CHEMSITRY NOTES Chapter 13. Electrons in Atoms

Transcription

1 CHEMSITRY NOTES Chapter 13 Electrons in Atoms Goals : To gain an understanding of : 1. Atoms and their structure. 2. The development of the atomic theory. 3. The quantum mechanical model of the atom. 4. Electron configurations NOTES: The different historical models are described as follows: Dalton's model of the atom - solid, tiny, indivisible particles. Thomson's model - often describe as the "plum pudding" model - electrons are scattered throughout the atom. Rutherford's model - includes the solid nucleus in the center of the atom. Niels Bohr's model - electrons are in fixed orbitals, or energy levels, at certain distances from the nucleus much the way the planets orbit the sun. Quantum mechanical model - the electrons have certain probabilities of being located at certain places and their positions are described according to four quantum numbers. The circle in the diagram below shows where the electrons have a 90% probability of being within the electron cloud. The quantum mechanical model is the most accepted model today. The diagrams you are probably most familiar with are that of the Bohr model, which is useful in teaching basic electron configuration (distribution in energy levels). Quanta (s. quantum) are the discrete amounts of energy that an electron absorbs as it moves up an energy level or releases when it moves down an energy level. Energy levels are certain regions surrounding the nucleus of an atom where electrons are likely to be found. The quantum mechanical model of the atom is a mathematical model which predicts the probability of electron locations and paths in electron clouds. The positions and orbits of electrons are referred to as energy states and are described by four quantum numbers : the principle quantum number (n) - indicates the energy level the electrons are in (there are seven energy levels, therefore n may equal 1,2,3,4,5,6 or 7) the orbital quantum number - indicates the shape of the atomic orbital (region in space where there is a high probability of finding an electron). the first four orbital numbers are indicated by the letters s (has a spherical shape), p (has a dumbbell shape), d, and f (d and f have complex shapes). magnetic quantum number - indicates the position of the orbital about the x, y, and z axes spin quantum number - indicates the direction of spin of an electron - clockwise or counterclockwise

2 These quantum numbers can be summarized as follows: Principle energy level Number of sublevels Type of sublevel (number of orbitals per sublevel) Total number of sublevels (= n 2 ) Total number of electrons(= 2n 2 ) n=1 1 1s (one orbital only) 1 2 n=2 2 2s (1 orbital), 2p (3 orbitals) 4 8 n=3 3 n=4 4 3s (1 orbital), 3p (3 orbitals), 3d (5 orbitals) 4s (1 orbital), 4p (3 orbitals), 4d (5 orbitals), 4f (7 orbitals) Note : Each quantum position as described by the first three quantum numbers can be occupied by only two electrons with opposite spins. For example, if the position of a pair of electrons are described as 4px, this means they are in the fourth energy level (n=4), and have a dumbbell shaped orbital (orbital quantum number is p) around the x axis (third quantum number). These two electrons in this orbital, as stated above, would have opposite spins. Electrons must be placed in the lowest possible energy levels first. This is referred to as the ground state of an atom - the state with the lowest possible energy level. There are three rules which govern electron configuration (the arrangement of the electrons of an atom into its orbitals). Aufbau (German for "building up') Principle - electrons enter orbitals of the lowest energy levels first. The following is the aufbau help diagram. The arrows indicate the lowest energy levels which need to be filled first. The 1s is always filled first then the 2s, 2p, 3s etc.. Note the 4s is filled before the 3d. Pauli Exclusion Principle - an atomic orbital may describe at most two electrons ; electrons in the same orbital must have opposite spins. Hund's rule - when electrons occupy orbitals of the same energy level (e.g. the three 3p orbitals or the five 4d orbitals), one electron will enter each orbital with a parallel spin until all orbitals have one electron. Then and only then will a second electron be added to those orbitals. The following diagram shows four elements and their electron configurations. They are explained below.

