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Chapter 7 Atomic Structure and Periodicity 7.1 Electromagnetic Radiation electromagnetic radiation: form of energy that acts as a wave as it travels includes: visible light, X rays, ultraviolet and infrared light, microwaves, and radio waves travel at a speed of 2.9979 x 108 m/s in a vacuum All forms are combined to form electromagnetic spectrum Electromagnetic Spectrum Low Energy High Energy Radio Micro Infrared Ultra- XGamma waves waves . violet Rays Rays Low High Frequency Frequency Long Short Wavelength Visible Light Wavelength - Page 139 “R O Y Frequency Increases Wavelength Longer G B I V” Parts of a wave Crest Wavelength Amplitude Origin Trough Electromagnetic radiation propagates through space as a wave moving at the speed of light. Equation: c = c = speed of light, a constant (2.998 x 108 m/s) (lambda) = wavelength, in meters (nu) = frequency, in units of hertz (hz or sec-1) Wavelength and Frequency Are inversely related • As one goes up the other goes down. Different frequencies of light are different colors of light. There is a wide variety of frequencies The whole range is called a spectrum The Wave-like Electron The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves. Louis deBroglie Electromagnic radiation propagates through space as a wave moving at the speed of light. Wave nature of electromagnetic Radiation wavelength: λ = Greek letter lambda distance between points on adjacent waves (consicutive peaks or troughs) in nm (109nm = 1m) frequency: = Greek letter nu number of wave cycles that passes a point in a second. 108 cycles/s= 108 s-1 =108 Hertz = 108 Hz in 1/second (Hertz = Hz) c c C = speed of light, a constant (3.00 x 108 m/s) = frequency, in units of hertz (hz, sec-1) = wavelength, in meters Long Wavelength = Low Frequency = Low ENERGY Short Wavelength = High Frequency = High ENERGY Wavelength Table Calculate the energy of red light vs. blue light. red light: 700 nm blue light: 400 nm red: hc E hc E blue: hc E 34 8 (6.62x1034 J s)( 3.00x108 m / s) ( 6 . 62 x 10 J s )( 3 . 00 x 10 m / s) E E 700x109 m 400x109 m E = 2.85 x 10-19 J E = 4.96 x 10-19 J sunburn????? uv 7.2 Nature of Matter Before 1900, scientists thought that matter and energy were totally different matter particles mass position energy wave massless delocalized In 1900 Matter and energy were seen as different from each other in fundamental ways. Matter was particles. Energy could come in waves, with any frequency. Max Planck found that as the cooling of hot objects couldn’t be explained by viewing energy as a wave. Explanation of atomic spectra When we write electron configurations, we are writing the lowest energy. The energy level, and where the electron starts from, is called it’s ground state - the lowest energy level. Changing the energy Let’s look at a hydrogen atom, with only one electron, and in the first energy level. Changing the energy Heat, electricity, or light can move the electron up to different energy levels. The electron is now said to be “excited” Changing the energy As the electron falls back to the ground state, it gives the energy back as light Changing the energy They may fall down in specific steps Each step has a different energy Ultraviolet The Visible Infrared further they fall, more energy is released and the higher the frequency. This is a simplified explanation! The orbitals also have different energies inside energy levels All the electrons can move around. What is light? Light is a particle - it comes in chunks. Light is a wave - we can measure its wavelength and it behaves as a wave 2 If we combine E=mc , c=, E = 1/2 mv2 and E = h, then we can get: = h/mv (from Louis de Broglie) called de Broglie’s equation Calculates the wavelength of a particle. Wave-Particle Duality J.J. Thomson won the Nobel prize for describing the electron as a particle. His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron. The electron is a particle! The electron is an energy wave! Confused? You’ve Got Company! “No familiar conceptions can be woven around the electron; something unknown is doing we don’t know what.” Physicist Sir Arthur Eddington The Nature of the Physical World 1934 The physics of the very small Quantum mechanics explains how very small particles behave • Quantum mechanics is an explanation for subatomic particles and atoms as waves Classical mechanics describes the motions of bodies much larger than atoms Heisenberg Uncertainty Principle It is impossible to know exactly the location and