Chapter 2 Atoms, Molecules, and Ions

1. Chapter 2Atoms, Molecules, and Ions Prof. Dr. Nizam M. El-Ashgar ChemistryDepartment, IUG

2. THE ATOMIC THEORY OF MATTER Historically: • Greek Philosophers: matter can be subdivided into fundamental particles. • Democritus (460–370 BC) and other early Greek philosophers described the material world as made up of tiny indivisible particles they called atomos, meaning “indivisible or uncuttable.” • The “atomic” view of matter faded for many centuries. • The notion of atoms reemerged in Europe during the seventeenth century. • That theory came from the work of John Dalton during the period from 1803 to 1807.

3. The Atomic Theory John Dalton’s 4 Postulates: • Each element is composed of extremely small particles called atoms. • All atoms of an element are identical to each other, but different than atoms of other elements. • Atoms of one element cannot be changed into atoms of different elements by chemical reactions; atoms are neither created or destroyed in reactions. • Compounds are formed when atoms of different elements combine in a definite ratio.

4. Hg(NO3)2(aq) + 2KI(aq) HgI2(s) + 2KNO3(aq) I) Law of Conservation of Mass Mass is neither created nor destroyed in chemical reactions. 3.25 g + 3.32 g = 6.57 g 4.55 g + 2.02 g = 6.57 g

5. II) the Law of Definite Proportions Different samples of a pure chemical substance always contain the same elements with constant proportion of elements by mass. By mass, water (H2O) is: 88.8 % oxygen 11.2 % hydrogen

6. III) Law of Multiple Proportions: • If 2 elements combine to form more than one compound, and if one element presents in fixed mass the masses of the second element exist in small whole number ratios. • mall whole number ratios: 2/1, 1/2, 3/2, 2/3, 3/4,,,,,,,, • This follows from the postulate that individual atoms enter into chemical combination.

7. The Discovery of Atomic Structure J. J. Thomson (1898—1903) • Postulated the existence of electrons using cathode-ray tubes. • Determined the charge-to-mass ratio of an electron. • The atom must also contain positive particles that balance exactly the negative charge carried by particles that we now call electrons.

8. Voltage source Thomson’s Experiment - + • Cathode rays = radiation produced when high voltage is applied across the tube. • The voltage causes negative particles to move from the negative electrode (cathode) to the positive electrode (anode).

9. The Electron (Cathode Rays) • Streams of negatively charged particles were found to emanate from cathode tubes, causing fluorescence. • J. J. Thomson is credited with their discovery (1897).

10. The Electron

11. If no Magnetic and Electric field applied: The cathode rays pass in straight direction. • Applying EF: The path of the cathode rays altered toward the +ve plate. • Applying MF: The cathode rays can be deflected in opposite direction. • By balancing both MF and EF: The rays pass in straight direction. Deflection of electron depends on three factors: 1) Strength of electric or magnetic field. 2) Size of negative charge. 3)Mass of the electron. • In 1897 Thomson determined the charge-to-mass ratio of an electron. Charge/Mass = 1.76 108 C/g. • C is a symbol for coulomb (SI of electric charge). • Either the charge or the mass of an electron would yield the other.

12. - Oil Drop Experiment (Millikan, 1868–1953):Applied a voltage to oppose the downward fall of charged drops and suspend them. • Voltage on plates place 1.602176 x 10-19 C of charge on each oil drop = electron charge.

13. Radioactivity • Radioactivity is the spontaneous emission of radiation by an atom. • It was first observed by Henri Becquerel. • Marie and Pierre Curie also studied it. Method: • A radioactive substance is placed in a lead shield containing a small hole so that a beam of radiation is emitted from the shield. • The radiation is passed between two electrically charged plates and detected.

14. Three types of radiation were discovered by Ernest Rutherford: •  particles: helium nucleus (+2 charge, large mass) •  particles: high speed electron (-ve) •  rays: high energy light, similar to X-rays (no charge). • The mass of an  -particle is 7300 times that of the electron • (similar to X-rays)

15. The Atom, circa 1900 • Early model: the “plum pudding” model. • Thompson: proposed a positive sphere of matter with negative electrons imbedded in it.

