Chapter 7 Electron Configurations & the Periodic Table General Chemistry I S. Imbriglio Part A: Electron Configurations • The arrangement of all of the electrons in an atom is called the electron configuration • Electron configurations can be used to explain: – Reactivity & properties of the elements – Trends in reactivity & properties (periodic table!) • The electron configuration of an atom is best investigated using electromagnetic radiation A. Electromagnetic (EM) Radiation 1. Electromagnetic (EM) Waves: oscillating perpendicular magnetic & electric fields that travel through space at the same rate (the “speed of light”: c = 3.00x108 m/s) - Unlike sound waves, electromagnetic waves require no medium for propagation eg. This allows the sun’s electromagnetic radiation to reach the earth as sunlight. http://ww2.unime.it/dipart/i_fismed/wbt/mirror/ntnujava/emWave/emWave.html a) Wavelength & Frequency • All EM waves can be described in terms of wavelength & frequency. – Wavelength ( - lambda): distance between adjacent crests (or troughs) in a wave a) Wavelength & Frequency - Frequency ( - nu): the number of complete waves passing a point in a given period of time (remember c = 3.00x108 m/s) Unit of frequency is the Hertz (Hz): 1 Hz = 1 s-1 = 1 “per second” a) Wavelength & Frequency • For EM radiation, frequency is related to wavelength by = c. • If you know one, you know the other. Calculate the frequency of an X-ray that has a wavelength of 8.21 nm. = c i. Electromagnetic Spectrum • The type of electromagnetic radiation is defined by its frequency & wavelength • Remember, as increases, decreases, and vice versa. b) Amplitude • The intensity of radiation is related to its amplitude. – Amplitude: height of the wave crest – In the visible portion of the spectrum, brighter light is light with a greater amplitude. c) Refraction • Classifying EM radiation (light) as a wave explains many fundamental properties of light. • Refraction: – When white light passes through a narrow slit & then through a glass prism, the light separates in a continuous spectrum. – The spectrum is continuous because each color merges into the next without a break – all wavelengths (or frequencies) of visible light are observed. d) Diffraction • Diffraction: Waves can add constructively or destructively to amplify or cancel each other. e) Black-Body Radiation • At high temperatures, matter emits electromagnetic radiation. • As the temperature increases, the maximum intensity of the emitted radiation increases in frequency. • The observed spectrum depends only on temperature & not on the particular elements present. b) Black-Body Radiation • THE EXPLANATION: According to classical physics, as the temperature of a solid increases, the atoms vibrate more vigorously – some of the vibrational energy is released as EM radiation. • THE PROBLEM: Using the classical picture of light as a wave, scientists were unable to explain the shape of the observed black-body radiation spectra. 3. Planck’s Quantum Hypothesis… THE ANSWER? • According to classical physics, the energy scale is continuous – there are no limitations on the amount of energy a system can gain or lose. • Planck proposed that variations in energy are discontinuous – energy changes occur only by discrete amounts. eg. The “Quantization” of Elevation Quantized (1 step = 1 quantum) Classical (continuous) 3. Planck’s Quantum Theory • For electromagnetic radiation of a certain frequency, the smallest amount of energy, called a quantum, is defined by the relationship: E = h (h = Planck’s constant = 6.626x10-34 Js) Energy can be absorbed or emitted only as a quantum, or some whole-number multiple of a quantum. 3. Planck’s Quantum Theory • According to Planck’s theory, the energy of one quantum of EM radiation is dependent on the frequency (and wavelength) of the radiation: E = h = hc/ • The energy per quantum increases as the frequency gets higher & the wavelength gets shorter. 3. Planck’s Quantum Theory • The energy per quantum increases as the frequency gets higher & the wavelength gets shorter. E = h = hc/ • Which has more energy – a quantum of microwave radiation ( = 1x10-2 m) or a quantum of infrared radiation ( = 1x10-6 m)? 3. Planck’s Quantum Theory • Planck proposed that vibrating atoms in a heated solid can absorb and emit EM radiation only in certain discrete amounts. • Planck’s quantum theory allowed him to successfully explain black-body radiation spectra, but his radical assertion that “energy is quantized” was difficult for the scientific community to accept. • Fortunately, five years after its inception, Einstein used Planck’s Quantum Theory to explain another well-known phenomenon called the photoelectric effect. 