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Chapter 6 - The Periodic Table and Periodic Law Objectives: Identify different key features of the periodic table. Explain why elements in a group have similar properties. Relate the group and period trends seen in the periodic table to the electron configuration of atoms Why this is important: The periodic table is one of the most useful reference tools available in chemistry. Understanding its organization and interpreting its data will aid in understanding chemistry concepts. www.privatehand.com/flash/elements.html Slide 2 Development of the Periodic Table In 2003, there were 118 elements known. The majority of the elements were discovered between 1735 and 1843. How do we organize all the different elements in a meaningful way that will allow us to make predictions about undiscovered elements? Slide 3 A variety of scientists tried to arrange the known elements to reflect the trends in chemical and physical properties but their systems did not allow for newly discovered elements to fit in their charts. 1869 Dmitri Mendeleev and Lothar Meyer separately arranged the elements in order of increasing atomic mass and into columns with similar properties That seemed to work to organize most of the elements. http://www.chemistrydaily.com/chemistry/upload/a/a1/Dmendeleev.jpg http://www.chemistryexplained.com/images/chfa_03_img0535.jpg Slide 4 Mendeleev is given more credit than Meyer because he published his findings first and he left spaces for elements that were not yet discovered. Some of the elements that he predicted were scandium, gallium, and germanium. In 1871, Mendeleev noted that arsenic (As) properly belonged underneath phosphorus (P) and not silicon (Si), which left a missing element underneath Si. He predicted a number of properties for this element. In 1886 Germanium (Ge) was discovered. The properties of Ge matched Mendeleevs predictions. Slide 5 Mendeleevs table was not completely correct. Arranging elements by atomic mass caused some elements to be put in the wrong groups so that the properties did not exactly match up 1913 English chemist Henry Moseley arranged elements in order of increasing atomic number Problems with order of elements were solved and there was a clear repeating pattern of properties of the elements in their groups. The PERIODIC LAW states there is a periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number. Periodic means: happening or reoccurring at regular intervals (definition from Websters Dictionary) http://www.rsc.org/education/teachers/learnnet/periodictable/scientists /moseley.jpg Slide 6 The Modern Table Boxes are arranged in order of increasing atomic # Elements are grouped into columns by similar properties Scientists keep adding elements that were discovered The final adjustment was when physicist Glenn Seaborg had the inner-transition elements pulled below the rest of the periodic table and into 2 separate rows (This occurred in the late 1940s) http://www.lbl.gov/Science-Articles/Research- Review/Magazine/1994/seaborgium-mag.html Slide 7 Horizontal rows are called periods There are 7 periods Slide 8 Vertical columns are called groups or families. Elements are placed in columns by similar properties. b/c of the similar numbers of valence e - they contain Slide 9 12 1314151617 18 3456789101112 The Different Groups of Elements 1 18 system used by all chemists A & B system is an older American System A elements are representative elements B elements are transition elements 1A2A 3A4A5A6A7A 8A 3B4B5B6B7B8B 1B2B Slide 10 Representative or Main Group Elements They have a wide range of physical & chemical properties. They have the whole range of possible valence electrons (1 to 8) Also called s and p block elements Here are some important groups: Slide 11 Group 1 (1A) contains the alkali metals (remember to NOT include hydrogen) Group 2 (2A) contains the alkaline earth metals Slide 12 Alkali Metals Called this because these metals react w/ water to form alkaline (basic) solutions. Highly reactive metals that lose their 1 valence electron to form 1+ ions Soft enough to be cut with a knife. They are stored in oil to prevent reactions with oxygen and water in the