Chapter 5 and 6 Notes Electromagnetic Radiation and Light

Transcription

1 Models of the Atom Chapter 5 and 6 Notes Electromagnetic Radiation and Light There were many different models over time Dalton-billiard ball model (1803) Thompson plum-pudding model (1897) Rutherford model of the atom (1911) Bohr uses quantized of the atom (1913) Quantum Mechanical Model of the Atom (1 Each new model contributed to the model we use today. Even our current Quantum Mechanical model, does not give us an exact model of how behave. The Bohr Model of the Atom Bohr used the simplest element,, for his model He proposed an electron is found in specific circular paths, or orbits around the nucleus Each electron orbit was thought to have a fixed level Lowest level-ground state Any Higher level- state The Bohr Model of the Atom cont. One electron is capable of many excited states (whenever an electron jumps to higher level) Quantum: specific amount of an electron can or lose when moving energy levels You can excite an electron with energy like electricity,, or magnets Problems with the Bohr Model OOPS!-Model only works with. Model did not account for the behavior of atoms WRONG: do not move around the nucleus in circular orbits STILL VERY HELPFUL!!! How do Neon Signs Work? They have gases in them. Explanation Step 1: an electron energy and moves to a energy level Step 2: electron drops back down to a energy level During drop it gives off called a photon Sometimes this energy is light (ROYGBIV) When a photon is emitted, energy is released. We can calculate the energy released using the equation: Application: Atomic Emission Spectrum Used to determine which elements are present in a sample Used to determine which elements are present in a star (because stars are gases) Each element has a spectrum Only certain are emitted because the energy released relates to a specific frequency Spectroscope A spectroscope is needed to see the atomic emission spectra, which acts similar to a prism, separating different of light 1

2 Electromagnetic Spectrum Electromagnetic spectrum is the range of all energies emitted from photons acting like. Electromagnetic Spectrum with Visible Light Spectrum Light Behaves like a. Behaves like a. Characteristics of a Wave Wavelength (lambda) shortest between equivalent points on a continuous wave [Unit = meters] Frequency (nu) the of waves that pass a given point per second [Unit = 1/second = s -1 = Hertz (Hz)] Crest point of a wave Trough point of a wave Amplitude (a) height from its origin to its crest (highest point) or trough (lowest point) [Unit = meters] Wavelength and Frequency Wavelength ( ) and frequency ( ) are related As wavelength goes up, frequency goes down As wavelength goes down, frequency goes up This relationship is proportional Wavelength and Frequency cont. c = Speed of light (c) = 3 x 10 8 m/s Calculate the wavelength ( ) of yellow light if its frequency ( ) is 5.10 x Hz. What is the frequency of radiation with a wavelength ( ) of 5.00 x 10-8 m? What region of the electromagnetic spectrum is this radiation? How Much Energy Does a Wave Have? Energy of a wave can be calculated Use the formula E= h E= Energy = frequency h = Planck s constant = x Joule. Sec Joule is a unit for energy (J) Energy and frequency are directly proportional, as frequency increases, energy Remember that the energy of a photon is E =h How much energy does a wave have with a frequency ( ) of 2.0 x 10 8 Hz? ( h = x J. s) 2

3 Visible Light, Frequency, and Energy Red wavelength ( ), smallest frequency ( ) Red frequency smallest ( ), least amount of energy (E) Violet smallest wavelength ( ), largest ( ) Violet frequency largest ( ), greatest amount of energy (E) Flame Test The flame test is a way to determine the present in a sample When placed in a flame, each element gives off a color Operates the same as neon signs; electrons are excited by and fall back down and give off different colors Current Model of the Atom Quantum Mechanical Model of the Atom Quantum Mechanical Model is the current description of electrons in atoms. It does not describe the electron s around the nucleus Quantum Mechanical Model is based on several ideas including: Schrodinger wave equation (1926) treats electrons as. Heisenberg uncertainty principle (1927) states that it is impossible to know both the and of a particle at the same time. Where do electrons live? Principal Energy Levels 1. Principal energy levels n =1 to. (Row # on the periodic table) The electron s principal energy level is based on its location around the nucleus. Electrons closer to the nucleus are at a energy level and have lower energy than those farther away from the nucleus Atomic Orbital An is a region of space in which there is a of finding an electron Orbitals necessarily spherical Energy Sublevels (also called orbitals) and Orbitals 1. Energy sublevels assigned letters,,, or f (smart people do fine) Energy sublevels correspond to a where the electron is likely to be found. 2. Orbitals describes the electron s (maximum of electrons per orbital) s sublevel has 1 orbital (2 electrons total) - spherical p sublevel has 3 orbitals (6 electrons total) dumb-bell shaped d sublevel has 5 orbitals (10 electrons total )-double dumb-bells f sublevel has 7 orbitals (14 electrons total) Electron Configurations Energy Levels, Sublevels, and Orbitals 1. Principal energy levels n, assigned values (Like floors in a hotel) 2. Energy sublevels- s, p, d, f (Type of suite in a hotel) (Orbitals are like the number of rooms in a suite) 1. s sublevel 1 orbital 2. p sublevel 3 orbitals 3. d sublevel 5 orbitals 4. f sublevel 7 orbitals 3. Orbitals electrons per orbital (Two people per room) Electron Configurations Electron configuration the of electrons in an atom. Example Sodium (Na) 1s 2 2s 2 2p 6 3s 1 Three rules determine electron configurations the Aufbau Principle, the Pauli Exclusion Principle Hund s rule The Aufbau Principle Each electron occupies the energy orbital available 3

