Chapter 2 Atoms, Ions and Molecules PDF

Chapter 2 Atoms, Ions and Molecules Dr. Hicham H. Dib Dr. Fatma Hussain  Lecture Outline 2.1 The Atomic Theory 2.2 The Structure of the Atom 2.3 Atomic Number, Mass Number, and Isotopes 2.4 The Periodic Table 2.5 Molecules and Ions 2.6 Chemical Formulas 2.7 N...

Chapter 2 Atoms, Ions and Molecules Dr. Hicham H. Dib Dr. Fatma Hussain  Lecture Outline 2.1 The Atomic Theory 2.2 The Structure of the Atom 2.3 Atomic Number, Mass Number, and Isotopes 2.4 The Periodic Table 2.5 Molecules and Ions 2.6 Chemical Formulas 2.7 Naming Compounds  Atomic Theory Atoms are incredibly small, yet they compose everything. Atoms are the pieces of elements. Properties of the atoms determine the properties of the elements. For any given compound the atoms present always in same ratio. Chemical Laws: 1- Law of Conservation of Mass. 2- Law of Definite Proportion. 3- Law of Multiple Proportions. Chemical laws: 1- Law of Conservation of Mass Antoine Lavoisier “Matter is neither created nor destroyed in a chemical reaction.” The total amount of matter present before a chemical reaction is always the same as the total amount after. The total mass of all the reactants is equal to the total mass of all the products.  Conservation of Mass Total amount of matter remains constant in a chemical reaction. 64 g 36 g  Example—A Student Places Table Sugar and Sulfuric Acid into a Beaker and Gets a Total Mass of 144.0 g. Shortly, a Reaction Starts that Produces a “Snake” of Carbon Extending from the Beaker and Steam Is Seen Escaping. If the Carbon Snake and Beaker at the End Have a Total Mass of 129.6 g, How Much Steam Was Produced? Total of reactants and beaker = 144.0 g. Conservation of mass says total of products and beaker must be 144.0 g. Mass of steam = 144.0 g − 129.6 g = 14.4 g. 2- Law of Definite Proportion Is also known as the Law of constant composition, in different pure samples of a compound always contain the same elements in the same proportion by mass.  Example— Show that Two Samples of Carbon Dioxide Are Consistent with the Law of Constant Composition. Given: Sample 1: 4.8 g O, 1.8 g C; Sample 2: 17.1 g O, 6.4 g C Find: proportion O:C Solution Map: element masses compound composition Relationships: composition = mass O : mass C Solution: Sample 1 Sample 2 4.8 g O 17.1 g O = 2.7 = 2.7 1.8 g C 6.4 g C Compare: Since both samples have the same proportion of elements, carbon dioxide shows constant composition. 3- Law of Multiple Proportions When two elements, A and B, form more than one compound, If the masses of one element are the same in the two samples, then the masses of the other element are in a ratio of small whole numbers. 1C 1O 2O 1:2 Example: Which of the following pairs of compounds cannot be used to prove the law of multiple proportions? A. C2H6; C2H4 B. SO2; SO3 C is correct C. CO2; SO2 D. NO; NO2 E. H2O; H2O2  The Discovery of Atomic Structure In Dalton’s view, the atom was the smallest particle possible. Many discoveries led to the fact that the atom itself was made up of smaller particles (subatomic particles) which are: 1- Cathode rays (electrons) 2- Radioactivity (α particles, β particles and γ-ray) 3- Nucleus (protons and neutrons).  The structure of Atom 1- Dalton’s Atomic Theory(1808) Elements are composed of very small particles called atoms. All atoms of a given element are identical, having the same size, mass and chemical properties. Compounds are composed of atoms of more than one element. The relative number of atoms of each element in a given compound is always the same. Law of Definite proportion Chemical reactions only involve the rearrangement, separation, or combination of atoms. Atoms are not created or destroyed in chemical reactions. Law of Conservation of Mass  2- J.J. Thomson (English physicist, 1856-1940) J.J. Thomson Studied the electrical discharges in evacuated tubes called ‘Cathode-ray tubes’ 1906 Nobel Prize in Physics. The electron discovery. Early Atom Characterization. e direction Cathode Ray Tube J.J. Thomson, measured mass/charge of e− Glass tube evacuated from air (-1.76 × 108 C/g) Anode Cathode - + He found particles transfer from Cathode to Anode  it must be negative particles High Voltage Why Cathode Rays deflected away from negative charged plate?  