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Chem 142: Chapter 2: Atoms, Molecules, and Ions

by: Teresa Ngo

Chem 142: Chapter 2: Atoms, Molecules, and Ions chem 142

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Teresa Ngo
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Atoms, Molecules, and Ions
Chemistry 142
Xiaosong Li
Class Notes
Chemistry, Chem, atoms, Molecules, ions




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This 9 page Class Notes was uploaded by Teresa Ngo on Sunday October 16, 2016. The Class Notes belongs to chem 142 at University of Washington taught by Xiaosong Li in Fall 2016. Since its upload, it has received 2 views. For similar materials see Chemistry 142 in Chemistry at University of Washington.


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Date Created: 10/16/16
BN Chapter 2: Atoms, Molecules, and Ions 2.1: The Early History of Chemistry  Greeks first to try to explain chemical changes (400BC) o All matter composed of 4 elements:  Fire, earth, water, and air o Matter?  Matter is continuous (infinitely divisible into smaller pieces) OR  Composed of small indivisible particles o Democritus  Atomos=atoms  Supported small indivisible particles o No definite answer b/c no experiments  Alchemy (next 2000 yrs) o Mystics obsessed w/ idea of turning cheap metal into gold o Discovery:  How to prepare mineral acids  Elements: mercury, sulfur, and antimony  Robert Boyle, the first "chemist" (1600's) o Measured the relationship b/t pressure & the volume of gases o Element=substance that cannot be broken down further  Combustion (17th-18th cent.) o Georg Stahl: phlogiston flowed out of burning material--stopped burning b/c container filled w/ phlogiston  Phlogiston: substance believed in the 18th cent. that exist in all combustible bodies; released in combustion o Joseph Priestley: vigorous combustion, low in phlogiston o Priestley's observations did not contradict Stahl's--diff. theories for WHY 2.2: Fundamental Chemical Laws  Combustion studied extensively (late 18th cent.) o CO , N , H , and O discovered 2 2 2 2  Antoine Lavoisier o Explained true nature of combustion o Measurement=essential to chemistry o Law of Conservation of Mass  Mass is neither created nor destroyed o Combustion involves O , not phlogiston 2  Life supported by process that involved O , similar to combustion… 2 o "Elementary Treatiste on Chemistry"  Joseph Proust o Law of Definite Proportion  A given compound always contains exactly the same proportion of elements by mass  Ex. Copper carbonate:  5.3 parts copper; 4 parts oxygen; 1 part carbon (mass)  John Dalton o Law of Multiple Proportions  When 2 elements form a series of compounds, the ratios of the masses of the second element that combine w 1g of the first element can always be reduced to small whole numbers  N 2 ,NO ,N O 4 2 2.3: Dalton's Atomic Theory Dalton's Model/ Modern Atomic Theory 1.  Each element is made up of atoms 2.  Atoms of a given element are identical  Atoms of diff. elements are different 3.  Chemical compounds formed when atoms combined A given compound always has the same relative  #'s and types of atoms 4.  Chem rxn =reorganization of atoms  Changes in the way they are bound  Atoms cannot be changed in a chem rxn  Dalton & Atomic Masses o First table of atomic masses  Many proved wrong, but still an impt step  Thought formula for water was OH  Water:  8g O: 1g H  OH = O atom 8x the mass of H  H2O = O atom 16x the mass of H  Decided nature would be as simple as possible  H mass of 1; O mass of 8  Joseph Gay-Lussac o 1809: performed experiments where he measured the volume of gases that reacted w/ one another (under same temp & pressure conditions)  2 volume H 2 , + 1 volume O 2 = 2 volume H 2 gas  1 volume Cl + 1 volume H = 2 volume HCl 2  Amedeo Avogadro o Avogadro's