Unit 3 Notes: Periodic Table Notes John Newlands proposed an organization system based on increasing atomic mass in 1864. He noticed that both the chemical and physical properties repeated every 8 elements and called this the Law of Octaves. In 1869 both Lothar Meyer and Dmitri Mendeleev showed a connection between 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 by increasing 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 undiscovered elements. After many new elements were discovered, it appeared that a number of elements were out of order based on their Properties. In 1913 Henry Mosley discovered that each element contains a unique number of Protons. By rearranging the elements based on Atomic Number, the problems with the Periodic Table were corrected. This new arrangement creates a periodic repetition of both physical and chemical properties known as the Periodic Law. Periods are the Rows Valence electrons across a period are in the same energy level Groups/Families are the Columns There are equal numbers of valence electrons in a group. 1
When elements are arranged in order of increasing _Atomic Number_, there is a periodic repetition of their physical and chemical properties Family (Group): Columns (vertical) ; tells the number of electrons in the _Outer Energy level, called Valence Electrons (only for representative elements) Period (Series): Rows (horizontal) ; tells the number of Energy Levels an atom has; the number of electrons Increases across a period Representative Elements: Groups 1A through 8A _ (called the s and p blocks) (Columns 1, 2, 13, 14, 15, 16, 17, and 18) Valence Electrons: e- in the outer most energy level ; farthest away from the nucleus (protons) ; the e- with the most reactive Energy; the e- involved with Bonding (transferring or sharing) Metals: most of the periodic table, located to the Left of the stair-step Properties- good conductors of _heat_ and _Electricity_; they also are Malleable ; Ductile ; _ High Density, BP and MP Nonmetals: to the Right of the stair-step, located in the upper corner of P.T._ Although five times more elements are metals than nonmetals, two of the nonmetals hydrogen and helium make up over 99 per cent of the observable Universe Properties- mostly _ Brittle, but a few _low luster and _poor conductors ; they have _ low density, low Melting Point and Boiling Point 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 between metals and nonmetals. This makes them good for manufacturing computer chips. 2
Valence electrons Valence electrons the electrons that are in the highest (outermost) energy level that level is also called the valence shell of the atom they are held most loosely The number of valence electrons in an atom determines: The properties of the atom The way that atom will bond chemically As a rule, the fewer electrons in the valence shell, the more reactive the element is. When an atom has eight electrons in the valence shell, it is stable. Our discussion of valence electron configurations leads us to one of the cardinal tenets of chemical bonding, the octet rule. The octet rule states that atoms become especially stable when their valence shells gain a full complement of valence electrons. For example, Helium (He) and Neon (Ne) have eight outer valence electrons in their outer shells which means it is completely filled, so they have a tendency to neither gain or lose electrons. Therefore, Helium and Neon, two of the so-called Noble gases or Inert gases Group # Group Name # of valence electrons 1 Alkali Metals 1 2 Alkaline Earth Metals 2 3-12 Transition Metals 1 or 2 13 Boron Group 3 14 Carbon Group 4 15 Nitrogen Group 5 16 Oxygen Group 6 17 Halogens 7 18 Noble Gases 8 The number of valence electrons increases as you go across the periodic table from left to right. 3
