CHM012-1ST TOPIC.pdf

CHM012 Chemistry TOPIC 1for Atomic Structure Engineers and Periodicity BILLACURA, M.D.  Electromagnetic Radiation Electromagnetic radiation (EMR) – is one of the ways that energy travel through space. The visible light is one type of EMR. Class...

CHM012 Chemistry TOPIC 1for Atomic Structure Engineers and Periodicity BILLACURA, M.D.  Electromagnetic Radiation Electromagnetic radiation (EMR) – is one of the ways that energy travel through space. The visible light is one type of EMR. Classification of electromagnetic spectrum. BILLACURA, M.D.  Electromagnetic Radiation Characteristics of Electromagnetic radiation (EMR) Relationship of wavelength and frequency: BILLACURA, M.D.  Example The brilliant red colors seen in fireworks are due to the emission of light with wavelengths around 650 nm when strontium salts such as Sr(NO3)2 and SrCO3 are heated. Calculate the frequency of red light of wavelength 6.50 x 102 nm. Practice Exercise The laser in an audio CD player uses light with a wavelength of 7.80x102 nm. Calculate the frequency of this light. BILLACURA, M.D.  Nature of Matter Max Planck found that energy can be gained or lost only in whole-number multiples of hv. Energy was found to be quantized, wherein a where h is called system can transfer energy in whole quanta or Planck’s constant, “packets”. Thus energy has a particle-like 6.626x10-34 Js properties. Einstein suggested that electromagnetic radiation can be viewed as a stream of “particles” called Photons. Where the energy of a photon is: BILLACURA, M.D.  Example The blue color in fireworks is often achieved by heating copper (I) chloride (CuCl) to about 1200oC. Then the compound emits blue light having a wavelength of 450 nm. What is the increment of energy (the quantum) that is emitted at 4.50x102 nm by CuCl? Practice Exercise Microwave radiation has a wavelength on the order of 1.0 cm. Calculate the frequency and the energy of a single photon of this radiation. BILLACURA, M.D.  Atomic Spectrum of Hydrogen Continuous spectrum (results when white light is passed through a prism) – contains all the wavelengths of visible light Line spectrum – each line corresponds to a discrete wavelength: Significance Only certain energies are allowed for the electron in the hydrogen atom. Energy of the electron in the hydrogen atom is quantized. Emission and absorption spectrum of hydrogen BILLACURA, M.D.  The Bohr Model Assumptions Electrons in an atom can only occupy certain orbits (corresponding to certain energies). Electrons in permitted orbits have specific “allowed” energies; these energies will not be radiated from the atom. Energy is only absorbed or emitted in such a way as to move an Electronic Transitions in the Bohr Model for the electron from one “allowed” energy Hydrogen Atom state to another; the energy a) An Energy-Level Diagram for Electronic Transitions defined by: E = hv. b) An Orbit-Transition Diagram, Which Accounts for the Experimental Spectrum BILLACURA, M.D.  The Bohr Model The energy absorbed or emitted from a single electron transition from one energy level to another: ΔE = change in energy of the atom (energy of the emitted photon) nfinal = integer; final distance from the nucleus ninitial = integer; initial distance from the nucleus Limitations of Bohr’s model It only works for hydrogen! Classical physics would result in an electron falling into the positively charged nucleus. Bohr simply assumed it would not! Circular motion is not wave-like in nature. Important ideas of Bohr’s model Electrons exist only in certain discrete energy levels. Energy is involved in the transition of an electron from one level to another. BILLACURA, M.D.  Quantum Numbers ❖Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies. ❖Each orbital describes a spatial distribution of electron density. ❖An orbital is described by a set of three quantum numbers. Principal Quantum Number (n) The principal quantum number, n, describes the energy level on which the orbital resides. The values of n are integers ≥ 1. These correspond to the values in the Bohr model. Angular Momentum Quantum Number (l) This quantum number defines the shape of the orbital. Allowed values of l are integers ranging from 0 to n − 1. We use letter designations to communicate the different values of l and, therefore, the shapes and types of orbitals. BILLACURA, M.D.  Quantum Numbers Magnetic Quantum Number (ml) The magnetic quantum number describes the three-dimensional orientation of the orbital. Allowed values of ml are integers ranging from −l to l: −l ≤ ml ≤ l Therefore, on any given energy level, there can be up to 1 s-orbital, 3 p-orbitals, 5 d- orbitals, 7 f-orbitals, and so forth. ❖ Orbitals with the same value of n form an electron shell. ❖ Different orbital types within a shell are subshells. BILLACURA, M.D.  