velocity of a particle. The better we know one, the less we know the other. Measuring changes the properties. True in quantum mechanics, but not classical mechanics Heisenberg Uncertainty Principle “One cannot simultaneously determine both the position and momentum of an electron.” Werner Heisenberg You can find out where the electron is, but not where it is going. OR… You can find out where the electron is going, but not where it is! It is more obvious with the very small objects To measure where a electron is, we use light. But the light energy moves the electron And hitting the electron changes the frequency of the light. After Before Photon Moving Electron Photon wavelength changes Electron velocity changes Fig. 5.16, p. 145 Light is a Particle? Energy is quantized. Light is a form of energy. Therefore, light must be quantized These smallest pieces of light are called photons. Photoelectric effect? Albert Einstein Energy & frequency: directly related. The energy (E ) of electromagnetic radiation is directly proportional to the frequency () of the radiation. Equation: E = h E = Energy, in units of Joules (kg·m2/s2) (Joule is the metric unit of energy) h = Planck’s constant (6.626 x 10-34 J·s) = frequency, in units of hertz (hz, sec-1) Nature of Matter Max Planck: a German physicist suggested that an object emits energy in the form of small packets of energy called quanta Quantum- the minimum amount of energy that can be gained or lost by an atom E h Planck’s constant (h): 6.626 x 10-34 J*s Nature of Matter Einstein proposed that radiation itself is really a stream of particles called photons Energy of each photon is : E photon hv also showed that energy has mass E mc 2 hc Nature of Matter c v c v E hv h mc hc mc E hc 2 E mc shows that anything with both mass and velocity has a corresponding wavelength 2 Nature of Matter In 1924, Louis de Broglie (French scientist) suggested that matter has both particle-like and wave-like characteristics h mv Main Ideas: matter and energy are not distinct energy is a form of matter larger objects are mostly particle-like smaller objects are mostly wave-like 7.3 The Atomic Spectrum of Hydrogen Spectroscopic analysis of the visible spectrum… White light …produces all of the colors in a continuous spectrum Atomic Spectra White light is made up of all the colors of the visible spectrum. Passing it through a prism separates it. • These are called the atomic emission spectrum • Unique to each element, like fingerprints! • Very useful for identifying elements Continuous Spectra White light passed through a prism produces a spectrum – colors in continuous form. The Continuous Spectrum ~ 650 nm ~ 575 nm ~ 500 nm ~ 480 nm ~ 450 nm The different colors of light correspond to different wavelengths and frequencies Continuous Emission Spectrum line-emission spectrum- series of wavelengths of light created when visible portion of light from excited atoms is shined through a prism scientists using classical theory expected atoms to be excited by whatever energy they absorbed continuous spectrum emission of continuous range of frequencies of EM radiation contains all wavelengths of visible light Spectroscopic analysis of the hydrogen spectrum… H receives a high energy spark H-H bonds Are broken and H atoms are excited …produces a “bright line” spectrum Line Spectra Light passed through a prism from an element produces a discontinuous spectrum of specific colors Hydrogen only four lines are observed Line Spectra The pattern of lines emitted by excited atoms of an element is unique = atomic emission spectrum • These are called the atomic emission spectrum • Unique to each element, like fingerprints! • Very useful for identifying elements H Line-Emission Spectrum light is emitted by excited H atoms when bond is broken in the diatomic molecule ground state- lowest energy state of an atom excited state- when an atom has higher potential energy than it has at ground state H Line-Emission Spectrum when an excited electron falls back to ground state, it emits photon of radiation the photon is equal to the difference in energy of the original and final states of electron since only certain frequencies are emitted, only certain energies are allowed for electrons in H atom Electron transitions involve jumps of definite amounts of bands This produces of light with definite energy. wavelengths. 