Name Per. Notes: Electrons and Periodic Table (text Ch. 4 & 5) NOTE: This set of class notes is not complete. We will be filling in information in class. If you are absent, it is your responsibility to get missing information from a fellow classmate or the chemistry website: http://othschem.weebly.com/ I. The Wave Nature of Light Light travels through space as a wave. Waves have three primary characteristics: Wavelength (λ): the distance between 2 consecutive crests or troughs. Often measured in meters. Frequency (f): the number of wave cycles (successive crests or troughs) that pass a given point per unit of time. Often measured in cycles per second ( s 1 or Hertz). Speed (c): a constant for all forms of light; c = 3.00 x 10 8 m/s The speed of light, wavelength, and frequency are related by this formula: Wavelength and frequency are proportional 1
II. The Particle Nature of Light Not all properties of light can be explained by the wave model. Max Planck and Albert Einstein refined our understanding of light to include its particle nature. Light can be described as a stream of particles called photons. The details of the particle nature of light will not be investigated in this class. We will only highlight the relationship between the energy of light and its frequency: h = Planck s constant = 6.626 x 10 34 J. s Energy and frequency are proportional. All atoms absorb and emit light when excited by electricity or heat Elements emit light in unique, distinct patterns called Spectra corresponds to the movement of electrons between Spectra are like fingerprints that can be used to an element (Flame Test Lab) III. The Electron Cloud Electrons are found outside the nucleus, in a region of space called the. Electrons are organized in of positive integer value (n = 1, 2, 3,...). Within each energy level are, designated by a letter: s, p, d, or f. 2
Each sublevel corresponds to a certain electron cloud shape, called an. Analogy: The electron cloud is like an apartment building. The energy levels are like floors in the apartment building. The sublevels are like apartments on a floor of the building. Just like there are different sizes of sublevels, there are different sizes of apartments: 1 bedroom, 2 bedroom, etc. The orbitals are like rooms within an apartment. The electrons are like people living in the rooms. What do the atomic orbitals look like? Some examples: How are they organized around the nucleus? Here is a picture of the orbitals that make up the first four sublevels. Each orbital can hold a maximum of 2 electrons. An s sublevel contains 1 s orbital. How many total electrons can fit in an s sublevel? A p sublevel contains 3 p orbitals. How many total electrons can fit in a p sublevel? A d sublevel contains 5 d orbitals. How many total electrons can fit in a d sublevel? An f sublevel contains 7 f orbitals. How many total electrons can fit in an f sublevel? 3
IV. Electron Configurations and Orbital Diagrams Three rules govern the filling of atomic orbitals: 1. The Aufbau Principle: Electrons enter orbitals of lowest energy first. The Aufbau order lists the orbitals from lowest to highest energy: ( Aufbau is from the German verb aufbauen: to build up) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 14 6d 10 7p 6 First # = energy level, letter = type of sublevel, superscript # = max # of electrons You will learn to use the Periodic Table to determine the Aufbau order. It is not to be memorized. 2. The Pauli Exclusion Principle: An atomic orbital may hold at most 2 electrons, and they must have opposite spins (called paired spins). 3. Hund s Rule: When electrons occupy orbitals of equal energy (such as three p orbitals), one electron enters each orbital until all the orbitals contain one electron with spins parallel. Second electrons then add to each orbital so that their spins are paired with the first electron in the orbital. An electron configuration uses the Aufbau order to show how electrons are distributed within the atomic orbitals. Examples of electron configurations: Element Total # of Electrons Electron Configuration Carbon Fluorine Magnesium Argon Orbital diagrams show with arrow notation how the electrons are arranged in atomic orbitals for a given element. Element Orbital Diagram Carbon Fluorine Magnesium Argon 4
Some elements have large numbers of electrons, making their electron configurations cumbersome to write out in full. The noble gas configuration was developed for use with larger elements, as a shorthand way of writing electron configurations. Use the noble gas (Group 18) in the period above the element to write its noble gas configuration. Examples: Element Full Electron Configuration Noble Gas Configuration Sr Mn Bi V. Valence electrons Valence electrons are the electrons in the outer energy level of an atom. They are like the front lines of an army, because they are the ones involved in chemical reactions. Valence electrons are shared or transferred during reactions. The number of valence electrons is directly responsible for an element s chemical behavior and reactivity. The number of valence electrons can be determined from electron configurations. Count all the electrons in the outer (highest) energy level for that atom: Li Be B C N O F Element Electron Configuration # Valence Electrons Electron Dot Structure Ne The number of valence electrons for the representative element groups (Group A elements or Groups 1-2 and 13-18) follows an easy pattern from 1 8 going across the Periodic Table from left to right. Notice that the elements in a group (also called a family) all have the same number of valence electrons. Example: Group 1 elements (Alkali metals) all have 1 valence electron. Therefore, these elements will behave in similar ways. Try this: Find an element on the periodic table that has chemical reactivity similar to sulfur. 5