3 Hydrogen has only one electron so it is in the lowest orbital, the 1s orbital. The upwards arrow indicates a clockwise spin. Hydrogen's electron configuration can then be described as 1s 1. This means there is one electron in the first energy level in the s orbital (spherical shape). Helium has two electrons. The first orbital will hold two electrons so its summary electron configuration is 1s 2. Note the arrows indicating opposite spins (Pauli exclusion principle). Lithium has three electrons so one has to move into the next lowest orbital, the 2s. Its summary electron configuration is 1s 2 2s 1. Nitrogen has seven electrons. The 1s and 2s fill up first (Aufbau principle). Now one electron will go into each of the 2p orbitals with the same spin according to the Hund's rule. Nitrogen's summary electron configuration is 1s 2 2s 2 2p 3. Oxygen's electron configuration is similar to nitrogen, only now there are four electrons to put into the three 2p orbitals. Is summary electron configuration is 1s 2 2s 2 3p 4. Flame tests are tests used to identify metal or metallic ions. The substance to be identified is usually put on a wire loop and then placed into a flame. The flame will then show a color which can be used to help identify the metal. The colors in fireworks are produced by metals or metallic ions added to the fireworks. The color comes from energy being released by the metallic atoms or ions. When energy is added to the atoms or ions electrons jump up energy levels as they absorb the energy. When they fall back down energy levels they release energy in the form of light of a characteristic wavelength or color for that element. For example barium gives off a pale green color and strontium gives off a brilliant red color. Spectroscopy is the study of spectra - the types of light (wavelengths) that are absorbed or emitted by elements. Spectra allow scientists to, in a sense, see into the atomic structure of matter. The colors emitted or absorbed indicate changes in the positions of electrons in their orbitals. Emission spectra are the spectra released by substances that have been excited by the addition of energy. Absorption spectra are spectra that show the wavelengths (colors) of light that are absorbed by substances. Spectroscopy is a very important tool in the hands of an analytical chemist. Qualitative data (what is present) can be determined by the wavelengths of light that are produced - much like the flame tests. Quantitative data (how much is present) can be determined by the brightness or thickness of the spectral bands that are produced.. The elements are arranged in order of increasing atomic number across the periods (horizontal rows). The periods are arranged so that the elements in the vertical columns (groups or families) have similar properties. This causes the properties of elements to change as you move horizontally from group to group across a period. Note that their are 7 periods - one for each energy level (principle quantum number). The periodic law states that when the elements are arranged according to increasing atomic number there is a periodic pattern in their physical and chemical properties. Although the periodic table is arranged by increasing atomic number (which indicates the number of protons), the electrons configuration is what really determines the physical and chemical properties of the elements. The periodic table can be divided into four groups based on electron configurations : The Noble gases (Group 0) - have their outermost s and p orbitals filled which creates a stable and non-reactive (inert) element. The representative elements - Group A elements - have their s and p orbitals being filled. These include : o Group 1A - Li, Na, K etc. - all very reactive with one electron in the outer s orbital o Group 2A - Be, Mg, Ca etc. - all quite reactive with 2 electrons filling their outer s orbital o Group 3A - Aluminum group - 3 electrons in outer energy level (2s and 1p) properties vary from metallic to metalloid

