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Name: Periodic Trends Reading Periodic Trends The arrangement of the periodic table reveals trends in the properties of the elements. A TREND is a predictable change in a particular direction. For example, thpere is a trend in the reactivity of the alkali metals as you move down Group 1. Each of the alkali metals react with water. However, the reactivity of the alkali metals varies. At the top of Group 1, lithium is the least reactive, sodium is more reactive, and potassium is still more reactive. In other words, there is a trend toward great reactivity as you move down the alkali metals in Group 1. Understanding a trend among the elements allows you to make predictions about the chemical behaviour of the elements. These trends in properties of the elements in a group or period can be explained in terms of electron organization. These trends can be predicted using the periodic table and can be explained and understood by analyzing the electron configurations of the elements. Elements tend to gain or lose valence electrons to achieve stable octet formation. Stable octets are seen in the inert gases, or noble gases, of Group 18 of the periodic table. In addition to this activity, there are two other important trends. First, electrons are added one at a time moving from left to right across a period. As this happens, the electrons of the outermost shell experience increasingly strong nuclear attraction, so the electrons become closer to the nucleus and more tightly bound to it. Second, moving down a column in the periodic table, the outermost electrons become less tightly bound to the nucleus. This happens because the number of filled principal energy levels increases downward within each group. These trends explain the periodicity observed in the elemental properties of atomic radius, reactivity, ionization energy, electron affinity, and electronegativity. Atomic Radius The exact size of an atom is hard to determine. An atom’s size depends on the volume occupied by the electrons around the nucleus, and the electrons do not move in well-defined paths. Rather, the volume the electrons occupy is though of as an electron cloud, with no clear-cut edges. In addition, the physical and chemical state of an atom can change the size of an electron cloud. Instead, to measure an atom’s size, we measure the atomic radius. This is the length that is half the distance between the nuclei (plural of nucleus) of two bonded atoms. Generally, the atomic radius decreases across a period from left to right and increases down a given group. The atoms with the largest atomic radii are located in Group I and at the bottom of groups. Moving from left to right across a period, electrons are added one at a time to the outer energy shell. Since the number of protons is also increasing, the nuclear charge (total charge of all the protons in the nucleus) increases across a period. This causes the atomic radius to decrease. Moving down a group in the periodic table, the number of electrons and filled electron shells increases, but the number of valence electrons remains the same. The outermost electrons in a group are farther from the nucleus as the number of filled energy shells increases. Therefore, the atomic radii increase. Atomic Radius 1. Define atomic radius 2. How do atomic radii vary as you go from left to right and from top to bottom among the elements in a periodic table? Why? 3. Of carbon, oxygen, and lithium, which has the smallest atomic radius? Why? 4. Of the elements gallium, potassium, and bromine, which one has the largest atomic radius? Name: Periodic Trends Reading 5. Which atom in the pair, sulfur and tellurium, has the smallest atom radius? Why? 6. Which atom in the pair magnesium and barium has the larger atomic radius? 7. The elements with the largest atomic radii are found in which corner of periodic table A. lower left-hand corner C. lower right-hand corner B. upper right-hand corner D. upper left-hand corner 8. Given the representation of a chlorine atom, which circle might represent an atom of Sulphur? 9. Using the above diagram of a chlorine atom, which circle might represent an atom of fluorine? 10. Explain why nitrogen is a smaller atom than both boron and phosphorous Reactivity Reactivity refers to how likely or vigorously an atom is to react with other substances. This is usually determined by how easily electrons can be removed (ionization energy) and how badly they want to take other atom’s electrons (electron affinity) because it is the transfer/interaction of electrons that is the basis of chemical reactions. Metals and non-metals each have their own trends. When we look at the metals, we see that the reactivity increases as you move down the group. However, as we move across a period, the reactivity decreases as you move from left to right. The farther to the left and down the periodic table that you go, the easier it is for electrons to be given or taken away, resulting in higher reactivity. This is a direct result of the nuclear charge and the attraction to the outermost electrons. The non-metals, however, show different trends. As you move from left to right in the period, we see that the reactivity increases. On the other hand, as you move down a group in the nonmetals, the reactivity decreases. The farther right and up you go on the periodic table, the higher the electron affinity, resulting in a more vigorous exchange of electrons. Where metals will easily give up an electron to gain stability, non-metals want to gain electrons. Reactivity 1. Define reactivity 2. What are the distinguishing properties of metals and non-metals? Give three example of each. 3. In reactions, do alkaline earth metals gain or lose electrons (how many)? Explain. 