Chemisty 1210, Week 2 Notes
Chemisty 1210, Week 2 Notes CHEM 1210
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Date Created: 09/04/16
WEEK 2 NOTES 8/29/16 – 9/2/16 Chapter 2: Atoms, Molecules, and Ions A. The Atomic Theory of Matter a. Dalton’s Atomic Theory i. Each element is composed of extremely small particles called atoms ii. All atoms of a given element are identical, but the atoms of one element are different from the atoms of all other elements iii. Atoms of one element cannot be changed into atoms of a different element by chemical reactions 1. Atoms are neither created nor destroyed in chemical reactions 2. Law of Conservation of Mass iv. Compounds are formed when atoms of more than one element combine 1. A given compound always has the same relative number and kinds of atoms a. Law of Constant Composition 2. Atoms combine in whole numbers a. Law of Multiple Proportions b. Dalton’s Atomic Theory explains the law of constant composition and the law of conservation of mass, and deduced the law of multiple proportions B. The Discovery of Atomic Structure a. Atoms i. Atomic mass unit (amu) = 1.6603 x 10^-‐24 g ii. Mass of atoms: 1 – 260 amu iii. Radii 1. Use nm iv. 1 Angstrom = 10^-‐10 m = 10^-‐8 cm b. Subatomic particles i. Make up atoms 1 ii. Electron 1. Charge = -‐1.6022 x 10^-‐19 C; relative = -‐1 2. m = 9.1094 x 10^-‐28 g = 5.486 x 10^-‐4 amu 3. atoms are neutral until electrons are gained or lost 4. number of electrons in an atom = number of protons in the atom 5. electrons are in constant motion 6. Can be traded back and forth between atoms iii. Proton 1. Charge = 1.6022 x 10^-‐19 C, relative = +1 2. m = 1.6726 x 10^-‐24 g = 1.0073 amu 3. make up about ½ of mass of an atom 4. number of protons in an atom of an element is that element’s atomic number 5. The number of protons do not change in a normal chemical reaction iv. Neutron 1. charge = 0 2. m = 1.6749 x 10^-‐24 g = 1.0088 amu 3. make up about ½ of mass of an atom v. Particles of the same charge repel each other; particles of different charges attract each other c. Cathode rays and Electrons i. The radiation produced when a high voltage is applied to electrodes in a tube ii. Cause certain materials to give off light (fluoresce) iii. J.J. Thompson 1. Discovered the electron by observing cathode rays as streams of negatively charged particles 2. Measured charge to mass ratio (coulombs) 2 iv. Robert Millikan-‐ Oil Drop Experiment 1. Determined the charge of an electron: 1.6022 x 10^-‐19 d. Radioactivity-‐ the spontaneous emission of radiation i. Ernest Rutherford found 3 types of radiation with his gold foil experiment ii. Alpha 1. Bent by an electric field 2. Fast moving particles 3. Have positive charge of +2; attracted to a negative charge iii. Beta 1. Bent by an electric field 2. High speed electrons 3. Radioactive equivalent of cathode rays 4. Have a charge of -‐1; attracted to a positive charge iv. Gamma 1. Unaffected by an electric field e. The Nuclear Model of the Atom i. Atoms contain a dense nucleus 1. Nucleus composed of protons and neutrons ii. Nucleus surrounded by electrons in mostly empty space 1. Electrons are in a constant motion iii. Atoms’ diameter = 10^-‐8 cm C. The Modern View of Atomic Structure a. Atomic Number (Z) i. Number of protons in an atom ii. Periodic table is arranged by increasing atomic number b. Mass Number (A) i. A = number of protons + number of neutrons ii. Is an integer c. Elemental/Atomic Symbol 3 i. Represents an element ii. EX: Zn = zinc iii. Number in top left corner is the element’s mass number iv. Number in the bottom left corner is the element’s atomic number d. Isotopes i. Atoms of the same element that have the same number of protons but a different number of neutrons ii. Have the same atomic number, but different mass number