Chemistry Notes PDF

WHAT IS CHEMISTRY? WHAT DO CHEMIST DO? HOW MIGHT KNOWLEDGE OF CHEMISTRY HELP YOU? What is universe made of? … ..  …. Chemistry is a science that studies .....  ….  …....

WHAT IS CHEMISTRY? WHAT DO CHEMIST DO? HOW MIGHT KNOWLEDGE OF CHEMISTRY HELP YOU? What is universe made of? … ..  …. Chemistry is a science that studies .....  ….  …. CH 2. All rights reserved. Dr. Gulacar, Department of Chemistry, University of California, Davis. Reproducing these notes or posting them on any online source is strongly prohibited. CONTENTS CONTENTS 2-1 Early Chemical Discoveries and the Atomic Theory 2-2 Electrons and Other Discoveries in Atomic Physics 2-3 The Nuclear Atom 2-4 Chemical Elements 2-5 Atomic Mass 2-6 Introduction to the Periodic Table 2-7 The Concept of the Mole and the Avogadro Constant 2-8 Using the Mole Concept in Calculations 1  Fundamental Chemical Laws Dalton’s Atomic Theory Fundamental Chemical Laws: Do atoms exist? The Law of Conservation of Mass (Antoine Lavoisier, 1774) The Law of Definite Proportions or the Law of Constant Composition (Joseph Proust, 1799) Law of Conservation of Mass If you mix silver nitrate (colorless) and potassium chromate, what would happen? 2  Learning Check: Magnesium burns in oxygen to form magnesium oxide. If 16.88 g of Mg are consumed and 28.00 g of MgO are produced, what mass of oxygen was consumed? A. 5.560 B. 11.12 C. 16.88 D. 44.88 A Conceptual Example Jan Baptista van Helmont (1579–1644) first measured the mass of a young willow tree and, separately, the mass of a bucket of soil and then planted the tree in the bucket. After five years, he found that the tree had gained 75 kg in mass even though the soil had lost only 0.057 kg. He had added only water to the bucket, and so he concluded that all the mass gained by the tree had come from the water. Explain and criticize his conclusion. Law of Definite Proportions In a given chemical compound, the elements are always combined in the same proportions by mass. If we decompose samples of water (a compound) into the elements oxygen and hydrogen, we always find that the ratio of oxygen to hydrogen, by mass, is 8 to 1. In other words, the mass of oxygen obtained is always eight times the mass of hydrogen. ….g hydrogen …g hydrogen 9g water 18g water …g oxygen …g oxygen 3  Learning Check In a sample of MgO, there are 16.89 g Mg and 11.11 g O. What mass of O would there be in a sample that contains 2.00 g of Mg? A. 1.11 B. 1.00 C. 1.32 D. 1.316 Dalton’s Atomic Theory Proposed in 1803 to explain the law of conservation of mass and law of definite proportions. 1. Matter is composed of atoms: tiny, indivisible particles. 2. All atoms of a given element are the same and differ from all other elements 3. Compounds are formed when atoms of different elements unite in fixed proportions. 4. A chemical reaction involves rearrangement of atoms. No atoms are created, destroyed, or broken apart. 2 Na(s) + 2 H2O(l)  H2(g) + 2 NaOH(aq) Dalton’s Theory Led to a Prediction  Law of Multiple Proportions When two or more different compounds of the same two elements are compared, the masses of one element that combine with a fixed mass of the second element are in the ratio of small whole numbers.  Example: Carbon with Oxygen Carbon monoxide  12.0 g C + 16.0 g O Carbon dioxide  12.0 g C + 32.0 g O  What is the ratio of O/C in each of these two compounds? 4  Law of Multiple Proportions (cont’d)  Four different oxides of nitrogen can be formed by combining 28 g of nitrogen with:  16 g oxygen, forming Compound I  48 g oxygen, forming Compound II  64 g oxygen, forming Compound III  80 g oxygen, forming Compound IV Early Experiments to Characterize the Atom Electricity and magnetism were used in the experiments that led to the current theory of atomic structure. 