1. Chapter 2/Unit 2: Matter is Made of Atoms Notes Chemistry CPA

2. Unit Goals After learning this chapter you will be able to: Identify “who” discovered each part of the atom Describe the historic and present models of the atom Label and describe the function of each part of the atom (nucleus, proton, neutron, electrons) Define and identify an isotope of any element Calculate the average atomic mass of elements as listed on the periodic table Describe the connection between light waves and electron energy

3. Atomic Theory Timeline

4. Atomic Models This model of the atom may look familiar to you. This is the Bohr model. In this model, the nucleus is orbited by electrons, which are in different energy levels. A model uses familiar ideas to explain unfamiliar facts observed in nature. A model can be changed as new information is collected.

5. The atomic model has changed throughout the centuries, starting in 400 BC, when it looked like a billiard ball →

6. Who are these men? In this lesson, we’ll learn about the men whose quests for knowledge about the fundamental nature of the universe helped define our views.

7. Democritus This is the Greek philosopher Democritus who began the search for a description of matter more than 2400 years ago. – He asked: Could matter be divided into smaller and smaller pieces forever, – Or was there a limit to the number of times a piece of matter could be divided? 400 BC

8. Atomos His theory: Matter could not be divided into smaller and smaller pieces forever, eventually the smallest possible piece would be obtained. This piece would be indivisible. He named the smallest piece of matter “atomos,” meaning “not to be cut.”

9. This theory was ignored and forgotten for more than 2000 years!

10. Why? The popular philosophers of the time, Aristotle and Plato, had a more respected, (and ultimately wrong) theory. Aristotle and Plato favored the earth, fire, air and water approach to the nature of matter. Their ideas were most believed because of their popularity as philosophers. The atomos idea was buried for approximately 2000 years.

11. Dalton’s Model In the early 1800s, the English Chemist John Dalton performed a number of experiments that eventually led to the acceptance of the idea of atoms.

12. Dalton’s Theory (The Four Postulates) 1.He deduced that all elements are composed of atoms. Atoms are indivisible and indestructible particles. 2.Atoms of the same element are exactly alike. 3.Atoms of different elements are different. 4.Compounds are formed by the joining of atoms of two or more elements.

13. . This theory became one of the foundations of modern chemistry.

14. Thomson’s Plum Pudding Model In 1897, the English scientist J.J. Thomson provided the first hint that an atom is made of even smaller particles.

15. Thomson Model He proposed a model of the atom that is sometimes called the “Plum Pudding” model. Atoms were made from a positively charged substance with negatively charged electrons scattered about, like raisins in a pudding.

16. Thomson Model Thomson studied the passage of an electric current through a gas. As the current passed through the gas, it gave off rays of negatively charged particles.

17. Thomson Model This surprised Thomson, because the atoms of the gas were uncharged. Where had the negative charges come from? Where did they come from?

18. Thomson concluded that the negative charges came from within the atom. A particle smaller than an atom had to exist. The atom was divisible! Thomson called the negatively charged “corpuscles,” today known as electrons. Since the gas was known to be neutral, having no charge, he reasoned that there must be positively charged particles in the atom. But he could never find them.

19. Rutherford’s Gold Foil Experiment In 1908, the English physicist Ernest Rutherford was hard at work on an experiment that seemed to have little to do with unraveling the mysteries of the atomic structure.

20. Rutherford’s experiment Involved firing a stream of tiny positively charged particles at a thin sheet of gold foil (2000 atoms thick)

21. – Most of the positively charged “bullets” passed right through the gold atoms in the sheet of gold foil without changing course at all. – Some of the positively charged “bullets,” however, did bounce away from the gold sheet as if they had hit something solid. He knew that positive charges repel positive charges.

23. This could only mean that the gold atoms in the sheet were mostly open space. Atoms were not a pudding filled with a positively charged material. Rutherford concluded that an atom had a small, dense, positively charged center that repelled his positively charged “bullets.” He called the center of the atom the “nucleus” The nucleus is tiny compared to the atom as a whole.

24. Rutherford Rutherford reasoned that all of an atom’s positively charged particles were contained in the nucleus. The negatively charged particles were scattered outside the nucleus around the atom’s edge.

25. Bohr Model In 1913, the Danish scientist Niels Bohr proposed an improvement. In his model, he placed each electron in a specific energy level.

26. Bohr Model According to Bohr’s atomic model, electrons move in definite orbits around the nucleus, much like planets circle the sun. These orbits, or energy levels, are located at certain distances from the nucleus.

27. Wave Model

28. The Wave Model Today’s atomic model is based on the principles of wave mechanics. According to the theory of wave mechanics, electrons do not move about an atom in a definite path, like the planets around the sun.

29. The Wave Model In fact, it is impossible to determine the exact location of an electron. The probable location of an electron is based on how much energy the electron has. According to the modern atomic model, at atom has a small positively charged nucleus surrounded by a large region in which there are enough electrons to make an atom neutral.

30. Electron Cloud: A space in which electrons are likely to be found. Electrons whirl about the nucleus billions of times in one second They are not moving around in random patterns. Location of electrons depends upon how much energy the electron has.

31. Electron Cloud: Depending on their energy they are locked into a certain area in the cloud. Electrons with the lowest energy are found in the energy level closest to the nucleus Electrons with the highest energy are found in the outermost energy levels, farther from the nucleus.

