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The Components of Matter Chapter Two AP Chemistry Element Simplest type of substance with unique physical and chemical properties. Consists of only one type of atom. An element cannot be broken down into any simpler substances by physical or chemical means. A few elements exist in nature bonded with like atoms, e.g. P4, S8, Se8. The diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2 Molecule Structure that consists of two or more atoms which are chemically bound together. Behaves as an independent unit. Which is an element and which is a molecule? i Molecule Element Compound A substance composed of two or more elements which are chemically combined. When a compound is formed, there is evidence of a chemical reaction. Evidence of a chemical reaction: •Energy change heat, light, sound, motion, electricity, etc. •Release of a gas •Color Change •Precipitate is formed Evidence of Rxns (Evidence of Rxns.ASF) Compounds have a fixed mass ratio … Water is always 2 g H and 16 g O. Properties of compounds are different from the individual elements that make them up. Na and Cl2 are different from NaCl. Mixture A group of two or more elements and/or compounds that are physically intermingled. The individual properties of the components of a mixture are retained. When a mixture is formed, there is no evidence of a chemical reaction. Compound 11 Mixture Components are present in one constant ratio only Components may be present in a range of ratios Components may only be separated chemically Components may be separated physically Components lose their identities Components retain their identities An energy change between components and products takes place No evidence of an energy change Which is a compound and which is a mixture? Mixture 12 Compound Distinguishing Elements, Compounds, and Mixtures at the Atomic Scale Theses scenes represent an atomic-scale view of three samples of matter. Describe each sample as an element, compound, or mixture. (a) mixture (b) element (c) compound The Law of Mass Conservation: Mass remains constant during a chemical reaction. reactant 1 + product reactant 2 total mass = total mass calcium oxide + carbon dioxide calcium carbonate CaO + CO2 CaCO3 56.08 g + 44.00 g 100.08 g Law of Definite (or Constant) Composition: No matter what its source, a particular chemical compound is composed of the same elements in the same parts (fractions) by mass. CaCO3 1 atom Ca 1 atom C 3 atoms O 40.08 amu 12.00 amu 3 x 16.00 amu 100.08 amu 40.08 amu 100.08 amu = 0.401 parts Ca 12.00 amu 100.08 amu = 0.120 parts C 48.00 amu 100.08 amu = 0.480 parts O • 100 g CaCO3 • …always contains –40g of Ca –12g of C –48g of O Analysis by Mass (grams/20.0 g) Mass Fraction (parts/1.00 part) 0.40 calcium 0.12 carbon 0.48 oxygen 8.0 g calcium 2.4 g carbon 9.6 g oxygen 1.00 part by mass 20.0 g Percent by Mass (parts/100 parts) 40% calcium 12% carbon 48% oxygen 100% by mass Calculating the Mass of an Element in a Compound Ammonium Nitrate How much nitrogen (N) is in 455kg of ammonium nitrate? Formula Mass of Cpd: Add masses from the P.T. 4 x H = 4 x 1.008 = 4.032 g 2 x N = 2 X 14.01 = 28.02 g 3 x O = 3 x 16.00 = 48.00 g 80.052 g Therefore g nitrogen/g cpd 28.02g N = 0.3500gN/g cpd 80.052g cpd Note: .3500 X 100 = 35% N This is % Composition 455kg x 1000g = 455,000g NH4NO3 kg 455,000g cpd x 0.3500g N g cpd = 1.59 x 105g nitrogen Or: 28.02 kg N 455 kg NH4NO3 X 80.052 kg NH4NO4 = 159 kg Nitrogen Calculating the Mass of an Element in a Compound Pitchblende (PB) is a cpd that contains U. Analysis shows that 84.2 g of pitchblende contains 71.4 g of U with O as the only other element. How many grams of U can be obtained from 102 kg of pitchblende? The mass ratio of U/PB will be the same no matter the amount of pitchblende. 102kg PB x 71.4kg U 84.2 kg PB = 86.5 kg U 86.5 kg U = 86500 g U Law of Multiple Proportions 003_MULTIPLEPROP.mov If elements A and B react to form two compounds, the different masses of B that combine with a fixed mass of A can be expressed as a ratio of small whole numbers. The Atomic Basis of the Law of Multiple Proportions CO, CO2 CO = 12g C + 16g O CO2 = 12g C + 32g O 16g O/32g O = 1/2 H2O, H2O2 H2O = 2g H + 16g O H2O2 = 2g H + 32g O 16g O/32g O = 1/2 NO2, N2O, N2O3, N2O5 Dalton’s Atomic Theory The Postulates 1. All matter consists of atoms. 