5 J. J. Thomson used the deflection of cathode rays and the magnetic field strength together, to find the cathode ray particle’s mass-to-charge ratio: me /e = –5.686 × 10–12 kg/C The ratio me/e for cathode rays is about 2000 times smaller than the smallest previously known me/e (for hydrogen ions).  George Stoney: names the cathode-ray particle the electron.  Robert Millikan: determines a value for the electron’s charge: e = –1.602 × 10–19 C 6  Thomson determined the mass-to-charge ratio; Millikan found the charge; we can now find the mass of an electron: me = 9.109 × 10–31 kg/electron  This is almost 2000 times less than the mass of a hydrogen atom (1.79 × 10–27 kg)  Thomson proposed an atom with a positively charged sphere containing equally spaced electrons inside.  He applied this model to atoms with up to 100 electrons.  Research into cathode rays showed that a cathode-ray tube also produced positive particles. Unlike cathode rays, these positive particles were ions. The metal of the cathode: M  e– + M+ 7 Mass spectrum of an element shows the abundance of its isotopes. What are the three most abundant isotopes of mercury? Mass spectrum of a compound can give information about the structure of the compound. The Modern View of Atomic Structure: An Introduction 8  Atom Structure Neutron – carries no charge, found in the nucleus, a bit heavier than a Electron – carries a negative -1 proton, about 1800 times heavier than charge, found outside the an electron ( mass=1.0087 amu) nucleus, about 1/1800 the (Chadwick discovered in 1932) mass of a proton (mass=5.486 x 10-4 amu) +1 +1 Proton – carries a positive charge, found in the nucleus Electron Cloud -1 (mass=1.0073amu) 0.1-0.5 Å Protons and (Angstroms) Neutrons 1 Å = 1 x 10-10 m 10-4 Å in diameter Charge is relative 1 amu = 1.66054 x 10-24 g Atomic Notation A q Z Sy e  Atomic number, Z = number of protons  An element is a substance whose atoms all contain the identical number of protons  Mass number, A = (number of protons) + (number of neutrons)  Charge, q =(number of protons) - (number of electrons) Examples Find the  number of protons ..  number of neutrons 80 Br- ..  number of electrons .. 35  Atomic number ..  Mass Number ..  Charge .. 9  Ions Anion: A negative ion (original atom gains electron) Cl, SO42 Cation: A positive ion (original atom loses electron) Mg2+, NH4+ Learning Check  Which of the following is false regarding Mg2+? a) Mg2+ is a cation b) Mg2+ has 14 electrons c) Mg2+ has 12 protons d) Mg2+ has 10 electrons Molecules and Ions 10  CHEMICAL COMPOUNDS  Molecular and ionic compounds (based on type of bonding): ◦ Molecular compounds involve shared electrons (among non-metals) and consist of electrically neutral, discrete particles (molecules) ◦ Ionic compounds involve electron transfer (between metals and non- metals) and charged particles (ions) Chemical Bonds  Covalent—electrons shared between atoms Octet (8) or Doublet (2) O H2O H H HYDROGEN AND OXYGEN Chemical Bonds Ionic—complete transfer of electrons from one atom to another - + 11+ 17+ 11+ and 10- = 1+ 17+ and 18- = 1- Na+ Cl- SODIUM AND CHLORINE 11  TYPES OF FORMULAS  Empirical formulas (C2H5)  consist of the lowest whole number ratio of atoms in a compound reduced by the largest common denominator  Ionic compounds have empirical formulas NaCl MgCl2 Al2(SO4)3 K2CO3  Molecular formulas (C4H10)  show the true number of atoms of each element in the formula of a compound  Molecular compounds can have either empirical and molecular formulas LEARNING CHECK Which of the following formulas is an empirical formula? A. C3H6O B. C2H6O2 C. C2H6 D. C6H12O6 E. C5H5 STRUCTURAL FORMULAS  Show how atoms are attached together and the order in which atoms are bonded.  