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Week 2 Notes

by: Victoria Mazur

Week 2 Notes CH-1213-32

Victoria Mazur

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Atoms, Molecules, and Ions
Chem 1
Dr. Dornshuld
Class Notes
Chemistry, Lecture Notes
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This 7 page Class Notes was uploaded by Victoria Mazur on Sunday August 28, 2016. The Class Notes belongs to CH-1213-32 at Mississippi State University taught by Dr. Dornshuld in Fall 2016. Since its upload, it has received 92 views. For similar materials see Chem 1 in Chemistry at Mississippi State University.


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Date Created: 08/28/16
Chapter 2: Atoms, Molecules, and Ions Definitions Atoms- the smallest unit of an element; can participate in chemical change Compound- consists of atoms of two or more elements combined in a small, whole number ratio Alpha Rays- positively charged particles; relatively massive Beta Rays- negatively charged particles Gamma Rays- not effected by the electric or magnetic fields Nucleus- the center of an atom Proton- positively charged; same magnitude of charge as an electron; larger mass −19 −24 (charge: +1.6022*0 C___ mass: 1.67262*0 g) Electron- negatively charged; small mass −19 −28 (charge: -1.6022* C___ mass: 9.10*0 g) Neutron- uncharged particles; roughly the same mass as a proton Neutral Atoms- no charge; equal numbers of protons and electrons Isotope- atoms of the same element with different masses Ions- atoms that exhibit charge; unequal between protons and electrons Cation- positive charge Anion- negative charge Atomic Number- the number of protons in the nucleus of an atom Mass Number- the number of protons and neutrons; “nucleons” Atomic Mass Unit- used to describe an atom’s mass; “amu” Molecular formula- a representation of a molecule using chemical symbols Structural formula- a representation of a molecule showing HOW they are connected Diatomic molecules- molecules with two of the same atoms bonded Compounds- two or more elements chemically combined Empirical formula- the simplest whole-number ratio of atoms or ions in a compound Isomers- compounds with the same formula but a different structure Spatial Isomers- compounds with the same formula and structure, with a different orientation in space Periods- rows in the periodic table Groups- columns in the periodic table Metals-elements that are shiny, malleable, electric and heat conductors Nonmetals- dull elements, poor electric and heat conductors Metalloids- mediocre heat and electricity conductors that possesses properties of metals and nonmetals Main group elements- elements in groups 1, 13-18; hydrogen is not included; sometimes group 12 is included Transition elements/metals- all elements in groups 3-11 Volatile metals- group 12 Monatomic ions- ions formed from a single atom Polyatomic ions- a group of atoms bonded together as discrete units with an overall charge Oxyanions- polyatomic ions that contain one or more oxygen atoms Ionic compounds- a compound of ions held by ionic bonds; electrically neutral; high melting point; not Conductive as a solid Molecular compounds- a compound held together by covalent bonds; low-boiling liquids; low melting solids; often seen as a gas Nomenclature- a collection of rules for naming things Binary compound- a compound containing only two elements Acid- a compound that produces hydrogen ions when dissolved in water Binary acid- an acid that is comprised of H and one other nonmetal Oxyacid- an acid containing H, O, and another element 2.1 Five Postulates of Dalton’s Atomic Theory 1. Matter is composed of atoms 2. An element consists of only one type of atoms, with a specific mass 3. The properties of each element is different than every other 4. *see compound 5. Atoms are not created or destroyed, only rearranged through chemical changes 2.2 Evolution of the Atomic Theory (experiments) Cathode Tube Ray Experiment  J J Thompson (late 1800s)  The particles in the Cathode rays were attracted to the positive electric and magnetic fields o Those particles thus have a negative charge 108 g−1  Mass-to-charge ratio -1.76* C  Particles were lighter than atoms Oil Drop Experiment  Robert A. Millikan (1909)  Added a charge to oil drops and changed their fall rate with electric fields  Determined the charge on each drop −19  The fundamental charge of an electron -1.6022* 10 C  Discovered the mass of electrons using Thompson’s discovery −19 8 Mass of electrons= charge/(mass-charge-ratio) = (-1.6022* 10 ) / ( -1.76*10 ) 10 −28 = 9.10* Radiation Experiments  Late 1890s  Radioactive material fives off radioactive rays: alpha, beta, and gama  The degree of change in direction shows how massive the alpha and beta rays are o Alpha- relatively massive o Beta- less massive than alpha rays Gold-Foil Experiment  Ernest Rutherford (early 1900s)  Directed alpha particles at a thin sheet of gold  Hypothesis: the alpha particles will flow right through the sheet, unchanged.  Alpha particles actually were dispersed non-uniformly, some even bouncing back. 2.3 Atomic Structure and Symbolism Why are isotopes interesting and possible? They have the same number of protons and electrons, but have different masses. The masses are different due to the number of neutrons. Atomic number 1 Mass number A (number of protons) H (protons + X neutrons) Hydrogen Z Atomic number 1.008 Element symbol The atomic mass is approximately equal to the mass number. It is AVERAGED of the values of the known masses of the isotopes of the element. ATOMIC MASS = ∑ (fractional abundance*isotopic mass) 2.4 Chemical Formulas Molecular formula examples CH₄ CH₃COOH Structural formula examples 2.5 The Periodic Table KEY Alkali Metals Alkaline Earth Metals Transition Metals Pnictogens(13-15) Chalcogens(16) Halogens (17) Noble Gasses Lanthanides Actinides Common Polyatomic Ions Charge Name Formula 1+ Ammonium NH₄⁺ 1- Acetate C₂H₃O₂⁻ 1- Cyanide CN⁻ 1- Hydroxide OH⁻ 1- Nitrate NO₃⁻ 1- Nitrite NO₂⁻ 1- Perchlorate CIO₄⁻ 1- Chlorate CIO₃⁻ 1- Chlorite CIO₂⁻ 1- Hypochlorite CIO⁻ 1- Permanganate MnO₄⁻ 1- Hydrogen HCO₃⁻ carbonate/bicarbo nate 2- Carbonate CO₃²⁻ 2- Peroxide O₂²⁻ 1- Hydrogen sulfate/ HSO₄⁻ bisulfate 2- Sulfate SO₄²⁻ 2- Sulfite SO₃²⁻ 1- Dihydrogen H₂PO₄⁻ phosphate 2- Hydrogen PHO₄²⁻ phosphate 3- Phosphate PO₄³⁻ 2.6 Molecular and ionic compounds Rules for naming Ionic Compounds  Binary compounds: the name of the metal followed by the anion with the ending -ide  Compounds containing polyatomic ions: same as monatomic ions, except you do not change the ending (-ide) 2.7 Chemical Nomenclature Compounds containing a metal ion with variable charge:  Transition metals can take on different charges  Use parenthesized roman numerals following the metal to indicate charge Compounds composed of two elements  Use prefixes to indicate number of atoms  Change endings to -ide Binary Acids  Change “hydrogen” to “hydro- “  Change the ending of the non-metal to -ic  Add the word acid to the very end Oxyacids  Omit “hydrogen”  Start with the root name of the anion  Replace -ate with -ic, or -ite with -ous  Add the word acid to the very end


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