Atoms The Building Blocks of Matter Chapter 3

OBJECTIVES The Atom: Philosophy to Science 3.1 Explain the law of conservation of mass, the law of definite proportions, and the law of multiple proportions. Summarize the five essential points of Dalton’s atomic theory. Explain the relationship between Dalton’s atomic theory and the 3 Laws

Foundations of Atomic Theory The Philosophers 400 B.C. – Particle Theory of Matter Greek Philosophers Democritus “atmos” - indivisible Aristotle matter is continuous, did not believe in atoms Neither view was supported by experiments until the 18 th century

Alchemy 700 B.C. – 1700’s Transmutation – metals are made of varying proportions of sulfur and mercury - achieving the right combination would produce gold Some used alchemy to make medicines

Balance Scale – Quantitative Analysis

The Laws of Chemistry Measuring the masses of elements and compounds it was observed that when elements react to form compounds they combine in fixed proportions by mass. Three basic laws of chemistry were proposed.

1. Law of Conservation of Mass Antoine Lavoisier observed that the mass of the reactants before the reaction and the mass of the products after the reaction are the same. Mass is neither created nor destroyed during ordinary chemical reactions or physical changes.

2. The Law of Definite Proportions Regardless of the size of the sample or the source of a chemical compound, it is composed of a fixed ratio of elements by mass.

3. Law of Multiple Proportions If two or more different compounds are made 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.

Examples: P. 87 #2

Dalton’s Atomic Theory 1808, an English school teacher and chemist, John Dalton, proposed a scientific explanation of these three laws based upon the idea of atoms of elements.

Dalton’s Atomic Theory 1. All matter is made of tiny particles called atoms. 2. Atoms of the same element are identical, those of different elements are different. 3. Atoms cannot be subdivided, created, or destroyed. 4. Atoms of different elements combine in small whole number ratios to form compounds 5. In chemical reactions, atoms are combined, separated or rearranged.

Dalton’s Atomic Theory How it Supported the laws: Conservation of mass - atoms are not created or destroyed Definite proportions – a given compound is always the same proportion of atoms Multiple proportions – in 2 different compounds made of the same elements the ratio of the second atoms to the first atoms is a small whole number

Modern Atomic Theory Dalton turned Democritus’ idea of the atom into a scientific theory that could be tested by experiment. Not all aspects of Dalton’s theory have been proven correct. The theory has been modified by new discoveries.

OBJECTIVES The Structure of the Atom 3.2 Define atom. Summarize the observed properties of cathode rays that led to the discovery of the electron. Summarize the experiment carried out by Rutherford and his co-workers that led to the discovery of the nucleus. List the properties of protons, neutrons, and electrons.

Atom The smallest particle of an element that retains the chemical properties of the element.

Discovery of the electron J. J. Thomson - English physicist. Used a cathode ray tube. It is a vacuum tube - all the air has been pumped out and replaced with a low pressure gas

Thomson’s Experiment Voltage source +- n An electric current was passed through the tube from the cathode (the negative electrode) to the anode (the positive electrode).

Voltage source Thomson’s Experiment n By adding an electromagnetic field he found that the moving particles were negative + -

Thomson’s Experiment Movement of a paddlewheel in the path of the electrodes led scientist to conclude that the rays have mass. By using the cathode ray tube, Thomson determined that electrons have a very high charge and a very low mass.

Thomson’s Model “plum pudding model” An atom made of negative particles surrounded by positive material with the mass and charges uniformly distributed

Millikan Oil Drop Experiment. In 1909, Robert Millikan, an American physicist, confirmed that the electron has the smallest possible negative charge and that all other negative charges are whole number multiples of the charge of the electron.

Rutherford’s experiment (1910) Ernest Rutherford - English physicist. Believed in the plum pudding model of the atom. Used radioactivity to test it Alpha particles - positively charged pieces given off by uranium Shot them at gold foil which can be made a few atoms thick

Lead block Uranium Gold Foil Florescent Screen

He expected the alpha particles to pass through without changing direction very much Because the positive charges were spread out evenly. Alone they were not enough to stop the alpha particles

He thought the mass was evenly distributed in the atom

What he got

Most of the particles passed through the foil indicating that the atom is mostly empty space. + Very few of the particles bounced back, but with great force,

What Rutherford concluded: + The atom is mostly empty space A small, very dense, positively charged core within the atom The nucleus

Protons and Neutrons In 1919 Rutherford discovered the proton. The neutron was discovered in 1932 by an English scientist, James Chadwick.

Composition of the Atom – Subatomic Particles Protons are subatomic particles located in the nucleus of the atom with high mass and a positive charge equal in magnitude to the negative charge of the electron. The nucleus also contains neutrons which are electrically neutral and have a mass ~ equal to a proton Electrons surround the nucleus in an electron cloud. They have very little mass and a negative charge Atoms are neutral due to the presence of equal numbers of protons and electrons.