4. The Photoelectric Effect • Certain metals exhibit a photoelectric effect – when illuminated by light of certain wavelengths (photo-), they emit electrons (-electric). • In order for the photoelectric effect to occur, the frequency of the light must be higher than a certain minimum value – called the threshold frequency. • Each photosensitive metal has a different threshold frequency. 4. The Photoelectric Effect • When light of a high enough energy (frequency) is used, the number of electrons ejected is proportional to the intensity of the light. • Light below the threshold frequency will not cause an electric current to flow – no matter how bright (intense) the light is. eg. Light meters use the photoelectric effect to measure the intensity (brightness) of light. http://jchemed.chem.wisc.edu/JCEDLib/WebWare/collection/open/JCEWWOR006/peeffect5.html a) Photons • Classical physics could not explain the existence of a threshold frequency, so Einstein turned to Planck’s Quantum Theory. • Einstein defined a quantum of electromagnetic radiation as a photon. • Einstein proposed that light could be thought of as a stream of photons with particle-like properties as well as wave properties. • For light of frequency : Ephoton = h = hc/ a) Photons • The photoelectric effect can be explained by assuming that light has particle-like properties: – Removing one electron from a photosensitive metal requires a certain minimum energy (Emin). – Each photon has an energy given by E = h. – Only photons with E > Emin have enough energy to knock an electron loose. – Photons of lower frequency (lower energy) do not have enough energy to knock an electron loose. a) Photons • If the intensity of light is proportional to the number of photons, then more intense light means more photons. • If each photon ejects an electron, then more photons means more electrons ejected. • The number of electrons ejected is proportional to the intensity of light. The Photoelectric Effect…Explained b) Wave-Particle Duality of Light • Depending on the circumstances, light (all EM radiation) can appear to have either wave-like or particle-like characteristics. • Both ideas are needed to fully explain light’s behavior in different phenomena. “It not only prohibits the killing of two birds with one stone, but also the killing of one bird with two stones.” - James Jeans, commenting on Einstein’s explanation of the photoelectric effect Nobel Prize Winners • Max Planck won the Nobel Prize for Physics in 1918 for his quantum theory. Blackbody radiation spectra explained • Albert Einstein won the Nobel Prize for Physics in 1921 for his theory on the quantized nature of light and how it relates to light’s interaction with matter (not for his theory of relativity!). Photoelectric effect explained 5. Line Emission Spectra • In the 1920s, another phenomenon was left unexplained by classical physics – the observance of atomic line emission spectra. • When a voltage is applied to a gaseous element at low pressure, the atoms absorb energy & become “excited.” • The “excited” atoms then emit the extra energy as EM radiation. 5. Line Emission Spectra • When this radiation is passed through a prism, a limited number of discrete colored lines are seen – a discontinuous spectrum. • This discontinuous spectrum is called a line spectrum, or a line emission spectrum. • Unlike black-body radiation, each element has a unique line emission spectrum. Why don’t these atoms emit continuous spectra? B. Bohr’s Hydrogen Atom: A Planetary Model • Classical physics could not explain the presence of line emission spectra. • Not long after Einstein used quantum theory to explain the photoelectric effect, Niels Bohr used quantum theory to explain the behavior of the electron in a hydrogen atom. • Bohr’s model provided the first explanation of the discontinuous line emission spectrum of hydrogen. B. Bohr’s Hydrogen Atom: A Planetary Model • Bohr assumed that the single electron in a hydrogen atom moves around the nucleus in a circular orbit. • Bohr applied quantum theory to his model by proposing that the electron is restricted to circling the nucleus in orbits of certain radii, each of which corresponds to a specific energy. • Thus, the energy of the electron is quantized, and the electron is restricted to certain energy levels – orbits. B. Bohr’s Hydrogen Atom: A Planetary Model B. Bohr’s Hydrogen Atom 1. Energy Levels (Orbits): – Each allowed orbit is assigned a principal quantum number (n = 1,2,3,…). – The energy of the electron and the radius of its orbit increase as the value of n increases. – An atom with its electron in the lowest energy level is said to be in the ground state. 