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Accelerated Chemistry Study Guide the Periodic Table, Chapter 5

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Accelerated Chemistry Study Guide the Periodic Table, Chapter 5 Accelerated Chemistry Study Guide The Periodic Table, Chapter 5 Terms, definitions, and people Dobereiner Newlands Mendeleev Moseley Periodic table Periodic Law group family period Page 1 of 38 alkali metal alkaline earth metal halogen noble gas metal nonmetal semimetal, metalloid s-, p-, d-, and f-blocks valence electrons (s and p) isoelectronic abbreviated electron configuration Page 2 of 38 periodic trend atomic radius ionic radius ionization energy, successive ionization energies electron affinity electronegativity Specific topics to know periodic trends within a period periodic trends within a family Page 3 of 38 The World of Chemistry Episode 7 - The Periodic Table 1. What two types of properties are described in the video? 2. What are some examples of physical properties? 3. How many elements are on the modern periodic table? How many of these can be found in nature? 4. Why do the symbols for some elements (such as iron) seem to have no relationship to their name? 5. What is meant by the atomic and mass number of an element? 6. Elements in the periodic table are arranged by increasing __________ number. 7. What is a group of elements? a period? 8. What are the ingredients used in the making of glass? What determines the color of glass? 9. What are alkali metals? Describe their reaction with water. 10. How does the size of an atom change - a. as you go down a group of elements? b. as you go from left to right in a period of elements? 11. Who developed the periodic table? 12. What did Mendeleev do for elements that had not yet been discovered? 13. How did Glenn Seaborg change the periodic table? 14. Why are the electrons in the outer shell of an atom important? Page 4 of 38 The Periodic Table Elements in the periodic table are arranged in order of increasing atomic number. Atomic number indicates the number of protons, and if the element is neutral the number of protons is equal to the number of electrons. The periodic table is arranged in four blocks: s, p, d, and f that group together elements with similar electron configurations. The horizontal rows are called periods. The vertical columns are called groups or families. Elements in the same column are chemically similar to each other because each element in the family has the same number of valence electrons. The number of valence electrons (outer shell electrons) is equal to the number of s and p electrons in the highest principal energy level. The Periodic Table consists of nonmetals, metalloids and metals, and elements become more “metallic” as you move from the right to the left. Periodic Trends. Make sure that in each case you can describe the trend (up/down and left/right) and also explain why the trend occurs. Atomic Size—the distance from the nucleus of an atom to the outermost electrons. It depends entirely on the outer shell electrons. Notice that though the nucleus makes up most of the mass of the atom, it takes up very little space. In a typical atom, the nucleus would be about the size of a marble on the 50-yard line of a football stadium with the electrons orbiting like crazed fans in the upper stands. Vertical trend: Atomic size increases moving down the periodic table Explanation: Atomic size increases going down because the principal energy level increases (2p, 3p, 4p for example). This means that the outer shell electrons have more energy and can move farther from the nucleus. Horizontal trend: Atomic size decreases when moving left to right. Explanation: As you move left to right along a period, the atomic size decreases because outer shell electrons (valence electrons) are added to the same principal energy level or lower (ex: 4s then 3d then 4p in period 4) while the number of protons increases. In other words, as the atomic number increases, the number of protons increases. Remember that protons attract electrons because protons are positive and electrons are negative. The fact that there are more protons causes the electrons to be drawn in towards the nucleus. 