air. Slide 13 Alkaline Earth Metals Called this because most of these metals react with oxygen to form compounds called oxides (the alchemists called them earths because of this) and the oxides react w/ water to form alkaline (basic) solutions Not as reactive (but do react easily) & harder than group 1 metals They lose their 2 valence electrons to form 2+ ions Slide 14 Groups 13 - 16 Named for the first element in each group. They do have mixed groupings of elements because each column contains nonmetals, metalloids, and metals. Many of the elements in these groups form various charges Slide 15 Group 17 (7A) contains the halogens Group 18 (8A) contains the noble gases Slide 16 Halogens Called this because halogen means salt formers b/c they react with metals to form salts (ionic compounds) Physically F & Cl are gases at room temp., Br is a liquid but it evaporates easily, and Iodine is a solid that sublimes easily Astatine is the odd-ball of the group b/c its radioactive w/ no known uses Chemically they are the most reactive nonmetals They have 7 valence e- so they will share or gain 1 e- and they tend to form 1- ions. Slide 17 Noble Gases Last naturally occurring elements to be discovered b/c they are colorless & unreactive Very stable with full valence electrons = 8 (except He w/ 2) With lots of energy you can get Xe, Kr and Ar compounds (There are no known He or Ne compounds) In 1962 the first compound of the noble gases was prepared: XeF 2, XeF 4, and XeF 6. To date the only other noble gas compounds known are KrF 2 and HArF. Slide 18 Transition elements (metals) d-block f-block Slide 19 u These are called the inner transition elements and they belong here Slide 20 Slide 21 Transition Metals Make up the majority of elements on the periodic table Have a wide variety of uses & effect the economy The variation of their physical properties is b/c of their electron configurations & b/c unpaired d-electrons can move into valence shells. The more unpaired d-electrons, the greater the hardness & higher the melting & boiling points Most lose electrons to become positively charged ions Cu, Ag, Au, Pt, and Pd are the only ones unreactive enough to be found alone in nature Slide 22 Inner Transition Metals Lanthanide series follow element lanthanium All silvery metals w/ high melting points Actinide series follow element Actinium All are radioactive & only 3 exist in nature Slide 23 HYDROGEN In a class by itself, it is a unique element Most often occurs as a colorless diatomic gas, H 2 It is placed in Group 1 b/c it has one valence e- and will easily lose its 1 electron when reacting w/ other nonmetals to become a 1+ ion (H + is a proton) But it shares many properties w/ the halogens and will sometimes gain e- when bonding w/ a metal to become a 1- ion (the hydride ion, H - ) It is the most abundant element in the universe (90% by mass) Slide 24 Metals, Nonmetals, and Metalloids Slide 25 Metals What are some common properties of metals? Have luster (shine) when smooth & clean. Good conductors of heat & electricity Most are solid @ room temp. Most are: Ductile = drawn into wires. Malleable = hammered into sheets. Most lose electrons to become cations Most of the elements on the periodic table are classified as metals Slide 26 Nonmetals What are some common properties of nonmetals? At room temp: Some are brittle & dull solids Gases Poor conductors = good insulators Most tend to gain e- to become anions They have a wide variety of melting & boiling points Slide 27 Metalloids or Semimetals Chemical & Physical Properties of both metals & non- metals Example: Si has a metallic luster but it is brittle. They are semiconductors b/c they do conduct electricity but not as well as metals Silicon (Si) & Germainum (Ge) are 2 most important for computer chips & solar panels Slide 28 Valence Electrons REVIEW These are defined as electrons in the atoms highest energy levels. Examples: Sodium Has 11 electrons but only one valence electron Cesium (Cs) Has 55 electrons but only one valence electron Bromine Has _____ electrons but only _____ valence electron Valence electrons help determine the chemical properties of an element and how it will bond to form compounds. Slide 29 