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Ch. 5 “The Periodic Table” Why is the Periodic Table important to me? • The periodic table is the most useful tool to a chemist. • It organizes lots of information about all the known elements. • You get to use it on every test. Pre-Periodic Table Chemistry … • …was a mess!!! • No organization of elements. • Imagine going to a grocery store with no organization!! • Difficult to find what you need. • Chemistry didn’t make sense. (CHAOS) During the nineteenth century, chemists began to categorize the elements according to similarities in their physical and chemical properties. The end result of these studies was our modern periodic table. How did chemists begin to organize the known elements? As the number of elements increased, chemists inevitably began to find patterns in their properties. Chemists used the properties of elements to sort them into groups Ex. Chlorine, bromine, and iodine have very similar chemical properties. Model of Triads In 1817, Johann Dobereiner classified some elements into groups of three, which he called triads. The elements in a triad had similar chemical and physical properties. 1780 - 1849 Law of Octaves In 1865, John Newlands suggested that elements be arranged in “octaves” because he noticed when he arranged the elements in order of increasing atomic mass certain properties repeated every 8th element. 1838 - 1898 Lightest to heaviest. A Li B Be C B D C E N F O G F A Na B Mg C Al D Si E P S F G Cl A K B Ca C ? D ? E As F Se G Br Li Be B C N O F Na Mg Al Si P S Cl K Ca ? ? As Se Br He called this the “Law of Octaves” because of its similarity to musical octaves John Newlands Law of Octaves Newlands' claim to see a repeating pattern was met with savage ridicule on its announcement. His classification of the elements, he was told, was as arbitrary as putting them in alphabetical order and his paper was rejected for publication by the Chemical Society. The Modern Periodic Table In 1869 Dmitri Mendeleev published a table of the elements organized by increasing atomic mass. He was trying to organize elements so his students could learn them more easily! 1834 - 1907 A. Mendeleev and Chemical Periodicity • Mendeleev placed known information of elements on cards (atomic mass, density, etc…). He arranged them in order of increasing atomic masses, certain similarities in their chemical properties appeared at regular intervals. Such a repeating pattern is referred to as periodic. At the same time, Lothar Meyer published his own table of the elements organized by increasing atomic mass. 1830 - 1895 • Both Mendeleev and Meyer arranged the elements in order of increasing atomic mass. • Both left vacant spaces where unknown elements should fit. So why is Mendeleev called the “Father of the Periodic Table” and not Meyer, or both? Could it be his dashing good looks?! Mendeleev published first! • Mendeleev left blank spaces in his table when the properties of the elements above and below did not seem to match. The existence of unknown elements was predicted by Mendeleev on the basis of the blank spaces. When the unknown elements were discovered, it was found that Mendeleev had closely predicted the properties of the elements as well as their discovery. and the elusive element 32… Predicted Properties Observed Properties Atomic weight 72 72.61 Density 5.5 g/cm3 5.32 g/cm3 Melting point 825 C 938 C Oxide formula RO2 GeO2 Density of oxide 4.7 g/cm3 4.70 g/cm3 Dates (predicted and found) 1871 1886 Color Dark gray Gray-white Gallium Germanium The Father of the Periodic Table After the discovery of these unknown elements between 1874 and 1885, and the fact that Mendeleev’s predictions were amazingly close to the actual values, his table was generally accepted. However, in spite of Mendeleev’s great achievement, problems arose when new elements were discovered and more accurate atomic weights were determined. By looking at our modern periodic table, can you identify what problems might have caused chemists a headache? Ar and K Co and Ni Te and I Th and Pa Remember This…?! In 1913, through his work with X-rays, Henry Moseley determined the actual nuclear charge (atomic number) of the elements*. He rearranged the elements in order of increasing atomic number. *“There is in the atom a fundamental quantity which increases by regular steps as we pass from each element to the next. This quantity can only be the charge on the central positive nucleus.” 