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Chapter 2

Chapter 2. Atoms, Molecules, and Ions. Preview. The chapter covers the following: Fundamental Chemical Laws and Atom. Modern View of Atomic Structure, Molecules, and Ions. Periodic Table . Naming Simple compounds, Ionic compounds, Formula from names. 2.1 The early history of chemistry.

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Chapter 2

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  1. Chapter 2 Atoms, Molecules, and Ions

  2. Preview The chapter covers the following: • Fundamental Chemical Laws and Atom. • Modern View of Atomic Structure, Molecules, and Ions. • Periodic Table. • Naming Simple compounds, Ionic compounds, Formula from names.

  3. 2.1 The early history of chemistry • Greeks proposed:Matter was composed of four fundamental substances:fire, earth, water, and air. • Greeks also considered: Substances are composed of small indivisible particles. Abdera & Leucippus used the term atomos (became atoms) to describe these ultimate particles. • Foundation of Modern Chemistry: In 1494 – 1555 the development of systematic metallurgy (extraction of metals from ores) of chemistry science was established by George Barer (German). • Truly Quantitative Chemical Experiment: In 1627 – 1691 the first truly quantitative experiments were performed by Robert Boyle who measured the relationship between pressure and volume of air. • Homework Read Assignment to answer the following: • What was Boyle’s experiment? • Who did discover the oxygen and its effect on combustion experiment?

  4. 2.2 Fundamental Chemical Laws • 1743 – 1794, Antoine Lavoisier explained the nature of combustion and like Boyle, he carefully measured the weight of reactant and product. He suggested that "Mass is neither created nor destroyed". • Lavoisier states that combustion involvedOxygenand life process used oxygen – similar to combustion. • Lavoisier published the first modern chemistry textbook "Elementary Treatise on Chemistry"

  5. 2.2 Fundamental Chemical Laws • 1754 – 1826, Joseph Proust determined the composition of various chemical compounds. Proust's Law or “Law of definite proportion” “a given compound always contains exactly the same proportion of elements by mass" • 1766 – 1844, John Dalton discovered that "different relative amounts of elements may also found for similar elements.”{Law of multiple proportions} • i.e. cpd. I is CO and cpd. II is CO2

  6. 2.2 Fundamental Chemical Laws Example 2.1: The hollowing data were collected for several compounds of nitrogen and oxygen: • Note: In fact, an infinite number of combination exists: • N2O , NO , NO2 • NO , NO2 , NO4 • N4O2 , N2O2 , N2O4

  7. 2.3 Dalton's Atomic Theory The modern phases of Dalton's ideas are: • Each element is made up of tiny particles called atoms. • The atoms of a given element are identical. • chemical compounds are formed when atoms combine with each other. • Chemical reactions involve reorganization of the atoms changes in the way they are bound together. Dalton's theory lead to: 1gm hydrogen + 8gm of oxygen  water he assume that water formula is "OH" and the mass of hydrogen is "1" and of oxygen is "8". Using the same concepts, Dalton's proposed the first table of atomic masses. It has been proved later that Dalton's table contain incorrect masses.

  8. 2.3 Dalton's Atomic Theory Gay-lussac (1778-1850) found experimentally that: Avogadro proposed that "at the same temperature and pressure, equal volumes of different gases contain the same number of particles". If Avogadro's hypothesis is correct, Gay-Lussaci result (figure 2.5): Note: this suggests that the formula for water is H2O, not OH as Dalton believed.

