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Ch 2.4 & 8: The Periodic Table Atomic Radius (pm) 250 200 150 100 50 0 0 5 10 Atomic Number 15 20 I. Creation of the Periodic Table The Person Who Organized the Elements Dmitri Mendeleev 1860’s, Russia Dmitri Mendeleev • taught chemistry in terms of properties (St. Petersburg University). • Mid 1800 - molar masses of elements were known….Wrote down the elements in order of INCREASING MASS • Found a pattern of repeating properties – elements with SIMILAR CHEMICAL AND PHYSICAL PROPERITES were GROUPED together (there were some discrepancies) Mendeleev’s Periodic Table Like A Calendar Mendeleev’s Periodic Table of Elements •Only 63 elements known at time •He left spaces on the table for undiscovered elements. •He predicted the properties of elements that had not yet been discovered. Example: Eka-aluminum He left a space for an undiscovered element below aluminum. He predicted the melting point to be low and the density to be 5.9 g/cm3. • It was later found and named Gallium with a density of 5.91 g/cm3 and melting point of 29.7oC. Gallium History Henry Moseley (1913, British) -Re-organized elements by increasing atomic number. -Resolved discrepancies in Mendeleev’s arrangement. Periodicity & Periodic Law 250 Atomic Radius (pm) PERIODICITY— reoccurrence (repetition) of properties at regular intervals. 200 150 100 50 0 PERIODIC LAW— when elements are arranged by atomic number, there is a periodic (i.e. cyclical or repeating) pattern in their physical and chemical properties. In other words, when elements are arranged by atomic number, elements with similar properties occur at regular intervals 0 5 10 Atomic Number 15 20 II. Organization of the Elements Metals, Nonmetals, & Metalloids • Metals (lustrous, conductive, malleable, ductile, lose e-/ form + CATIONS) • Nonmetals (dull, insulating, brittle solids or gases & liquids, gain e-/ form - ANIONS) • Metalloids (both metallic & non-metal properties) Metals Nonmetals Metalloids Row Series Period The periods (rows) of the periodic table indicate the valence shell. 1 2 3 4 5 6 7 • Horizontal rows are called Periods or Series There are 7 periods Column Group Family • Vertical columns are called Groups. • Elements are placed in columns by similar properties. • Also called families 1A • The elements in the groups in the s and p block are called the 2A representative elements3A 4A 5A 6A 7A 8A0 •Main Group or Blocks Representative Elements •Transition Metals •Inner Transition or Rare Earth Metals Chemical Reactivity • Families – Similar valence e- within a group result in similar chemical properties Valence Electrons • The valence electrons occupy the OUTERMOST (physically farthest away from the nucleus) shell of an atom. • Valence electrons play a key role in the chemical properties of an element. Periodicity Explained • The valence shell fills up in a regular pattern • The valence shell electron configuration REPEATS • The properties of atoms therefore repeat when placed in order of Atomic Number H 1 1s1 3 1s22s1 Li Na 11 1s22s22p63s1 1s22s22p63s23p64s1 K 19 Rb 37 Cs 55 Fr 87 1s22s22p63s23p64s23d104p65s1 1s22s22p63s23p64s23d104p6 5s24d10 5p66s1 1s22s22p63s23p64s23d104p65s24d105p6 6s24f145d106p67s1 He 2 1s 1s22s22p6 1s22s22p63s23p6 2 Ne 10 Ar 18 1s22s22p63s23p64s23d104p6 Kr 1s22s22p63s23p64s23d104p65s24d105p6 Xe 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p6 36 54 Rn 86 s- block s1 s2 • Alkali metals all end in s1 • Alkaline earth metals all end in s2 • really have to include He but it fits better later. • He has the properties of the noble gases (because the first shell is filled when 1s2) Transition Metals -d block d1 d2 d3 d5 d5 d6 d7 d8 d9 d10 The p-block p1 p 2 p3 p4 p5 p6 f - block • inner transition elements f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 Group 1- Alkali Metals • 1 valence electron (ns1), forms 1+ ion when it loses an e-. • VERY REACTIVE WITH WATER! Metal + H2O H2 + MOH (a base), reactivity increases from top to bottom • Note: Hydrogen, a nonmetal, is located in the first column because it has one valence electron. Note: • Sodium and Potassium are stored in oil to keep them from reacting with oxygen and water in the air. • Cesium is stored in glass tubes of argon gas( an inert gas). Group 2- Alkaline Earth Metals • Have 2 valence electrons. Beryllium • Form 2+ ions when 2 eare lost. Magnesium Calcium Strontium Barium Radium Group 2- Alkaline Earth Metals Reactions with Water • Be does not react with water. • Mg reacts with hot water. • Ca, Sr, Ba react easily with cold water. • Which way along the group does reactivity