1 of 3 Page 5 of 38 Ionic size -- the distance from the nucleus of an ion to the outermost electrons. The trends in atomic radius are different from the trends in ionic radius. It is important to keep in mind that the number of electrons is different for ions than it is for neutral elements. It is possible to predict how much bigger or smaller an ion will be by comparing its relative number of protons and electrons and comparing that number to a neutral element with the same number of protons or electrons. Ionization energy—the energy required to remove an electron. In equation form, this looks like: X X+ + e- Notice that this equation looks reasonable for metals: Na Na+ + e- Mg Mg+ + e- Sn Sn+ + e- And should look unreasonable for nonmetals: Cl Cl+ + e- O O+ + e- Vertical trend: ionization energy decreases moving down the periodic table Explanation: as with ionic radius, electrons are in higher principal energy levels moving down the PT. This means that they are of higher energy and farther from the nucleus. This makes them easier to remove. There is another factor at work here: inner shell electrons (the ones that aren’t the outer shell or valence electrons) get in the way of the protons’ ability to hang on to the electrons. This is called the shielding effect. Horizontal trend: ionization energy increases going left to right () on the PT Explanation: this occurs for the same reason as the horizontal trend in atomic size. More protons, same principal energy level—electrons are held more strongly and harder to remove. Electronegativity: the tendency for an atom to attract electrons to itself when forming a bond with another element. The vertical and horizontal trends for electronegativity (and their explanations) are the SAME as those for ionization energy. There is an exception: fluorine, not helium, has the highest electronegativity on the periodic table and similarly the halogens are the group with the highest electronegativity, not the noble gases. This is because the noble gases have a full outer shell of electrons and do not tend to attract additional electron density. The halogens, on the other hand strongly attract electrons in order to fill their outer shell (remember that they have seven valence electrons, but would be most stable with eight) 2 of 3 Page 6 of 38 Electron affinity: the energy change that occurs when an atom gains an electron. This is how electron affinity can be expressed as an equation: Ne (g) + e- Ne- electron affinity = 29 kJ/mol F (g) + e- F- electron affinity = -328 kJ/mol Notice that the reaction is highly exothermic (releases energy, which you know because the heat change is negative) in the case of fluorine and endothermic in the case of neon. This makes sense because fluorine is made much more stable when it gains an electron, so it releases heat as it becomes more stable. The gain of an electron is favored for fluorine. In some ways, electron affinity can be viewed as the opposite of ionization energy, but there are some ways in which they are not completely opposite. For instance, the noble gases have a high ionization energy but low and positive electron affinity 3 of 3 Page 7 of 38 Page 8 of 38 Regions of the Periodic Table: families, states of matter Regions of the Periodic Table: s-, p-, d-, f-blocks Page 9 of 38 Regions of the Periodic Table: metals, nonmetals, metalloids Valence electrons Page 10 of 38 Periodic trends: atomic size, ionization energy Periodic trends: electron affinity Page 11 of 38 Periodic trends: electronegativity Periodic trends: summary Page 12 of 38 Page 13 of 38 Page 14 of 38 Page 15 of 38 Page 16 of 38 Page 17 of 38 Page 18 of 38 ChemQuest 16 Information: Shielding FIGURE 1: “Bohr Diagrams” of boron, carbon and nitrogen nucleus charge = +5 nucleus charge = +6 nucleus charge = +7 Boron Carbon Nitrogen Because the nucleus is positively charged, it exerts an attractive force on the electrons. However, the three electrons in boron’s outer energy level do not feel the full +5 attraction from the 5 protons in boron’s nucleus. Before the +5 attraction gets to the outer energy level it gets partially cancelled (or “shielded”) by the two electrons in the first energy level. The two electrons in the first energy level weaken the attractive force by two. Therefore to the outer energy level it only “feels” like a +3 charge rather than a +5 charge from the nucleus. Consider the diagram of carbon. An electron in the outer energy level only “feels” a charge of +4 coming from the nucleus because the two electrons in the first energy level shield two of the positive charges from the nucleus. Critical Thinking Questions 1. How large is the charge that the second energy level of nitrogen “feels” from the nucleus? 2. Why does the first energy level in each of the three above diagrams only contain two electrons? 3. How many electrons can fit in the second energy level of any atom? 4. How many electrons can fit in the third energy level? Page 19 of 38 5. How many energy levels does aluminum have? How many electrons should be in each energy level? 6. Draw a Bohr diagram for aluminum similar to those above. 7. Explain why the second energy level of aluminum only feels a +11 attraction instead of a +13 attraction from aluminum’s electrons.
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