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Chapter 12 Electrons in Atoms Introduction The view of the atom as a positively charged nucleus (protons and neutrons) surrounded by electrons is useful for visualizing the basic structure of an atom However, this model of the atom explains only a few simple properties of atoms. This chapter will include models of atomic structure with an emphasis on the electrons in atoms 12.1 The Development of Atomic Models For about 50 years after the time of Dalton (1766-1844) the atom was considered to be a solid indestructible mass The discovery of subatomic particles shattered every theory that included indestructible atoms 12.1 Thompson Model After discovering the electron, Thomson described a “plum pudding” model of the atom in which an atom was a ball of positive charge containing electrons. 12.1 Rutherford Model After learning of the nucleus, Rutherford proposed the nuclear atom in which electrons surround a dense nucleus most of the atom being empty space. Later experiments showed that the nucleus is composed of protons and neutrons 12.1 Bohr Model In 1913, Niels Bohr proposed the “planetary model” of the atom in which the electrons move in orbits around the nucleus. 12.1 Bohr Model Bohr proposed that electrons in a particular path have a fixed energy that keeps them from falling into the nucleus. energy level – the region around the nucleus where the electron is likely to be moving Ladder Analogy 12.1 Bohr Model Electrons can jump from one energy to another by gaining or losing just the right amount of energy Electrons cannot exist between energy levels A quantum of energy is the amount of energy required to move an electron to the next highest energy level. The higher the energy level the farther the electron is from the nucleus (usually) Energy levels are more closely spaced further from the nucleus The higher the energy level the easier it is for the electron to escape 12.2 Quantum Mechanical Model Quantum mechanical model states that the atom has no definite shape and that electrons do not have precise orbits 12.2 Quantum Mechanical Model In 1926, Erwin Schrödinger used the new quantum theory to write and solve a mathematical equation to describe the location and energy of an electron in a hydrogen atom The modern description for electrons in the atom, the quantum mechanical model, comes from the mathematical solution to Schrödinger’s equation. Primarily mathematical – has few analogies in the visible world 12.2 Quantum Mechanical Model Features of the quantum mechanical model: restricts the energy of electrons to certain values but does not define the exact path taken by the electron estimates the probability of finding the electron within a given region called a cloud (electron cloud) the electron can be found within this cloud 90% of the time 12.3 Atomic Orbitals The quantum mechanical model energy levels with principle quantum numbers (n) Each principle quantum number refers to a major (principle) energy level in an atom n=1, 2, 3, 4, etc. The average distance of the electron from the nucleus increases with increasing values of n. 12.3 Atomic Orbitals Within each principle energy level, the electrons occupy energy sublevels The number of energy sublevels is the same as the principle quantum number 12.3 Atomic Orbitals The quantum mechanical model describes the position of an electron with cloud shapes (based on probability) These cloud shapes are called atomic orbitals Atomic orbitals are represented with letters: s, p, d, f, g Regions where there is a low probability to find electrons are called nodes Atomic orbitals have different characteristic shapes s orbitals are spherical p orbitals are dumbell shaped d and f orbitals are more complex and harder to visualize 12.3 Atomic Orbitals 12.3 Atomic Orbitals Table 12.1 Summary of Principle Energy Levels, Sublevels, and Orbitals Principle Energy Level # of Sublevels Type of Sublevel n =1 1 1s (1 orbital) n=2 2 2s (1 orbital), 2p (3 orbitals) n=3 3 3s (1 orbital), 3p (3 orbitals), 3d (5 orbitals) n=4 4 4s (1 orbital), 4p (3 orbitals), 4d (5 orbitals), 4f (7orbitals) 12.3 Atomic Orbitals Each orbital can hold a maximum of 2 electrons The maximum number of electrons that can occupy a principle energy level is given by the formula 2n2 where n = principle quantum number Example: How many electrons can be found in the 3rd principle energy level? A: 2n2 = (2)(32) = 18 electrons 12.4 Electron Configurations In all natural phenomena, change proceeds toward the lowest possible energy state. High-energy systems are unstable Unstable systems lose energy to become more stable In the atom, electrons and the nucleus interact to make the most stable arrangement possible. 12.4 Electron Configurations The ways in which electrons are arranged around the nuclei of atoms are called electron configurations. Three rules or principles are used to determine the electron configuration of atoms: 1. 2. 3. Aufbau Principle Pauli Exclusion Principle Hund’s Rule 12.4 Aufbau Principle Electrons enter orbitals of lowest energy first Orbitals within a sublevel of a principle energy level are always of equal energy See Figure 12.7, page 330 12.4 Aufbau Diagram 12.4 Pauli Exclusion Princple An atomic orbital may describe at most two electrons To occupy the same orbital, two electrons must have opposite spins (the electrons spin must be paired) or they would repel each other Electron spin can clockwise or counterclockwise and is represented by vertical arrows that point up or down Hund’s Rule When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain on electron with parallel spins (arrows in the same direction). Second electrons then add to each orbital so that their spins are paired with those of the first electrons in the orbital. 12.4 Electron Configurations You can use a shorthand method for showing the electron configuration of an atom This involves writing the energy level and symbol for every sublevel occupied by an electron A superscript indicates the number of electrons occupying that sublevel. See Table 12.2 for examples 12.4 Example 1 Use Figure 12.7 to write electron configurations and orbital diagrams for these atoms. a. phosphorus b. nickel 12.4 Practice Problems 8. Arrange the following sublevels in order of decreasing energy: 2p, 4s, 3s, 3d, and 3p. 9. Write electron configurations for atoms of the following elements. How many unpaired electrons do these atoms have? a. boron b. fluorine 12.5 Exceptional Electron Config. Like most rules, there are exceptions to the rules for determining electron configurations. Cr, Cu, Ag, Au, and Pt are five of the 14 elements that fill their electron orbitals a little differently than all the other elements These particular elements are more stable with an electron moving into 3d and making 4s half filled rather than a filled 4s shell, this is due to their geometric structure 12.5 Exceptional Electron Config. Following the Aufbau diagram, this would be the electron configuration for Cr and Cu: This is the correct electron configuration: