Models of the Atom Section 13.1 The story of how the atomic theory has evolved over time…. John Dalton The atom is a solid indivisible mass. He had 4 key ideas: Page 107 All elements are composed of tiny indivisible particles called the atom. Atoms of the same element are identical. Atoms of different elements are different. Atoms combine chemically with one another in simple whole-number ratios. During chemical reactions, atoms are separated, joined, or rearranged. Atoms are never created nor destroyed. JJ Thomson Plum-pudding model. Sometimes called the Chocolate Chip Cookie Model The atoms has negatively charged electrons stuck into a lump of positively charged material, like chocolate chips stuck in a cookie dough. Did not address protons and neutrons. Ernest Rutherford Discovered the nucleus. Showed that most of an atom’s mass is concentrated in a small, positively charged region called the nucleus. Electrons resided on the outside. Did not address how electrons were arranged. Neils Bohr Electrons are arranged on concentric circular paths, or orbits around the nucleus. Solar system model or planetary model. Gave us the idea of definite energy levels. Quantum Mechanical Model – Our Currently Accepted Model Erwin Schrodinger Primarily a mathematical model using quantum mechanics It addresses “probabilities” of finding an electron at any instant in an area called “electron clouds”. Introduced the ideas of Principal Energy Levels and Sublevels of energies. The electron clouds take certain shapes, represented by the s,p,d,f subatomic orbitals. Principal Energy Levels Just like the Bohr model, the Quantum Mechanical Model designates energy levels of electrons by means of principal quantum numbers called (n). Principal Energy Levels refers to a major region where electrons are most likely to be found. They are assigned values in order of increasing energy: n=1, n=2, n=3, etc. Sublevels Within each principal energy level, the electrons occupy energy sublevels. The number of sublevels within each principal energy level is the same as the principal quantum number. How many sublevels does the 4th principal energy level have? Atomic Orbitals The regions in which electrons are likely to be found are called atomic orbitals. Letters denote the atomic orbitals S-shape orbitals are spherical P-shape orbitals are hour-glass shapes D-shape orbitals have clover-leaf shapes Draw an example of each into your notes. Exploring further… The lowest principal energy level (n=1) has only one sublevel, called 1s. The second principal energy level (n=2) has 2 sublevels, the 2s and 2p. The 2p is higher in energy and consists of 3 p orbitals. Let’s show an easy way to remember this order of filling electrons… The Electron Pyramid The s orbitals have 1 spatial orientation, therefore can hold 2 electrons The p orbitals have 3 spatial orientations, therefore can hold 6 electrons The d orbitals have 5 spatial orientations, therefore can hold 10 electrons The f orbitals have 7 spatial orientations, therefore can hold 14 electrons. Electrons Fill following 3 simple rules Aufbau principle: Electrons enter the lowest energy level first. Pauli Exclusion Principle: An atomic orbital may describe at most 2 electrons, both spinning in opposite directions. Hund’s Rule: When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with parallel spins. Classifying Elements by Electron Configuration Of the three major subatomic particles, the electron plays the most significant role in determining the physical and chemical properties of an element The arrangement of elements in the periodic table depends on these properties Elements can be classified into four categories according to their electron configurations 1. 2. 3. 4. The noble gases The representative elements The transition metals The inner transition metals The Noble Gases These are elements in which the outermost s and p sublevels are filled Belong to Group 0 (Group 8A) Configurations Helium Neon Argon Krypton 1s2 1s22s22p6 1s22s22p63s23p6 1s22s22p63s23p64s23d104p6 The Representative Elements In these elements, the outermost s or p sublevel is only partially filled Usually called the Group A elements and include the noble gases in some definitions Configurations Lithium Sodium Potassium Carbon Silicon Germanium 1s22s1 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p2 1s22s22p63s23p2 1s22s22p63s23p64s23d104p2 The Transition Metals These are metallic elements in which the outermost s sublevel and nearby d levels contain electrons Called the Group B elements Configurations Zinc 1s22s22p63s23p64s23d10 Tungsten 1s22s22p63s23p64s23d104p65s24d105p66s24f145d4 Lead 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p2 The Inner Transition Metals These are metallic elements in which the outermost s sublevel and nearby f sublevel generally contain electrons Characterized by the filling of f orbitals Configurations Uranium 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f4 Curium 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s25f8 S1 S2 1 2 S2 3 Period 5 d1 d2 “S” Block 4 p1 d3 d4 d5 d6 d7 d8 d9 p2 p3 p4 p5 d1 0 “P” Block “D” Block 6 7 Period F1 6 7 F2 F3 F4 F5 F6 F7 F8 “F” Block F9 F10 F11 F12 F13 F14 p6 Practice Write the electron configurations for the following elements: Nitrogen 1s22s22p3 Nickel 1s22s22p63s23p64s23d8 Vanadium 1s22s22p63s23p64s23d3 Practice What are the symbols for all the elements that have the following outer configurations? S2 Helium (He) Beryllium (Be) Magnesium (Mg) Calcium (Ca) Strontium (Sr) Barium (Ba) Radium (Ra) S2P5 Fluorine (F) Chlorine (Cl) Bromine (Br) Iodine (I) Astatine (At) S2D2 Titanium (Ti) Zirconium (Zr) Hafnium (Hf) Rutherfordium (Rf) Exceptional Electron Configurations Chromium and Copper have exceptional electron configurations. They fill their d sublevel completely, leaving their 4s partially filled. Much more stable this way! Write them correctly into your notepacks… Cr: 1s22s22p63s23p63d54s1 Cu: 1s22s22p63s23p63d104s1 Physics and the Quantum Mechanical Model This section studies the electron as a property of light. Electrons travel as waves and are made of particles of light called photons According to the wave model, light consists of ELECTROMAGNETIC RADIATION. Electromagnetic Spectrum This form of energy includes Gamma rays X-rays Ultraviolet rays Visible light Infrared rays Radar FM TV Shortwave AM Electromagnetic Spectrum Every element emits light when it is excited by the passage of electric discharge through its gas or vapor. The atoms first absorb energy, then lose the energy as they emit light. Electromagnetic Spectrum Electrons are said to move from their GROUND STATE (lowest energy level) to and EXCITED STATE (higher energy level). When the electron falls back to its lower energy, it emits a PHOTON of energy, and can be seen in the visible spectrum. Electromagnetic Spectrum Passing the light emitted by an element through a prism gives the ATOMIC EMISSION SPECTRUM of the element. Because each atom has a unique electron arrangement, each atom emits a unique wavelength during this process. This wavelength falls within the visible spectrum. Kernel Structures The kernel is a structure used to shorten an electron configuration. A kernel is an inert gas symbol in brackets that stands in place of all of the filled orbitals contained in the inert gas. Examples Section 14.1 Classification of the Elements By Electron Configuration Classifying Elements by Electron Configuration Of the three major subatomic particles, the ELECTRON plays the most significant role in determining the properties of an element. The arrangement of elements in the PERIODIC TABLE depends on these properties. Elements can be classified into 4 categories: The Noble Gases : These are elements in which the outermost s and p sublevels are filled. Write for Helium, Neon, Argon, Krypton The representative elements: In these elements, the outermost s and p sublevel is only partially filled. Write for Lithium, Sodium, Potassium, Carbon, Silicon, Germanium Elements can be classified into 4 categories: The transition metals: These are metallic elements in which the outermost s sublevel and nearby d sublevel contain electrons. Write for Zinc and Zirconium. The inner transition metals: These are metallic elements in which the outermost s sublevel and nearby f sublevel generally contain electrons. 14.2 Periodic Trends Atomic radius – ½ the distance between the nuclei of two like atoms in a diatomic molecule. Group Trends Atomic size generally increases as you go down a group on the periodic table. Adding additional energy levels! Periodic Trends Atomic size generally decrease as you move from left to right across a period. Same energy level increasing nuclear charge pulls electrons closer to nucleus. Ionization Energy An ion – a charged atom that results from either losing or gaining an electron. Ionization Energy – The energy required to overcome the attraction of the nuclear charge and remove an electron from a gaseous atom. (The ease of losing an electron and forming a +1 charge) Ionization Energy First ionization energy – the energy needed to remove the first electron from an atom. Second ionization energy – the energy needed to remove the second electron from an atom, etc. Ionization Energy Group Trends: The first ionization energy generally decreases as you move down a group on the periodic table. The size of the atoms increases, so the outermost electron is farther from the nucleus and will be more easily removed. Ionization Energy Periodic Trends – For the representative elements, the first ionization energy generally increases as you move from left to right across a period. Increasing nuclear charge makes it more difficult to remove an electron. Ionic Size The atoms of METALLIC elements have low ionization energies. They form POSITIVE ions easily. By contrast, the atoms of NONMETALLIC elements readily form NEGATIVE ions. Trends in Ionic Size Positive ions are always smaller than the neutral atoms from which they form. They lose their outer shell electrons Negative ions are always larger than the neutral atoms from which they form. This is because the effective nuclear attraction is less for an increased number of electrons. Trends in Electronegativity Electronegativity – the tendency for the atoms of the element to attract electrons when they are chemically combined with atoms of another element. Electronegativity generally DECREASES as you go down a group. As you go across a period from left to right, the electronegativity of the representative elements INCREASES. Electronegativity The electronegativity of cesium, the least electronegative element is 0.7 The electronegativity of fluorine, the most electronegative element, is 4.0 Electronegativity values help predict the type of bonding that can exist between atoms in compounds, either IONIC OR COVALENT bonds. Summary of Periodic Trends Using page 406, create a summary of periodic trends into your notes.