Example 1. For principal quantum level n = 5, determine the number of allowed subshells (different values of l), and give the designation of each. For n = 5, the allowed values of l run from 0 to 4 (n – 1 = 5 – 1). Thus the subshells and their designations are: l=0 l=1 l=2 l=3 l=4 5s 5p 5d 5f 5g 2. For l = 2, determine the magnetic quantum numbers (ml) and the number of orbitals. magnetic quantum numbers = –2, – 1, 0, 1, 2 number of orbitals = 5 BILLACURA, M.D.  s Orbital The value of l for s orbitals is 0. They are spherical in shape. The radius of the sphere increases with the value of n. For an ns orbital, the number of peaks is n. For an ns orbital, the number of nodes (where there is zero probability of finding an electron) is n – 1. As n increases, the electron density is more spread out and there is a greater probability of finding an electron further from the nucleus. Topic 1: Atomic Structure and Periodicity BILLACURA, M.D.  p Orbital The value of l for p orbitals is 1. They have two lobes with a node between them. BILLACURA, M.D.  d Orbital The value of l for a d orbital is 2. Four of the five d orbitals have four lobes; the other resembles a p orbital with a doughnut around the center. BILLACURA, M.D.  f Orbital Very complicated shapes. Seven equivalent orbitals in a sublevel, l = 3 BILLACURA, M.D.  Orbital Energy – Hydrogen Atom For a one-electron hydrogen atom, orbitals on the same energy level have the same energy. Chemists call them degenerate orbitals. BILLACURA, M.D.  Orbital Energy – Many-electron Atom As the number of electrons increases, so does the repulsion between them. Therefore, in atoms with more than one electron, not all orbitals on the same energy level are degenerate. Orbital sets in the same sublevel are still degenerate. Energy levels start to overlap in energy (e.g., 4s is lower in energy than 3d.) BILLACURA, M.D.  Quantum Numbers Spin Quantum Number (ms) In the 1920s, it was discovered that two electrons in the same orbital do not have exactly the same energy. The “spin” of an electron describes its magnetic field, which affects its energy. This led to the spin quantum number, ms. The spin quantum number has only two allowed values, +½ and –½. BILLACURA, M.D.  Pauli Exclusion Principle No two electrons in the same atom can have exactly the same energy. Therefore, no two electrons in the same atom can have identical sets of quantum numbers. This means that every electron in an atom must differ by at least one of the four quantum number values: n, l, ml, and ms. BILLACURA, M.D.  Electron Configuration The way electrons are distributed in an atom is called its electron configuration. 4p 5 The most stable organization is the lowest possible energy, called the ground state. Each component consists of ❖a number denoting the energy level; ❖a letter denoting the type of orbital; ❖a superscript denoting the number of electrons in those orbitals. BILLACURA, M.D.  Orbital Diagrams Each box in the diagram represents one orbital. Half-arrows represent the electrons. The direction of the arrow represents the relative spin of the electron. Example The orbital diagram of oxygen: Total no. of electrons: 8 Electron configuration, O: 1s22s22p4 BILLACURA, M.D.  Hund's Rule “For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.” This means that, for a set of orbitals in the same sublevel, there must be one electron in each orbital before pairing and the electrons have the same spin, as much as possible. BILLACURA, M.D.  Condensed Electron Configuration Elements in the same group of the periodic table have the same number of electrons in the outer most shell. These are the valence electrons. The filled inner shell electrons are called core electrons. These include completely filled d or f sublevels. We write a shortened version of an electron configuration using brackets around a noble gas symbol and listing only valence electrons. BILLACURA, M.D.  Filling of Orbitals in the Periodic Table We fill orbitals in increasing order of energy. Different blocks on the periodic table correspond to different types of orbitals: s = blue, p = pink (s and p are representative elements); d = orange (transition elements); f = tan (lanthanides and actinides, or inner transition elements) BILLACURA, M.D. BILLACURA, M.D.  Electron Configuration Anomalies Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row. ❖For instance, the electron configuration for chromium is [Ar] 4s1 3d5 rather than the expected [Ar] 4s2 3d4. ❖This occurs because the 4s and 3d orbitals are very close in energy. ❖These anomalies occur in f-block atoms with f and d orbitals, as well. BILLACURA, M.D.  Example Give the electron configurations for sulfur (S) and cadmium (Cd) Solution: Sulfur is element 16 and resides in Period 3, where the 3p orbitals are being filled. Since, sulfur is the fourth among the “3p elements,” it must have four 3p electrons. Its configuration is: S: 1s22s22p63s23p4 or [Ne]3s23p4 Cadmium is element 48 and is located in Period 5 at the end of the 4d transition metals. It is the tenth element in the series and thus has 10 electrons in the 4d orbitals, in addition to the 2 electrons in the 5s orbital. The configuration is: Cd: 1s22s22p63s23p64s23d104p65s24d10 or [Kr]5s24d10 BILLACURA, M.D.  