1	Chapter 7 Periodic Properties of the elements
2	Periodic Table Periodic table constructed in 1880’s by Mendeleev and Meyer Before subatomic particles were known (no atomic numbers!!). Elements in same group have similar chemical properties.
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5	Atomic Numbers Mosely (1913) frequency of x-rays emitted by different elements when hit with high energy electrons were different for each element. Could relate the frequency to a number called atomic number. Atomic number, Z: number of protons in nucleus. Also equal to number of positive charges in the nucleus. Effective nuclear charge, Z_eff: outer-lying electrons are shielded by inner-lying electrons effectively reducing amount of positive charge to Z_eff= Z – S.
6	Effective nuclear charge Influences various properties of the atom. The greater the effective nuclear charge, the greater the positive charge on the electron which tends to pull it closer to the nucleus.
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8	Periodic Properties Periodic table: some properties of elements repeat themselves as you move down the periodic table. These properties are called periodic properties. If you know how these vary, you can tell something about the properties of elements based solely on their position in the periodic table.
9	Atomic radius Atoms are considered to be little “billiard balls” with a particular radius. Two bonded spherical atoms will have a distance between them that is twice the radius of these billiard balls. Atomic radius: defined as exactly half the bond distance between two bonded atoms of the same element.
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11	Atomic radius Two factors influence atomic radii: Outer n shell: As n increases, average distance of the electron from the nucleus increases, atomic radius will increase. Z_eff: As effective nuclear charge increases, the electrons get pulled in closer to the nucleus, atomic radius will decrease Atomic radius patterns As you move from left to right in the same period, atomic radii tend to decrease due to increasing effective nuclear charge. As you move from top to bottom in the same group, atomic radii tend to increase due to filling of higher n shells.
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13	Ionic Radius When an atom loses an electron to form a cation, the resulting cation is much smaller. When an atom gains an electron to form an anion, the resulting ion is much larger. This can be observed in Ionic Radii.
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15	Isoelectronic Species 2 species with the same number of electrons are called isoelectronic. In an isoelectronic series of ions, the largest is the one with the largest negative charge and the smallest is the one with the largest positive charge.
16	Ionization Energy Ionization Energy, I: defined as the energy required to remove an electron from a gaseous atom or ion. Mg(g) ¾® Mg⁺(g) + e¯(g) DE = I₁(Mg) = 738 kJ/mol Mg⁺(g) ¾® Mg²⁺(g) + e¯(g) DE = I₂(Mg) = 1450 kJ/mol Mg²⁺(g) ¾® Mg³⁺(g) + e¯(g) DE = I₃(Mg) = 7730 kJ/mol An atom can have as many ionization energies as it has electrons
17	Ionization Energy (cont.) I₁< I₂< I₃…
18	First Ionization Energy, I₁ As you move from left to right in the same period, I₁ tends to increase due to increasing effective nuclear charge. As you move from top to bottom in the same group, I₁ tends to decrease due to filling of higher n shells.
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21	Electron Configuration of Ions Start with neutral atom ground state electron configuration For cations: Remove one electron for each positive charge. Electrons are removed from outer n shells first. (4s electrons are removed before 3d electrons) For anions: Add one electron for each negative charge. Electons are added according to the Aufbau procedure.
22	Electron Configuration of Ions What are the ground state electron configurations of the following ions: Mg²⁺; N^3-; Mn²⁺, Fe³⁺, Br¯
23	Noble gas electron configuration Noble gases are “extra” stable because they have filled s and p subshells called a closed shell. Many ions have the same configuration which is called a noble gas electron configuration. The group charges for A groups are noble gas electron configurations.
24	Electron Affinity EA is defined as the energy required to add an electron from a gaseous atom or ion. Cl(g) + e¯(g) ¾® Cl¯(g) DE = –349 kJ/mol For most elements, the electron affinity is a negative number. Most negative electron affinity is for Cl. As you move from left to right in the same period, EA’s tend to get more negative.
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27	Metal Have a luster, most are silver colored Most are solids, one liquid metal Malleable and ductile Good conductors of electricity and heat Have low ionization energies so they form cations easily. Metal oxides are ionic bases.
28	Metal oxides Ionic oxides that contain the O^2- ion which is a base. For metal oxides that dissolve in water Metal oxide + water ¾® metal hydroxide Li₂O(s) + H₂O(l) ¾® 2 LiOH(aq) Metal oxides react in acid solution in an acid base reaction: Metal oxide + acid ¾® salt + water MgO(s) + 2 HCl(aq) ¾® MgCl₂(aq) + H₂O(l)
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30	Nonmetals Do not have a luster (dull) with various colors. Brittle (some are hard and some are soft) Poor conductors of heat and electricity. Tend to have more negative electron affinities which makes them form anions easily Nonmetal oxides are covalent acids.
31	Nonmetal oxides Covalent compounds that do not contain any ions!! They are acidic when dissolved in water: CO₂(g) + H₂O(l) ¾® H₂CO₃(aq) SO₃(g) + H₂O(l) ¾® H₂SO₄(aq) They react with bases in an acid-base reaction: Nonmetal oxide + base ¾® salt + water CO₂(g) + NaOH(aq) ¾® Na₂CO₃(aq) + H₂(l)
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33	Basic reaction between metals and nonmetals For the main groups, a metal and a nonmetal will react to form a salt with the metal cation and the nonmetal anion. Sr(s) + Cl₂(g) ¾® SrCl₂(s) 6 Li(s) + N₂(g) ¾® 2 Li₃N(s) 4 Al(s) + 3 O₂(g) ¾® 2 Al₂O₃(s)
34	Metalloids Elements that border the stepwise line: Si, Ge, As, Sb, Te Properties lie between metallic and nonmetallic. Silicon: very pure silicon has a luster but is a poor conductor and is brittle.
