Atoms, Molecules, and Ions • Classification and Compositions of Matter • Atomic Structures • Ancient Philosophy • Dalton’s Atomic Theory • Isotopes • Periodic Table • Molecules and Ions • Types of Compounds • Naming Compounds
Matter According to Ancient Philosophy • Matter is composed of four basic elements: Earth, water, wind, and fire • Two schools of thought emerged during Greek Civilization: • Aristotle and his followers believed that matter to be infinite – not composed of discrete unit. • Democritus and Leucippus believed that matter is made of discrete units called “atomos” that is indivisible.
The “Development” of Theory • Many early 18th Century chemists the combustion process; They “observed” that when a piece of wood burn, the mass of the ash formed is apparently less than that of the original wood. So, what happen to the rest of the mass of the wood?
The Phlogiston Theory • This is their theory to explain combustion: • Materials burn because they contains a substance called phlogiston; • During combustion the phlogiston is lost; • Thus, the mass of ash is less than the wood. Really? What did they forget to do before coming to that conclusion?
Experimental Science versus Philosophy • Antoine Lavoisier (1743-1794) performed quantitative experiments to study combustion processes. The results showed that in all combustion reactions the masses of products to be greater than those of the material being burned. He concluded that: • The Phlogiston theory was incorrect; • Combustion involves oxygen gas – the gain in mass during combustion is due combination with oxygen; • Mass is conserved during chemical reactions.
Experimental Science versus Philosophy • Joseph Proust (1754-1826) performed numerous experiments to analyze the compositions of compounds, and found that: a given compound has a constant composition (in mass %) regardless of its origin or sample size.
The Law of Constant Composition Examples: • Copper carbonate is always composed (by mass %) of 51.4% Cu, 9.7% C, and 38.9% O . (b) Sodium chloride is always composed of 39.34% Na and 60.66% Cl, by mass.
Two Fundamental Laws of Matter • The Law of Conservation of Mass: During chemical reactions, the total mass of substances is conserved. (Mass is neither created or destroyed during chemical reaction.) • The Law of Definite Proportions: A given compound always contains the same types of elements chemically combined in a fixed proportion by mass, regardless of its origin.
Dalton’s Atomic Theory (1805) 1. Elements are made up of discrete, indivisible particles, called atoms; 2. Atoms of the same element are identical, but are different for different elements; 3.A compounds is formed when atoms of different elements combined in simple whole number ratios; 4. The smallest unit of a given compound always contains the same number and type of atoms; 5. Atoms are not created or destroyed during chemical reactions.
Principle of Chemical Combination • Law of Multiple Proportion: When two elements react to form more than one type of compounds, there exist a simple ratio of the masses of one of the elements that combine with a fixed mass of the other element in these compounds; Example: Carbon reacts with oxygen to form two compounds, X and Y. In X, 1.00 g of carbon combines with 1.33 g of oxygen, and in Y, 1.00 g of carbon combines with 2.66 g of oxygen. The mass ratio of oxygen in compounds X and Y is 1:2. If X = CO, then Y = CO2
Exercise #1: Law of Multiple Proportions • Sulfur reacts with fluorine to form three different compounds, A, B and C. In compound-A, 1.000 g of sulfur combines with 1.185 g of fluorine; in compound-B, 2.370 g of fluorine was found for every gram of sulfur, and in compound-C, the mass ratio of fluorine to sulfur is 3.556-to-1. Show that these data conform with the law of multiple proportions. Derive the formula of compounds A, B, and C.
Principle of Chemical Combination • Gay-Lussac’s Law of Combining Volumes: In reactions that involve gaseous reactants and products, there exist “simple ratios” of their volumes measured under the same temperature and pressure. Examples: • 1 volume of hydrogen reacts with 1 volume of chlorine to form 2 volumes of hydrogen chloride; • 2 volumes of hydrogen reacts with 1 volume of oxygen to form 2 volumes of water vapor.