iii. Some mass gets lost in isotopic masses because it gets converted to energy when the atoms come together e. Atomic Weight i. Weighted average of the isotopes of an atom ii. Given in amu iii. The average atomic mass 1. Doesn’t give you the actual weight of the atom, but gives you an average D. Periodic Table a. The arrangement of elements in order of increasing atomic number i. Modern periodic law b. Places elements with similar properties in vertical columns c. Groups or families i. Vertical columns ii. Elements within a group share similar properties iii. Columns labeled at the top by either numbers, letters, or roman numerals iv. Representative Elements (Main group elements) 1. The first 2 columns and the last 6 columns 2. 1A-‐ 8A 3. 1A = alkali metals (except for hydrogen) 4 4. 2A = alkaline earth metals 5. 7A = halogens 6. 8A = noble gases v. Transition Metals 1. The middle 10 columns 2. 1B-‐ 8B d. Periods i. Horizontal rows ii. The 2 long rows below the main body of the periodic table are inner transition elements 1. Lanthanides – part of row 6 2. Actinides -‐ part of row 7 e. The stair step semimetals are transitional metals E. Metals vs Nonmetals a. Metals i. Solids 1. Except Hg ii. Metallic luster iii. Malleable and ductile 1. Can be pounded into shapes and drawn out into a wire iv. Conduct heat and electricity well v. Oxides (compounds): 1. Nonvolatile (ionic compounds) 2. High melting 3. MgO, Na2O b. Nonmetals i. Gases or solids 1. Except Br ii. Variety of color and appearance 5 iii. Solids are brittle iv. Not good conductors 1. Insulators v. Oxides: 1. Volatile 2. Low melting 3. CO, CO2, SO2 F. Molecules and Molecular Compounds a. Molecular formula i. Chemical formula that indicates the actual numbers of atoms in a molecule ii. Subscripts are always some integer multiple of the subscripts in the empirical formula b. Empirical formula i. Chemical formula that gives only the relative number of atoms of each type in a molecule ii. The smallest whole number ratio of atoms c. Molecular elements i. Diatomic molecule-‐ a molecule made up of 2 atoms 1. EX: H2… has 2 H atoms bonded together 2. All of the halogens are diatomic ii. Polyatomics 1. A molecule made up of more than 2 atoms 2. EX: P4, S8, O3 d. Molecular Compounds i. Compounds composed of molecules containing more than one type of an atom (contain atoms of different elements) ii. EX: CO2, H2O e. Structural formula 6 i. Shows which atoms are attached to which ii. Demonstrates structure G. Ions and Ionic Compounds a. Ion-‐ a charged particle that is formed when electrons are added or removed from an atom i. Cation-‐ positively charged ion 1. Metals tend to lose electrons and become cations ii. Anion-‐ negatively charged ion 1. Nonmetals tend to gain electrons and become anions b. Polyatomic ion-‐ atoms joined as a molecule but have a net charge i. NH^4+, SO4^2-‐ c. Predicting ionic charges i. Many atoms gain/lose electrons to end up with the same number of electrons as the noble gas closest to them in the periodic table 1. Representative elements (1A-‐ 8A) ii. Elements in the same isoelectronic series have the same electron configuration iii. Periodic Table 1. Cation groups a. Charge = group number i. Except for H b. Special cations: i. Al = +3 ii. Zn = +2 iii. Ag = +1 2. Anion groups a. Charge = group number – 8 3. 