Chapter 6 - The Periodic Table and Periodic Law
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1 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.
2 Development of the Periodic Table In 2003, there were 118 elements known. The majority of the elements were discovered between 1735 and How do we organize all the different elements in a meaningful way that will allow us to make predictions about undiscovered elements?
3 1869 Dmitri Mendeleev and Lothar Meyer separately arranged the elements in order of increasing atomic mass and into columns with similar properties. 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 Mendeleev s predictions.
4 Mendeleev s 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 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 Webster s Dictionary) /moseley.jpg
5 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) Review/Magazine/1994/seaborgium-mag.html
6 Horizontal rows are called periods There are 7 periods
7 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!
8 1 2 1A 2A 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 A 4A 5A 6A 7A 18 8A B 4B 5B 6B 7B 8B 8B 8B 1B 2B
9 Representative or Main Group Elements Wide range of physical & chemical properties. The whole range of possible valence electrons (1 to 8) Also called s and p block elements Here are some important groups:
10 Group 1 (1A) contains the alkali metals (remember to NOT include hydrogen) Group 2 (2A) contains the alkaline earth metals
11 Alkali Metals 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.
12 Alkaline Earth Metals 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.
13 Groups Named for the first element in each group. Have mixed groupings of elements because each column contains nonmetals, metalloids, and metals. Many of the elements in these groups form various charges.
14 Group 17 (7A) contains the halogens Group 18 (8A) contains the noble gases
15 Halogens Halogen means salt formers b/c they react with metals to form salts (ionic compounds). 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 radioactive w/ no known uses. They are the most reactive nonmetals! 7 valence e- share or gain 1 e- and they tend to form 1- ions.
16 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. (No known He or Ne compounds) In 1962 the first compound of the noble gases was prepared: XeF 2, XeF 4, and XeF 6. (flexible e- arrangments b/c f orbitals) To date the only other noble gas compounds known are KrF 2 and HArF.
17 Transition elements (metals) d-block f-block
18 These are called the inner transition elements and they belong here
19 Transition Metals Majority of elements on the periodic table. Wide variety of uses & affect the economy! Variation of physical properties 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. Lose electrons to become positively charged ions. Cu, Ag, Au, Pt, and Pd are the only ones found alone in nature.
20 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.
21 HYDROGEN In a class by itself, it is a unique element! Most often occurs as a colorless diatomic gas, H 2. 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 - ). Most abundant element in the universe (90% by mass).
22 Metals, Nonmetals, and Metalloids
23 Metals What are some common properties of metals? Have luster (shine) when smooth & clean. Good conductors of heat & electricity. Most are 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.
24 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.
25 Metalloids (or Semimetals) Chemical & Physical Properties of both metals & nonmetals. 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!
26 Section Periodic Trends Objective: - Compare period & group trends for shielding, atomic radius, ionic radius, ionization energy, & electronegativity. Shielding (or screening): 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. Concept of shielding will play a large role in a lot of the trends.
27 Trend within the period (left to right): Generally decreases Why: Shielding: Number of energy levels & core e- stays the same, but nucleus is increasing. Increased attraction between nucleus and valence e-. Trend down a group: Generally increases Why: Number of energy levels & core e- increases. Valence e- farther from nucleus & more blocked by inner e-.
28 Atomic Radius: ½ the distance between adjacent nuclei of identical atoms. Trend within the period (left to right): Generally decreases Why: Number of energy levels & core e- stays the same, but nucleus is increasing. Increased attraction between nucleus and valence e-. This attraction pulls the e- closer to the nucleus and makes the atom smaller. Trend down a group: Increases Why: Number of energy levels increases & core e- increases. Each energy level is larger than the next. Valence e- farther from the nucleus and more blocked by the inner e-.
29
30 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
31 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
32 Ionic Radius: Distance between the nucleus and the outermost electron in ions (can t be determined directly). Trend between atom & ion: Cations are smaller than original atom. (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. (Adding negative e-, adds to repulsion between valence e-, pushing them apart.) Trend within the period (left to right): Representative Elements Decreases. Cations: 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. Same reason as atomic radii trend.
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34 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 Mg 2+ or O 2-7. Fe 2+ or Fe 3+
35 Ionization Energy: Energy required to remove an electron from a gaseous atom (also called First Ionization Energy, I 1 ) Na (g) kj Na + (g) + e - The second ionization energy, I 2, is the energy required to remove the next available electron: NOTICE: Na + (g) kj Na 2+ (g) + e - Ionization Energy increases for each electron removed from the same element. The larger ionization energy, the more difficult it is to remove the electron.
36 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.
37 Trend within the period: Increases Why: Electrons are more difficult to remove from smaller atoms. Closer to the nucleus and increased nuclear charge. Trend down a group: Decreases Why: Electrons are easier to remove from large atoms. Farther away from the nucleus so less energy is needed to remove them. Notice the trend in ionization energy is inversely related to trends in atomic radii.
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39
40 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
41 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
42 ELECTRONEGATIVITY: Ability of an atom to attract electrons in a chemical bond to itself. Chemist Linus Pauling set electronegativities on a scale. 0.7 (Cs) to 4.0 (F) Used to help determine types of bonding (ionic or covalent) that are occurring in a compound. Noble gases are not usually given electronegativity values. Trend within the period: Increases Why: Atoms become smaller, so shared electrons are closer to the nucleus. Trend down a group: Decreases Why: Atoms become larger, so shared electrons are farther from the nucleus.
43 Electronegativity
44 Examples put each set in order by increasing electronegativity: 1. Na, Li, Rb, K, Fr 2. Cl, Ca, F, P, Mg, S, K
45 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
46 Review: 1. As you move across a period, left to right, describe what generally happens (decreases, increases, or remains the same) to: a. Number of valence electrons b. Ionization energy c. Atomic radius 2. Give a brief explanation for your answers to a-c.
47 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.
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