Chem131, Week 1 notes
Chem131, Week 1 notes CHEM131
Popular in Chemistry I - Fundamentals of General Chemistry
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This 5 page Class Notes was uploaded by Christina Notetaker on Monday September 5, 2016. The Class Notes belongs to CHEM131 at University of Maryland - College Park taught by John Ondov in Fall 2016. Since its upload, it has received 37 views. For similar materials see Chemistry I - Fundamentals of General Chemistry in Chemistry at University of Maryland - College Park.
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Date Created: 09/05/16
Chem 131 (8/29/20169/1/2016) Chapter 1: Atoms 1.1 chemistry study of matter; composition and properties; chemical and physical reactions; and energy associated with changes matter mass and volume substance specific instance of matter 1.2 Classifying matter state and composition state depends on particles’ relative positions and interaction strength (relative to temperature) ➢ Solids fixed volume + shape; high attractions between atoms ➢ Liquids fixed volume, not shape; weaker attractions b/w atoms ➢ Gases u nfixed volume + shape; weakest attractions b/w atoms composition types of substances making up samples of matter: chemical properties interact with substances (eg. solubility, flammability) physical properties measure and describe a state of matter (eg. density, mass, boiling point) ***change in structures leads to change in properties 1.3 SI base units ( “systeme internationale” or International System of Units) Quantity Unit Symbol length meter m mass kilogram kg time second s electric current ampere A amount of substance* mole mol temperature Kelvin K luminous intensity candela cd *chemical substance that contains ions, atoms, molecules, protons, etc. SI prefixes T G M k h da | d c m μ n p Tera giga mega kilo hecto deca | deci centi mili micro nano pico 1012 109 106 103 102 101 | 101 102 103 106 109 1012 The great monkey king has died | drinking chocolate milk, μoving nine palaces 1.4 (((Historical theories on matter by ancient philosophers and scientists))) 1.5 Laws Establishing Modern Atomic Theory 1789, law of conservation of mass matter is neither created nor destroyed 1797, law of definite proportions a sample has the same proportions as the elements that compose it regardless of how it is created 1804, law of multiple proportions two elements that are combined to form two different compounds produce two masses; are expressed in a small ratio of whole numbers 1808, atomic theory: 1. matter is composed of indivisible atoms 2. all atoms of the same element have the same mass and properties 3. atoms combine to form compounds 4. during chemical reactions atoms of an element change the way they are bound together and do not change into atoms of another element 1.6 Electron 1800s, Thomson observed properties of particles in cathode rays: they travel in straight lines; they are independent of the composition of materials of origin (the cathode); and they carry a negative electrical charge electrical charge results in attractive and repulsive forces, aka electrostatic forces electric field area around charged particle where forces exist Electron charge 1909, Millikan calculated negative charge by electric field strength and mass of floating oil drops 19 charge of electron (1.6 x 10 C) Thomson determined charge:mass ratio by measuring deflection required to halt freefall of oil 8 drops in magnetic field (1.76 x 10 C/g) [ charge * charge:mass ratio = mass of electron e] (1.6 x 10 C) x (g/1.76 x 10 C) = (9.10 x 10 g) mass of electron e (9.10 x 10 g) 28 1.7 Atom Structure Thomson proposed a plumpudding model: 1909, Rutherford performed the gold foil experiment and proposed the nuclear theory of the atom: 1. nucleus contains most of the atom’s mass and charge 2. electrons occupy most of the atom’s volume 3. number of electrons = number of protons = electrically neutral atom neutrons neutral particles located in nucleus; its mass ≈ proton; no electrical charge 1.8 Subatomic Particles Protons, Neutrons, Electrons proton charges are equal in magnitude but opposite in sign to electron charges number of protons defines element >atomic number # of protons, symbol Z isotopes atoms with same number of protons, but different number of neutrons natural abundance percentage of a given isotope eg. 90.48% neon atoms have 10 neutrons, 0.27% 11 neutrons, 9.25% 12 neutrons ***the greater the natural abundance, the fewer the number of neutrons mass number sum of neutrons + protons, symbol A >[ = protons + neutrons] Isotope symbol: number of electrons = number of protons = atomic number ions atoms become charged particles by losing or gaining electrons cations positively charged ions anions negatively charged ions ***the number of neutrons does not affect the atom’s size b/c it is too small within the nucleus 1.9 atomic mass the weighted average of masses with individual isotopes (amu) Eg. Chloride (Cl) atomic mass is 35.45amu and has isotopes Cl35 and Cl37 atomic mass = 0.7577 (34.97 amu) + 0.2423 (36.97 amu) = 35.45 amu Equation: atomic mass = Σ (fraction of isotope n) x (mass of isotope n) = (fraction of isotope 1 x mass of isotope 1) + (fraction of isotope 2 x mass of isotope 2) + (fraction of isotope 3 x mass of isotope 3) + etc... Or average weight = (A x atomi1 ass ) + (A x a1mic ma2 ) + etc... 2 A = (% abundance/ 100)
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