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Unit 4A Glow in the Dark 4.1 History of Atomic Theory **The ______________ is defined as the smallest particle of an element that retains the properties of that element.** History 400 BC Democritus, a Greek philosopher expressed his idea that matter was made of very small, indivisible particles that he named “ATOMOS” 1803 John Dalton Assumptions of the theory: 1 All matter is made up of _________________ particles called atoms. 2 Atoms of the ___________ element are identical in properties. Atoms of different elements are different in properties. 3 Atoms of different elements combine in whole number ratios to form _____________________. 4 Chemical _______________ involve the rearrangement of atoms. Atoms cannot be created, divided, destroyed or changed into other types of atoms Diagram of Dalton’s Atomic Model tiny solid ball that could not be broken up into parts. What was wrong with Dalton’s Theory? o The Discovery of Subatomic Particles led to the idea that the atom is not “INDIVISIBLE.” o Atoms of the same element can have different masses(isotopes) Discovery of Subatomic Particles 1897 Thompson discovers the ______________ through the use of a cathode ray tube. He knows it is ________________ charged and has an extremely small mass. Goldstein determines there are positive particles called protons and have a mass 1837 times heavier than an electron. Plum Pudding Model (ball of positive charge with negatively charged particles evenly distributed) 1910 Rutherford discovers the nucleus through the gold foil experiment and that atoms are mostly _____________________ . Rutherford’s Atomic Model (sphere with dense middle center called the _________________ with electrons dispersed around it. 1932 Chadwick confirms the ______________________ which has a mass similar to the _____________________and no charge. They are located in the nucleus. 1913 Bohr performed experiments with hydrogen and light. Electrons are in levels according to how much energy they have and that only certain energy amounts are allowed. Think of the energy levels as rungs of a _________________. The farther away an energy level is from the nucleus, the more _____________ it attains! The first level can hold 2 electrons, then the next two levels can each hold 8 and then levels farther out can hold 18. Electrons can move from one energy level to the next by gaining or losing energy(quanta). Ground State: An electron is as close to the nucleus as it can get. Excited State: An electron in a higher energy level than it should be. Drawing of simple Bohr Models When assigning electrons, a max of 2 electrons are placed in the first shell, up to 8 in the 2nd shell, up to 18 in the 3rd shell, etc. Only 8 electrons can be placed in the 3rd shell at first, then 2 electrons will move into the 4th shell and the remaining of the 18 will be placed back in the 3rd shell for a total of 18. **( Just know this : it will be explained in the next section) Valence electrons are the outermost electrons is found in the ___________________ level of the atom. Using the drawn diagrams, what is the number of valence electrons for H _______ He ______ Li ________ Na _______ Fe ________ Lewis Dot Diagrams A way to show the valence electrons in an atom. The symbol represents the nucleus and inner core electrons. The dots represent the valence electrons. Only 2 dots per side. Only 4 sides. Max. of 8 valence electrons. Add dots one at a time on each side until each side is full. Then add a second dot to make a pair when needed. Exception is Helium. It can look like this: He: Examples H Li Be B C N O F Ne Fe 1920 Modern Atomic Theory Schrodinger developed the Quantum Mechanical Model (Modern view of the atom) Modern atomic theory uses calculus to show how electrons act as both particle & _____________ These equations show the most probable location of electrons in the atom (known as atomic ______________________________) 4.2 Atomic Structure Subatomic Particles and their Properties Particle Symbol Location Electrical Charge Actual mass in grams +1 Relative Mass in a.m.u* 1 amu -1 1 amu 1.6710-27 kg 0 .00055 amu 9.1010-31 kg 1.6710-27 kg p+ n0 e*a.m.u : atomic mass unit 1 amu (“atomic mass unit”) = 1.66 10-27 kg What are the 2 regions of the atom as of now? Nucleus: dense center containing ___________________ and _____________________. Electron Cloud: region surrounding nucleus containing electrons and mostly ________________________. Counting Subatomic Particles Mass Number The number of protons and neutrons in a nucleus is called the mass number. Round atomic mass to a whole number to get an element’s mass number. mass # = _________________ + _________________ Nitrogen’s mass number is 14. Atomic Number Every atom has a different number of protons. The number of protons determines the identity of the atom The atomic number shows the number of protons. ___________________= protons Nitrogen’s atomic number is 7! Calculating Neutrons To calculate the number of neutrons, subtract the atomic number from the mass number # _______________________ = mass # - atomic # 14 – 7 = 7 neutrons Charges A neutral atom has the ________________number of protons and electrons. An ion has a charge. Overall Charge = protons - electrons Example: How many electrons does Br-1 have? How mane electrons does Al+3 have? Element Information Nuclear Symbol ____ p ____ e ____ n Hyphen Notation: copper – 65 ___________ Example Nuclear Symbol Hyphen Notation Atomic # Magnesium25 Mass # Charge Proton Neutron Electron 126 82 +2 82 How do Atoms Differ: Isotopes Isotopes are