7.4 The Bohr Model Niels Bohr (Danish physicist) in 1913 Developed a quantum model for H atom that explained the emission line spectrum Electron moves around the nucleus only in certain allowed circular orbits, in which it has a certain amount of energy The Bohr model Energy level of an electron analogous to the steps of a ladder The electron cannot exist between energy levels, just like you can’t stand between steps on a ladder A quantum of energy is the amount of energy required to move an electron from one energy level to another Niels Bohr Developed the quantum model of the hydrogen atom. He said the atom was like a solar system. The electrons were attracted to the nucleus because of opposite charges. Didn’t fall in to the nucleus because it was moving around. The Bohr Atom He didn’t know why but only certain energies were allowed. He called these allowed energies energy levels. Putting Energy into the atom moved the electron away from the nucleus. From ground state to excited state. When it returns to ground state it gives off light of a certain energy. The Model: Summary Space around nucleus is divided into spherical (circualr) paths (orbits) each has a number called “Principal Quantum number” The electron can exist only in one of these orbitals but not in between Orbits possess fixed size and energy, therefore electron has a definite energy characteristic of its orbit R Z E E Energy of electron n J R Rydberg constant 2.180x10 particle 2 H n orbit 2 -18 H En 2.180 X 10 23 18 J 6.02 X 10 particles 1kJ X X 1312kJ / mol particle 1mol 1000J 1312kJ En n mole 2 Orbits allowed for electron are those in which electron has an angular momentum= nh 1 An electron can pass only from one bit to another. Absorption or emission will occur Energy of the outermost orbit is zero The Bohr Atom n=4 n=3 n=2 n=1 Bohr Model To create an accurate model, he had to use quantum theory instead of classical created an equation used to calculate the energy levels available to electrons in a certain atom: E 2.178 10 18 2 Z J( 2 ) n where n= integer and Z=atomic number negative sign makes the energy more negative the closer it is to the nucleus Bohr Model Can gain energy by moving to a higher energy level Can lose energy by moving to lower energy level Bohr Model a photon is released that has an energy equal to the difference between the initial and final energy orbits Bohr Model equation can be used twice to find the ∆E when an electron moves energy levels E 2.178 10 E [2.178 10 18 J( Zf 2 nf 2 E 2.178 10 18 2 Z J( 2 ) n )] [2.178 10 18 1 1 J( 2 2 ) nf ni 18 2 Zi J ( 2 )] ni Bohr Model can wavelength of photon released by using hc E E=0 is set at an distance of ∞ away from the nucleus and becomes more negative as the electron comes closer to the nucleus E 2.178 10 18 1 J( ) 0 Example 1 1 1 Calculate the 18 E 2.178 10 J ( 2 2 ) energy required to nf ni move the hydrogen electron from n=1 1 1 18 E 2.178 10 J ( 2 2 ) 1.633 10 18 J to n=2. Find the 2 1 wavelength of radiation that had to m 34 (6.626 10 J s)( 2.9979 ) hc s be absorbed by the E 1.633 10 18 J electron. 9 10 nm 1.216 107 m 121.6nm 1m Calculate the energy required to remove the electron from the hydrogen atom in its ground state. Example 2 E 2.178 10 18 E 2.178 10 E 2.178 10 18 1 1 J( 2 2 ) nf ni 18 1 1 J( 2 ) 1 J (0 1) 2.178 10 18 Energy was absorbed by the electron so the value of ∆E value is positive. J The Bohr Model Doesn’t work. Only works for hydrogen atoms. Electrons don’t move in circles. The quantization of energy is right, but not because they are circling like planets. Bohr Model problems: did not work for other atoms did not explain chemical behavior of atoms Heisenberg’s Uncertainty Principle According to de Broglie: Electron behaves like a wave It is possible to specify the position of a wave at a particular instant? Energy, wavelength and amplitude can be determined But exact position is impossible to be determined The electron cannot be imagined as : moving particle In a path of the same radius (well defined orbits) Thus, location, direction and speed of motion of a particle cannot be determined Then Bohr Model had to be “Abandoned Heissenberg Uncertainty Principle “It is impossible to determine both the position and momentum of a subatomic particle (such as the electron) with arbitrarily high accuracy” The effect of this principle is to convert the laws of physics into statements about relative, instead of absolute, certainties. Heisenberg Uncertainty Principle we cannot know the exact position and momentum (motion) of the electron as more is known about position, less is known about momentum uncertainties are inversely proportional h x ( m ) 4 where ∆x: uncertainty in position ∆m : uncertainty in mometum minimum uncertainty is h/4 7.5 The Quantum Mechanical Model Exactt position of electron can not be defined Exact bath of electron about nucleus can not be defined Werner Heisenberg, Louis de Broglie and Erwin Schrodinger made the approach called “Quantum Mechanics” They assumed that the electron is a standing wave The Quantum Mechanical Model Waves are associated with electrons Information about energies of electrons and their positions are obtained from studying the associated waves Description of electron is based upon “ Probability of finding a particle within a given region of space” “ but not on the exact position” Schrödinger Equation Wave equation describing electron as being a wave The amplitudes (height), , of electron wave at various points of space are calculated commonly called “wave function” provides information about the allowable energies for an electron in H atom. corresponds to a certain energy and describes a region around nucleus “Orbital” where the electron having that energy may be found Orbital: Region around the nucleus where the electron can be expected to be found The Function 2 2 describes the probability of the position of the electron at a particular point 2 Probablity of finding a particle in a given region of space 2 Electric charge density at a given region of space Orbital A 3 dimensional space around a nucleus in which electrons are most likely to be found Shape represents electron density (not a path the electron follows) Each orbital can hold up to 2 electrons. Thus, The charge can be assumed to be spread out as a charge cloud by rapid motion of electron The cloud is denser in some regions than others The probability of finding electron in a given region in space is proportional to the density of the cloud Meaning of Wave Function the wave function itself does not have concrete meaning the square of the wave function represents the probability of finding an electron at a certain point easily represented as probability distribution where the deepness of color indicates the probability Meaning of Wave Function (a) electron density map probability of finding an electron is highest at short distances from nucleus (b) calculated probability of finding an electron at certain distances from nucleus in the 1s orbital 7.6 Quantum Numbers There are many solutions to Schroedinger’s equation for H atom Each solution is a wave function called Orbital. Each solution can be described with quantum numbers that describe some aspect of the solution. Schrödinger’s equation requires 3 quantum numbers 7.6 Quantum Numbers Quantum numbers specify the properties of atomic orbitals and of electrons in orbitals the first three numbers come from the Schrödinger equation and describe: main energy level shape orientation 4th describes state of electron 1st Quantum Number Principal Quantum Number: n Main energy level (or shell) occupied by electron. They are called atomic orbitals regions where there is a high probability of finding an electron. values are all positive integers >0 (1,2,3,…) As n increases size of orbital is larger electron has higher energy the electron’s average distance from the nucleus increases Principal Quantum Number Maximum number of electrons that can fit in an energy level: 2n2 st 1 Energy Quantum Number nd 2 Quantum Number Angular Momentum Quantum Number: l indicates the shape of the orbital (sublevel or subshell) the number of possible shapes (or l values) for an energy level is equal to n the possible values of l are 0 and all positive integers less than or equal to n -1 l has integer values from 0 to n-1 l = 0 is called s l = 1 is called p l =2 is called d l =3 is called f l =4 is called g 2nd Quantum Number s orbitals: 1: s spherical l value of 0 1st occur at n=1 2nd Quantum Number p orbitals: 3 2px, 2py, 2pz dumbbellshaped l value of 1 1st occur at n=2 for n>2, shape is same but size increases nd 2 Quantum Number d orbitals: 5: 3dxz, 3dyz, 3dxy, 3dx2-y2, dz2 mostly cloverleaf l value of 2 1st occur at n=3 for n>3, same shape but larger size nd 2 Quantum Number f orbitals: 7 types various shapes l value of 3 begin in n=4 nd 2 Quantum Number Other shapes can exist in energy levels as long as they follow the rules g (l=4) starts in 5 with 9 orbitals h (l=5) starts in 6 with 11 orbitals, etc but no known elements have electrons in them at ground state nd 2 Level Quantum Number Sublevels Sublevels 0 1 2 0 1 2 0 1 0 3 3rd Quantum Number Magnetic Quantum Number: ml indicates the orientation of an orbital