Atoms: The building blocks of matter Chapter 3

3.1 The Atom Democritus First to support the idea of the atom as the smallest particle of matter Circa 400 BC Competing with other philosophers of the time as well Idea falls flat because he doesn’t have any proof to support his claim

3.1 The Atom John Dalton (1808) Atomic Theory All matter is composed of tiny particles called atoms Atoms of a given element are identical in size, mass, and other properties; atoms of different elements are not the same Atoms cannot be subdivided, created, or destroyed Atoms of different elements combine in simple whole-number ratios to form compounds In reactions, atoms are combined, separated, or rearranged

3.1 The atom Dalton’s Atomic Theory Idea was supported because his points were backed by three scientific laws Law of Conservation of Mass Matter is neither created nor destroyed but conserved Law of Definite Proportions Compounds are contain the same elements in exactly the same proportions by mass regardless of the size of sample Law of Multiple Proportions Two compounds are made of the same two elements, then the ratio of the masses of the first element when combined with a set mass of the second will always be small whole numbers

3.1 The atom Additional links http://www.teachertube.com/video/law-of-conservation-of-energy-and-mass-chemis- 302255 http://www.teachertube.com/video/atom-theory-rap-268688

3.2 The structure of the Atom Dalton defined the atom as a small, indivisible particle that makes up all matter Scientist in the late 1800’s began to test the points of Dalton’s Theory Found that atoms were composed of smaller particles and that atoms of the same element might not always be identical Subatomic particles are protons, neutrons, and electrons

Discovery of electron JJ Thompson (1897) Discovery was made through an experiment known as the cathode ray experiment Changed model to the Plum Pudding Model of the atom

Cathode ray experiment http://www.teachertube.com/video/cathode-ray-tube-experiment-132346

Charge of electron Thompson test the tube with various elements and concluded that all contained electrons Robert Millikan (1909) Measured the charge of the electron Based upon this information, the charge to mass ratio of the electron was found to be about 1/1837 the mass of a hydrogen atom Since atoms are neutral, there must also be a positive charge Since the electron’s mass is so small, there must be another particle to make up its mass

Discovery of the nucleus Ernest Rutherford (1911) Discovery was made by an experiment known as the Gold Foil Experiment Assisted by Hans Geiger and Ernest Marsden Changed the model of the atom based upon his findings

Gold Foil Experiment https://youtu.be/_uEFKG122dA From the experiment, Rutherford was able to conclude that the majority of the mass of the atom is located in a dense, positively charged center Also concluded that the volume of the nucleus is relatively small compared to the volume of the atom

Composition of the nucleus Nucleus is composed of two subatomic particles Protons and neutrons These two subatomic particles have relatively equal masses Atoms have equal number of protons and electrons Forces in the Nucleus Short range force between proton-proton, proton-neutron, and neutron-neutron Known as Nuclear Force

3.3 Counting Atoms Atomic Number Isotopes Number that identifies the number of protons By default, also identifies the number of electrons for neutral atom Isotopes Atoms of the same element but have different number of neutrons

Counting atoms Mass number Atomic Mass Sum of the protons and neutrons for a single atom of an element Atomic Mass Weighted average of the isotopes for all the atoms of an element

Designating isotopes Methods for identifying isotopes Hyphen Notation Name of element – mass number Nuclear Symbol Mass number over atomic number then the symbol of the element Number of neutrons = mass number – atomic number

Calculating atomic mass Weighted average of the isotopes Mass number of each isotope multiplied by the percentage that the isotope exists in nature then added together EX. Copper-63 (69.15%) and Copper-65 (30.85%) 63 x .6915 = 43.5645 65 x .3085 = 20.0525 43.5645 + 20.0525 = 63.617 (63.62)

Relating mass to number of atoms Mole Amount of a substance that contains the same number of particles as Carbon-12 Avogadro’s number The number of particles in 1 mole which is equal to 6.02 x 1023 Molar Mass Number of grams equal to 1 mole of a substance For an element, this is equal to the atomic mass expressed in grams

Basic Conversions