Chemistry Science 09 Review!! Science 10 CT00D01 Topics from Science 09 Atomic Theory Subatomic Composition The Periodic Table Chemical Bonding Nomenclature and Chemical Reactions Atomic Theory Major Contributors: Dalton, Thomson, Rutherford, Bohr John Dalton England, 1808 Atomic Theory which states: An atom cannot be broken down into smaller parts. Atoms of the same element are exactly alike. Atoms of different elements are different. Atoms can combine in definite whole-number ratios to form compounds. Dalton Model Diagram: Shading is used in the sphere to show that its solid throughout. Every substance is made up of atoms Atoms are indestructible and indivisible Atoms can neither be created or destroyed J.J. Thomson England, 1897 Discovery Atoms contain electrons Thomson found the charge-to-mass ratio of an electron. Thomson Model Diagram This model is referred to as the Plum Pudding model. Thomson believed that the atom consisted of a positive sphere (the pudding) with electrons embedded in it (the raisins) Ernest Rutherford New Zealand, 1911 Discovery Atoms are made up of mostly empty space Atoms contain a small, dense, positively- charged nucleus. Gold Foil Experiment by Ernest Rutherford Positively charged radiation directed towards a thin sheet of gold foil Odd results: most radiation went through the foil, some scattered at wide angles, and some shot back at him! Rutherfords Model Diagram + Empty Space Dense, positively charged nucleus Niels Bohr Denmark, 1913 Discovery: Electrons contain specific amounts of energy and orbit the nucleus in specific paths, call energy levels Electrons must gain energy to move to a higher energy level or lose energy to move to a lower energy level Bohr Model Diagram The small particles in the center represent the protons and neutrons in the nucleus. e-e- e- Nucleus contains protons (+) and neutrons (o) Electrons exist in energy levels, 2-8-8, then. Subatomic Composition Atoms consist of: Protons, neutrons, electrons Isotopes, ions, anions, cations Three charges of an atom Mass of p & n are ~1000x greater than mass of e (making it insignificant) First look at the Periodic Table: What do the numbers mean? Atomic Number (p) Atomic Symbol A r (Relative At. Mass) For most common isotope, round it off. (For H would be 1 = 1p & 0n) How many of each are there? Protons: always the atomic # Electrons: same as p if neutral If -1, has one more e than p If +1, has one less e than p Neutrons: Depends on isotope For most common, round off the atomic mass to find the mass number, then use the formula A r = p + n Mass Number = protons + neutrons Atomic Number Number of protons contained in an atom (p) Determines the element Every element of that kind has the same number of protons Mass Number Number of protons and neutrons in an atom (p + n) Can change for various isotopes Ions Differing number electrons (e) Changes overall charge Ca t ions (positive) Anions ( negative ) Isotopes Differing number of neutrons (n) Changes Mass Number Isotopes An isotope is an element that contains a different number of neutrons and protons. Affects the molar mass of the element Longhand notation Element Mass Number Carbon 12 Carbon 14 Shorthand notation X Mass# Atomic # Charge The Periodic Table Layout: families, groups, etc Trends: Atomic number, mass, valence, charge Electronegativity, ionization energy, atomic radius, reactivity, metallic property HHe LiBeBCNOFNe NaMgAlSiPSClAr KCaSeTiVCrMnFeCoNiCuZnGaGeAsSeBrKr RbSrYZrNbMoTcRuRhPdAgCdInSnSbTeIXe CsBaLaHfTaWReOsIrPtAuHgTlPbBiPoAtRn FrRaAcRfDbSgBhHsMtDs CePrNdPmSmEuGdTbDyHoErTmYbLu ThPaUNpPuAmCmBkCfEsFmMdNoLr PERIODS - Similarities: The number of outer electron shells. HHe LiBeBCNOFNe NaMgAlSiPSClAr KCaSeTiVCrMnFeCoNiCuZnGaGeAsSeBrKr RbSrYZrNbMoTcRuRhPdAgCdInSnSbTeIXe CsBaLaHfTaWReOsIrPtAuHgTlPbBiPoAtRn FrRaAcRfDbSgBhHsMtDs CePrNdPmSmEuGdTbDyHoErTmYbLu ThPaUNpPuAmCmBkCfEsFmMdNoLr GROUPS Similarities: The number of electrons in the outer shell. Common reactivity, bonding, chemical and physical properties. H He LiBe BCNOFNe NaMg AlSiPSClAr KCaScTiVCrMnFeCoNiCuZnGaGeAsSeBrKr RbSrYZrNbMoTcRuRhPdAgCdInSnSbTeIXe CsBaLaHfTaWReOsIrPtAuHgTlPbBiPoAtRn FrRaAcRfDbSgBhHsMtDs CePrNdPmSmEuGdTbDyHoErTmYbLu ThPaUNpPuAmCmBkCfEsFmMdNoLr Metaloids Alkaline Earths Alkali Metals Transition Metals Non-metals Weak/Poor Metals Halogens Noble Gases Actinides Lanthanides HHe LiBeBCNOFNe NaMgAlSiPSClAr KCaSeTiVCrMnFeCoNiCuZnGaGeAsSeBrKr RbSrYZrNbMoTcRuRhPdAgCdInSnSbTeIXe CsBaLaHfTaWReOsIrPtAuHgTlPbBiPoAtRn FrRaAcRfDbSgBhHsMtDs CePrNdPmSmEuGdTbDyHoErTmYbLu ThPaUNpPuAmCmBkCfEsFmMdNoLr METALIC PROPERTIES Similarities: An elements relative ability to conduct energy in the form of heat or electricity. the stair Metals Non metals Metaloid s Valence Electrons The number of electrons in the outer shell of an atom Valence Electrons Periodic Law When the elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties Atomic Radius A measure of the distance from the center of the nucleus to the outer-most electron (Bottom Left, Fr has the largest Atomic Radius) Electronegativity An atoms ability or affinity to gain another electron. Trend (Top right, excluding the noble gasses, F is the most electronegative element) Table of Electronegativities Ionization Energy The energy required to lose the outer-most electron from an element. (Top right, He has the greatest ionization energy) Reactivity An atoms general ability to undergo a chemical reaction. The Alkali Metals are the most electronegative family/group. (Bottom Left, Fr is the most reactive) Highly reactive readily participates in chemical reactions Un-reactive does not readily participate in chemical reactionsniac Atomic Radius: A measure of the distance from the center of the nucleus to the outer-most electron Electronegativity: An atoms ability or affinity to gain another electron. Ionization Energy: The energy required to lose the outer-most electron from an element. Reactivity: An atoms general ability to undergo a chemical reaction. Trends of the Periodic Table 4 major trends: top right or bottom left? H He LiBe BCNOFNe NaMg AlSiPSClAr KCaSeTiVCrMnFeCoNiCuZnGaGeAsSeBrKr RbSrYZrNbMoTcRuRhPdAgCdInSnSbTeIXe CsBaLaHfTaWReOsIrPtAuHgTlPbBiPoAtRn FrRaAcRfDbSgBhHsMtDs CePrNdPmSmEuGdTbDyHoErTmYbLu ThPaUNpPuAmCmBkCfEsFmMdNoLr Electronegativity (F) & Ionization Energy (He) Atomic Radius (Fr) & Reactivity (Fr) Chemical Bonding Inter- vs Intra-molecular forces Ionic vs Covalent Molecular Interactions Inter-molecular Forces Interaction between molecules that hold it together in a network. Intra-molecular Forces Forces that hold groups of atoms together and make them function as a unit Intra-molecular Forces: Bonding Forces that hold groups of atoms together and make them function as a unit. Ionic bonds transfer of electrons Ionic bonds Covalent bonds sharing of electrons Covalent bonds Metallic Bonding sea of electrons Metallic Bonding Which elements form which type of bond? Metal / nonmetal = Ionic NaCl, MgBr 2 Nonmetal / nonmetal = covalent CO 2, CH 4, H 2 O, NO Group like metals = metalic Fe, Ti, Mg Review: Why do atoms bond? To satisfy the octet rule? Yes, but to be more specific, atoms share electrons in order to complete their outer electron shell making them more stable as they are then in a lower