hypothesis  At the same temp & pressure, equal volumes of diff. gases contain the same # of particles  Assumption: distances b/t particles in a gas are v great compared to size of particles  Interpretation of Gay-Lussac's results:  2 volume H 2 + 1 volume O 2 = 2 volume H 2 (g)  2 molecules H 2 + 1 molecule O 2 = 2 molecules of H 2  Assumption: hydrogen, oxygen, and chlorine are all diatomic (two- atom) molecules  Not accepted by most  Most thought that only atoms of diff. elements could attract each other to form molecules; no "affinity" towards ea other  No agreement for formulas of elements =period table chaos o All interpreted in diff. ways depending on assumptions about formulas of elements & compounds 2.4: Cannizzaro's Interpretation  First International Chemical Congress (1860) in Karlsruhe, Germany by August Kekeulé o To come up with common set of atomic masses o Stanislao Cannizarro  Presented his ideas so clearly (lots of evidence) that virtually everyone accepted it  Cannizzaro's 2 main beliefs o Compounds contained whole number of atoms (as Dalton predicted) o Avogadro's hypothesis was correct  Hydrogen gas = H 2 molecules = relative molecular mass is 2 o Proved by Avogadro's hypothesis & Gay-Lussac's results of combining volumes of gases  Measured the relative molecular masses for other gaseous substances o Compare mass of 1L of given gas to 1L of H 2 (g) o Same temp & pressure o Ex. Ratio of masses of 1L O 2 : 1L H 2 is 32:2 OR 16:1  Carbon problem o Carbon dioxide: relative molecular mass is 44 o 27% carbon: (0.27)(44g)= 12g carbon/44g carbon dioxide o 44-12 = 32g oxygen  Oxygen relative mass = 16g o CO 2  12g carbon + 2 (16g oxygen) = 44g  HOWEVER  C O2 2 = 12g carbon represents 2 carbon atoms--> 6g relative mass  C O3 2 = relative mass of 4g o Cannizarro addressed this problem by obtaining relative molecular masses of many other compounds containing carbon  Data shows: always a multiple of 12  Relative mass of hydrogen = [(% hydrogen)/100] x relative molecular mass 2.5: Early Experiments to Characterize the Atom  The Electron o J.J Thomson  Studied cathode-ray tubes (electrical discharges in partially evacuated tubes)  Cathode ray - "ray" produced after high voltage is applied to the tube  Ray produced at the negative electrode and repelled by the negative pole of an applied field = ray was a stream of negatively charged particles (electrons)  Charge-to-mass ratio  e = 8 C/g −1.76x10 m  E = charge on the electron in coulombs  M = electron mass (g)  Plum pudding model  Believed that since electrons could be produced from electrode made of various types of metals, all atoms must contain electrons  Since atoms known to be electrically neutral --> atoms must also contain some positive charge  Postulated that an atom consisted of a diffuse cloud of positive charge w/ negative electrons embedded randomly in it  Plum pudding b/c pudding = positive cloud; electrons=raisins o Robert Millikan  Experiments w/ charged oil drops to determine magnitude of electron charge 31  Mass of electron = 9.11x10 kg  Radioactivity o Late 19th cent. : discovered that certain elements produce high-energy radiation o Antoine Henri Becquerel  Accidentally found that the image of a piece of mineral containing uranium could be produced on a photographic plate in the absence of light  Radioactivity: spontaneous emission of radiation by the uranium o Three types of radioactive emission:  Gamma (γ) - high energy "light"  Beta (β) - high-speed electron  Alpha (α) - 2+ charge  Twice that of an electron and w/ the opposite sign  Mass of α particle is 7300x that of an electron  The Nuclear Atom o Ernest Rutherford  Experiment to test Thomson's plum pudding model  Direct α particles at a thin sheet of metal foil  If correct/expected results: massive α particles should crash through the