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Electrons in Atoms Democritus (400 B.C.) • Proposed that matter was composed of tiny indivisible particles • Not based on experimental data • Greek: atomos Alchemy (next 2000 years) • Mixture of science and mysticism. • Lab procedures were developed, but alchemists did not perform controlled experiments like true scientists. John Dalton (1807) British Schoolteacher based his theory on others’ experimental data Billiard Ball Model atom is a uniform, solid sphere Henri Becquerel (1896) Discovered radioactivity spontaneous emission of radiation from the nucleus Three types: alpha () - positive beta () - negative gamma () - neutral J. J. Thomson (1903) Cathode Ray Tube Experiments beam of negative particles Discovered Electrons negative particles within the atom Plum-pudding Model J. J. Thomson (1903) Plum-pudding Model positive sphere (pudding) with negative electrons (plums) dispersed throughout Ernest Rutherford (1911) Gold Foil Experiment Discovered the nucleus dense, positive charge in the center of the atom Nuclear Model Rutherford’s Gold Foil Experiment (a) The results that the metal foil experiment would have yielded if the plum pudding model had been correct. (b) Actual results. Ernest Rutherford (1911) Nuclear Model dense, positive nucleus surrounded by negative electrons Niels Bohr (1913) Bright-Line Spectrum tried to explain presence of specific colors in hydrogen’s spectrum Energy Levels electrons can only exist in specific energy states Planetary Model Niels Bohr (1913) Bright-line spectrum Planetary Model electrons move in circular orbits within specific energy levels Erwin Schrödinger (1926) Quantum mechanics electrons can only exist in specified energy states Electron cloud model orbital: region around the nucleus where e- are likely to be found Erwin Schrödinger (1926) Electron Cloud Model (orbital) dots represent probability of finding an enot actual electrons James Chadwick (1932) Discovered neutrons neutral particles in the nucleus of an atom Joliot-Curie Experiments based his theory on their experimental evidence James Chadwick (1932) Neutron Model revision of Rutherford’s Nuclear Model Electromagnetic Radiation Electromagnetic radiation – radiowaves, X-rays, microwaves, infrared waves, visible light, ultraviolet waves and gamma rays. All electromagnetic radiation travel at the speed of light (c = 3.0 x 108 m/s) in a vacuum. The different wavelengths of electromagnetic radiation. The electromagnetic spectrum. Physics and the Quantum Mechanical Model Amplitude – wave’s height from the origin to the crest. Wavelength (l)– distance between the crests. Frequency (u)– number of wave cycles to pass a given point per unit of time. Water wave (ripple). Physics and the Quantum Mechanical Model Frequency and wavelength are inversely proportional. As frequency increases, wavelength decreases, and vice versa, but their product will always equal the speed of light. c = lu SI units for frequency are cycles per second is a hertz (Hz), or 1/seconds (1/s or s-1). Relationship Between Wavelength and Frequency Physics and the Quantum Mechanical Model What is the frequency of light that has a wavelength of 550 nm? (1m = 109 nm or 1 nm = 10-9 m)? What is the wavelength of light, in cm, that has a frequency of 9.60 x 1014 Hz (1/s)? What is the frequency of light (Hz) that has a wavelength of 740 nm (1m = 109 nm or 1 nm = 10-9 m)? Physics and the Quantum Mechanical Model Sunlight splits into a spectrum of colors when it passes through a prism. Colors of the spectrum include red, orange, yellow, green, blue, indigo and violet. Red light has the longest wavelength and the lowest frequency, while violet light has the shortest wavelength and the highest frequency. Dispersion of White Light By a Prism A photon of red light (relatively long wavelength) carries less energy than a photon of blue light (relatively short wavelength) does. Physics and the Quantum Mechanical Model Every element emits light after it absorbs energy. The light that is emitted (atomic emission spectra) is different for every element, and differs from white light because it is not continuous. Max Planck said that color changes can be explained if you assume that the energy of a substance changes in small increments. Emission (line) Spectra of Some Elements Emission (line) Spectra of Some Elements (cont’d) Emmision (line) Spectra of Some Elements (cont’d) Physics and the Quantum Mechanical Model Planck showed that the amount of radiant energy (E) absorbed or emitted by a substance is proportional to the frequency of the radiation. E = hu h is Planck’s constant (6.626 x 10-34 J s) Any attempt to increase or decrease the energy of a system by a fraction of h times u will fail because energy is only emitted or absorbed in quanta, or bunches of energy. Planck’s Constant Examples