Hon Ch 6 Periodic Table 2020 PDF

Chapter 6 “The Periodic Table” Section 6.1 Organizing the Elements ◼ OBJECTIVES: ❑ Explain how elements are organized in a periodic table. Section 6.1 Organizing the Elements ◼ OBJECTIVES: ❑ Compare early and modern periodic tables. Section 6.1 Organizing the Eleme...

Chapter 6 “The Periodic Table” Section 6.1 Organizing the Elements ◼ OBJECTIVES: ❑ Explain how elements are organized in a periodic table. Section 6.1 Organizing the Elements ◼ OBJECTIVES: ❑ Compare early and modern periodic tables. Section 6.1 Organizing the Elements ◼ OBJECTIVES: ❑ Identify three broad classes of elements. Section 6.1 Organizing the Elements ◼ A few elements, such as gold and copper, have been known for thousands of years - since ancient times ◼ Yet, only about 13 had been identified by the year 1700. ◼ As more were discovered, chemists realized they needed a way to organize the elements. Section 6.1 Organizing the Elements ◼ Chemists used the properties of elements to sort them into groups. ◼ In 1829 J. W. Dobereiner arranged elements into triads – groups of three elements with similar properties ❑ One element in each triad had properties intermediate of the other two elements  Mendeleev’s Periodic Table ◼ By the mid-1800s, about 70 elements were known to exist ◼ Dmitri Mendeleev – a Russian chemist and teacher ◼ Arranged elements in order of increasing atomic mass ◼ Thus, the first “Periodic Table”  Mendeleev ◼ Heleft blanks for yet undiscovered elements ❑ When they were discovered, he had made good predictions ◼ But, there were problems: ❑ Such as Co and Ni; Ar and K; Te and I  A better arrangement ◼ In 1913, Henry Moseley – British physicist, arranged elements according to increasing atomic number ◼ The arrangement used today ◼ The symbol, atomic number & mass are basic items included  The Periodic Law says: ◼ When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties. ◼ Horizontal rows = periods ❑ There are 7 periods ◼ Vertical column = group (or family) ❑ Similar physical & chemical prop. ❑ Identified by number or number & letter (IA, IIA)  Areas of the periodic table ◼ Three classes of elements: ◼ 1) metals, 2) nonmetals, and ◼ 3) metalloids 1) Metals: electrical conductors, have luster, ductile, malleable 2) Nonmetals: generally brittle and non- lustrous, poor conductors of heat and electricity  Metals ◼ Conductors ◼ Lose electrons ◼ Malleable and ductile  Nonmetals ◼ Brittle ◼ Gain electrons ◼ Covalent bonds Semi-metals or Metalloids  Areas of the periodic table ◼ Some nonmetals are gases (O, N, Cl); some are brittle solids (S); one is a fuming dark red liquid (Br) ◼ Notice the heavy, stair-step line? 3) Metalloids: border the line-2 sides ❑ Properties are intermediate between metals and nonmetals Section 6.2 Classifying the Elements ◼ OBJECTIVES: ❑ Describe the information in a periodic table. Section 6.2 Classifying the Elements ◼ OBJECTIVES: ❑ Classify elements based on electron configuration. Section 6.2 Classifying the Elements ◼ OBJECTIVES: ❑ Distinguish representative elements and transition metals.  Squares in the Periodic Table ◼ The periodic table displays the symbols and names of the elements, along with information about the structure of their atoms: Atomic number and atomic mass Black symbol = solid; red = gas; blue = liquid (from the Periodic Table on our classroom wall)  Electron Configurations in Groups 1) Noble gases are the elements in Group 8A (also called Group18) Previously called “inert gases” because they rarely take part in a reaction; very stable = don’t react Noble gases have an electron configuration that has the outer s and p sublevels completely full Noble Gases  Electron Configurations in Groups 2) Representative Elements are in Groups 1A through 7A or 1,2,13-17. Display wide range of properties, thus a good “representative” Some are metals, or nonmetals, or metalloids; some are solid, others are gases or liquids Their outer s and p electron configurations are NOT completely