* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
Download chapter5
Condensed matter physics wikipedia , lookup
Periodic table wikipedia , lookup
Molecular Hamiltonian wikipedia , lookup
Computational chemistry wikipedia , lookup
Bremsstrahlung wikipedia , lookup
Bent's rule wikipedia , lookup
Electrical resistivity and conductivity wikipedia , lookup
Molecular orbital wikipedia , lookup
Metastable inner-shell molecular state wikipedia , lookup
Resonance (chemistry) wikipedia , lookup
Chemical bond wikipedia , lookup
Marcus theory wikipedia , lookup
Jahn–Teller effect wikipedia , lookup
Low-energy electron diffraction wikipedia , lookup
Wave–particle duality wikipedia , lookup
Atomic nucleus wikipedia , lookup
Photoredox catalysis wikipedia , lookup
Gaseous detection device wikipedia , lookup
Oxidative phosphorylation wikipedia , lookup
Extended periodic table wikipedia , lookup
Rutherford backscattering spectrometry wikipedia , lookup
Electron transport chain wikipedia , lookup
X-ray fluorescence wikipedia , lookup
Hydrogen atom wikipedia , lookup
Metallic bonding wikipedia , lookup
Auger electron spectroscopy wikipedia , lookup
Photoelectric effect wikipedia , lookup
X-ray photoelectron spectroscopy wikipedia , lookup
Molecular orbital diagram wikipedia , lookup
Atomic orbital wikipedia , lookup
Light-dependent reactions wikipedia , lookup
Photosynthetic reaction centre wikipedia , lookup
Electron-beam lithography wikipedia , lookup
Chapter 5 Electrons in Atoms 5.1 Models of the atom Rutherford's atomic model could not explain the chemical properties of elements due to the lack of explanation regarding the "electron cloud". Electrons – determine the chemical properties of an atom ( i.e. is it highly or slightly reactive? or unreactive?). These properties cannot be understood without knowing where the electrons are, and how they are arranged around the nucleus The Bohr Model In 1913, Neils Bohr, a student of Rutherford, refined Rutherford's atomic model in order to explain why atoms of all elements give off light when heated. He proposed a planetary model of the atom with the electrons orbiting around the nucleus in a specific circular paths. Each electron has an energy level. Each energy level of the electron can be thought of as rungs on a ladder. The energy levels closest to the nucleus are like rungs of a ladder closest to the ground. The rungs at the top of the ladder are analogous to the energy levels furthest from the nucleus. People can go up a ladder by going from rung to rung the way electrons in an atom can go from energy level to energy level. A person can never stand between the rungs of a ladder and electrons must always be at a specific energy level as well (and not between them). In order for a person to go from rung to rung, they must move a specific distance. In order for electrons to move from energy level to energy level, they must gain or lose a certain amount of energy. The amount of energy is called a quantum of energy. quantum - amount of energy required to move an electron from one energy level to another energy level The amount of energy an electron gains is not always the same. The energy levels in an atom are not equally spaced like a normal ladder. The Bohr model was successful in understanding hydrogen's electron, but failed to account for the energy emitted by electrons in other atoms The Quantum Mechanical Model - The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus. Electrons only have a probability of being in a certain location, the same way the exact location of a fast moving propeller blade at any time cannot not be determined. In the quantum mechanical model, the probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloud. The cloud is more dense where the probability of finding the electron is high. Atomic Orbitals- are regions of space in which there is a high probability of finding an electron. (a) Principal Energy levels (sometimes called shells) -The energy levels of electrons are labeled by principal quantum numbers (n), which are integers starting at n = 1, 2, 3, 4, and so on. (b) Energy sublevels - For each principal energy level there may be several orbitals with different shapes that have different energy levels. These various energy levels within the principal energy level are called sublevels. **Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found. **Different atomic orbitals are denoted by letters. The s orbitals are spherical, and p orbitals are dumbbell-shaped, and the d orbitals are mostly clover shaped. 5.2 Electron arrangement in Atoms Electron Configurations (Shown in the bottom of each elements box on the periodic table)- the way that electrons are arranged in various orbitals around the nuclei of atoms (electrons and protons always interact to make the most stable arrangement possible, which is the configuration with the lowest amount of energy). Ground state - when an atom's electrons are in the lowest possible energy state Three main rules for finding/writing electron configurations 1. Aufbau Principle - electrons occupy the orbitals of lowest energy first. Each box below represents an atomic orbital The orbitals for any sublevel of a principal energy level are always of equal energy The s sublevel always has the lowest energy in the principal energy level Notice that the 4s orbital has a lower energy than the 3d orbital 2. Pauli Exclusion principle - only two electrons can occupy an orbital and the two electrons must have opposite spins (this is called paired electrons) 3. Hund's Rule - states that electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible. ( i.e. when filling orbitals, only one electron can be placed in orbitals of equal energy until they are all filled, only then can they be paired up) Ex. Consider the electron configuration of oxygen It has 8 electrons. The orbital of lowest energy is 1s (only one atomic