Lecture 02 (Chapter 2) Atoms, Molecules, and Ions

1. Lecture 02 (Chapter 2)Atoms, Molecules, and Ions

2. Dalton’s Atomic Theory • John Dalton (1766-1844) used experimental data to propose one of the first theories to explain the properties of matter. • Dalton’s Atomic Theory, based on the idea that matter is discontinuous, is described by 4 postulates. • Dalton’s Atomic Theory helped to explain other observations, including the law of constant composition.

3. Dalton’s Atomic Theory 1. Matter is composed of atoms. An atom is the smallest unit of an element that has all the properties of that element. 2. An element is composed entirely of one type of atom. 3. A compound contains atoms of two or more different elements. The relative number of atoms of each element in a compound is always the same. 4. Atoms do not change identity in chemical reactions; only the way in which they are joined together changes.

4. Laws Related to Dalton’s Atomic Theory • Law of constant composition: All samples of a pure substance contain the same elements in the same proportions by mass. • From Dalton’s third assumption. • Law of multiple proportions: When the same elements form more than one compound, the masses of one element that combines with a fixed mass of a second element are in a ratio of small whole numbers. • From Dalton’s third assumption. • Law of Conservation of Mass: There is no detectable change in mass when a chemical reaction occurs. • From Dalton’s fourth assumption.

5. Atomic Composition and Structure • Experiments over many years showed that atoms are not simple particles, but are composed of the subatomic particles listed below: • Electrons • Protons • Neutrons

6. Cathode Rays; Discovery of electrons • Experiments in the late 1800’s found that the application of high voltage across a partially evacuated tube produced cathode rays.

7. Electrons • In 1897, J. J. Thomson demonstrated that cathode rays were negatively charged by applying magnetic and electric fields to cathode rays, which deflected their path, and confirmed their negative charge. • Cathode rays are electrons, negatively charged particles that are one of the components of an atom.

8. Millikan Oil Drop Experiment • 1913, Robert Millikan performed experiments to measure charge of electron. • Oil drops formed by injector were charged by capturing electrons. Based on rate the drops moved in absence or presence of electrical field, charge of the electron was calculated to be 1.602 x 10-19 coulombs.

9. The Nuclear Model of the Atom • 1899, Ernest Rutherford conducted experiments that involved passing alpha particles (from radioactive decay) through thin metal targets. • Observed that some particles deflected at very large angles.

10. The Nuclear Model of the Atom • Rutherford concluded that the results of the scattering experiment required that atoms consist of: • a nucleus that is very small compared to the atom, has a high positive charge and contains most of the mass of the atom. • the remainder of the space in an atom contains enough electrons to give a neutral atom.

11. Protons and Neutrons • Rutherford proposed that the hydrogen nucleus was a fundamental particle called the proton, which has apositive charge equal in magnitude to the negative charge of the electron. • Protons account for charge on nuclei of all atoms. • Proton mass (1.673 x 10-27 kg) is 1836 times that of the electron. • However, the number of protons in a nucleus accounted for half or less of the nuclear mass. • Scientists inferred there must be a massive, neutral particle also present in the nucleus. • This neutral particle is called the neutron; its mass is almost the same as that of the proton.

12. Particles in the Atom

13. Definitions • Atomic number (Z) is the number of protons in the nucleus of an atom (see periodic table). • Mass number (A) is the sum of the numbers of protons and neutrons in the nucleus (may indicate different isotopes, i.e., same Z/different A of same element).

14. Applications of Atomic and Mass Numbers • On the periodic table, the atomic number is written as a whole number above the symbol F. • In the written description, fluorine is said to have 9 protons (the atomic number is the number of protons). • In the symbol, the number 9 is written in the atomic number or Z (lower left) position. • Note: The periodic table does not show the mass number for an individual atom. It lists an average mass number for a collection of atoms! Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

15. Isotopes • Isotopes are atoms that have the same number of protons in the nucleus but different numbers of neutrons. That is, they have the same atomic number but different mass numbers. • Because they have the same number of protons in the nucleus, all isotopesof the same element have the same number of electrons outside the nucleus. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