Element Lithium Germanium Sulfur Symbol Li Ge S Group # 1A(1) 4A(14) 6A(16) # of valence e- 1 4 6 Period # 2 4 3 # of E levels 2 4 3 Type of element M ML NM Periodic 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- are _added_but _pulled closer to the nucleus - as you go down a group, you add energy levels 2. Ionization Energy: the amount of E needed to _remove _ an electron - Increases from left to right across a period - Decreases from top to bottom down a group Why? 4
- as you go across a period, e- feel stronger attraction from nucleus (protons), _Energy to remove e-, Ionization E necessary as you go down a group, Energy_, _Decreases_ to remove outermost 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 group Why? - as you go across a period, e- feel more attraction from nucleus _Protons 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 electrons 5
- smaller than corresponding neutral atom Why? - fewer e-, so it s _easier_ for protons to pull in remaining e- Anions: Negative ions; nonmetal atoms that _gain_ electrons - larger than corresponding neutral atom Why? - _more_ e-, so it s harder_ for protons to pull in outermost e- Shielding: The ability of the _inner (lower levels)_ electrons to _shield (reduce)_ the pull of 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 the position of an electron, cannot determine both. Schrödinger's Equation Developed an equation that treated the hydrogen atom's electron as a wave. o Only limits the electron's energy values, does not attempt to describe the electron'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 position of an electron. Orbitals The location of an electron in an atom is described with 4 terms. 6
o Energy Level Described by intergers. The higher the level, the more energy an electron has to have in order to exist in that region. o Sublevels energy levels are divided into sublevels. The # of sublevels contained 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 2 electrons. o Spin Electrons can be spinning clockwise (+) or counterclockwise ( ) within the orbital. Periodic Table Activity: Complete the table on page 21 with the information found on pages 18 20. When complete color each group in a 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 ranked elements by increasing atomic mass. Moseley (1911): Periodic table arranged by atomic number Top table: Metals, nonmetals, and metalloids Metals: Explain the electron sea theory, and as you explain each of the properties below, discuss how they are explained by the electron sea theory. Also make sure to explain that these are general properties and may not be true for all metals. o Malleable: Can be pounded into sheets. o Ductile: Can be drawn into wires o Good conductors of heat and electricity o High density (usually) o High MP and BP (usually) o Shiny o Hard 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 nonmetals. o Brittle o Poor conductors of heat and electricity o Low density o Low MP and BP (many are gases)! Metalloids: The bonding in metalloids is between that of metals and nonmetals, so metalloids have properties 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. 7
o Metalloids are semiconductors of electricity somewhere between metals and nonmetals. This makes 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 o o Elements within families/groups tend to have similar physical and chemical properties. They have similar chemical and physical properties because they have similar electron configurations. Example: Li = [He] 2s 1, Na = [Ne] 3s 1 each has one electron in the outermost energy level. Explain that s and p electrons in the outermost energy level are responsible for the reactions that 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 that they are only relevant for s and p electrons. Do several examples. Periods: Elements in the same rows of the periodic table o 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 the elements within a period similar to one another. The closer elements are to each other in the same period, the closer are their chemical and 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, usually refers 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 metals o Most reactive group of metals o Flammable in air and water o Form ions with +1 charge o 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 metals o Reactive, but less so than alkali metals o React in air and water (show Ca reacting in water) o Form ions with +2 charge o Low MP and BP, but higher than alkali metals (MP of Ba= 302º C, Mg = 649ºC o Soft, but harder than alkali metals o Low density, but higher than that of alkali metals (Ca = 1.55, Mg = 1.74). Groups 3 12: (Outer) transition metals o 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 8