Periodicity Periodicity is the repetitive pattern of a property for elements based on atomic number. The following properties are discussed in this chapter: Sizes of atoms and ions Ionization energy Electron affinity First, we will discuss a fundamental property that leads to may of the trends, effective nuclear charge. BILLACURA, M.D.  Effective Nuclear Charge Many properties depend on attractions between valence electrons and the nucleus. Electrons are both attracted to the nucleus and repelled by other electrons. The forces an electron experiences depend on both factors. The effective nuclear charge, Zeff, is found this way: Zeff = Z − S where Z is the atomic number and S is a screening constant, usually close to the number of inner electrons. Effective nuclear charge is a periodic property: ❖ It increases across a period. ❖ It decreases down a group. BILLACURA, M.D.  Atomic Radius The bonding atomic radius is half the internuclear distance when atoms are bonded. The bonding atomic radius tends to — decrease from left to right across a period (Zeff ↑). — increase from top to bottom of a group (n ↑). BILLACURA, M.D.  Ionic Radius Determined by interatomic distances in ionic compounds Ionic size depends on the nuclear charge. the number of electrons. the orbitals in which electrons reside. Cations are smaller than their parent atoms: The outermost electron is removed and repulsions between electrons are reduced. Anions are larger than their parent atoms: Electrons are added and repulsions between electrons are increased. BILLACURA, M.D.  Ionic Radius – Isoelectronic Series In an isoelectronic series, ions have the same number of electrons. Ionic size decreases with an increasing nuclear charge. An Isoelectronic Series (10 electrons) Note increasing nuclear charge with decreasing ionic radius as atomic number increases O2– F– Na+ Mg2+ Al3+ 1.26 Å 1.19 Å 1.16 Å 0.86 Å 0.68 Å BILLACURA, M.D.  Electron Configuration of Ions Cations: The electrons are lost from the highest energy level (n value). Example: Li+ is 1s2 (losing a 2s electron). Fe2+ is 1s22s22p63s23p63d6 (losing two 4s electrons). Anions: The electron configurations are filled to ns2np6. Example: F– is 1s22s22p6 (gaining one electron in 2p). BILLACURA, M.D.  Example Predict the trend in radius for the following ions: Be2+, Mg2+, Ca2+, and Sr2+ Solution: All of these ions are formed by removing two electrons from an atom of a Group 2A element In going from beryllium to strontium, we are going down the group, so the sizes increase: BILLACURA, M.D.  Ionization Energy, I The ionization energy is the minimum energy required to remove an electron from the ground state of a gaseous atom or ion. The first ionization energy is that energy required to remove the first electron. The second ionization energy is that energy required to remove the second electron, etc. Note: the higher the ionization energy, the more difficult it is to remove an electron! It requires more energy to remove each successive electron. When all valence electrons have been removed, it takes a great deal more energy to remove the next electron. BILLACURA, M.D.  Ionization Energy - Trends 1. I1 generally increases across a period. 2. I1 generally decreases down a group. 3. The s- and p-block elements show a larger range of values for I1. (The d- block generally increases slowly across the period; the f-block elements show only small variations.) Factors that Influences Ionization Energy Smaller atoms have higher I values. I values depend on effective nuclear charge and average distance of the electron from the nucleus. BILLACURA, M.D.  Ionization Energy – Irregularities The trend is not followed when the added valence electron in the next element enters a new sublevel (higher energy sublevel); is the first electron to pair in one orbital of the sublevel (electron repulsions lower energy). BILLACURA, M.D.  Electron Affinity Electron affinity is the energy change accompanying the addition of an electron to a gaseous atom: Cl + e− ⎯⎯→ Cl− It is typically exothermic, so, for most elements, it is negative! Not much change in a group. Across a period, it generally increases. Three notable exceptions include the following: 1) Group 2A: s sublevel is full 2) Group 5A: p sublevel is half-full 3) Group 8A: p sublevel is full Note: For Group 8A the electron affinity for many of these elements is positive (X– is unstable). BILLACURA, M.D.  EXERCISES 1. Which of the following would require more energy to remove an electron? Why? Sodium vs. Chlorine Lithium vs. Cesium 2. Which element has the larger second ionization energy? Why? Lithium vs. Beryllium 3. Which of the following should be the larger atom? Why? Sodium vs. Chlorine Lithium vs. Cesium 4. Which is larger? Why? The hydrogen 1s orbital The lithium 1s orbital 5. Which is lower in energy? Why? The hydrogen 1s orbital The lithium 1s orbital BILLACURA, M.D.  EXERCISES Explain why the graph of ionization energy versus atomic number (across a row) is not linear. Where are the exceptions? BILLACURA, M.D.