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Chapter 2 Atoms, Molecules, and Ions Learning a Language • When learning a new language: • Start with the alphabet • Then, form words • Finally, form more complex structures such as sentences • Chemistry has an alphabet and a language; in this chapter, the fundamentals of the language of chemistry will be introduced Outline • • • • • • Atoms and Atomic Theory Components of the Atom Introduction to the Periodic Table Molecules and Ions Formulas of Ionic Compounds Names of Compounds The Language of Chemistry • This chapter introduces the fundamental language of chemistry • Atoms, molecules and ions • Formulas • Names The Structure of Matter • Atoms • Composed of electrons, protons and neutrons • Molecules • Combinations of atoms • Ions • Charged particles Laws of Chemical Composition Conservation of Mass - The total mass remains constant during a chemical reaction. EX: A bicycle rusting Law of Definite Proportions Law of Definite Proportions : All samples of a compound have the same composition; that is, all samples have the same proportions, by mass, of the elements present Ex: Water always contains: ~89% oxygen ~11% hydrogen John Dalton Atoms and Atomic Theory • An element is composed of tiny particles called atoms • All atoms of the same element have the same chemical properties • In an ordinary chemical reaction • There is a change in the way atoms are combined with each other • Atoms are not created or destroyed • Compounds are formed when two or more atoms of different element combine • Atoms are indivisible and indestructible Figure 2.1 - John Dalton and Atomic Theory Law of Multiple Proportions Law of Multiple Proportion: When two or more different compounds of the same two elements are compared, the masses of one element that combine with the fixed mass of the second element are in the ratio of small whole numbers. Figure A – The Law of Multiple Proportions Two different oxides of chromium Law of Multiple Proportions • Atomic theory raised more questions than it answered • Could atoms be broken down into smaller particles • 100 years after atomic theory was proposed, the answers were provided by experiment Fundamental Experiments • J.J. Thomson, Cavendish Laboratories, Cambridge, England • Ernest Rutherford • McGill University, Canada • Manchester and Cambridge Universities, England Figure 2.2 – J.J. Thomson and Ernest Rutherford Electrons J.J. Thomson, 1897 • First evidence for subatomic particles came from the study of the conduction of electricity by gases at low pressures • Rays emitted were called cathode rays • Rays are composed of negatively charged particles called electrons • Electrons carry unit negative charge (-1) and have a very small mass (1/2000 the lightest atomic mass) The Electron and the Atom • Every atom has at least one electron • Atoms are known that have one hundred or more electrons • There is one electron for each positive charge in an atom • Electrical neutrality is maintained Figure 2.3 – Cathode Ray Apparatus Robert Millikan and the Oil Drop Experiment, 1909 • Produced small oil drops, acquire an electric charge, measure velocity of a falling droplet with and without an electric field • Obtained the charge on an electron, which coupled with Thomson’s work, allowed the calculation of the mass of an electron. • Calculated the charge of an electron: -1.602 * 10-19 coulomb (C) • Calculated the mass of an electron: 9.109 * 10-31 kg Protons and Neutrons – The Nucleus • Ernest Rutherford, 1911 • Bombardment of gold foil with alpha particles (helium atoms minus their electrons • Expected to see the particles pass through the foil • Found that some of the alpha particles were deflected by the foil • Led to the discovery of a region of heavy mass at the center of the atom Rutherford’s Gold Foil Experiment • to prove or disprove Thomson’s model • alpha particles (Helium atoms, 2 protons and 2 neutrons) shot at gold foil • most particles went through, but some came back at different angles • “It is about as incredible as if you had fired a 15-inch shell at a piece of tissue paper and it came back and hit you.” -ER • discovered the nucleus (positive core of an atom) Figure 2.4 – Rutherford Backscattering Thomson’s Plum Pudding Model • aka chocolate chip cookie model • positive charge uniformly distributed