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, jewelry Inner Transition Metals: Lanthanides and actinides o 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 Group Group 14: Carbon Group Group 15: Nitrogen Group Group 16: Oxygen Group Group 17: Halogens o 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 X 2 o Form ions with 1 charge o Used in water treatment and chemical production Cl 2 was used as a chemical weapon in World War I. Group 18: Noble Gases o Highly unreactive o 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 element o Diatomic H 2 N 2 O 2 F 2 Cl 2 Br 2 I 2 o o Can form either a +1 or 1 charge Relatively unreactive unless energy is added (under most conditions) it can form explosive mixtures with oxygen (as it did in the Hindenburg explosion) 9
Groups on the Periodic Table Summary Sheet: Group Examples of Words used Location on Periodic Table Group 1, Group 3-12, etc Metals, Non-Metals, Metalloids? Common Charge(s)? Metal +1 Reactivity Highly reactive, unreactive Interesting Information It can be cut with a plastic knife Example: Element s Name Number of Valance Electrons Alkali Metals 1 M +1 Y N Any Name in Family 1 1 Alkaline Earth Metals 2 M +2 Y N Any Name in Family 2 2 Transition Metals (Outer) 3-12 M +2 N N Any Name in Family 3-12 2 Inner Transition Metals 3 (atomic # 58-71, 90-103) M +2 N N Any Name atomic number 58-71, 90-103 2 Halogens 17 NM -1 Y Y Any Name in Family 17 7 Noble Gases 18 NM 0 N NA Any Name in Family 18 8 Hydrogen 1 M +1 Y NA Hydrogen 1 10
P E R I O D Li 3 Lithium 22.98977 Na 11 Sodium 39.0983 K 19 Potassium 85.4678 Rb 37 Be 4 Beryllium 24.305 Mg 12 Periodic Table of the Elements Magnesium 3 4 5 6 7 8 9 10 11 12 40.08 44.9559 47.88 50.9415 51.996 54.9380 55.847 58.9332 58.69 63.546 65.39 Ca 20 Calcium 87.62 Sr 38 Sc 21 Ti 22 Scandium Titanium 88.9059 91.224 Y 39 Groups O Alkali Metals O Alkali Earth Metals O Boron Group O Carbon Group O Hydrogen Group 1 1.00794 Atomic Mass 28.0855 Mass numbers in parenthesis are those of the most stable or most common isotope. H Symbol Si 1 Atomic Number 14 Nonmetals Hydrogen 2 Silicon Name 13 14 15 16 17 6.941 9.01218 O Halogen s 10.81 12.01 14.0067 15.9994 18.998403. O Inner Transition Metals Metals..... O Metaloids O Nitrogen Group O Noble Gasses O Oxygen Group O Transition Metals Zr 40 V 23 Cr 24 Mn 25 Fe 26 Vanadium Chromium Manganese Iron 92.9064 95.94 (98) 101.07 Nb 41 Mo 42 Transition Elements Tc 43 Ru 44 Co 27 Cobalt 102.906 Rh 45 Ni 28 Nickel 106.42 Pd 46 Cu 29 Copper 107.868 Ag 47 Zn 30 Zinc 112.41 Cd 48 B 5 Boron 26.98154 Al 13 C 6 Carbon 28.0855 Si 14 Aluminum Silicon 69.72 72.59 Ga 31 Gallium 114.82 In 49 Ge 32 N 7 Nitrogen 30.97376 P 15 Germanium Arsenic 118.71 121.75 Sn 50 O 8 Oxygen 32.06 S 16 Phosphorus Sulfur 74.9216 78.96 As 33 Sb 51 Se 34 Selenium 127.60 Te 52 F 9 Fluorine 35.453 Cl 17 Chlorine 79.904 Br 35 Bromine 127.60 I 53. 