1. Energy Levels (Orbits) En = _ 2.179x10-18 J n2 (n = 1, 2, 3, …) • The allowed energies of an electron (orbit) in a hydrogen atom are restricted by the principal quantum number (n), according to the equation above. • The negative sign is a result of Bohr’s choice to define En = 0 when n = . a) Excited State vs. Ground State • Transitions Between Levels: Electrons can move from one energy level to another – An electron must absorb energy to transition from a lower energy level to a higher energy level – Energy is emitted when an electron transitions from a higher energy level to a lower energy level • When an electron absorbs energy and moves to a higher energy level, that atom is said to be in an excited state. http://www.upscale.utoronto.ca/GeneralInterest/Harrison/BohrModel/Flash/BohrModel.html a) Excited State vs. Ground State • Absorb energy to move to a higher energy orbit. • Emit energy to move to a lower energy orbit. a) Excited State vs. Ground State • When an “excited” electron returns to the ground state, energy is emitted as a photon with an energy corresponding to the difference in energy between the two levels. • In the Bohr model, n= is the excited state in which enough energy has been added to completely separate the electron from the proton – Bohr arbitrarily assigned this state as having E = 0 (hence the negative energy values). 2. Explanation of Line Spectra • Bohr’s model of the hydrogen atom can be used to explain the line emission spectrum of hydrogen: E = Efinal - Einitial |E| = h 2. Explanation of Line Spectra • Using Bohr’s equation for allowed energies in a hydrogen atom: E = Ef - Ei E = _ 2.179x10-18 J nf2 E = 2.179x10-18 J x _ _ 2.179x10-18 J ni2 1 _ 1 ni2 nf2 Only certain energies of light (E ) can be absorbed or emitted by electrons in a hydrogen atom. 2. Explanation of Line Spectra • Now, coupling that equation with E = h allows us to describe the frequencies of light that can be absorbed or emitted by an electron in a hydrogen atom. E = h where E = |E| • The frequencies () determined by this equation correlate with the frequencies of light observed in the line emission spectrum of hydrogen. • The discrete lines in the line emission spectra correspond to photons of specific frequencies that are emitted when electrons relax from higher energy levels to lower energy levels Using Bohr’s model, calculate the frequency of the radiation released by the transition of an electron in a hydrogen atom from the n = 5 level to the n = 3 level. Using Bohr’s model, calculate the wavelength of the radiation absorbed by a hydrogen atom when the electron undergoes a transition from the n = 4 to n = 5 level. C. Quantum Mechanical Model of the Atom • By the early 1920s, the theory of the WaveParticle Duality of light had been accepted, but a young scientist named Louis De Broglie was ready to shock the scientific community with another hypothesis. • De Broglie proposed that matter can exhibit wave-like properties. eg. Electrons exhibit diffraction similar to that observed with light. 1. De Broglie: Matter as Waves • De Broglie proposed that a particle of mass m moving at speed v will have a wave nature consistent with a wavelength given by the equation: = h/mv • Large (macroscale) objects have wavelengths too short to observe. • Small (nanoscale) objects have longer & more readily observable wavelengths. a) Quantum Mechanics • Current ideas about atomic structure are based on De Broglie’s theory. • The treatment of atomic structure using the wave-like properties of the electron is called quantum mechanics (or wave mechanics) • In contrast to Bohr’s precise atomic orbits, quantum mechanics provides a “less certain” picture of the hydrogen atom. b) Wave Equation & Wave Functions • In 1926, Erwin Schrödinger used De Broglie’s theory to develop an equation (Schrödinger’s wave equation) describing the locations & energies of the electron in a hydrogen atom. • Acceptable solutions to Schrödinger’s wave equation are called wave functions (). • Unlike Bohr’s model, these wave functions do not describe the exact location of an electron. b) Wave