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Chapter Summary: pg. 62

Ch. 2 – The Periodic Table 2.1 The Periodic Table 

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1869: Dmitri Mendeleev (Russian chemist) o Published 1st version of his Periodic Table of the Elements  Showed that ordering the known elements according to atomic weight revealed a pattern of periodically recurring physical & chemical properties Periodic table has since been revised o Henry Moseley’s (physicist) work used to organize the elements based on increasing atomic number (the # of protons in an element) rather than atomic weight o Using revised table, many properties of elements that hadn’t been discovered could be predicted o Periodic table creates a visual representation of the periodic law  Periodic law: the chemical & physical properties of the elements are dependent, in a periodic way, upon their atomic numbers Modern periodic table arranges the elements into periods (rows) & groups/families (columns), based on atomic number o Periods:  7 periods representing the principal quantum numbers n = 1 through n = 7 for the s& p-block elements  Each period is filled sequentially  Each element is a given period has 1 more proton & 1 more electron than the element to its left (in their neutral states) o Groups contain elements that have the same electronic configuration in their valence shell & share similar chemical properties  Valence electrons: the electrons in the valence shell  Farthest from the nucleus  Have the greatest amt of potential energy  Can become involved in chemical bonds with the valence electrons of other atoms o Bc of higher PE & held less tightly by the nucleus o \ valence shell electrons largely determine the chemical reactivity & properties of the element  Elements with similar valence electron configurations generally behave in similar ways, as long as they’re the same type (metal, nonmetal, or metalloid) Roman numeral above each group represents the # of valence electrons elements in that group have in their neutral state o Roman number combined with the letter A or B to separate the elements into 2 larger classes: 1. A elements  Representative elements  Groups IA through VIIIA  Valence electrons in orbitals of either s or p subshells  The Roman numeral & letter designation determine the electron configuration o An element in Group VA has 5 valence electrons with the configurations s2p3 2. B elements  Nonrepresentative elements  Include: a) Transition elements: valence electrons in the s & d subshells b) Lanthanide & actinide series: valence electrons in the s & f subshells  May have unexpected electron configurations o Chromium  4s13d5 o Copper  4s13d10 Modern IUPAC ID system: groups are #ed 1 to 18 & aren’t subdivided into Group A & Group B elements

2.2 Types of Elements 1

1. Metals o Left side & middle of periodic table o Properties: a) Lustrous (shiny) solids  Except for mercury – liquid under standard conditions b) High melting points & densities  There are exceptions o Lithium – density ~ ½ that of water c) Ability to be deformed w/o breaking  Malleability: the ability of metal to be hammered into shapes  Ductility: the ability of metal to be pulled or drawn into wires d) Atomic level:  Low effective nuclear charge  Low electronegativity (high electropositivity)  Large atomic radius  Small ionic radius  Low ionization energy  Low electron affinity  *All these characteristics are manifestations of the ability of metals to easily give up electrons o Valence electrons of all metals are only loosely held to their atoms (they’re free to move)  Makes metals good conductors of heat & electricity o Include: a) Active metals  Valence electrons found in s subshell  Alkali & alkaline earth metals are both metallic in nature bc they easily lose electrons from the s subshell of their valence shells b) Transition metals (Group B elements)  Many have 2+ oxidation states (charges when forming bonds with other atoms)  Valence electrons found in s & d subshells  Some are relatively nonreactive – makes them ideal for the production of coins & jewelry o Copper, nickel, silver, gold, palladium, and platinum  Copper wire – exhibits luster, malleability, and ductility  Used as a wire bc it also exhibits good heat & electrical conductivity c) Lanthanide & actinide series of elements  Valence electrons found in s & f subshells 2. Nonmetals o Predominantly on upper right side of periodic table o Properties:  Generally brittle in solid state  Show little/no metallic luster  Atomic level:  High ionization energies  High electron affinities  High electronegativities  Small atomic radii  Large ionic radii  Usually poor conductors of heat & electricity  *All these characteristics = manifestations of the inability of nonmetals to easily give up electrons  Less unified in their chemical & physical properties than the metals o Carbon = stereotypical nonmetal  Retains solid structure but is brittle, nonlustrous, and generally a poor conductor of heat & electricity