1887 - 1915 Henry Moseley His research was halted when the British government sent him to serve as a foot soldier in WWI. He was killed in the fighting in Gallipoli by a sniper’s bullet, at the age of 28. Because of this loss, the British government later restricted its scientists to noncombatant duties during WWII. Periodic Law When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties. The Current Periodic Table • Mendeleev wasn’t too far off. • Now the elements are put in rows by increasing ATOMIC NUMBER!! • The vertical columns are called groups or families and are labeled from 1 to 18 (modern) • or in A & B Groups (with Roman numerals) Groups…Here’s Where the Periodic Table Gets Useful!! • Elements in the same group have similar chemical and physical properties!! • (Mendeleev did that on purpose.) Why?? • They have the same number of valence electrons. • They will form the same kinds of ions. Groups in the Periodic Table Elements in groups react in similar ways! Periodsrows in the Periodic Table The horizontal are called periods and are labeled from 1 to 7. All elements in a period have the same number of energy levels (= to period #) 11 1 2 3 4 5 6 7 Energy Levels n=1 n=2 n=3 n=4 In addition to Group Labels, many of the groups have Family Names Group 1A: Alkali Metals lithium Cutting sodium metal potassium Alkali Metals • They are the most reactive metals. • They react violently with water. • Alkali metals are never found as free elements in nature - they are always in compounds with other elements. • Only 1 valence electron • Soft metals • Must be stored under mineral oil, etc. Group 2A: Alkaline Earth Metals Magnesium Group 2A: Alkaline Earth Metals Only 2 valence electrons Too reactive to be uncombined in nature. calcium strontium barium Group 7A: The Halogens 7 valence electrons All non-metals Very reactive All physical states represented Colored gases (always poisonous!) Occur as diatomic molecules when pure fluorine F2 Cl2 chlorine Br2 bromine Iodine I2 The Noble Gases Noble Gases • Noble Gases are colorless gases that are extremely un-reactive.(inert) • They are inactive because their outermost energy level is full. (8 valence electrons – except He which has 2) • Having 8 valence electrons is low in energy and, therefore, very stable. Hydrogen • The hydrogen square sits atop Family IA, but it is not a member of that family. Hydrogen is in a class of its own. (An orphan?) • Like the Alkali metals, it only needs to lose one electron to be stable. (but it is not a metal!) • Sometimes it’s shown above 7A. • Like the Halogens, it only needs to gain one electron to have the stable Noble Gas electron configuration. (but it is not a Halogen!) Hey Cameron, why are those elements by themselves on the bottom of the Periodic Table?! I’ll handle this one, Cam! If they weren’t put on the bottom, the Periodic Table wouldn’t fit very nicely on a page! In fact, the table would look like this. In fact, we have Glen Seaborg to thank for the fact that my Periodic Table doesn’t stick out of my notebook in a truly tasteless manner! Glenn T. Seaborg After co-discovering 10 new elements, in 1944 he moved 14 elements out of the main body of the periodic table to their current location below the Lanthanide series. These became known as the Actinide series. 1912 - 1999 “I was warned at the time that it was professional suicide to promote this idea, which has since been called one of the most significant changes in the periodic table since Mendeleev’s 19th century design. Luckily, I stuck to my guns and have seen the actinide concept become the foundation for many significant discoveries in heavy element research.” Seaborgium Glenn T. Seaborg He is the only person to have an element named after him while still alive. "This is the greatest honor ever bestowed upon me - even better, I think, than winning the Nobel Prize." 