in a sphere • electrons imbedded in the sphere YUMMY!!!! Rutherford's Model • Nucleus in the center of the atom containing protons • Electrons in the space surrounding the nucleus Nuclear Particles 1. Protons • Mass nearly equal to the H atom • Positive charge 2. Neutrons • Mass slightly greater than that of the proton • No charge Mass and the Atom • More than 99.9% of the atomic mass is concentrated in the nucleus • The volume of the nucleus is much smaller than the volume of the atom Table 2.1 – Subatomic Particles Protons and neutrons are located at the center of an atom (at the nucleus). Electrons are dispersed around the nucleus. Terminology • Mass number, A = • Number of protons plus number of neutrons • Atomic number, Z= • Number of protons in the atom Nuclear symbolism A Z X • A is the mass number • Z is the atomic number • X is the chemical symbol 19 1 F 9 Isotopes • Isotopes are two atoms of the same element • Same atomic number • Different mass numbers • Number of neutrons is A-Z • Number of neutrons differs between isotopes • Isotopes & Their Uses • Heart scans with radioactive technetium-99. • 99 43Tc • Emits gamma rays Isotopes of hydrogen • 1H, 2H, 3H • Hydrogen, deuterium, tritium • Different masses 1 H 1 1 proton and 0 neutrons, protium 2 H 1 proton and 1 neutron, 1 deuterium 3 H 1 proton and 2 neutrons, 1 tritium radioactive Note that some of the ice is at the bottom of the glass – this is 2H2O Isotopes Atoms can be represented using the element’s symbol and the mass number (A) and atomic number (Z): A E Z 35 Cl 17 37 Cl 17 Isotopes • Example 2.1 Write the atomic symbols for the following species: • a. the isotope of carbon with a mass of 13 • b. the nuclear symbol when Z= 92 and then number of neutrons = 146 Answers: 13 C 6 238 U 92 Ions • An electrically charged particle comprised of one or more atoms Example 2.2 • Write the atomic symbols for the following: • a. a species having 16 protons, 16 neutrons and 18 electrons • b. the phosphide ion (symbol = P) with an overall charge of -3 Answers: 32 S 2- 16 31 P 315 An atomic mass unit (amu) is defined as exactly onetwelfth the mass of a carbon-12 atom • 1 u = 1.66054 × 10–24 g • The atomic mass of an element is the weighted average of all of the isotopes an element has • Weighted average is the addition of the contributions from each isotope • percent abundance is the percent or fraction of each isotope found in nature. • Example 2.3 Determine the average atomic mass of magnesium which has three isotopes with the following masses: 23.98 (78.6%), 24.98 (10.1%), 25.98 (11.3%). Radioactivity • Radioactive isotopes are unstable • These isotopes decay over time • Emit other particles and are transformed into other elements • Radioactive decay is not a chemical process! • Particles emitted • High speed electrons: β (beta) particles • Alpha (α) particles: helium nuclei • Gamma (γ) rays: high energy light Nuclear Stability • Nuclear stability depends on the neutron/proton ratio • For light elements, n/p is approximately 1 • For heavier elements, n/p is approximately 1.4/1 • The belt of stability Figure 2.5 – The Nuclear Belt of Stability Introduction to the Periodic Table The Periodic Table: Elements Organized • Know location and description of: • groups or families • periods or series • metals, metalloids, nonmetals and their properties • main group elements • transition metals • lanthanides and actinides Periods and Groups • Vertical columns are groups • IUPAC convention: use numbers 1-18 Families with Common Names • • • • Alkali Metals, Group 1 Alkaline Earth Metals, Group 2 Halogens, Group 17 Noble Gases, Group 18 Importance of Families • Elements within a family have similar chemical properties • Alkali metals are all soft, reactive metals • Noble gases are all relatively unreactive gases; He, Ne and Ar do not form compounds Group 1A: Alkali Metals Li, Na, K, Rb, Cs Reaction of potassium + H2O Cutting sodium metal Group 2A: Alkaline Earth Metals Be, Mg, Ca, Sr, Ba, Ra Magnesium Magnesium oxide Group 3A: B, Al, Ga, In, Tl • Al resists corrosion Gallium is one of the few metals that can be liquid at room temp. Group 4A: C, Si, Ge, Sn, Pb Quartz, SiO2 Diamond Group 5A: N, P, As, Sb, Bi Ammonia, NH3 White and red phosphorus Group 6A: O, S, Se, Te, Po Sulfuric acid dripping from snot-tite in cave in Mexico Elemental S has