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Chem131, Week 1 notes

by: Christina Notetaker

Chem131, Week 1 notes CHEM131

Christina Notetaker

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About this Document

These notes cover the first week of the semester. 8/31/16- 9/2/16
Chemistry I - Fundamentals of General Chemistry
John Ondov
Class Notes
General Chemistry, 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  10​12 10​9 10​6 10​3 10​2 10​1 | 10​­1 10​­2 10​­3 10​­6 10​­9 10​­12  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 a1​mic ma2​ ​ ) + etc...  2​ A = (% abundance/ 100) 


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