"Chapter 6 Atomic Theory"
Chapter 6 Atomic Theory Greek Idea • Democritus – Matter is made up of indivisible particles • Dalton – One type of atom for each element Thomson’s Model • Discovered electrons • Atoms were made of positive “stuff” • Negative electron embedded inside the atom • “Plum-Pudding” model – [Blueberry muffin model ] Rutherford’s Model • Discovered dense positive piece at the center of the atom – Nucleus – Electrons moved around – Mostly empty space Bohr’s Model – Why don’t the electrons fall into the nucleus? – Electrons move like planets around the sun • In circular orbits at different levels – Amounts of energy separate one level from another Bohr’s Model Nucleus Electron Orbit Energy Levels Bohr’s Model Energy Levels Fifth • Further away from the Increasing energy Fourth nucleus means Third more energy Second • There is no “in First between” energy Nucleus Bohr’s Model The Bohr model shows electrons circling around the nucleus in definite orbits or paths. The electrons orbits the nucleus much like planets circle the sun. Electrons move from one orbit to another. The further away the orbit from the nucleus means more energy. The electrons can not move “in between” orbits; electrons can only exist in an orbit at definite energy levels. Energy Levels The energy absorbed or released when electrons change energy levels is in the form of electromagnetic radiation (light). Energy – When electrons “jump” orbits (move up or down) energy is Energy needed or released respectively needed • Energy is in the form of light 2 – It was through the study of light that several advances in science were made • Physicists determined the properties of light 1 • The Quantum Mechanical model Energy was produced released Electromagnetic Radiation – Light is just one component of the Electromagnetic Spectrum • Composed of waves of different energies – Waves travel through empty space as well as through air and other substances – Electromagnetic radiation has a dual "personality“ • Acts like waves and particles (photons) • The photons with the highest energy correspond to the shortest wavelengths Electromagnetic Spectrum Electromagnetic Radiation Parts of the Wave Crest Wavelength Origin Amplitude Trough Parts of a Wave Crest – Origin Wavelength • the base-line of energy – Crest Origin • Highest point on a wave Amplitude – Trough • Lowest point on a wave Trough – Amplitude • Distance from origin to crest – Wavelength • Distance from crest to crest (abbreviated “l” - lambda) – Frequency • The number of waves passing a given point per second • Units are cycles/sec or hertz (Hz) Frequency and the Spectrum – Frequency and wavelength are inversely related n=C l Frequency = speed of light / wavelength – Different frequencies of light have different energy levels –E = hν (Energy = Planck’s constant x frequency) • High frequency light has high energy (violet light) • Low frequency light has low energy (red light) – White light is made up of all the colors of the visible spectrum, and all frequencies therein – The whole range of colours is called a continuous spectrum; one colour leads into the next with no break And now back to the atom … – Electrons occupy the lowest energy levels, making the atom stable (low energy content; ground state) – When electrons interact with energy (photon) they may absorb it and move away from the nucleus • Referred to as an electron transition – The electrons are no longer in the ground state; they are in an excited state – The electrons can return to the ground state by releasing quanta of energy • Energy released is of a definite quantity • The energy of electron transitions are quantized (fixed), quantum of energy is released Catch and Release Quantized energy is Quantized energy is absorbed emitted (released) – Electron is excited – Electron jumps to a into a higher energy lower energy level level Bohr Model Emission Spectra – Light of different frequencies have different characteristic colours • Low energy, low frequency light is seen as red light; high frequency, high energy light is seen as violet light – All possible jumps from one energy level to another can occur at the same time because of the number of atoms present, each giving off its characteristic colour – The colour that we see is a mixture of colours • Can be separated by using a diffraction grating – Breaks up the predominant colour into bright lines of specific