Condensed structural formulas also show the structures but with less detail (e.g.,CH3COOH)  They can represent how a group of atom is attached to another atom: CH3CH(CH3)CH3 or CH(CH3)3  Line-angle (line structural) formulas represent organic compounds without showing C and H explicitly.  A carbon atom exist where a line ends or meets another line  # of H atoms needed to complete each carbon atom’s four bonds are assumed to be present 12  An Introduction to the Periodic Table PERIODIC TABLE  The Periodic Table summarizes chemical and physical properties of the elements  The first Periodic Tables were arranged by increasing atomic mass  The Modern Periodic table is arranged by increasing atomic number:  Elements are arranged in  numbered rows called periods  vertical columns called groups or families (group labels vary) Chalcogens Alkali Metals Noble Gases The Periodic table Alkaline Earths Main Group Halogens Transition Metals Main Group Lanthanides and Actinides 13  VALENCE ELECTRONS 8A 1A (ELECTRONS IN THE OUTERMOST ORBIT =GROUP#) 2A 3A 4A 5A 6A 7A  effect the way an atom bonds.  determine many properties of the element.  this is why elements within a group usually have similar properties. Elements can be metals, non- metals, or metalloids PROPERTIES OF METALS METALS  Reflect light (have metallic luster)  Can be hammered or rolled into thin sheets (are malleable) and can be drawn into wire (are ductile)  Are solids at room temperature (except Hg)  Conduct electricity and heat (How?)  Sea of electrons NONMETALS  Lack the properties of metals  Tend to pulverize when struck with a hammer  Non-conductors of electricity and heat  Many are gases, a few solids, and one liquid (Br) METALLOIDS  Have properties between metals and nonmetals  Found along the diagonal: B  At 14  Naming Simple Compounds Writing Ionic Compound Formulas Write the formulas for cation Aluminum sulfide and anion Check to see if the charges are balanced Aluminum phosphate Ionic compounds must have a neutral charge overall (The total positive charge must equal the total Ammonium sulfate negative charge) If not, balance them with subscripts. Use parentheses if you need more than one polyatomic ion THE NAMES AND SYMBOLS OF THE MOST COMMON ELEMENTS Element Symbol Element Symbol Aluminum Al Lithium Li Antimony (stibium)* Sb Magnesium Mg Argon Ar Manganese Mn Arsenic As Mercury (hydrargyrum) Hg Barium Ba Neon Ne Bismuth Bi Nickel Ni Boron B Nitrogen N Bromine Br Oxygen O Cadmium Cd Phosphorus P Calcium Ca Platinum Pt Carbon C Potassium (kalium) K Chlorine Cl Radium Ra Chromium Cr Silicon Si Cobalt Co Silver (argentum) Ag Copper Cu Sodium (natrum) Na Fluorine F Strontium Sr Gold (aurum) Au Sulfur S Helium He Tin (stannum) Sn Hydrogen H Titanium Ti Iodine I Tungsten (wolfram) W Iron (ferrum) Fe Uranium U Lead (plumbum) Pb Zinc Zn *Where appropriate, the original name is shown in parentheses so that you can see the sources of some of the symbols. 