orbital). It can hold two electrons of opposite spin. The first electron is placed in arrow up, while the next electron is placed in arrow down. This leaves 6 remaining electrons that must fill higher energy electron orbitals. The next highest sublevel is 2s which can hold two electrons with opposite spin as well. This leaves 4 remaining electrons which must fill the next highest energy sublevel: 2p. The 2p sublevel contains 3 atomic orbitals that can hold a maximum of six electrons, but we only have 4 electrons remaining to place in. Each orbital must receive 1 electron pointing in the up direction before any electrons are paired. This leaves one electron remaining which must be placed into the first orbital that already contains one electron in the opposite direction or spin. To the right of the above orbital diagrams, the short hand method for writing electron configurations is shown, which shows the number of electrons in the sublevel as a superscript. For instance, Na is shown as 1s22s22p63s1 There are exceptions to the orbital filling diagrams for specific elements - this occurs after vanadium (Z=23) because the d orbital fills after 4s and will take an electron from the 4s orbital to exactly half fill or completely fill the d orbital - this configuration is more stable Incorrect Correct Ex. Cr 1s22s2 2p63s23p63d44s2 --------------> Cr 1s22s22p63s23p63d54s1 Ex. Cu 1s22s22p63s23p63d94s2 ---------------> Cu 1s22s2sp63s23p63d104s1 More exceptions occur in the higher energy levels due to the energy decrease between principal energy levels difference at higher principal quantum numbers (n). Section Review Questions 1. Write the electron configuration and the expanded electron configuration of chlorine. 2. Write the electron configuration and the expanded electron configuration of titanium. 3. Write the electron configuration and the expanded electron configuration of bromine. 5.3 Physics and Quantum Mechanical Model Light is electromagnetic energy and is part particle and part wave. The color of light corresponds to a specific wavelength or frequency of electromagnetic energy. Where: red light has the longest wavelength (above 700 nm) and a relatively low energy (frequency) violet light has a shorter wavelength (around 400 nm) and a relatively high energy or frequency Atomic Spectra - When atoms absorb energy in the form of excessive heat or electrical charge, electrons move into higher energy levels. This is called the excited state as opposed to ground state Ex. Na 2-8-1 ----------------------------------------> 2-7-1-1 ground heat/electricity excited state state These electrons then lose energy by emitting light when they return to lower energy levels. When this light is passed through a diffraction grating or a spectroscope the wavelength/frequencies of light are separated into bands of light called an atomic emission spectrum. Every element has its own unique atomic emission spectrum. The same way no two people have the same fingerprints, no two elements have the same emission spectrum. Ex. Helium Atomic emission spectrums are used to identify elements Ex. The spectrums of hydrogen, mercury, and neon are compared below. Each element always look the same! Ex. Use the four bright-line spectra data of lithium, hydrogen, helium and sodium to determine the elements in the unknown sample. Bohr used physics and mathematics to explain the bright-line spectrum of hydrogen. Based on his data, he proved that hydrogen’s 1 electron could produce all 4 spectral lines in the bright line spectrum by showing that the electron normally existed in a “ground state” energy level. But when heated, the electron would become excited. E absorbed -----------------> Ground state E released ----------------> Excited state ground state + light (spectral line) or photon Explanation of a all four spectral lines Therefore – The Bohr model proposed that all atoms had electrons orbiting the nucleus in electron shells (also called principle energy levels). These shells exist even when electrons are not present to fill them. The electrons always remained in the ground state or lowest possible E level unless they were excited by E. Bohr also proposed that only 2 e- could fit in the first energy level, 8 in the 2nd, and 8 in the 3rd. Valence electrons – electrons in the highest occupied energy level of an atom Valence electrons determine an atoms chemical properties, and are typically the only electrons involved in chemical bonding. They are the "glue" that holds atoms together The electron configurations of all atoms are shown in the bottom of the elements box in the periodic table The number farthest to the right in these configurations is the number of valence electrons for the atoms for each element Quantum Mechanics - A new branch of physics that takes into account the wave and particle nature of matter. This discovery of "matter waves" has lead to the modern wave-mechanical model of the atom Main ideas are: 1. all moving objects have wavelike behavior 2. wavelengths of large objects is unobservable but the wavelengths of small particles like subatomic particles exhibit wavelengths that are easily detectable. 3. Heisenburg uncertainty which applies to very small objects like electrons states: It is impossible to know both the position and velocity of a particle at the same time. Section Review Questions 1. What is the ground state electron configuration of calcium? Write a possible excited electron configuration for calcium? 2. Which of the following electron configurations is an excited state for magnesium 2-8-2 2-7-4 2-8-3 2-7-3 3. Describe how an atomic emission spectrum is produced. 4. What is an excited electron? 5. What is a ground state electron? 6. Do any different elements have the same emission spectrum? explain. 7. What element could have a possible electron configuration of 2-8-17-3? Is this a ground state electron configuration