Interpretation of Gay-Lussac’s Experiments • According to Avogadro’s law: “under same temperature and pressure, equal volumes of gases contain the same number of molecules” • 1 L of hydrogen + 1 L of chlorine 2 L of hydrogen chloride • This implies: 1 H-molecule + 1 Cl-molecule 2 HCl molecules. Hydrogen and chlorine molecules must be diatomic (2 atoms per molecule), and the reaction may be written as follows: H2(g) + Cl2(g) 2 HCl(g)
Interpretation of Gay-Lussac’s Experiments • 2 L of hydrogen + 1 L of oxygen 2 L of water vapor implies: 2 H-molecules + 1 O-molecule 2 water molecules. a) Hydrogen and oxygen gases contains diatomic molecules (H2 and O2), and water has the formula H2O. b) The above reaction can be represented by the equation: 2H2(g) + O2(g) 2 H2O(g)
Discovery of Cathode Ray • In 1895, J.J. Thomson discovered cathode-ray while studying the flow of electric current through a vacuum.
Characteristics of Cathode Ray • The ray originates from the cathode plate; • It is composed of negatively charged particles - it bends in an electric or a magnetic field in the direction similar to negatively charged particles; • The charge-to-mass ratio of the cathode ray particles is constant at -1.76 x 108 C/g, regardless of the materials used as the cathode; • Conclusion: cathode ray is a beam of negatively charged particles that we now called electron.
Thomson’s Model of Atom • J.J. Thomson proposed the “Plum-pudding” model: (a) Atom is composed of a diffused mass (like a cotton ball) of positive charge, with electrons loosely embedded on its surface; (b) The number of electrons present is such that the total negative charge is equal to the magnitude of positive charges in the atom.
Rutherford’s Atomic Model • The Nuclear Model: 1. Atom contains nucleus, composed of protons and neutrons; 2. The nucleus is much, much smaller than the atom; 3. Electrons occupy the vast “empty space” surrounding the nucleus; 4. The mass of atom is concentrated in the nucleus; 5. Proton or neutron is almost 2000 times larger and more massive than electron;
Relative and Absolute Masses • Proton: 1.007276 amu; 1.673 x 10-27 kg. • Neutron: 1.008665 amu; 1.675 x 10-27 kg. • Electron: 0.000549 amu; 9.109 x 10-31 kg.
Relative and Absolute Charges • Proton = +1; +1.602 x 10-19 C; • Neutron = 0; • Electron = -1; -1.602 x 10-19 C;
Isotopes • Atoms of the same element that have different masses; • Atoms having the same number of protons but different number of neutrons; • Atoms with the same atomic number (Z) but different mass number (A); Atomic number (Z) = number of protons; Mass number (A) = # of protons + # of neutrons; Number of neutrons = (A – Z)
Exercise #2: Isotope Symbols Write the symbols of isotopes that contain the following: (a) 10 protons, 10 neutrons, and 10 electrons. (b) 12 protons, 13 neutrons, and 10 electrons. (c) 15 protons, 16 neutrons, and 15 electrons. (d) 17 protons, 18 neutrons, and 18 electrons. (e) 24 protons, 28 neutrons, and 21 electrons. (a) 20Ne; (b) 25Mg2 +; (c) 31P; (d) 35Cl-; (e) 52Cr3+
Exercise #3: Isotopes Indicate the number of protons, neutrons, and electrons in each isotope with the following symbols. (a) 60Ni (b) 239Pu4+ (c) 79Se2- Answers: (a) 28 protons, 32 neutrons, and 28 electrons; (b) 94 protons, 145 neutrons, and 90 electrons; (c) 34 protons, 45 neutrons, and 36 electrons.
Molecules and Ions • Molecule: A neutral particle that contains two or more atoms bound together (chemically) by covalent bonds. • Ion: Electrically charged particle, either positive (called cation) or negative (called anion) An atom may lose one or more electrons to form a cation, or may gain electrons to form anions.
Periodic Table • The modern Periodic Table is divided into 18 columns (groups) and 7 rows (periods). • Groups are numbered 1 – 18 in the IUPAC configuration, or 1A – 8A and 1B – 8B in the ACS configuration. • In each period, elements are arranged left-to-right in increasing atomic number; • Within each group, elements share similar chemical and physical characteristics.