colours representing electron transitions Emission Spectra Hydrogen Emission Spectrum Iron Emission Spectrum The Quantum Mechanical Model – Bohr’s model of the atom introduced the concept of quantum energy levels, but couldn’t explain how electrons are arranged in atoms – Louis de Broglie (1892-1987) • Proposed that if waves can have particle-like behaviour, particles of matter can behave like waves [under appropriate conditions] • Suggested that as an electron moves about the nucleus, an appropriate wavelength is associated with it – Just a few years later, experimental evidence supported de Broglie … The Quantum Mechanical Model – Werner Heisenberg (1901-1976) • Limit to how precisely we can know both location and momentum of any object [only important with subatomic particles] – Impossible to know both exact momentum and exact location of electron at any point in time » Heisenberg Uncertainty Principle – Erwin Schröndinger (1887-1961) • Proposed an equation which leads to a series of wave functions – Related the probability of finding the electrons to a particular volume of space (3D) » Schröndinger’s wave equation The Plot for atomic orbital • The orbital is classified by its shape. • S orbital a sphere. • P orbital is a dumbbell shape. Atomic Orbitals – Represents the likelihood of finding an electron at a particular point • Densest near the nucleus; less dense with increasing distance – Indicates most probable location around the nucleus – Scientists arbitrarily draw the orbital surface to contain 90% of the total probability distribution – To assign relative sizes and energies to orbitals, the quantum mechanical model assigns Principal Quantum Numbers (n) • “n” specifies the atom’s major energy levels [Principal Energy Levels] Atomic Orbitals – Within each energy level the complex math of Schrödinger’s equation describes several shapes • Each type of shape is identified as a sublevel containing a specific atomic orbital – Energy level 1 (n=1) has one sublevel that contains one “s” orbital • Every succeeding energy level contains an s sublevel with an s orbital • The s orbital is spherically shaped – Each s sublevel is identified by its principal quantum number n • Called 1s, 2s, 3s, … – Each s sublevel has one s type orbital, and can hold 2 electrons Atomic Orbitals – Energy level 2 (n=2) has 2 sublevels, the s and p sublevel (The sublevel name comes from the type of orbital it contains.) • Every succeeding energy level contains a p sublevel with three p orbitals – The p orbital is shaped like a dumbbell and labelled depending on its orientation • One p orbital is oriented along the x axis (px) • One p orbital is oriented along the y axis (py) • One p orbital is oriented along the z axis (pz) – Each p orbital can hold 2 electrons, so the p sublevel can hold a maximum of six electrons Atomic Orbitals – The third energy level (n=3) has three sublevels • s sublevel with one s orbital • p sublevel with three p orbitals • d sublevel with five d orbitals – The d orbitals are more complex in shape and are oriented along planes not axes – Each of the five orbitals can hold two electrons so the d sublevel can hold a maximum of ten electrons Atomic Orbitals – The fourth energy level (n=4) has four sublevels • s sublevel with one s orbital • p sublevel with three p orbitals • d sublevel with five d orbitals • f sublevel with seven f orbitals – The f orbital are even more complex in shape – Each of the seven orbitals can hold two electrons so the f sublevel can hold a maximum of fourteen electrons Orbitals Summary – Atomic orbital pictures – Atomic orbital density representations Principal Number of Number of Total number of quantum number sublevels orbitals electrons (n) n n n2 2n2 1 1 (s) 1 2 2 2 (s, p) 4 8 3 3 (s, p, d) 9 18 4 4 (s, p, d, f) 16 32 Electron Configuration – The Electron Configuration is the arrangement of the electrons within an atom • The atomic orbitals do not fill up in a neat order; the energy levels overlap – Large complex orbital shapes cause the electron to be, on average, further from the nucleus than the simple orbital shapes from greater energy levels – Arrangements follow specific principles: • aufbau principle – Electrons enter the lowest available energy level or sublevel first – This causes difficulties because of the overlap of orbitals of different energies Electron Configuration – Pauli Exclusion Principle • At most 