorine "wants" one electron to fill up its outer shell of electrons. Carbon will provide that electron as long as fluorine shares one of its electrons with carbon because carbon also wants to attain 8 outer electrons.

You can see how four fluorine atoms can connect to one carbon atom. They come together to form carbon tetrafluoride, which is a gas used in refrigeration systems. With nine protons, fluorine's pull on that electron is stronger than carbon's pull because carbon only has 6 protons. This uneven sharing of the electrons means fluorine is slightly negative because the electrons are pulled closer to it. Carbon is partially positive because its outer electrons are closer to the fluorine atoms and not canceling out its own positive charge. For this reason, they call this a polar covalent bond. "Covalent" because the electrons are shared and "polar" because there's a + and - charge separation (like the ends of a pole). Roll cursor over the image to see the + and - charged poles.

 

The electron can change shape in order to maximize space between it and other electrons. Electrons are always moving and their exact position is not set but based on probability. In other words, the orbits we've been showing are their most probable position. But there is a chance that electrons can be elsewhere at least for a moment. There's probably several electrons in your body that may somewhere else in the room at this moment. Again, electrons are amazing little "particles."

In the above movie, you see carbon with its six electrons orbiting it. Again, this is a simplistic view of the electrons. First, let's change the inner two electrons to the spherical electron cloud that better represents them. The next two electrons also occupy a spherical shell. The next two electrons, however, form the shape of four lobes with each electron being two lobes. Sometimes when carbon bonds with other atoms, these four outer electrons will blend their shapes to form four equally shaped lobes that project out in four directions. The arrangement of the lobes match a four sided pyramid called a tetrahedron. A tetrahedron allows the electrons maximum separation in 3 dimensions.

Lewis Dot Structure
vs.
Costello Hole Punch Structure
 The carbon on the right is my invention of showing the outer electrons of the atom and its vacancies. The usual way of showing this is in the left image. The dots represent the 4 outer electrons of carbon. They call this a Lewis Dot Structure, named after the chemistry instructor, Gilbert Lewis, who used these images to help his students remember how many outer electrons elements have and how they might bond. My image shows the protons, the inner electrons, the outer electrons, and the vacancies used for sharing electrons. Because I show vacancies, I call it the Hole Punch Structure.
Below are five ways of representing a carbon atom. The far right considers a more modern view of the electron showing them in orbitals.
IONIC BONDING

Ionic Bonding: Who bonds with whom? The non-metals in the upper right have a strong pull on electrons and, in contrast, metals tend to let go of their electrons easily. So when they get together, you can guess what happens. It's like putting a greedy person together with a generous person. You can guarantee money will leave the generous person and end up owned by the greedy person. The same thing is true about non-metals and metals. In contact with each other, the metals are usually giving away one or more of their electrons to the non-metals. When this happens, the metals become positive ions because they lost one or more the negative electrons. The non-metals become negative ions because they gained electrons. The metal and non-metal will attract each other because they have opposite charges. That's ionic bonding.

The exception are the noble (inert) gases. They already have 8 outer electrons, which is a very stable configuration. So they won't share any electrons and therefore won't bond with other elements.

The halogens (group 17) are one electron short of a stable 8 configuration (called an octet). That makes them the most "greedy" of all elements. In contrast the metals in groups 1 & 2 are the metals that give up their 1 or 2 electrons the easiest.

So when the halogens come in contact with the metals from the first two groups, step back because there's going to be a violent reaction as electrons transfer from these metals to the halogen elements.

Here chlorine gas is reacting with sodium metal. After all the fireworks, all that's left is NaCl that you can place on your French fries (assuming you have equal numbers of chlorine and sodium atoms present).

Here's what's happening at the atomic level. Sodium atoms have one electron in the outer orbit (s orbital). If sodium loses the electron it will have a stable configuration of 8 outer electrons (an octet). Chlorine gas travels in pairs. Chlorine has 7 outer electrons, so there's one vacancy. To have 8 outer electrons, each chlorine shares one of its electrons with the other chlorine atom. That's a covalent bond. Each sodium gives its outer electron to one of the chlorine atoms. Now the chlorine ions have 8 outer electrons without having to share. That's even more stable.