18 4.00260 He 2 Helium 20.179 Ne 10 Neon 39.948 Ar 18 Argon 83.80 Kr 36 Krypton 131.29 Xe 54. Rubidium 132.905 Strontium 137.33 Yttrium 138.906 Zirconium Niobium 178.49 180.948 Molybdenum Technetium Ruthenium Rhodium 183.85 186.207 190.2 192.22 Palladium 195.08 Silver 196.967 Cadmium 200.59 Indium 204.383 Tin 207.2 Antimony 208.980 Tellurium (209) Iodine (210) Xenon (222) Cs 55 Ba 56 La 57 Hf 72 Ta 73 W 74 Re 75 Os 76 Ir 77 Pt 78 Au 79 Hg 80 Tl 81 Pb 82 Bi 83 Po 84 At 85 Rn 86 Cesium (223) Barium 226.025 Lanthanum 227.028 Hafnium (261) Tantalum (262) Tungsten (263) Rhenium (262) Osmium (265) Iridium (266) Platinum (269) Gold (272?) Mercury (277?) Thallium (?) Lead (289?) Bismuth Polonium (289?) Astatine Radon (293?) Fr 87 Ra 88 Ac 89 Rf 104 Db 105 Sg 106 Bh 107 Hs 108 Mt 109 Ds 110 Uuu 111 Uub 112 Uut 113 Uuq 114 Uuh 116 Uuo 118 Francium Radium Actinum Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Dormstadtium Unununium Ununbium Ununtrium Ununquadium Ununhexium Ununoctium 140.12 140.908 144.24 (145) 150.36 151.96 157.25 158.925 162.50 164.930 167.26 168.934 173.04 174.967 Lanthanoid Series Ce 58 Pr 59 Nd 60 Pm 61 Sm 62 Eu 63 Gd 64 Tb 65 Dy 66 Ho 67 Er 68 Tm 69 Yb 70 Lu 71 Cerium 232.038 Praseodymium Neodymium Promethium Samarium 231.036 238.029 237.048 (244) Europium (243) Gadolinium Terbium (247) (247) Dysprosium Holmium (251) (252) Erbium (257) Thulium (258) Ytterbium (259) Lutetium (260) Actinoid Series Th 90 Pa 91 U 92 Np 93 Pu 94 Am 95 Cm 96 Bk 97 Cf 98 Es 99 Fm 100 Md 101 No 102 Lr 103 Thorium Protactinium Uranium Neptunium Plutonium Americium 11 Curium Berkelium CaliforniumEinsteinium Fermium Mendelevium Nobelium Lawrencium
Orbital Diagrams Energy Level Indicates relative sizes and energies of atomic orbitals. Whole numbers, ranging from 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<f Orbitals Represented by m l 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. 12
Energy Level Sublevels Present # of Orbitals Total # of Orbitals in Energy Level Total # of Electrons in Energy Level 1 s 1 1 2 2 s, p 1, 3 4 8 3 s, p, d 1, 3, 5 9 18 4 s, p, d, f 1, 3, 5, 7 16 32 Orbital 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 orbital that can receive it. To determine which orbital will have the lowest energy, look to the periodic table. o Hund s Rule- Orbitals of equal energy must each contain one electron before electrons begin pairing. o Pauli Exclusion Principle- If two electrons are to occupy the same orbital, 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 sublevel 13
There are two ways of representing the electron distribution among the various orbitals of an atom: 1. Electron configuration An electron configuration consists of the symbol for the occupied subshell with a superscript indicating the number of electrons in the subshell. The electron configuration for sodium (atomic number 11) is 1s 2 2s 2 2p 6 3s 1 The large numbers represent the energy level. The letters represent the sublevel. The superscript numbers indicate the number of electrons in the sublevel. 2. Orbital diagram An orbital diagram consists of a box representing each orbital and a half arrow representing each electron. The orbital diagram below is for sodium (atomic number 11) Condensed Configurations For large atoms, showing all the electrons with an electron configuration or orbital diagram can become quite complex. Since it is the outermost electrons that are largely responsible for chemical behavior, we can condense the electron configuration and orbital diagram to focus on those electrons. Outer-shell electrons, those involved in chemical bonding, are called valence electrons. Those electrons below the outer shell, inner-shell electrons, are usually referred to as core electrons. The electron configuration and orbital diagram can be condensed by beginning with the nearest (before the atom) noble gas symbol in brackets to represent the core electrons, then showing the valence electrons as usual. Sodium's complete electron configuration is 1s 2 2s 2 2p 6 3s 1 The same electron configuration in condensed form becomes [Ne]3s 1 The complete orbital diagram for sodium is The same orbital diagram in condensed form becomes 14
Orbital Diagrams 1s 2 2s 2 2p 4 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 1s 2 2s 2 2p 6 3s 2 3p 1 S 1s 2 2s 2 2p 6 3s 2 3p 4 As 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 Mn 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 N 1s 2 2s 2 2p 3 Sc 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 15