2. Atoms of one element cannot be converted into atoms of another element. 3. Atoms of an element are identical in mass and other properties and are different from atoms of any other element. 4. Compounds result from the chemical combination of a specific ratio of atoms of different elements. Dalton’s Model of the Atom Explains the Mass Laws Mass conservation Atoms cannot be created or destroyed postulate 1 or converted into other types of atoms. postulate 2 Since every atom has a fixed mass, postulate 3 during a chemical reaction atoms are combined differently, and therefore, there is no mass change overall. Definite Composition Atoms are combined in compounds in specific ratios and each atom has a specific mass. postulate 3 postulate 4 Each element has a fixed fraction of the total mass in a compound. Multiple Proportions Atoms of an element have the same mass and atoms are indivisible. postulate 3 postulate 1 When different numbers of atoms of elements combine, they must do so in ratios of small, whole numbers. Visualizing the Mass Laws Theses scenes represent an atomic-scale view of a chemical reaction. Which of the mass laws: mass conservation, definite composition, or multiple proportions is (are) illustrated? Mass conservation illustrated if number of each atom before and after reaction remains constant. Definite composition illustrated by formation of compounds that always have the same atom ratio. Different compounds made of same elements have small whole number ratios of those elements illustrates multiple proportions. Seven purple and nine green atoms in each circle, mass conserved. One compound formed has one purple and two green, definite composition. Law of multiple proportions does not apply. Properties of Cathode Rays. Observation Thompson’s Experiment Conclusion Ray bends in magnetic field Consists of charged particles Ray bends toward positive plate in electric field Consists of negative particles Ray is identical for any cathode Particles found in ALL matter Determined Q/m as 1.756 x 108 C/g Credited with discovering the electron + anode - cathode Cathode Thompson’s Model of the Atom + + + + + + + + + + + + + + Millikan’s Experiment Millikan In a replication of Millikan’s oil drop experiment you observe the following charges, measured in coulombs, on different oil drops: -9.6 x 10-19 -1.328 x 10-17 -7.84 x 10-18 -4.16 x 10-18 -4.64 x 10-18 -1.408 x 10-17 -8.16 x 10-18 -1.28 x 10-17 -1.488 x 10-17 -1.456 x 10-17 -1.6 x 10-18 -8.0 x 10-19 -1.12 x 10-18 -1.104 x 10-17 -4.16 x 10-18 -1.6 x 10-19 -3.84 x 10-18 -6.4 x 10-19 -3.2 x 10-19 -1.28 x 10-18 -4.8 x 10-19 -6.24 x 10-18 -9.44 x 10-18 -5.44 x 10-18 -1.104 x 10-17 -1.28 x 10-17 -6.24 x 10-18 -2.72 x 10-18 -1.232 x 10-17 -1.28 x 10-17 What should you conclude is the charge of an individual electron? -9.6 x 10-19 (6) -1.328 x 10-17 (83) -7.84 x 10-18 (49) -4.16 x 10-18 (26) -4.64 x 10-18 (29) -1.408 x 10-17 (88) -8.16 x 10-18 (51) -1.28 x 10-17 (80) -1.488 x 10-17 (93) -1.456 x 10-17 (91) -1.6 x 10-18 (10) -8.0 x 10-19 (5) -1.12 x 10-18 (7) -1.104 x 10-17 (69) -4.16 x 10-18 (26) -1.6 x 10-19 (1) -3.84 x 10-18 (24) -6.4 x 10-19 (4) -3.2 x 10-19 (2) -1.28 x 10-18 (8) -4.8 x 10-19 (3) -6.24 x 10-18 (39) -9.44 x 10-18 (59) -5.44 x 10-18 (97) -1.104 x 10-17 (69) -1.28 x 10-17 (80) -6.24 x 10-18 (39) -2.72 x 10-18 (17) -1.232 x 10-17 (77) -1.28 x 10-17 (80) The smallest charge on an oil -19 drop was 1.60 X 10 C. Thus he concluded that the charge of an electron was 1.60 X 10-19 coulombs! determined by J.J. Thomson and others 1.6 x 10-19 1 gram CX 8 1.756 X 10 C = 9.11 X -28 10 g Rutherford’s Experiment Simple Rutherford Complex Rutherford Rutherford 1. The atom is mostly empty space! 2. The charge is concentrated in the center. Credited with discovering the nucleus Rutherford’s Model of the Atom Chadwick’s Experiment 1. The charge on the particle was found to be zero. 2. The mass of the particle was found to be 1.67 x 10-24 g, or the same as the p+. Credited with discovering the neutron Chadwick’s Model of the Atom General Features of the Atom Today. •The atom is an electrically neutral, spherical entity composed of a positively charged central nucleus surrounded by one or more negatively charged electrons. •The atomic nucleus consists of protons and neutrons. The Modern Reassessment of the Atomic Theory 1. All matter is composed of atoms. The atom is the smallest body that retains the unique identity of the element. 