increase? Notes: • There is Mg in chlorophyll C55H72O5N4Mg • Calcium is in your bones, chalk, limestone, toothpaste, pearl (all as calcium carbonate). Transition Metals • Groups 3-12 are the transition metals • They have metallic properties such as high luster, good conductivity and are relatively non-reactive. • They form very colorful compounds. p-block metals (Groups 13-18) • These are main group elements. • The groups are sometimes named for the first element in the group. For example, the carbon family • Group 17 is the halogens. They are highly reactive and like to react with metals to form salts. Why would they like to react with group 1 elements? • Group 18 is the noble gases. They have filled valence shells and are non-reactive. PERIODIC & GROUP TRENDS Atomic Radius (aka Atomic Size) } Radius •First problem: Where do you start measuring. •The electron cloud doesn’t have a definite edge. •They get around this by measuring more than 1 atom at a time. •Commonly measure in pm (ex. I is 140 pm) • Atomic Radius = half the distance between two nuclei of a diatomic molecule. Trends in Atomic Size • Influenced by two factors. – # of Shells – “Shielding” • More shells- electrons are further away from nucleus. – Charge on nucleus – Zeff • More charge (# of protons) pulls electrons in closer. Atomic Radius • Atomic Radius K Atomic Radius (pm) 250 Na 200 Li 150 100 Ar Ne 50 0 0 5 10 Atomic Number 15 20 GROUP TREND: Atomic Size (atomic radii) H Li Na K Rb • As we go down a group • Each atom adds another shell • So the atoms logically get bigger. PERIODIC TREND: Atomic Size (atomic radii) • As you go across a period the radius gets smaller. 1. Same number of shells (shielding). 2. More protons = increased nuclear charge. 3. Increased attraction for outermost electrons. • So… the atoms should get smaller. Na Mg Al Si P S Cl Ar Increases DOWN Atomic Radius Decreases RIGHT Atomic Radius • Why larger going down? – Higher shells have larger orbitals – Shielding - core e- block or screen the attraction between the nucleus and the valence e- • Why smaller to the right? – Increased nuclear charge without additional shielding pulls e- in tighter Shielding Effect Valence + nucleus Kernel electrons block the attractive force of the nucleus from the valence electrons - - - Electron Shield “kernel” electrons Electrons Examples • Which atom has the larger radius? Be or Ba Ca or Br Example: Which is larger: a lithium atom or a fluorine atom? A lithium atom Which is larger: an arsenic atom or a sulfur atom? An arsenic atom First Ionization Energy –The amount of energy (J) required to completely remove one e- from a GASEOUS atom (a measure of how strongly the atom is holding its electrons) © 1998 LOGAL *Ionization—removal of electrons 1st: Atom A +I1 Ion A+ + free e2nd: Ion A+ + I2 Ion A+2 + free e- Ionization Energy 1st Ionization Energy (kJ) • First Ionization Energy He 2500 Ne 2000 Ar 1500 1000 Li 500 Na K 0 0 5 10 Atomic Number 15 20 e- e- GROUP trend: Ionization Energy (IE) • As you go down a group first IE decreases because…. • The electron is further away…because there are more shells between the nucleus and the e- thus…. • Shielding the e- from + nucleus. • IT IS EASIER (takes LESS energy) to remove an electron Symbol IE (J) Li Be B C N O F Ne 520 900 800 1086 1402 1314 1681 2080 Periodic trends: Ionization Energy (IE) e- e- • All the atoms in the same period have the same number of shells. • Same number of shells = same shielding, but the number of protons increase, therefore…. • Increasing nuclear charge …helps pull e- in tighter, therefore it is harder to remove. • So IE generally increases from left to right. Ionization Energy • Why opposite of atomic radius? – In small atoms, e- are close to the nucleus where the attraction is stronger. SMALL atoms will require MORE energy to remove e-; LARGE atoms will require LESS energy to remove e- • Why small exceptions within each period? – Stable e- configurations (s2, s2p3) don’t want to lose e- (will elevate the IE more than what’s expected from the trend) Examples • Which atom has the higher 1st I.E.? N or Bi Ba or Ne SUCCESSIVE Ionization Energy Large jump in I.E. occurs when a CORE e- is removed. (Write in the equations below) Mg Core e- 1st I.E. 2nd I.E. 3rd I.E. 736 kJ 1,445 kJ 7,730 kJ Successive Ionization Energy Large jump in I.E. occurs when a CORE eis removed. Al Core e- 1st I.E. 2nd I.E. 3rd I.E. 4th I.E. 577 kJ 1,815 kJ 2,740 kJ 11,600 kJ Electronegativity (EN) • The relative ability of an atom IN A BOND to attract electrons to itself (Linus Pauling) Which has greater electronegativity in each