Chapter 3 Atoms and Ionic Bonds LEARNING OBJECTIVES SUMMARIES 1. Periodic trends: Know how each of the following are affected by the size of the valence shell (n), the nuclear charge (# of protons), and ionization (# of electrons) a. Ionization energy Ionization is the removal of an electron from a nucleus. It takes energy to do this, since the negatively charged electron is attracted to the positively charged nucleus. There are two major factors that influence how much energy this takes. Factor #1 the size of the valence shell (n) If an electron is closer to the nucleus, it s going to be harder to remove (higher ionization energy). So when the valence shell of the outermost electron decreases, then the ionization energy increases. Factor #2 nuclear charge If a nucleus has more positive charge, then it will attract the negatively charged electrons more. So when the number of protons increase across a period (row) of the periodic table, in general the ionization energy increases. b. Electron affinity Electron affinity is the amount of energy gained by adding an electron to an atom. Since this is a favorable process, it s sometimes reported as a negative number, but what matters is the magnitude of the number. If the attraction between the nucleus and the electron is greater, than the electron affinity will increase in magnitude.
The same two factors that affect ionization energy affect electron affinity. Smaller valence shell size (Quantum number n) and larger nuclear charge both mean a greater attraction between the electron being added and the nucleus, resulting in a larger electron affinity. However, you must remember that the noble gases already have a full valence shell, so the electron being added has to go into the NEXT shell. c. Atomic/ionic radii The dependence of the radius of an atom based on valence shell size is simple: smaller valence shell size means a smaller radius. For nuclear charge, an increase in the number of protons causes the nucleus to hug its electrons a bit tighter, decreasing the radius. 2. Describe the difference between an ionic bond and a covalent bond Covalent bonding involves the sharing of electrons, typically between non-metals, while ionic bonding involves the transfer of electrons. Ionic bonds usually involve a metal and a non-metal. Since metals do not hold their electrons very tightly (low ionization energy) and non-metals do (high electron affinity), metals can lose their electrons to form a positively charged cation, while the non-metal that gains the electron forms a negatively charged anion.
3. For ionic compounds, translate between a chemical formula and the chemical name An ionic compound is a neutral compound in which the total number of positive charges must equal the total number of negative charges. Take calcium chloride (CaCl 2 ), for example. The calcium has a charge of +2, and each chlorine has a charge of -1. So there are two chlorines per calcium to make an overall neutral compound. Here are some other simple examples. Total charge = 0 = +2 + 2(-1) To translate from the chemical formula to the name, remember that the cation is always named before the anion, and the suffix ide is added to the anion. So CaCl 2 is calcium chloride, and MgO is magnesium oxide. Determining the charge of an ion The charge on a metal or a non-metal depends on its group number. Many transition metals can have multiple possible charges as shown below, and the charges are always specified in roman numerals by the name. For example, CuCl 2 = copper (II) chloride CuCl = copper (I) chloride. Common charges of transition metal cations
4. Know the name and formula of common polyatomic ions You need to memorize the name, formula, and charge of each of the following polyatomic ions. Make notecards. 5. Understand the origin and application of the octet rule for main group elements The octet rule Main group elements tend to undergo reactions that leave them with eight outer-shell electrons. The octet rule was applied above when naming ionic compounds. For example, that is why alkali metals lose a single electron to form a cation with a +1 charge, and why halogens gain a single electron to form an anion with a -1 charge. The octet rule results from the influences affecting ionization energies and electron affinities already discussed. Metals, which occur towards the left of the periodic table, have relatively low nuclear charges for their given valence shell size. So it s easy to remove electrons until they form a noble gas electron configuration. The next electron removed would have to be one in a lower valence shell, and therefore closer to the nucleus and much harder to remove. Non-metals, which occur towards the right of the periodic table, have relatively high nuclear charges for their given valence shell size. So it s favorable to gain electrons until they form a noble gas electron configuration. The next electron gained would have to occupy a higher valence shell, and therefore further from the nucleus with a much lower electron affinity.
6. Learn the structure of the periodic table, define groups and periods, and identify the alkali metal, alkali earth metals, halogens, main-group elements, transition metals, and the locations of metals, non-metals, and semi-metals Group = column Period = row Semi-metals = boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb), tellurium (Te), and astatine (At)