16. Rutherford’s Experiment:(Discovery Of Neucleus) • Used uranium to produce alpha particles • Aimed alpha particles at gold foil by drilling hole in lead block • Since the mass is evenly distributed in gold atoms alpha particles should go straight through. • Used gold foil because it could be made atoms thin

17. +

18. Explanation: • Atom is mostly empty. • Small dense, positive piece at center in the nucleus. • Alpha particles are deflected by it if they get close enough. • Proton (p) has opposite (+) charge of electron (-) • Mass of p is 1840 times the mass of e-(1.67 x 10-24 g). • Protons were discovered by Rutherford in 1919. • Neutrons were discovered by James Chadwick in 1932. +

19. Since some particles were deflected at large angles, Thomson’s model could not be correct. This led to the nuclear view of the atom.

20. Subatomic Particles • Protons and electrons are the only particles that have a charge. • Protons and neutrons have essentially the same mass. • The mass of an electron is so small we ignore it.

21. The angstrom is a convenient non-SI unit of length used to denote atomic dimensions. 1 Å = 1 x 10 –10 m Density of nucleus:1013–1014 g/cm3 A matchbox full of material of such density would weigh over 2.5 billion tons!

22. SAMPLE EXERCISE 2.1 • The diameter of a US dime is 17.9 mm, and the diameter of a silver atom is 2.88 Å . How many silver atoms could be arranged side by side across the diameter of a dime? Conversion Factors: 1 Å  10 -10 m 1 Ag atom  2.88 Å

23. Practice Exercise The diameter of a carbon atom is 1.54 Å. (a) Express this diameter in picometers. (b) How many carbon atoms could be aligned side by side in a straight line across the width of a pencil line that is 0.20 mm wide? Answer: (a) 154 pm, (b) 1.3 ×106C atoms

24. Atomic Numbers, Mass numbers and Isotopes Atomic Number (Z): Number of protons in an atom’s nucleus. Equivalent to the number of electrons around an atom’s nucleus Mass Number (A): The sum of the number of protons and the number of neutrons in an atom’s nucleus Isotope: Atoms with identical atomic numbers but different mass numbers

25. All atoms of the same element have the same number of protons: The atomic number (Z)

26. Isotopes are atoms of the same element with different masses. Isotopes have different numbers of neutrons. Isotopes have different numbers of neutrons, but the same number of protons. 11 6 12 6 13 6 14 6 C C C C Isotopes C11 C  12 C  13 C  14

27. 14 12 C C 6 6 Isotopes carbon-12 mass number 6 protons 6 electrons 6 neutrons atomic number carbon-14 mass number 6 protons 6 electrons 8 neutrons atomic number

28. SAMPLE EXERCISE 2.2 • How many protons, neutrons, and electrons are in (a) an atom of 197Au: From PT: Z = 79 P= 79 e = 79 n = 197-79 =118 (b) an atom of strontium-90? From PT: Z = 38 P= 38 e = 38 n = 90-38 = 52

29. Practice Exercise: How many protons, neutrons, and electrons are in (a) a 138Ba atom, (b) an atom of phosphorus-31? Answer: (a) 56 protons, 56 electrons, and 82 neutrons; (b) 15 protons, 15 electrons, and 16 neutrons.

30. Atomic Weights: The Atomic Mass Scale : Early: It was related to H mass. • Consider 100 g of water: • Upon decomposition 11.1 g of hydrogen and 88.9 g of oxygen are produced. • The mass ratio of O to H in water is 88.9/11.1 = 8. • Therefore, the mass of O is 2 x 8 = 16 times the mass of H. • If H has a mass of 1, then O has a relative mass of 16. • We can measure atomic masses using a mass spectrometer. 1H mass = 1.6735 x 10–24 g 16O mass = 2.6560 x10–23 g.

31. Atomic mass units (amu): are convenient units to use when dealing with extremely small masses of individual atoms. • The amu is 1/12 the mass of one 12C atom. 1 amu = 1.66054 x 10–24 g 1 g = 6.02214 x1023amu • By definition, the mass of 12C is exactly 12 amu. • Now all the present atoms are assigned according to C-12 isotopes • A 24Mg atom has a mass approximately twice that of the 12C atom, so its mass is 24 u. • A 4He atom has a mass approximately 1/3 that of the 12C atom, so its mass is 4 u. • 1H atom has a mass of 1.0078 amu. • 16O atom has a mass of 15.9949 amu.

32. Mass Spectrophotometer Atomic and molecular masses can be measured with great accuracy with a mass spectrometer.

33. Atomic weight • Most elements occur in nature as mixtures of isotopes. • We can determine the average atomic mass of an element, usually called the element’s atomic weight. • We average the masses of isotopes to give average atomic masses.