compound? Electronegativities of Some Elements Element F Cl O N S C H Na Cs Pauling scale 4.0 3.0 3.5 3.0 2.5 2.5 2.1 0.9 0.7 Note • Most electronegative element is F (EN 4.0) • Least electronegative stable element is Cs (EN 0.7) PERIODIC TREND: Electronegativity (EN) e- e- • From left to right atoms become smaller, with greater nuclear charge. e - are attracted by the + charged nucleus. •therefore EN will increase from left to right GROUP TRENDS: Electronegativity (EN) e- • Down a group atoms become larger, and have greater nuclear charge. • e - are attracted by the increased +charge but shielding also increases …which has a greater influence!! • therefore EN will decrease as you go down a group or family e- Electronegativities F Cs Probably the most important trend. (1) Useful for explaining chemical reactivity. (2) It is also useful for categorizing the elements. •Metals are electropositive and lose electrons, •nonmetals are electronegative and gain electrons, •metalloids have intermediate electronegativities Electronegativity & BONDING • When two atoms have similar electronegativities, they pull on the electrons equally and form a NONPOLAR COVALENT BOND. Example, H-H. • When the atoms are different, the atoms pull on the electrons unevenly, and form either a POLAR COVALENT BOND (if they’re not very different, and are still sharing, but not equally) or an IONIC BOND (if they are very different an the electron is completely transferred). Example, HCl and NaCl LARGE EN DIFFERENCES IONIC COMPOUNDS/ BONDING (full transfer of e-) Ions Ion • Atom or group of atoms that have lost or gained electrons & have a charge • Charges not balanced; differing number of protons and electrons Ions are formed when atoms gain or lose electrons. NEVER ADD OR SUBTRACT PROTONS! Cation (ca+ion) • positively charged ion • formed when neutral atoms LOSE electrons ex. Determine + : a. K + #p b. 2+ Zn c. +3 Al and #e Anion (Nion) • Negatively charged ion • formed when neutral atoms GAIN electrons ex. Determine : a. Cl b. -2 S c. 3P + #p and #e Ion charges can be determined from the periodic table: metals → cations non-metals → anions OCTET RULE – atoms will gain or lose e- or bond to form a full valence shell (usually 8 e-, except for He). ex. Cation or anion? Charge?: a. Nitrogen b. Sodium c. Sulfur Naming Cations 1. Group IA / 2A & Ag, Cd, Zn, & Al: Use element name 2. d-block and p-block: Most use Roman numerals for charge ex. Name the following cations: 2+ a. Mg b. 2+ Cu c. +3 Fe Naming Anions Add “-ide” to element name. ex. Name the following anions: 2a. O b. F c. -3 N Ionic Radius • Ionic Radius Cations (+) lose esmaller Anions (–) gain elarger © 2002 Prentice-Hall, Inc. Ion Sizes + Li,152 pm Li + , 78 pm 3e and 3p 2e and 3 p Forming a cation. • CATIONS are SMALLER than the atoms from which they come. • **Losing a valence electron is not • just the lose of an e• But the lose of the valence shell!!! Ion Sizes F, 71 pm F - , 133 pm 9e and 9p 10 e and 9 p Forming an anion. • ANIONS are LARGER than the atoms from which they come. • The addition of an electron increases the repulsion in the valence shell and so size INCREASES. • Trends in ion sizes are the same as atom sizes. Ionic Radii Trend • same general trend as atomic radii • anions are larger than neutral atom • cations are smaller than neutral atom Trends in Ion Sizes Figure 8.13 Examples • Which particle has the larger radius? S or 2S Alor 3+ Al Periodic & Group trends: Electron Affinity(EA) • The energy change associated with adding an electron to a gaseous atom. • e - are attracted by a + charge. – Therefore more protons more attraction – But remember “shielding” • Easiest to add to group 7A. • Gets them to full energy level. • Increase from left to right atoms become smaller, with greater nuclear charge. • Decrease as we go down a group. Periodic & Group trends: Electron e- Affinity(EA) • From left to right atoms become smaller, with greater nuclear charge. e - are attracted by the + charged nucleus. •therefore EA will increase from left to right e- Periodic & Group trends: Electron e- Affinity(EA) • Down a group atoms become larger, and have greater nuclear charge. • e - are attracted by the increased +charge but shielding also increases …which has a greater influence!! • therefore EA will decrease as you go down a group or family e- Nuclear charge increases Shielding increases Atomic radius increases Ionic size increases Ionization energy decreases Electronegativity decreases Summary Shielding is constant Atomic Radius decreases Ionization energy increases Electronegativity increases Nuclear charge increases