Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
CHEMISTRY CHAPTER 11 & 12 THE ELEMENTS OF THE PERIODIC TABLE 1 Organization of the Elements • • • • Metals vs. non-metals. Listing according to mass. Listing according to properties Listing according to Chemistry – MO ratio; reactions with Hydrogen, halogens etc. Organization of the Elements • Dmitri Mendeleev (1860) – Established the Periodic Table – Stated the Periodic Law: – “Elements were placed on the table according to their atomic masses, with columns established by periodic recurring similar properties.” Organization of the Elements Mendeleev's Periodic Table •Vertical columns in atomic weight order •Mendeleev made some exceptions to place elements in rows with similar properties •Tellurium and iodine's places were switched Organization of the Elements •Horizontal rows have similar chemical properties •Missing Elements -Gaps existed in Mendeleev’s table •Mendeleev predicted properties of the “yet to be discovered” elements •Scandium, Germanium and Gallium agreed with predictions Unanswered Questions 1. Why didn't some elements fit in order of increasing atomic mass? 2. Why did elements exhibit periodic behavior? Moseley and the Periodic Table (1911) •Protons and Atomic Number •X ray experiments revealed a way to determine the number of protons in the nucleus of an atom •The periodic table was found to be in atomic number order, not atomic mass order • The tellurium-iodine anomaly was explained The Periodic Law •The physical and chemical properties of the elements are periodic functions of their atomic numbers The Modern Periodic Table •Discovery of the Noble Gases 1868 - Helium discovered as a component of the sun, based on the emission spectrum of sunlight 1894 - William Ramsay discovers argon 1895 - Ramsay finds helium on Earth 1898 - Ramsay discovers krypton and xenon 1900 - Freidrich Dorn discovers radon The Lanthanides Early 1900's the elements from cerium (#58) to lutetium (#71) are separated and identified Organization of the Elements The Modern Periodic Table – Listed according to atomic numbers. Organization of the Elements • The Modern Periodic Table – Vertical columns = Families • Group 1 (1A) = Alkali Metals • Group 2 (2A) = Alkaline Earth Metals Organization of the Elements •Group 17 (7A) Halogens •Group 18 (8A) Noble Gases •Group 3-12: Transitional Metals •Horizontal Row: Periods Periodic Table and econfiguration • • • • • Alkali - s1 Alkaline Earth - s2 Halogens - p5 Noble gases - p6 Valence Electrons= electrons in outer most shell. – The outermost shell electrons are the most easily removed. Periodic Table and econfiguration – The outermost (valence) electrons. – In the outermost s and p orbitals. • What is the maximum number of valence orbitals? FOUR • What is the maximum number of valence electrons? EIGHT Periodic Trends • Metallic properties: – Decrease across a period.(left to right) – Increase going down a column. • Ionization Energies: – Increase across a period. .(left to right) – Decrease going down a column. Ionization Energy •The energy required to remove one electron from a neutral atom of an element, measured in kilojoules/mole (kJ/mol) A + energy -> A+ + e- Ionization Energy Trends •Ionization energy of main-group elements tends to increase across each period (Left to Right) •Atoms are getting smaller, electrons are closer to the nucleus •Ionization energy of main-group elements tends to decrease as atomic number increases in a group Ionization Energy Trends •Atoms are getting larger, electrons are farther from the nucleus •Outer electrons become increasingly more shielded from the nucleus by inner electrons •Metals have a characteristic low ionization energy •Nonmetals have a high ionization energy •Noble gases have a very high ionization energy Ionization Energy Decreases Metallic Properties Increase Metallic Ionization Properties Energy Decrease Increases Periodic Trends • Atomic radii: – Decrease across a period. – Increase going down a column. •Atomic Radius -One half the distance between nuclei of identical atoms that are bonded together Ionic Radii Ionic radii: –Positive: Decrease across a period. Predominately Metals –Negative: Decrease across a period. Predominately non-metals Cations -Positively charged ions -Smaller than the corresponding atom -Less shielding of electrons Ionic Radii Anions -Negatively charged ions -Larger than the corresponding atom -Greater electron-electron repulsion -Ion size tends to increase downward within a group This table shows the relative sizes of atoms based on increasing ATOMIC RADIUS Size is dependent on energy level and electrons filling level. As you fill energy level across the period, the atomic radius decreases Argon fills the 3p orbitals and Potassium starts the 4s (larger sized) orbital. Calcium has more electrons filling level four shell which shrinks the atom size (negative shell attracts to positive nucleus As Atoms lose electrons the radius of the ion gets smaller As elements receive electrons the radius of the ion gets larger Atomic Radii Increase Atomic Radii Decrease Electronegativity • A measure of the ability of an atom in a chemical compound to attract electrons • Electronegativities tend to increase across a period • Electronegativities tend to decrease down a group or remain the same Example: F(4.0) vs. Cl(3.0), O(3.5) vs. Ba(0.9) • Elements that do not form compounds are not assigned electronegativities Table of Electronegativities Electronegativity Trends •Nonmetals have characteristically high electronegativity •Highest in the upper right corner •Metals have characteristically low electronegativity •Electron Electron Affinity Affinity: The energy change that occurs when an electron is acquired by a neutral atom, measured in kJ/mol •Most atoms release energy when they acquire an electron A + e- -> A- + energy (exothermic) •Some atoms must be forced to gain an electron A + e- + energy -> A - (endothermic) Electron Affinity Trends 1. Halogens have the highest electron affinities 2. Metals have characteristically low electron affinities 3. Affinity tends to increase across a period a. Irregularities are due to the extra stability of half-filled and filled