2. Atoms of one element cannot be converted into atoms of another element in a chemical reaction. Elements can only be converted into other elements in nuclear reactions. 3. All atoms of an element have the same number of p+ and e-, which determines the chemical behavior of the element. Isotopes of an element differ in the number of n0, and thus in mass number. A sample of the element is treated as though its atoms have an average mass. 4. Compounds are formed by the chemical combination of two or more elements in specific ratios. Properties of the Three Key Subatomic Particles Charge Mass Location Name(Symbol) Relative Absolute(C)* Relative(amu)† Absolute(g) in the Atom Proton (p+) 1+ +1.60218x10-19 1.00727 1.67262x10-24 Nucleus Neutron (n0) 0 0 1.00866 1.67493x10-24 Nucleus 1- -1.60218x10-19 Electron (e-) 0.00054858 9.10939x10-28 * The coulomb (C) is the SI unit of charge. † The atomic mass unit (amu) equals 1.66054x10-24 g. Outside Nucleus The AMU • Atomic Mass Unit is a relative unit and does not represent the mass in real masses. • 1 amu = 1/12 C-12 = 1.6606 X 10-24 g. Modern terminology now uses the unit of Daltons (D) in place of AMU’s. Atomic Notation Mass Number (p+ + n0) A X Z Atomic Symbol Z & A are integers Atomic Number (p+ ) Example: Manganese 55 Mn 25 A Z X p+ = 25 e- = 25 n0 = 30 Example: Chlorine 35 Cl 17 A Z X p+ = 17 e- = 17 n0 = 18 Ions • Ions are charged atoms. • In many cases, chemical bonding is the attraction between oppositely charged particles. (Ionic Bonding) How do atoms become charged ions? • By the addition or removal of e-. (Neutrons are neutral and would have no effect; changing protons would change the element. The only particle that is charged and of low enough energy to be removed is the electron.) Example: Lithium Ion 7 Li 3 A Z X +1 p+ = 3 e- = 2 n0 = 4 Example: Selenide Ion 79 -2 Se 34 A Z 59 X p+ = 34 e- = 36 n0 = 45 Isotopes • Atoms of the same element which differ slightly in atomic mass. (A) –Same element - same properties. • Electrons are considered massless and the number of protons (Z) can not change. • Therefore the different masses (A) are caused by different numbers of neutrons. Determining the Number of Subatomic Particles in the Isotopes of an Element PROBLEM: Silicon (Si) has three naturally occurring isotopes: 28Si, 29Si, and 30Si. Determine the number of protons, neutrons, and electrons in each silicon isotope. PLAN: Mass number (A), protons + neutrons, is given for the listed isotopes. Atomic number (Z), number of protons, for each element is given in the periodic table and equal to the number of electrons. Number of neutrons is determined using equation 2.2. SOLUTION: The atomic number of silicon is 14. Therefore 28Si has 14p+, 14e- and 14n0 (28-14) 29Si has 14p+, 14e- and 15n0 (29-14) 30Si has 14p+, 14e- and 16n0 (30-14) How do we know there are different isotopes? • A mass spectrograph is used to deflect ions of an element as they move through a magnetic field. http://www.colby.edu/chemistry/OChem/DEMOS/MassSpec.html Mass Spectrometer and Its Data a A a Heaviest Isotope Lightest Isotope Why are the atomic masses different from the A? • The atomic weights (masses) are the average for all the isotopes. • B has two isotopes … 10B, 11B. 10B= 10.0129 and 11B= 11.0093. Are these the masses (A) on the periodic table? No! Why not? Because 10B is found as 20.00% and 11B @ 80.00% A stream of B ions are shot through a magnetic field. The following is observed. Magnet Ion Gun 10B 11B Ion Detector Screen Calculation • Average Atomic Mass = %1 (X1) + %2(X2) + etc 100% (20.0)10.0129 + (80.0)11.0093 100 = 10.81 Another Type • An element (Sb) has an Atomic Mass of 121.76 amu • There are only two naturally occuring isotopes. • Sb –121 @ 120.904 amu • Sb – 123 @ 122.904 amu • Since percentages are based on 100, assume a 100 particle sample. 121 X = Sb (100 - X) = 123Sb Calculate their percent abundances. 