Equation & Wave Functions • The square of a wave function (2) gives the probability of finding an electron in a particular infinitesimally small volume of space in an atom. • Because we are treating electrons as waves (not particles) we cannot pinpoint the specific location of an electron. • Instead, mathematical solutions to the wave functions give 3-dimensional shapes (orbitals) within which electrons can usually be found. b) Wave Equation & Wave Functions • These 3-D orbitals (probability clouds) take the place of Bohr’s simple well-defined orbits in the modern model of the atom. We don’t know exactly where the electrons are. • This “less certain” model is justified by an important principle of science established in 1927. 2. Heisenberg’s Uncertainty Principle • It is impossible to determine the exact location and the exact momentum of a tiny particle like an electron. – The very act of measurement would affect the position and momentum of the electron because of its very small size and mass. – The collision of an electron with a high-energy photon (required to locate the electron) would change the momentum of the electron. – The collision of an electron with a low-energy photon would not provide much information about the location of the electron. 2. Heisenberg’s Uncertainty Principle • A macroscale analogy… High Shutter Speed Low Shutter Speed Can judge location, but not speed. Can judge speed, But not location D. Quantum Numbers & Atomic Orbitals • According to quantum mechanics, each electron in an atom can be described using four quantum numbers: – – – – n l ml ms Principal Quantum Number Angular Momentum Quantum Number Magnetic Quantum Number Electron Spin Quantum Number – The first three numbers describe the atomic orbital in which the electron resides & the fourth differentiates electrons that are in the same atomic orbital. 1. Principal Quantum Number (n) • The principal quantum number (n) has only integer values, starting with 1: n = 1, 2, 3, 4, . . . a) The value of n corresponds to the Principal Electron Shell that the orbital is in. b) The principal electron shell is the major factor in determining the energy of the electron(s) in that orbital – a higher n value means a higher energy. 2. Angular Momentum Quantum Number (l) • The angular momentum quantum number (l ) is an integer that ranges from zero to a maximum of n – 1: l = 0, 1, 2, 3, . . . (n – 1) a) The value of l indicates the subshell that the orbital is in (within the larger energy shell). n = 1; l=0 (1 subshell) n = 2; l = 0 or 1 (2 subshells) n = 3; l = 0, 1 or 2 (3 subshells) n = 4; l = 0, 1, 2 or 3 (4 subshells) 2. Angular Momentum Quantum Number (l) • Each subshell (l) is designated with a letter: b) Each letter (s, p, d, f) symbolizes a subshell containing one specific type of orbital with a unique shape. eg. All s orbitals are spherical (l = 0) & all p orbitals are shaped like dumbbells (l = 1) – more on this in a minute. s orbital p orbital 2. Angular Momentum Quantum Number In the third principle shell, there is one s subshell containing one s orbital, one p subshell containing three p orbitals & one d subshell containing five d orbitals.. In the second principle shell, there is one s subshell containing one s orbital & one p subshell containing three p orbitals. In the first principle shell, there is one s subshell containing one s orbital. Within a p or d subshell, how do you distinguish between the individual orbitals? 3. Magnetic Quantum Number (ml ) • The magnetic quantum number (ml) can have any integer value between l and - l, including zero: ml = l, . . . , +1, 0, -1, . . . , - l a) The magnetic quantum number (ml) is related to the directional orientation of the orbital. eg. There are three possible p orbitals – each pointing along a different axis in space. 3. Magnetic Quantum Number (ml ) (s) l = o; (p) l = 1; (d) l = 2; (f) l = 3; ml ml ml ml =0 = -1,0,1 = -2,-1,0,1,2 = -3,-2,-1,0,1,2,3 eg. There is only one type of directional orientation for any given s orbital in an l = 0 subshell because ml must equal 0. (1 s orbital) (3 p orbitals) (5 d orbitals) (7 f orbitals) 3. Magnetic Quantum Number (ml ) • There are three different p orbitals in every l = 1 subshell because ml = -1,0,1. Each of the three p orbitals is pointed along a different axis (x,y,z). 3. Magnetic Quantum Number (ml ) • There are five different d orbitals in every l = 2 subshell because ml = -2,-1,0,1,2. Four of the five d orbitals are pointed along a different axis. The fifth has a slightly different shape. 