filled  ◼ Elements in the 1A-7A groups are 1A called the representative elements 8A or main group elements 2A 3A 4A 5A 6A 7A outer s or p filling  Electron Configurations in Groups 3) Transition metals are in the “B” columns of the periodic table Groups 3-12 Electron configuration has the outer s sublevel full, and is now filling the “d” sublevel Examples are gold, copper, silver Transition metals  Electron Configurations in Groups 4) Inner Transition Metals are located below the main body of the table, in two horizontal rows Electron configuration has the outer s sublevel full, and is now filling the “f” sublevel Formerly called “rare-earth” elements, but this is not true because some are very abundant Inner Transition Metals The group B are called the transition elements These are called the inner transition elements, and they belong here Group 1A are the alkali metals (but NOT H) Group 2A are the alkaline earth metals H ◼ Group 8A are the noble gases ◼ Group 7A is called the halogens H 1s1 1 Do you notice any similarity in these Li 1s22s1 configurations of the alkali metals? 3 1s22s22p63s1 Na 11 1s22s22p63s23p64s1 K 19 1s22s22p63s23p64s23d104p65s1 Rb 37 1s22s22p63s23p64s23d104p65s24d10 Cs 5p66s1 55 Fr 1s22s22p63s23p64s23d104p65s24d105p66s 87 24f145d106p67s1 Do you notice any similarity in the 1s2 He 2 configurations of the noble gases? 1s22s22p6 Ne 10 1s22s22p63s23p6 Ar 18 1s22s22p63s23p64s23d104p6 Kr 36 1s22s22p63s23p64s23d104p65s24d105p6 Xe 1s22s22p63s23p64s23d104p65s24d10 54 5p66s24f145d106p6 Rn 86 Elements s1 in the s - blocks s2 He ◼ Alkali metals all end in s1 ◼ Alkaline earth metals all end in s2 ❑ really should include He, but it fits better in a different spot, since He has the properties of the noble gases, and has a full outer level of electrons. Transition Metals - d block Note the change in configuration. s1 s1 d 1 d 2 d 3 d 5 5 6 7 8 10 d d d d d d 10 The P-block p1 p2 p3 p4 p5 p6 F - block ◼ Called the “inner transition elements” f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14  1 2 3 Period 4 Number 5 6 7 ◼ Each row (or period) is the energy level for s and p orbitals. ◼ The “d” orbitals fill up in levels 1 less than the period number, so the first d is 3d even though it’s in row 4. 1 2 3d 3 4 5 4d 6 5d 7 1 2 3 4 5 6 7 4f 5f ◼ f orbitals start filling at 4f, and are 2 less than the period number  Details ◼ Valence electrons- the electrons in the outermost energy level. (not d or f). ◼ Core electrons- the inner electrons. Patterns in Electron Configuration ◼ Noble Gases-all have 8 valence electrons ◼ 1s2 2s2 2p6 ◼ 1s2 2s2 2p6 3s2 3p6 ◼ Alkali metals – all have 1 valence electron ◼ Alkaline-Earth Metals – 2 valence electrons ◼ Halogens – 7 valence electrons. Ions ◼ Atoms form ions to achieve noble gas configuration (become stable). ◼ Alkali Metals-lose 1 e¯ positive ion ◼ Alkaline-Earth- lose 2 e¯ positive ion ◼ Halogens –gain 1 e¯ -negative ion +1+2 -3 -2 -1  Section 6.3 Periodic Trends ◼ OBJECTIVES: ❑ Describe trends among the elements for atomic size.  Section 6.3 Periodic Trends ◼ OBJECTIVES: ❑ Explain how ions form.  Section 6.3 Periodic Trends ◼ OBJECTIVES: ❑ Describe periodic trends for first ionization energy, ionic size, and electronegativity. Trends in Atomic Size ◼ Very difficult to measure the radius of an atom (too small). ◼ Easier to compare the radii. ◼ Scientists measure more than 1 atom at a time.  Atomic Size } Radius ◼ Measurethe Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule. ALL Periodic Table Trends ◼ Influenced by three factors: 1. Energy Level ❑ Higher energy levels are farther away from the nucleus. 2. Charge on nucleus (# protons) ❑ More charge pulls electrons in closer. (+ and – attract each other) ◼ 3. Shielding effect (blocking effect?)  