2 electrons per orbital – Electrons have the same charge and will repel each other – The repulsion can be reduced if the electrons have different spins (paired spins) leading to opposite magnetic force fields – Hund’s Rule • When electrons occupy orbitals of equal energy (orbitals within the same sublevel) they don’t pair up until the sublevel is half full 7p 6d 7s 6p 5f 5d 6s 5p 4f 4d 5s Increasing energy 4p 4s 3d 3p 3s 2p 2s 1s n=4 With more sublevels the energy level n=3 gets wider (leads to an overlap). n=2 n=1 Electron Configuration Electron Configuration 7p 6d 5f ← 7s 6p 5d 4f ← 6s 5p 4d ← 5s 4p 3d ← 4s 3p ← 3s 2p ← 2s ← 1s ← Start Determining Electron Configuration – Let’s determine the electron configuration for phosphorus • Need to account for 15 electrons – The first two electrons go into the 1s orbital (1s2) • Remember electrons have opposite spins – The next two electrons go into the 2s orbital (2s2) – The next six electrons go to the 2p sublevel (2p6) • 3 paired spins. – Next is the 3s orbital with 2 electrons (3s2) • Total is now 12, only 3 more to go – The next three electrons go 3px1, 3py1, 3pz1 – The Electron Configuration for phosphorus is: 1s2 2s2 2p6 3s2 3px1, 3py1, 3pz1 1s2 2s2 2p6 3s2 3p3 (not showing the individual p orbitals) Determining Electron Configuration – Try for: Electron Configuration • Na 7p 6d 5f ← •C 7s 6p 5d 4f ← • Cl 6s 5p 4d ← 5s 4p 3d ← – Start off with “How many 4s 3p ← electrons to place?” 3s 2p ← 2s ← 1s ← Start – Add electrons to the orbitals starting at 1s and following through the table Valence Shell and Valence Electrons – The highest numbered energy level is given a special name: the valence shell • The electrons in this valence shell are referred to as valence electrons – Only the outer s, and p electrons Valence Shell and Valence Electrons Electron Configuration As electrons fill the sublevels, we 7p 6d 5f ← see: 7s 6p 5d 4f ← • 1s2 2s2 2p6 3s2 all fill up first 6s 5p 4d ← • 3p6 4s2 now we’re in the 4th 5s 4p 3d ← level, so we start in a new valence 4s 3p ← shell 3s 2p ← 2s ← • 3d10 4p6 5s2 again, a new shell 1s ← Start • The d and f electrons never “see” valence because a new level is started before they’re filled Valence Shell and Valence Electrons – The highest numbered energy level is given a special name: the valence shell • The electrons in this valence shell are referred to as valence electrons – Only the outer s, and p electrons – Identify each element, state the valence shell and the number of valence electrons for: Ne a.1s2 2s2 2p6 K b. 1s2 2s2 2p6 3s2 3p6 4s1 Ge c. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2 d. 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 Sr Exceptions to the Electron Configuration – Orbitals fill in order of lowest to highest energy • The most stable sublevel or energy level is “full” • The next most stable is a “half full” – Adding electrons can change the energy of the orbital or sublevel – “More than half filled” sublevels or “less than half” filled sublevels are higher in energy than those that are exactly half full • The most stable is a full sublevel (least energy) – Sometimes electrons will be promoted/demoted to help the atom stay in the lowest energy state (most stable) possible Electron Exceptions – Copper has 29 electrons so we expect 1s22s22p63s23p64s23d9 – The electron configuration for the copper is also 1s22s22p63s23p64s13d10 • One 4s electron has been promoted to a 3d orbital, leaving a valence shell with only 1 electron and a full 3d – Copper exists in both states • 1s22s22p63s23p64s23d9 • 1s22s22p63s23p64s13d10 Copper Ions – Because copper can exist in both states, it can form 2 different cations •Cu2+ – 1s22s22p63s23p64s03d9 4s2 electrons are lost • Cu+ – 1s22s22p63s23p64s03d10 4s1 electron is lost – Many transition metals can form multiple cations Electron Exceptions – Iron with 26 electrons has an electron configuration of 1s22s22p63s23p64s23d6 – When reacting, it can lose the 2 valence 4s2 electrons creating a 2+ cation • 1s22s22p63s23p64s03d6 – Iron can also lose one 3d electron in addition, to yield a half full 3d sublevel 3d5, creating a 3+ cation • 1s22s22p63s23p64s03d5 – most stable Silver (Ag) – Electron configuration should be 1s22s22p63s23p64s23d104p65s24d9 • It’s actually 1s22s22p63s23p64s23d104p65s14d10 – Results in a lower overall energy, and higher stability – Full d sublevel; partially filled s sublevel 5p 4f E 4d 5s