around the nucleus has values from +l -l specifies the exact orbital that the electron is contained in each orbital holds maximum of 2 electrons total number of orbitals is equal to n2 for an energy level number of possible ml values for a certain subshell is equal to 2l + 1 rd 3 Quantum Number Energy Level (n) 1 2 3 4 Sublevels in Level # Orbitals in Sublevel Total # of Orbitals in Level s s 1 1 1 4 p 3 s 1 p 3 d 5 s 1 p 3 d f 5 7 9 16 th 4 Quantum Number Spin Quantum Number: ms indicates the spin state of the electron only 2 possible directions only 2 possible values: +½ and -½ paired electrons must have opposite spins maximum number of electrons in an energy level is 2n2 Quantum Numbers: type n = principle quantum no. values 1,2,3,etc meaning shell (period) l = angular momentum quant. # or azimuthal q# 0,1,2,3... s,p,d,f,g,h.. subshell ml = magnetic q. # 0,1,2,3,.. orbital ms = spin q. # ½ spin These 4 Quantum numbers give the general location of electrons within an atom and the general shape of the orbital in which they reside. Electrons Allowed All electrons in the same sublevel have the same energy. All 2s electrons have the same energy. All 2p electrons have the same energy which is slightly higher than the energy of the 2s electrons s sublevel 2 electrons p sublevel 6 electrons d sublevel 10 electrons f sublevel 14 electrons 7.9 Polyelectronic Atoms Kinetic energy - as the electrons move around the nucleus Potential energy - from their attraction to nucleus Potential energy - from their repulsion to each other Electron Correlation Problem can’t find the exact location of electrons can’t find the specific repulsions between electrons so we must treat each electron as if it has an average amount of attraction to nucleus and repulsion to other electrons Electron Shielding occurs when an electron is not attracted to the nucleus because of electrons in lower energy levels repelling it. Penetration Effect all orbitals in the same energy level do NOT have the same amount of energy ( are not degenerate) Es < E p < Ed < E f the amount of energy in each sublevel is determined by its average distance from the nucleus n=4 4 subshells f f f f f f f f (7 orbitals) d (5 orbitals) d p (3 orbitals) d px d d py s (1 orbital) d pz n = 4, l = 0, ml = 0 s n=3 3 subshells d (5 orbitals) d p (3orbitals) d px d py n = 4, l = 2, ml = +2 d d pz + s (1 orbital) n=2 2 subshells p (3 orbitals) s (1 orbital) n=1 (1st shell) 1 subshell s (1 orbital) s px py pz n = 2, l = 1, ml = -1 s subshell s subshell only holds 2 electrons Gives: "electron Address" = (n)energy , (l)shape of house (orbital), (ml) which , spin n=4 4 subshells f f f f f f f f (7 orbitals) d (5 orbitals) d p (3 orbitals) d px d d py s (1 orbital) d pz s n=3 3 subshells d (5 orbitals) d p (3orbitals) d px d py d d pz + s (1 orbital) n=2 2 subshells p (3 orbitals) s (1 orbital) n=1 (1st shell) 1 subshell s (1 orbital) l = 3 ml =-3,-2,-1,0,1,2,3 n = 4 l = 2 ml =-2,-1,0,1,2 l = 1 ml =-1,0,1 l=0 ml =0 l=2 n=3 l=1 l=0 s px py pz n=2 s subshell s subshell only holds 2 electrons Gives: "electron Address" = (n)energy , (l)shape of house (orbital), (ml) which , spin n =1 n=4 4 subshells f f f f f f f f (7 orbitals) d (5 orbitals) d p (3 orbitals) d px d d py s (1 orbital) d Pauli exclusion principle: no 2 e-’s in the same atom can have the same 4 quantum numbers. pz s n=3 3 subshells d (5 orbitals) d p (3orbitals) d px d py d d pz + s (1 orbital) n=2 2 subshells p (3 orbitals) s (1 orbital) n=1 (1st shell) 1 subshell s (1 orbital) s px py Each compartment (orbital) can only hold two electrons. pz s subshell s subshell only holds 2 electrons Gives: "electron Address" = (n)energy , (l)shape of house (orbital), (ml) which , spin Hund’s rule: each orbital of a subshell will get 1 e- of parallel spin before they are paired. n = 2, l =1, ml =+1, ms = -1/2 7.11 Electron Arrangement in Atoms The way electrons are arranged in various orbitals around the nuclei of atoms. Three rules tell us how: 1) Aufbau principle - electrons enter the lowest energy first. • This causes difficulties because of the overlap of orbitals of different energies – follow the diagram! 