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Periodic Table Chapter 5 History Organization Introduction to Bonding Trends Terms • The Periodic Law States that the physical & chemical properties of the elements are periodic functions of their atomic numbers. • The Periodic Table An arrangement of the elements in order of their atomic numbers so that elements with similar properties fall in the same columns. • The Modern Periodic Table Is organized by electron configuration. Periodic Table & electron configurations Blocks of the Periodic Table s p d f Video Clip: Periodic Table by cassiopeiaproject • https://youtu.be/5MMWpeJ5dn4 Extended Periodic Table History 1817-Dobereiner’s Triads Groups of 3 elements with similar properties, the atomic mass of second was an average of all three. EX: Ca = 40 Sr = 88 Ba = 137 History 1863- Newland’s Law of Octaves If elements are arranged by increasing atomic mass, their properties repeat every 8 elements. Chapter 5 – Section 1: History of the Periodic Table Mendeleev and Periodicity • The first periodic table of the elements was published by Russian chemist Dmitri Mendeleev. • Mendeleev left empty spaces in his table and predicted elements that would fill 3 of the spaces. • By 1886, all 3 of these elements had been discovered. History 1869- Mendeleev’s First Periodic Table He arranged all the known elements in rows by increasing mass AND he grouped elements with similar properties in columns. Brilliant in spite of discrepancies: Co & Ni, Te & I Brilliant because of his predictions: ekasilicon = Ge Video Clip: Periodic Table of Elements - Chemistry: A Volatile History - BBC • https://youtu.be/nsbXp64YPRQ Chapter 5 – Section 1: History of the Periodic Table Mosley and the Periodic Law • In 1911, the English scientist Henry Moseley discovered that the elements fit into patterns better when they were arranged according to atomic number, rather than atomic weight. • The Periodic Law states that the physical and chemical properties of the elements are periodic functions of their atomic numbers. History 1913- Moseley discovers protons Nuclear charge (protons) increases by one for each element! Atomic number, not atomic mass, is the basis for organization! Organization Groups, Periods and Blocks • Elements in the periodic table are arranged into vertical columns, called groups or families, that share similar chemical properties. • Elements are also organized horizontally in rows, or periods. GROUPS • 18 vertical columns (families) with family names • same outer e- configuration. • similar chemical properties. • labeling columns has changed over time… Labeling Groups… then & now PERIODS • 7 horizontal rows • corresponds to main Energy level that is filled • length of row determined by sublevels that are filling Video Clip: The (truly) Periodic Table by scienceoffice.org • https://youtu.be/xd4-Uy2FLWc BLOCKS The representative elements of the Main Groups: s & p blocks only Group 1: Alkali Metals • Group 1 elements are called alkali metals. • Alkali metals have a silvery appearance and are soft enough to cut with a knife. • They are extremely reactive and are not found in nature as free elements. • They must be stored under oil or kerosene. Group 2: Alkaline Earth Metals • Elements in group 2 are known as the alkaline earth metals. • Group 2 metals are harder, denser and stronger than alkali metals, and have higher melting points. • Less reactive than group 1, but still too reactive to be found in nature as free elements. Group 17: Halogens • Elements in group 17 are known as the halogens. • Halogens are the most reactive nonmetals, reacting vigorously with metals to form salts •Most halogens exist in nature as diatomic molecules (i.e. F2, Cl2, Br2 and I2.) Group 18: Noble Gases • Elements in group 18 are known as noble gases. • They are completely nonreactive and don’t form compounds under normal conditions. • A new group was added to the periodic table in 1898 for the noble gases. d-block: Transition Metals • Elements in the d-block are called transition metals. • They have typical metallic properties such as conduction of electricity and high luster. • Less reactive than group 1 and 2 elements. • Some (i.e. platinum & gold) are so unreactive they usually don’t form compounds. f-block: Lanthanides & Actinides • Elements in period 6 of the f-block are called lanthanides (or rare-earth). • Lanthanides are shiny metals similar in reactivity to alkaline earth metals. • Elements in period 7 of the f-block are called actinides. • Actinides are all radioactive, and many of them are known only as man-made elements. Classification & Location of the elements Metals: left of staircase Nonmetals: right of staircase Semimetals: touch the staircase (except Al) Why are H & He sometimes separated from Periodic Table? • Hydrogen has one valence electron so it is normally placed in group 1. • Hydrogen has a second home at the top of group 17 because it can gain an electron as the halogens do. • Helium has two valence electrons like group 2 but is clearly a noble gas because its outer energy level is full. Introduction to Bonding PHYSICAL PROPERTIES OF METALS luster, conductivity, solid, malleable, ductile 3 or fewer valence e- = defines a metal Introduction to Bonding CHEMICAL PROPERTIES OF METALS….HOW THEY FORM BONDS Lose valence e- to nonmetals → Ionic compounds Bond with each other → Metallic bonds Mix with other metals → Alloys Introduction to Bonding PHYSICAL PROPERTIES OF NONMETALS No luster, poor conductors, not malleable or ductile 5 -8 valence e- = defines a nonmetal Can be soft solids: C, P, S, Se, I2 Can be liquids: Br2 Can be gases: N2, Cl2, H2, O2, F2, & the noble gases All gases are nonmetals but not all nonmetals are gases! Introduction to Bonding CHEMICAL PROPERTIES OF NONMETALS …. HOW THEY FORM BONDS Gain valence e- from metals → Ionic compounds (EX: salt: NaCl) Share valence with themselves→ Diatomic Molecules (EX: N2, H2 ) Share valence with other nonmetals → Molecules (EX: H2O, CO2) Chapter 5 – Section 3: Electron Configuration and Periodic Properties Atomic Radii • Atomic radius – one-half the distance between the nuclei of identical atoms that are bonded together. Group 1 •Atomic radii tend to increase as you go down a group because electrons occupy successively higher energy levels farther away from the nucleus. Chapter 5 – Section 3: Electron Configuration and Periodic Properties Atomic Radii (continued) • Atomic radii tend to decrease as you go across a period because as more electrons are added they are pulled closer to the more highly charged nucleus with more protons. Period 2 Atomic Radius (Size) Chapter 5 – Section 3: Electron Configuration and Periodic Properties Atomic Radii Sample Problem Of the elements Mg, Cl, Na, and P, which has the largest atomic radius? Explain. Solution: Na has the largest radius. All of the elements are in the 3rd period, and atomic radii decrease across a period. Chapter 5 – Section 3: Electron Configuration and Periodic Properties Ionization Energy • An ion is an atom of group of bonded atoms that has a positive or negative charge. • The energy required to remove an electron from a neutral atom of an element is called the ionization energy (IE). • Ionization energy tends to increase across each period because a higher nuclear charge more strongly attracts electrons in the same energy level. Chapter 5 – Section 3: Electron Configuration and Periodic Properties Ionization Energy (continued) • Ionization energy tends to decrease down each group because electrons farther from the nucleus are removed more easily. Chapter 5 – Section 3: Electron Configuration and Periodic Properties Ionization Energy Sample Problem Consider two elements, A and B. A has an IE of 419 kJ/mol. B has an IE of 1000 kJ/mol. Which element is more likely to be in the s block? Which will be in the p block? Which is more likely to form a positive ion? Solution: Element A is most likely to be in the s-block since IE increases across the periods. Element B would most likely lie at the end of a period in the p block. Element A is more likely to form a positive ion since it has a much lower IE than B. Ionization Energy Chapter 5 – Section 3: Electron Configuration and Periodic Properties Electron Affinity and Electronegativity • Electron affinity is the energy change that occurs when an electron is acquired by a neutral atom. • Electronegativity is a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound. • Electronegativity applies to atoms in a compound, while electron affinity is a property of isolated atoms. Chapter 5 – Section 3: Electron Configuration and Periodic Properties Electron Affinity and Electronegativity (continued) • Electron affinity and electronegativity both tend to increase across periods, and decrease (or stay the same) down a group. Electron Affinity ELECTRONEGATIVITY Chapter 5 – Section 3: Electron Configuration and Periodic Properties Electronegativity Sample Problem Of the elements Ga, Br, and Ca, which has the highest electronegativity? Explain . Solution: All of these elements are in the fourth period. Br has the highest atomic number and is farthest to the right in the period. Br would have the highest electronegativity since electronegativity increases across a period. Periodic Trends • Atomic Radius, and Metallic Character • Factors affecting size: Distance of outer e- from nucleus, shielding effect of inner e- and size of the ecloud with each sublevel makes atoms larger down a group. Nuclear charge across the period pulls e- cloud in and makes the atoms smaller. Periodic Trends • Ionic Radius • Factors affecting size: Metals lose valence e- the ion is smaller = previous noble gas! Nonmetals gain valence eion is larger = next noble gas! Periodic Trends First Ionization Energy: E req’d to remove outer eElectron Affinity: Atom’s attraction for e- Metals low IE low e- affinity Nonmetals high IE high e- affinity Periodic Trends-Bonded Atoms Electronegativity: The ability of an atom to attract e- in a chemical bond. Fluorine is assigned highest value = 4.0 Nonmetals high EN high e- affinity Metals low EN low e- affinity START CH.7 valence electrons • Occupy highest principal energy level • Responsible for chemical properties = bonding • Elements in a group have similar properties b/c valence e- same. • Outer E level has e- in s & p sublevels only • Outer E level full when s & p sublevels are full = noble gas configuration Ion formation in atoms Ion: an atom that has a charge b/c it has lost or gained e The Octet (Duet) Rule: atoms will gain, lose or share e- in order to acquire a full set of valence e- = 8 The Noble Gas Rule: atoms attain the nearest noble gas configuration when they become ions. Types of Ions CATIONS are positive ions ANIONS are negative ions Metals form monatomic cations To name: use name of element Na+1 Mg+2 magnesium ion…) Al+3 Si+4 (sodium ion, Types of Ions Nonmetals form monatomic anions To name: use name of element + ide N-3 O-2 F-1 (nitride ion, oxide ion, P-3 S-2 Cl-1 (phosphide, sulfide, chloride) fluoride ion) Transition Metals: d block • 1 or 2 valence e• Several of these metals form more than one type of ion. (MULTIPLE OXIDATION STATES) • How? Will lose the valence s e- then the d e- one at a time. • To name, use roman numerals Cu+1 ion Cu+2 ion ion Fe+2 iron II copper II ion Fe+3 iron III copper I Lower left of p block = metals • 3 to 5 valence e• Several of these metals form more than one type of ion. • How? Will lose all the valence s & p electrons OR just the p electrons. • To name, use roman numerals Tl+1 ion Tl+3 thallium I ion Sn+2 tin II thallium III ion Sn+4 tin IV Types of Ions Polyatomic ions: groups of covalently bonded atoms that carry a charge. NH4+1 ammonium ion SO4-2 sulfate ion OH-1 hydroxide ion CN-1 cyanide ion PO4-3 phosphate ion CO3-2 carbonate ion NO -1 nitrate ion