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CHEM 1031 (8)
Chapter 2

CHEM 1031 Chapter 2: Chapter 2 Atoms and Elements
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Department
Chemistry
Course
CHEM 1031
Professor
James D.Bloxton
Semester
Fall

Description
Chapter 2 Atoms and Elements Atoms are the smallest identifying unit of an element. There are about 91 different naturally occurring elements. Scientists have succeeded in making 20 synthetic elements. Antoine Lavoisier formulated the Law of Conservation of Mass. In a chemical reaction, matter is neither created nor destroyed. The total mass of the reactants will not change in the product. Joseph Proust formulated the Law of Definite Proportions. All samples of a given compound, regardless of their source or how they were prepared have the same proportions of their constituent elements. The law is sometimes called the law of constant composition. Mass Ratio= 16.0 𝑔 = 8.0 or 8.1 2.0 𝑔 𝐻 This ratio of water decomposing will have an oxygen and hydrogen product with a mass ratio that will hold for any sample of pure water regardless of its origin. John Dalton formulated the Law of Multiple Proportions. When elements (A &B) form two different compounds the masses of element B that combine with 1 gram of A can be expressed as a ratio of small whole numbers. The ratio of the masses of B that reacted with a fixed mass of A would always be small whole numbers. John Dalton also formulated the Atomic Theory. • Each Element is composed of tiny indestructible particles called atoms. • All atoms of a given element have the same mass and other properties that distinguish them from atoms of other elements. • Atoms combine in a simple, whole number ratios to form compounds. • Atoms of one element can’t change into atoms of another elements. In a chemical reaction, atoms only change the way that they are bound together with other atoms. The Discovery of the Electron: J.J Thomson preformed an experiment using cathode tubes showing that particles traveled in straight lines, theywereindependent of thecomposition material from thecathodeandtheycarried 8 a negative charge. He measured the electron’s negative charge to be -1.76 x 10 Coulombs per gram. Millikan’s Oil Drop Experiment: The Charge of the Electron: Millikan calculated the charge on oil droplets falling in an electric field. He found that it was always -1.60 x 10 -1C, the charge of a single electron. 𝑀𝑎𝑠𝑠 Charge x = Mass 𝐶ℎ𝑎𝑟𝑔𝑒 -19 𝑔 -28 -1.60 x 10 C x −1.76 x 108 = 9.10 x 10 g Rutherford’s Gold Foil Experiment: Rutherford tried to conform the plum pudding model. He used α particles, which proved it wrong instead. In the experiment Rutherford directed the positively charged α particles at an ultrathin sheet of gold foil. Majority of particles passed through the foil but some particles were deflected. He proposed the nuclear theory of the atom. • Most of the atom’s mass and all of its positive charge are contained in small core called the nucleus. • Most of the volume of the atom is empty space, throughout where tiny, negatively charged electrons are dispersed. • There are as many negatively charged electrons outside the nucleus as there as positively charged protons within the nucleus, so that the atom is electrically neutral. James Chadwick demonstrated that the unaccounted mass was due to neutrons, the neutral charged particles in the nucleus. Subatomic Particles: Protons, Neutrons, and Electrons in Atoms. Elements: Defined by Their Numbers of Protons: The number of protons in an element is the atomic number. It is given the symbol Z as well. Atomic Mass is just the average of the natural abundance of the isotopes. Isotopes: When the number of Neutrons Varies All atoms of a given element have the same number of protons but sometimes they don’t have the same number of neutrons. Isotopes- atoms with the same number of protons but different numbers of neutrons are called i
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