Why is nacl brittle

Look again at the last diagram:. Now imagine what would happen if you replaced the cesium ion with the smaller sodium ion. Sodium ions are, of course, smaller than cesium ions because they have fewer layers of electrons around them.

You still have to keep the chloride ions in contact with the sodium. The effect of this would be that the whole arrangement would shrink, bringing the chloride ions into contact with each other - and that introduces repulsion. Any gain in attractions because you have eight chlorides around the sodium rather than six is more than countered by the new repulsions between the chloride ions themselves.

When sodium chloride is coordinated, there are no such repulsions - and so that is the best way for it to organize itself. Which structure a simple compound like NaCl or CsCl crystallizes in depends on the radius ratio of the positive and the negative ions. That puts it in the range where you get coordination. Sodium chloride is taken as typical of ionic compounds, and is chosen rather than, say, cesium chloride, because it is found on every syllabus at this level.

There are strong electrostatic attractions between the positive and negative ions, and it takes a lot of heat energy to overcome them. Ionic substances all have high melting and boiling points. Differences between ionic substances will depend on things like:. Brittleness is again typical of ionic substances.

Imagine what happens to the crystal if a stress is applied which shifts the ion layers slightly. Ions of the same charge are brought side-by-side and so the crystal repels itself to pieces! Many ionic solids are soluble in water - although not all. It depends on whether there are big enough attractions between the water molecules and the ions to overcome the attractions between the ions themselves. Positive ions are attracted to the lone pairs on water molecules and co-ordinate dative covalent bonds may form.

Water molecules form hydrogen bonds with negative ions. This is also typical of ionic solids. The attractions between the solvent molecules and the ions are not big enough to overcome the attractions holding the crystal together. Solid sodium chloride does not conduct electricity, because there are no electrons which are free to move.

When it melts, sodium chloride undergoes electrolysis, which involves conduction of electricity because of the movement and discharge of the ions. In the process, sodium and chlorine are produced. This is a chemical change rather than a physical process. The positive sodium ions move towards the negatively charged electrode the cathode. When they get there, each sodium ion picks up an electron from the electrode to form a sodium atom.

These float to the top of the melt as molten sodium metal. And assuming you are doing this open to the air, this immediately catches fire and burns with an orange flame. The movement of electrons from the cathode onto the sodium ions leaves spaces on the cathode.

The power source the battery or whatever moves electrons along the wire in the external circuit to fill those spaces. That flow of electrons would be seen as an electric current the external circuit is all the rest of the circuit apart from the molten sodium chloride.

Meanwhile, chloride ions are attracted to the positive electrode the anode. When they get there, each chloride ion loses an electron to the anode to form an atom. These then pair up to make chlorine molecules. Chlorine gas is produced. Overall, the change is. The new electrons deposited on the anode are pumped off around the external circuit by the power source, eventually ending up on the cathode where they will be transferred to sodium ions.

Molten sodium chloride conducts electricity because of the movement of the ions in the melt, and the discharge of the ions at the electrodes. Both of these have to happen if you are to get electrons flowing in the external circuit. If we added another layer of caesium ions, you could similarly work out that each chloride ion was touching eight caesium ions.

The chloride ions are also 8-co-ordinated. Overall, then, caesium chloride is co-ordinated. The final diagram in this sequence takes a slightly tilted view of the structure so that you can see how the layers build up. These diagrams are quite difficult to draw without it looking as if ions of the same charge are touching each other.

They aren't! Diagrams of ionic crystals are usually simplified to show the most basic unit of the repeating pattern. For caesium chloride, you could, for example, draw a simple diagram showing the arrangement of the chloride ions around each caesium ion:.

By reversing the colours green chloride ion in the centre, and orange caesium ions surrounding it , you would have an exactly equivalent diagram for the arrangement of caesium ions around each chloride ion. Note: These diagrams are difficult enough to draw convincingly on a computer. Trying to draw them freehand in an exam is seriously difficult. If you are doing a syllabus which wants you to know about the structure of caesium chloride, take a careful look at past exam papers and mark schemes to see exactly what sort of diagrams if any you need to use, and then practise them so that you can draw them quickly and well.