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Charcoal composed of carbon

3. Metalloids (semimetals) o Stair-step (“staircase”) group of elements separating the metals & nonmetals o Include:  Boron (B)  Silicon (Si)  Germanium (Ge)  Arsenic (As)  Antimony (Sb)  Tellurium (Te)  Polonium (Po)  Astatine (At) o Share some characteristics with both metals & nonmetals  Electronegativities & ionization energies lie b/w those of metals & nonmetals  Physical properties (densities, melting points, and boiling points) vary widely – can be combinations of metallic & nonmetallic characteristics  Silicon (Si) has a metallic luster but is brittle & a poor conductor  Make good semiconductors due to their partial conductivity of electricity o Reactivities are dependent on the elements with which they’re reacting  Boron (B) behaves like a nonmetal when reacting with sodium (Na) & like a metal when reacting with fluorine (F)

2.3 Periodic Properties of the Elements  

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Organized to represent visually the periodicity of chemical & physical properties of the elements 3 key rules that control how valence electrons work in an atom : 1. Electrons & protons are added 1 at a time as from left to right across a period  As the positivity of the nucleus increases, the electrons surrounding the nucleus (incl. e–‘s in valence shell) experience a stronger electrostatic pull toward the center of the atom  Causes the electron cloud (outer boundary defined by valence shell electrons) to move closer & bind more tightly to the nucleus o Effective nuclear charge (Zeff): the electrostatic attraction b/w the valence shell electrons & the nucleus  A measure of the net positive charge experienced by the outermost electrons  This pull is somewhat mitigated by nonvalenced electrons that reside closer to the nucleus  For elements in the same period, Zeff increases from left to right  Factors that determine effective nuclear charge:

2. Principal quantum number increases by 1 for each element moving down a given group  Valence electrons are increasingly separated from the nucleus by a greater number of filled principle energy levels (aka inner shells)  Result of increased separation: reduction in the electrostatic attraction b/w the valence electrons & the positively charged nucleus  Outermost electrons held less tightly as the principal quantum number increases  Moving down a group – the increased shielding created by the inner shell electrons cancels the increased positivity of the nucleus  Zeff = more or less constant among the elements within a given group o Despite this, valence electrons are held less tightly to nucleus as move down a group  Due to the increased separation b/w valence electrons & the nucleus 3. Element can gain or lose electrons in order to achieve a stable octet formation representative of the noble (inert) gases (Group VIIIA or Group 18)  Octet rule  Elements (esp. ones with biological roles) tend to be most stable with 8 electrons in their valence shell Atomic & Ionic Radii A. Atomic Radius: the size of a neutral element o Equal to ½ distance b/w the centers of 2 atoms of an element that are briefly in contact with each other  Distance b/w 2 centers of circles in contact is akin to a diameter o Can’t be measured by examining a single atom  the electrons are constantly moving around, making it impossible to mark the outer boundary of the electron cloud o RULE #1 – Electrons & protons are added 1 at a time as from left to right across a period  Electrons added only to outermost shell & # of inner-shell electrons remains constant  increasing positive charge of the nucleus pulls the outer electrons more closely inward & holds them more tightly  Zeff increases left to right across a period  \ atomic radius decreases from left to right across a period o RULE #2 – increasing principal quantum number implies that valence electrons will be found farther away from nucleus bc # of inner shells is increasing, separating the valence shell from the nucleus  Atomic radius increases down a group (even though Zeff remains essentially constant)  Within each group: largest atom at bottom  Within each period: largest atom in Group IA (Group 1)  Largest atomic radius  cesium (Cs, 260 pm)  Smallest atomic radius  helium (He, 25 pm)  *Francium typically not considered bc it’s exceptionally rare in nature o Essentially opposite that of all other periodic trends  Others increase going up & to the right  Atomic radius increases going down & to the left

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Empirical Atomic Radius (pm)

