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Take a look at the periodic table. Generally, there are two types of elements:. ..and nonmetals. metals…. Please note: H is considered a NONMETAL. Since bonds occur between two atoms. It should make sense that there are three types of bonds:. bonds between two metal atoms….
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Take a look at the periodic table. Generally, there are two types of elements: ..and nonmetals metals… Please note: H is considered a NONMETAL
Since bonds occur between two atoms It should make sense that there are three types of bonds: bonds between two metal atoms… …and bonds between one metal atom and one nonmetal atom bonds between two nonmetal atoms… Na Na Na Cl Cl Cl
Let’s first look at this type of bond: the one that occurs between a metal and a nonmetal. Recall that metal atoms have only 1, 2 or 3 loosely held valence electrons which they tend to give up rather easily. Na Cl And that nonmetals have 5, 6 or 7 tightly held valence electrons and tend to gain electrons to fulfill their octets. As you would expect, the Na atom gives up its electron and becomes a positively charged ion. …and the Cl atom gains an electron and becomes a negatively charged ion.
Let’s first look at this type of bond: the one that occurs between a metal and a nonmetal. Recall that metal atoms have only 1, 2 or 3 loosely held valence electrons which they tend to give up rather easily. Na Cl And that nonmetals have 5, 6 or 7 tightly held valence electrons and tend to gain electrons to fulfill their octets. As you would expect, the Na atom gives up its electron and becomes a positively charged ion. …and the Cl atom gains an electron and becomes a negatively charged ion. Since these ions have opposite charges…
Let’s first look at this type of bond: the one that occurs between a metal and a nonmetal. Recall that metal atoms have only 1, 2 or 3 loosely held valence electrons which they tend to give up rather easily. And that nonmetals have 5, 6 or 7 tightly held valence electrons and tend to gain electrons to fulfill their octets. As you would expect, the Na atom gives up its electron and becomes a positively charged ion. …and the Cl atom gains an electron and becomes a negatively charged ion. Since these ions have opposite charges…they are attracted to one another. This is known as an “ionic bond.”
But this doesn’t just happen with one Na atom and one Cl atom, it happens with several trillion of each… and the ions that form arrange themselves so that each positive ion is totally surrounded by negative ions… and vise-versa…
row after row… layer upon layer… to make what is known as a 3-dimension crystal lattice structure.
The crystal shown below contains only 200 ions (100 Na’s and 100 Cl’s), and if it were “life-size,” it would be way too tiny to see, even with the most powerful microscope. Since the ions are in a 1:1 ratio, the formula for this ionic compound is simply NaCl
Ionic compounds are often referred to as “salts.” The most common salt, of course, is the one we have been talking about – sodium chloride (NaCl) – the same one you use on your food: but many other salts also have this same structure.
There are thousands of different salts (ionic compounds): KF, NaBr, CuO, FeBr2, Li3P, Co2S3… Any compound that forms between a metal and a nonmetal is likely to be an ionic compound. Notice that not all these compounds are 1:1 salts, some are 1:2 or 3:1 or even 2:3… These salts would obviously have different structures than the one shown below:
Here’s a different kind of crystal structure – one that a 1:2 compound might have. Can you see that there are twice as many yellow ions as there are blue ions in the pattern? - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - - 2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+
Do ionic compounds have high or low melting points: Now do parts 1 & 2 in the Bonding Lab. These parts examine some of the properties of ionic substances. Return to this tutorial when you are finished with parts 1 & 2.
So, what was happening at the atomic level when we were melting the NaCl crystals? In the diagram below, it looks as though the ions are not moving at all, but… they are actually moving all the time, but not nearly fast enough to overcome the strong ionic bonds holding them in their rigid crystal structure. So all they can do is just vibrate in place. But then we started heating the salt crystals up…
And that made them move faster and faster… Eventually, the ions move fast enough to start overcoming the ionic bonds, and when that happens,
And that made them move faster and faster… Eventually, the ions move fast enough to start overcoming the ionic bonds, and when that happens, they start to move around each other… and that’s what melting is all about. For NaCl, the melting point is around 800*C! That’s hot! Ionic substances have such high melting points because ionic bonds are so strong and hard to break. Then we cooled the molten salt down by pouring it into the large beaker and…
that made the ions freeze back into their rigid crystal structure again. Then we tried bending the crystal and found that it was quite brittle. Why?
Imagine taking a tiny hammer and hitting the ionic crystal. That would push down one block of ions…
And when that happens, neg. ions would end up next to neg. ions and pos. next to pos. This would cause a strong repulsive force…
Which would cause the crystal to break apart. Hence, ionic crystals tend to be quite brittle: that is, they tend to break rather than bend.
Covalent bonds are the bonds that form between nonmetals, and they involve electrons being shared between two atoms. C C Sometimes this goes on and on in all directions to make what is called a network covalent crystal …like diamond: C C C C C C Other times, it joins together just a small distinct number of atoms and that creates a molecule C C C C C …like water: C C C We’ll come back to molecular substances in a little while… C C C The sand that we considered melting was a network covalent crystal. H O H
The structure for sand (SiO2) is shown below. The red atoms are Si’s and the blue atoms are O’s. But please note, these are NOT positive and negative ions bonded together as NaCl was (ionically). These are atoms which are bonded by sharing electrons between them (covalently). And covalent bonds are much stronger than ionic! To melt sand would require breaking covalent bonds… And that takes higher temperatures than a burner flame can reach.