  9. 2.3 Dalton's Atomic Theory • Jakob Berzelius (1779 – 1848) use the same Avogadro's concept. • He discovered the elements cerium, selenium, silicon, and thorium. • He developed the modern symbols for the elements. e.g. Old new symbols 1) Silver ﴿ Ag 2) Lead 5Pb 3) Tin 4Sn 4) Sea salt Θ NaCl

  10. 2.4 Early Experiments to Characterize the Atom The concept of atoms was a good idea, Then scientists began to wonder about the nature of the atom. • What is an atom made of? • How do the atoms of various elements differ? EXP # 1: J. Thomson (1898 – 1903) use the cathode-ray-tube: • He observed an existence of –ve rays produced at the negative electrode and repelled by –ve pole • He determine the charge –to- mass ratio of this particles: e/m = -1.76 x 108 c/g This particle, now called “Electrons”

  11. 2.4 Early Experiments to Characterize the Atom EXP # 2: Millikan does another experiment to determine the mass of this –ve particles (electrons). The experiment used mainly to determine the magnitude of the electron charge and using e/m to get m- value: Mass of electron (me) = 9.11 x 10-13 kg

  12. Radioactivity 2.4 Early Experiments to Characterize the Atom Henri Becquerel (1896) found accidentally that piece of mineral containing Uranium could produce image on photographic plates. • Uranium emitted high radiation and this is called Radioactivity. • Studies on radioactive material showed three types of emission are expected: • g rays : high energy light • β particles : -ve particles = electrons • a particles : +ve particles has two +ve, and mass 7300 times mass of electron

  13. The nuclear Atom Ernest Rutherford (1911) did many experiment to explore radioactivity found the a-particles reflected and not completely pass through a metal sheet as expected by Thomson's model of the atom is correct [plum pudding made]. Moved electron at a distance = 10-8 cm Nucleus Diameter = 10-13 cm Piece of nucleus material with pea size has mass equal 250 million tons!!.

  14. Mass Charge 31 - Electron 9.11x10 kg 1 - 27 - kg Proton 1.67x10 1+ 27 - Neutron 1.67x10 kg None 2.5 The Modern View of Atomic Structure Rutherford’s Model of the Atom: • More studies on nucleus showed that nucleus contain protons (+ve particles) and neutrons (neutral particles) and electrons moving around. • Mass and charges of all of these are: atomic radius ~ 100 pm = 1 x 10-10 m nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m

  15. A X Mass Number Element Symbol Z Atomic Number 2 3 1 H (D) H (T) H 1 1 1 2.5 The Modern View of Atomic Structure • The atomic components are related and con be derived from the element symbols notation: Isotopes • Different mass number between two elements have the same atomic number are called "isotopes". • Isotopes: have similar chemical properties. "Z" = number of protons = number of electrons for neutral atoms. "A" = total number of protons and neutrons.

  16. H2 H2O NH3 CH4 2.6 Molecules and Ions • Atoms combined with each others to form compounds. • Forces that hold atoms together in compounds are called chemical bonds. Only, electrons participate in bonding. • There are two types of chemical bonding: • Electron sharing: this lead to covalent bond formation, and the resulted compound is called a molecule. There are two ways to present the molecules: (a) Chemical Formula: (Simplest way) e.g. H2O, CO2, NH3, C2H6O, etc (b) Structural Formula: (more information) It may or may not indicate the actual shape of molecule.

  17. 11 protons 11 electrons 11 protons 10 electrons Na+ Na 17 protons 18 electrons 17 protons 17 electrons Cl- Cl 2.6 Molecules and Ions • Electron sharing: • Attraction Among Ions: this lead to ionic bonding and the formed compound is called ionic composure ionic solid, or salt. Ions are atoms or group of atoms that have a net (+ve) or (-ve) charge. e.g. cation – ion with a positive charge: If a neutral atom loses one or more electrons it becomes a cation. anion – ion with a negative charge: If a neutral atom gains one or more electrons it becomes an anion. 2.5

  18. 2.6 Molecules and Ions • Using Molecular Modeling Computer software one can build many molecules. 2.5

  19. Hydrogen (name) Sodium = [original name is Natriam] 2.7 An Introduction to the Periodic Table The simple periodic table (figure 2.21) contains symbols that represent elements. • Element symbols are abbreviations based on the element name or the original name e.g.