35	Metallic Character The more an element acts like a metal, the more metallic character it has. Metallic character is a periodic property As you move from left to right in the same period, metallic character decreases. As you move from top to bottom in the same group, metallic character increases.
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37	Group 1A: Alkali Metals Soft, metallic solids. Silvery metallic luster and high thermal and electrical conductivity. “Alkali” derived from the Arabic word for ashes. Na, K: among the most abundant elements on Earth. They are found in salts as the metal cation.
38	Alkali Metals Elements have low densities and melting points. The elements are extremely reactive with reactivity increasing as you move down the periodic table. Low ionization energies. Chemistry dominated by the element’s desire to lose an electron to form a +1 cation.
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40	Alkali Metals React with nonmetals to produce a salt with the metal cation and the nonmetal anion: 2 K(s) + H₂(g) ¾® 2 KH(s) 2 Na(s) + S(s) ¾® Na₂S(s) React with water to produce the metal hydroxide and hydrogen gas: 2 Na(s) + 2 H₂O(l) ¾® 2 NaOH(aq) + H₂(g)
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44	Alkali metals Lithium reacts with oxygen to form lithium oxide: 4 Li(s) + O₂(g) ¾® 2 Li₂O(s) Sodium reacts with oxygen to form sodium peroxide: 2 Na(s) + O₂(g) ¾® Na₂O₂(s) K, Rb, and Cs react with oxygen to form the metal superoxide: K(s) + O₂(g) ¾® KO₂(s)
45	Group 2A: Alkaline Earth Metals Metallic solids that are harder and more dense than the alkali metals. They are less reactive than group 1A but still pretty reactive. Chemistry is dominated by the element’s tendency to lose 2 electrons to form +2 cations.
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47	Alkaline Earth Metals Reaction with water Be: does not react with water Mg: reacts with steam (hot water) to form magnesium oxide and hydrogen: Mg(s) + H₂O(g) ¾® MgO(s) + H₂(g) Ca, Sr, Ba: react with water to form the metal hydroxide and hydrogen: Sr(s) + 2 H₂O(l) ¾® Sr(OH)₂(aq) + H₂(g)
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49	Alkaline Earth Metals React with most nonmetals to form a salt with a +2 cation and the nonmetal anion. Mg²⁺ and Ca²⁺ are essential for living organisms. Skeletons: Ca²⁺ salts!!
50	Flame tests Many group 1A and group 2A emit a specific color when burned in a flame. Li⁺- crimson red Mg²⁺-not in visible Na⁺- yellow-orange Ca²⁺-brick red K⁺- lilac (purple) Sr²⁺-crimson red Ba²⁺- yellow-green These colors are used in fireworks.
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52	Nonmetals: Hydrogen Only element that is placed in two groups: Group 1A and Group 7A. Chemistry is not like either group. H₂ is a nonmetal with a large ionization energy unlike Group 1A. At high pressures, H₂ can be metallic. Hydrogen reacts with active metals to form metal hydrides which contains H¯ ion.
53	Group 6A: chalcogens or the oxygen family From top to bottom go from nonmetallic elements to metallic. O, S, Se: nonmetals Te: metalloid Po: metal, radioactive, and very rare.
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55	Elemental Forms of Group 6A Oxygen Most stable elemental form: O₂(g) Another elemental form: O₃(g) called ozone Formed from oxygen gas in electrical discharges 3 O₂(g) ¾® 2 O₃(g) Ozone is extremely toxic. It is an allotrope of oxygen. Allotropes: different forms of the same element.
56	Oxygen Extremely reactive. Chemistry dominated by element’s desired to gain two electrons to form O^2- (oxide ion) Other oxygen anions: O₂^2- peroxide ion; O₂¯ superoxide. Both are extremely reactive and decompose to form O^2- and O₂.
57	Sulfur Exists in several allotropic forms. Most stable elemental form is a yellow solid that consists of S₈ molecules. Sulfur tends to gain two electrons to form the sulfide ion. Not as reactive as oxygen.
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59	Group 7A: Halogens Halogen derived from Greek meaning “salt-formers”. Found as –1 anions in salts. All are nonmetals. F₂, Cl₂ are diatomic gases. F₂ is pale yellow. Cl₂ is yellow-green. Br₂ is a volatile red-brown liquid I₂ is a purple solid.
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62	Halogens Chemistry dominated by the element’s tendency to gain one electron to form an anion of charge –1 called a halide ion. Reactivity decreases as you go down the group. F₂ reacts with almost anything to form fluoride. Very difficult and dangerous to work with.
63	Chlorine gas Most industrially useful of halogens Reacts with water to form hypochlorite ion: Cl₂(g) + H₂O(l) ¾® HCl(aq) + OCl¯(aq) This reaction is used to disinfect water supplies and bleach colors out of paper and other things.
64	Halogen reactions React with hydrogen to form hydrogen halides: H₂(g) + Cl₂(g) ¾® 2 HCl(g) HCl, HBr, HI: strong acids, HF: weak acid.
65	Group 8A: The Noble Gases Nonmetals Exists as extremely unreactive monatomic gases. Before 1962, it was thought that these gases did not react with anything so they were called the Inert Gases. In 1962, Neil Bartlett was able to form a compound of xenon (Xe) with fluorine so these gases are not inert. Name was changed to the Noble Gases.
66	Group 8A Since 1962, compounds of Xe with F, N, and O have been prepared as well as compounds with Kr with F.
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