After those two electrons have transferred, the sodium atoms are now positively charged (+1) because the protons in the nucleus out number the electrons in the atom by one.  The chlorine atoms are now negatively charged (-1) because each gained an extra electron.

All atoms are now ions because they have a charge.  The two sodium ions will push away from each other because they are both positive. The two chloride ions will push away from each other because both are negative.  However, since sodium and chlorine are now oppositely charged, they will be attracted to each other. That attraction is a bond appropriately called an ionic bond.
Ionic bonds don't share electrons; one atom has taken one or more electrons completely away from the other atom. They will become oppositively charged and will then stick together (bond) because of the electrical attraction. In a crystal of NaCl, the sodium and chlorine atoms will alternate so as to minimize repulsion between like charges. In 3 dimensions, NaCl will be stack in cubes with Na and Cl alternating. (roll cursor over image to see lattice).

Na+ and Cl- are both very stable because they have the 8 outer electron configuration (octet). That's why our world has trillions of tons of salt (NaCl).

So what does this octet look like?
The octet is a stable configuration. It contains a full s orbital (2 electrons) and three full p orbitals (6 electrons). Here are 4 ways of showing the stable 8 outer electrons (octet rule). In the upper left, my image shows a chlorine that has eight outer electrons. The red one is an extra one, which makes the chlorine a negative one charge. The Lewis dot structure shows Cl with 8 dots around it. The brackets show that the chlorine has a negative one charge, meaning it has an extra electron that wasn't its own. The lower drawings show the s & p orbitals. The p orbitals align with the x, y, and z axes. The left image shows a more definite boundary for the orbitals (which make them easy to see). The right image is a plotting of the probability of where the electron is, which is more cloud-like. When full, each of these 4 orbitals have two electrons each giving us the octet.

In the earlier tutorial on Types of Chemical Reactions, we focused on the building block aspect of reactions. In this tutorial on bonds, we can focus on the force & energy aspect of bonds involved in these reactions.

Types of Chemical Reactions
Synthesis (Combination)
Decomposition
Single Replacement/Displacement
Double Replacement/Displacement
Oxidation (Combustion)

2Al + 3O2 -> 2Al2O3
 Aluminum reacts with oxygen to make aluminum oxide. Since aluminum loses all 3 of its outer electrons to oxygen, it forms an ionic bond and is very strong.

 Synthesis: Inorganic compounds
A metal with a non-metal. They form ionic bonds because electron(s) will leave the metal and be captured by the non-metal. That will make them oppositely charged, which then draws them together. We covered the NaCl example. Oxygen is the most common example. Metals in ores or jewels are combined with oxygen.

Synthesis: Organic compounds
All organic compounds have carbon in them. As covered earlier, carbon has 4 outer electrons, so it can share these 4 while accepting four. The bonds are covalent bonds since it's a sharing of the electrons. More specifically it's a polar covalent bond because, as mentioned before, the fluorines have a much stronger pull on the shared electrons than does carbon. The molecule, as you might guess, is carbon tetrafluoride.

The fluorine bonds are stronger than oxygen bonds, which is why this compound won't burn. Oxygen can't break the fluorine-carbon bond.

Here is carbon tetrafluoride again but written using Lewis dot structures. The top row are the reactants. The fluorine atoms travel in pairs because they each have 7 outer electrons, but overlapping one of their electrons with the other fluorine atom gives both a virtual octet, which is more stable than traveling alone. The bottom left molecule shows how carbon shares one electron with each of the fluorine atoms while at the same time borrowing one from each fluorine. Here again, all atoms get a virtual octet, which is a stable configuration. Normally the sharing of one electron each is shown as a bar (bottom right compound). I added the blue color to help keep track of carbon's electrons, but normally there's no color shown.