Major Classifications of Elements • Main group elements: • Group 1A (1): the alkali metals; • Group 2A (2): the alkaline Earth metals; • Groups 3A (13), 4A (14), 5A (15), and 6A (16), • Group 7A (17): the halogens, and • Group 8A (18): the noble gases. • Transition metals: Groups 3B (3) – 2B (12) ; contains heavy metals. • Metalloids (semi-metals): B, Si, Ge, As, Sb, Te, Po, At
Characteristics of Metals and Nonmetals • Metals: • Mainly solid, except for mercury; have shiny appearance; • Good conductors of heat and electricity; • Malleable and ductile; • Nonmetals: • Mainly gases, one (bromine) is a liquid, and a few solids; • Generally poor conductors of electricity; • Solids are generally brittle and not lustrous. • Metalloids (semi-metals): • Very hard (covalent network) solids; • physically look like metals, but chemically behave like nonmetals;
More on Classifications of Elements • Lanthanide series: Elements after lanthanum (La): Ce, Pr, Nd, Pm, Sm, Eu, Gd, Tb, Dy, Ho, Er, Tm, Yb, and Lu; • Actinide series: • Elements after actinium (Ac): Th, Pa, U, Pu, Am, Cm, Bk, Cf, Es, Fm, Md, No, and Lr; • Mostly synthesized in particle accelerators and all are radioactive;
Nomenclature • Type-I (ionic) compounds: contain cations with fixed charges: Group 1 and Group 2 metals, and aluminum. • Type-II (ionic) compounds: contain cations with variable charges: transition metals, In, Sn, Tl, Pb, or any metals from the lanthanide or actinide series. • Type-III (molecular) compounds: contain only nonmetals or metalloids and nonmetals;
Type-I (Ionic) Compounds • Binary compounds: NaCl: sodium chloride; MgF2: magnesium fluoride; Al2O3: aluminum oxide; • Compounds containing polyatomic ions: CaSO4: calcium sulfate; NaHCO3: sodium hydrogen carbonate; KNO3: potassium nitrate
Type-II (Ionic) Compounds • Binary compounds: FeCl2: iron(II) chloride; FeCl3: iron(III) chloride; CrO: chromium(II) oxide; Cr2O3:chromium(III) oxide; • Compounds containing polyatomic ions: Co(NO3)2: cobalt(II) nitrate; Co(NO3)3: cobalt(III) nitrate; Pb(C2H3O2)2: lead(II) acetate; Pb(C2H3O2)4: lead(IV) acetate;
Type-II Compounds Naming System FormulaStock SystemOld System CrO chromium(II) oxide chromous oxide Cr2O3chromium(III) oxide chromic oxide Fe(NO3)2Iron(II) nitrate Ferrous nitrate Fe(NO3)3Iron(III) nitrate Ferric nitrate
Exercise #4: Formula Write the formulas of the following compounds: (a) Aluminum nitrate (b) Barium chromate (c) Magnesium carbonate (d) Iron(III) chloride (e) Lead(II) acetate (f) Nickel(II) sulfate hexahydrate
Exercise #4 Answers (a) Aluminum nitrate=Al(NO3)3 (b) Barium chromate = BaCrO4 (c) Magnesium carbonate = MgCO3 (d) Iron(III) chloride = FeCl3 (e) Lead(II) acetate= Pb(C2H3O2)2 (f) Nickel(II) sulfate hexahydrate = NiSO46H2O
Exercise #5: Nomenclature Name the following compounds: (a) Ca(OH)2 (b) Cr(NO3)3 (c) FeSO4 (d) NaHCO3 (e) KH2PO4 (f) CuCl22H2O
Exercise #5 Answers (a) Ca(OH)2 = Calcium hydroxide (b) Cr(NO3)3 = Chromium(III) nitrate (c) FeSO4 = Iron(II) sulfate (d) NaHCO3 = Sodium hydrogen carbonate (e) KH2PO4 = Potassium dihydrogen phosphate (f) CuCl22H2O = Copper(II) chloride dihydrate
Naming Binary Acids (Acids without oxygen in the formula): Hydro + first syllable of anion + ic acid HF = hydrofluoric acid (weak) HCl = hydrochloric acid (strong) HBr = hydrobromic acid (strong) HI = hydroiodic acid (very strong) H2S = hydrosulfuric acid (weak) HCN = hydrocyanic acid (very weak)