34. Example: Naturally occurring C consists of: Atomic mass: 12.0 amu 13.00335 amu Abundance: 98.93 % 1.07 % • The average mass of C is: (0.9893)(12 amu) + (0.0107)(13.00335 amu) = 12.01 amu. • Atomic weights are listed on the periodic table • Because in the real world we use large amounts of atoms and molecules, we use average masses in calculations. 13 C 12 C 6 6

35. HW: Chlorine has two naturally occurring isotopes: Cl-35 with an isotopic mass of 34.969 amu, and Cl-37 with an isotopic mass of 36.966 amu. The atomic weight of chlorine is 35.5 . What is the relative abundances of the two isotopes? Atomic mass 34.969 amu 36.966 amu Abundance: X 100-X The solve for X,,,,, 37 35 Cl Cl 17 17

36. Elements in Periodic Tables

37. Sample Exercise 2.3 Magnesium has three isotopes, with mass numbers 24, 25, and 26. Write the complete chemical symbol (superscript and subscript) for each of them. (b) How many neutrons are in an atom of each isotope? The numbers of neutrons in an atom of each isotope are therefore 12, 13, and 14, respectively. 24 25 26 Mg Mg Mg 12 12 12

38. Practice Exercise Give the complete chemical symbol for the atom that contains 82 protons, 82 electrons, and 126 neutrons.

39. Practice Exercise Three isotopes of silicon occur in nature: 28Si (92.23%), which has an atomic mass of 27.97693 amu; 29Si (4.68%), which has an atomic mass of 28.97649 amu; and 30Si (3.09%), which has an atomic mass of 29.97377 amu. Calculate the atomic weight of silicon. Answer: 28.09 amu

40. Periodic Table • It is a systematic catalog of the elements. • Elements are arranged in order of atomic number.

41. Periods: The rows on the periodic chart. • Columns: The groups on the periodic chart. Elements in the same group have similar chemical properties.

42. Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of properties and reactivity.

43. General Arrangement • Metallic elements, or metals: Are located on the left-hand side of the periodic table (most of the elements are metals). Properties: Shiny luster, malleable, ductile, and lustrous and are good thermal and electrical conductors, solids at RT (except Hg ). • Nonmetallic elements, or nonmetals: Are located in the top right-hand side of the periodic table. Properties: at RT brittle solids (C), liquids (Br), gas (Ne), dull in appearance, and do not conduct heat or electricity well, . • Metalloids:Are located at the interface between the metals and nonmetals (dteplike) except Al, Po, At. Elements with properties similar to both metals and nonmetals These include the elements B, Si, Ge, As, Sb and Te

44. Names of Some Groups in the Periodic Table Halogens: salt formers Chacogenes: ore formers These five groups are known by their names.

46. Important Families of Elements Representative Elements: The elements in the A groups (1,2, 13-18). Transition Metals: The elements in B groups (3-12). Inner Transition Metals: The two rows of metals (lanthanides and actinides) set at the bottom of the periodic table.

47. Sample Exercise 2.5 Which two of the following elements would you expect to show the greatest similarity in chemical and physical properties: B, Ca, F, He, Mg, P? Solution: Ca and Mg should be most alike because they are in the same group (2A, the alkaline earth metals).

48. Practice Exercise Locate Na (sodium) and Br (bromine) on the periodic table. Give the atomic number of each, and label each a metal, metalloid, or nonmetal.

49. 2.6 Molecules and Molecular Compounds • Only Nobel gases are found in nature as isolated atoms. • Most matter is composed of molecules or ions. • A molecule:consists of two or more atoms bound tightly together. Types: 1) Molecular Elements: Same two or more atoms combined with each others. • Diatomic molecules: made up of two same kind of atoms. Examples: N2, O2, F2, Cl2 , , Br2,I2 • Polyatomic molecules: made up of more than two atoms of same atoms. Examples: S8, P4, O3

50. H2 H2O NH3 CH4 Allotropes: Different forms of an element, which have different chemical formulas Allotropes differ in their chemical and physical properties. Examples: Ozone (O3) and “normal” Oxygen (O2) C (diamond) and C (Graphite), C60 2) Molecular compounds: Composed of molecules which contain more than one type of nonmetallic atoms. Examples: - Diatomic: HCl, CO. - Polyatomic: H2SO4,CO2, H2O2, NH3 Molecules