a ring structure. Group 7A: Halogens F, Cl, Br, I, At Group 8A: Noble Gases He, Ne, Ar, Kr, Xe, Rn • Horizontal rows are periods • First period is H and He • Second period is Li-Ne • Third Period is Na-Ar Arrangement of Elements • Periods • Arranged by increasing atomic number • Families • Arranged by chemical properties Metals and Nonmetals • Diagonal line starting with B separates the metals from the nonmetals • Elements along this diagonal have some of the properties of metals and some of the properties of nonmetals • Metalloids • B, Si, Ge, As, Sb, Te Blocks in the Periodic Table • Main group elements • 1, 2, 13-18 • Transition Metals • 3-12 • Post-transition metals • Elements in groups 13-15 to the right of the transition metals • Ga, In, Tl, Sn, Pb, Bi • Lanthanides and Actinides • Below the main part of the table A Look at the Sulfur Group • Sulfur (nonmetal), antimony (metalloid) and silver (metal) Biological View of the Periodic Table • “Good guys” • Essential to life • Carbon, hydrogen, oxygen, sulfur and others • “Bad guys” • Toxic or lethal • Some elements are essential but become toxic at higher concentrations • Selenium Figure 2.8 – Biologically Important and Toxic Elements Mendeleev • Dmitri Mendeleev, 1836-1907 • Arranged elements by chemical properties • Left space for elements unknown at the time • Predicted detailed properties for elements as yet unknown • Sc, Ga, Ge • By 1886, all these elements had been discovered, and with properties similar to those he predicted Moseley • Arranged the periodic table according to atomic number Molecule • Two or more atoms may combine to form a molecule • Atoms involved are often nonmetals • Covalent bonds are strong forces that hold the atoms together • Molecular formulas • Number of each atom is indicated by a subscript • Examples • Water, H2O • Ammonia, NH3 Empirical and Molecular Formulas Empirical formula: the simplest whole number ratio of elements in a compound • A molecular formula : gives the number of each kind of atom in a molecule. • Example: Molecular formula of glucose – C6H12O6 The elemental ratio C:H:O is 1:2:1, so the empirical formula is CH2O Structural Formulas • Structural formulas show the bonding patterns within the molecule Structural Formulas • Condensed structural formulas suggest the bonding pattern and highlight specific parts of a molecule, such as the reactive group of atoms Ball and Stick Models Ions • When atoms or molecules lose or gain electrons, they form charged particles called ions • Na Na+ + e• O + 2e- O2• Positively charged ions are called cations • Negatively charged ions are called anions • There is no change in the number of protons in the nucleus when an ion forms. Example 2.3 Polyatomic Ions • Groups of atoms may carry a charge; these are the polyatomic ions • OH• NH4+ Ionic Compounds • Compounds can form between anions and cations • Sodium chloride, NaCl • Sodium cations and chloride ions associate into a continuous network Forces Between Ions • Ionic compounds are held together by strong forces • Electrostatic attraction of + and – for each other • Compounds are usually solids at room temperature • High melting points • May be water-soluble Solutions of Ionic Compounds • When an ionic compound dissolves in water, the ions are released from each other • Strong electrolytesPresence of ions in the solution leads to electrical conductivity • When molecular compounds dissolve in water, no ions are formed • NonelectrolytesWithout ions, solution does not conduct electricity Figure 2.12 – Electrical Conductivity Formulas of Ionic Compounds • Charge balance • Each positive charge must have a negative charge to balance it • Calcium chloride, CaCl2 • Ca2+ • Two Cl- ions are required for charge balance Noble Gas Connections • Atoms that are close to a noble gas (group 18) form ions that contain the same number of electrons as the neighboring noble gas atom • Applies to Groups 1, 2, 16 and 17, plus Al (Al3+) and N (N3-) Cations of Transition and Post-Transition Metals • Iron • Commonly forms Fe2+ and Fe3+ • Lead • Commonly forms Pb2+ and Pb4+ Copper forms either copper(I) or copper(II) ions. Titanium forms both titanium(II) and titanium(IV) ions. What is the charge on a zirconium(IV) ion? Polyatomic Ions • There are only two common polyatomic cations • NH4+ and Hg22+ • All other common polyatomic