4 o Group 4A - Carbon group - 4 electrons in outer energy level (2s and 2p) - properties vary from nonmetallic to metalloid to metallic down the group o Group 5A - Nitrogen group - 5 electrons in outer energy level (2s and 3p) - properties vary from nonmetallic to metalloid to metallic o Group 6A - Oxygen group - 6 electrons in outer energy level (2s and 4p) - properties vary from nonmetallic to metalloid o Group 7A - Halogens - all have 7 electrons in the outer energy level (2s and 5p) - properties vary from nonmetallic to metalloid. Very reactive due to the outer energy level being almost filled. The transition metals - elements whose d orbitals are being filled - found in the "d-block." These are also called the Group B elements The Inner transition metals - These are the Lanthanide and Actinide series, element whose f orbitals are being filled. The s, p, d, and f groups can be identified on the diagram below. The f block (inner transition metals) is usually shown separated and below the rest of the table. The accepted theory which explains the origin of the elements and all matter in the universe is the Big Bang theory. It states that the universe began with an explosion of tremendous energy. This energy was converted into matter, according to Einstein's equation E=mc 2. At first all matter was in the form of quarks. As the universe expanded it cooled allowing matter to condense and form the lightest elements first and then the heavier elements. The representative groups of elements : Noble gases - Group 0, helium, neon, argon, krypton, xenon and radon o inert (unreactive) because of stable electron configuration (filled s and p orbitals) o helium is used in weather balloons o helium and neon are used to create artificial, unreactive environments (less soluble than nitrogen and therefore less likely to cause the bends o other noble gases are used to create unreactive environments in flashbulbs or aluminum welding Alkali metals - Group 1A, lithium, sodium, potassium, rubidium, cesium and francium o very reactive (one electron away from a filled s and p orbital) o low density o low melting point o good electrical conductivity o react with water to form strong bases (sodium hydroxide, lithium hydroxide etc.) Alkaline earth elements - Group 2A, beryllium, magnesium, calcium, strontium, barium and radium o very reactive (2 electrons away from a filled s and p orbital) o react with water to form hydroxides o used to form metal alloys Aluminum group - Group 3A - 3 electrons in outer energy level (2s and 1p) properties vary from metallic to metalloid o aluminum is the most useful metal of this group being lightweight and strong to make boats, aircraft etc. Group 4A - Carbon group - 4 electrons in outer energy level (2s and 2p) - properties vary from nonmetallic to metalloid to metallic down the group o diamond and graphite are forms of pure carbon o silicon and germanium are semiconductors used in electronics

5 o tin and lead are useful metals Group 5A - Nitrogen group - 5 electrons in outer energy level (2s and 3p) - properties vary from nonmetallic to metalloid to metallic o nitrogen and phosphorus are elements necessary to form proteins and nucleic acids in living things Group 6A - Oxygen group - 6 electrons in outer energy level (2s and 4p) - properties vary from nonmetallic to metalloid o oxygen is the most abundant element on the earth o sulfur has many industrial uses (sulfuric acid is the most widely used industrial chemical) Group 7A - Halogens - all have 7 electrons in the outer energy level (2s and 5p) - properties vary from nonmetallic to metalloid. Very reactive due to the outer energy level being almost filled. o iodine is used as an antiseptic o chlorine is a bleaching agent and disinfecting agent o fluorine, as the fluoride ion, is used to maintain the health of our teeth o fluorine is used to make teflon The Development of a New Atomic Model I. Properties of Light A. Electromagnetic Radiation 1. Many types of EM waves a. visible light b. x-rays c. ultraviolet light d. infrared light e. radio waves 2. EM radiation are forms of energy which move through space as waves a. Move at speed of light (1) x 10 8 m/s b. Speed is equal to the frequency times the wavelength c = λν (1) Frequency (ν) is the number of waves passing a given point in one second (unit = 1/s) (2) Wavelength (λ) is the distance between peaks of adjacent waves (unit = m) c. Speed of light is a constant, so λν is also a constant (1) λ and ν must be inversely proportional B. Light and Energy - The Photoelectric Effect 1. The Photoelectric Effect a. Electrons are emitted from a metal when light shines on the metal 2. Radiant energy is transferred in units (or quanta) of energy called photons a. A photon is a particle of energy having a mass of zero and carrying a quantum of energy b. A quantum is the minimum amount of energy that can be lost or gained by an atom 3. Energy of a photon is directly proportional to the frequency of radiation a. E = hν (h is Planck s constant, x J.s )

6 4. Wave-Particle Duality a. Energy travels through space as waves, but can be thought of as a stream of particles (Einstein) II. The Hydrogen Line Spectrum A. Ground State 1. The lowest energy state of an atom B. Excited State 1. A state in which an atom has a higher potential energy than in its ground state C. Bright line spectrum 1. Light is given off by excited atoms as they return to lower energy states 2. Light is given off in very definite wavelengths 3. A spectroscope reveals lines of particular colors (with a specific wavelength and frequency) Bright line spectrum of Hydrogen III. The Bohr Model of the Atom A. Electron Orbits, or Energy Levels 1. Electrons can circle the nucleus only in allowed paths or orbits (like planets around the sun) 2. The energy of the electron is greater when it is in orbits farther from the nucleus 3. The atom achieves the ground state when atoms occupy the closest possible positions around the nucleus 4. Electromagnetic radiation is emitted when electrons move closer to the nucleus