Name: Date: Period: Honor Code: Electron Configuration WS Give the COMPLETE electron configuration for the following elements: 1. Ar = 1s 2 2s 2 2p 6 3s 2 3p 6 2. P = 1s 2 2s 2 2p 6 3s 2 3p 3 3. Fe 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 4. Ca = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 5. Br = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 6. Mn = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 7. U = 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 3 6d 1 16
Electron Configurations and Oxidation States Electron configurations are shorthand for orbital diagrams. The electrons are not shown in specific orbitals nor are they shown with their specific spins. Draw the orbital diagram of oxygen: The electron configuration should be: 1s 2 2s 2 2p 4 Manganese (25) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 Arsenic (33) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 Promethium (61) 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 4 5d 1 The Noble Gas shortcut can be used to represent the electron configuration for atoms with many electrons. Noble gases have a full s and p and therefore can 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] 4s 2 3d 5 Arsenic (33) As = [Ar] 4s 2 3d 10 4p 3 Promethium (61) Pm = [Xe] 6s 2 4f 4 5d 1 17
Valence electrons, or outer shell electrons, can be designated by the s and p sublevels in the highest energy levels Write the noble gas shortcut for Bromine Br = [Ar]4s 2 3d 10 4p 5 Write only the s and p to represent the valence level. Br = 4s 2 4p 5 This is the Valence Configuration. Bromine has 7 valence electrons. Silicon [Ne] 3s 2 3p 2 3s 2 3p 2 4 valence electrons Uranium [Rn] 7s 2 5f 4 6d 1 7s 2 2 valence electrons Lead [Xe] 6s 2 4f 14 5d 10 6p 2 6s 2 6p 2 4 valence electrons Octet Rule and Oxidation States The octet rule states the electrons need eight valence electrons in order to achieve maximum stability. In order to do this, elements will gain, lose or share electrons. Write the Valence configuration for oxygen O = 2s 2 2p 4 6 valence electrons Oxygen will gain 2 electrons to achieve maximum stability O 2 = 2s 2 2p 6 8 valence electrons o Now, oxygen has 2 more electrons than protons and the resulting charge of the atom will be 2 o 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 electrons to 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). 18
(-2) 19
Electron Configuration and Oxidation States Worksheet Give the noble gas shortcut configuration for the following elements: 1. Pb 2. Eu Eu = [Xe] 6s 2 4f 6 5d 1 3. Sn Sn = [Kr] 5s 2 4d 10 5p 2 4. As As = [Ar] 4s 2 3d 10 4p 3 Give ONLY the outer shell configuration for the following elements: 1. Ba 6s 2 2. Po 6s 2 6p 4 3. S 3s 2 3p 4 4. F 2s 2 2p 5 Au 6s 2 Cm 7s 2 20
Periodic Trends- Review Notes Shielding: As you go down the periodic table, the number of shells increases which results in greater electron electron repulsion. o The more shells there are, the further from the nucleus the valence electrons are. o Therefore, more shielding means the electrons are _Less_ attracted to the nucleus of the atom. Atomic Radius is defined as half the distance between adjacent nuclei of the same element. o As you move DOWN a group an entire energy level is added with each new row, therefore the atomic radius increases_(larger)_. o As you move LEFT-TO-RIGHT across a period, a proton is added, so the nucleus more strongly attracts the electrons of a atom, and atomic radius decreases (smaller). Ionic Radius is defined as half the distance between adjacent nuclei of the same ion. o For cation an electron was lost and therefore the ionic radius is smaller than the atomic radius. o For anion an electron was gained and therefore the ionic radius is larger than the atomic radius. o As you move down a group an entire energy level is added, therefore the ionic radius increases. o As you move left-to-right across a period, a proton is added, so the nucleus more strongly attracts the electrons of a atom, and ionic radius decreases. 21
