Presentation Transcript

1. Atoms: The Building Blocks of Matter

   Chapter 3

2. Section 1 The particle theory of matter was supported as early as 400 b.c. by certain Greek thinkers, such as Democritus. He called nature’s basic particle an atom, based on the Greek word meaning “indivisible.” Aristotle was part of the generation that succeeded Democritus. His ideas had a lasting impact on Western civilization, and he did not believe in atoms. He thought that all matter was continuous, and his opinion was accepted for nearly 2000 years. Neither the view of Aristotle nor that of Democritus was supported by experimental evidence, so each remained speculation until the eighteenth century. Then scientists began to gather evidence favoring the atomic theory of matter.
3. Foundations of Atomic Theory Nearly all chemists in late 1700s accepted the definition of an element as a substance that cannot be broken down further Knew about chemical reactions Great disagreement as to whether elements always combine in the same ratio when forming a specific compound
4. Law of Conservation of Mass With the help of improved balances, investigators could accurately measure the masses of the elements and compounds they were studying This lead to discovery of several basic laws Law of conservation of mass states that mass is neither destroyed nor created during ordinary chemical reactions or physical changes
5. Law of Definite Proportions Was quickly followed by the law of definite proportions A chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound Each of the salt crystals shown here contains exactly 39.34% sodium and 60.66% chlorine by mass.
6. Law of Multiple Proportions Two elements sometimes combine to form more than one compound For example, the elements carbon and oxygen form two compounds, carbon dioxide and carbon monoxide Consider samples of each of these compounds, each containing 1.0 g of carbon In carbon dioxide, 2.66 g of oxygen combine with 1.0 g of carbon In carbon monoxide, 1.33 g of oxygen combine with 1.0 g of carbon The ratio of the masses of oxygen in these two compounds is exactly 2.66 to 1.33, or 2 to 1
7. 1808 John Dalton Proposed an explanation for the law of conservation of mass, the law of definite proportions, and the law of multiple proportions He reasoned that elements were composed of atoms and that only whole numbers of atoms can combine to form compounds His theory can be summed up by the following statements
8. Dalton’s Atomic Theory All matter is composed of extremely small particles called atoms. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties. Atoms cannot be divided, created or destroyed. Atoms of different elements combine in simple whole-number ratios to form chemical compounds. In chemical reactions, atoms are combined, separated, or rearranged.
9. Conservation of Mass Explained Chemical reactions involve the combination, separation, or rearrangement of atoms During these processes atoms are not divided, created, or destroyed
10. Definite Proportions Explained A given chemical compound is always made of the same combination of atoms
11. Multiple Proportions Explained Example of carbon oxides, the ratio of 2-to-1 ratio of oxygen masses results because carbon dioxide always contains twice as many atoms of oxygen (per atom of carbon) as does carbon monoxide
12. Modern Atomic Theory Dalton turned Democritus’sidea into a scientific theory which was testable Not all parts of his theory have been proven correct Ex. We know atoms are divisible into even smaller particles We know an element can have atoms with different masses
13. Section 2 The Structure of the Atom
14. All atoms consist of two regions Nucleus very small region located near the center of an atom In the nucleus there is at least one positively charged particle called the proton Usually at least one neutral particle called the neutron Surrounding the nucleus is a region occupied by negatively charged particles called electrons
15. Discovery of the Electron Resulted from investigations into the relationship between electricity and matter Late 1800s, many experiments were performed  electric current was passed through different gases at low pressures These experiments were carried out in glass tubes known as cathode-ray tubes
16. Cathode Rays and Electrons Investigators noticed that when current was passed through a cathode-ray tube, the opposite end of the tube glowed Hypothesized that the glow was caused by a stream of particles, which they called a cathode ray The ray traveled from the cathode to the anode when current was passed through the tube
17. Observations 1. An object placed between the cathode and the opposite end of the tube cast a shadow on the glass 2. A paddle wheel placed on rails between the electrodes rolled along the rails from the cathode toward the anode These facts supported the existence of a cathode ray The paddlewheel experiment showed that a cathode ray had sufficient mass to set the wheel in motion
18. 3. Cathode rays were deflected by a magnetic field in the same way as a wire carrying electric current, which was known to have a negative charge 4. The rays were deflected away from a negatively charged object
19. A paddle wheel placed in the path of the cathode ray moves away from the cathode and toward the anode. The movement of the wheel led scientists to conclude that cathode rays have mass.