7 B. Energy transitions 1. Energies of atoms are fixed with definite quantities 2. Energy transitions occur in jumps of discrete amounts of energy 3. Electrons only lose energy when they move to a lower energy state C. Shortcomings of the Bohr Model 1. Doesn't work for atoms larger than hydrogen (with more than one electron) 2. Doesn't explain chemical behavior The Quantum Model of the Atom I. Electrons as Waves and Particles A. Louis debroglie (1924) 1. Electrons have wavelike properties 2. Consider the electron as a wave confined to a space that can have only certain frequencies B. The Heisenbery Uncertainty Principle (Werner Heisenberg ) 1. "It is impossible to determine simultaneously both the position and velocity of an electron a. Electrons are located by their interactions with photons b. Electrons and photons have similar energies c. Interaction between a photon and an electron knocks the electron off of its course C. The Schrodinger Wave Equation 1. Proved quantization of electron energies and is the basis for Quantum Theory a. Quantum theory describes mathematically the wave properties of electrons 2. Electrons do not move around the nucleus in "planetary orbits" 3. Electrons exist in regions called orbitals a. An orbital is a three-dimensional region around the nucleus that indicates the probable location of an electron (a 90% probability of finding an electron) II. Atomic Orbitals and Quantum Numbers Quantum Numbers specify the properties of atomic orbitals and the properties of the electrons in orbitals A. Principal Quantum Number (n) 1. Indicates the main energy levels occupied by the electron (distance from the nucleus) (size) 2. Values of n are positive integers a. n=1 is closest to the nucleus, and lowest in energy 3. The number of orbitals possible per energy level (or "shell") is equal to n 2 B. Angular Momentum Quantum Number (l) 1. Indicates the shape of the orbital 2. Number of orbital shapes = n a. Shapes are designated s, p, d, f (sublevels)

8 C. Magnetic Quantum Number (m) 1. The orientation of the orbital around the nucleus a. s orbitals have only one possible orientation b. p orbitals have three, d have five and f have 7 possible orientations D. Spin Quantum Number 1. Indicates the fundamental spin states of an electron in an orbital 2. Two possible values for spin, +1/2, -1/2 or 3. A single orbital can contain only two electrons, which must have opposite spins

9 Electron Configurations I. Writing Electrons Configurations A. Rules 1. Aufbau Principle a. An electron occupies the lowest-energy orbitals first 2. Pauli Exclusion Principle a. No two electrons in the same atom can have the same set of four quantum numbers (opposite spins) 3. Hund's Rule a. Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin B. Orbital Notation 1. Unoccupied orbitals are represented by a line, (or a box or circle) a. Lines are labeled with the principal quantum number and the sublevel letter 2. Arrows are used to represent electrons a. Arrows pointing up and down indicate opposite spins C. Electron Configuration Notation 1. The number of electrons in a sublevel is indicated by adding a superscript to the sublevel designation Hydrogen = 1s 1 Helium = 1s 2 Lithium = 1s 2 2s 1 II. Survey of the Periodic Table A. Elements of the Second and Third Periods 1. Highest occupied energy level a. The electron containing energy level with the highest principal quantum number 2. Inner shell electrons a. Electrons that are not in the highest energy level 3. Octet a. Highest energy level s and p electrons are filled (8 electrons) b. Characteristic of noble gases, Group Noble gas configuration a. Outer main energy level fully occupied, usually (except for He) by eight electrons b. This configuration has extra stability B. Elements of the Fourth Period 1. Irregularity of Copper a. Expected: [Ar] 4s 2 3d 9 b. Actual: [Ar] 4s 1 3d Several transition and rare-earth elements borrow from smaller sublevels in order to half fill larger d sublevels. Half filled and fully filled sublevels are particularly stable.