Chapter 3 Atoms: The Building Blocks of Matter Honors Chemistry 412
Foundations of Atomic Theory Democritus Greek Philosopher 460-370 B.C. Stated Matter could be divided into smaller & smaller particles until it could no longer be divided. Called these Particles Atomos (indivisible) Not based on any physical evidence, only thought.
Chemical Reaction The transformation of a substance or substances into one or more new substances In the 1790 s, technology enabling more quantitative analysis let to the discovery of several basic laws
Law of Conservation of Mass Mass is neither created nor destroyed in any ordinary chemical reaction or physical change A + B AB + 1 a.m.u. + 3 a.m.u. 4 a.m.u.
Law of Definite Proportions A compound always contains the same elements in the same proportions by mass, regardless of the size of the sample OR the source of the compound. The mass ratio of A to B will always be the same! In this case 1:3 or 25% to 75% + 1 a.m.u. + 3 a.m.u. 4 a.m.u.
Law of Definite Proportions Calculations Hydrogen and Oxygen have a mass ratio of 1:16. What is the mass of oxygen needed to form with 14g of Hydrogen? What is the total mass of this compound? 1 16 = 14 x xx = 16 14 xx = 224 or 16 1 = x 14 MMMMMMMM oooo cccccccccccccccc = 14gg + 224gg = 238 gg
Law of Definite Proportions Calculations Magnesium and Oxygen have a mass ratio of 3:2. What is the mass of oxygen needed to form with 20g of Magnesium? What is the total mass of this compound? 3 2 = 20 xx xx = 13.3 13.3g +20g = 33.3g MgO
Law of Multiple Proportions If two or more different compounds are composed of the same two elements, then 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. CO 2 ; 1g to 2.66g CO; 1g to 1.33g Ratio of oxygen would 2:1.
John Dalton In 1808, Dalton proposed explanations for these laws If elements were made of atoms, then only whole numbers of atoms could combine to form compounds His entire theory is often summarized into 5 basic statements
Dalton s Atomic Theory 1. All matter is composed of extremely small particles called atoms. 2. **Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties 3. **Atoms cannot be subdivided, created, or destroyed. 4. Atoms of different elements combine in simple, wholenumber ratios to form chemical compounds. 5. In chemical reactions, atoms are combined, separated, or rearranged.
Modern Atomic Theory Dalton was able to transform the idea that Democritus had into a testable theory Not all of the aspects of Dalton s theory have proven to be 100% correct We now know that atoms are divisible through nuclear reactions We will see that atoms of the same element can have different masses
Modern Atomic Theory Atom The smallest particle of an element that retains the chemical properties of that element Two regions: Nucleus Very small region at the center of the atom Contains at least one proton (+) and usually one or more neutrons (0) Large region surrounding the nucleus Contains electrons (-)
Modern Atomic Theory How do we know all of this?
J.J. Thomson One of many scientists experimenting with a a Cathode Ray Tube (CRT) Electric current passed through low pressure gas in a glass tube. Produced a ray composed of particles that always traveled from the cathode (-) to the anode (+) Rays were deflected away from negatively charged objects
J.J. Thomson Was able to confirm the hypothesis that the particles in the cathode ray were negatively charged Was also able to calculate the ratio between the mass and the charge of the particles Concluded that the ray was composed of identical negatively charged particles which came to be known as electrons Credited with the plum pudding model of the atom
Robert Millikan Millikan s Oil-Drop Experiment Proved the mass and charge of all electrons are identical. By varying the charge on different drops, he noticed that the charge was always a multiple of -1.6 x 10-19 C.
Ernest Rutherford Gold Foil Experiment Carried out by Marsden and Geiger Proved the existence of the nucleus Bombarded thin sheets of metal with (+) charged alpha particles. Results of the Particle deflections Most traveled through Some Deflected in Both Directions Very Few Deflected Back
Rutherford vs. Thomson Plum Pudding Nuclear
Results of the Gold Foil Experiment Atoms Contain a Nucleus Central part of the atom. Very small and very massive = Very dense. Essentially all the mass of an atom Positively charged
E. Goldstein Contents of the Nucleus Used canal rays in a cathode ray tube to prove the existence of protons. Positively charged particles that had significant mass moved towards the cathode. Sir James Chadwick Neutrons, were neutrally charged particles found in the nucleus to add mass to the atom and to act as nuclear glue.
Summary of the Atom Nucleus Protons Positively charged particle found in the nucleus Gives mass to the atom. Neutrons Neutrally charged particle found in the nucleus. Gives mass to the atom and acts as nuclear glue Held together by nuclear forces short-range forces that hold the nucleus together; prot0n-proton, proton-neutron, and neutron-neutron Electrons Negatively charged particles that give the atom its size.
Properties of Subatomic Particles Particle Symbols Relative Charge Mass Number Relative Mass Actual Mass Electron e - -1 0 0.00055 u 9.11 x 10-28 g Proton p + +1 1 1.00728 u 1.67 x 10-24 g Neutron n 0 0 1 1.00866 u 1.68 x 10-24 g
Describing Atoms Atomic Number (Z) Number of protons contained in the nucleus of that element s atoms Equivalent to the number of electrons in a neutral atom of that element Found on the periodic table (whole number!)