However! This occurs in 2 sections. The cations form the first group, and the anions form the second group. Isoelectronic Ions: Ions of different elements that contain the same number of electrons. Ionization energy is defined as the energy required to remove the first electron from an atom. o As you move down a group atomic size increases, allowing electrons to be further from the nucleus, therefore the ionization energy decreases. o As you move left to right across a period, the nuclear charge increases, making it harder to remove an electron, thus the ionization energy increases. Electronegativity is defined as the relative ability of an atom to attract electrons in a electron cloud by the nucleus. o As you move down a group atomic size increases, causing available electrons to be further from the nucleus, therefore the electronegativity decreases. o As you move left to right across a period, the nuclear charge increases, making it easier to gain an electron, thus the electronegativity increases. Reactivity is defined as the ability for an atom to react/combine with other atoms. o With reactivity we must look at the metals and non metals as two separate groups. Metal Reactivity metals want to lose electrons and become cations o As you move down a group atomic size increases, causing valence electrons to be further from the nucleus, therefore these electron are more easily lost and reactivity decreases. o As you move left to right across a period, the nuclear charge increases, making it harder to lose electrons, thus the reactivity increases. Non metal Reactivity non metals want to gain electrons and become anions o As you move down a group atomic size increases, making it more difficult to attract electrons, therefore reactivity decrease. o As you move left to right across a period, the nuclear charge increases, making it easier to attract electrons, thus the reactivity increases. 22
Periodic Table : What is the Trend? Definition Trend Atomic Size (Atomic Radius) Radius is defined as half the distance between adjacent nuclei of the same element. Electronegativity Ability of an atom to attract electrons See above Ionization Energy Energy required to remove an e- from an atom See above Metal Having the characteristics of a metal Non Metal Having the characteristics of a nonmetal Shielding This describes the decrease in attraction between an electron and the nucleus in any atom with more than one electron shell. As more electrons are between the valence electrons and the nucleus the more shielded the outer electrons are from the nucleus. 23
Periodic Trends Worksheet 1. Explain why a magnesium atom is smaller than both sodium AND calcium. It is smaller than Na because it has more protons and smaller than Ca because is has less energy levels. 2. Would you expect a Cl - ion to be larger or smaller than a Mg 2+ ion? Explain You would expect Cl - to be larger because of the electron to proton ratio and Mg +2 now has the second energy level as its outer level. 3. Explain why the sulfide ions (S 2- ) is larger than a chloride ion (Cl - ). It is larger because of the electron to proton ratio. S -2 has two more electrons than protons and Cl - only has one more. 4. Compare the ionization energy of sodium to that of potassium and EXPLAIN. It would require less ionization energy for K to loss an electron than Na. K has more energy levels and the valence electrons are further from the nucleus. 5. Explain the difference in ionization energy between lithium and beryllium. They are the same energy level, but Be is slightly smaller so the valence electron are closer to the nucleus so it would have a higher ionization energy. 6. Order the following ions from largest to smallest: Ca 2+, S 2-, K +, Cl -. Explain your order. S 2-, Cl -, K +, Ca 2+ It is because of the electron to proton ratio. 7. Rank the following atoms/ions in each group in order of decreasing radii and explain your ranking for each (larger to smaller). a. I, I - I -, I b. K, K + K, K + c. Al, Al 3+ Al, Al 3+ 8. Which element would have the greatest electron affinity: B or O? Explain. Hint: a positive electron affinity means that the element wants to form a negative charge. It would be O. Because O wants to gain two electrons to achieve noble gas configuration. While B wants to lose three electrons. 24