Decomposition: For decomposition reactions it's nice to have compounds whose bonds can break fairly easily. This is the ammonium dichromate decomposition reaction shown in the Type of Reactions tutorial. Regarding bonds, you can tell from the products which bonds are the strongest. Apparently N2, H2O, and Cr2O3 have stronger bonds. Water is a product which means the hydrogens came off of the ammonium (NH3) and the oxygens came off of dichromate (Cr2O7). You could deduce that the bonds in water (oxygen to hydrogen) are stronger than the nitrogen to hydrogen bonds in ammonium (NH4+) and some of the oxygen to chromium bonds in the dichromate (Cr2O7).

Also, since 3/4 of the world's surface is water, we could say the bonds in water must be strong (stable). So a lot reactions produce water because once hydrogen and oxygen come together to make water, it's hard to break those bonds.

Single Replacement/Displacement

This is the example reaction given in the Types of Reaction tutorial.
Zn + FeCl2(aq) -> Fe + ZnCl2(aq)
The (aq) means it's dissolved in water. That's always a clue that it's an ionic bond. Also because it's a metal combined with a non-metal, that's an ionic bond because the metal loses its electron(s) to the non-metal. Ionic bonds are usually quite strong. For example, you can heat FeCl2 up to 1300F and it will melt but not decompose. However, put a few drops of water on it at room temperature, and the bonds will break. So ionic bonds are interesting. They are very resistant to heat but many will be easily broken by water.

Here's an example with salt. You can heat it to 3000F and it will melt but not decompose. The sodium and chlorine atoms will stay together. However, if you put a little water on it, the bonds start breaking immediately. That's because water has a plus end and a minus end (polar). The minus end by the oxygen pulls on the positive sodium ions. The positive side of water (hydrogen side) pulls on the negative chloride ions. Because water is in motion, it carries the sodium and chlorine ions away.
Below is the double replacement example given in the tutorial Types of Reactions. The (aq), again, means the compound is currently dissolved in water. The bottom equation shows the ions separately, which is more the way they would be in water. Before being dissolved in water, barium chloride was held together by an ionic bond. Likewise, the magnesium ion bonded to the sulfate polyatomic ion with an ionic bond (+ and - attraction). Remember non-metals bonded with non-metals share their electrons. So it's covalent bonding. When both barium chloride and magnesium sulfate were dissolved in water and mixed, the barium ion has a chance to bump into the sulfate ion. This time this ionic bond is strong enough to keep water from pulling it apart. So particles of barium sulfate form in the water and drop to the bottom. This solid that comes from compounds that had been dissolved in the water is called a precipitate.
Here is the sulfate ion shown with the Lewis Dot structure (except I've color coded the dots). The red dots, of course, are the electrons of oxygen (6 per oxygen). Sulfur can share its six electrons, but it's two short of sharing with the fourth oxygen. That's why there needs to be two extra electrons that came from somewhere else (black dots). That also gives the whole polyatomic ion a minus 2 charge.

Oxidation (Combustion)
This is where bond strength makes combustion possible. The 3-carbon hydrocarbon is propane. When it burns, bonds between the hydrogen and carbon atoms are broken and replaced with bonds to oxygen atoms. Notice that the bond energy of oxygen to hydrogen (in water) is higher than the bond energy of hydrogen to carbon (in propane) 111 kcal vs 98.7 kcal. Especially notice the double bond energy of carbon to oxygen is more than twice as strong as the bonds between carbon atoms (192,000 calories per mole vs. 82,900 calories per mole). As bonds are broken, energy is absorbed. As bonds are formed, energy is released. This shows that the bonds forming as oxygen connects with hydrogen and carbon atoms produces more energy than what was needed to break the bonds. Therefore, we've got fire (heat & light). The propane heater in this hot air balloon depends on the bond energies of C=O and O—H being higher than C—H and C—C.
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Since April. 23, 2008