Unit 3 Notes: Periodic Table Notes

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Unit 3 Notes: Periodic Table Notes John Newlands proposed an organization system based on increasingatomic mass in 1864. He noticed that both the chemical and physical properties repeated every 8elements and called this the Law of Octaves . In 1869 both Lothar Meyer and Dmitri Mendeleev showed a connectionbetween atomic mass and an element’s properties. Mendeleev published first, and is given credit for this. He also noticed a periodic pattern when elements were ordered byincreasing Atomic Mass . By arranging elements in order of increasing atomic mass into columns,Mendeleev created the first Periodic Table. This table also predicted the existence and properties of undiscoveredelements. After many new elements were discovered, it appeared that a number ofelements were out of order based on their Properties . In 1913 Henry Mosley discovered that each element contains a uniquenumber of Protons . By rearranging the elements based on Atomic Number , theproblems with the Periodic Table were corrected. This new arrangement creates a periodic repetition of both physical andchemical properties known as the Periodic Law .Periods are the RowsGroups/Families are the ColumnsValence electrons across a period arein the same energy levelThere are equal numbers of valenceelectrons in a group. When elements are arranged in order of increasing Atomic Number ,there is a periodic repetition of their physical and chemical properties1

Family (Group): Columns (vertical) ; tells the number of electronsin the Outer Energy level, called Valence Electrons (onlyfor representative elements) Period (Series): Rows (horizontal) ; tells the number of EnergyLevels an atom has; the number of electrons Increasesacross a period Representative Elements: Groups 1A through 8A (called the s and pblocks) (Columns 1, 2, 13, 14, 15, 16, 17, and 18) Valence Electrons: e- in the outer most energy level ; farthest awayfrom the nucleus (protons) ; the e- with the most reactiveEnergy; the e- involved with Bonding (transferred or shared) Metals: most of the periodic table, located to the Left of the “stair-step”Properties- good conductors of heat and Electricity ; they also areMalleable ; Ductile ; High Density, BP and MP Nonmetals: to the Right of the “stair-step”, located in the upper corner ofP.T. Although five times more elements are metals than nonmetals, twoof the nonmetals—hydrogen and helium—make up over 99 per centof the observable Universe Properties- mostly Brittle , but a few low luster and poorconductors ; they have low density, low MP and BP Metalloids: also called semi-metals , located along the “stair-step” Properties - similar to both metals and nonmetals Some metalloids are shiny (silicon), some are not (gallium)Metalloids tend to be brittle, as are nonmetals.Metalloids tend to have high MP and BP like metals.Metalloids tend to have high density, like metals.Metalloids are semiconductors of electricity – somewhere betweenmetals and nonmetals. This makes them good for manufacturingcomputer chips.2

ElementLithiumGermaniumSulfurSymbolLiGeSGroup #1A(1)4A(14)6A(16)# of valence e-146Period #243# of E levels243Type of elementMMNMPeriodic Trends:1. Atomic Size- Decreases from left to right across a period (smaller)- Increases from top to bottom down a group (larger)Why?- as you go across a period, (same energy level ), e- areadded but pulled closer to the nucleus- as you go down a group, you add energy levels2. Ionization Energy: the amount of E needed to remove / lose an electron- Increases from left to right across a period- Decreases from top to bottom down a groupWhy?3

- as you go across a period, e- feel stronger attraction from nucleus(protons) ,Energy to remove e-, Ionization E necessaryas you go down a group, more Energy , Decreases to removeoutermost e- because they are further away from the Nucleus (protons)3. Electronegativity: the tendency for an atom to attract electrons;exclude Noble Gases!- Increases from left to right across a period (except Noble Gases)- Decreases from top to bottom down a groupWhy?- as you go across a period, e- feel more attraction from nucleusProtons to pull in more e- as you go down a group, more shielding from inner e-,hinders the nucleus ability to attract more e-4. Ionic Size:Cations: positive ions; metal atoms that lose electrons4