VI. Ions An ion is a charged particle that forms when an atom loses or gains electrons. Atoms lose or gain electrons in order to bond with other atoms and become more stable (less reactive). Remember: atoms are neutral because the # of protons (+) = the # of electrons (-) When an atom loses electrons, this results in a positively charged ion (cation). When an atom gains electrons, this results in a negatively charged ion (anion). Example: Sodium atoms have 11 p+ and 11 e-. If a sodium atom loses one electron, its charge will be +1 because it then has 11 p+ but only 10 e-. Element Ion formed Lost or gained electrons? How many? Sodium Na + Chlorine Cl Magnesium Mg 2+ Nitrogen N 3 Oxygen O 2 Aluminum Al 3+ A pattern exists with the Representative Elements (Groups 1, 2, 13, 15-17). With your teacher s assistance, assign the following charges on your periodic table to the Groups: Group 1: 1+ (or just +) Group 2: 2+ Group 13: 3+ Group 15 3 Group 16 2 Group 17 1 (or just ) Group 18 elements are the Noble Gases; for our purposes, they do not form ions. For Groups 3-12 and 14, ion charges can vary within the group, and even vary for one metal (ex. sometimes Cu is +1 and sometimes it is +2), so we cannot assign a charge for these groups as we did above. We will go into this in more detail later. You must memorize the charges for these 3 ions: (they are always the same) Ag is 1+ Zn is 2+ Cd is 2+ ***Note that the formulas and charges for polyatomic ions can be found on the back of your Periodic Table, with the exception of the phosphite ion: PO3 3- (memorize this one!) 6
VII. Ions and Electron Configuration Electron configurations for ions must be adjusted, based on the number of valence electrons lost or gained to form the ion. Element e- dot structure Lose or gain e-? How many? Ca.. Ca lose 2 O Ion formula w/ charge Ca 2+ Al F K N Now, for each of the above elements, write the electron configuration for the atom, and then adjust the configuration for the ion by losing or gaining the proper number of valence electrons. Element Electron config for ATOM Electron config for ION Ca Ca: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 Ca 2+ : 1s 2 2s 2 2p 6 3s 2 3p 6 (lost 2) O Al F K N Ionic compound form when a metal loses electrons to a nonmetal, and the resulting oppositely charged ions attract. With your teacher s assistance, make electron dot diagrams for the following pairs of elements, showing how electrons transfer to form the compound. 1) calcium and fluorine 2) sodium and nitrogen 3) aluminum and oxygen 7
VIII. Periodic Trends Trends in properties of the elements that follow a pattern down a group and across a period in the periodic table. 1. Trends in Atomic Radius Group trend: atomic radius going down a group. As you move down a group, energy levels are added, thus increasing the size of the electron cloud, so the atoms get larger. Periodic trend: atomic radius going left to right across a period. [NOTE: from now on, across a period will refer to the left-to-right direction] The atoms are getting heavier across a period. So why don t they get larger in volume? The reason has to do with increasing nuclear charge. With an increase in the # of protons from left to right, the nucleus gets a more powerful + charge. So it pulls the electron cloud in tighter. 2. Trends in Ionic Radius Positive ions are always smaller than the neutral metal atoms from which they were formed. When metal ions form, typically the outer energy level will be emptied, resulting in an overall smaller electron cloud. Also, the excess of protons compared to electrons draws the remaining electrons in closer. Example: Na vs. Na+ Negative ions are always larger than the neutral nonmetal atoms from which they were formed. There is more repulsion in the cloud due to the added electrons, therefore making it spread out, but there are no extra protons to pull it closer. Example: F vs. F - Group trend: Ionic radius (for both cations and anions) generally as you move down a group. This is mainly due to succeeding being filled. Periodic trend: Ionic radius (for both cations and anions) generally as you move across a period from left to right. This is mainly due to increasing. ***NOTE: anions within one period are larger than the cations within one period. This is due to the fact that cations have lost an energy level. 8
3. Trends in Ionization Energy (electron-losing) Ionization energy: the measure of the energy required to remove an electron from the outermost energy level of an atom. Group trend: ionization energy going down a group This is due to the shielding effect: an electron in the outer energy level of a large atom is easier to remove because it is well-shielded from the pull of the nucleus by the inner electrons. Diagram of the shielding effect: Periodic trend: ionization energy going across a period This is due to nuclear charge -- across a period, nuclear charge increases, so it becomes more difficult to remove an electron (held tighter). Note that this periodic trend supports the idea that metals have a much greater tendency to lose electrons than nonmetals do. 4. Trends in Electronegativity Electronegativity: the tendency of an atom to attract electrons to itself when it is chemically bonded with another element. Electronegativity is a numerical scale which can be used to predict whether atoms will form ionic or covalent bonds in molecules. In H2O, oxygen is more electronegative than hydrogen, so it pulls the electrons closer, and thus obtains a partially negative charge. Diagram of water molecule: Group trend: electronegativity down a group. Larger atoms have more energy levels, so it is harder for them to attract electrons to the nucleus (shielding effect). Periodic trend: electronegativity across a period. Nonmetallic character increases across a period, and nonmetals attract electrons more than metals do, because of increasing nuclear charge. 9
Summary of Periodic Trends: Atomic radius and ionic radius increase down a group, and decrease across a period. Ionization energy and electronegativity decrease down a group, and increase across a period. Write in the patterns that the trends follow on this periodic table: 1 Group 1 2 Periodic Table 13 14 15 16 17 18 2 3 4 3 4 5 6 7 8 9 10 11 12 5 6 7 *Be sure that you can label and describe all of the following on a Periodic Table: (Study your gallery walk and the Periodic Table you colored) Groups and Periods Metals Nonmetals Metalloids Alkali metals Alkaline earth metals Transition metals Lanthanides Actinides Halogens Noble Gases 10