5A = -‐3, 6A = -‐2, 7A = -‐1 iv. Ionic Compounds 7 1. Compound made of cations and anions a. Oppositely charged ions are held together by electrostatic attractions 2. Compounds of metals and nonmetals a. Electrically neurtral 3. Crystalline solids (salts) 4. EX: Ca^2+ and CO3^2-‐ form CaCO3 (calcium carbonate) 5. Compounds are NOT molecules a. They are formula units 6. Empirical formula-‐ chemical formula that shows the simplest ratio of ions in a compound H. Naming Inorganic Compounds a. Organic compounds contain carbon and hydrogen, all other compounds are considered inorganic b. Chemical nomenclature-‐ system of naming substances c. Monatomic ions i. Cations 1. Formed from metal atoms: name of the elements followed by “ion” 2. EX: K^+ is potassium ion ii. Anions 1. Add “-‐ide ion” to the end of the element’s name 2. EX: Br^-‐ is bromide ion d. Stock System i. Metals that can have more than 1 possible charge 1. Transition and representative metals ii. Use roman numerals in parenthesis after the element’s name to represent the ion’s charge iii. EX: Fe ^3+ …. Iron (III); Sn ^4+ …. Tin (IV) 8 e. Polyatomic Ions i. Group of chemically bonded atoms with an overall charge ii. Cations 1. EX: NH4 ^+ ammonium ion 2. H3O^+ hydronium ion 3. Hg2^2+ mercury (I) ion iii. Anions 1. Ending in “-‐ide” a. EX: OH^-‐ hydroxide ion 2. Others a. EX: C2H3O2^-‐ acetate ion iv. Oxyanions-‐ polyatomic anions that contain oxygen 1. Suffixes a. “-‐ate” (most common) b. “-‐ite” (same charge but 1 fewer O atom) 2. Prefixes a. “per-‐“ i. 1 more O atom than “-‐ate” b. “hypo-‐“ i. 1 less O atom than “-‐ite” 3. EX: a. ClO4^-‐ perchlorate b. ClO3^-‐ chlorate c. ClO2^-‐ chlorite d. ClO^-‐ hypochlorite e. Cl^-‐ chloride f. ***Overall charge remains the same 4. Addition of H^+ to a -‐2 or -‐3 oxyanion a. Results still charged (anions) 9 b. EX: i. HCO3^-‐ bicarbonate/ hydrogen carbonate ii. HPO4^2-‐ monohydrogen phosphate 1. Mono-‐ is usually dropped iii. H2PO4^-‐ dihydrogen phosphate 5. Acids a. Acid-‐ H^+ produces a neutral compound when combined with an anion b. Not ionic i. But ionize in H2O to produce H^+ (H3O^+) c. EX: HCl (g) H2O H^+ (aq) + Cl^-‐ (aq) d. Binary (2 element) acids i. Hydrogen + nonmetal 1. –ide à hydro-‐ -‐ic acid a. EX: HF (aq) hydrofluoric acid 2. Per-‐ -‐ate à per-‐ -‐ic acid 3. –ate à -‐ic acid 4. –ite à -‐ous acid 5. hypo-‐ -‐ite à hypo-‐ -‐ous acid I. Binary Molecular Compounds a. Compounds contain 2 different elements i. Nonmetals or nonmetals and semimetals b. Element further left in the periodic table (closest to the metals) is usually listed first i. Exception when the compound contains oxygen and any halogen (except for fluorine), then oxygen is always written last c. If both elements are in the same group, the lower one is listed first (less electronegative) d. The name of the second element is given an –ide ending 10 e. Greek prefixes are used to indicate the number of atoms of each element i. Mono-‐ is never used with the first element 1. When the prefix ends with an a or an o and the second element begins with a vowel, the prefix is dropped ii. Prefixes: 1. Mono-‐ 2. Di-‐ 3. Tri-‐ 4. Tetra-‐ 5. Penta-‐ 6. Hexa-‐ 7. Hepta-‐ 8. Octa-‐ 9. Nona-‐ 10. Deca-‐ iii. EX: N2O4 Dinitrogen tetroxide iv. SO3 Sulfur trioxide J. Simple Organic Compounds a. Hydrocarbons-‐ compounds that contain only carbon and hydrogen b. Alkanes-‐ simplest class of hydrocarbons i. Each carbon is bonded to 4 other atoms ii. Methane CH4 iii. Ethane C2H6 iv. Propane C3H8 c. Functional groups i. Hydrogens are replaced with functional groups-‐ specific groups of atoms ii. Alkenes 1. Ethene CH2 (double bond) CH2 iii. Alkynes 11 1. Ethyne CH (triple bond) CH d. Alcohols i. Replace hydrogen atom with an –OH group ii. Methanol, ethanol, propanol 1. Are structural isomers of each other (have same chemical formula but different arrangements of atoms) e. Ethers i. Dimethyl ether, diethyl ether ii. Functional group isomers with alcohols f. Aldehydes i. C (double bond) O group … carbonyl ii. EX: HCHO formaldehyde g. Ketones i. EX: CH3COCH3 acetone h. Carboxylic acids i. Carboxyl group ii. EX: H-‐CO2H formic acid iii. CH3CO2H acetic acid i. Amines i. organic analogues of NH3 (ammonia) ii. EX: CH3NH2 methyl amine 12
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