121.76 = 120.904 (x) + 122.904(100-x) 100 121X = 57.000% and 123X = 43.000% • Silicon has three isotopes. Masses are Si-28 @ 27.977, Si-29 @ 28.977 and Si-30 @ 29.974 amu. The % abundance for the middle isotope is 4.70%. Calculate the % abundances for the two remaining isotopes. •Si-28 = 92.2% and Si-30 = 3.1% The Modern Periodic Table. Organization of the Elements 06 Track 6.wma • 1869: Dmitri Mendeleev –Published an organizational scheme for the elements. –He called it the Periodic Table. • Organized his elements by atomic mass in two rows of 7 elements and two rows of 17 elements. • Noticed that the properties of the elements repeated so he put elements that behaved the same in the same vertical column. • Left blank spots on his periodic table and predicted the discovery and properties of several undiscovered elements. __________________________ • Problems were found with the table. • Ar -- K Te -- I Co -- Ni • If Mendeleev’s Table was to be correct, these elements should have been reversed. • Elements could not be reversed because of their properties. • This lead to a new version of the Periodic Table. The Modern Periodic Law • Henry Mosely - The properties of the elements are periodic when they are arranged in increasing order of their Atomic Number or Protons. “Patterns” on the Periodic Table. • Columns: Group or family - elements with similar properties. • Several groups have descriptive names. Alkali Metals - Group IA • React with water to form an alkaline or basic solution. • Forms +1 ions (Cations). “s” sublevel electrons. Base Na(S) + H2O(L) --> NaOH(aq) + H2(g) Alkaline Earth Metals - Group IIA • Elements also react with water to form a base. •Forms +2 ions (Cations). •“s” sublevel electrons. Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g) GroupI_And_II_Rxn.rm Halogens - Group VIIA • “Salt Formers” Forms -1 ions (Anions). “p” sublevel electrons. Na(S) + Cl2(g) --> NaCl(S) PhysicalPropsofHalogens.MOV “Salt” Noble Gases - Group VIIIA • Gases that are highly unreactive because their electron configuration is very stable. • Does not form ions. • “p” sublevel electrons completed. Chalcogens - Group VIA • any of the elements oxygen, sulfur, selenium, and tellurium. Forms -2 ions (Anions). “p” sublevel electrons. 2Ca(s) + O2 (g) → 2CaO(s) The Periodic Table of the Elements H Li Be NaMg K Ca Sc Ti R Sr Y Zr b Cs Ba La Hf Fr Ra Ac Rf B C N Al Si P V Cr Mn Fe Co Ni Cu Zn Ga Ge As NbMo Tc R Rh Pd Ag Cd In Sn Sb Ta W Re uOs Ir Pt Au Hg Tl Pb Bi Du Sg Bo HaMe O S Se Te Po He F Ne Cl Ar Br Kr I Xe At Rn Ce Pr Nd PmSm Eu Gd Tb Dy Ho Er TmYb Lu Th Pa U Np PuAmCmBk Cf Es FmMd No Lr The Alkali Metals The Halogens The Alkaline Earth Metals The Noble Gases The Periodic Table of the Elements H Li Be NaMg K Ca Sc Ti Rb Sr Y Zr Cs Ba La Hf Fr Ra Ac Rf B C N Al Si P V CrMn Fe Co Ni Cu Zn Ga Ge As Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Ta W Re Os Ir Pt Au Hg Tl Pb Bi Du Sg Bo Ha Me O S Se Te Po F Cl Br I At Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu Th Pa U Np PuAmCm Bk Cf Es FmMd No Lr Boron family Nitrogen family Carbon Family Chalcogens He Ne Ar Kr Xe Rn Transition Metals – B •Generally form +2 cations •Many are multivalent •Formed by d-sublevel eScandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc 21 22 23 24 25 26 27 28 29 30 Sc Ti V Cr Mn Fe Co Ni Cu Zn 44.956 4s23d1 47.88 4s2 3d2 50.942 4s2 3d3 51.996 4s1 3d5 54.938 4s2 3d5 55.847 4s23d6 58.933 4s23d7 58.633 4s2 3d8 63.546 4s1 3d10 65.39 4s23d10 Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium 39 40 41 42 43 44 45 46 47 48 Y Zr Nb Mo Tc Ru Rh Pd Ag Cd 88.906 5s24d1 91.224 5s2 4d2 92.906 5s1 4d4 95.94 5s1 4d5 (98) 5s2 4d5 101.07 5s14d7 102.906 5s14d8 106.42 5s1 4d9 107.868 5s1 4d10 112.411 5s24d10 Lanthanum Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury 57 72 73 74 75 76 77 78 79 80 La Hf Ta W Re Os Ir Pt Au Hg 138.906 6s25d1 178.49 6s2 5d2 180.948 6s2 5d3 183.85 6s2 5d4 186.207 6s2 5d5 190.2 6s25d6 192.22 6s25d7 195.08 6s1 5d9 196.967 6s1 5d10 200.59 6s25d10 Actinium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Ununbium 89 104 105 106 107 108 109 110 111 Ac Rf Db Sg Bh Hs Mt Ds Rg 112 Uub 227.028 7s26d1 (261) 7s2 6d2 (262) 7s2 6d3 (263) 7s2 6d4 (262) 7s2 6d5 (265) 7s26d6 (266) 7s26d7 (281) 7s1 6d9 (280) 7s1 6d10 (285) 7s26d10 Inner Transition Metals •Lanthanide series •Actinide series •Formed by f-sublevel eCerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium 58 59 60 61 62 63 64 65 66 67 68 69 70 71 Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 140.115 6s24f15d1 140.908 6s24f3 144.24 6s2 4f4 (145) 6s2 4f5 150.36 6s2 4f6 151.965 6s2 4f7 157.25 6s2 4f75d1 158.925 6s2 4f9 162.50 6s2 4f10 164.930 6s2 4f11 167.26 6s2 4f12 168.934 6s2 4f13 173.04 6s24f14 174.967 6s24f145d1 Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium 90 91 92 93 94 95 96 97 98 99 100 101 102 103 Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr 232.038 7s26d2 231.036 7s25f26d1 238.029 7s2 5f36d1 237.048 7s2 5f46d1 (244) 7s2 5f6 (243) 7s2 5f7 (247) 7s2 5f76d1 (247) 7s2 5f9 (251) 7s2 5f10 (252) 7s2 5f11 (257) 7s2 5f12 (258) 7s2 5f13 (259) 7s25f14 (260) 7s25f146d1 Lutetium The Periodic Table of the Elements H Li Be B C N Na Mg Al Si P K Ca Sc Ti V CrMn Fe Co Ni Cu Zn Ga Ge As Rb Sr Y Zr NbMo Tc Ru Rh PdAg Cd In Sn Sb Cs Ba La Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Fr Ra Ac Rf Du Sg Bo Ha Me O S Se Te Po He F Ne Cl Ar Br Kr I Xe At Rn Transition Metals Ce Pr Nd PmSm Eu Gd Tb DyHo Er TmYb Lu Th Pa U Np PuAmCmBk Cf Es FmMd No Lr Lanthanides: Rare Earth Elements Actinides Period • Row on the periodic table. • There are 7 periods on the periodic table. –One period for each “layer” or energy level of electrons. The Periodic Table of the Elements H Li Be Na Mg K Ca Sc Rb Sr Y Cs Ba La Fr Ra Ac Ti Zr Hf Rf B C N Al Si P V Cr Mn Fe Co Ni Cu Zn Ga Ge As NbMo Tc Ru Rh Pd Ag Cd In Sn Sb Ta W Re Os Ir Pt Au Hg Tl Pb Bi Du Sg Bo Ha Me O S Se Te Po F Cl Br I At He Ne Ar Kr Xe Rn Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er TmYb Lu Th Pa U Np PuAmCm Bk Cf Es FmMdNo Lr Metals Non Metals Semi - metals Metalloids A Biological Periodic Table. Bonding • Most elements occur combined … very few are found in elemental form. • Noble gases are found uncombined. • Au, Ag,Cu, and Pt do not react easily and are often found uncombined. •Some elements combine with themselves … Diatomics S8 and P4 Types of Chemical Formulas A chemical formula is comprised of element symbols and numerical subscripts that show the type and number of each atom present in the smallest unit of the substance. An empirical formula indicates the relative number of atoms of each element in the compound. It is the simplest type of formula. The empirical formula for hydrogen peroxide is HO. A molecular formula shows the actual number of atoms of each element in a molecule of the compound. The molecular formula for hydrogen peroxide is H2O2. A structural formula shows the number of atoms and the bonds between them, that is, the relative placement and connections of atoms in the molecule. The structural formula for hydrogen peroxide is H-O-O-H. H2O Ammonia - NH3 Methane - CH4 Salt Crystal Cl Na Single Bond Double Bond Which molecule would be hardest to pull apart? Triple Bond Types of Bonds Ionic: attraction of oppositely charged ions resulting from the transfer of e-’s Reaction is typically between a metal and a nonmetal Covalent: attraction based on localized epair sharing between two atoms Reaction is typically between two nonmetals Metallic: attraction based on delocalized e-’s and metal ions Reaction is typically between two metals Ionic Bond • Electrons transferred—usually between metals and nonmetals. • Metals lose e-’s and form + ions (cations) • Nonmetals gain e-’s and form – ions (anions) • Metals and nonmetals will lose or gain e-’s to form a noble gas configuration • Na → Na+1 + e- (isoelectronic with Ne) • Cl + e- → Cl-1 (isoelectronic with Ar) •Sodium is isoelectronic with Ne •Iso = same •Chlorine is isoelectronic with Ar •Almost exclusively in ionic bonding, the energy of the newly formed compound is MUCH lower than the original elements. The Relationship Between Ions Formed and the Nearest Noble Gas. The Periodic Table of the Elements Most Probable Oxidation State +1 +4 0 +3 -4 -3 -2 -1 He H +2 Li Be Na Mg +3 K Ca Sc Rb Sr Y Cs Ba La B C N O F Ne +1 +2 Al Si P S Cl Ar +4 +5 Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge A Se Br Kr Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn sSb Te I Xe Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn Fr Ra Ac Rf Du Sg Bo Ha Me +3 +3 Ce Pr Nd PmSm Eu Gd Tb Dy Ho Er TmYb Lu Th Pa U Np Pu AmCmBk Cf Es FmMd No Lr Common Monoatomic Ions Cations Charge Formula 1+ 2+ 3+ H+ Li+ Na+ K+ Cs+ Ag+ Mg2+ Ca2+ Sr2+ Ba2+ Zn2+ Cd2+ Al3+ Name hydrogen