Examples Look at the electron configuration of each element. What do you notice? Na[Ne]3s 1 K[Ar]4s 1 Cs[Xe]6s 1 F1s 2 2s 2 2p 5 Cl[Ne]3s 2 3p 5 Br[Ar]4s 2 3d 10 4p 5 Ne1s 2 2s 2 2p 6 Ar[Ne] 3s 2 3p 6 Kr[Ar]4s 2 3d 10 4p 6 Most elements in the same group have the same ending electron configuration = same number of valence electrons Slide 30 Valence electrons and groups: Group 1 elements have 1 valence electron Group 2 elements have 2 valence electrons Group 13 elements have 3 valence electrons Group 14 elements have 4 valence electrons Group 15 elements have 5 valence electrons Group 16 elements have 6 valence electrons Group 17 elements have 7 valence electrons Group 18 elements have 8 valence electrons (except He, it only has 2) Most of the transition metals have 2 valence electrons, there are a lot of exceptions!! Slide 31 Depicting Valence Electrons Review Electron dot structures used to visually represent valence electrons in a shorthand method. We use an elements symbol to show what element we are talking about and the dots represent the atoms valence electrons. In writing these structures the dots are placed one at a time on the four sides of the symbol and then paired up until all are used. Slide 32 Examples of electron dot structures: Magnesium Sulfur Rubidium (Rb) Bromine Oxygen Slide 33 Section 6.3 - Periodic Trends Objectives: Compare period & group trends for shielding, atomic radius, ionic radius, ionization energy, & electronegativity Slide 34 Shielding (or screening) What does this mean? The valence e- are blocked from the full positive charge of the nucleus (effective nuclear charge) by the inner (core) e- As the average number of core e- increases, the effective nuclear charge decreases This idea of shielding will play a large role in a lot of the trends Slide 35 Mg Slide 36 Trend within the period (left to right): Generally decreases Why: B/c the number of energy levels & core e- stays the same but the nucleus is increasing This increase the attraction between the nucleus and valence e- Trend down a group: Generally increases Why: B/c the number of energy levels & core e- increases This makes the valence e- farther from the nucleus and more blocked by the inner e- Shielding Slide 37 Atomic Radius the distance between adjacent nuclei of identical atoms either in crystal form (metals) or in molecular form (nonmetals) Trend within the period (left to right): Generally decreases Why: B/c the number of energy levels & core e- stays the same but the nucleus is increasing This increase the attraction between the nucleus and valence e- This attraction pulls the e- closer to the nucleus and makes the atom smaller Ex: Na vs. S Slide 38 Trend down a group: Increases Why: B/c the number of energy levels increases & core e- increases Each energy level is larger than the next This makes the valence e- farther from the nucleus and more blocked by the inner e- - EX: Na vs. K Slide 39 Slide 40 Examples Place each group of elements in order of increasing atomic radius: 1.S, Al, Cl, Mg, Ar, Na 2.K, Li, Cs, Na, H 3.Ca, As, F, Rb, O, K, S, Ga Slide 41 Examples Place each group of elements in order of increasing atomic radius: 1.S, Al, Cl, Mg, Ar, Na Ar < Cl < S < Al < Mg < Na 2.K, Li, Cs, Na, H H < Li < Na < K < Cs 3.Ca, F, As, Rb, O, K, S, Ga F < O < S < As < Ga < Ca < K < Rb Slide 42 Ionic Radius The distance between the nucleus and the outermost electron in ions (cant be determined directly) Trend between atom & ion and Why: Cations are smaller than original atom b/c losing e- the atom has unequal positive charge that attracts the valence e- closer to the nucleus Anions are larger than original atom and cations b/c adding negative e- adds to the repulsion between other valence e-, pushing them apart Slide 43 Ionic Radius Continued Trend within the period (left to right): Representative Elements Cations the size decreases Anions the size drastically increases compared to the positive ions and then decreases across the period Trend down a group: Increases for both cations & anions Why: Same reason as atomic radii trend Slide 44 Slide 45 Electron Configurations of Ions Cations: electrons removed from orbital with highest principle quantum number, n, first: Li 1s 2 2s 1 Li + 1s 2 Fe [Ar]3d 6 4s 2 Fe 3+ [Ar]3d 5 Anions: electrons added to the orbital with highest n: F 1s 2 2s 2 2p 5 F 1s 2 2s 2 2p 6 Slide 46 Write electron configurations for the following ions: 1.Al 3+ 2.S 2- 3.Li + 4.Br - 5.Fe 2+ 6.Fe 3+ Slide 47 Write electron configurations for the following ions: 1.Al 3+ 1s 2 2s 2 2p 6 2.S 2- [Ne]3s 2 3p 6 3.Li + 1s 2 4.Br - [Ar]4s 2 3d 10 4p 6 5.Fe 2+ [Ar]3d 6 6.Fe 3+ [Ar]3d 5 Slide 48 For ions of the same charge, ion size increases down a group. All the members of an isoelectronic series have the same number of electrons. As nuclear charge increases in an isoelectronic series the ions become smaller: O 2- > F - > Na + > Mg 2+ > Al 3+ Slide 49 Examples Choose the larger species in each case: 1.Na or Na + 2.Br or Br - 3.N or N 3- 4.O - or O 2- 5.Mg 2+ or Sr 2+ 6.Mg 2+ or O 2- 7.Fe 2+ or Fe 3+ Slide 50 Examples Choose the larger species in each case: 1.Na or Na + 2.Br or Br - 3.N or N 3- 4.O - or O 2- 5.Mg 2+ or Sr 2+ 6.Mg 2+ or O 2- 7.Fe 2+ or Fe 3+ Slide 51 Ionization Energy: The energy required to remove an electron from a gaseous atom (also called First Ionization Energy, I 1 ) Na (g) + 496 kJ Na + (g) + e - The second ionization energy, I 2, is the energy required to remove an electron from a (1+) gaseous ion: Na + (g) + 4562 kJ Na 2+ (g) + e - NOTICE: Ionization Energy increases for each electron removed from the same element The larger ionization energy, the more difficult it is to remove the electron. Slide 52 Variations in Successive Ionization Energies There is a sharp increase in ionization energy when a core electron is removed. Notice the large increase after the last valence electron is removed. This chart can be used to determine the number of valence electrons in an atom of an element. Slide 53 Trends within the periods: Increases Why: Electrons are harder to remove from smaller atoms because they are closer to the nucleus and there is an increased nuclear charge Trends down a group: Decreases Why: Electrons are easier to remove from large atoms because they are farther away from the nucleus so there is less energy needed to remove them. Notice the trends in ionization energy is inversely related to trends in atomic radii. Slide 54 Slide 55 Slide 56 Examples Put each set in order of increasing first ionization energy: 1.P, Cl, Al, Na, S, Mg 2.Ca, Be, Ba, Mg, Sr 3.Ca, F, As, Rb, O, K, S, Ga Slide 57 Examples Put each set in order of increasing first ionization energy: 1.P, Cl, Al, Na, S, Mg 2.Ca, Be, Ba, Mg, Sr 3.Ca, F, As, Rb, O, K, S, Ga 1. Na < Al < Mg < S < P < Cl 2.Ba < Sr < Ca < Mg < Be 3.Rb < K < Ca < Ga < As < S < O < F Slide 58 ELECTRONEGATIVITY Electronegativity: The relative ability of atoms that are bonded to attract electrons in the chemical bond to itself Chemist Linus Pauling set electronegativities on a scale from 0.7 (Cs) to 4.0 (F). Values are calculated from ionization energies and electron affinities. They are used to help determine types of bonding (ionic or covalent) that are occurring in a compound. Noble gases are not usually given electronegativity values Slide 59 Trends within the periods: Increases Why: Atoms become smaller so the shared electrons are closer to the nucleus in small atoms Trends down the groups: Decreases Why: Atoms become larger so shared electrons are farther from the nucleus in large atoms Slide 60 Electronegativity Slide 61 Examples put each set in order by increasing electronegativity: 1.Na, Li, Rb, K, Fr 2.Cl, Ca, F, P, Mg, S, K Slide 62 Examples put each set in order by increasing electronegativity: 1.Na, Li, Rb, K, Fr 2.Cl, Ca, F, P, Mg, S, K 1.Fr < Rb < K < Na < Li 2.K < Ca < Mg < P < S < Cl < F Slide 63 Review: 1.As you move across a period, left to right, describe what generally happens (decreases, increases, or remains the same) to: a.The number of valence electrons b.The ionization energy c.The atomic radius 2. Give a brief explanation for your answers to a-c. Slide 64 3. Identify the element from the clues given: a.This element has a smaller atomic radius than phosphorous, it has a smaller ionization energy than fluorine and is chemically similar to iodine b.This element has the smallest ionization energy of any element in Period 4.