thin foil  Minor deflections  Results: most α particles passed straight through; many deflected at large angles; some reflected, never hit detector  Proved:  Center of concentrated positive charge that contained most of atom's mass  Nuclear atom : atom w/ nucleus  Nucleus: dense center of positive charge  Most passed directly through b/c atom is mostly open space  Deflected α particles = had "close encounter" w/ nucleus  Reflected α particles = "direct hit" on nucleus  Electrons moving around nucleus at a distance that is large relative to nuclear radius 2.6: The Modern View of Atomic Structure: An Introduction  Atom 13 o Consists of a tiny nucleus w/ diameter of ~ 10 cm o Electrons move about nucleus at avg. distance of ~ 8 from nucleus 10 cm o Chemistry of an atom mainly result from its electrons  Protons o Positive charge equal in magnitude to electron's negative charge  Neutrons o Neutral charge o Virtually same mass as proton but no charge Properties of subatomic particles Name Symbol Charge (in Approximate mass multiples of e-) (amu) Electron e- -1 0.0005 Proton p+ +1 1.0 Neutron 0 0 1.0 n  Nucleus o Small size (compared to overall size of atom) o Extremely high density  Nucleus accounts for almost all of the atom's mass  ex. Like a pea having the mass of 250 mill tons!  If all atoms are composed of these same components, why do different atoms have different chemical properties? o Based on number and arrangement of the electrons o Electrons constitute most of the atomic volume & intermingle when atoms combine to form molecules  # of electrons possessed greatly affects its ability to interact w/ other atoms  Therefore, atoms of diff. elements, w/ diff. #'s of protons & electrons, show diff. chemical behavior  Isotopes o Atoms w/ the same number of protons but diff. numbers of neutrons  Atoms of the same element that differ in mass o Show almost identical chemical properties o In nature, most elements contain a mixture of isotopes o Note: Isotope mass/weight doesn't affect # protons/electrons--only charge changes electrons 12 13 C  C , C,∧14 ,allhavesameelectron¿ AX Z  Atomic number (Z) = # protons (number of electrons too, if molecule is neutral)  Mass number (A) = total # protons & neutrons  Element (X)  Electrical Force o Use Coulomb's Law (strength of electric force b/t objects that have an electric charge) 1. The electric force gets stronger as distance b/t charged objects gets smaller 2. The electric force gets stronger as the magnitude of either charge increases  Positive/negative 2.7 Molecules & Ions  Chemical bonds o Forces that hold atoms together in compounds  Covalent bonds o Sharing electrons  Molecule o Collection of atoms  Chemical formula o Simplest representation o Symbols used to indicate types and relative #'s of atoms o Write symbols in order of how they appear on the periodic table (left to right)  Exception: write Hydrogen after all elements except those from Group 6A and 7A  Write hydrogen in front of those with ONLY hydrogen and elements from 6A and 7A  Structural formula o Individual bonds shown o May/may not indicate actual shape of molecule  Dashed lines = behind the plane of the pg.  Wedge = in front of the plane of the pg. o Space-filling model  Shows both relative sizes of the atoms in the molecule & their spatial relationships o Ball-and-stick model  Ionic bond o Force of attraction b/t oppositely charged ions o Ion  Atom/group of atoms that has net positive or negative charge  Cation  Positive ion  Lost electron  Anion  Negative ion  Gained electron  Cations & anions attract ea. Other o Ionic solid aka salt  Solid consisted of oppositely charged ions  Can consist of simple ions or  Polyatomic ions  Molecular vs. Ionic Compounds o Based on chemical elements in it:  Molecular: nonmetals  Ionic: metals & nonmetals 2.8: An Introduction