Subatomic particles Electron Proton Neutron NameSymbolCharge Relative mass Actual mass (g) e-e- p+p+ n0n x x

Subatomic Particles NUCLEUS ELECTRONS PROTONS NEUTRONS NEGATIVE CHARGE POSITIVE CHARGE NEUTRAL CHARGE ATOM QUARKS

Homework P. 74 # 1-5

OBJECTIVES Counting Atoms 3.3 Explain what isotopes are. Define atomic number and mass number, and describe how they apply to isotopes. Given the identity of a nuclide,determine its number of protons,neutrons, and electrons. Define mole in terms of Avogadro’s number, and define molar mass. Solve problems involving mass in grams, amount in moles, and number of atoms of an element.

Atomic Number Atomic Number (Z) = number of protons in the nucleus of atom Atomic Number Identifies the element # of protons determines kind of atom # protons = # electrons in the neutral atom

Atomic Number – The identity 3 Li Lithium [He]2s1 The atomic number in this periodic-table entry reveals that an atom of lithium has three protons in its nucleus

Mass Number mass # = protons + neutrons always a whole number © Addison-Wesley Publishing Company, Inc.

Isotopes Atoms of the same element with different numbers of neutrons –so different mass numbers. Mass # Atomic # Nuclear symbol: Hyphen notation: carbon-12

Isotopes © Addison-Wesley Publishing Company, Inc.

Designating Isotopes Isotopes of Hydrogen and Helium – p. 77

Isotopes Chlorine-37 atomic #: mass #: # of protons: # of electrons: # of neutrons:

Symbols of Nuclides Find the # p + # n 0 # e - Atomic number Mass Number F 19 9

Symbols of Nuclides n Find n Find the # p+p+p+p+ # n0n0n0n0 # e-e-e-e- –Atomic –Atomic number –Mass –Mass Number Br 80 35

Nuclides n if an element has an atomic number of 34 and a mass number of 78 what is the –number of protons –number of neutrons –number of electrons –Complete symbol

Nuclides n if an element has 91 protons and 140 neutrons what is the –Atomic number –Mass number –number of electrons –Complete symbol

Nuclides n if an element has 78 electrons and 117 neutrons what is the –Atomic number –Mass number –number of protons –Complete symbol

Using the periodic table for nuclides How many protons, neutrons, and electrons are in an atom of Uranium-235 Uranium-238

Homework p. 85 # 2,3

Relative Atomic Mass The mass of an atom expressed in atomic mass units is called the atomic mass of the atom. 1 p= amu 1 n = amu 1 e - = amu © Addison-Wesley Publishing Company, Inc. atomic mass unit (amu or u) 1 amu= 1 / 12 the mass of a 12 C atom

Average Atomic Mass weighted average of all naturally occurring isotopes of an element Periodic Table shows the average Avg. Atomic Mass

Avg. Atomic Mass Average Atomic Mass 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 amu

Avg. Atomic Mass Average Atomic Mass Ex. copper consists of 69.17% copper-63, which has an atomic mass of u, and copper-65, which has an atomic mass of u amu

Magnesium has three isotopes % magnesium - 24 with a mass of amu, 10.00% magnesium - 25 with a mass of amu, magnesium - 26 with a mass of amu. What is the average atomic mass of magnesium?

The Mole Avogadro

What is the Mole? A counting number (like a dozen) Avogadro’s number 1 mol =  items A large amount!!!!

1 mole of pennies would cover the Earth 1/4 mile deep! 1 mole of basketballs would fill a bag the size of the earth!

Molar Mass Mass of 1 mole of an element or compound. Molar mass tells the … grams per mole (g/mol) Use average atomic mass on the periodic table – same # different unit

Molar Mass Examples carbon aluminum zinc g/mol g/mol g/mol

Molar Conversions molar mass (g/mol) MASS IN GRAMS MOLES NUMBER OF PARTICLES  (particles/mol) Mole

Molar Conversions – grams to mol How many moles of carbon are in 26.0 g of carbon? 26.0 g C 1 mol C g C = 2.16 mol C

Molar Conversions – mol to grams What is the mass in grams of 3.50 mol of Cu? 3.50 mol Cu g Cu 1 mol Cu = 222 g Cu

Molar Conversions Using Avogadro’s number How many atoms are in 2.50 moles of lead? 2.50 mol  atoms 1 mol = 1.51  atoms Pb

Molar Conversions Using Avogadro’s number How many moles of Ag are in 3.01  atoms of silver? 3.01  atoms 1 mol Ag = mol Ag  1023 atoms

Molar Conversions Find the mass of 2.1  atoms of Copper  atoms 1 mol  atoms = 222 g Cu g 1 mol