What is the energy of a photon with a frequency of 2.94 x 1015 cycles per second (s-1 or Hz)? What is the energy of a light particle with a wavelength of 675 nm? Homework Problem Examples What is the wavelength, in nm, of light with a frequency of 9.5 x 109 s-1? ( 1 m = 109 nm) How much energy is contained in a photon with a wavelength of 5.17 x 10-4 m? Planck’s Revelation Showed that light energy could be thought of as particles for certain applications Stated that light came in particles called quanta or photons Particles of light have fixed amounts of energy The energy of the photon is directly proportional to the frequency of light Higher frequency = More energy in photons Physics and the Quantum Mechanical Model Photons – light energy. The energy of photons is quantized according to the equation E = hu. Light was therefore thought to have a dual wave-particle behavior to explain all of its characteristics. Electromagnetic radiation (a beam of light) can be pictured in two ways: as a wave and as a stream of individual packets of energy called photons. Bohr’s Model Energy of an electron is related to the distance electron is from the nucleus Energy of the atom is quantized atom can only have certain specific energy states called quantum levels or energy levels when atom gains energy, electron “moves” to a higher quantum level when atom loses energy, electron “moves” to a lower energy level lines in spectrum correspond to the difference in energy between levels Bohr’s Model Atoms have a minimum energy called the ground state The ground state of hydrogen corresponds to having its one electron in an energy level that is closest to the nucleus Energy levels higher than the ground state are called excited states the farther the energy level is from the nucleus, the higher its energy To put an electron in an excited state requires the addition of energy to the atom; bringing the electron back to the ground state releases energy in the form of light (a) A sample of H atoms receives energy from an external source. (b) The excited atoms (H) can release the excess energy by emitting photons. When an excited H atom returns to a lower energy level, it emits a photon that contains the energy released by the atom. Hydrogen atoms have several excitedstate energy levels. Each photon emitted by an excited hydrogen atom corresponds to a particular energy change in the hydrogen atom. Bohr’s Model Distances between energy levels decreases as the energy increases light given off in a transition from the second energy level to the first has a higher energy than light given off in a transition from the third to the second, etc. 1st energy level can hold 2 electrons (e-1), the 2nd 8e-1, the 3rd 18e-1, etc. farther from nucleus = more space = less repulsion Models of the Atom Energy level – region around the nucleus where the electron is likely to be found. Think of steps on a ladder. Essentially, you must be in one energy level or another, you can’t be between energy levels, just like you can’t stand in mid-air between the steps of a ladder. The difference between continuous and quantized energy levels can be illustrated by comparing a flight of stairs with a ramp. Models of the Atom Energy levels are not equally spaced. The further away an electron is from the nucleus, the easier it becomes to pull that electron off of that particular atom. Erwin Schrodinger – in 1926, he came up with a new way of describing the energy and location of an electron, called the quantum mechanical model, which is a mathematical method. Models of the Atom The quantum mechanical model does not say that electrons take exact paths around the nucleus, but that it estimates the probability (likelihood) of finding an electron in a certain position. If the electron cloud is very dense, it is more likely that you will find the electron there, then if the electron cloud is less dense. The probability map, or orbital, that describes the hydrogen electron in its lowest possible energy state. Orbitals Orbital – area where an electron is likely to be found. usually use 90% probability to set the limit three-dimensional Orbitals are defined by three integer terms called the quantum numbers. Each electron also has a fourth quantum number to represent the direction of spin Models of the Atom Principal quantum number (n) – designates the energy level of the electrons. n will always be an integer. The distance from the nucleus increases as n increases. Within each energy level, electrons occupy energy sublevels. The number of energy levels (n) is always the same as the number of sublevels. Models of the Atom Sublevel – part of an energy level. 