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CHMA10H3 (188)
Lecture

Chapter 1-4

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Department
Chemistry
Course
CHMA10H3
Professor
Heinz- Bernhard Kraatz
Semester
Fall

Description
Chapter1 Conservation of Mass- Antoinne Lavoisier Definite Proportions- Joseph Proust; elements composing a given compound always occur in fixed/definite proportions in all samples of the compound (vs components in a mixture possessing any proportions whatsoever);Ammonia contains 14g of nitrogen for every 3g of hydrogen, therefore mass ratio=14g/3g or 4.7 Law of multiple proportions- Dalton; when two elements form two different compounds the mass of element B that combines with 1g of element A can be expressed as a ratio of small whole numbers: Mass ratio of carbon dioxide is 2.67:1, therefore 2.67g of oxygen reacts with 1g of carbon to form carbon dioxide. Carbon monoxide has 1.33g of oxygen to every 1g of carbon so the ratio is 1.33:1. Mass of oxygen to 1g of carbon in carbon dioxide/ mass of oxygen in carbon monoxide=2. Atomic Theory- Dalton; each element is composed of atoms, all atoms of a given element have the same chemical/physical properties, atoms combine in simple whole number ratios to form compounds, atoms of one element cannot change into atoms of another element (atoms only change the way they are bound together in chemical reactions) J.J Thomson- using cathode ray tube observed that cathode rays travelled from cathode to anode. He found these particles travelled in straight lines, were independent of the composition of the cathode material and carried a negative electrical charge. Thomson measure the charge to mass ratio of the ray particles and found they were 2000 times lighter than hydrogen discovery of the electron Milikan- Oil drop; sprayed oil into fine droplets using atomizer, and negatively charged them by bombarding the chamber with electrons. Instead of falling under the influence of gravity, the drops were repelled by the negative plate discovery of the charge of an electron Mass Spectometry- atoms injected into apparatus, vaporized, ionized; fast moving electrons knock electrons off of atoms causing positively charged ions; charged ions pass through and deflected off a plate; amount of effort in deflection is used to calculate mass Rutherford- developed nuclear theory (gold foil) atom is mainly empty space Chapter3 - Properties of compounds are generally different than those of the elements that compose them - Electrostatic forces exist between charged particles - Ionic: metal and non - Covalent: two or more non (electron sharing to form molecular compound) Formulas - Empirical gives relative number of atoms in each compound (Benzene=CH) - Molecular gives actual number of atoms in a compound (Benzene=C6H6) - Structural uses lines to give sense of molecular geometry - Ball and stick/space filling are 3D - Pure substances are either elements (atomic or molecular) or compounds (ionic or molecular) - Atomic elements exist in nature with single atoms as their base (helium) - Molecular elements do no usually exist in nature with single atoms as their base. Most are diatomic (BrINHOFPSCSn) - Molecular compounds contain two or more covalently bonded non-metals (H2O) - Ionic compounds consist of cations and anions, whose basic unit is a formula unit (NaCl); formula reflects smallest whole number ratio - Switch from common to systematic names for universal info conveyed through names (muriatic vs hydrochloric acid) - ClO Hypochlorite ClO 2 Chlorite - ClO - Chlorate ClO 4 Perchlorate Ionic Compounds - Hydrated ionic compounds (water can usually be removed by heating) - Anhydrous salt: salt without water added - ½H2O- hemihydrate, 6H2O=hexahydrate - Smallest repeat unit is a water molecule Molecular compounds - Formula cannot be readily determined from its elements because the same elements can combine different compounds (carbon dioxide vs carbon monoxide) Molecular ions: F-, Br-, N- etc Naming Acids - Release H+ ions in water - HCl(g) is hydrogen chloride, HCl(aq) is hydrochloric acid - Binary acids only contain 2 elements, oxyacids contain oxygen Acids containing ions ending with ide often become Cl - chloride HCl hydrochloric acid - hydrofluoric F Fluoride HF acid S 2- Sulfide H S hydrosulfuric 2 acid Acids containing ions ending with ate usually become CH C32 - Acetate CH C3 H 2 acetic acid CO 2- carbonate H CO carbonic acid 3- 2 3 BO 3 Borate H 3O 3 boric acid NO 3- Nitrate HNO 3 nitric acid SO 2- Sulfate H SO sulfuric acid 4 - 2 4 ClO 4 perchlorate HClO 4 perchloric acid 3- phosphoric PO 4 phosphate H 3O 4 acid MnO 4- permanganate HMnO 4 permanganic acid CrO 42- chromate H 2rO 4 chromic acid - ClO 3 chlorate HClO 3 chloric acid Acids containing ions ending with ite usually become ClO - chlorite HClO chlorous acid 2-
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