lithium sodium potassium cesium silver magnesium calcium strontium barium zinc cadmium aluminum Anions Charge Formula 1- HFCl Br I- Name hydride fluoride chloride bromide iodide 2- O2S2 - oxide sulfide 3- N 3- nitride Listed by charge; those in boldface are most common Metals With Several Oxidation States Element Copper Cobalt Iron Manganese Tin Lead Ion Formula Systematic Name Common Name Cu+1 copper(I) cuprous Cu+2 copper(II) cupric Co+2 cobalt(II) Co+3 cobalt (III) Fe+2 iron(II) ferrous Fe+3 iron(III) ferric Mn+2 manganese(II) Mn+3 manganese(III) Sn+2 tin(II) stannous Sn+4 tin(IV) stannic Pb+2 lead(II) Pb+4 lead(IV) Coulomb’s Law The magnitude of the electrostatic force between two point electric charges is directly proportional to the product of the magnitudes of each charge and inversely proportional to the square of the distance between the charges. k ⋅ Q1 ⋅ Q 2 F= 2 d charge1× charge2 Forceα 2 Distance charge1× charge2 Energyα Distance Covalent Bond • Electrons shared—usually between two nonmetals. • e-’s are attracted to both nuclei • More common than ionic bonds • e-’s are shared in pairs • Localized—does not extend to other molecules Covalent Bonding Formation of a Covalent Bond between Two Hydrogen Atoms Covalently bonded substances exist as individual entities. Ionic substances are continuous arrays of oppositely charged ions. No individual “molecules” exist in ionic compounds. Metallic Bond • Electrons shared among all of the atoms in the substance • e-’s form an electron sea The reason metals deform. metal is deformed Structure of Metals Polyatomic Ions • A charged ion whose atoms are covalently bonded groups • Behave as a monatomic –ion (anion) when joined with a metal. • Most end in the suffixes –ate or –ite • Oxygen is the prevalent atom Polyatomic Ions •Many important ions: Carbonate CO3-2 Nitrate NO3Nitrite NO2+ Ammonium NH4 … A Polyatomic Ion Some Common Polyatomic Ions Formula Name Formula Cations NH4+ H3O+ Anions Cont. Ammonium Hydronium Anions CH3COOCNOHClOClO2ClO3ClO4NO2NO3- Name Acetate Cyanide Hydroxide Hypochlorite Chlorite Chlorate Perchlorate Nitrite Nitrate MnO4CO3-2 HCO3CrO4-2 Cr2O7-2 O2-2 PO4-3 HPO4-2 H2PO4SO3-2 SO4-2 HSO4- Permanganate Carbonate Hydrogen Carbonate Chromate Dichromate Peroxide Phosphate Hydrogen Phosphate Dihydrogen Phosphate Sulfite Sulfate Hydrogen Sulfate Writing Ionic Formulas Write the symbol of the positive metal ion or polyatomic group first and include the oxidation number. • Write the symbol of the negative non-metal ion or polyatomic group second with the oxidation number. Using the LCD, write subscripts to balance the oxidation numbers and remember: The algebraic sum of all charges must be ZERO. The formulas must be in the simplest whole number ratio. The Criss -Cross Method • Caution: This method does not always produce the SIMPLEST WHOLE NUMBER ratio. Aluminum and sulfur: Al+3 S-2 Al2 S3 Hence Al2+3S3-2 EXAMPLES • Aluminum and chlorine Al+3 Cl-1 Al+3Cl3-1 • Magnesium and phosphorus +2 -3 Mg P +2 -3 Mg3 P2 Metal Ions With Multiple Oxidation Numbers. Numbers Lead (IV) with oxygen: +4 -2 Pb O +4 -2 Pb2 O4 However the answer is: Pb+4O2-2 … The simplest whole number ratio. Lead (II) with oxygen: +2 -2 +2 -2 Pb O gives Pb2 O2 ...or Pb+2O-2 after it is reduced. Polyatomic Ions • Groups of atoms that act like a single ion. • If more than one polyatomic group is needed, the group must be wrapped in parathensis. Try these: • calcium acetate Ca+2(C2H3O2-1)2 • aluminum hydroxide +3 -1 Al (OH )3 • aluminum sulfate Al2+3(SO4-2)3 • ammonium phosphate +1 -3 (NH4 )3PO4 Naming Ionic Compounds • IONIC COMPOUNDS = METAL + NON METAL 1. Write the name of the positive metal first 2. Write the ROOT of the negative non metal then drop the last syllable, sometimes two, and add IDE . • Try… Try… Al2O3 –Aluminum Aluminum oxide oxide • Try… NaCl –Sodium Sodium Chloride Chloride • Try … Ba3P2 –Barium Barium Phosphide Phosphide Naming ionic compounds with metals that have more than one oxidation number. • Use the “ROMAN “ method. • Iron has two oxidation numbers, Fe+2 and Fe+3. • … then Fe+2Cl2-1 is named Iron(II Iron(II) II) chloride 1 +3 • … and Fe Cl3 is named Iron(III Iron(III)chloride III)chloride • … and PbCl4 is named Lead(IV)chloride Naming Compounds With Polyatomic Ions • 1. Write the name of the positive metal ion or polyatomic ion. • 2. Write the name of the negative polyatomic ion. Examples • (NH4+1)3PO4-3 • Ammonium Phosphate • Fe2(SO4)3 • Fe2+3(SO4-2)3 • Iron(III) Sulfate Hydrated Compounds • Some ionic salts often trap water molecules within their structures. • Mechanically held water in a specific ratio by weight is a “hydrated” compound. • CuSO4 . 