Atoms: The Building Blocks of Matter Ch. 3

Early Atomic Theory Democritus: coined the word atom (“indivisible”) Thought nature’s particle was an atom Aristotle: didn’t believe atoms existed Thought all matter was continuous Neither theory was supported by scientific evidence. Evidence was not gathered until the 18th century.

Early Atomic Theory Late 1700s: all scientists believed in the concept of an element Elements combined to form compounds. Chemical reaction: combination of a substance or substances into one or more new substances Improved equipment helped chemists to more accurately measure experiments

Basic Laws of Chemistry Law of conservation of mass: mass is neither created nor destroyed during ordinary chemical rxns or physical changes Law of definite proportions: a chemical compound contains the same elements in exactly the same proportions by mass regardless of sample size Law of multiple proportions: If two or more different compounds are composed of the same two elements, then the ratio of masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers

Dalton’s Atomic Theory English Schoolteacher (1808) Proposed an explanation for the three laws described above Theory: All matter is composed of atoms. Atoms of a given element are identical. Atoms cannot be subdivided, 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.

How does Dalton’s law explain chemical laws? Law of conservation of mass: mass is not changed – it is rearranged, separated, or combined Law of definite proportions: a given chemical compound is always composed of the same combination of atoms Law of multiple proportions: ratio of different masses of the same compound

Modern Atomic Theory Not all Dalton’s statements have been proven to be correct. We can divide atoms. Atoms of a given element can have different masses (isotopes). Important remaining concepts: All matter is composed of atoms Atoms of any one element differ in properties from atoms of another element

Atomic Structure Atom: smallest particle of an element that retains the chemical properties of that element All atoms have two regions: Nucleus: Proton: positively charged particle Neutron: neutral particle Electrons: negatively charged particles

Discovery of the Electron Resulted from investigations between electricity and matter (late 1800s) Many experiments were conducted using cathode-ray tubes: pass electric current through a tube of gas Noticed that when electricity was passed through tube, opposite end glowed. They thought the glow was from a stream of particles (a cathode ray). They found these rays were deflected by negatively charged objects

Charge/Mass of Electron Robert A. Millikan: found the charge of an electron Mass of an electron: 9.109 x 10-31 kg. Two inferences made on atomic structure: Because atoms are electrically neutral, they must contain a positive charge to balance the negative electrons. Because electrons have so much less mass than atoms, atoms must contain other particles that account for most of their mass.

Atomic Number All atoms are composed of the same basic particles Atoms of different elements have different numbers of protons Atomic number: number of protons of each atom in that element

Isotopes Some elements have different numbers of neutrons. Example: Hydrogen Protium: 1 proton, 1 electron Deuterium: 1 proton, 1 neutron, 1 electron Tritium: 1 proton, 2 neutrons, 1 electron Isotopes: atoms of the same element that have different masses Mass number = number of protons + number of neutrons

Designating Isotopes Most isotopes don’t have unique names like hydrogen does Usually identified by their mass number Two ways to identify them Hyphen notation Hydrogen-3 = Tritium Number at the end represents the mass number Nuclear symbol: 168O = oxygen Top number is the mass number, bottom number is the atomic number.

Sample Problem A How many protons, electrons, and neutrons are there in an atom of chlorine-37?

Atomic Mass Masses of atoms are very small…kind of cumbersome to use grams Atomic mass unit: 1/12 the mass of a carbon-12 atom Example: Hydrogen – 1.08 amu (located on periodic table) Average atomic mass: weighted average of the atomic masses of the isotopes of an element

Average Atomic Mass Box with two sizes of marbles: 100 marbles 25% = 2.00g each 75% = 3.00g each 100 marbles 25 marbles x 2.00 g = 50 g 75 marbles x 3.00 g = 225 g 50+225 = 275 g 275/100 = 2.75 g (average mass of the marbles)

Calculating Average Atomic Mass Copper 69.15% copper-63 (amu = 62.929601) 30.85% copper 65 (amu = 64.927794) Average atomic mass (62.929601 x 0.6915) + (64.927794 X 0.3085) = 63.55 amu (always round atomic mass to 2 decimal places before using in any calculation)

Sample Problem (p.82 Table 4) Hydrogen, carbon, oxygen

The Mole Definition: the amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon-12. A counting unit – just like a dozen is. Avogadro’s number: the number of particles in exactly 1 mole of a pure substance (6.022 x 1023) 1 mole = 6.022 x 1023 atoms

Molar Mass The mass of one mole of pure substance Units: g/mol Numerically equivalent to atomic mass Look on the number underneath each element of the periodic table CONVERSIONS! Remember to put whatever you want to find on top.

Sample Problem B What is the mass in grams of 3.5 mol of the element copper (Cu)?

Sample Problem C A chemist produced 11.9 g of aluminum, Al. How many moles of aluminum were produced?

Sample Problem D How many moles of silver, Ag, are in 3.01 x 1023 atoms of silver?

Sample Problem E What is the mass in grams of 1.20 x 108 atoms of copper, Cu?