Unit 3 Test Review: Give the Orbital Diagram for the following elements: 1. Chromium 2. Nitrogen Give the COMPLETE electron configuration for the following elements: 3. Argon 1s 2 2s 2 2p 6 3s 2 3p 6 4. Phosphorous 1s 2 2s 2 2p 6 3s 2 3p 3 Give the Noble Gas electron configuration for the following elements: 5. Plutonium Pu = [Rn] 7s 2 5f 5 6d 1 6. Mercury Hg = [Xe] 6s 2 4f 14 5d 10 7. Complete the table. Element Total # of electrons Valence Configuration Gain or Lose e - How Many? Ion Symbol New Valence Configuration Total # of e - Phosphorous 15 3s 2 3p 3 G 3 P -3 3s 2 3p 6 18 Chlorine 17 3s 2 3p 5 G 1 Cl -1 3s 2 3p 6 18 Cesium 55 6s 1 L 1 Cs +1 5s 2 5p 6 54 Lithium 3 2s 1 L 1 Li +1 1s 2 2 Give the 4 quantum numbers for the last electron of the following elements: 8. Phosphorous n=3, l=1, ml=1, ms= +1/2 9. Manganese n= 3, l=2, ml=2, ms= +1/2 10. Silver n= 4, l=2, ml=1, ms= -1/2 11. Promethium n= 4, l=3, ml=0, ms= +1/2 12. Iodine n= 5, l=1, ml=0, ms= -1/2 25
Determine if the following sets of quantum numbers would be allowed in an atom. If not, explain why and if so, identify the corresponding atom. 1 13. n = 2, l = 1, m l = 0, m s = + Yes 2 14. n = 4, l = 0, m l = 2, m s = - 2 1 No, because orbital 0 only has sub level 0 15. n = 1, l = 1, m l = 0, m s = + 2 1 No, because l must be one less than n Give the element with the LARGER radius, ionization energy, electronegativity and reactivity. ELEMENTS ATOMIC RADIUS IONIZATION ENERGY ELECTRONEGATIVITY Sodium and Aluminum Na Al Al Chlorine and Iodine I Cl Cl Oxygen and Fluorine O F F Magnesium and Calcium Ca Mg Mg Circle the element / ion with the larger radius. 16. Mg or Mg 2+ 18. S or S 2-20. N 3- or F - 17. Sr 2+ or Br - 19. Cl - or Mg 2+ 21. B or F For each of the following families, give their relative reactivity, the number of valence electrons, and at least one additional piece of information (such as how they are found in nature or what other group the generally react with). 22. Alkaline Earth Metals Very reactive, s 2, most in the earth s crust 23. Alkali Metals Very reactive, s 1, they will react in air and with water 24. Halogens Very reactive, s 2 p 5, they form salts 25. Noble Gases non reactive, s 2 p 6, they are gases at room temperature 26
Matching (1 point each): Match the description in Column B with the correct term in Column A. Write the letter in the blank provided. Each term matches with only one description, so be sure to choose the best description for each term. Not all descriptions will be used. Column A Column B A 26. Alkaline Earth Metal D 27. Transition Metal F 28. Alkali Metal A. located in the second column B. solid or liquid mixture of two or more metals C. horizontal row of elements I _ 29. Noble Gases D. located in columns 3-12 K 30. Halogen C 31. Period E 32. Ionization Energy H 33. Valence Electron E. energy required to remove an e- from an atom F. located in the first column G. ability of an atom to attract electrons G 34. Electronegativity I. located in column 18 J 35. Group H. an electron in the outermost shell of an atom J. vertical column of elements K. located in column 17 _D_ 36. Elements in a family or group in the periodic table often share similar properties because a. They look alike. b. They are found in the same place on Earth. c. They have the same physical state. d. Their atoms have the same number of electrons in their outer energy level. _B 37. Groups 3-12 are commonly referred to as a. Alkali metals. b. Transition metals. c. Lanthanides. d. Actinides. C_ 38. Which of the following elements has the highest electronegativity? a. Ca b. Cu c. Br d. As _B_ 39. An atom is neutral because the number of a. Electrons equals the number of neutrons. b. Electrons equals the number of protons. c. Protons equals the number of neutrons. d. None of the above. 27