2) Pauli Exclusion Principle - at most 2 electrons per orbital - different spins Electron Configurations Electron configuration describes the distribution of electrons among the various orbitals in the atom The spdf notation uses numbers to designate a principal shell and the letters to identify a subshell; a superscript number indicates the number of electrons in a designated subshell Orbital Diagram An orbital diagram uses boxes to represent orbitals within subshells and arrows to represent electrons: Each box has arrows representing electron spins; opposing spins are paired together EOS Rules for Electron Configurations Electrons occupy the lowest available energy orbitals Pauli exclusion principle – no two electrons in the same atom may have the same four quantum numbers •Orbitals hold a maximum of two electrons •spins must be opposed Rules for Electron Configurations For orbitals of identical energy, electrons enter empty orbitals whenever possible – Hund’s rule Electrons in half-filled orbitals have parallel spins EOS Rules for Electron Configurations Capacities of shells (n) and subshells (l) EOS Rules for Electron Configurations Subshell filling order ... Each subshell must be filled before moving to the next level Rules for Electron Configurations Subshell filling order ... Each subshell must be filled before moving to the next level Aufbau Principle Aufbau is German for building up. As the protons are added one by one, the electrons fill up hydrogen-like orbitals. Fill up in order of energy levels. The Aufbau Principle A hypothetical building up of an atom from the one that precedes it in atomic number (Z = 1) H 1s1 (Z = 3) Li 1s22s1 [He]2s1 (Z = 2) He 1s2 Abbreviated electron configuration (Z = 3) Li 1s22s1 EOS The Aufbau Principle ... [He]2p2 [He]2p3 [He]2p4 [He]2p5 [He]2p6 EOS Increasing energy 7s 6s 5s 7p 6p 5p 4p 4s 6d 5d 4d 3d 3p 3s 2p 2s 1s aufbau diagram 5f 4f Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli To show the different direction of spin, a pair in the same orbital is written as: Quantum Numbers Each electron in an atom has a unique set of 4 quantum numbers which describe it. 1) 2) 3) 4) Principal quantum number Angular momentum quantum number Magnetic quantum number Spin quantum number Electron Configurations 3) Hund’s Rule- When electrons occupy orbitals of equal energy, they don’t pair up until they have to. Let’s write the electron configuration for Phosphorus We need to account for all 15 electrons in phosphorus Increasing energy 7s 6s 5s 7p 6p 6d 5d 5p 4d 4p 3s 2s 1s 4f 3d 4s 3p 5f The first two electrons go into the 1s orbital 2p Notice the opposite direction of the spins only 13 more to go... Increasing energy 7s 6s 5s 7p 6p 6d 5d 5p 4d 4p 5f 4f 3d 4s 3p 3s 2p 2s 1s The next electrons go into the 2s orbital only 11 more... Increasing energy 7s 6s 5s 7p 6p 5p 4p 4s 6d 5d 4d 5f 4f 3d 3p 3s 2p 2s 1s • The next electrons go into the 2p orbital • only 5 more... Increasing energy 7s 6s 5s 7p 6p 5p 4p 4s 6d 5d 4d 5f 4f 3d 3p 3s 2p 2s 1s • The next electrons go into the 3s orbital • only 3 more... Increasing energy 7s 6s 5s 4s 3s 2s 1s 7p 6p 5p 4p 6d 5d 4d 5f 4f 3d 3p • The last three electrons go into the 3p orbitals. 2p They each go into separate shapes (Hund’s) • 3 unpaired electrons • = 1s22s22p63s23p3 Orbital Diagram for A Nitrogen Atom 7N 1s 2s 2p 3s Orbital Diagram for A Fluorine Atom 9F 1s 2s 2p 3s Orbital Diagram for A Magnesium Atom 12Mg 1s 2s 2p 3s 8O 1s 2s 2p 3s Write the orbital diagram for the electrons in an iron atom 26Fe 1s 2s 2p 3d 3s 3p Orbitals fill in an order Lowest energy to higher energy. Adding electrons can change the energy of the orbital. Full orbitals are the absolute best situation. However, half filled orbitals have a lower energy, and are next best • Makes them more stable. • Changes the filling order Write the electron configurations for these elements: Titanium - 22 electrons 1s22s22p63s23p64s23d2 Vanadium - 23 electrons 1s22s22p63s23p64s23d3 Chromium - 24 electrons 1s22s22p63s23p64s23d4 (expected) But this is not what happens!! Chromium is actually: 1s22s22p63s23p64s13d5 Why? This gives us two half filled orbitals (the others are all still full) Half full is slightly lower in energy. The same principal applies to copper. Copper’s electron configuration Copper has 29 electrons so we expect: 1s22s22p63s23p64s23d9 But the actual configuration is: 1s22s22p63s23p64s13d10 This change gives one more filled orbital and one that is half filled. Remember 9 d these exceptions: 4 d, Electron Configurations of Ions Anions: gain e– to complete the valence shell Example: - EOS Electron Configurations of Ions Cations: lose e– to attain a complete valence shell Example: (Z = 11) Na (Z = 11) Na+ EOS Electron Configurations of Ions Cations formed from transition metals lose e– from the highest principal energy level (n) EOS 7.10 The history of the Periodic Table Developed independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev (1870”s). Didn’t know much about atom. Put in columns by similar properties. Predicted properties of missing elements. The aufbau principle and Periodic Table 7s 7p 7d 7f 6s 6p 6d 6f 5s 5p 5d 5f 4s 4p 4d 4f 3s 3p 3d 2s 2p 1s • 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 56 • 38 20electrons 4212 Details of the Periodic Table Elements in the same column have the same electron configuration. Elements in columns have similar properties. Similar properties because of electron configuration. Noble gases have filled energy levels. Transition metals are filling the d Valence electrons- the electrons in the outermost energy levels (not d). Core electrons- the inner electrons. Hund’s Rule- The lowest energy configuration for an atom is the one have the maximum number of of unpaired electrons in the orbital. C 1s2 2s2 2p2 Valence Electrons and Core Electrons Valence electrons are those with the highest principal quantum number Sulfur has six valence electrons EOS Valence Electrons and Core Electrons Electrons in inner shells are called core electrons Sulfur has 10 core electrons EOS Periodic Relationships We can deduce the general form of electron configurations directly from the periodic table The Periodic Table Exceptions 2 2 Ti = [Ar] 4s 3d 2 3 V = [Ar] 4s 3d 1 5 Cr = [Ar] 4s 3d 2 5 Mn = [Ar] 4s 3d Half filled orbitals. Scientists aren’t sure of why it happens 1 10 same for Cu [Ar] 4s 3d Irregular configurations of Cr and Cu Chromium steals a 4s electron to make its 3d sublevel HALF FULL Copper steals a 4s electron to FILL its 3d sublevel More exceptions 2 1 Lanthanum La: [Xe] 6s 5d 2 1 1 Cerium Ce: [Xe] 6s 4f 5d Promethium Pr: [Xe] 6s2 4f3 5d0 2 7 1 Gadolinium Gd: [Xe] 6s 4f 5d 2 14 1 Lutetium Pr: [Xe] 6s 4f 5d We’ll just pretend that all except Cu and Cr follow the rules. Main Group (Representative) and Transition Elements Elements in which the orbitals being filled in the aufbau process are either s or p orbitals of the outermost shell are called main group (Representative) elements “A” group designation on the periodic table The first 20 elements are all main group elements Transition Elements In transition elements, the sublevel (shell) being filled in the aufbau process is in an inner principal shell (d or f) d-Block transition elements: Electrons enter the d-sublevels. f-Block transition elements: d sublevels are completely filled. Electrons enter f-sublevels Lanthanides: electrons fill 4f subleve Actinides: electrons fill 5f sublevels Periodic Relationships General form of electron configurations can be deduced directly from the periodic table The Periodic Table Periodic Relationships Transition Elements Completely filled and halffilled sublevels are more energetically favorable configurations 7.12 Periodic Trends in Atomic Properties Ionization Energy Ionization energy the energy required to remove an electron form a ground state atom in the gaseous phase A (g) + energy A+ (g) + e Highest energy electron removed first. First ionization energy (I1) is that required to remove the first electron. Second ionization energy (I2) - the second electron, etc. Successive ionization Energies Continual removal of electrons increases ionization energy greatly B B+ + e– I = 801 kJ mol–1 B+ B+2 + e– I = 2427 kJ mol–1 B+2 B+3 + e– I = 3660 kJ mol–1 B+3 B+4 + e– I = 25,025 kJ mol–1 B+4 B+5 + e– I = 32,822 kJ mol–1 Ionization Energy: The energy required to completely remove an e- from an atom in its gaseous state. Mg(g) Mg1+ + eMg1+(g) Mg2+ + e- 1st ionization energy 2nd ionization energy Question: Which of the above ionizations would have the highest ionization energy and why? Trends in ionization energy for Mg • • • I1 = 735 kJ/mole I2 = 1445 kJ/mole I3 = 7730 kJ/mole The effective nuclear charge increases as electrons are removed It takes much more energy to remove a core electron than a valence electron because there is less shielding. Ionization Energy if one electron is removed, the positive charge binds the electrons more tightly so 2nd ionization energy must be higher the largest jump in energy is when you remove a core electron instead of valence Ionization Energy First ionization energies across Periods and Groups Across Periods and Groups Generally from left to right, I1 increases because there is a greater nuclear charge with the same shielding. As you go down a group I1 decreases because electrons are farther away from nucleus Ionization Energy Across Period: requires more energy to remove an electron so