106 Sg 1912 - 1999 Seaborgium 271 There are many ways that we can break the Periodic Table up into sections! Metals Metals Metals Metals are good conductors of heat. That's why a branding iron is made from metal. The heat transfers quickly to the animal's hide. Metals also conduct electricity. Notice that the Tesla coil sparks seek out metallic objects because they conduct electricity better than the nonmetallic materials such as wood or soil. Metals are also malleable and can be bent or hammered into various shapes. Metals • • • • What comes to mind? Most elements are metals Loosely held valence e-’s Properties of metals: 1. Good conductors of heat and electricity (p) 2. High density (p) 3. High melting points (p) 4. Luster (p) 5. Malleable (p) 6. Ductile (p) 7. 1, 2, or 3 valence electrons Nonmetals Nonmetals do not conduct heat well. The insulating tiles from the Space Shuttle are made from fibers of silicon and oxygen (silica=sand). Nonmetals • • Opposite of metals Properties of nonmetals: 1. Dull (no luster) 2. Do not conduct heat/elec. 3. Not ductile 4. Not malleable 5. All phases 6. Have 5, 6, or 7 valence electrons • Form many compounds with metals Metalloids (Semi-Metals) • Means “metal-like” • Dividing line between metals and nonmetals • Al is the exception • Properties of both metals and nonmetals Four Main Categories of Elements • Noble Gases- group 18 or 0 or 8A – s & p sublevels filled – 8 valence __s2…__p6 – Inert- not reactive- because of full outer shell of electrons • Representative Elements also called main group elements- Groups 1A-7A – s & p partly filled – Includes alkali metals, alkaline earth metals, and halogens 1A The 2A elements in the A groups 8A 0 are called the representative 3A 4A 5A 6A 7A elements Four Main Categories of Elements • Transition Metals – – Unfilled inner shells – outermost s & inner d sublevels contain electrons – Hard & brittle • Inner transition metals– outermost s & nearby f sublevel contain electrons – Lanthanides (4f) and actinides (5f) Representative Representative Noble Gases Inner Transition Elements Using the Diagonal Rule is just so bothersome! I wish there was an easier way to figure out electron configurations! Oh, but there is! Watch this! 1 2 H 1 1s1 Li 1s22s1 3 22s22p63s1 1s Na 3 11 K 4 19 1s22s22p63s23p64s1 Rb 22s22p63s23p63d104s24p65s1 5 1s 37 Cs 1s22s22p63s23p63d104s24p64d105s2 5p66s1 6 55 22s22p63s23p63d104s24p64d104f145s25p65d1 1s Fr 06s26p67s1 7 87 Group 2A Be 1s22s2 Mg 1s22s22p63s2 Ca 1s22s22p63s23p64s2 Sr 1s22s22p63s23p64s23d104p65s2 s-block 1 H 2 He 2 3 Li 4 Be 3 11 Na 12 Mg Always for row you are on! 4 19 K 20 Ca 5 37 Rb 38 Sr 6 55 Cs 56 Ba 7 87 Fr 88 Ra Group 3A B Al Ga 1s22s22p1 1s22s22p63s23p1 1s22s22p63s23p64s23d104p1 Always for row you are on! The P-block p 1 p2 p3 p4 p5 p6 6.2 Electron Configurations in Groups • In atoms of the Group 1A elements below, there is only one electron in the highest occupied energy level. It’s always s1 for the row it’s on! • In atoms of the Group 4A elements below, there are four electrons in the highest occupied energy level. Always s2p2 for the row they’re on! 6.2 Electron Configurations in Groups – The Noble Gases • In atoms of the Group 8A elements below, there are eight electrons in the highest occupied energy level. Except for He, always s2p6 for the row they are on! Chemical elements in d-block Always 1 level in (back) from the row you’re on! Group → 3 4 5 6 7 8 9 10 11 12 4 21 Sc 22 Ti 23 V 24 Cr 25 Mn 26 Fe 27 Co 28 Ni 29 Cu 30 Zn 3d 5 39 Y 40 Zr 41 Nb 42 Mo 43 Tc 44 Ru 45 Rh 46 Pd 47 Ag 48 Cd 4d 6 71 Lu 72 Hf 73 Ta 74 W 75 Re 76 Os 77 Ir 78 Pt 79 Au 80 Hg 5d 7 103 104 105 106 107 108 109 110 111 112 Lr Rf Db Sg Bh Hs Mt Ds Rg Uub 6d ↓ Period F - block inner transition elements Always go back 2 energy levels from the row you’re on! f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 4f 5f Blocks of Elements s-block d-block p-block f-block Periodic Table e- configuration from the periodic periodic table (To be covered in future chapters) 1 IA 18 VIIIA 2 IIA 1 H 1s1 2 Li Be 2s1 2s2 Na Mg 3s1 3s2 3 4 5 6 7 13 IIIA 3 IIIB 4 IVB Sc 3d1 Rb 5s1 Ca 4s2 Sr 5s2 Y 4d1 V Ti Cr Mn Fe Co 3d2 3d3 4s13d5 3d5 3d6 3d7 Zr Nb Mo Tc Ru Rh 4d2 4d3 5s14d5 4d5 4d6 4d7 Cs 6s1 Ba 6s2 La 5d1 Hf Ta W Re Os 5d2 5d3 6s15d5 5d5 5d6 Fr 7s1 Ra 7s2 Ac Rf 6d1 6d2 K 4s1 5 VB 6 VIB 