20. Observations led to the hypothesis that the particles that compose cathode rays are negatively charged Strongly supported by a series of experiments carried out in 1897 by the English physicist Joseph John Thomson
21. He was able to measure the ratio of the charge of cathode-ray particles to their mass He found that this ratio was always the same, regardless of the metal used to make the cathode or the nature of the gas inside the cathode-ray tube Thomson concluded that all cathode rays are composed of identical negatively charged particles, which were later named electrons
22. Charge and Mass of the Electron Thomson’s experiment revealed that the electron has a very large charge for its tiny mass Mass of the electron is about one two-thousandth the mass of the simplest type of hydrogen atom (the smallest atom known) Since then found that the electron has a mass of 9.109 × 10−31 kg, or 1/1837 the mass of the hydrogen atom
23. Thomson’s Atom
24. 1909 Robert A. Millikan Confirmed that the electron carries a negative electric charge Because cathode rays have identical properties regardless of the element used to produce them, it was concluded that electrons are present in atoms of all elements Cathode-ray experiments provided evidence that atoms are divisible and that one of the atom’s basic constituents is the negatively charged electron
25. Two more inferences about the atom 1. Because atoms are electrically neutral, they must contain a positive charge to balance the negative electrons 2. Because electrons have so much less mass than atoms, atoms must contain other particles that account for most of their mass
26. Discovery of Atomic Nucleus 1911 by New Zealander Ernest Rutherford and his associates Hans Geiger and Ernest Marsden Bombarded a thin, gold foil with fast-moving alpha particles (positively charged particles with about four times the mass of a hydrogen atom) Assume mass and charge were uniformly distributed throughout atoms of gold foil (from Thomson’s model of the atom) Expected alpha particles to pass through with only slight deflection
27. What Really Happened… Most particles passed with only slight deflection However, 1/8,000 were found to have a wide deflection Rutherford explained later it was “as if you have fired a 15-inch artillery shell at a piece of tissue paper and it came back and hit you.”
28. Explanation After 2 years, Rutherford finally came up with an explanation The rebounded alpha particles must have experienced some powerful force within the atom The source of this force must occupy a very small amount of space because so few of the total number of alpha particles had been affected by it The force must be caused by a very densely packed bundle of matter with a positive electric charge Rutherford called this positive bundle of matter the nucleus
29. Rutherford had discovered that the volume of a nucleus was very small compared with the total volume of an atom If the nucleus were the size of a marble, then the size of the atom would be about the size of a football field But where were the electrons? Rutherford suggested that the electrons surrounded the positively charged nucleus like planets around the sun He could not explain, however, what kept the electrons in motion around the nucleus
30. Rutherford’s Atom
31. Composition of Atomic Nucleus Except hydrogen, all atomic nuclei made of two kinds of particles Protons Neutrons Protons = positive Neutrons = neutral Electrons = negative Atoms are electrically neutral, so number of protons and electrons IS ALWAYS THE SAME
32. The nuclei of atoms of different elements differ in the number of protons they contain and therefore in the amount of positive charge they possess Thus, the number of protons in an atom’s nucleus determines that atom’s identity
33. Forces in the Nucleus Usually, particles that have the same electric charge repel one another Would expect a nucleus with more than one proton to be unstable When two protons are extremely close to each other, there is a strong attraction between them Nuclear forces short-range proton-neutron, proton-proton, and neutron-neutron forces that hold the nuclear particles together
34. Section 3 – Counting Atoms
35. Atomic Number Atoms of different elements have different numbers of protons Atomic number (Z) number of protons in the nucleus of each atom of that element Shown on periodic table Atomic number identifies an element 3 Li Lithium 6.941 [He]2s1
36. Isotopes Hydrogen and other atoms can contain different numbers of neutrons Isotope atoms of same element that have different masses
37. Isotopes of Hydrogen
38. Mass Number Mass number total number of protons and neutrons in the nucleus of an isotope
39. Naming Isotopes Isotopes are usually identified by specifying their mass number There are two methods for specifying isotopes Mass number is written with a hyphen after the name of the element Tritium is written as hydrogen-3 Refer to this method as hyphen notation Shows the composition of a nucleus as the isotope’s nuclear symbol Uranium-235 is written as 23592U The superscript indicates the mass number and the subscript indicates the atomic number