state of energy. But how do we know what type of bonding will occur between two atoms? Metal Metalloids Nonmetals IIIIIIIVVVIVIIVIII Transition metals Lewis Dot Diagrams Lewis Dot Diagrams are used in both ionic and covalent bonding Simple Rules for Bonding Duet Rule: H and He require 2 electrons for stability Octet Rule: All other elements will have 8 valence electrons for stability In order to achieve this elements can steal (ionic) or share (covalent) electrons Ions form when atoms lose or gain electrons. Atoms with few valence electrons tend to lose them to form cations. Atoms with many valence electrons tend to gain electrons to form anions NeN Na F Na + N 3- F-F-F-F- O O 2- Mg Mg 2+ CationsAnions 41 Ionic Bonds Ionic bonds result from the attractions between positive and negative ions. Ionic bonding involves 3 aspects: 1. loss of an electron(s) by one element, 2. gain of electron(s) by a second element, 3. attraction between positive and negative 42 Ionic Bonds Electrons are transferred Electronegativity differences are Greater than for covalent The formation of ionic bonds is always exothermic! A covalent/molecular bond is a chemical bond in which two or more electrons are shared by two atoms. Why should two atoms share electrons? FF + 7e - FF 8e - F F F F Lewis structure of F 2 lone pairs single covalent bond Covalent Bonds 8e - H H O ++ O HH O HHor 2e - Lewis structure of water Double bond two atoms share two pairs of electrons single covalent bonds O C O or O C O 8e - double bonds Triple bond two atoms share three pairs of electrons N N 8e - N N triple bond or Covalent Bonds Linear 180 o Covalent Bonds Trigonal Planar 120 o Covalent Bonds Tetrahedral o Covalent Bonds Pyramidal (Tetrahedral) 107 o Covalent Bonds Bent (Tetrahedral) o Covalent Bonds Boron Trifluoride Non-Polar (even) F B F Trigonal planar shape: BF 3 violates the octet rule Covalent Bonds .. HNH H Ammonia NH 3 Polar (uneven) Covalent Bonds Trigonal Bipyramidal / Octahedral 120 o 90 o Covalent Bonds Explain intra/inter molecular bonding, the difference between covalent/ionic bonds, how you determine bonding type, and types of bonds: Review Covalent Bonds Chemical Reactions & Nomenclature Nomenclature for Ionic, covalent, acids Types of Chemical Reactions Balancing Look at the First Element!! A metal = MUST BE IONIC Does it have a polyatomic? Does it have a transition metal? A nonmetal = MUST BE COVALENT Hydrogen = MUST BE AN ACID Binary Acid? Oxoacid (Polyatomic)? first element Metal = Ionic NO prefixes The simplest whole number ratio is generally the ionic formula. (empirical formula) Binary NaCl MgBr 2 Li 2 S Polyatomic Na 2 CO 3 Mg(NO 3 ) 2 Li 3 PO 4 Transition FeCl 3 Ni 3 (PO 3 ) 2 ZrSO 4 Polyatomic Ions ( Common 11) PO 4 3- PO 3 3- OH - NO 2 - NO 3 - NH 4 + SO 3 2- SO 4 2- CO 3 2- ClO 3 - ClO 2 - First element nonmetal = Covalent Greek Prefixes Mono (1), di (2), tri (3), tetra (4), penta (5), etc. First nonmetal keeps element name Change ending of second nonmetal to ide P 2 O 5 CCl 4 CO 2 COSiO 2 NO 2 CF 4 SF 6 PF 5 NO First element hydrogen = Acid Binary H and another element Use hydro- prefix Use ic suffix HClHFHBr H 2 SHI Oxoacid / Poly H and 2 or more NO prefix If poly ate -ic If poly ite -ous HNO 3 H 3 PO 4 H 3 PO 3 H 2 SO 4 Reaction Types Synthesis/Composition (A + B AB) Decomposition (AB A + B) Single Replacement (A + BC B + AC) Double Replacement (AB + CD CB + AD) Combustion (C x H y + O 2 CO 2 + H 2 O) Balancing Equations Conservation of Mass There must be the same amounts of each element on each side __ C 3 H 8 + ___ O 2 ___ CO 2 + ___ H 2 O ___ KClO 3 ___ KCl + ___ O 2 __ N 2 O 3 + __ H 2 O __ HNO 2