We’ve talked about strong crystals like NaCl that are held together by ionic bonds between metal and nonmetals. And we’ve talked about even stronger crystals like SiO2 that are held together by covalent bonds between two nonmetals. Now let’s talk about metallic crystals. These are made up of metal atoms bonded together to other metal atoms. 2+ 2+ 2+ According to current theories, metallic bond- ing occurs when a cluster of metal atoms… Zn Zn Zn 2+ 2+ 2+ all get rid of their valence electrons: Zn Zn Zn That gives the atoms a sort of pos. charge… 2+ 2+ 2+ Zn Zn Zn But these delocalized electrons form a kind of “sea of electrons” which acts like a glue to hold these positive metal atoms together…. 2+ 2+ 2+ 2+ Zn Zn Zn
But this “sea of electrons” makes for a weak bond when compared to ionic bonding, and especially when compared to network covalent bonding. This is why it was so much easier to melt the zinc than to melt either the salt in Part 1 of the lab or the sand in part 3. Metals in general have lower melting points than ionic or network covalent substances. There are about 15 different metals that melt at temperatures below 500*C. 2+ 2+ 2+ Zn Zn Zn 2+ 2+ 2+ Gallium (Ga) has a melting point of just 34*C. That means a piece of gallium would literally melt in your hand! Zn Zn Zn 2+ 2+ 2+ Zn Zn Zn And of course there’s mercury (Hg) which has such a low melting point (-40*C) that it is a liquid at room temperature! 2+ 2+ 2+ 2+ Zn Zn Zn
And because they are not locked into specific energy levels, this “sea of electrons” is able to absorb and then re-emit all frequencies of visible light… This accounts for why metals are shiny like the Zn sample you observed in the lab today. And the fact that there are already electrons flowing throughout the crystal accounts for why metals are such good conductors of electricity. Best of all, this model of metallic bonding helps explain why metals are so bendable. 2+ 2+ 2+ Zn Zn Zn Remember that with the ionic NaCl crystals, when a hammer pushed a block of them down, it made everything change: instead of pos ions being next to neg ions, they were next to other pos ions, and neg ions were next to neg ions, and this made the crystal break apart and be brittle. 2+ 2+ 2+ Zn Zn Zn 2+ 2+ 2+ Zn Zn Zn 2+ 2+ 2+ 2+ Zn Zn Zn
But can you see what would happen if a hammer knocked a block of metal atoms down one notch in the diagram below? Nothing would really change. The positively charged atoms that were next to positively charged atoms before would now be next to new positively charged atoms and the “sea of electrons” that holds them all together would still be holding them all together. So the blocks of atoms would simple slip by each other without ever break- ing any of the bonds between them. 2+ 2+ 2+ 2+ Zn Zn Zn Zn 2+ 2+ 2+ 2+ Zn Zn Zn Zn This is what makes metals so bendable, and this is why the piece of zinc you worked with today bent rather than shattered when you applied pressure to it. 2+ 2+ 2+ 2+ Zn Zn Zn Zn 2+ 2+ 2+ 2+ 2+ Zn Zn Zn Zn
We mentioned earlier that covalent bonds can give rise to network covalent crystals like diamond or sand that are extremely hard and have extremely high melting points. But they can also give rise to molecules like wax (C25H52) and water (H2O), C C C C C which have very low melting points. C C C C C How can a substance contain covalent bonds (the strongest bond type of them all) and have a low melting point? C C C C C C C H O H
Simple: when you melt a molecular substance like wax or ice, you are not breaking any of the covalent bonds. Instead, you are just breaking the attractions between the molecules, and they are much, much weaker than any of the bonds we have talked about so far. C C C C C C C Look at the diagram of diamond above: all it contains are strong covalent bonds, so of course it is going to be difficult to melt it. C C C C C C C C But look at the diagram of ice at right: see how there are strong covalent bonds (in black) holding the H and O atoms together as molecules. C C H H H H O O O O But there are also much weaker bonds between the molecules (in green), and those are what get broken when ice melts. H H H H
In other words, when ice melts it does this: NOT this: That should make sense, because when ice melts into water, it still has the chemical formula: H2O. H H O O H H H H O H O H H H H H O O O O H H H H
But why are some molecular solids brittle like ice, and other are soft and somewhat bendable like wax? This happens because some molecules arrange themselves into very specific crystal arrangements in the solid state. Ice, for example looks like this: Notice the solid lines here are the covalent bonds between the H and O atoms, and the dotted lines are the much weaker bonds between the molecules. But other molecules, like wax: are too bulky or stringy to form neat, precise crystal arrangements. Instead they pack together randomly, like cooked spaghetti noodles:
When solids have their particles arranged randomly, they are referred to as “amorphous” meaning they have no definite form. It should make sense then that ice would be hard and brittle because of its precise crystal arrangement, but wax would softer because of its amorphous structure. One more thing about ice: Notice the big open spaces between the molecules. This gives ice an unusually low density. Most solids are more dense than their liquids. You saw that today in the lab. But H2O is an exception: its solid is actually less dense than its liquid. That’s why ice floats in water.
So, you actually learned about five different types of solids today. The concept map below might help you put all this in perspective Weakest………………….......Strongest Amorphous molecular solids like wax (C25H52) soft and with very low melting points Ionic solids like salt brittle, high melting points (think: electrons being transferred) Metallic solids like zinc (Zn), bendable, low melting points (think: sea of delocalized electrons)) Network covalent solids like sand, very hard and very high melting points (think: electrons being shared) Crystalline molecular solids like ice(H2O) brittle and with very low melting points (think: strong covalent bonds within the molecules, but much weaker bonds between the molecules)