5 H2O; 36.0% water. •BaCl2 . 2 H2O; 14.7% water. How are these hydrates named? Numerical Prefixes for Hydrates and Binary Covalent Compounds 1 = mono 5 = penta 9 = nona 2 = di 6 = hexa 10 = deca 3 = tri 7 = hepta 4 = tetra 8 = octa Hence: CuSO4 . 5 H2O is: Copper (II) sulfate pentahydrate BaCl2 . 2 H2O is: Barium chloride dihydrate Naming Covalent Compounds. “Greek “ Method For the 1st element … name that element using a prefix indicating the number of atoms present. For the 2nd element … name the second element using a prefix indicating the number of atoms present and add root and -ide ending. P2O5 is named: diphosphorus pentaoxide N2O3 is named: dinitrogen trioxide P2S is named: diphosphorus monosulfide SiO2 is named: silicon dioxide … when there is one atom of the first element the prefix mono is not used. Then CO & CO2 are …. carbon monoxide and carbon dioxide The Following May Have Mistakes … Correct Them • SO2 is sodium dioxide –sulfur dioxide • Cl2O is chlorine dioxide –dichlorine monoxide • SF10 is disulfur decafluoride –sulfur decafluoride •MgCl2 is magnesium dichloride –magnesium chloride (ionic) Something is wrong with the second part of each statement. Provide the correct name or formula. Ba(C2H3O2)2 = barium diacetate Barium is always a +2 ion and acetate is -1. The “di-” is unnecessary. Sodium sulfide = (Na)2SO3 An ion of a single element does not need parentheses. Sulfide is S2-, not SO32-. The correct formula is Na2S. Iron(II) sulfate = Fe2(SO4)3 Since sulfate has a 2- charge, only 1 Fe2+ is needed. The formula should be FeSO4. Cesium carbonate = Cs2(CO3) The parentheses are unnecessary. The correct formula is Cs2CO3. Basic Oxoanions • Combos With Two Possibilities: The most common are SO4-2 and SO3-2 and are named sulfate and sulfite. The ion with the greater # of oxygen atoms has the ate ending. The lesser # of oxygen atoms has the -ite ending! Two other common oxoanions: NO3- and NO2- Naming2Seriesof2Oxyanions.swf Nitrate and nitrite More Basic Oxoanions Ions • Combos With Four Possibilities: – This would be any halogen (group 7A) with 4,3,2,1 oxygen atoms respectively. BrO4-1 = perbromate BrO3-1 = bromate BrO2-1 = Bromite BrO-1 = hypobromite NamingASeriesof4Oxyanions.swf NaClO3 = Sodium chlorate KNO2 = Potassium Nitrite BaSO3 = Barium Sulfite LiClO4 = Lithium Perchlorate Ca(BrO)2 = Calcium Hypobromite NaClO3 Sodium Chlorate Al(IO2)3 Aluminum Iodite RbClO2 Rubidium Chlorite KBrO3 Potassium Bromate CsClO Cesium Hypochlorite LiIO4 Lithium Periodate Calcium Nitrate = Ca(NO3)2 Ammonium Sulfite = (NH4)2SO3 Lithium Nitrite = LiNO2 Potassium Hypochlorite = KClO Lithium Perbromate = LiBrO4 Magnesium Hypoiodite = Mg(IO)2 Naming Acids • Two types: • 1. Binary acids -two elements •Forms when certain gaseous compounds dissolve in water. Naming Binary Acids Hydro + root + ic HCl = hydro + chlor + ic hydrochloric acid HBr = hydrobromic acid HF = hydrofluoric acid H2S = hydrosulfuric acid 2. Ternary (Oxoacid )acids- those with three or more elements. • Naming ternary: –use the polyatomic name but change the: -ate to ic -ite to ous Or … Anion “-ate” suffix becomes an “-ic” suffix in the acid. Anion “ -ite” suffix becomes an “ -ous” suffix in the acid. The oxoanion prefixes “hypo-” and “per-” are retained. Examples: HClO3 is named after the chlorate ion, and thus becomes chloric acid. HBrO4 is perbromic acid and is named after the perbromate ion. HFO = Hypoflourous acid H2SO4 = Sulfuric acid Nitric acid = HNO3 Chlorous acid = HClO2 Determining Names and Formulas of Anions and Acids Problem: Name the following anions and give the names and a) I - formulas of the acid solutions derived from them: b) BrO2c) SO3-2 d) NO3e) CN - a) The anion is Iodide; and the acid is Hydroiodic acid, HI b) The anion is Bromite; and the acid is Bromous acid, HBrO2 c) The anion is Sulfite; and the acid is Sulfurous acid, H2SO3 d) The anion is Nitrate; and the acid is Nitric acid, HNO3 e) The