1st energy level has 1 sublevel (“s” sublevel) 2nd energy level has 2 sublevels (“s” and “p” sublevels) 3rd energy level has 3 sublevels (“s”, “p”, and “d” sublevels) 4th energy level has 4 sublevels (“s”, “p”, “d” and “f” sublevels) An illustration of how principal levels can be divided into sublevels. Principal level 2 shown divided into the 2s & 2p sublevels. Models of the Atom Atomic orbitals – areas where electrons are likely to be found. s orbital – spherical in shape, only 1 s orbital per sublevel. p orbital – dumbbell shaped, 3 p orbitals per sublevel. d orbital – 5 d orbitals per sublevel. f orbital – 7 f orbitals per sublevel. The relative sizes of the 1s and 2s orbitals of hydrogen. The 2p orbitals. The five 3d orbitals. Models of the Atom In any orbital, there can be a maximum of two electrons. The maximum number of electrons that can occupy an energy level is given by the formula 2n2, where n is the # of the energy level. 1st energy level up to 2 electrons 2nd energy level up to 8 electrons 3rd energy level up to 18 electrons 4th energy level up to 32 electrons Quick Review Max of 2 electrons per orbital “s” sublevel – 1 orbital per sublevel (up to 2 total electrons) “p” sublevel – 3 orbitals per sublevel (up to 6 total electrons “d” sublevel – 5 orbitals per sublevel (up to 10 total electrons) “f” sublevel – 7 orbitals per sublevel (up to 14 total electrons) The orbitals being filled for elements in various parts of the periodic table. Electron Arrangements in Atoms Electron configuration – the way in which electrons are arranged in energy levels outside of the nucleus. Orbital notation – a way of showing the electron configuration using arrows to represent each electrons and boxes to represent each orbital. Electron Arrangements in Atoms Rules Aufbau principle – electrons enter orbitals of lowest energy first. Pauli exclusion principle – an atomic orbital may hold at most two electrons. Electrons within the same orbital have opposite spins. Hund’s rule – one electron must be put in each orbital of a sublevel before any one orbital can have two electrons in it. Orbital Notations When writing orbital notations, use one arrow to represent each electron. Electrons must enter the lowest energy sublevel possible before moving to a higher energy sublevel Even if you don’t have enough electrons to fill each orbital of a sublevel, you must still show that those orbitals exist. The total number of arrows (electrons) must be equal to the atomic # for each element. Types of Electrons in Arrangements Shared electrons – orbitals where there are two electrons (arrows) with opposite spins. Unshared electrons – when an orbital only has one electron in it. Shared pair of electrons – any orbital that contains two electrons. A box diagram showing the order in which orbitals fill to produce the atoms in the periodic table. Each box can hold two electrons. Orbital Notations Write the orbital notation for oxygen. Write the orbital notation for aluminum. Write out the orbital notation for cobalt The orbitals being filled for elements in various parts of the periodic table. Electron Arrangements in Atoms Electron Configurations – the way in which electrons are arranged around the nucleus of an atom. Each configuration has 3 parts: 2 1s “1” represents the energy level, “s” represents the sublevel, and “2” represents the number of electrons in that sublevel The total of superscripts is equal to the atomic number for the element. Electron Arrangements in Atoms 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p The orbitals being filled for elements in various parts of the periodic table. Electron Configurations Which element is represented by the following electron configuration: 1s22s22p63s23p6 1s22s22p63s23p64s23d104p65s24d105p66s2 4f145d106p67s1 Electron Configurations Write the electron configuration for the following elements: Sulfur Gallium Thorium Platinum Electron Configurations What is wrong with each of the following electron configurations? 1s22s22p63s23p63d104s24p5 1s22s22p63s23p64s23d104p65s24d105p66s25d106p3 1s22s22p63s23p64s23d84p65s1 Noble Gas Configurations Noble gas configurations are used as a shorthand for long electron configurations. Find the noble gas before the element you are writing the configuration for, put it in brackets, and then start with the next s sublevel to fill out the rest of the configuration. The orbitals being filled for elements in various parts of the periodic table. The periodic table with atomic symbols, atomic numbers, and partial electron configurations. Noble Gas Configurations Write the noble gas configuration for the following elements: Sulfur Iron Thorium Platinum Noble Gas Configurations What element is represented by the following noble gas configuration: [Kr]5s24d105p2 [Ar]4s2 [Xe]6s24f145d6 Noble Gas Configurations What is incorrect about the following noble gas configurations? [Ar]2s22p2 [Kr]4d10 [At] 7s24f146d7 Electron Configuration Elements in the same column on the Periodic Table have Similar chemical and physical properties Similar