32. IndivisibleElectronNucleusOrbit Electron Cloud Greek X Dalton X Thomson X Rutherford X X Bohr X X X Wave X X X

33. Atomic Structure

34. What’s in an Element? Each atom can be classified as one of the elements on the periodic table. Each element has the atomic number (which is the # of protons), the element symbol, the element name, and the average atomic mass.

35. The parts of an atom Parts of an atom: – Nucleus – Electron Cloud Subatomic particles: – Protons (in the nucleus) – Neutrons (in the nucleus) – Electrons (outside the nucleus in the electron cloud)

36. The Nucleus The “core” of the atom Protons and neutrons are found in the nucleus of an atom Contains 99.9% of the mass of the atom The MASS NUMBER – is the mass of the nucleus – a sum of the mass of the protons + the mass of the neutrons: Mass number = mass of protons + mass of neutrons

37. The Atomic Mass Unit In an atom, mass is measured in “atomic mass units”, or “amu”. Protons have a mass of 1 amu Neutrons have a mass of 1 amu Electrons have a mass of 1/1837 of an amu – so small, we don’t even count it! So…which particles give the atom it’s “mass”? ____________________

38. Protons Positively charged subatomic particle in an atom’s nucleus Gives the atom it’s “identity” Has a mass of 1 amu The number of protons is the atom’s ATOMIC NUMBER: #PROTONS = ATOMIC NUMBER

39. Neutrons A subatomic particle that has no charge Found in the nucleus of the atom Contributes to the mass of an atom Has a mass of 1 amu # of neutrons = mass number - #protons

40. Electrons Are found in the “electron cloud” which is the space outside of the nucleus. Are negatively charged particles Have almost no mass Are responsible for all chemical reactions and bonding that happens with other atoms, which means they give the atom its chemical properties In an electrically neutral atom: # protons = #electrons

41. More on the nucleus… …a nucleus generally has an equal number of protons and neutrons. When a nucleus has a different number of neutrons than protons, it is called an “isotope”.

42. Isotopes Atoms of the same element with different mass numbers. Different number of neutrons Mass # Atomic # Nuclear symbol: carbon-12; carbon-13Hyphen notation: carbon-12; carbon-13 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem 12 6 C

43. Calculating Mass # Practice RECALL: MASS # = # PROTONS + # NEUTRONS Mass numbers are always WHOLE NUMBERS (they aren’t the decimal numbers found on the periodic table) Practice: Element# protons# neutronsMass # Carbon6 8

44. Isotope Practice Chlorine-37 – atomic #: – mass #: – # of protons: – # of electrons: – # of neutrons: 17 37 17 20 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Cl 37 17 37 17 Cl

45. Isotopes

46. Isotope notation Element name – Mass number Examples: Carbon-14, Chlorine - 37 Practice: # protons# neutronsMass # Carbon - ____6 7 8

47. Shorthand Notation Mass number SYMBOL OF ELEMENT Atomic number Examples: Write the shorthand notation for the following Isotopes: Carbon-14Oxygen-18Magnesium-25

48. Average Atomic Mass The number on the periodic table is an average of all of the masses of all of the isotopes that exist in nature Based on percent abundance of an isotope’s occurrence in nature.

49. Example

50. Average Atomic Mass Equation Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Avg. Atomic Mass = (mass)(%) + (mass)(%) 100

51. Average Atomic Mass Practice Problem EX: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16 O, 0.04% 17 O, and 0.20% 18 O. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Avg. Atomic Mass = (16)(99.76) + (17)(0.04) + (18)(0.20) 100 = 16.00 amu

52. Average Atomic Mass EX: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine-35 and 2 are chlorine-37. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Avg. Atomic Mass = (35)(8) + (37)(2) 10 = 35.40 amu

53. Atomic Structure Recap: ATOMS ATOMS – Differ by number of protons IONS IONS – Differ by number of electrons ISOTOPES ISOTOPES – Differ by number of neutrons

55. Electrons are found in the electron cloud The cloud has regions of space called energy levels The first energy level holds 2 electrons The second energy level holds 8 electrons. The third energy level holds 18 electrons

56. Valence Electrons Are found furthest from the nucleus Dictate the physical and chemical properties of an element Use the periodic table to determine the number of valence electrons. All atoms want 8 valence electrons

57. Examples of electron filling

58. Lewis Dot Diagram A way to illustrate the number of valence electrons – Use one dot for each valence electron – Place the dot around each side of the symbol before pairing the electrons – The symbol represents the nucleus plus all the inner electrons for the element.

59. Lewis Dot Diagram Practice H O N F Ne

60. Electrons and Light Electrons are normally in the ground state When the atom is given energy the electrons move to the excited state. When the electrons lose this energy they fall back to the ground state and emit (give off) light. Each element has a unique emission spectrum

63. Electromagnetic Spectrum Electromagnetic Radiation – – A broad range of energetic emissions – made up of photons Photons – bundles of energy – Travel like waves – Move at the speed of light = 3.0 x 10 8 m/s – Electromagnetic waves do not require a medium to move

65. Parts of the wave Amplitude – the height of the wave Wavelength – the distance between the two successive waves Frequency – the number of waves that pass a given reference point per second

66. Wave Units Wavelength = lamda unit is nanometer Frequency = nu in units of 1/s or s -1

67. What is the difference between these waves?