40. Sample Problem How many protons, electrons, and neutrons are there in an atom of chlorine-37? atomic number = number of protons = number of electrons mass number = number of neutrons + number of protons mass number of chlorine-37 − atomic number of chlorine = number of neutrons in chlorine-37 mass number − atomic number = 37 (protons plus neutrons) − 17 protons = 20 neutrons
41. Practice Problems How many protons, electrons, and neutrons are in an atom of bromine-80? Answer 35 protons, 35 electrons, 45 neutrons 2. Write the nuclear symbol for carbon-13. Answer 136C 3. Write the hyphen notation for the element that contains 15 electrons and 15 neutrons. Answer phosphorus-30
42. Relative Atomic Masses Because atoms are so small, scientists use a standard to control the units of atomic mass  carbon-12 Randomly assigned a mass of exactly 12 atomic mass units (amu) 1 amu is exactly 1/12 the mass of a carbon-12 atom Atomic mass of any atom determined by comparing it with mass of C-12 Ex. H  1/12 C-12, so 1 amu
43. Average Atomic Masses Average atomic mass the weighted average of the atomic masses of the naturally occurring isotopes of an element Ex. You have a box containing 2 size of marbles 25% are 2.00g each 75% are 3.00g each 25 x 2.00g = 50g 75 x 3.00g = 225g 50g + 225g = 275g Divide by 100 – average marble mass = 2.75g
44. Method 1 25 x 2.00g = 50g 75 x 3.00g = 225g 50g + 225g = 275g Divide by 100 – average marble mass = 2.75g Method 2 25% = 0.25 75% = 0.75 (2.00g x 0.25) + (3.00g x 0.75) = 2.75g
45. Calculating Average Atomic Mass Copper 69.17% Cu-63 – 62.929599 amu 30.83% Cu-65 – 64.927793 amu (0.6917 x 62.929599) + (0.3083 x 64.927793) = 63.55 amu
46. Relating Mass to Numbers of Atoms The relative atomic mass scale makes it possible to know how many atoms of an element are present in a sample of the element with a measurable mass Three very important concepts provide the basis for relating masses in grams to numbers of atoms The mole Avogadro’s number Molar mass
47. The Mole SI unit for an amount of substance (like 1 dozen = 12) Mole (mol)  amount of a substance that contains as many particles as there are atoms in exactly 12g of C-12
48. Avogadro’s Number Avogadro’s number the number of particles in exactly one mole of a pure substance 6.022 x 1023 How big is that? If 5 billion people worked to count the atoms in one mole of an element, and if each person counted continuously at a rate of one atom per second, it would take about 4 million years for all the atoms to be counted
49. Molar Mass Molar mass the mass of one mole of a pure substance Written in unit g/mol Found on periodic table (atomic mass) Ex. Molar mass of H = 1.008 g/mol
50. Gram/Mole Conversions WHEN YOU SEE GRAMS, USE GRAMS/MOLE!
51. Some concepts to consider: 1. 1.000 mole is 6.022x 1023 2. 1 atom has a mass equal to its atomic mass in AMU 3. 1 mole of atoms has a mass equal to its atomic mass in g 4. The molecular mass of a molecule equals the sum of the masses of its atoms 5. 1 molecule has a mass equal to the molecular mass in AMU 6. 1 mole of molecules has a mass equal to the molecular mass in g
52. molar mass of element = x = x Amount of element in moles Mass of element in grams Number of atoms of element = x 6.022x10E23 = x
53. Example How many grams of helium are there in 2 moles of helium? 2.00 mol He x = ? g He 2.00 mol He x = 8.00 g He
54. Practice Problems What is the mass in grams of 3.50 mol of the element copper, Cu? 222 g Cu What is the mass in grams of 2.25 mol of the element iron, Fe? 126 g Fe What is the mass in grams of 0.375 mol of the element potassium, K? 14.7 g K What is the mass in grams of 0.0135 mol of the element sodium, Na? 0.310 g Na What is the mass in grams of 16.3 mol of the element nickel, Ni? 957 g Ni
55. 1. A chemist produced 11.9 g of aluminum, Al. How many moles of aluminum were produced? 0.441 mol Al 2. How many moles of calcium, Ca, are in 5.00 g of calcium? 0.125 mol Ca 3. How many moles of gold,Au, are in 3.60 × 10−10 g of gold? 1.83 × 10−12 mol Au
56. Conversions with Avogadro’s Number How many moles of silver, Ag, are in 3.01 x 1023 atoms of silver? Given: 3.01 × 1023 atoms of Ag Unknown: amount of Ag in moles Ag atoms × = moles Ag 3.01x1023atomsAg x =0.500 mol Ag
57. Practice problems 1. How many moles of lead, Pb, are in 1.50 × 1012 atoms of lead? 2.49 × 10−12 mol Pb 2. How many moles of tin, Sn, are in 2500 atoms of tin? 4.2 × 10−21 mol Sn 3. How many atoms of aluminum, Al, are in 2.75 mol of aluminum? 1.66 × 1024 atoms Al
58. 1. What is the mass in grams of 7.5 × 1015 atoms of nickel, Ni? 7.3 × 10−7 g Ni 2. How many atoms of sulfur, S, are in 4.00 g of sulfur? 7.51 × 1022 atoms S 3. What mass of gold,Au, contains the same number of atoms as 9.0 g of aluminum,Al? 66 g Au