increases because electrons added in the same energy level do not shield electrons from nuclear charge Down Group: requires less energy to remove electron so decreases because the valence electrons are farther away from protons attracting them It is not that simple Zeff changes as you go across a period, so will I1 Half filled and filled orbitals are harder to remove electrons from. here’s what it looks like. Atomic number First Ionization energy Atomic number First Ionization energy Atomic number First Ionization energy Ionization Energy Ionization Energy Ionization Energy Electron Affinity Electron Affinity – the energy change when an electron is added to a gaseous neutral atom A + e- A- + energy Cl(g) + e- Cl-(g) E = -349 kJ/mol • Electron affinities are expressed as negative because the process is exothermic Electron affinity values Electron Affinity Electron Affinity Across Period: releases more energy so number increases (gets more negative) because electrons added in the same energy level do not shield electrons from nuclear charge Down Group: releases less energy so number decreases (gets less negative) because the electrons being added are farther away from the attracting protons Electron Affinity Atomic Radii Size of orbitals can not be specified exactly, neither can the size of atom Atomic Radii – half the distance between the nuclei of identical atoms that are bonded together Atomic Radii Across Period: atoms get smaller because of the increased number of protons attracting the electrons the electrons added in the same energy level do not shield electrons from nuclear charge Down Group: atoms get larger increases because the energy levels being added to the atom Atomic Radii Properties Ionic Radii The ionic radius of each ion is the portion of the distance between the nuclei occupied by that ion Ionic Radii Cations are smaller than the atoms from which they are formed – the nucleus attracts the remaining electrons more strongly Anions are larger than the atoms from which they are formed – the greater number of electrons repel more strongly EOS Isoelectronic Configurations • Elements that all have the same number of electrons For isoelectronic species, the greater the nuclear charge, the smaller the species Effective nuclear charge Atomic and Ionic Radii Atomic/Ionic Radii Summary of Periodic Trends 7.13 Properties of a group: Alkali metals Information contained in the Periodic Table I Groups n of representative elements exhibit similar chemical f properties that change in a regular way Each o group member has same valence electron configuration r It is the number and type of valence electrons that m an atoms chemistry. determine a get the electron configuration from the You can Periodic t Table Metals i lose electrons have the lowest IE Non metalsgain electrons most negative electron o affinities. n Various groups have special names Special Names for Groups in the Periodic Table Metals, Nonmetals, and Metalloids Metals have a small number of electrons in their valence shells and tend to form positive ions Except for hydrogen and helium, all s-block elements are metals All d- and f-block elements are metals EOS Metals, Nonmetals, and Metalloids Atoms of a nonmetal generally have larger numbers of electrons in their valence shell than do metals, and many tend to form negative ions Nonmetals are all pblock elements and include hydrogen and helium Metalloids have properties of both metals and nonmetals EOS The Alkali Metals Do not include hydrogen- it behaves as a nonmetal On going down the group there is : decrease in IE increase in radius Increase in density decrease in melting point They loose their valence electron readily. Thus they behave as reducing agents They react readily with nonmetals Reducing ability Lower IE< better reducing agents Cs>Rb>K>Na>Li works for solids, but not in aqueous solutions. In solution Li>K>Na Why? It’s the water -there is an energy change associated with dissolving Potassium Reacts Violently with Water Hydration Energy It is exothermic for Li+ -510 kJ/mol for Na+ -402 kJ/mol for K+ -314 kJ/mol Li is so big because of it has a high charge density, a lot of charge on a small atom. Li loses its electron more easily because of this in aqueous solutions The reaction with water Na and K react explosively with water Li doesn’t !!! Even though the reaction of Li has a more negative H than that of Na and K Na and K melt easily compared to Li. Thus they will have greater contact with water than Li H does not tell you speed of reaction