4. Explain the group trend for reactivity of metals? 5. The halogen group wants to gain one electron, and the alkali metal group wants to lose one electron. Explain why this is. 6. Why is it easier for oxygen to gain electrons than it is for selenium? Name: Periodic Trends Reading Ionization Energy When atoms have equal amount of protons and electrons, they are electrically neutral. But when enough energy is added, individual electrons can be removed from an atom. The neutral atom then becomes a positively charged ion. The ionization energy, or ionization potential, is the energy required to completely remove an electron from a gaseous atom or ion. The closer and more tightly bound an electron is to the nucleus, the more difficult it will be to remove, and the higher its ionization energy will be. The first ionization energy is the energy required to remove one electron from the parent atom. The second ionization energy is the energy required to remove a second valence electron from the +1 ion to form the +2 ion, and so on. Successive ionization energies increase. The second ionization energy is always greater than the first ionization energy. Ionization energy tends to decrease down a group. Each element has more occupied energy levels than the one above it does. Therefore, the outermost electrons are furthest from the nucleus near the bottom of a group. Similarly, as you move down a group, each successive element contains more electrons in the energy levels between the nucleus and the outermost electrons. These inner electrons cause the outermost electrons to be held less tightly to the nucleus. Ionization energy, however, tends to increase as you move from left to right across a period. From one element to the next in a period, the number of protons and electrons increase by one each. The additional proton increases the total charge of the nucleus (nuclear charge). The additional electron is added to the same outer energy level as the previous element in the period. However, a higher nuclear charge attracts the outer electrons more strongly. Therefore, they are held onto “tighter” and need more energy to be removed. Ionization Energy 7. Define ionization energy. 8. How does ionization energy vary as you go from left to right across the periodic table? 9. How does ionization energy vary as you go from top to bottom on the periodic table? 10. Of the elements calcium, beryllium or magnesium, which has the highest ionization energy? Why? 11. Which atom in each of the following pairs has higher ionization energy? a. S or P b. Rb or Sr c. Mg or Al d. O or S e. Mg or Ba Name: Periodic Trends Reading Electron Affinity Electron affinity reflects the ability of an atom to accept an electron. It is the energy change that occurs when an electron is added to a gaseous atom. Atoms with stronger effective nuclear charge have greater electron affinity. Some generalizations can be made about the electron affinities of certain groups in the periodic table. The Group 2 elements, the alkaline earths, have low electron affinity values. These elements are relatively stable because they have filled s subshells. Group 17 elements, the halogens, have high electron affinities because the addition of an electron to an atom results in a completely filled shell. Group 18 elements, noble gases, have electron affinities near zero, since each atom possesses a stable octet and will not accept an electron readily. Elements of other groups have low electron affinities. Electron Affinity 1. What is meant by electron affinity? 2. How does electron affinity vary as you go from left to right across the periodic table? 3. How does electron affinity vary as you go down a group in the periodic table? 4. Which element should have the higher electron affinity, selenium or tellurium? Why? 5. Which element should have the higher electron affinity, arsenic or antimony? Why? Electronegativity Electronegativity is a measure of the attraction of an atom for the electrons in a chemical bond. You can think of it as how evenly an electron will be shared between the atoms. The more electronegative an atom is, the more it will want to hang on to the electron (it will “hog” the electron). The higher the electronegativity of an atom, the greater its attraction for bonding electrons. Electronegativity is related to ionization energy. Electrons with low ionization energies have low electronegativites because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted on electrons by the nucleus. In a group, the electronegativity decreases as atomic number increases, as a result of increased distance between the valence electron and nucleus (greater atomic radius). An example of an electropositive (i.e., low electronegativity) element is cesium; an example of a highly electronegative element is fluorine. Electronegativity 6. Define electronegativity. 7. How does electronegativity vary as you go from left to right across the periodic table? 8. How does electronegativity vary as you go from top to bottom on the periodic table? 9. Which element is the most electronegative? Why? 10. Which element is the least electronegative (that means the most electropositive)? Why? p