Electron Configurations and Periodic Properties

Electron Configurations and Periodic Properties Name:___________________ Guided Inquiry Period:____ Objective: You have learned how the atoms are arranged and how to determine the electron configurations for atoms. Now we need to look at the relationship between electron configurations and periodic trends. Please put the answers into your own words. Atomic Radius 1. Define atomic radius: Atomic radius is the distance from the nucleus of an atom to the outermost electron shell or orbital of the atom 2. What is the trend for atomic radius as you go from left to right on the periodic table? The trend for atomic radius as you go from left to right on the periodic table generally decreases. 3. What causes the trend for atomic radius as you go from left to right on the periodic table? The trend for atomic radius as you go from left to right on the periodic table is caused by increasing nuclear charge. As you move from left to right across a period, the number of protons in the nucleus increases, leading to a greater attractive force pulling the electrons closer to the nucleus. This results in a decrease in atomic radius. 4. What is the trend for atomic radius as you go from top to bottom on the periodic table? The trend for atomic radius as you go from top to bottom on the periodic table generally increases. 5. What causes the trend for atomic radius as you go from top to bottom on the periodic table? The trend for atomic radius as you go from top to bottom on the periodic table is caused by the addition of energy levels or electron shells. Each successive row or period in the periodic table represents the addition of a new electron shell, further away from the nucleus. With increasing energy levels, the outermost electrons are farther from the nucleus, resulting in an increase in atomic radius. Additionally, the shielding effect of inner electron shells reduces the effective nuclear charge experienced by the outermost electrons, allowing them to spread out more, further contributing to the increase in atomic radius. Ionization Energy 1. Define ionization energy: Ionization energy is the energy required to remove an electron from an atom or ion in its gaseous state. 2. What is the trend for ionization energy as you go from left to right on the periodic table? The trend for ionization energy as you go from left to right on the periodic table generally increases. 3. What causes the trend for ionization energy as you go from left to right on the periodic table? The trend for ionization energy as you go from left to right on the periodic table is caused by increasing effective nuclear charge. As you move from left to right across a period, the number of protons in the nucleus increases, leading to a greater attractive force pulling the electrons closer to the nucleus. This makes it more difficult to remove an electron, resulting in higher ionization energies. 4. What is the trend for ionization energy as you go from top to bottom on the periodic table? The trend for ionization energy as you go from top to bottom on the periodic table generally decreases. 5. What causes the trend for ionization energy as you go from top to bottom on the periodic table? The trend for ionization energy as you go from top to bottom on the periodic table is caused by increasing atomic size. As you move down a group or family in the periodic table, the number of electron shells or energy levels increases, leading to greater distance between the outermost electrons and the nucleus. This reduces the attractive force between the nucleus and the outermost electrons, making it easier to remove an electron and resulting in lower ionization energies. Additionally, the shielding effect of inner electron shells further reduces the effective nuclear charge experienced by the outermost electrons, contributing to the decrease in ionization energy. Electron Configurations and Periodic Properties Guided Inquiry Ionic Radius Name:___________________ Period:____ 1. Define cation: A cation is a positively charged ion formed when an atom loses one or more electrons. 2. Define anion: An anion is a negatively charged ion formed when an atom gains one or more electrons. 3. What types of ions do metals tend to form? What happens to the radius of this type of ion compared to its atom? Metals tend to form cations by losing electrons to achieve a stable electron configuration, typically by reaching a filled valence shell. When a metal atom loses electrons to form a cation, its radius decreases compared to the neutral atom. This is because the loss of electrons reduces the electron-electron repulsion, allowing the remaining electrons to be pulled closer to the nucleus, thereby reducing the atomic radius. 4. What types of ions do nonmetals tend to form? What happens to the radius of this type of ion compared to its atom? Nonmetals tend to form anions by gaining electrons to achieve a stable electron configuration, typically by filling or completing their valence shell. When a nonmetal atom gains electrons to form an anion, its radius increases compared to the neutral atom. This is because the addition of electrons increases the electron-electron repulsion, causing the outermost electron shell to expand further from the nucleus, thereby increasing the atomic radius. Electronegativity 1. Define electronegativity: Electronegativity is the tendency of an atom to attract electrons towards itself in a chemical bond. 2. What is the trend for electronegativity as you go from left to right on the periodic table? The trend for electronegativity as you go from left to right on the periodic table generally increases. 3. What causes the trend for electronegativity as you go from left to right on the periodic table? The trend for electronegativity as you go from left to right on the periodic table is caused by increasing effective nuclear charge. As you move from left to right across a period, the number of protons in the nucleus increases, leading to a greater attractive force pulling electrons towards the nucleus. This results in higher electronegativity values because atoms become more efficient at attracting electrons. 4. What is the trend for electronegativity as you go from top to bottom on the periodic table? The trend for electronegativity as you go from top to bottom on the periodic table generally decreases. 5. What causes the trend for electronegativity as you go from top to bottom on the periodic table? The trend for electronegativity as you go from top to bottom on the periodic table is caused by increasing atomic size. As you move down a group or family in the periodic table, the number of electron shells or energy levels increases, leading to greater distance between the outermost electrons and the nucleus. This reduces the effective nuclear charge experienced by the outermost electrons, making it less likely for atoms to attract additional electrons strongly, resulting in lower electronegativity values. Practice Problems 1. Place these elements in order of increasing atomic radius: Cs, Bi, Rn, Au, Hf Hf < Au < Bi < Cs < Rn 2. Place these elements in order of decreasing ionization energy: Li, O, C, K, Ne, F F < Ne < O < C < Li < K 3. Place these elements in order of increasing electronegativity: Sr, Cl, Ra, In, As Ra < Sr < As < In < Cl