sublevels Electron Affinity Trends 4. Electron affinity tends to decrease down a group 5. Second electron affinities are always positive (endothermic) Chemical Families • Alkali Metals: – Very reactive metals – All form 1+ ions and are found in nature as ions. Never found pure in nature – Sodium and Potassium are the most common in the family - Soft, silvery metals of low density and low melting points - Highly reactive, The Properties of a Group: the Alkali Metals • Easily lose valence electron (Reducing agents) •React violently with water Large hydration energy •React with halogens to form salts Chemical Families – React readily with halogens to form common salts • Example: NaCl (table salt) – React readily with water to form basic solutions (alkali), hydrogen and Energy. Chemical Families • Alkaline Earth Metals – 2nd most reactive metallic family. – All form 2+ ions. - Denser, harder, stronger, less reactive than Group 1 - Too reactive; not found pure in nature – Mg and Ca are the most common family members. Chemical Families • Halogens Group 17 – Very reactive non-metals, each having SEVEN valence electrons. – Valence Electrons: The electrons available to be lost, gained, or shared in the formation of chemical compounds (outer most) – Question: In what orbitals are the valence electrons? s and p Chemical Families • Halogens React readily with Alkali metals to form salts (salt formers). • React with hydrogen to form hydrogen halides (strong acids). Chemical Families • Halogens Group 17 are: – Natural occurring as diatomic molecules. – all COLORED. – all (1-) ions. Chemical Families • Noble Gases: Group 18 –Also known as “inert gases,” they are characteristically unreactive. Noble gases have EIGHT valence electrons. Chemical Families • Noble Gases: Group 18 – First found to react in 1962, but formed unstable compounds. Chemical Families • Transitional Metals – Columns 3-12 – Have more than one Oxidation state. – Most common: Iron – Includes “precious metals” – Compound often highly colored. Chemical Families – Gemstones most often are transitional metal compounds. • Metaloids: In between metals and non-metals – Include semiconductors • Non-Metals: Carbon, Oxygen, Nitrogen families Chemical Families – Allotropic forms are known of both metaloids and non-metals: same element having different structures. • Inner transitional elements: (Rare Earth Elements) – Lanthanide series (Atomic # 57-70) – Actinide series (#89-102) Properties of Metals • Metals are good conductors of heat and electricity • Metals are malleable • Metals are ductile • Metals have high tensile strength • Metals have luster Increases for successive electrons taken from the same atom Tends to increase across a period Electrons in the same quantum level do not shield as effectively as electrons in inner levels Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove Tends to decrease down a group Half of the distance between nucli in covalently bonded diatomic molecule "covalent atomic radii" Periodic Trends in Atomic Radius Radius decreases across a period Increased effective nuclear charge due to decreased shielding Radius increases down a group Addition of principal quantum levels Determination of Atomic Radius: Ionization of Magnesium Mg + 738 kJ Mg+ + eMg+ + 1451 kJ Mg2+ + eMg2+ + 7733 kJ Mg3+ + e- Affinity tends to increase across a period Affinity tends to decrease as you go down in a period Electrons farther from the nucleus experience less nuclear attraction Some irregularities due to repulsive forces in the relatively small p orbitals Electron Affinity - the energy change associated with the addition of an electron Ionization Energy Decreases Metallic Properties Increase Metallic Ionization Properties Energy Decrease Increases 5-2 Electron Configuration and the Periodic Table I. Periods and the Blocks of the Periodic Table A. Periods 1. Horizontal rows on the periodic table 2. Period number corresponds to the highest principal quantum number of the elements in the period B. Sublevel Blocks 1. Periodic table can be broken into blocks corresponding to s, p, d, f II. Blocks and Groups A. s-Block, Groups 1 and 2 1. Group 1 - The alkali metals a. One s electron in outer shell b. Soft, silvery metals of low density and lo c. Highly reactive, never found pure in natur 2. Group 2 - The alkaline earth metals a. Two s electrons in outer shell b. Denser, harder, stronger, less reactive t c. Too reactive to be found pure in nature B. d-Block, Groups 3 - 12 1. Metals with typical metallic properties 2. Referred to as "transition" metals 3. Group number = sum of outermost s and C. p-Block elements, Groups 13 - 18 1. Properties vary greatly a. Metals (1) softer and less dense than d-block meta (2) harder and more dense than s-block me b. Metalloids (1) Brittle solids with some metallic and properties (2) Semiconductors c. Nonmetals (1) Halogens (Group 17) are most react D. f-Block, Lanthanides and Actinides 1. Lanthanides are shiny metals similar i 2. Actinides a. All are radioactive b. Plutonium (94) through Lawrencium (10 5-3 Electron Configuration and Periodic P V. Valence Electrons A. Valence Electrons 1. 2. Main group element valence electrons are outermost energy level s and p sublevels Group # 1 2 13 14 15 16 17 18 Number of valence Electrons 1 2 3 4 5 6 7 8 Periodic Trends – Positive: Increase down a column. – Negative: Increase down a column. B. 1. C. 1. a. 2. Ionization Energy Tends to increase across d- and f-B Ion Formation and Ionic Radii Electrons are removed from the out Most d-block elements form 2+ ions Ions of d- and f-Blocks are cations D. Electronegativity 1. Characteristically low electronegativity of metals 2. Electronegativity increases as atomic radius decreases B. Trends 1. Atomic radius tends to decrease acr 2. Atomic radii tend to increase down a VII. Periodic Properties of the d- and fBlock Elements A. Atomic Radii 1. Smaller decrease in radius across a period within the d- Block than within the main-group elements a. Added electrons are partially shielded from the increasing positive nuclear charge b. Slight increase at the end of the dBlock is due to electron-electron repulsion 2. Little change occurs in radius across