4. Shells (n), Subshells (l ) & Orbitals (ml ): A Summary 4. Shells (n), Subshells (l ) & Orbitals (ml ): A Summary This picture shows all of the orbitals in the first three electron shells (n = 1,2,3). State whether an electron can be described by each of the following sets of quantum number. If a set is not possible, state why not. a) n = 2, l = 1, ml = -1 b) n = 1, l = 1, ml = +1 c) n = 4, l = 3, ml = +3 d) n = 3, l = 1, ml = -3 Replace the question marks by suitable responses in the following quantum number assignments. a) n = 3, l = 1, ml = ? b) n = 4, l = ?, ml = -2 c) n = ?, l = 3, ml = ? Provide the three quantum numbers describing each of the three p orbitals in the 2p subshell. n 2px 2py 2pz l ml 5. Electron Spin Quantum Number (ms) • The first three quantum numbers (n, l, ml) fully characterize all of the orbitals in an atom. • But, one more quantum number is necessary to describe all of the electrons in an atom. • This is because every orbital can hold two electrons. 5. Electron Spin Quantum Number (ms) • The spin quantum number (ms) can have just one of two values (+1/2 & -1/2). • Each electron exists in one of two possible spin states. - The “spinning” electron induces an external magnetic field. Opposite spins induce opposing magnetic fields. 5. Electron Spin Quantum Number (ms) • When two electrons have the same ms quantum number, those spins are said to be parallel. • When two electrons in the same orbital have different ms quantum numbers, those electrons are said to be paired. Parallel spins Paired spins a) Pauli Exclusion Principle • The Pauli Exclusion Principle states that no more than two electrons can be assigned to the same orbital in an atom & those two electrons must have opposite spins. • In other words: – No two electrons in the same atom can have the same set of four quantum numbers (n, l, ml, ms). – If two electrons occupy the same orbital, their spins must be paired (+1/2 & -1/2). Quantum Numbers: A Macroscale Analogy • n • l - indicates which train (shell) - indicates which car (subshell) • ml - indicates which row (orbital) • ms - indicates which seat (spin) No two people can have exactly the same ticket (sit in the same seat). For n = 1, determine the possible values of l. For each value of l, assign the appropriate letter designation & determine the possible values of ml. n=1 How many orbitals in shell n = 1? How many electrons possible? For n = 2, determine the possible values of l. For each value of l, assign the appropriate letter designation & determine the possible values of ml. n=2 How many orbitals in shell n = 2? How many electrons possible? For n = 3, determine the possible values of l. For each value of l, assign the appropriate letter designation & determine the possible values of ml. n=3 # of Orbitals? # of Electrons? For n = 4, determine the possible values of l. For each value of l, assign the appropriate letter designation & determine the possible values of ml. Provide the four quantum numbers describing each of the two electrons in the 3s orbital. n l ml ms E. Electron Configurations • The electron configuration of an atom is the complete description of the orbitals occupied by all of its electrons – eg. The electron in a ground state hydrogen atom occupies the 1s orbital • There are several ways to represent electron configurations. . . 1. Representations of Electron Configuration • In most cases, it is sufficient to write a list of all of the occupied subshells and indicate the number of electrons in each subshell with a superscript. H 1s1 C 1s2 2s2 2p2 Ar 1s2 2s2 2p6 3s2 3p6 1. Representations of Electron Configuration a) Expanded Electron Configuration: In some cases, it is more informative to write a list of each occupied orbital and indicate the number of electrons in each orbital. N 1s2 2s2 2p3 versus N 1s2 2s2 2p1 2p1 2p1 The expanded configuration indicates that there is one electron in each of the three 2p orbitals – the original configuration doesn’t. 1. Representations of Electron Configuration b) An orbital box diagram goes one step further by also illustrating the spins of the elctrons. P 1s2 2s2 2p2 2p2 2p2 3s2 3p1 3p1 3p1 P 1s 2s 2p 3s 3p The orbital box diagram indicates that the three electrons in the 3p subshell all have parallel (unpaired) spins. i) Hund’s Rule • In the last example, we saw that: – Atoms can have half-filled orbitals – the electrons in the half-filled orbitals tend to have parallel spins • Hund’s Rule: The most stable arrangement of electrons in the same subshell has the maximum number of unpaired electrons, all with the same spin • In other words, electrons pair only after each orbital in a subshell is occupied. Write the expanded electron configuration and the box orbital diagram for oxygen (1s2 2s2 2p4). O O O 1s2 2s2 2p4 Write the expanded electron configuration and the box orbital diagram for boron (1s2 2s2 2p1). B B B 1s2 2s2 2p1 1. Representations of Electron Configuration c) When you get deeper into the periodic table, electron configurations can be abbreviated by using noble gas notation. - The noble gases are the elements in group 8A (He, Ne, Ar, Kr, Xe, Rn) - Each noble gas has a filled outer subshell (enough electrons to fill its highest energy subshell) c) Noble Gas Notation • Electron Configurations of Noble Gases [He] [Ne] [Ar] [Kr] = 1s2 = 1s2 2s2 2p6 = 1s2 2s2 2p6 3s2 3p6 = 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 • To use noble gas notation, write the symbol for the preceding noble gas [in brackets] to represent all of the electrons in its electron configuration. • Add the rest of the electrons at the end. c) Noble Gas Notation • Write the following electron configurations using noble gas notation: O 1s2 2s2 2p4 Si 1s2 2s2 2p6 3s2 3p2 Now we know how to write electron configurations. How do we know what the ground state electron configuration for an element is??? 2. Ground State Configuration • Afbau Principle: Every atom has an infinite number of possible electron configurations. • For an atom in its ground state, electrons are found in the energy shells, subshells & orbitals that produce the lowest energy for the atom. • Other configurations correspond to excited states. 2. Ground State Configuration • In other words, when deciding where to “put” the electrons in the ground state, always start filling the lowest energy orbitals first. • In general: – Orbital energy increases as n increases – Within the same shell (n), orbital energy increases as l increases (E: s<p<d<f) a) Order of Subshell Filling • The electron configurations of the first ten elements illustrate this point. a) Order of Subshell Filling • In general, subshells are filled in order of increasing n + l value • If two orbitals have the same value for n + l, fill the subshell with lowest n value first a) Order of Subshell Filling i) Using the Periodic Table • You don’t have to memorize the order of the subshells, just use the periodic table! • Start at H & move through the table in order until the desired element is reached. Notice: (n – 1)d orbitals are filled after ns and before np orbitals. a) Order of Subshell Filling i) Using the Periodic Table • Write the electron configuration for Al. Al Al Ne Al a) Order of Subshell Filling i) Using the Periodic Table • Write the electron configuration for As. As As Ar As a) Order of Subshell Filling i) Using the Periodic Table Write the electron configuration for Sn. Sn Sn Kr Sn ii) Transition Metals • Remember, (n – 1)d orbitals are generally filled after ns orbitals and before np orbitals. • There are some exceptions: – When it is possible to half-fill or fill the (n-1)d shell, the ns subshell can be left half-filled – This is an example of Hund’s Rule. The ns and (n-1)d orbitals are very close in energy, so the more parallel spins, the better. ii) Transition Metals 4s Sc [Ar]3d14s2 - 4s filled before 3d 3d 4s Ti [Ar]3d24s2 3d 4s V [Ar]3d34s2 3d ii) Transition Metals Cr might expect… 4s [Ar]3d44s2 3d Physical properties indicate that this is not the electron configuration. It is actually… 4s Cr [Ar]3d54s1 3d Notice the 3d subshell is half-filled. This configuration maximizes unpaired electrons - Hund’s Rule. ii) Transition Metals Having a filled subshell is also energetically favorable, so copper has an unexpected configuration… 4s Cu [Ar]3d104s1 3d The energetic stability gained from having either a filled or a half-filled subshell has an effect on the reactivity of different elements. iii) Magnetic Properties • The electron configuration of an atom determines its magnetic properties. • In atoms (or ions) with completely filled shells, all of the electron spins are paired, so their individual magnetic fields effectively cancel each other out. • Such substances are called diamagnetic. iii) Magnetic Properties • Atoms (or ions) with unpaired electrons (parallel spins) are attracted to a magnetic field. – More unpaired electrons, stronger attraction. • Such substances are called paramagnetic. eg. Metallic nickel is paramagnetic: 4s Ni 3d iii) Magnetic Properties • Ferromagnetic substances are permanent magnets. – Spins of electrons in a cluster of atoms are aligned in same direction, regardless of external magnetic field – Metals in the iron, cobalt & nickel groups exhibit ferromagnetism iv) Valence Electrons • The atomic electron configuration of an element determines the chemical reactivity of that element, but it is not the total number of electrons that is important. • If that were the case, each element would have unique reactivity & we would not observe periodicity in atomic trends and reactivity. • How do we explain the trends in the periodic table? Valence Electrons! iv) Valence Electrons • When considering the principal electron shells (n = 1,2,3,…), there are two types of electrons: – Core Electrons: electrons in the filled “inner” shell(s) of an atom – Valence Electrons: electrons in the unfilled “outer” shell of an atom • All elements in the same group have similar chemical properties because they have the same number of valence electrons in their outer shell! iv) Valence Electrons • For elements in the first three periods: – The core electrons are those in the preceding noble gas configuration. – The additional electrons in the outer shell are the valence electrons. eg. B 1s2 2s2 2p1 B [He]2s2 2p1 Core: 1s2 (Shell with n = 1) Valence: 2s2 2p1 (Shell with n = 2) iv) Valence Electrons Cl Cl Core: Valence: iv) Valence Electrons • For elements in the fourth period and below in groups 3A – 7A, the filled d subshells are also part of the core, even though they are not included in the noble gas configuration. Se Se Core: Valence: iv) Valence Electrons • In each A group, the number of valence electrons is equal to the group number. 1 8 2 # of valence electrons 3 4 5 6 7 p-block s-block v) Lewis Dot Symbols • The number of valence electrons in an atom is directly related to its reactivity. • Gilbert Lewis came up with a way to represent an element & its valence electrons. One dot equals one valence electron. 3. Ion Electron Configurations • The number of valence electrons determines the type of cation (+) or anion (-) an atom will form. • When s- and p-block elements form ions, electrons are removed or added such that a noble gas configuration is achieved. • The ions are said to be isoelectronic with the noble gas – more on this later. • In general: – Metals lose electrons; form cations – Non-metals gain electron; form anions a) Cations Li Li+ 1s22s1 1s2 = [He] (loses an electron) - Group 1 metals form cations with +1 charge. Mg 1s22s22p63s2 (loses two electrons) Mg2+ 1s22s22p6 = [Ne] - Group 2 metals form cations with +2 charges. b) Anions O 1s22s22p4 (gains two electrons) O2- 1s22s22p6 = [Ne] - Group 6 elements form anions with -2 charge. F F- 1s22s22p5 (gains one electron) 1s22s22p6 = [Ne] - Group 7 elements (halogens) form anions with -1 charge. Periodic Trends • When Mendelev created the first periodic table, he organized the elements based on similarities in chemical properties & reactivity. • Now, we can use the electron configurations of the elements to explain the trends in the periodic table. – Atomic & Ionic Radii – Ionization Energy – Electron Affinity F. Atomic Radii • The atomic radius of an atom is defined as one-half of the internuclear distance between two of the same atoms in a simple diatomic molecule. - In this simplified picture, we assume that each atom is spherical & the radius is the distance from the center to the edge of the sphere. 1. Trends: Atomic Radii • The size (radius) of an atom is determined by two main factors: a) Principal Quantum Number (n): the larger the principal quantum number (n), the larger the orbitals - As you move down a group in the periodic table, the atomic radii of the atoms increase because n increases. 1. Trends: Atomic Radii b) Effective Nuclear Charge (Z*): the nuclear positive charge experienced by outer-shell electrons in a many-electron atom – Outer-shell electrons are shielded from the full nuclear positive charge (Z) by the inner-shell electrons (electron-electron repulsion) – The effective nuclear charge (Z*) felt by an outer-shell electron is less than the actual charge of the nucleus (Z). 1. Trends: Atomic Radii b) Effective Nuclear Charge: • As Z* increases, the outer electrons are pulled closer to the nucleus & the atomic radius decreases. • Z* increases across a period in the periodic table (additional attraction to nucleus stronger than electron-electron repulsion/shielding). • Atomic Radius decreases across a period in the periodic table. 