atoms of the same element with a different number of ____________________. Most elements contain a mixture of 2 or more isotopes. Each one having its own mass and abundance. Isotope Atomic number Protons Neutrons Electrons Mass (a.m.u) Lithium-6 Lithium-7 Lithium-8 You Try ? ( Be Careful: Might not be most common isotope. Never use the given atomic mass on the periodic table unless absolutely necessary!!) What the number of protons, neutrons and electrons in each atom? 19 57 F 9 204 Fe Hg 26 80 How to calculate Average Atomic Mass? 35 37 Cl 17 35.45 amu WHY? Cl 17 Average atomic mass is the weighted _____________________ of the masses of all naturally occurring isotopes. Equation: Average atomic mass = (% abundance of isotope x mass of 1st isotope) + (% abundance of isotope x mass of 2nd isotope) + ……… Example: Element x has 2 natural isotopes. Calculate the average atomic mass. 1st isotope has a mass of 10.012 a.m.u with 19.91% abundance. 80.09% of the 2 nd element has a mass of 11.009 a.m.u. You Try! Calculate the average atomic mass if copper if it has 2 isotopes. 69.11% has a mass of 62.93 a.m.u and the rest has a mass of 64.93 a.m.u. 4.3 Electron Structure Where are the electrons? 1. Within the Electron Cloud are “Energy Levels” There are ______ on the periodic table. The period number on the periodic table corresponds to the energy level. Energy levels are also called ___________________. Ca is in energy level _______ Cl is in energy level _____________ 2. Within each energy level are “Subshells” Subshells are a set of orbitals with equal _______________ “s” subshell spherical shaped there is only ______ orientation (position) = orbital represented in the periodic table as groups 1A and 2A + helium first seen in the 1st energy level maximum ______ electrons “p” subshell dumbbell shaped there are ______ orientations(positions) = orbitals represented in the periodic table as groups 3A- 8A first seen in the 2nd energy level maximum ____ electrons “d” subshell four lobed shaped there are _______ orientations(positions)= orbitals represented in the periodic table as the transition metals, group 3B – 2B first seen in the 3rd energy level maximum ____ electrons “f” subshell too complex of a shape to name there are _______ orientations(positions)= orbitals represented in the periodic table as the inner transition metals, lower block first seen in the 4th energy level maximum ____ electrons 3. Within each sublevel are “Orbitals” An orbital is defined as an area of high ______________ of the electron being located. Each orbital can hold ___________ electrons To calculate the total number of electrons in an energy level, use 2(n2) Summary n2 2(n2) Energy Level # of subshells subshell Total # of orbitals Total # of electrons 1 s 1 2 1 2 s&p 1&3=4 8 3 s&p&d 1&3&5=9 18 2 3 4 s & p & d & f 1 & 3 & 5 & 7 = 16 32 4 Electron Configuration Is an _________________ of an electron Electrons must be placed in the ____________ possible energy levels first (ground state) 4p1 3 rules that govern electron configuration 1. Aufbau Principle: electrons must be fill the lowest available subshells and orbitals before moving to the next higher energy subshell/orbital. Filling Order: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d Can use the periodic table as a guide or memorize filling order An orbital can ____________ within different energy levels __________is lower in energy than 3d. 2. Hund’s rule: place electrons in unoccupied energy orbitals of the same energy level before doubling up. Example: You need to add 3 electrons to a p subshell. 3. Pauli Exclusion Principle: two electrons in the same orbital must have opposite ___________. Example: you need to add 4 electrons to a p subshell Electron Configuration 1. Determine the number of electrons to place 2. Follow Aufbau Principle for filling order 3. Fill in subshells until they reach their max (s = 2, p = 6, d = 10, f = 14) or use periodic table as a guide. 4. The total of all the superscripts is equal to the number of electrons Example: Write electron configuration for S Write electron configuration for K Write electron configuration for Ti Orbital Notation use boxes or lines for orbitals and arrows for electrons. Examples Write the boxes & arrows (orbital notation) for Cl Write the orbital notation for N Write the orbital notation for Fe Shorthand Notation Noble gas is used to represent the core (inner) electrons and only the valence shell is shown. 1. Determine the number of electrons to place 2. Determine which noble gas to use 3. Start where the noble gas left off and write electron configuration for the valence electrons Example: Write the shorthand notation for Br. Write the shorthand notation for Fe. Write the shorthand notation for Ba Exceptional (special) Configurations Elements up to vanadium (V) follow Aufbau principle Half-filled or completely filled d & f sublevels have lower energies and are more stable than partially filled d’s and f’s. This means that an atom can “borrow” one of its “s” electrons from the previous orbital to become more stable. Example: Ag ___ ___ ___ ___ ___ ___ 5s 4d becomes ___ ___ ___ ___ ___ ___ 5s 4d Because the 4d sublevel is now full, the atom is at a lower energy state and therefore more stable. Electron Configuration for Ions o Determine the number of electrons to place. Positive ions lose electrons; negative ions gain electrons. o Follow Aufbau Principle for filling order o Fill in subshells until they reach their max (s = 2, p = 6, d = 10, f = 14) or use periodic table as a guide. Example: Write the electron Configuration for the sulfide ion, S-2