to the Periodic Table  Metals o Physical properties:  Efficient conduction of heat & electricity  Malleability (can be hammered into thin sheets)  Ductility (pulled into wires)  Lustrous appearance o Chemical properties:  Tend to lose electrons to form positive ions  Nonmetals o Physical properties:  Lack properties that characterize metals o Chemical properties:  Tend to gain electrons to form anions in reactions w/ metals  Bond together to form covalent bonds  Groups or families o Vertical columns o *Have similar chemical properties o Alkali metals (1A)  Very active elements that readily form 1+ charge w/ nonmetals o Alkaline earth metals (2A)  Form ions w/ 2+ charge when react w/ nonmetals o Halogens (7A)  Form diatomic molecules o Noble gases (8A)  Exist under normal conditions as monoatomic (single-atom) gases  Little chemical reactivity o Elements in the same group are most similar to ea other  & metals are more similar to ea other than nonmetals, & vice versa  Periods o Horizontal row  Standard states o All metals are solids  Except mercury; Hg(l) o Most nonmetals are simple solids  Except: noble gases, seven diatomic nonmetals, phosphorus, and sulfur  Diatomic nonmetals are mostly gases, but iodine is a solid and bromine is a liquid 2.9: Naming Simple Compounds  Binary compounds o Compounds composed of 2 elements  Binary ionic compounds (Type I: Ionic) o Contains: positive cation from metal element and negative anions from a nonmetal element o Rules: 1. The cation always named first, anion second 2. A monatomic (one atom) cation has same name as its parent element  Na+ = sodium 3. A monatomic anion is named by taking the first part of the element name and adding -ide  Chloride  Binary Compounds (Type II: Ionic) o Contains: a metal that can form more than one type of cation--must have Roman numeral in its name 1. Most common in compounds containing transition metals o The ion w/ the higher charge has a name ending in -ic, and the one w/ the lower charge has a name ending in -ous 2+¿=ferrousion ¿ 1. 3+¿= ferricion;Fe Fe ¿  Ionic Compounds w/ Polyatomic Ions o Ionic compounds that contain polyatomic ions are not binary compounds o Oxyanions  Polyatomic anion made from a central atom chemically bonded to one or more atoms of oxygen  Two members in a series of oxyanions:  Name of the one w/ smaller # of oxygen atoms ends in -ite  Name of the one w/ the larger # of oxygen atoms ends in -ate 2−¿ 2−¿  Ex. Sulfite ( ¿ Sulfate ( ¿ S O 3 S O 4  More than 2 members in a series of oxyanions:  Hypo- Less than  Per- More than  Prefixes to name member w/ fewest & most oxygen atoms 2 less O atoms 1 less O atom Most common oxoanion 1 more O atom −¿ −¿ −¿ −¿ ¿ ¿ ¿ ¿ ClO ClO 2 ClO 3 ClO 4 Hypochlorite chlorite chlorate perchlorate  Binary Compounds (Type III: Covalent-Contains Two Nonmetals) o Binary covalent compounds  Formed b/t two nonmetals  Use prefixes  Be sure to include prefix for second molecule o Do not contain ions but are named very similarly to binary ionic compounds o Rules:  First element in the formula is named first, using full element name (and prefix)  Second element is named as if it were an anion  Prefixes are used to denote numbers of atoms present  Prefix mono- is never used for naming the first element  Still used in second name  Baron oxide is NOT correct--> baron monoxide is o Avoid awkward pronunciation: often drop the final o or a of the prefix when element begins with vowel  Ex. Dinitrogen tetroxide, not dinitrogen tetraoxide  Formulas from Names o Ex. Calcium hydroxide 2+¿  Calcium forms only ions C a ¿  Hydroxide is OH-  2 OH- needed to make compound neutral  Ca (OH ) 2  Acids o Molecule w/ one or more H+ ions attached to an anion o Dissolved in water, certain molecules produce solution w/ free H+ ions (protons) o Depend on whether the anion