1. Trends: Atomic Radii Increasing Atomic Radius Main Group Elements Decreasing Atomic Radius G. Ionic Radii Periodic trends in ionic radii parallel the trends in atomic radii within the same group. Ionic radii increase as you move down a group in the periodic table (n increases). G. Ionic Radii 1. Cations: The radius of a cation is always smaller than that of the atom from which it is derived. – Nuclear charge (Z) remains the same. – Electron-electron repulsion (shielding) decreases. – Effective nuclear charge (Z*) increases. G. Ionic Radii 2. Anions: The radius of an anion is always larger than that of the atom from which it is derived. – Nuclear charge (Z) remains the same. – Electron-electron repulsion (shielding) increases. – Effective nuclear charge (Z*) decreases. a) Isoelectronic Ions • Atoms or ions with identical electron configurations are said to be isoelectronic. • In general: – Anions in a given period are isoelectronic with the noble gas in the same period. – Cations in a given period are isoelectronic with the noble gas in the preceding period. a) Isoelectronic Ions a) Isoelectronic Ions • When comparing isoelectronic ions, the radii depend on the number of protons in the nucleus (Z) • The more protons in the nucleus, the smaller the radius. a) Isoelectronic Ions Isoelectronic Ions O2- F- Na+ Mg2+ Ionic Radius (pm) 140 133 102 66 Number of Protons 8 9 11 12 Number of Electrons 10 10 10 10 Rank the following series in order of increasing atomic or ionic radii. (1 = smallest, 3 = largest) Series 1: Al O P Series 2: Be+2 Mg+2 Sr+2 Series 3: O-2 F- Na+ H. Ionization Energies • The first ionization energy of an atom is the minimum energy required to remove the highest energy (outermost) electron from the neutral atom in the gas phase. • The larger the I.E., the harder it is to remove the electron. eg. The first ionization energy of lithium is illustrated by the following equation Li (g) Li+ (g) + e1s22s1 1s2 E = 520 kJ/mol 1. Trends: Ionization Energies • Ionization energies tend to decrease as you move down a group in the periodic table. – Size - it is easier to remove an electron that is further from the nucleus. • Ionization energies tend to increase as you move across a period in the periodic table. – Effective Nuclear Charge – As Z* increases, it becomes harder to remove an electron. 1. Trends: Ionization Energies 1. Trends: Ionization Energies • Ionization energies do not increase smoothly across the periods in the periodic table. In general: - It is easier to remove an electron if it results in the formation of a filled or half-filled subshell. - It is harder to remove an electron from a filled or half-filled subshell. 1. Trends: Ionization Energies eg. In contrast to periodic trends, the ionization energy of nitrogen is higher than the ionization energy of oxygen. N N+ O O+ 1. Trends: Ionization Energies In contrast to periodic trends, the ionization energy of beryllium is higher than that of boron. Why? Be Be+ B B+ 2. Subsequent Ionization Energies • The first, second & third ionization energies are the energies associated with removing the first, second & third highest energy electrons in an atom. • Ionization energies increase with each successive electron removed because Z* increases (same number of protons, less electron repulsion). 2. Subsequent Ionization Energies • Ionization energies can be used to explain why Li forms Li+ cations and Be forms Be2+ cations. I. Electron Affinities • The electron affinity of an element is the energy change resulting from an electron being added to an atom to form a 1- anion. – Electron affinity is the measure of the attraction an atom has for an additional electron. eg. The electron affinity of fluorine is illustrated by: F (g) + e- F- (g) E = EA = -328 kJ/mol 1s22s22p5 1s22s22p5 Fluorine readily accepts an electron to gain a stable noble gas configuration (filled n = 2 shell). I. Electron Affinities • Most electron affinities are <0 (favorable). • Electron affinities are generally 0 (unfavorable) for atoms with filled subshells (Groups 2A & 8A). Knowing the Trends… • You should be able to explain the trends & use the periodic table to predict relative magnitudes for the following properties: – Atomic Radius – Ionic Radius – Ionization Energy • The trends in electron affinities are less regular, but you should be able to explain differences in EAs based on filled/unfilled subshells.