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Chapter 2

McMurry and Fay Chemistry Chapter 2 A summary of Chapter 2 of the McMurry and Fay textbook required for the course. Great for studying for tests!

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Department
Chemistry
Course
CHM135H1
Professor
C.Scott Browning
Semester
Fall

Description
Chapter 2: Atoms, Molecules, and Ions 2.1 – Conservation of Mass and the Law of Definite Proportions - Robert Boyle: The first to define an element as a substance which could not be chemically broken down into anything simpler. - Joseph Priestley: Prepared and isolated oxygen gas by heating mercury(II) oxide. - Antoine Lavoisier: Discovered that the products of combustion are equal to the mass of the starting reactants. o Law of Conservation of Mass: Mass is neither created nor destroyed in chemical reactions. - Joseph Proust: o Law of Definite Proportions: Different samples of a pure chemical substance always contain the same proportion of elements by mass. o Elements combine in specific proportions and not in random proportions. 2.2 – Dalton’s Atomic Theory and the Law of Multiple Proportions - John Dalton proposed a new theory of matter: o Elements are made up of tiny particles called atoms. o Each element is characterized by the mass of its atoms. Atoms of the same elements have the same mass, but those of different elements have different masses. o Chemical combination of elements to make different substances occurs when atoms join together in small whole-number ratios. Fractional parts of atoms are not involved in chemical reactions. o Chemical reactions only rearrange the way that atoms are combined; the atoms themselves don’t change. Atoms are chemically indestructible. - Law of Multiple Proportions: Elements can combine in different ways to form different substances, whose mass ratios are small, whole-number multiples of each other. o Eg) Comparison of N:O ratios in NO and NO2 2.3 – The Structure of Atoms: Electrons - Cathode-ray tubes: o Sealed glass vessel with no air. o Two thin pieces of metal called electrodes are sealed inside. o When a high voltage is applied across the electrode, an electric current is sent from the negatively charged cathode to the positively charged anode. o If there is still air or gas in the tube, the current can be seen as a cathode ray. o If there is a hole at the end of the anode, a bright spot of light can be seen when the electric current hits. - J.J. Thomson proposed that cathode rays must consist of negatively charged particles known as electrons. o This was because the beam was produced at the negative electrode. o All metals must have electrons since they are used to make many electrodes. - Thomson believed that the deflection of an electron beam by a nearby magnetic or electric field depended on three factors: o The strength of the field: the greater the strength, the greater the deflection. o The size of the charge on the electron: greater size means greater interaction with the fields, which results in greater deflection. o The electron mass: the lighter the electron, the more easily it can be deflected. 8 - Thomson could determine the ratio of the electrons charge to its mass (1.758820 x 10 C/g). - R.A. Millikan determined the mass of an electron (m = 9.109382 x 10 ).8 2.4 – The Structure of Atoms: Protons and Neutrons - Matter is neutral overall, therefore if it could give of negatively charged particles, there must be positively charged particles as well. - Rutherford directed positive alpha particles toward a thin gold foil, and found that most went through, a small amount was deflected at an angle, and a few bounced back toward the source. o The atom is mostly empty space. o Must have a concentrated central core, known as the nucleus.  The nucleus contains most of the mass and positive charges. o Particles were reflected only if it encountered a large positive mass. - It was soon discovered that the nucleus was composed of both protons and neutrons. -24 o Protons: positively charged particle with a mass of 1.672622 x 10 g o Neutrons: carry no charge, and have a mass of 1.674927 x 10 g24  The number of neutrons is not related to the number of protons. 2.5 – Atomic Number - Elements are different because of the number of protons they are made up of. o The atomic number is the number of protons of the element, as well as the number of electrons. - Mass number: the sum of the protons and neutrons in an atom. - Atoms may have different masses depending on the number of neutrons that they have. - Isotopes: atoms with the same atomic number but a different mass number. o Mass number is the left superscript, and the atomic number is the left subscript. o The number of neutrons can be found by subtracting the number of protons from the atomic mass. o Most isotopes have similar behaviours since neutrons have very little effect on the atom’s chemical properties. 2.6 – Atomic Mass - Atomic mass unit (amu) is the unit used in biological work because the values are often much to small to be convenient. -24 o 1 amu = 1.660539 x 10 o The atomic mass of an atom is in atomic mass units. - Atomic mass: the weighted average of the isotopic masses of all its naturally occurring isotopes. o Atomic mass = (Mass of isotope A)(Abundance of isotope A) + (Mass of isotope B)(Abundance of isotope B)…etc 2.7 – Compounds and Mixtures - Matter is classified as either a pure substance or a mixture. - A pure substance can be an element or a chemical compound. o Chemical compound: pure substance formed when two or more elements combine to create a ne
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