4 Like filling the hotel from the bottom up Pauli Exclusion Principle A maximum of electrons may occupy a single orbital Like only two people sharing one hotel room Hund s Rule If two or more orbitals of energy are available, electrons will occupy them with the same spin, before filling them in pairs with opposite spins A spin is denoted with an up or down arrow to fill orbitals This is like trying to find your own room in the same suite before having to share a room with someone else Writing Electron Configurations Aufbau diagram for sodium (Na) which has 11 electrons Na electron configuration1s 2 2s 2 2p 6 3s 1 Exceptions to Electron Configurations Copper and chromium are exceptions to the principle. Element Should be Actually is Copper 1s 2 2s 2 2p 6 3s 2 3p 6 3d 4 4s 2 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1 Chromium 1s 2 2s 2 2p 6 3s 2 3p 6 3d 9 4s 2 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 Some configurations violate the Aufbau Principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations Valence electrons Valence electrons are electrons in the orbitals. For A group elements the number corresponds to number of valence electrons. Electron-dot structures Element s symbol surrounded by representing the valence electrons Noble Gas Configuration What are Noble Gases? Noble gases are found in group The elements are called noble because they are non-reactive and very. The do not tend to form compounds Complete Electron Configuration What is the electron configuration for Ne? Ne: What is the electron configuration for Mg? Mg: What do both electron configurations have in common? 1s 2 2s 2 2p 6 = [Ne] Noble Gas Configuration (Abbreviated Configuration) Using neon s configuration and then adding magnesium s extra electrons we can get the noble gas configuration. Ne: 1s 2 2s 2 2p 6 = [Ne] Mg: 1s 2 2s 2 2p 6 3s 2 Noble gas configuration Mg: Only use noble gases in the brackets. Which Noble Gas is Used? To figure out which noble gas to use find the noble gas that is closest to the element without going over in atomic number Which noble gas is closest without going over? Rb : 4

5 Cl : Ra : What About the Other Electrons? To know what to write for the other electrons that are not included in the noble gas, understanding the periodic table is important. The periodic table is organized by blocks according to the energy Blocks of the Periodic Table There are s, p, f, and d blocks of the periodic table which correspond to the energy sublevels. s Block Elements Write the closest noble gas without going over in brackets. Use the row number to get the energy level. Count the number of electrons until you get to the element in the s block. Mg Try other s-block elements. Write the noble gas configuration of the following elements Cs Ca Ba p block elements (Between 5-18) Write the closest noble gas without going over in brackets. Use the row number to get the energy level. Write s 2 after the row number because you have to go through the s-block to get to the p-block. Write the row number again The write p and then count the number of p electrons you must get through to get to your element as a superscript Si: Try other p-block elements. Write the noble gas configuration of the following elements N : S : Cl : d block elements (Between 21-48) Write the closest noble gas without going over in brackets. Use the row number to get the energy level. Write s 2 after the row number because you have to go through the s-block to get to the d-block. Write one less than the row number (d-block elements are always one less than the row number)**d for down one row number Then write d and count the number of d electrons you must get through to get to your element as a superscript Co: Try other d-block elements. Write the noble gas configuration of the following elements Ti : Zn : Mn : p block elements (Between 31-53) Write the closest noble gas without going over in brackets. Use the row number to get the energy level. Write s 2 after the row number because you have to go through the s-block to get to the p-block. 5

