The Atomic Theory of Matter

1. The Atomic Theory of Matter • John Dalton: • Each element is composed of atoms • All atoms of an element are identical. • In chemical reactions, the atoms are not changed. • Compounds are formed when atoms of more than one element combine. • Dalton’s law of multiple proportions: When two elements form different compounds, the mass ratio of the elements in one compound is related to the mass ratio in the other by a small whole number.

2. Sodium solid chlorine gas exothermic reaction results in NaCl

3. The Discovery of Atomic Structure • The ancient Greeks were the first to postulate that matter consists of indivisible constituents. • Later scientists realized that the atom consisted of charged entities. • Cathode Rays and Electrons • A cathode ray tube (CRT) is a hollow vessel with an electrode at either end. • A high voltage is applied across the electrodes.

4. JJ at work • The voltage causes negative particles to move from the negative electrode to the positive electrode. • The path of the electrons can be altered by the presence of a magnetic field.

5. Cathode Rays and Electrons • The amount of deflection of the cathode rays depends on the applied magnetic and electric fields. • In turn, the amount of deflection also depends on the charge to mass ratio of the electron. • In 1897, Thomson determined the charge to mass ratio of an electron to be 1.76  108 C/g. • Goal: find the charge on the electron to determine its mass.

6. The Discovery of Atomic Structure Cathode Rays and Electrons

8. The Discovery of Atomic Structure • Cathode Rays and Electrons • Consider the following experiment: • Oil drops are sprayed above a positively charged plate containing a small hole. • As the oil drops fall through the hole, they are given a negative charge. • Gravity forces the drops downward. The applied electric field forces the drops upward. • When a drop is perfectly balanced, the weight of the drop is equal to the electrostatic force of attraction between the drop and the positive plate.

9. The Discovery of Atomic Structure Cathode Rays and Electrons

10. Charge and mass of electrons …. • Using this experiment, Millikan determined the charge on the electron to be 1.60  10-19 C. • Knowing the charge to mass ratio, 1.76  108 C/g, Millikan calculated the mass of the electron: 9.10  10-28 g. • With more accurate numbers, we get the mass of the electron to be 9.10939  10-28 g.

11. Radioactivity Products of decay Radioactivity

12. The Discovery of Atomic Structure • The Nuclear Atom • From the separation of radiation we conclude that the atom consists of neutral, positively, and negatively charged entities. • Thomson assumed all these charged species were found in a sphere.

13. Rutherford carried out the following experiment

14. Conclusion: Nuclear Atom • In order to get the majority of -particles through a piece of foil to be undeflected, • 1. the majority of the atom must consist of a low mass, • 2. diffuse negative charge - the electron. • To account for the small number of high deflections of the -particles • 3. the center or nucleus of the atom must consist of a dense positive charge.

16. Rutherford modified Thomson’s model as follows: • assume the atom is spherical but the positive charge must be located at the center, with a diffuse negative charge surrounding it.

17. Atomic Theory of Matter The nucleus of an atom is composed of two different kinds of particles: protons and neutrons. • Nuclear structure; Isotopes • An important property of the nucleus is its positive electric charge. • A proton is the nuclear particle having a positive charge equal to that of the electron’s (a “unit” charge) and a mass more than 1800 times that of the electron’s • The number of protons in the nucleus of an atom is referred to as its atomic number (Z).

18. An element is a substance whose atoms all have the same atomic number. • The neutronis a nuclear particle having a mass almost identical to that of a proton, but no electric charge. • The mass number is the total number of protons and neutrons in a nucleus. • A nuclide is an atom characterized by a definite atomic number and mass number. • The shorthand notation for a nuclide consists of its symbol with the atomic number as a subscript on the left and its mass number as a superscript on the left.

19. Isotopes are atoms whose nuclei have the same atomic number but different mass numbers; that is, the nuclei have the same number of protons but different numbers of neutrons. Chlorine, for example, exists as two isotopes: chlorine-35 and chlorine-37. The fractional abundance is the fraction of a sample of atoms that is composed of a particular isotope.

20. Mass SpectrophotometerGas Chromatography Mass Spectrometry: How Does It Work?

22. Atomic Weights Calculate the atomic weight of boron, B, from the following data: ISOTOPE ISOTOPIC MASS (amu) FRACTIONAL ABUNDANCE B-10 10.013 0.1978 B-11 11.009 0.8022

23. Atomic Weights Calculate the atomic weight of boron, B, from the following data: ISOTOPE ISOTOPIC MASS (amu) FRACTIONAL ABUNDANCE B-10 10.013 0.1978 B-11 11.009 0.8022 B-10: 10.013 x 0.1978 = 1.9805 B-11: 11.009 x 0.8022 = 8.8314 10.8119 = 10.812 amu ( = atomic wt.)

24. Atomic Weights • Dalton’s Relative Atomic Masses • Since Dalton could not weigh individual atoms, he devised experiments to measure their masses relative to the hydrogen atom. • Hydrogen was chosen as it was believed to be the lightest element. Daltons assigned hydrogen a mass of 1. • For example, he found that carbon weighed 12 times more than hydrogen. He therefore assigned carbon a mass of 12.

25. Atomic Weights • Dalton’s Relative Atomic Masses • Dalton’s atomic weight scale was eventually replaced in 1961, by the present carbon–12 mass scale. • One atomic mass unit (amu) is, therefore, a mass unit equal to exactly 1/12 the mass of a carbon–12 atom. • On this modern scale, the atomic weight of an element is the average atomic mass for the naturally occurring element, expressed in atomic mass units.

26. The Periodic Table

27. The Periodic Table • Some of the groups in the periodic table are given special names. • These names indicate the similarities between group members: • Group 1A: Alkali metals. • Group 2A: Alkaline earth metals. • Group 6A: Chalcogens. • Group 7A: Halogens. • Group 8A: Noble gases.

28. Molecules and Molecular Compounds • Molecules and Chemical Formulas • Molecules are assemblies of two or more atoms bonded together. • Each molecule has a chemical formula. • The chemical formula indicates • which atoms are found in the molecule, and • in what proportion they are found. • Compounds formed from molecules are molecular compounds. • Molecules that contain twoatoms bonded together are called diatomic molecules.

29. Diatomic molecules

30. Molecules and Molecular Compounds • Molecular and Empirical Formulas • Molecular formulas

31. Molecular and Empirical Formulas • Empirical formulas • give the relative numbers and types of atoms in a molecule. • the lowest whole number ratio of atoms or moles of atoms in a molecule • Examples: H2O, CO2, CO, CH4, HO, CH2. • 2 atoms of hydrogen / 1 atom of oxygen • 2 moles of hydrogen / I mole of oxygen

32. Chemical FormulasMolecular and Ionic Substances • The chemical formula of a substance is a notation using atomic symbols with subscripts to convey the relative proportions of atoms of the different elements in a substance. Consider the formula of aluminum oxide, Al2O3. This formula implies that the compound is composed of aluminum atoms and oxygen atoms in the ratio 2:3.

33. A molecule is a definite group of atoms that are chemically bonded together – that is, tightly connected by attractive forces. • Molecular substances • A molecular substance is a substance that is composed of molecules, all are nonmetals. • A molecular formula gives the exact number of atoms of elements in a molecule. • Structural formulas show how the atoms are bonded to one another in a molecule.

35. An ion is an electrically charged particle obtained from an atom or chemically bonded group of atoms by adding or removing electrons. • Anion  • Cation  • The formula of an ionic compound is written by giving the smallest possible whole-number ratio of different ions in the substance. • Empirical formula • The formula unit of the substance is the group of atoms or ions explicitly symbolized by its formula. • Ionic substances

36. Chemical compounds are classified as organic or inorganic. • Naming simple compounds • Organic compounds are compounds that contain carbon combined with other elements, such as hydrogen, oxygen, and nitrogen. • Inorganic compounds are compounds composed of elements other than carbon. • Ionic = inorganic • Most ionic compounds contain metal and nonmetal atoms; for example, NaCl. • You name an ionic compound by giving the name of the cation followed by the name of the anion.

37. Most of the main group metals form cations with the charge equal to their group number. • Rules for predicting charges on monatomic ions • Most transition elements form more than one ion, each with a different charge.

38. Ions and Ionic Compounds Predicting Ionic Charge

40. Chemical Substances; Formulas and Names • Rules for naming monatomic ions • Monatomic cations are named after the element. For example, Al3+ is called the aluminum ion. • If there is more than one cation of an element, a Roman numeral in parentheses denoting the charge on the ion is used. • usually transition elements. • The names of the monatomic anions use the stem name of the element followed by the suffix – ide. For example, Br- is called the bromide ion.

41. Naming Binary Compounds • NaF - • LiCl - • MgO - • CaS • AlP • KI • RbCl • CsBr • BaO • SrS

42. Chemical Substances; Formulas and Names • Polyatomic ions • A polyatomic ion is an ion consisting of two or more atoms covalently bonded together and carrying a net electric charge. • Cations formed from non-metals end in -ium. • Example: NH4+ ammonium ion.

43. NH4+ - Ammonium OH- - Hydroxide CN- - Cyanide SO42- - Sulfate ClO4- - Perchlorate O22- - Peroxide PO43- - Phosphate CO32- - Carbonate HCO3- - Bicarbonate Plus what is on your “stuff you should know for the AP test sheet” Ions You Should Know

44. Oxy anions • General rule: • If there are two oxyion forms of an anion • NO2 and NO3 • the one with the most oxygen ends in –ate • If that anion forms an acid the –ate is changed to –ic • Ic I ate it all! Refers to the concept that –ate has all of the oxygens (the most) • NO3 nitrate ion  HNO3 Nitric Acid • The lower oxygenated ion ends with –ite and forms an acid that ends with –ous • NO2 Nitrite ion  HNO2 Nitrous Acid

45. More than two forms of oxy ions? • If there are more than two forms then the prefixes • Hyper and Hypo are used • ClO4- HyperChlorate ion • ClO3- Chlorate ion • ClO2-Chlorite ion • ClO- Hypochlorite ion

48. Polyatomic anions containing oxygen with additional hydrogens • named by adding • hydrogen or bi- (one H), dihydrogen (two H), etc., to the name as follows: • CO32- is the carbonate anion • HCO3- is the hydrogen carbonate (or bicarbonate) anion. • H2PO4- is the dihydrogen phosphate anion. • NaHCO3 Sodium bicarbonate

49. Binary molecular compounds • Binary compounds composed of two nonmetals are usually molecular and are named using a prefix system. • The name of the compound has the elements in the order given in the formula. • You name the first element using the exact element name. • Name the second element by writing the stem name of the element with the suffix “–ide.” • If there is more than one atom of any given element, you add a prefix.

50. Here are some examples of prefix names for binary molecular compounds. • SF4 sulfur tetrafluoride • ClO2 chlorine dioxide • SF6 sulfur hexafluoride • Cl2O7 dichlorine heptoxide