- smaller than corresponding neutral atomWhy?- fewer e-, so it’s easier for protons to pull in remaining eAnions: Negative ions; nonmetal atoms that gain electrons- larger than corresponding neutral atomWhy?- more e-, so it’s harder for protons to pull in outermost eShielding:The ability of the inner (lower levels) electrons to shield (reduce) the pullof the protons on the outer (higher levels) electrons.“Shielding effect” increase as you add Energy levels (move down a group)Quantum Model Notes Heisenberg's Uncertainty Principle- Can determine either the velocity or theposition of an electron, cannot determine both. Schrödinger's Equation - Developed an equation that treated the hydrogenatom's electron as a wave.o Only limits the electron's energy values, does not attempt to describe theelectron's path. Describe probability of finding an electron in a given area of orbit. The Quantum Model- atomic orbitals are used to describe the possible positionof an electron.Orbitals The location of an electron in an atom is described with 4 terms.5

o Energy Level- Described by intergers. The higher the level, the more energyan electron has to have in order to exist in that region.o Sublevels- energy levels are divided into sublevels. The # of sublevelscontained within an energy level is equal to the integer of the energy level.o Orbitals- Each sublevel is subdivided into orbitals. Each orbital can hold 2electrons.o Spin- Electrons can be spinning clockwise ( ) or counterclockwise (-) withinthe orbital.Periodic Table Activity:Complete the table on page 21 with the information found on pages 18-20. When complete color each group ina different color in the periodic table.The Periodic Table Notes:Historical development of the periodic table: Highlights Mendeleev (1869): Put the elements into columns according to their properties. Generally rankedelements by increasing atomic mass. Moseley (1911): Periodic table arranged by atomic numberTop table: Metals, nonmetals, and metalloids Metals: Explain the electron sea theory, and as you explain each of the properties below, discuss howthey are explained by the electron sea theory. Also make sure to explain that these are generalproperties and may not be true for all metals.o Malleable: Can be pounded into sheets.o Ductile: Can be drawn into wireso Good conductors of heat and electricityo High density (usually)o High MP and BP (usually)o Shinyo Hard Nonmetals: Explain how the bonds between the atoms are highly localized, causing each of theproperties below. Again, emphasize that these are general properties and may not be true for allnonmetals.o Brittleo Poor conductors of heat and electricityo Low densityo Low MP and BP (many are gases)! Metalloids: The bonding in metalloids is between that of metals and nonmetals, so metalloids haveproperties of both.o Some metalloids are shiny (silicon), some are not (gallium)o Metalloids tend to be brittle, as are nonmetals.o Metalloids tend to have high MP and BP like metals.o Metalloids tend to have high density, like metals.6

Metalloids are semiconductors of electricity – somewhere between metals and nonmetals. Thismakes them good for manufacturing computer chips.Structure of the periodic table Families/groups (the terms are synonymous and will be used interchangeably)o These are elements in the same columns of the periodic table.o Elements within families/groups tend to have similar physical and chemical properties.o They have similar chemical and physical properties because they have similar electronconfigurations. Example:Li [He] 2s1, Na [Ne] 3s1 – each has one electron in the outermostenergy level.o Explain that s- and p-electrons in the outermost energy level are responsible for the reactionsthat take place. Valence electrons: The outermost s- and p-electrons in an atom. Show them how to find the number of valence electrons for each atom and explain thatthey are only relevant for s- and p- electrons. Do several examples.oPeriods: Elements in the same rows of the periodic tableo Elements in the same period have valence electrons in the same energy levels as one another.o Though you’d think this was important, it has very little effect on making the properties of theelements within a period similar to one another. The closer elements are to each other in the same period, the closer are their chemicaland physical properties. Other fun locales in the periodic table:o Main block elements: These are the s- and p- sections of the periodic table (groups 1,2, 13-18)o Transition elements: These are the elements in the d- and f-blocks of the periodic table. The term “transition element”, while technically referring to the d- and f-blocks, usuallyrefers only to the d-block. Technically, the d-block elements are the “outer transition elements” Technically, the f-block elements are the “inner transition elements”Major families in the periodic table:(Show them examples of these elements – if available – and color each family as I discuss their properties) Group 1 (except for hydrogen) – Alkali metalso Most reactive group of metalso Flammable in air and watero Form ions with 1 chargeo Low MP and BP (MP of Li 181º C, Na 98º C)o Soft (Na can be cut with a knife)o Low density (Li 0.535, Na 0.968) Group 2: Alkaline earth metalso Reactive, but less so than alkali metalso React in air and water (show Ca reacting in water)o Form ions with 2 chargeo Low MP and BP, but higher than alkali metals (MP of Ba 302º C, Mg 649ºCo Soft, but harder than alkali metalso Low density, but higher than that of alkali metals (Ca 1.55, Mg 1.74). Groups 3-12: (Outer) transition metalso Note: These are general properties and may vary from transition metal to transition metal!There are many exceptions to each of these rules!o Stable and unreactive.o Hard 7