6 Write one less than the row number (d-block elements are always one less than the row number)**d for down one row number Then write d and count the number of d electrons you must get through to get to your element as a superscript Write the row number again and p and count over the number of p electrons until you get to your element Br: Try other p-block elements. Write the noble gas configuration of the following elements Sn : Se: The Modern Periodic Table Early Periodic Table Atomic Number In 1913 Henry Mosley discovered that each element contained a unique number of protons in the nuclei Arranged elements in order of atomic. Resulted in a clear periodic pattern of properties. Periodic Law There is a periodic repetition of chemical and physical of elements when arranged in increasing atomic number (increasing number of protons) is called the periodic Modern Periodic Table Organized in columns called or families Rows are called Group A representative elements (1A- ) Group B - elements (1B-8B) Classification of Elements Three classifications for elements metals, nonmetals, and metalloids (semimetals) Metals Properties of Metals shiny, smooth, clean solids (except mercury) conductors of heat and electricity High 6 High melting and boiling points bended or pounded into sheets Ductile drawn into Groups of Metals metals group 1A except H Alkaline earth metals group Alkali metals and alkaline earth metals are chemically reactive Transition metals group elements Inner transition metals Lanthanide Actinide Organizing by Electron Configuration Group number for group A elements represents the number of electrons Atoms in the same group have similar chemical properties because they have the same number of valence electrons Alkali Metals Electron configurations for alkali metals Lithium [He]2s 1 Sodium 1s 2 2s 2 2p 6 3s 1 [Ne]3s 1 Potassium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 [Ar]4s 1 Rubidium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 [Kr]5s 1 What do the four configurations have in common? They have a electron in their outermost energy level They all have one valence electron, thus similar chemical properties Alkaline Earth Metals Electron configuration for alkaline earth metals Beryllium [He]2s 2 Magnesium [Ne]3s 2 Calcium [Ar]4s 2 Strontium [Kr]5s 2

7 All alkaline earth metals have valence electrons, thus similar chemical properties. Nonmetals Gases or brittle, dull looking solids conductors of heat and electricity Usually have lower densities, melting point, and boiling point than metals. Groups of nonmetals Halogens Noble gases Noble Gases Noble gases Group Called inert gases because they rarely take part in a reaction He 1s 2 Ne 1s 2 2s 2 2p 6 Ar 1s 2 2s 2 2p 6 3s 2 3p 6 Kr 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 Because noble gases have completely filled s and p sublevels, they do not react with other elements Metalloids (Semimetals) Physical and chemical properties similar to both metals and nonmetals They are metallic-looking solids Relatively good electrical conductivity. Used in glasses, alloys, and semiconductors The six elements commonly recognized as metalloids are boron, silicon, germanium, arsenic, antimony, and tellurium. Polonium and astatine are sometimes classified as metalloids Do the Trends w/s first! Periodic Trends Atomic Radius Defined as of the distance between two bonding atom s nuclei Atomic Radius Across a Period Atomic radius generally in size as you move left to right across the period positive charge in the nucleus pulls the electrons of the same energy level in. Atomic Radius Down a Group Atomic radius as you move down a group Orbital size increases as you move down a group with increasing energy level Larger orbitals means that outer electrons are from the nucleus. This increased distance offsets the greater pull of the increased nuclear charge. As additional orbitals between the nucleus and the outer electrons are occupied, the inner electrons shield the outer electrons from the pull of the nucleus this is called. 7

8 Cation and Anion An ion is a positively or negatively charged atom that gains or loses an. A cation loses electrons and produces a charge An anion gains electrons and produces a charge Ionic Radius - Cations Groups 1A, 2A, 3A, and other metals electrons and form cations. When atoms lose electrons they become The electron lost will be a valence electron leaving a completely empty outer orbital Protons in nucleus can pull fewer electrons tighter Ionic Radius - Anions Group 5A, 6A, and 7A tend to electrons and form anions When atoms gain electrons and form negatively charged ions, they become. Protons in nucleus have more electrons to pull and cannot pull in as tight Do Ionization and Electronegativity w/s First! Ionization Energy The energy required to an electron from a gaseous atom Indication of how strongly an atom s nucleus holds onto its electron Groups 1A, 2A, and 3A tend to have low ionization energies because they want to lose electrons. Ionization Energy Trends Across a Period Ionization energy generally as you move left to right Across a period electrons are added to the same energy level (same distance away from the nucleus), yet the nuclear charge is increasing across a period increasing the attraction to the electrons. Ionization Energy Trends Down a Group Ionization energy as you move down a group Down a group electrons are added to a higher energy level (farther distance away from the nucleus), making it easier to remove an electron Octet Rule Sodium atom 1s 2 2s 2 2p 6 3s 1 Sodium ion 1s 2 2s 2 2p 6 (Sodium atom lost 1 electron) Neon 1s 2 2s 2 2p 6 Sodium ion has the same electron configuration as neon Octet rule states that atoms gain, lose, or share electrons to acquire a full set of valence electrons (to be like a noble gas) Electronegativity Indicates an element s ability to electrons in a shared chemical bond fluorine (F) is the most electronegative element Cesium (Cs) and francium (Fr)are the least electronegative Noble gases do not tend to have an electronegativity number since they tend not to form Trends with Electronegativity Electronegativity as you move left-to-right across a period Electronegativity as you move down a group 8