Atomic Number

B. Ionic Radii: dependent on how the element ionizes based on its element type & group # o Must make 2 generalizations to determine ionic radii: 1. Metals lose electrons & become positive; nonmetals gain electrons & become negative 2. Metalloids can go in either directions, but tend to follow the trend based on which side of the metalloid they fall on  Silicon (Si) behaves more like a nonmetal  Germanium (Ge) behaves more like a metal  Boron (B) behaves as a metal & tellurium (Te) behaves as a nonmetal o Under varying conditions, these metalloids can have opposite behaviour o Nonmetals close to the metalloid line:  Group # dictates that they require more electrons than other nonmetals to achieve the electronic configuration seen in Group VIIIA (Group 18  Gain electrons while their nuclei maintain the same charge  Possess a larger ionic radius than their counterparts closer to Group VIIIA o Metals:  Trend is similar (^) but opposite  Closer to metalloid line:  Have more electrons to lose to achieve the electronic configuration seen in Group VIIIA o Ionic radius = dramatically smaller than that of other metals  Closer to Group IA:  Have fewer electrons to lose o Less drastic reduction in radius during ionization

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*Size of atoms & their ions in pm

Ionization Energy (IE ) (aka Ionization Potential)  The energy req’d to remove an electron from a gaseous species  Removing an electron from an atom always requires an input of heat  endothermic process  Greater the atom’s Zeff or closer the valence electrons are to the nucleus, the more tightly bound they are o Makes it more difficult to remove 1+ electrons, increasing the ionization energy  Ionization energy increases from left to right across a period & from bottom to top in a group  Subsequent removal of a 2nd or 3rd electron requires increasing amts of energy  the removal of more than 1 electron means that the electrons are being removed from an increasingly cationic (positive) species o First ionization energy: the energy necessary to remove the 1st electron o Second ionization energy: the energy necessary to remove the 2nd electron from the univalent cation (X+) to form the divalent cation (X2+) o So on… o Example:

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−¿ +¿ ( g ) +e ¿ ¿ Mg ( g ) → Mg

first ionization energy = 738 kJ/mol

−¿ 2+¿ ( g) +e ¿ +¿ ( g) → Mg ¿ ¿ Mg

second ionization energy = 1450 kJ/mol

Active metals: elements in Groups IA & IIA (Groups 1 & 2) (e.g., lithium & beryllium) o Bc such low ionization energies o Don’t exist naturally in their neutral forms  Always found in ionic compounds, minerals, or ores o Formation of a stable, filled valence shell is the result of either:  Loss of 1 electron from the alkali metals (Group IA)  Loss of 2 electrons from the alkaline earth metals (Group IIA) Halogens: Group VIIA (Group 17) elements o Don’t typically give up their electrons  generally anions in their ionic form

First ionization energy (eV)

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Values for second ionization energies are: o Disproportionally larger for Group IA monovalent cations (like Na +)  *Bc removing 1 electron from a Group IA metal results in a noble gas-like electron configuration o Generally not much larger for Group IIA or subsequent monovalent cations (like Mg+) Group VIIIA (Group 18) elements (noble/inert gases) = least likely to give up electrons o Already have a stable electron configuration  Unwilling to disrupt that stability by giving up an electron o Among the elements with the highest ionization energies First ionization energy (IE) will always be smaller than second IE, which will always be smaller than third IE o Degree to which the IE increases provides clues about the identity of the atom o If losing a certain # of electrons give an element a noble-gas like electron configuration, then removing a subsequent electron will cost much more energy

Electron Affinity  The energy dissipated by a gaseous species when it gains an electron o Exothermic process (–DHrxn) – expels energy in the form of heat  Reported as a positive number  electron affinity refers to the energy dissipated o If 200 kJ/mol of energy is released, DHrxn = –200 kJ/mol & electron affinity is 200 kJ/mol  Essentially the opposite concept of from ionization energy  The stronger the electrostatic pull (the higher the Zeff) b/w the nucleus & the valence shell electrons, the greater the energy release will be when the atom gains the electron o Electron affinity:  Increases across a period from left to right  Decreases in a group from top to bottom  Bc valence shell is farther away from the nucleus as principal quantum number increases  Groups IA & IIA (Groups 1 & 2): v low electron affinities o Prefer to give up electrons to achieve the octet configuration of noble gas in previous period  Group VIIA (Group 17) elements: v high electron affinities o Need to gain only 1 electron to achieve the octet configuration of the noble gases (Group VIIIA or Group 18) in same period o Halogens = most “greedy” group of elements in terms of electrons Electron  Group VIIIA (Group 18)affinity (noble gases): electron affinities on the order of zero (eV) o They already possess a stable octet & can’t readily accept an electron  Most metals have low electron affinity values

