Chapter 3 The Atom

The Atom Democritus (450 B.C.) proposed that all matter is made up of tiny, indivisible particles. (atomos) 3-2

Laws Law of Conservation of Mass: mass is neither created nor destroyed during ordinary chemical rxns or physical changes. 3-3

Laws Law of Definite Proportions: compounds contain the same elements in exactly the same proportions by mass regardless of the size of the sample or the source of the compound. 3-4

Laws Law of Multiple Proportions: If two or more different compounds are composed of the same two elements, the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers. 3-5

Daltons Atomic Theory In 1808 John Dalton, an English schoolteacher, came up with an atomic theory to explain these laws. Many of the tenets of his theory still hold true today. 3-6

Dalton’s Atomic Theory ~ Each element is composed of atoms. ~ Atoms of a given element are identical, and different than those of any other element. ~ A given compound forms by combination of two or more different atoms, always in the same relative numbers and kinds of atoms. ~ Atoms are neither created nor destroyed in any chemical reaction, only rearranged. 3-7

The Structure of the Atom Atom: the smallest particle of an element that retains the chemical identity of that element. 3-8

The Discovery of the Electron Thomson (1897) an English physicist who discovered electrons using his famous cathode ray tube experiment. He determined the charge to mass ratio. 3-9

The Discovery of the Electron

The Discovery of the Electron Millikan (1909) an American physicist who determined the charge of electrons using his famous oil drop experiment. The mass of an e- is approximately 1/2000 the mass of an atom. 3-11

The Discovery of the Electron

The Discovery of the Electron Based on these discoveries, two inferences were made: Because atoms are electrically neutral, they must contain a positive charge to balance the negative electrons. Because e- are so light in mass compared to atoms, atoms must contain other particles that account for most of it’s mass. 3-13

Plum Pudding Atomic Model

Discovery of the Nucleus 1911, Ernest Rutherford conducted experiments with radioactive materials that released only positively charged alpha particles…. 3-15

Discovery of the Nucleus

Discovery of the Nucleus Rutherford concluded atoms have a dense core with a positive charge. Rutherford’s Atomic Model:

Discovery of the Nucleus Rutherford’s student, Neils Bohr, came up with a way to explain the location of e- in the atom:

Discovery of the Nucleus The electron cloud is the current accepted model of the atom.

The Atom Except for H, all atomic nuclei contain protons and neutrons. A proton has a + charge equal in magnitude to the neg. charged e-. Atoms contain equal numbers of p+ and e-. Neutrons are electrically neutral. P+ and no have almost identical masses, electrons weigh 1836 times less. 3-20

The Atom Nuclear Forces: short range p+ to p+, p+ to no and no to no forces hold the nuclear particles together. Atomic Number (Z): the number of protons in each atom of a particular element. ☺The atomic number identifies the element! 3-21

Isotopes Isotopes: atoms of the same element that have different masses due to different numbers of neutrons. Mass Number: the total number of protons and neutrons that make up the nucleus of an isotope ~ Isotopes are written with the mass number written after the element name or symbol with a hyphen: ex. Uranium-235 or U-235 3-22

Isotopes Uranium-235 or U-235 Mass number – atomic number = number of neutrons 235 (protons + neutrons) – 92 protons = 143 neutrons This info could also be portrayed using a nuclear symbol: 235U 92

Isotopes ~ In nature, elements are almost always found as a mixture of isotopes ~ Isotopes have essentially identical chemical properties ~Isotopes with more neutrons have a higher mass and are often descibed as “heavy” 3-24

Isotopes Nuclide: a general term for a specific isotope of an element.

Atomic Mass Atomic Mass Unit (AMU): one amu is exactly 1/12 of the mass of a carbon-12 atom.(p+ = 1amu, no = 1 amu, e- = 0 amu) Average Atomic Mass: the weighted average of the atomic masses of the naturally occurring isotopes of an element. 3-26

Calculating the Average

Moles Mole: the amount of a substance that contains as many particles as there are atoms in exactly 12g of carbon-12.