anion is Cyanide; and the acid is Hydrocyanic acid, HCN Naming Alkanes Alkanes are hydrocarbons that are called “saturated” hydrocarbons, they contain only single bonds, no multiple bonds! Alkanes have the general formula: C n H 2n+2 Alkanes are found in three distinct groups: a) Straight chain hydrocarbons b) Branched chain hydrocarbons c) Cyclic hydrocarbons Each carbon atom has four bonds to others atoms ! The names for alkanes all end in – ane. The First 10 Straight-Chain Alkanes Name Formula Structural Formulas Methane Ethane CH4 C2H6 Propane Butane C3H8 C4H10 CH3-CH2-CH3 Pentane Hexane C5H12 C6H14 CH3-(CH2)3-CH3 H H C H H H H H C C H H H CH3-(CH2)2-CH3 CH3-(CH2)4-CH3 Heptane Octane C7H16 C8H18 Nonane Decane C9H20 C10H22 CH3-(CH2)5-CH3 CH3-(CH2)6-CH3 CH3-(CH2)7-CH3 CH3-(CH2)8-CH3 Branched Chain Hydrocarbons Cyclic Hydrocarbons Calculating the Molecular Mass of a Compound PROBLEM: Using the data in the periodic table, calculate the molecular (or formula) mass of: (a) tetraphosphorous trisulfide PLAN: (b) ammonium nitrate Write the formula and then multiply the number of atoms by the respective atomic masses. Add the masses for the compound. SOLUTION: (a) P4S3 molecular mass = (4 x atomic mass of P) + (3 x atomic mass of S) = (4 x 30.97 amu) + (3 x 32.07 amu) = 220.09 amu (b) NH4NO3 molecular mass = (2 x atomic mass of N) + (4 x atomic mass of H) + (3 x atomic mass of O) = (2 x 14.01 amu) + (4 x 1.008 amu) + (3 x 16.00 amu) = 80.05 amu Naming Compounds from Their Depictions PROBLEM: Each circle contains a representation of a binary compound. Determine its formula, name, and molecular (formula) mass. PLAN: Each compound contains only two elements. Find simplest whole number ratio of one atom to the other to determine formula, name, and mass. SOLUTION: (a) There is 1 sodium (brown) for every fluorine (green), so the formula is NaF. formula mass = (1x atomic mass of Na) + = (1x atomic mass of F) 22.99 amu + 19.00 amu = 41.99 amu (b) There are 3 fluorines (green) for every nitrogen (blue), so the formula is NF3. molecular mass = = (3x atomic mass of F) + (1x atomic mass of N) (3x 19.00 amu) + 14.01 amu = 71.01 amu The Distinction Between Mixtures and Compounds. S2- Fe2+ Physically mixed therefore can be separated by physical means; in this case by a magnet. Allowed to react chemically therefore cannot be separated by physical means. Mixtures Heterogeneous mixture: has one or more visible boundaries between the components. Homogeneous mixture: has no visible boundaries because the components are mixed as individual atoms, ions, and molecules. Solutions: A homogeneous mixture is also called a solution. Solutions in water are called aqueous solutions, and are very important in chemistry. Although we normally think of solutions as liquids, they can exist in all three physical states. Solutions • A solution that mixes completely is miscible. (alcohol and water, gases) • Solutions that mix only within certain limits are expressed as a solubility. (sugar in water, carbon in iron) • Substance that is dissolved is the solute. • Substance that does the dissolving is the solvent. 170 Nine Types of Solutions 171 Solute Gas Gas Gas Liquid Liquid Liquid Solvent Gas Liquid Solid Gas Liquid Solid Example Air Carbonated water Catalytic converter Water in air Alcohol in water Mercury in gold Solid Gas Smoke Solid Solid Liquid Solid Salt in water Copper in zinc Tools of the Laboratory Basic Separation Techniques Filtration: Separates components of a mixture based upon differences in particle size. Normally separating a precipitate from a solution, or particles from an air stream. Filtration Crystallization: Separation is based upon differences in solubility of components in a mixture. Crystallization Distillation: separation is based upon differences in volatility. Extraction: Separation is based upon differences in solubility in different solvents (major material). Centrifugation Separates components of a mixture based upon differences in density. 176 Precipitation Separates components of a mixture based upon the formation of a solid. 177 Chromatography: Separation is based upon differences in solubility in a solvent versus a stationary phase. Procedure for Column Chromatography Separation by Gas - Liquid Chromatography PaperChromatographyofInk.MOV The End