Thomson said ‘since e- can be produced from different metals then all atoms must contains e-’ It is known that atoms are neutral. Because of that Thomson assumed that atoms also contain positive charges. He assumes atoms as a cloud of +ve charge with – ve electrons dispersed in it. 3- R. A. Millikan (British,1868-1953) He measured e- charge and mass of e- 1923 Nobel Prize in Physics He sprayed oil with atomizer Some of these small drops fall dawn from the upper small hole He determine the mass of oil drop from its velocity Next he ionized the chamber gas using X-ray source The electron charge and mass discovery. Millikank’s e- charge = -1.60 x 10-19 C Millikank’s e- mass = 9.10 x 10-28 g 4- A.H. Becquerel (French, 1845-1923) Nobel Prize in 1901. He discovered the three types of Radioactive Elements 1. α-particles (positively charged +2, i.e. 2 × 1.60 × 1019C ). 2. β-particles high speed e- (negatively charged, like electrons). 3. γ-rays high energy light (uncharged). 5- Rutherford (New Zealand,1871-1937) He did most of his work in England Nobel prize 1908. The discovery of proton and nucleus. He tested Thomson Model He directed  particles at a thin Sheet of metal - If there were no big deflection then Thomson model correct - If there were big deflection then Thomson model wrong He found some  particles not only deflected but even scattered or bounced back, So he assume a new model: - most of the atom are empty space - the atom’s +ve charge are concentrated in the center   particle velocity ~ 1.4 x 107 m/s (~5% speed of light) 1. atoms positive charge is concentrated in the nucleus 2. proton (p) has opposite (+) charge of electron (-) 3. mass of p is 1840 x mass of e- (1.67 x 10-24 g) Rutherford’s Model of the Atom - Atomic radius ~ 100 pm = 1 x 10-10 m - nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m - P + ve charge = e –ve charge = 1.6 × 10-19 C - N are neutral and its mass  mass of P 1.67 x 10-24 g Mass of atom  mass of nucleus -Nucleus are very dense. Atomic number, Mass number & Isotopes Atomic number = number of protons in nucleus Mass number = number of protons + number of neutrons = atomic number + number of neutrons = Isotopes Atomic mass = is an average mass of all isotopes  Isotopes Isotopes: atoms of the same element has the same number of P but different numbers of N  Example What is the atomic number of boron, B? 5 What is the atomic mass of silicon, Si? 28.09 amu How many protons does a chlorine atom have? 17 How many electrons does a neutral neon atom have? 10 Will an atom with 6 protons, 6 neutrons and 6 electrons be electrically neutral? Yes Will an atom with 27 protons, 32 neutrons, and 27 electrons be electrically neutral? Yes Will an Na atom with 10 electrons be electrically neutral? No  Introduction to periodic Table monatomic diatomic Elements are listed in order of Alkali Earth Metal increasing atomic number rather than relative atomic mass Alkali Metal Noble Gas Halogen Period Group  Metal left side Elements Nonmetals upper right side Metalloids B, Si, As, Te, Ge, Sb, Po  Periodic Table Metals: Good conductor of heat and electricity, malleability (can be hammered), ductility (can be wired), shiny, tend to lose e- to form +ve ions, form ionic bond with non-metals. Non-Metals: Poor conductor of heat and electricity, can not be hammered, can not be wired, not shiny, tend to gain e- to form -ve ions, form covalent bond with nonmetal. Metalloid: Intermediate properties between metals and non-metals For example Silicon: shiny, conduct electricity and does not conduct heat well.  Rows on the periodic table are called periods. Columns on the periodic table are called groups. Elements in the same group have similar chemical properties.  Molecules and Ions 1- Molecules A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical bonds H2 H2O NH3 CH4 A diatomic molecule contains only two atoms H2, N2, O2, Br2, HCl, CO A polyatomic molecule contains more than two atoms O3, H2O, NH3, CH4 A Covalent bond bond formed by sharing of e-( non-metalnon- metal) An Ionic bond bond formed by attraction between ions (metal  non- metal) 2- Ions An ion is an atom, or group of atoms, that has a net positive or negative charge. Cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. 11 protons 11 protons Na 11 electrons Na+ 10 electrons Anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. 17 protons 17 protons Cl 17 electrons Cl- 18 electrons A monatomic ion contains only one atom Na+, Cl-, Ca2+, O2-, Al3+, N3- A polyatomic ion contains more than one atom OH-, CN-, NH4+, NO3-  Chemical Formulas A Chemical formula express the composition of molecules and ionic compounds in terms of chemical symbol. Show exact number H2 O C6H12O6 N2 H4 1) Exact number 2) How atoms are bonded to one another H2 O CH2O NH2  Classifying Materials Atomic elements = Elements whose particles are single atoms.  Examples: He , Cu Molecular elements = Elements whose particles are multi- atom molecules.  