AP Chemistry Unit 1- Atomic Structures and Properties

Unit 1 - Atomic Structures and Properties Photoelectron Spectroscopy ● Each peak is a subshell ● Determine relative height for number of electrons first peak is 1s² Cations are smaller than neutral atom, and anions are larger than neutral atom because additional electrons mean more repulsive for Here’s the reason why these trends happen: ● Atomic radius: From left to right on the periodic table the atomic radius decreases. Although we add protons across the periodic table (protons also considered nuclear charge), the size of the atom will decrease due to effective nuclear charge. Effective nuclear charge is a pull toward the charged proton. Since the effective nuclear charge increases, the valence electrons are drawn closer to the nucleus, decreasing the size of an atom. Similarly, atomic radius increases as we go down the table. This is because there are more sublevels added to the atom and due to shielding effect. Shielding effect is when the electron and the nucleus in an atom have a decrease in attraction because the electrons are further away due to increased sublevels. o For example, Na is bigger than Cl because Na has less effective nuclear charge! ● Ionic radius: Cations (positively charged ions) are generally smaller than the atoms that they came from because they have lost valence electrons. The amount of protons always stays the same so this means that since all of those protons are pulling in on the fewer electrons that are there (effective nuclear charge) the atomic radius is significantly smaller. In anions we are adding electrons. There are more electrons on the outside energy levels. Since we have the same number of protons and more electrons, the protons cannot fill as effectively so the radius gets larger. ● Ionization energy: Ionization energy is the energy needed to remove an electron. This is strongly associated with atomic radius. Ionization energy increases from left to right because they have a greater effective nuclear charge. Ionization energy decreases as you go down the table because it’s easier to remove the electrons due to increased shielding effect. o However, there are exceptions: Li has greater Ionization Energy than Be because Li’s core electron is being removed. ● Electronegativity: the ability of an atom to attract electrons toward itself. Electronegativity is a scale. For the AP test, F has the highest electronegativity of 4.0. Electronegativity increases across the period because of increased effective nuclear charge (want more electrons). On the other hand, electronegativity decreases as you go down the table. ● Electron affinity: the opposite of electronegativity. It is the energy change when an atom becomes negatively charged. Electron affinity is always negative. There isn’t a solid trend for electron affinity compared to others. However, this is the general trend: electron affinity becomes more negative from left to right. Electron affinity becomes more positive when you go down the table. Aufbau Principle: electrons fill the lowest energy orbitals (not levels) first. Use this diagram to know which orbitals to fill first. Start from 1s, 2s, then 2p, and so on. With this principle, it helps us to write out an electron configuration: For example, Oxygen’s configuration would be 1s2 2s2 2p4 This diagram tells us that 2 electrons can be filled in the s orbitals, 6 electrons in the p orbitals, 10 electrons in the d orbitals, and 14 electrons in the f orbitals. There’s a shorthand version for writing electron configuration. Let’s write an electron configuration for different elements: Oxygen: [He] 2s2 2p4 Potassium: [Ar]4s1 When you are writing a shorthand version, you can only include a noble gas in the bracket. For ions, you have to remove electrons that are outside leaves first. For example, an ion electron configuration for Mg2+ would be [Ne] 3s1 Let’s do another example: Ion electron configuration for Ag: [Kr]4d105s1. The 5s orbital lost an electron because it’s a more outer sublevel that 4d. You can check if your electron configuration is right by counting the numbers in the exponent. Remember that the electron configuration for Oxygen is [He] 2s2 2p4? In order to check if this configuration is right, you can add 2 and 4, which would give you 6, and this is correct because there are 6 valence electrons for Oxygen. Hund's rule states that each sublevel should have one electron before any others are doubled up (for example, all of the 2p subshells should have at least one arrow in it before moving onto another subshell). Pauli Exclusion Principle states that no two electrons in the same atom can have identical values for all four of their quantum numbers.