  20. 2.7 An Introduction to the Periodic Table • Most of the elements are metals. They have characteristic physical properties e.g. • High heat and electric conduction. • Malleability (hammered to sheet) • Ductility (pulled into wires) • Chemically metals tend to lose electrons to form +ve-ions. Fe2+ , Fe3+ , Na+ ,K+ , Ca2+ • Upper right – hand corner contains nonmetals. They lack the physical properties of metals. • Nonmetals tend to gain electrons to for –ve-ions. Cl-, F-, O2-, S2-… • Nonmetals tend to bond with each other by forming covalent bonds.

  21. Periodic Table arranged in: • vertical columns called groups or families e.g. alkali metals, Noble gases, halogen …etc. • Horizontal rows called periods e.g. first period, second period …etc. • Noble gases group exist normally as monoatomic (Single atoms) because they have little chemical reactivity. • all elements per group have similar chemical properties e.g. gp I: alkali metals only form +ve ions with one +ve charge. gp II: Alkali earth metals form +ve ions with two +ve charge. 2.4

  22. Noble Gas Halogen Alkali Earth Metal Period Alkali Metal Group

  23. 2.8 Naming Simple Compounds: • The following systematic naming will be for both ionic and covalent compounds: • Binary ionic compounds (type I): Containpositive ionNegative ion Cation anion Rule: Cation named first as name of element. Anion: named after cation, root + ide. e.g. Na Cl Sodium Chloride.

  24. Binary ionic compounds (type I): • Binary Ionic Compounds (Type II): • Many metals (transition metals) form more than one type of ionic compound e.g. Fe Cl2 and Fe Cl3. To name these compounds we have: Rule: • Determine the charge on the metal. • Use Roman numeral to indicate the charge of the cation (I, II, III, IV,…). • Use same naming as type I. e.g. Fe Cl2 Fe Cl3 Fe2+ Fe3+ Iron (II) Chloride Iron (III) Chloride

  25. Figure 2.22: The common cations and anions

  26. Example 2.4 Give the systematic name of the following compounds. a. CuCl b. HgO c. Fe2O3 d. MnO2. Note: for elements that will not give more than one ionic species DON'T use Roman Numeral. e.g. Ag Ag+ only Al Al3+ onl Zn Zn2+ only gp I elements gp II elements

  27. Note: For polyatomic ion formulas Special names are given and must be memorized.

  28. Binary ionic compounds (type I): • Binary Ionic Compounds (Type II): • Binary compounds (Type III): Compounds formed from Nonmetalsand give Covalent molecules. Naming is similar to type I for RHS-element and LHS-element + prefixes to denote the number of atoms present. (mono, di, tri, tetra, penta, heca, etc..) e.g. N2O Dinitrogen monoxide NO nitrogen monoxide

  29. Figure 2.23: A flowchart for naming binary compounds Figure 2.24: Overall strategy for naming chemical compounds 2.8 Naming Simple Compounds:

  30. Example 2.8 Give the systematic name for each of the following : a. P4O10 b. Nb2O5 c. Li2O2 d. Ti(NO3)4 Tetra phosphorus decaoxide. Niobium (V) oxide. Lithium peroxide (O22- peroxide). Titanium (lV) nitrate.

  31. 2.8 Naming Simple Compounds Acids: Acids are the molecules that produce H+ when they dissolved in water (chapter 4). There are two types of acids: (a) H+ + anion does not contain oxygen. Names given by adding prefix (hydro-) and suffix (-ic). e.g. HCl Hydrochloric acid H2S Hydrosulfuric acid (b) H+ + anion contain oxygen. Naming by using root + -ic …, or, root + -ous … or, Depending on the name of the anion. e.g. anion ends is –ate: acid -ic SO42-Sulfate Sulfuric acid anion ends is – ite acid -ous SO32- Sulfite  Sulfurous acid

  32. 2.8 Naming Simple Compounds Figure 2.25: A flowchart for naming acids. An acid is best considered as one or more H+ ions attached to an anion.

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