Examples: H2, N2, O2, F2, Cl2, Br2, I2 Molecular compounds = Compounds whose particles are molecules made of only nonmetals or nonmetals + metalloids.  Examples: HCl , H2O , C3H8O Ionic compounds = Compounds whose particles are cations and anions (metal and nonmetal).  Examples: NaCl, KCl Classify Each of the Following as Either an Atomic Element, Molecular Element, Molecular Compound, or Ionic Compound Aluminum, Al = Atomic element Aluminum chloride, AlCl3 = Ionic compound Chlorine, Cl2 = Molecular element Acetone, C3H6O = Molecular compound Carbon monoxide, CO = Molecular compound Cobalt, Co = Atomic element  1- Ionic compound Ionic compounds consist of a combination of cations and an anions The sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero The ionic compound NaCl Write the Formula of a Compound Made from Aluminum Ions and Oxide Ions 1. Write the symbol for the metal Al+3 column 3A cation and its charge. 2. Write the symbol for the O2- column 6A nonmetal anion and its charge. 3. Charge (without sign) becomes Al+3 O2- subscript for the other ion. 4. Reduce subscripts to smallest whole-number ratio. Al2O3 5. Check that the total charge of the cations cancels the total Al = (2)∙(+3) = +6 charge of the anions. O = (3)∙(-2) = -6 Example: What Are the Formulas for Compounds Made from the Following Ions? K+ with N3- K3N Ca+2 with Br- CaBr2 Al+3 with S2- Al2S3 35 Example 1—Write the type and number of atom for Each of the Following Compounds. Fe3O4 4 O, 3 Fe Type of atoms: 2 No. of atoms: 7 Mg3(PO4)2 3 Mg, 2 P, 8O Type of atoms: 3 No. of atoms: 13 Example 2—Which formula represents the greatest total number of atoms? A) Al(C2H3O2)3 B) Al2(Cr2O7)3 C)Pb(HSO4)4  Naming Ionic Compounds Often a metal + nonmetal. Metal Cations take their names from the element, for example: Na (sodium) Na+ (sodium ion/cation) Anion (nonmetal), usually we add “-ide” to element name (Group 4A, 5A, 6A and 7A) K Opotassium oxide BaCl barium chloride © McGraw-Hill Education. 2-37  Ionic Compounds Transition metal + nonmetal Indicate charge on metal with Roman numerals FeCl 2𝐶𝑙 − 2 so Fe is + 2 iron (II) chloride FeCl 3𝐶𝑙 − 3 so Fe is + 3 iron (III) chloride © McGraw-Hill Education. 2-38  Common Inorganic Ions Table 2.3 Names and Formulas of Some Common Inorganic Cations and Anions Cation Anion aluminium  Al 3  bromide  Br   ammonium  NH 4  carbonate  CO32  barium  Ba 2  Chlorate  ClO3  cadmium  Cd 2  chloride  Cl   calcium  Ca 2  chromate  CrO42  cesium  Cs   cyanide  CN   chromium  III  or chromic  Cr 3  dichromate  Cr 2O72  cobalt  U  or cobaltous  Co 2  dihydrogen phosphate  H 2 PO4  copper  I  or cuprous  Cu   fluoride  F   copper  II  or cupric  Cu 2  hydride  H   hydrogen  H   hydrogen carbonate or bicarbonate  HCO3  iron  II  or ferrous  Fe 2  hydrogen phosphate  HPO4  iron  III  or ferric  Fe3  hydrogen sulfate or bisulfatc  HSO4  lead  II  or plumbous  Pb 2  hydroxide  OH   lithium  Li   iodide  I   magnesium  Mg 2  nitrate  NO3  © McGraw-Hill Education. 2-39  Common Inorganic Ions Cation Anion manganesc  II  or mangauous  Mn 2  nitride  N 3  mcrcury  I  or mercurous  Hg 22  * nitrite  NO2  mcrcury  II  or mercuric  Hg 2  oxide  O 2  potassium  K   permanganate  MnO4  rubidium  Rb   peroxide  O22  silver  Ag   phosphate  PO43  sodium  Na   sulfate  SO42  strontium  Sr 2  sulfide  S 2  tin  II  or stannous  Sn 2  sulfite  SO32  zinc  Zn 2  thiocynate  SCN   *The word "carbide" is also used for the anion 𝐶 © McGraw-Hill Education. 2-40  Home work #2 Name the following compounds: a Cu NO or copper II nitrate b KH PO c NH ClO Write chemical formulas for the following compounds: a) mercury(I) nitrite b) cesium sulfide c) calcium phosphate © McGraw-Hill Education. 2-41  2- Molecular Compounds Table 2.4 Greek Prefixes Used in Naming Molecular Compounds Nonmetals or nonmetals + metalloids Meani Prefix ng Nonmetals are written in order from Table 1. mono- 1 di- 2 1- Carbon dioxide CO2 tri- 3 tetra- 4 2- Acetone C3H6O2 penta- 5 3- Exceptions: NaOH hexa- 6 hepta- 7 octa- 8 nona- 9 deca- 10 Table 5.1 1 Order of Listing Nonmetals C P N H S I Br Cl O F in Chemical Formulas  Examples of Molecular Compounds HI hydrogen iodide 𝑁𝐹 nitrogen trifluoride 𝑆𝑂 sulfur dioxide 𝑁 𝐶𝑙 dinitrogen tetrachloride 𝑁𝑂 nitrogen dioxide 𝑁𝑂 dinitrogen monoxide © McGraw-Hill Education. 2-43 Example 1: Write chemical formulas for the following molecular compounds: a) carbon disulfide CS2 b) disilicon hexabromide Si2Br6 © McGraw-Hill Education. 2-44 1- 2- 3- © McGraw-Hill Education. 2-45 © McGraw-Hill Education. 2-46 © McGraw-Hill Education. 2-47