o High MP and BP (Fe 1535º C, Ti 1660º C).o High density (Fe 7.87, Ir 22.4)o Form ions with various positive charges (usually include 2 and several others)o Used for high strength/hardness applications, electrical wiring, jewelryInner Transition Metals: Lanthanides and actinideso Lanthanides (4f section) Also called the rare earth metals, because they’re rare. Usually intermediate in reactivity between alkaline earth metals and transition metals. High MP and BP Used in light bulbs and TV screens as phosphors.o Actinides (5f section) Many have high densities Most are radioactive and manmade Melting points vary, but usually higher than alkaline earth metals. Reactivity varies greatly Used for nuclear power/weapons, radiation therapy, fire alarms.Group 13: Boron GroupGroup 14: Carbon GroupGroup 15: Nitrogen GroupGroup 16: Oxygen GroupGroup 17: Halogenso The most highly reactive nonmetals.o Highly volatile – F and Cl are gases, Br is a volatile liquid, and I is an easily sublimed solid.o Strong oxidizers – they readily pull electrons from other atoms.o Diatomic – form molecules with formula of X2o Form ions with -1 chargeo Used in water treatment and chemical production – Cl2 was used as a chemical weapon in WorldWar I.Group 18: Noble Gaseso Highly unreactiveo Used to provide the atmosphere in situations where you don’t want chemical reactions to occur(light bulbs, glove boxes, etc).Hydrogen – “The Weirdo”o Has properties unlike any other elemento Diatomic – H2o Can form either a 1 or -1 chargeo Relatively unreactive unless energy is added (under most conditions) – it can form explosivemixtures with oxygen (as it did in the Hindenburg explosion)8

Groups on the Periodic Table Summary Sheet:GroupExamplesof WordsusedAlkaliMetalsAlkalineEarthMetalsLocation onPeriodic TableGroup 1,Group 3-12,etcMetals, tal 1Highlyreactive,unreactiveInterestingInformationIt can be cutwith a plasticknifeM 1YNAny Name inFamily 11M 2YNAny Name inFamily 22M 2NNAny Name inFamily 3-122212Example:Number of ValanceElectronsElement’s ls3 (atomic #58-71, 90103)M 2NNAny Nameatomicnumber 58-71,90-103Halogens17NM-1YYAny Name inFamily 177NobleGases18NM0NNAAny Name inFamily 1881M 1YNAHydrogen1Hydrogen9

GroupsO Alkali MetalsO Alkali Earth MetalsO Boron GroupO Carbon GroupO HydrogenO Halogen sO Inner Transition MetalsO MetaloidsO Nitrogen GroupO Noble GassesO Oxygen GroupO Transition MetalsGroup 11.00794H.12Hydrogen6.9419.01218Li.Be iumMagnesium39.098340.08344.95594547.88Periodic Table of the ElementsAtomic MassMass numbers in parenthesis are those of themost stable or most common isotopeHe um18.9984038Carbon Nitrogen26.9815455.847.6Boron1614.00675Transition 964.00260SiSymbolAtomic 2Scandium88.9059Titanium Vanadium Chromium .906NickelCopper106.42107.868Zinc112.41Gallium um Strontium132.905137.33YttriumZirconium Niobium MolybdenumTechnetium Ruthenium Rhodium n207.2Antimony )Barium Lanthanum nciumRadiumActinum Rutherfordium Dubnium Seaborgium Bohrium140.12Lanthanoid Series144.24Mercury(277?)Uuu 118112Hassium Meitnerium Dormstadtium Unununium Ununbium Ununtrium 36465666768697071.Cerium232.038Actinoid Series140.908Gold(272?)Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium 79899100101102103. Protactinium Uranium Neptunium Plutonium AmericiumThorium10CuriumBerkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium

Orbital DiagramsEnergy Level Indicates relative sizes and energies of atomic orbitals. Whole numbers, rangingfrom 1 to 7. The energy level is represented by the letter n.Sublevels Number of sublevels present in each energy level is equal to the n. Sublevels are represented by the letter l. In order of increasing energy:s p d fOrbitals Represented by ml S Sublevel- Only 1 orbital in this sublevel level. P Sublevel- 3 orbitals present in this sublevel.o Each orbital can only have 2 electrons. D Sublevel- 5 orbitals present in this sublevel. F Sublevel- 7 orbitals present in this sublevel.11