contains oxygen o Simple acid rules have a lot in common w rules for naming ionic compounds o Focus on anion left behind if all acidic hydrogens were taken away from a molecule of the compound  Acidic hydrogens usu. written first in the chemical formulae of simple acids  Note: acids lose their acidic hydrogens as hydrogen cations (H+) o Binary acid (no oxygen):  Contains: hydrogen & just one other nonmetal element  Anion will be an atomic ion HYDRO + anion root + ic acid ATOMIC ANION BINARY ACID Formula Name (root highlighted) −¿ Fluoride F ¿ −¿ ¿ Chloride Cl 2−¿ SFormula Name (anion root) ¿ HF Hydrofluoric acid S HCl Hydrochloric acid Hydrosulfuric acid  Ex. HCl H 2 (g) dissolved in H O2 = hydrochloric acid o Oxyacid (yes oxygen):  Contains: hydrogen, oxygen, and one other element (usu. nonmetal)  Anion will be an oxoanion OXOANION ROOT + IC/OUS + ACID  -ic ending if the name of the oxoanion ends in -ate  -ous ending if the name of the oxoanion ends w -ite OXOANION OXYACID Formula Name (root Formula Name (anion root) highlighted) −¿ Nitrate Nitric acid ¿ HN O 3 N O 3 −¿ Nitrite Nitrous acid ¿ HN O 2 N O 2 Perbromate Perbromic acid −¿ ¿ HBrO 4 BrO 4 Hypoiodite Hypoiodous acid −¿ ¿ HIO IO Sulfate Sulfuric acid 2−¿ H 2 O 4 SO ¿ 4 2−¿ Sulfite H SO Sulfurous acid ¿ 2 3 SO 3 2−¿ ¿  Ex. SO 4 ¿ H 2O co4tainsthesulfateanion¿ 3−¿ P O ¿ 4 ¿ H PO contains the phosphateanion¿ 3 4 o Anion w/ irregular name o Name of anion ends w/ -ide, name the acid the way you name binary acids o Name of anion ends w/ -ate, name the acid the way you name oxyacids ANION ACID Formula Name (root highlighted)Formula Name (anion root) −¿ Cyanide Hydrocyanic acid ¿ HCN C N −¿ Thiocyanate Thiocyanic acid ¿ HSCN SC N −¿ Acetate HC H CO Acetic acid ¿ 3 2 C H 3O 2 2−¿ Oxalate H C O Oxalic acid 2 2 4 C O ¿ 2 4  Acid Salts o Contains: atomic cations and polyatomic anions o Each anion is an oxoanion w/ one or more extra hydrogens  MgHPO 3 2−¿  Mg+ cations; HPO3- anions HPO = ¿ w/ one extra H+) 3 PO 3  Name: Magnesium hydrogen phosphite o The extra H+'s on the anion are acidic hydrogens  Released as hydrogen cations (H+) when the anion dissolves in water  That makes the anion of this salt an acid --> acid salt o Follows general pattern of naming ionic compounds  [cation name] [anion name] NUMBER PREFIX + HYDROGEN + OXOANION NAME  NaH PO --> when 2 hydrogens are present, do not include "hydrogen" 2 4  Replace with "bi" --> sodium biphosphate  Add "bi" for an anion made of an oxoanion w/ enough extra H+ to give it an overall anion charge of -1  Ionic compounds w/ common polyatomic ions o Use net charge on ea. formula unit--> compound must = zero (neutral charge) o Cation is an atomic ion of a metal element:  Its name is just the name of the metal element, followed by the cation charge in parentheses when the metal can form more than one cation  Most of the metals that can form more than one cation are transition metals  *Note: Leave out charge when cation is formed from atom of group 1A or group 2A metal, or from atoms of the elements Al, Zn, or Cd  Has a definite charge; doesn't change o Anion is an atomic ion of a nonmetal element:  Its name is formed from the name of the nonmetal element by replacing everything after the first syllable with -ide o Oxoanions:  The name of the most common oxoanion is first syllable of name of the central element plus the suffix -ate, and the name of the less common oxoanions are made by adding diff prefixes & suffixes  Most common (below): chlorate  *Note: all oxoanions w/ same central elemtn have the same charge SOME IONIC COMPOUNDS Cation Anion Empirical formula Name of compound +¿ −¿ KCl O Potassium perchlorate K ¿ ClO ¿ 4 4 2+¿ −¿ Ca (CN ) Calcium cyanide ¿ C N ¿ 2 C a +¿ −¿ Ammonium nitrate ¿ ¿ N H N4O 3 N H 4 N O 3


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