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Ionic Compounds Lecture 1 - Introduction to Ions

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Ions 1Preview The lectures in this unit cover the formation of ions and ionic compounds, properties of ions and nomenclature of ionic compounds. This lecture covers the formation of atomic ions and the existence of polyatomic ions. Ions Compounds form when two or more atoms combine chemically to produce a new pure substance. There are two ways they can do this, and they both involve trying to achieve a full outer shell of valence electrons. The first method is to pick up or give away valence electrons. This allows the atom to either fill its current (neutral) valence shell or to backtrack to empty it, resulting in the previously filled shell holding the outermost electrons. When a compound gains or loses electrons, it forms an ion. I. Main Group Elements An ion is simply an atom which is carrying a charge (positive or negative). A positively charged ion is a cation. A negatively charged ion is an anion. We can calculate the charge if we know the number of extra or fewer electrons an atom has than normal. Or if we know the charge, we can also calculate the number of electrons the atom has. This was covered previously in our lecture on atoms and the periodic table. In that lecture we discussed how electrons are distributed in shells. We write their electron configurations following this general format: 1s22s22p63s23p5. If you pay attention to how those electrons are distributed in shells you find that the noble gases, which are very stable and unreactive, have full shells. That is the answer to why they are unreactive: any additional electrons must be placed in a higher energy shell and any removed would cause the current shell to be only partially full. A full shell bestows some stability upon the atom that contains it. Noble gases have a measure of natural stability because they naturally have full valence shells. Other atoms try to gain that stability in their electron configuration by gaining or losing electrons to obtain a full shell. Since all noble gases have 8 valence electrons in their outer shells we call this the Octet Rule. In order to obtain a full shell an atom must either gain or lose its valence electrons to obtain a total of 8. This gives it the full shell electronic configuration of the noble gas that comes before it or after it on the periodic table. The several exceptions to this rule are H, He, Li, and Be. For them a full shell may consist of only 2 electrons: that is the full level 1 shell with a configuration of 1s2. For example: Fluorine is in group 7 and has 7 valence electrons. Its electronic configuration looks like this: 1s22s22p5 If you count up the number of electrons in the 2 shell you get 7. The closest noble gas to fluorine is Neon, which has 8 valence electrons. Its electronic configuration looks like this: 1s22s22p6 Again, count and you find 8 valence electrons.Ions 2 If fluorine can gain 1 more electron in its valence shell it will have the same number of electrons as neon. And because electrons fall into the lowest energy level shells it will have the same electronic configuration as neon. This gives it the same relative electronic stability as neon! So fluorine naturally always forms ions with a charge of -1. We follow this argument with the octet rule across the periodic table. All atoms in group 7 have 7 valence electrons. If they want 8 they must gain one electron; so when atoms from group 7 form ions they have a charge of -1. Atoms in group 6 have 6 valence electrons. If they want 8 they must gain two electrons; so when atoms from group 6 form ions they have a charge of -2. Atoms in group 5 have 5 valence electrons. If they want 8 they must gain three electrons. So when atoms from group 5 form ions they have a charge of -3. Let’s stop for a minute and look at the other side of the periodic table. Alkali metals in the first group have 1 valence electron each. There are two ways to achieve a full shell: they can add 7 electrons, which will bring them up to a full 8; or they can remove 1 electron, which pushes them back to the previous full shell. For example, sodium has the following electronic configuration: 1s22s22p63s1. The two closest noble gases with full shells are Neon and Argon. Their configurations are respectively 1s22s22p6 and 1s22s22p63s23p6. To gain the same electronic configuration as the closest noble gases sodium may either remove its one electron in the 3s shell which leaves it with a full 2 shell; or sodium may gain 7 electrons in its level 3 shell. Removing 1 electron is much easier than adding 7. So group 1 elements naturally form +1 ions. Likewise, group 2 atoms have 2 valence electrons: to get to 8 they may either add 6 or lose 2. Again it is easier to lose 2 than gain 6. So every element in this column (group 2) naturally forms +2 ions. Following the same logic, group 3 elements make +3 ions, and group 4 atoms with 4 valence electrons are stuck in the middle. They may either gain 4 electrons or lose 4 electrons, but that’s a lot to lose or gain so usually group 4 atoms do not form ions at all. By comparing valence electrons to the octet rule we can tell from a glance at the periodic table what kind of ion a main group element forms. Ca? That is in group 2, so 2+. P? That is in group 5, so 3-. Li? That is in group 1, so 1+. By the way: this is the reason why Hydrogen is listed in group 1 on your periodic table. It shares more in common with the halogens when it comes to elemental properties; but Hydrogen forms a +1 ion more often than not. Refer to the periodic table below to see how we predict the typical charge of an ion formed from each of the main group element atoms.Ions 3 --------------------------------------------------------------------------------------------------------------------- Take a minute to look at your Ionic Compounds Worksheet and complete section I. This has you tell how many valence electrons an element has and also determine what charge ion it forms. -------------------------------------------------------------------------------------------------------------------- II. Transition Metal Ions It is not as easy to predict ionic charges for atoms in the transition block. The charges for transition metals do not uniformly follow a pattern on the periodic table like main group ions and furthermore, each of the transition metals may have more than one type of ion which forms! For example, some transition


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