valence shell electron configurations Same numbers of valence electrons Same orbital types Different energy levels Valence electrons – outer energy level “s” and “p” sublevel electrons or electrons that are furthest away from the nucleus Noble Gas Configurations & their relation to the Periodic Table Lithium – [He]2s1 Fluorine – [He]2s22p5 Sodium – [Ne]3s1 Chlorine – [Ne]3s23p5 Potassium – [Ar]4s1 Bromine-[Ar]4s23d104p5 Rubidium – [Kr]5s1 Iodine-[Kr]5s24d105p5 s1 1 2 3 4 5 6 7 s2 p1 p2 p3 p4 p5 s2 p6 d1 d2 d3 d4 d5 d6 d7 d8 d9 d10 f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 Periodic Trends Atomic radius – distance from the nucleus of an atom to its valence electrons. The radius tells that size of the atom. Moving from left to right across a period, atomic radius decreases. Electrons within the same energy level don’t have as great of an effect on one another as electrons from different energy levels. Trend in Atomic Size Increases down column valence shell farther from nucleus Decreases across period left to right adding electrons to same valence shell valence shell held closer because more protons in nucleus Periodic Trends Moving down a group, atomic radius increases. The valence electrons get further and further from the nucleus because you are adding more energy levels. Therefore the radius of the atom increases. Representation of Atomic Radii of the Main-Group Elements Periodic Trends Example: Put the following elements in order of increasing atomic radius: Zn, Sc, Se, K, Cs, O Example: Put the following elements in order of decreasing atomic radius: F, Cd, Ba, Ge, W, Cl The orbitals being filled for elements in various parts of the periodic table. Periodic Trends Ionization energy – the energy required to remove an electron from an atom (1st ionization energy). Removing an electron creates a charge imbalance, so a cation (positive ion) is formed. 2nd Ionization energy – the energy required to remove two electrons from an atom. Periodic Trends Moving from left to right across a period, ionization energy increases. Within the same energy level electrons experience an increasing pull from the nucleus, so it takes more energy to remove them. Periodic Trends Moving down a group, ionization energy decreases. The valence electrons feel less and less pull from the nucleus as they get further from the nucleus. Periodic Trends 2nd ionization energy is always greater than the 1st ionization energy. When you remove an electron from an atom the number of protons becomes greater than the number of electrons. The remaining valence electrons move closer to the nucleus, making it harder to pull them off the atom. Periodic Trends As electrons are removed, ionization energy increases gradually until an energy level is empty, then it makes a big jump. Pulling an electron off of a alkali metal (Group 1 elements) is easy. Trying to pull an electron off of a noble gas (Group 18 elements) takes much more energy. Periodic Trends Example: Put the following elements in order of increasing ionization energy: Sr, Cr, As, S, Rb, Cu Example: Put the following elements in order of decreasing ionization energy: O, V, K, P, Ga, Fr The orbitals being filled for elements in various parts of the periodic table. Periodic Trends Which of the following elements will have a very large second ionization energy? Third ionization energy? Na, Al, Ne, Mg, Si The orbitals being filled for elements in various parts of the periodic table. Periodic Trends Ionic radius – similar to atomic radius but it is the radius for an ion instead of an atom. Positive ions are always smaller than their neutral atoms, and negative ions are always larger than their neutral atoms. As you go down a group, ionic radius increases. Comparison of Atomic and Ionic Radii Periodic Trends As you go from left to right across a period, positive ions decrease in size. Negative ions also decrease as you go across a period, but they start off being much larger than positive ions. Periodic Trends Put the following ions in order of increasing ionic radius: Hint: If all of the ions have the same number of electrons, than the one with the highest number of protons has the smallest radius. Na+1, Al+3, N-3, F-1, O-2, Mg+2 Periodic Trends Electronegativity – how strongly the nucleus of an atom attracts the electrons of other atoms in a bond. Nonmetals tend to gain electrons when they form bonds, and have higher electronegativities than metals, which tend to lose electrons, when they form bonds. Periodic Trends Moving from left to right across a period, electronegativity increases. Moving down a group, electronegativity decreases. Electronegativities of the Elements Periodic Trends Put the following elements in order of increasing electronegativity: Fe, Si, O, Ba, Ca, Cs Put the following elements in order of decreasing electronegativity: Se, F, Ag, Pt, Fr, Sb