ions are anions Table 2.2 – Polyatomic ions Oxoanions • When a nonmetal forms two oxoanions • -ate is used for the one with the larger number of oxygens • -ite is used for the one with the smaller number of oxygens • When a nonmetal forms more than two oxoanions, prefixes are used • per (largest number of oxygens) • hypo (smallest number of oxygens) Polyatomic Ions • Oxoanions: the anions are composed of oxygen and one other element • Ex:SO42- (sulfate), NO2- (nitrite) , MnO4(permanganate) • two oxoanions of the same element • The anion with the smaller number of oxygens uses the roots of the element plus “ite” • The higher number use the root plus “ate” • Ex: SO32- sulfite, NO2- nitrite, PO3-3 phosphite • SO42- sulfate,NO3- nitrate, PO4-3 phosphate Table 2.3 – Oxoanions of Nitrogen, Sulfur and Chlorine 2.6 Names of Compounds - Cations Monatomic cations take the name from the metal from which they form • Na+, sodium ion • K+, potassium ion If more than one charge is possible, a Roman numeral is used to denote the charge • Fe2+ • Fe3+ iron(II) ion iron(III) ion Metals with multiple oxidation states • Two methods: Stock method and “classical” method • Stock system: • metal name and the oxidation state in Roman numbers in parenthesis • Ex: Fe 2+ = iron(II) • Form compound by balance charge of metal with correct number of nonmetals • Ex: CoCl3 = cobalt(III) chloride Classical Method The name of the metal ion that has the lower charge ends in “ous”, the higher charge ends in “ic” Names of Compounds - Anions • Monatomic anions are named by adding –ide to the stem of the name of the element from which they form • Oxygen becomes oxide, O2• Sulfur becomes sulfide, S2• Polyatomic ions are given special names (see table 2.3, p. 39) Ionic Compounds • Combine the name of the cation with name of the anion • Cr(NO3)3, chromium(III) nitrate • SnCl2, tin(II) chloride Write the formulas for: • rubidium bromide strontium hydroxide • barium nitride cobalt (II) sulfate • cobalt (II) bromide calcium phosphate • strontium sulfide tin (IV) phosphate Write the formulas for: • • • • RbBr Ba3N CoBr2 SrS Sr(OH) 2 CoSO4 Ca3(PO4) 2 Sn3(PO4) 4 Example 2.6 Example: Name the following binary ionic compounds • Metal nonmetal compound name • KI potassium iodine potassium iodide • Li2S lithium sulfur lithium sulfide • Mg3N2 magnesium nitrogen magnesium nitride Write the names for: • AlCl3 (NH4)2S K2Cr2O7 Ca3P2 Al(NO2)3 KMnO4 FeI2 PbO2 Write the names for: • aluminum chloride • potassium dichromate • aluminum nitrite ammonium sulfide calcium phosphide potassium permanganate • Iron II idodide Lead (IV) oxide Binary Molecular Compounds • Unlike ionic compounds, there is no simple way to deduce the formula of a binary molecular compound • Systematic naming 1. The first word is the name of the first element in the formula, with a Greek prefix if necessary • Mono is never used on the first word 2. The second word consists of • • • The appropriate Greek prefix The stem of the name of the second element The suffix -ide Writing Formulas of Binary Molecular Compounds • Find the number associated with the prefixes and put it after the element. • One is never shown. Table 2.4 - Greek Prefixes Names of Binary Compounds The lines trace a continuous path from boron (B) to fluorine (F). The element closer to the beginning of this path is generally written first in the formula of a binary molecular compound. Some Examples • Binary nonmetallic compounds • N2O5, dinitrogen pentaoxide • N2O4, dinitrogen tetraoxide • NO2, nitrogen dioxide • N2O3, dinitrogen trioxide • NO, nitrogen oxide • N2O, dinitrogen oxide • Common names • H2O, water • H2O2, hydrogen peroxide Common Molecular Compounds Example 2.7 Write names or formulas for the following: • B2O3 • AsO5 • As2O7 tetraphosphorus pentachloride dihyrdogen monoxide Answers: • Diboron trioxide • Arsenic pentoxide • Arsenic heptoxide P4Cl5 H2O Acids • Acids ionize to form H+ ions • Hydrogen and chlorine • As a molecule, HCl is hydrogen chloride • When put in water, HCl is hydrochloric acid Common Acids Oxoacids • Two common oxoacids • HNO3, nitric acid • H2SO4, sulfuric acid Oxoacids of Chlorine Example 2.8 Hydrates A hydrate is an ionic compound in which the formula unit includes a fixed number of water molecules associated with cations and anions Examples: BaCl2 . 2 H2O CuSO4 . 5 H2O