1Total # ofOrbitals inEnergy Level1Total # ofElectrons inEnergy Level2s, p1, 3483s, p, d1, 3, 59184s, p, d, f1, 3, 5, 71632Energy LevelSublevelsPresent# of Orbitals1s2Orbital Diagrams An orbital diagram shows the arrangement of electrons in an atom. The electrons are arranged in energy levels, then sublevels, then orbitals.Each orbital can only contain 2 electrons. Three rules must be followed when making an orbital diagram.o Aufbau Principle- An electron will occupy the lowest energy orbitalthat can receive it. To determine which orbital will have the lowest energy, look tothe periodic table.o Hund’s Rule- Orbitals of equal energy must each contain oneelectron before electrons begin pairing.o Pauli Exclusion Principle- If two electrons are to occupy the sameorbital, they must be spinning in opposite directions. Energy Levels (n) determined by the ROWS Sub Levels (s,p,d,f)- determined by the sections Orbitals - determined by the # of columns per sublevel12

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Orbital 22p63s23p1 S1s22s22p63s23p4 As1s22s22p63s23p64s23d104p3 Mn1s22s22p63s23p64s23d5 N Sc1s22s22p31s22s22p63s23p64s23d114

Name: Date: Period: Honor Code:Electron Configuration WSGive the COMPLETE electron configuration for the following elements:226261. Ar 1s 2s 2p 3s 3p226232. P 1s 2s 2p 3s 3p22626263. Fe 1s 2s 2p 3s 3p 4s 3d226264. Ca 1s 2s 2p 3s 3p 4s222626210226262555. Br 1s 2s 2p 3s 3p 4s 3d 4p6. Mn 1s 2s 2p 3s 3p 4s 3d7. U 1s22s22p63s23p64s23d104p6 5s24d105p66s24f145d106p67s25f36d115

Electron Configurations and Oxidation States Electron configurations are shorthand for orbital diagrams. The electrons arenot shown in specific orbitals nor are they shown with their specific spins. Draw the orbital diagram of oxygen: The electron configuration should be:1s22s22p4 Manganese (25)1s22s22p63s23p64s23d5 Arsenic (33)1s22s22p63s23p64s23d104p3 Promethium (61)1s22s22p63s23p64s23d104p65s24d105p66s24f4 The Noble Gas shortcut can be used to represent the electron configurationfor atoms with many electrons. Noble gases have a full s and p and thereforecan be used to represent the inner shell electrons of larger atoms. For example: Write the electron configuration for Lead. Write the electron configuration for Xenon. Substitution can be used: Manganese (25)Mn [Ar] 4s23d5 Arsenic (33)As [Ar] 4s23d104p3 Promethium (61)Pm [Xe] 6s24f45d116

Valence electrons, or outer shell electrons, can be designated by the s and psublevels in the highest energy levels Write the noble gas shortcut for BromineBr [Ar]4s23d104p5 Write only the s and p to represent the valence level.Br 4s24p5 This is the Valence Configuration. Bromine has 7 valence electrons. Silicon[Ne] 3s23p23s23p24 valence electrons Uranium[Rn] 7s25f46d17s22 valence electrons Lead[Xe] 6s24f145d106p26s26p24 valence electronsOctet Rule and Oxidation States The octet rule states the electrons need eight valence electrons in order toachieve maximum stability. In order to do this, elements will gain, lose or shareelectrons. Write the Valence configuration for oxygenO 2s22p4- 6 valence electrons Oxygen will gain 2 electrons to achieve maximum stability-2O 2s22p6- 8 valence electronso Now, oxygen has 2 more electrons than protons and the resulting charge ofthe atom will be -2o The symbol of the ion formed is now O-2. Elements want to be like the Noble Gas family, so they will gain or lose electronsto get the same configuration as a noble gas. When an element gains or losses an electron, it is called an ion . An ion with a positive charge is a cation (lost electrons) . An ion with a negative charge is an anion (gained electrons) .17

-( 2)18

Electron Configuration and Oxidation States WorksheetGive the noble gas shortcut configuration for the following elements:1

chemical properties known as the _ Periodic Law _. . Nonmetals: Explain how the bonds between the atoms are highly localized, causing each of the properties below. Again, emphasize that these are general properties and may not be true for all . o Note: These are general properties and may vary from transition metal to transition metal!

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