Shielding ◼ The electron on the outermost energy level has to look through all the other energy levels to see the nucleus. ◼ Second electron has same shielding, if it is in the same period  What do they influence? Energy levels and Shielding have an effect on a GROUP Nuclear charge has an effect on a PERIOD #1. Atomic Radius - Group trends ◼ As we increase the H atomic number (or go Li down a group)... ◼ each atom has Na another energy level, ◼ so the atoms get K bigger. Rb #1. Atomic Radius - Period Trends ◼ Going from left to right across a period, the size gets smaller. ◼ Electrons are in the same energy level. ◼ But, there is more nuclear charge. ◼ Outermost electrons are pulled closer. Na Mg Al Si P S Cl Ar  Rb K Period 2 Atomic Radius (pm) Na Li Kr Ar Ne H 3 10 Atomic Number  #2. Trends in Ionization Energy ◼ Ionization energy is the amount of energy required to completely remove an electron from a gaseous atom. ◼ Removing one electron makes a 1+ ion. ◼ The energy required to remove only the first electron is called the first ionization energy. Ionization Energy ◼ The second ionization energy is the energy required to remove the second electron. ❑ Always greater than first IE. ◼ The third IE is the energy required to remove a third electron. ❑ Greater than 1st or 2nd IE.  Table 6.1, p. 173 Symbol First Second Third H 1312 He 2731 5247 Li 520 7297 11810 Be 900 1757 14840 B 800 2430 3569 C 1086 2352 4619 N 1402 2857 4577 O 1314 3391 5301 F 1681 3375 6045 Ne 2080 3963 6276 Symbol First Second Third H 1312 Why did these values 2731 5247 increase so much? He Li 520 7297 11810 Be 900 1757 14840 B 800 2430 3569 C 1086 2352 4619 N 1402 2857 4577 O 1314 3391 5301 F 1681 3375 6045 Ne 2080 3963 6276 What factors determine IE ◼ The greater the nuclear charge, the greater IE. ◼ The greater the shielding (occupied energy levels) the lower IE Ionization Energy - Group trends ◼ As you go down a group, the first IE decreases because... ◼ There is more shielding (more occupied energy levels). ◼ The electron is farther away from the attraction of the nucleus. Ionization Energy - Period trends ◼ All the atoms in the same period have the same shielding (same number of occupied energy levels). ◼ But, increasing nuclear charge. ◼ So IE generally increases from left to right. ◼ A few exceptions.  He ◼ He has a greater IE First Ionization energy than H. ◼ Both elements have the same shielding H since electrons are only in the first level ◼ But He has a greater nuclear charge Atomic number First Ionization energy He Li has lower IE than H more shielding H farther away These outweigh Li the greater nuclear charge Atomic number First Ionization energy He Be has higher IE than Li same shielding H Be greater nuclear charge Li Atomic number  He B has lower IE First Ionization energy than Be same shielding H Be greater nuclear charge B Remove p Li electrons Atomic number  First Ionization energy H He Li Be B C Atomic number First Ionization energy He N H C Be B Li Atomic number First Ionization energy He ◼ Oxygen breaks N the pattern, because removing H C O an electron Be minimizes electron B Li repulsions Atomic number First Ionization energy He N F H C O Be B Li Atomic number  He Ne ◼ Ne has a lower First Ionization energy N F IE than He ◼ Both are full, H C O ◼ Ne has more Be shielding ◼ Greater distance B Li Atomic number  He Ne Na has a lower First Ionization energy N F IE than Li Both are s1 H C O Na has more Be shielding B Greater Li distance Na Atomic number  First Ionization energy Atomic number Electronegativity ◼ applies when an atom is in a compound NOT alone ◼ Electronegativity – measure of how strongly an atom attracts electrons when it is in a compound ◼ Fluorine (the most electronegative element) is assigned a 4.0 and then all the others were determined by comparison Electronegativity Period Trend ◼ Metals are at the left of the table. ◼ They let their electrons go easily ◼ Thus, low electronegativity ◼ At the right end are the nonmetals. ◼ They need more electrons. ◼ Try to take them away from others ◼ High electronegativity. The arrows indicate the trend: Ionization energy and Electronegativity INCREASE in these directions Atomic size and Ionic size increase in these directions: