John Dalton’s Atomic Theory – 1803

1. Atoms, Molecules, and IonsThe Evolution of the Atomic Model (from John Dalton to the Modern Theory – just the big kahunas!) • John Dalton’s Atomic Theory – 1803 Dalton stated a group of assumptions to explain the nature and behavior of chemical systems. These became known as Dalton’s Atomic Theory and he proposed this the year 1803.

2. Dalton’s Atomic Theory Four Assumptions: • All substances are composed of small, dense particles called ATOMS. • Atoms of a given substance are identical in mass, size and shape. • An atom is the smallest part of an element that enters into a chemical reaction. • Molecules are produced by a combination of atoms.

3. Dalton’s Atomic Theory • Dalton’s Atomic Theory is often called the Billiard Ball Model. His theory, however, fails to explain many different types of behavior in chemical reactions. • It was through Dalton’s work that evidence of the following two subatomic particles were discovered: ELECTRON: negatively charged particles. PROTON: positively charged particles.

4. Diagram of Dalton’s Concept of the Atom “Billiard Ball Model”

5. 2. Thomson’s Concept of Atoms • JJ Thomson, early 1900’s • Thomson proposed a “new and improved” model using Dalton’s theory as a foundation (remember – people thought Dalton was crazy when he said in the early 1800’s that everything was made of atoms!!) • Thomson proposed that atoms consist of a solid bulk of positive charge with electrons dispersed throughout. His model is known as the Plum Pudding Model.

6. Diagram of Thomson’s Atom “Plum Pudding Model”

7. 3. Ernest Rutherford - 1911 • Rutherford discovered that the positive charge (the proton) and the mass was concentrated in the center of the atom (called the NUCLEUS). • He postulated that the electrons were moving at fast speeds around the nucleus but were contained by a certain boundary.

8. 3. Diagram of Rutherford’s Atom “Empty Space Model”

9. 4. Neil Bohr’s Concept of the Atom - 1913 • Bohr proposed that electrons are arranged in definite energy levels(shells) and follow a definite orbit. • Bohr’s concept or model is known as the “satellite” or “solar system” model of the atom .

10. Diagram of Bohr’s Atom • Solar system model

11. 5. Modern Concept of the Atom – 1920’s to present • The modern theory states that the electrons have wave-like properties as they orbit around the central nucleus. The + charged nucleus is surrounded by electrons with definite energy levels (called orbitals). • The paths of the e- are described in terms of the probability of being found in certain regions. The e- do not follow a prescribed path. • The modern concept of the atom is known as the “wave-mechanical” model of the atom.

12. Diagram of the Modern Atom “wave mechanical” model

13. Evolution of the atomic model

14. Chemical reactions are represented by both words and symbols •WORD: Zinc and sulfur yields zinc sulfide • EQUATION

15. Law of Conservation of Mass • Or the Law of Conservation of Matter • The law states that in ordinary chemical reactions, the mass of the system remains constant. Zn + S → ZnS 65.4g + 32.1 g = 97.5 g

16. Law of Conservation of Energy • The heat lost by the system (reaction) is equal to the heat gained by the surroundings (or, in ordinary chemical reactions, the energy of the system remains constant). ENDOTHERMIC: heat energy is absorbed during a chemical reaction (it gets colder) EXOTHERMIC: heat energy is lost during a chemical reaction (it gets hotter!) 95% of all reactions are exothermic!!

17. Heat + nitrogen + oxygen → nitric oxide • q + N2 + O2 → 2NO (q is one of the symbols used for heat or energy) If q is written on the left side, this indicates the reaction is endothermic (since energy is being absorbed). If q is written on the right side of the equation, the reaction is exothermic (since heat is being released to the surroundings)

18. Law of Definite Composition • When elements combine and form specific compounds, they do so in definite proportions by mass. Zn + S → ZnS 65.4g + 32.1 g = 97.5 g Zinc and sulfur will always combine in a definite fixed ratio of 65.4 parts to 32.1 parts by mass. If Zn or S were present in any other ratio, the one in excess would remain unchanged or unused. The excess would remain unreacted!!

19. Law of Multiple Proportions • When two elements combine and form more than one compound, the masses of one element that combine with a fixed mass of the other are in the ratio of small, whole numbers. • Look at the formulas of C to O in your notes: notice that in none of those formulas do you see C1.2O2.67 or C.98O3.11 • Formulas are always WHOLE NUMBERS!

20. Gay-Lussac’s Law of Combining Volumes • Gases react chemically with a volume of small, whole numbers!! • 1 volume H2 + 1 volume Cl2 → 2 volumes HCl • S + O2 → SO2 1 L + 1 L = 1 L • 2H2 + O2 → 2H2O 2L + 1 L = 2 Liters (it’s all about the small, whole numbers baby!)

21. Discovery of the Electron • JJ Thomson studied electrical discharges in partially, evacuated tubes called cathode –ray tubes (does anyone know the connection to our TV’s?) NOTE: this is the after picture!!!

22. More Cathode Ray fun!!! • Since the ray was attracted to the positive electrode, Thomson called the stream of particles electrons. • Thomson also determined the charge-to-mass ratio of the electron to be: e/m = -1.76 x 108 Coulombs/gram (where e = charge of the e- in Coulombs and m = mass in grams) WHY is the e/m ratio important? It showed that the electron has a NEGATIVE CHARGE!

23. Radioactivity • The French scientist Henri Becquerel found that a piece of Uranium produced its image on a photographic plate. So, as a third arm was growing out of his ribs, he figured out that this U had some weird energy coming from it (kind of like a chicken patty sandwich) • RADIOACTIVITY: the spontaneous emission of radiation (particles with lots of energy!!)

24. Three Types of Radiation

25. Discovery of the Nucleus • ERNEST RUTHERFORD – most famous for his discovery of the nucleus. • Rutherford used a radioactive source that emitted alpha (α)particles (these were his “bullets”), a piece of VERY THIN gold foil, a fluorescent screen, and a lead block. • His experiment would be similar to if you took your genuine Red Ryder BB gun and shot BB’s at a brick wall. What would you expect to happen?

26. Diagram of Rutherford’s Experiment

27. Rutherford’s Interpretation

28. Rutherford’s 3 Assumptions • The slightly deflected α particle had a close encounter with the “positive center” of the atom. • Most of the atom is empty space (because most of the α particles passed through) • The alpha particles that were completely deflected hit head on with the nucleus (because likes repel – both were + charged)

29. Major Subatomic Particles

30. A,Z,X Method (atomic number, atomic weight, etc.) MASS NUMBER (or atomic weight) – the Total number of protons and neutrons In the atom. SYMBOL ATOMIC NUMBER –the number of protons (and also = to the number of electrons if the atom is neutral!)

31. What if it is an ION ? • An ion is an atom with a net positive or a net negative charge. +1 A positive ion means that the element has lost electrons (in this case 1 electron). Na now has one more proton than electrons. A negative ion means that element has gained electrons (it now has more electrons than protons).

32. Isotopes • An isotopes are atoms with the same number of protons but different numbers of neutrons. • Examples of the isotopes of Hydrogen: • protium deuterium tritium • p+ = p+ = p+ = • no = no = no =

33. Write the symbol of the element which has 20 protons and a mass number of 40, and then write the formula of an isotope of this element.

34. Formulas …can you say pay attention to this! • CHEMICAL FORMULA: The symbol for the elements are used to indicate the types of atoms present and the subscripts are used to show the relative numbers of atoms. • Example of a chemical formula: CO2

35. STRUCTURAL FORMULAS • A formula showing the individual bonds (using lines to show the bonds). • H2O CH4

36. IONS • ION: an atom with a positive or negative charge. • CATION: a positive ion Na+1 Cu+2 Al+3 •ANION: a negative ion C2 H3 O2-1 BO3-3 SO4-2

37. Introduction to the fabulous PERIODIC TABLE!!! • Identify the following sections on a blank periodic table: • Metals nonmetals noble gases • Hydrogen alkali metals alkali earth metal • Halogens metalloids transition met. • Al family Carbon family Nitrogen family • Oxygen family Groups IA-VIIIA • Rare Earth elements • Lanthanide Series Actinide series

38. Four Classifications of Elements • METALS a. shiny luster b. good conductors; poor insulators c. malleable – can be hammered into thin sheets d. ductile – can be drawn into a thin wire e. All are solids except Ga and Hg

39. Properties of Metal (con’d) f. all metals have 1 – 3 electrons in the outer shell g. all metals lose electrons during chemical change.

40. 2. Nonmetals a. very brittle but pretty colors b. poor conductors; good insulators c. nonmetals are solids, liquids and gases d. all nonmetals have 5 – 7 e- in the outer shell e. nonmetals gain e- in chemical reactions

41. 3. Metalloids • Also called semimetals • Metalloids have properties of both metals and nonmetals.

42. 4. Noble or Inert Gases a. all are gases (no way!) b. the noble gases form NO compounds c. they are unreactive – they have 8 electrons in the outer shell (except Helium which only has and only needs 2 electrons)

43. SPECIFIC TYPES OF METALS 1. ALKALI METALS – very active metals that form ions with a +1 charge. Group IA. 2. ALKALI EARTH METALS – reactive (but not as much as Group IA) metals that form +2 ions. These are the Group IIA. These are often called the Fireworks Metals!!

44. Specific Types of Nonmetals • HALOGENS – Group VIIA. The word Halogen means “very active nonmetal”; all halogens react with metals to form salts containing ions with a -1 charge. • NOBLE OR INERT GASES – Group VIIIA ; nonreactive gases/elements

45. Naming Compounds BINARY IONIC COMPOUNDS (FORMULAS) • Binary compounds are composed of 2 elements • The components of a binary ionic compound are a monoatomiccationand a monoatomic anion. (what does monoatomic mean?) • Binary compounds end in the suffix –ide • Ionic compounds are electrically neutral.

46. Writing Formulas e. In writing a formula, we must exactly balance the positive charge of the cation with the negative charge of the anion (the net charge should be 0). f. We use the Crisscross (makes you want to “…jump, jump”) method” Ie. Potassium chloride

47. Writing Binary Ionic Formulas • Calcium bromide • Iron III oxide • Calcium sulfide

48. 2. Naming Binary Ionic Compounds • Binary ionic compounds are named by writing the name of the cation followed by the anion (ending in –ide). • When the cation has more than one possible ionic charge, it is important to use the Roman numeral.. • DON’T overkill the use of the Roman numerals. The representative (Group A) metals only have one charge so a Roman numeral is NOT needed.

49. Naming Binary Ionic Compounds • AlI3 • FeO • Cu2S • CaSe

50. 3. Ternary Ionic Compounds • Ternary ionic compounds contains atoms of three or more different elements. • Ternary ionic compounds usually contain one or more polyatomic ions. • First, write down the symbol of the ions • Second, “crisscross”(without jumping) the charges. • Third, use parenthesis whenever a polyatomic ion needs to be taken 2, 3 or 4 times.