Moles Avogadro’s Number: the number of particles in exactly one mole of a pure substance ~ 6.02214179 x 1023 (we’ll use 6.02 x 1023)

Moles 602,200,000,000,000,000,000,000 If you had 6.022 x 1023 pennies and gave away 1 million a day to every person on earth, it would take you 3000 years to distribute all your money!! 3-30

Molar Mass Molar Mass: the mass of one mole of a pure substance. ~ the mass of single atoms or molecules is measured in amu’s. The mass of a mole of the same substance is numerically the same, with the units g/mol. Ex. H20 = H x 2 = 1.01 x 2 = 2.02 + O x 1 = 16.00 x 1 = +16.00 18.02 3-31

Molar Mass 1) What is the molar mass of BaCl2? 2) What is the molar mass of NaI? 1) Ba = 1 x 137.33 g/mol = 137.33 Cl = 2 x 35.45 g/mol = 70.90 208.23 g/mol 2) Na = 1 x 22.99 g/mol = 22.99 I = 1 x 126.90 g/mol = 126.90 149.89 g/mol 3-32

Molar Mass This photograph shows one mole of solid (NaCl 58.44g/mol), liquid (H2O 18.02 g/mol), and gas (N2 28.02 g/mol). 3-33

Mass/Mole Conversions When given the number of moles, you can find the grams by: Moles x _g__ = grams mole Ex. 5.0 moles of H2O = X g 5.0 moles x 18.02g = 90. g H2O in 5.0 mol moles 3-34

Mass/Mole Conversions Moles x _g_ = moles mol Now try these problems: 3) 8.32 moles of barium chloride equals how many grams? 4) 20.1 moles of sulfur dioxide equals how many grams? 3-35

Mass/Mole Conversions 3) 8.32 moles of barium chloride equals how many grams? 8.32 moles BaCl2 x 208.23 g/mol = 1730 grams BaCl2 4) 20.1 moles of sulfur dioxide equals how many grams? 20.1 moles SO2 x 64.07 g/mol = 1290 g SO2 3-36

Mass/Mole Conversions

Mass/Mole Conversions When given the amount in grams, you can calculate the number of moles by: g x mol = moles g Ex. 11.2 g NaCl = X moles 11.2 g NaCl x 1.0 mol NaCl = 0.191 mols NaCl 58.44g 3-38

Mass/Mole Conversions g x mol = moles g Now try these problems: 5) 50.56 g of sodium chloride equals how many moles? 6) 329.8 g of ammonia equals how many moles? 3-39

Mass/Mole Conversions 5) 50.56 g of sodium chloride equals how many moles? 50.56 g NaCl x mole = 0.8652 moles NaCl 58.44g 6) 329.8 g of ammonia equals how many moles? 329.8 g NH3 x mole = 19.35 moles NH3 14.03 g 3-40

Particle/Mole Conversions You can also calculate between moles and number of particles: (1.0 moles = 6.02 x 1023 particles) To enter this number into your calculator, punch in 6.02 EE button (one time) 23. 3-41

Particle/Mole Conversions Ex. 2.59 moles of marble (CaCO3) contains how many molecules? 2.59 mol CaCO3 x 6.02 x 1023 molecules = 1.56 x1024 1.0 mol CaCO3 molecules *Particles can be molecules, atoms or formula units 7) How many molecules are in 5.0 moles of carbon dioxide? 5.0 mol CO2 x 6.02 x 1023 molecules = 1.0 mole CO2 3.01 x 1024 molecules CO2 3-42

Particle/Mass Conversions Ex. What is the mass of 3.25 x1023 molecules of nitrogen? 3.25 x1023 N2 x 1.0 mol x 28.02 g = 15.1 g N2 6.02 x 1023 1.0 mol 8) How many molecules are 57.36 g of NaCl? 57.36 g NaCl x 1.0 mol x 6.02 x 1023 = 5.911 x 1023 58.44 g 1.0 mol molecules 3-43

Ch. 3 The End!