15  PERIODIC TABLE NAMING SIMPLE CHEMICAL COMPOUNDS Ionic (metal and nonmetal) Covalent (2 nonmetals) Metal Nonmetal First Second nonmetal nonmetal Forms Forms Single Polyatomic only one more than Negative Ion positive one positive Ion ion ion Use the Use element Use the name Use the Before Use a prefix name of name followed of the name of element name before element by a Roman element, but polyatomic use a prefix element name numeral to end with ide ion (ate or to match and end show the charge Ite) subscript with ide ACIDS 16  If the metal forms only one positive ion, the cation name is the English name for the metal  Example: Ca2+ = calcium ion=calcium Na+ = sodium ion = sodium Al3+ = aluminum ion = aluminum  Stem of the name of the element + suffix –ide Chemical Name as Name as Symbol Stem First Element Second Element O ox- oxygen oxide S sulf- sulfur sulfide N nitr- nitrogen nitride P phosph- phosphorus phosphide F fluor- fluorine fluoride Cl chlor- chlorine chloride Br brom- bromine bromide I iod- iodine iodide Single-charge cations Multiple-charge cations Elemental anions 1+ H He 1 1 2+ 3+ 3- 2- 1- 2 Li Be B C N O F Ne 2 3 4 5 6 7 8 9 10 Na Mg Al Si P S Cl Ar 3 11 12 1+ 2+ 13 14 15 16 17 18 K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr 4 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I Xe 5 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 Cs Ba Hf Ta W Re Os Ir Pt Au Hg Tl Pb Bi Po At Rn 6  55 56 72 73 74 75 76 77 78 79 80 81 82 83 84 85 86 Fr Ra Rf Db Sg Bh Hs Mt 7  87 88 104 105 106 107 108 109 La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 Ac Th Pa U Np Pu Am Cm Bk Cf Es Fm Md No Lr 89 90 91 92 93 94 95 96 97 98 99 100 101 102 103 17  Transition Metals Chromium Cr2+, Cr3+  Ions formed by Manganese Mn2+, Mn3+ transition metals Iron Fe2+, Fe3+ (Group IIIB – VIIIB) Cobalt Co2+, Co3+ and post-transition Mercury Hg22+, Hg2+ metals: Copper Cu+, Cu2+ Gold Au+, Au3+ Post-transition Metals Tin Sn2+, Sn4+ Lead Pb2+, Pb4+ STOCK SYSTEM : Naming cations with multiple charges A. To name given the formula: 1. Figure out charge on cation. 2. Write name of cation. 3. Write Roman numerals in ( ) to show cation’s charge. 4. Complete the name of the compound 4. Find the charge of cation and name the formula. FeO Fe2+ O2– iron (II) oxide Fe2O3 2 Fe3+ 3 O2– iron (III) oxide Learning Check A. CuCl3 B. Cu3Cl C. CoCl3 D. Co3Cl 1. Write symbols for the two types of ions. 2. Figure out the charges 3. Balance charges to write formula. 18  NAMES OF THE POLYATOMIC ANIONS DO NOT CHANGE OR APPEND ANYTHING Ion Name Ion Name NH4+ ammonium ion CO32- carbonate ion OH- hydroxide ion H3O+ hydronium ion NO2- nitrite ion SO32- sulfite ion NO3- nitrate ion SO42- sulfate ion ClO2- chlorite ion CrO42- chromate ion ClO3- chlorate ion Cr2O72- dichromate ion PO43- phosphate ion BrO41- BrO31- BrO21- BrO1- Perbromate ion Bromate ion Bromite ion Hypobromite ion CO42- CO32- CO22- CO2- Carbonate ion ClO4 1- ClO3 1- ClO2 1- ClO1- Chlorate ion IO41- IO31- IO21- IO1- Iodate ion NO41- NO31- NO21- NO1- Nitrate ion PO53- PO43- PO33- PO23- Phosphate ion SO52- SO42- SO32- SO22- Sulfate ion 1 more oxygen “normal” 1 less oxygen 2 less oxygen 19 Naming Covalent (Molecular) Compounds number of atoms prefix P2O5 = ……. 1 mono CO2 = ………… 2 di 3 tri CO = ……………… 4 tetra N2O = ……………….. 5 penta 6 hexa 7 hepta 8 octa 9 nona 10 deca  Rule  The Latin number of the 1st element + name of the first element + the Latin number of the 2nd element + the name of the second element +suffix (-ide)  Only use mono on second element - LEARNING CHECK Which is the correct formula for nitrogen triiodide? A. N3I B. NI3 C. NIO3 D. N(IO3)3 E. none of the above 20  NAMING ACIDS  Acids produce H+ when dissolved in water.  Halic Acids: “Hydro(nonmetal)ic acid”  HCl = hydrochloric acid  H2S(aq) = hydrosulfuric acid  HI(aq) = hydroiodic acid  HBr(aq) = hydrobromic acid  HF(aq) = hydrofluoric  Oxyacids:  -ate: (polyatom)ic acid  H2SO4, based on sulfate: ……… acid  -ite: (polyatom)ous acid  H2SO3, based on sulfite, ……. acid  prefixes are added if present in the ion  HClO4, based on perchlorate, ………. acid 21