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Electronegativity  A measure of the attractive force that an atom will exert on an electron in a chemical bond  The greater the electronegativity of an atom, the more it attracts electrons within a bond  Electronegativity values are related to ionization energies: o The lower the ionization energy, the lower the electronegativity o The higher the ionization energy, the higher the electronegativity o Exceptions: 1st 3 noble gases  Despite their high ionization energies, they have negligible electronegativity bc they don’t often form bonds  Electronegativity value = relative measure o Different scales used to express it  Pauling electronegativity scale (most common scale)  Ranges from 0.7 for cesium (least electronegative/most electropositive) element to 4.0 for fluorine (most electronegative element) o Cs = largest, least electronegative, lowest ionization energy, least exothermic (lowest) electron affinity o F = smallest, most electronegative, highest ionization energy, most exothermic (highest) electron affinity  Electronegativity: o Increases across a period from left to right o Decreases in a group from top to bottom

Pauling Electronegativity Values

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Right Atomic radius ¯ Ionization energy  Electron affinity  Electronegativity 

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Bottom Atomic radius  Ionization energy ¯ Electron affinity ¯ Electronegativity ¯

*Atomic radius is always opposite the other trends *Ionic radius is variable

2.4 The Chemistry of Groups 1. Alkali Metals (IA) o Possess most of the classic physical properties of metals  Except: their densities are lower than those of other metals o Only 1 loosely bound electron in their outermost shells o Zeff values are v low  Largest atomic radii of all the elements in their respective periods  Low ionization energies  Low electron affinities  Low electronegativities o Easily lose 1 electron to form univalent cations o React readily with nonmetals (esp. the halogens, e.g., NaCl) o

Rxn of Sodium with Water Group IA Due to their high reactivity with water & air, most alkali metals are stored in mineral oilmetals react violently with water, forming strong bases

2. Alkaline Earth Metals (IIA) o Possess many properties characteristic of metals o Share most of the characteristics of alkali metals  Except: they have slightly higher Zeff’s  slightly smaller atomic radii o 2 electrons in valence shell  Both easily removed to form divalent cations o

Alkali metals + alkaline earth metals = active metals  They’re so reactive that they aren’t naturally found in their elemental (neutral) state

3. Chalcogens (VIA) o Eclectic group of nonmetals & metalloids o Not as reactive as halogens  Crucial for normal biological functions  

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Many of the molecules for metabolism utilize lighter nontoxic elements from the chalcogen group (oxygen & sulfur) Many of the heavier chalcogens are toxic metals

Each have 6 electrons in their valence electron shell Generally small atomic radii & large ionic radii  bc of proximity to the metalloids

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Oxygen = most important element in this group  It’s one of the primary constituents of water, carbohydrates, and other biological molecules Sulfur = important component of certain aa’s & vitamins Selenium = important nutrient for microorganisms & has role in protection from oxidative stress Remainder of group = primarily metallic & generally toxic to living organisms At high concentrations, many of these elements (no matter how biologically useful) can be toxic or damaging

4. Halogens (VIIA) *frequently tested on MCAT o Highly reactive nonmetals with 7 valence electrons o Desperate to complete their octets by gaining 1 additional electron o Physical properties are variable  Standard conditions: range from gaseous (F2 & Cl2) to liquid (Br2) to solid (I2) forms o Chemical reactivity is more uniform o V high electronegativities & electron affinities  especially reactive toward the alkali & alkaline earth metals o Fluorine (F) has highest electronegativity of all the elements o So reactive that they’re not naturally found in their elemental state but rather as ions (halides) or diatomic molecules 5. Noble o o o o o o o

(Inert) Gases (VIIIA) Have minimal chemical reactivity due to their filled valence shells High ionization energies Little/no tendency to gain or lose electrons No measurable electronegativities (at least for He, Ne, and Ar) Extremely low boiling points Exist as gases at room temp Commercial niche as light sources due to their lack of reactivity (e.g., fluorescent signs)

6. Transition Metals (IB to VIIIB or 3 to 12) o Considered to be metals  Low electron affinities  Low ionization energies  Low electronegativities o V hard o High melting & boiling points o Tend to be quite malleable o Good conductors due to the loosely held electrons that progressively fill the d-orbitals in their valence shells o Unique p...