Isotopes Atoms that have the same number of protons but differing numbers of neutrons AKA nuclides Does not affect the charge, but will affect the overall mass of the atom Most elements consist of a mixture of isotopes Hydrogen:
Identification of Isotopes Mass Number The total number of protons and neutrons in a particular isotope Not reported on the periodic table Atomic Symbol Notation: X = atomic symbol Mass Form: Element Name Mass # mass # atomic # X
Hydrogen s Isotopes 3 isotopes of hydrogen Protium (Hydrogen 1) Accounts for more than 99% of hydrogen on the Earth Deuterium (Hydrogen 2) Less than 1% of hydrogen atoms Tritium (Hydrogen 3) Radioactive, can be found in nature or prepared artificially
Isotope Form Practice Write the element uranium with a mass number of 235 in atomic and mass form: Atomic form: 235 92 U Mass Form: Uranium 235
Subatomic Particle Practice How many protons, neutrons and electrons do the following contain? 198 Carbon 13 Au 79 Carbon p + = 6 e - = 6 n o = 7 Gold p + = 79 e - = 79 n 0 = 119
25 28 25 53 25 Manganese-53 56 83 56 139 56 Barium-139 74 106 74 180 74 Tungstun-180 76 115 76 191 76 Osmium-191 30 33 30 63 30 Zinc-63 e - n 0 p + Mass # Atomic # Mass Form Symbolic Form Os 191 76 Zn 63 30 W 180 74 Ba 139 56 Mn 53 25
Measuring the Mass of Atoms Atomic Mass Units (amu) Relative scale used to measure the mass of atoms and subatomic particles Based upon the mass of the Carbon-12 isotope 6 protons + 6 neutrons given a mass of 12 amu Protons 1 amu Neutrons 1 amu Electrons 0 amu
Average Atomic Mass Weighted average of all isotopic masses of a given element Takes into account not only the different masses, but also the abundance of the isotope Calculation: AAM = [Fractional abundance Isotopic Mass] + [FA IM] + Percent abundance in decimal form (divide by 100)
Average Atomic Mass 28.09 amu
Calculate the average atomic mass: 58.73 amu
Identify the following element: AAM = 83.80 amu Krypton
Chemical Measurements Atomic Mass The relative mass of a single atom. Reported on the periodic table in amu For this class, these may be rounded to the whole number with the exception of: Cu = 63.5 amu Cl = 35.5 amu Formula Mass Sum of the atomic masses for each atom contained in a compound.
Calculating Formula Mass Determine the formula mass of sulfuric acid, H 2 SO 4. 2(1 amu) + 1(32 amu) +4(16 amu)= 98 amu
The Mole The number of atoms of an element equal to the number of atoms in exactly 12.0 g of carbon-12. Referred to as a counting number Dozen =? The mole is a specific number as well: 1 Mole = 6.02 x 10 23 atoms Known as Avogadro s Number
Molar Mass Mass in grams of 1 mole of any substance. Equivalent to the formula mass of a compound and to the atomic mass of an element. Units are expressed as grams per mole (g/mol) S = 32 g/mol H 2 SO 4 = 98 g/mol
Mole Equivalencies 1 mol = # grams from PT Use when you want to convert to/from mass 1 mol = 6.02 x 10 23 particles Use with atoms, molecules, or formula units 1 mol = 22.4 L of gas at STP (Molar Volume) STP = Standard Temperature and Pressure (0 C and 1 atmosphere of pressure) Use when you want to convert to/from volume of a gas
Mole Conversions We can use the mole unit as a stepping stone between other units: Volume (L) Molar Volume: 22.4L = 1 mole Moles Particles: Atoms Molecules Formula Units Mass (g) Avogadro s #: 6.02x10 23 particles = 1 mole Molar Mass: # g (from PT) = 1 mole
Using the Mole as a Conversion Factor How many grams would be contained in 6.75 mol Li 2 O? Molar mass = 2(7) + 1(16) = 30 g mol 6.75 molli 30 g Li 2O O 1 mol Li2O 2 g Li O = 202.50 2
Mole Conversions 7.9 x 10 24 formula units Na 2 CO 3 =? mol 24 2 3 7.9x10 f. units Na2CO3 = 23 6.02x10 f. units Na2CO 3 0.36 mol Al=? atoms 1.54 mol CO 2 =? L (@STP) 1 mol Na CO 23 6.02x10 atoms Al 0.36 mol Al 2.2x10 1 mol Al = 240 L N 2 =? mol (@STP) 240 L N 1 mol N = 10.71mol 2 2 N2 22.4 L N 2 1.54 mol CO 22.4 L CO = 23 13molNa atoms 2 2 34.50 L CO2 1 mol CO 2 Al 2 CO 3
2 Step Mole Conversions 4.0x10 How many grams would be contained in 4.0x10 24 molecules CH 4? 24 molec.ch 4 1molCH4 23 6.02x10 molec.ch 4 16 gch 1molCH 4 4 = 106 g CH 4 How many liters of CO 2 would be contained in 2.73 g of gas (@STP)? 2.73 g CO 2 1molCO 2 22.4 L CO 44gCO 2 1molCO2 1.39 L CO 2 2 =
Practice 2 Step Conversions 1.8 g NaHCO 3 =? formula units 1.8 g NaHCO 15.9 L N 3 1 mol NaHCO 84 g NaHCO 22 = 1.3x10 15.9 L N 2 O 5 =? Grams (@ STP) 2 O 5 1 mol N2O 22.4 L N2O 5 5 3 3 23 6.02x 10 f. units NaHCO 1 mol NaHCO3 f.units NaHCO 108 g N 1 mol N 2 2 O O 5 5 3 = 76.66 g N 2 O 5 3 9.3x10 9.3x10 23 molecules NO 2 =? L 23 molec. NO 2 1 mol NO2 23 6.02x10 molec. NO 2 22.4 L NO 1 mol NO 2 2 = 34.6 L NO 2