If you haven't got any past papers and mark schemes, follow this link to the syllabuses page to find out how to get them if you are doing a UK-based exam. When attractions are set up between two ions of opposite charges, energy is released. The more energy that can be released, the more stable the system becomes. That means that the more contact there is between negative and positive ions, the more stable the crystal should become.

If you can surround a positive ion like caesium with eight chloride ions rather than just six and vice versa for the chloride ions , then you should have a more stable crystal. So why doesn't sodium chloride do the same thing? Now imagine what would happen if you replaced the caesium ion with the smaller sodium ion.

Sodium ions are, of course, smaller than caesium ions because they have fewer layers of electrons around them. You still have to keep the chloride ions in contact with the sodium. The effect of this would be that the whole arrangement would shrink, bringing the chloride ions into contact with each other - and that introduces repulsion. Any gain in attractions because you have eight chlorides around the sodium rather than six is more than countered by the new repulsions between the chloride ions themselves.

When sodium chloride is co-ordinated, there are no such repulsions - and so that is the best way for it to organise itself. Which structure a simple compound like NaCl or CsCl crystallises in depends on the radius ratio of the positive and the negative ions. That puts it in the range where you get co-ordination. At this point the negative ions will touch each other again even with co-ordination.

A new arrangement known as co-ordination then becomes necessary. This is beyond any syllabus that I am currently tracking.

Sodium chloride is taken as typical of ionic compounds, and is chosen rather than, say, caesium chloride, because it is found on every syllabus at this level.

There are strong electrostatic attractions between the positive and negative ions, and it takes a lot of heat energy to overcome them. Ionic substances all have high melting and boiling points. Differences between ionic substances will depend on things like:. Magnesium oxide has exactly the same structure as sodium chloride, but a much higher melting and boiling point.

If the ions are smaller they get closer together and so the electrostatic attractions are greater. Rubidium iodide, for example, melts and boils at slightly lower temperatures than sodium chloride, because both rubidium and iodide ions are bigger than sodium and chloride ions.

The attractions are less between the bigger ions and so less heat energy is needed to separate them. Brittleness is again typical of ionic substances. Imagine what happens to the crystal if a stress is applied which shifts the ion layers slightly. Many ionic solids are soluble in water - although not all. It depends on whether there are big enough attractions between the water molecules and the ions to overcome the attractions between the ions themselves.

Positive ions are attracted to the lone pairs on water molecules and co-ordinate dative covalent bonds may form. Water molecules form hydrogen bonds with negative ions. Note: The bonding in hydrated metal ions is covered in the page on co-ordinate bonding. The bonding between negative ions like chloride ions and water molecules is covered in the page on hydrogen bonding.

This is also typical of ionic solids. If it does leave a scratch, proceed with some additional tests because some synthetic diamonds will also scratch glass. Get a glass full of water and simply drop your diamond into the glass. If the diamond is real, it will drop to the bottom of the glass due to the high density of the stone. A fake diamond will have rainbow colors that you can see inside the diamond.

The best way to tell a cubic zirconia from a diamond is to look at the stones under natural light: a diamond gives off more white light brilliance while a cubic zirconia gives off a noticeable rainbow of colored light excessive light dispersion. Heat the stone over a lighter for 30 seconds and immediately drop it into a glass of water. A real diamond will be unaffected by dramatic changes in temperature and it will simply fall to the bottom of the glass.

Fake diamonds however will shatter immediately. Always attempt to find the owner if possible, or turn the item in to the police. Most states will allow finders to keep the property if the owner does not show up to claim it after a certain time. A flawless raw diamond may be worth more than a cut diamond with a low clarity grade. Color: Most colorless or white diamonds have natural yellow or brown tints of color in them. Kimberlite, also called blue ground, a dark-coloured, heavy, often altered and brecciated fragmented , intrusive igneous rock that contains diamonds in its rock matrix.

It has a porphyritic texture, with large, often rounded crystals phenocrysts surrounded by a fine-grained matrix groundmass.