7 VIIB Db Sg Bh 6d3 7s16d5 6d5 8 9 VIIIB 14 IVA 15 VA 16 VIA 17 VIIA B 2p1 •B C N O 1 2 3 •2p 2p 2p 2p4 F 2p5 Ne 2p6 Al 3p1 Si 3p2 Cl 3p5 Ar 3p6 He 1s2 10 11 IB 12 IIB Ni 3d8 Cu 4s13d10 Ni 4d8 5s14d10 Zn Ga Ge 3d10 4p1 4p2 Cd In Sn 10 4d 5p1 5p2 As Se Be 4p3 4p4 4p5 I Sb Te 5p3 5p4 5p5 Kr 4p6 Xe 5p6 Hg Tl Pb 5d10 6p1 6p2 Bi Po At 6p3 6p4 6p5 Rn 6p6 Ir Ni 7 5d 5d8 Hs Mt 6d6 6d7 Ag Au 6s15d10 S P 3 3p 3p4 Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of properties when the elements are in order of increasing atomic number. Explaining Periodic Trends Why a property is higher/lower, bigger/smaller, etc.! 1. Nuclear charge- the number of protons in the nucleus. More protons = increased nuclear charge so increased attraction between the nucleus and electrons. Think of the nucleus as a magnet – each extra proton makes the magnet more powerful at attracting electrons & holding them tight! 2. Shielding- lessens the attractive force of the nucleus for the valence electrons – caused by electrons in energy levels between the nucleus and the valence electrons Shielding increases as you go down a group because there are more energy levels (more core electrons). Shielding stays the same as you move across a period because the number of energy levels is staying the same. Which atom has more shielding? (A) K or Ca (B) Na or K Which atom is smaller? (A) N or P (B) Li or K Metallic Character Atomic Radius • Atomic Radius- half the distance between the nuclei of two atoms of the same element in a diatomic molecule Atomic Radius • Trend for atomic size– Down a group, size increases • Occurs because # of energy levels increases *Makes a BIG difference in size!! • shielding also increases. – Across a period, size decreases • # of protons increases (nuclear charge increases), pulling electrons closer shielding doesn’t change because electrons are added to the same energy Atomic Radius Ionization Energy • Ionization Energy- energy needed to remove an electron from an atom. • Outer shell electrons are easier to remove than ‘core’ electrons so it takes less energy to remove them! Highest toward upper right corner of PT since these atoms are smaller & their valence electrons are closer to the nucleus -so held more tightly Trends in Ionization Energy Periodic Trends • Ionization energy – Down a group- decreases – because electrons are held more loosely due to increased # of energy levels & increased shielding – Across a period- increases because electrons are held more tightly due to increased nuclear charge (increased # of protons in the nucleus) Ionization Energy Ionization Energy • There are big jumps in ionization energy whenever you try to remove an electron from an inner energy level! Ionization Energy Electron Affinity • Electron affinity of an element is the energy given off when an atom (in the gas phase) gains an electron to form an ion – Example: F(g) + e- F-(g) – Ho (ENERGY) = -328.0 kJ/mol Trends in Electron Affinity • It decreases down a group, because electron shielding blocks some of the attraction from the nucleus • It increases across a period, because nuclear charge increases, attracting electrons more strongly. Periodic Trends • Electronegativity- tendency for the atoms of the element to attract electrons when the atoms are part of a compound • Noble gases- no electronegativity valuesdon’t form compounds • In general, metals have low EN and nonmetals have high EN. • The actual amount of EN an atom has is indicated by a number on the Pauling Electronegativity Scale that goes from 0 to 4. • Dr. Linus Pauling set up this scale and gave the element having the greatest EN an arbitrary number of 4, and he assigned numbers to the others relative to this element. • Flourine is the most electronegative element at 4. (3.98) and Francium is the least electronegative at 0.7. Periodic Trends • Electronegativity Trends– Down a group – decreases- electron shielding results in less attraction for electrons by the nucleus – Across a period- increases- higher atomic number and consistent electron shielding result in more attraction for electrons • Electronegativity allows you to predict bond type: covalent- polar vs. nonpolar and ionic General Trends in the Periodic Table: Atomic & Ionic Radii Ionization Energy (IE) Electron Affinity (EA) Electronegativity (c) IE, EA, and c are useful concepts used to characterize different types of bonding and estimate bond energies. 