16. Symbols of Isotopes • Oxygen also has three isotopes, containing 8, 9, and 10 neutrons respectively. The symbols are: • Since the value of Z, and the symbol, both identify the element, Z is often omitted from the symbol:

17. Elemental Notation Q: How to represent element X with 4 p+ and 5 n. Q: How to represent lead-208? Q: How many p+, e-, n? 82, 82, 126

18. Ions • In many chemical reactions, atoms gain or lose electrons, producing charged particles called ions. • A cation has a positive charge and forms when an atom loses one or more electrons (Na+, Pb2+). • An anion has a negative charge and forms when an atom gains one or more electrons (Cl-, O2-). • In many cases, the periodic table tells us whether an atom tends to lose or gain electrons, and how many. Common atomic ions you should know: • H+, Na+, K+, Mg2+, Ca2+, Fe2+, Fe3+, Ag1+, Pb2+, • N3-, P3-, O2-, S2-, F-, Cl-, Br-

20. Atomic Number and Atomic Mass • The atomic number and atomic mass for each element is given on the periodic table. 38 Atomic number Sr Atomic mass 87.62

21. Test Your Skill • Write the symbols for the particles containing:(a) 8 protons, 9 neutrons, 10 electrons(b) 13 protons, 14 neutrons, 13 electrons

22. Test Your Skill • Write the symbols for the particles containing:(a) 8 protons, 9 neutrons, 10 electrons(b) 13 protons, 14 neutrons, 13 electrons Answer: (a) (b)

23. Example: Components of Ions • Fill in the blanks.Symbol Atomic number ____Mass number ____Charge ____no. of protons ____no. of neutrons ____no. of electrons ____

24. The Atomic Mass Unit (u) • A relative mass scale has been established to express the masses of atoms. • The atomic mass unit (u) is 1/12 the mass of one 12C atom. Experimentally to three significant digits: 1 u = 1.66 x 10-27 kg • The masses of both the proton and the neutron are approximately 1 u. • 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.

25. Atomic Mass and Mass Number • Factors other than the mass of the protons and neutrons affect the mass of atoms, so the actual mass of atoms are not whole numbers. (24Mg = 23.98504 u; 4He = 4.002603 u) • When the accurate atomic mass of an atom is rounded to a whole number, it equals the mass number. • Why are the numbers different?

26. How Isotopes Determine Atomic Mass • The atomic massof an element is the relative mass of an average atom of the element expressed in atomic mass units. • Many elements have more than 1 isotope (e.g. – 12C, 13C, 14C). • Abundance of isotopes are not evenly distributed. • Weighted atomic mass of Carbon (12C, 13C only) = (0.98882*12u) + (0.01108 * 13.300335u) = 12.011u. C 12 12 C 6 AMU 12 u 13.300335 u 6 12.011 12.011

27. The Mass Spectrometer • A mass spectrometer is used to measure the relative masses of isotopes, neon in this figure.

28. Mass Spectrum of Neon The mass spectrum of neon shows that the element is a mixture of three isotopes.

29. Natural Distribution of Isotopes • About 75% of the elements occur in nature as mixtures of isotopes. • Usually, the relative abundance of isotopes of an element is the same throughout nature. • In all natural samples of Li, 7.42% of the atoms are 6Li and the remaining 92.58% are 7Li.

30. Determining Atomic Mass • A specific example of the use of the equation is shown below for the element boron that consists of 19.78% boron-10 with a mass of 10.01 u and 80.22% boron-11 with a mass of 11.01u. • This calculated value is seen to agree with the value given in the periodic table. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

31. Example: Calculating Atomic Mass • A mass spectrometer was used to determine that gallium is 60.11% 69Ga (isotopic mass = 68.9256 u) and 39.89% 71Ga (isotopic mass = 70.9247 u). Calculate the atomic mass of Ga.

32. The Periodic Table • Proposed independently by Dimitri Mendeleev and Lothar Meyer. • Periodic table: arranges the elements in rows that place elements with similar properties in the same column. • Period: a horizontal row • Group: a column - contains chemically similar elements http://web.sbu.edu/chemistry/wier/atoms/meyer.html; http://www.chemistry.co.nz/mendeleev.htm

33. The Periodic Table

34. Important Groups of Elements • Metal: a material that is shiny and is a good electrical conductor; metallic elements are on the center and left side of the periodic table. • Nonmetal: an element that is typically a nonconductor; nonmetals are in the top right part of the periodic table. • Metalloid: an element that has properties of both metals and nonmetals.

35. Important Groups 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.

36. Important Groups of Elements • Alkali Metals: soft, reactive metals in group 1A. • Alkaline Earth Metals: elements in group 2A. • Halogens (salt formers): reactive nonmetals in group 7A. • Noble Gases: the stable, largely inert, gases in group 8A.

37. Molecules • A molecule is a combination of atoms joined so strongly that they behave as a single particle. • The simplest molecules are diatomic - they contain two atoms.

38. Elements • If all the atoms in a molecule are the same, the substance is an element.

39. Molecules • If two or more elements form a molecule, it is a molecular compound.

40. Molecular Formulas • A molecular formula gives the number of every type of atom in the molecule. • The elements present in the molecule are identified by their symbols. • A subscript number follows each symbol, giving the number of atoms of that element present in the molecule; the subscript is omitted if only one atom of the element is present. • A structural formula shows how the atoms are connected in the molecule.

41. Molecular Formulas

42. Molecular Mass • The relative mass of a molecule in atomic mass unitsis called the molecular mass of the molecule. • Because molecules are made up of atoms, the molecular mass of a molecule is obtained by adding together the atomic massof all the atoms in the molecule. • The formula for a molecule of water is H2O. This means one molecule of water contains two atoms of hydrogen, H, and one atom of oxygen, O. The molecular mass of water is then the sum of two atomic massesof H and one atomic mass of O: • MM = 2(at. wt. H) + 1(at. wt. O) • MM = 2(1.01 u) + 1(16.00 u) = 18.02 u Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

43. Molecular Mass • The clear liquid is carbon disulfide, CS2. It is composed of carbon (left) and sulfur (right). What is the molecular weight for carbon disulfide? • Answer: MW = 1(atomic weight C) + 2(atomic weight S) 12.01 u + 2(32.07 u) = 76.15 u Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

44. Example: Calculate Molecular Mass Values • One substance present in smog is dinitrogen tetroxide (N2O4). Calculate its molecular mass. • What is the molecular mass of the fuel propane (C3H8 )?

45. Chemical Bonding • Ionic bond: Attractive force that holds ions of opposite charge together. • Involves transfer of e- from one component to the other. • Occurs between positively-charged metal (loses 1 or more e-) and non-metal atom or molecule (gains 1 or more e-). • Usually satisfies octet rule • Common to inorganic chemistry • Covalent bond: Formed by sharing of electrons. • Occurs between: • Two non-metals • Nonmetal and metalloid • Two metalloids • Usually satisfies octet rule • Common to organic chemistry

46. Ionic Compounds • An ionic compound is composed of cations and anions joined to form a neutral species. • Each cation is surrounded by several anions and vice versa.

47. Formulas of Ionic Compounds • The formula of an ionic compound is an empirical formula that uses the smallest whole number subscripts to express the relative numbers of ions. • The relative numbers of ions in the empirical formula balances the charges to zero. • The formula of sodium chloride is NaCl, because the 1+ ions have to be present in a 1:1 ratio. • The formula of sodium oxide is Na2O, because the charge of the Na+ and O2- ions balance to zero in a 2:1 ratio.

48. Formulas of Ionic Compounds • The position of an element in the periodic table can be used to determine the charges of some ions. • The metallic elements in Groups 1A, 2A, 3B, and Al (Group 3A) all form cations with a charge equal to the Group number. • The nonmetals in Groups 6A, 7A, and N in group 5A form anions with a charge of 2-, 1- and 3-, respectively.

49. Charges on Common Ions

50. Example: Ionic Compounds Formulas • Write the empirical formulas of the compound formed by(a) the cation of Ca and the anion of Br.(b) the cation of Al and the anion of O.