6.3 Summary of Trends Increases Decreases Ionization Ionicof size energy Size Size Electronegativity Atomic Nuclear Shielding of anions cations Size Charge Increases DecreasesConstant IONS • Remember – Atoms are neutral • But…atoms can gain or lose electrons (*# of protons NEVER changes during reactions!) • IONS are atoms or groups of atoms with a charge. • To tell the difference between an atom and an ion, look to see if there is a charge in the superscript! • Examples: Na Ca I O • Na+ Ca+2 I- O-2 •When an atom loses an electron it gets a positive charge (because it now has more protons than electrons) Mg --> Mg+2 + 2 eWhen an atom gains an electron it forms a negative ion (because it now has more electrons than protons) F + e- --> F- Atom versus Ion Forming Cations & Anions A CATION forms when an atom loses one or more electrons. Mg --> Mg2+ + 2 eNow has 12 protons & 10 electrons An ANION forms when an atom gains one or more electrons F + e- --> FNow has 9 protons & 10 electrons – Metals have 1, 2, or 3 valence electrons so tend to lose electrons (to get an octet)- forming cations. (+ charge) – Non-Metals have 5, 6, or 7 valence electrons so tend to gain electrons (to get an octet) - forming anions. (charge) Periodic Trends • Ionic Radii Trends – Cations- smaller than neutral atom because fewer electrons result in greater attraction by nuclei – Anions- larger than neutral atom because more electrons result in less attraction by nuclei – Within period- size decreases – Down a group – size increases Forming cations Forming anions EXAMPLE What would the charge be on a sodium ion? Since sodium in in Group IA it has 1 valence eand so it would LOSE an electron It goes from 11 protons & 11 electrons to 11 protons & 10 electrons So it gets a charge of +1 Remember an electron is negatively charged. When an atom loses electrons it forms positively charged ions. When electrons are gained negatively charged ions form EXAMPLE How would you write the symbol for the sodium CATION? Na +1 How many outer electrons does sodium have before it loses one? It has 1…remember the group number! 5 6.3 Section Quiz – 1. Which of the following sequences is correct for atomic size? • • • • Mg > Al > S Li > Na > K F>N>B F > Cl > Br 6.3 Section Quiz – 2. • • • • Metals tend to gain electrons to form cations. gain electrons to form anions. lose electrons to form anions. lose electrons to form cations. 6.3 Section Quiz – 3. Which of the following is the most electronegative? • • • • Cl Se Na I The Periodic Table Summary of Trend • Periodic Table and Periodic Trends • 1. Electron Configuration 3. Ionization Energy: Largest toward NE of PT 4. Electron Affinity: Most favorable NE of PT 2. Atomic Radius: Largest toward SW corner of PT Periodic Table: electron behavior West (South) Mid-plains East (North) METALS Alkali Alkaline Transition These elements tend to give up e - and form CATIONS METALLOID These elements will give up e- or accept e- NON-METALS Noble gas Halogens Calcogens These elements tend to accept e - and form ANIONS • The periodic table can be classified by the behavior of their electrons 1 IA 1 2 IIA 13 IIIA 2 3 4 5 6 7 3 IIIB 4 IVB 5 VB 6 VIB 7 VIIB 8 9 VIIIB 10 11 IB 12 IIB 14 IVA 15 VA 16 VIA 17 VIIA 18 VIIIA ELEMENTS THAT EXIST AS DIATOMIC MOLECULES Remember: BrINClHOF These elements only exist as PAIRS. Note that when they combine to make compounds, they are no longer elements so they are no longer in pairs! Valence electrons 1 2 3 4 5 6 7 8 Select an element ( = Internet link ) Configuration of Ions Ions of representative elements have noble gas configuration Na is 1s22s22p63s1 Forms a 1+ ion - 1s22s22p6 Same configuration as neon Metals form ions with the configuration of the noble gas before them - they lose electrons This can explain why metals are shiny. This is the surface of copper at a ridiculously high magnification. The surface shows a lake of electrons along with ripples. The two islands are imperfections on the surface. Most likely a couple of atoms that aren't copper.