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General Chemistry, Chapter 2 Reading!

by: Madelyne Crawford

General Chemistry, Chapter 2 Reading! CHEM-111

Marketplace > Campbell University > Chemistry > CHEM-111 > General Chemistry Chapter 2 Reading
Madelyne Crawford


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These notes are an overview of what is in Chapter 2. There is a total of five pages of detailed information that I think we will be quizzed on next Wednesday, September 7. A Quizlet is linked to th...
General Chemistry
Dr. Kesling
Class Notes
Chemistry, Moles, Molecules, atomic, Theory, Brown, mass
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This 6 page Class Notes was uploaded by Madelyne Crawford on Wednesday August 31, 2016. The Class Notes belongs to CHEM-111 at Campbell University taught by Dr. Kesling in Fall 2016. Since its upload, it has received 109 views. For similar materials see General Chemistry in Chemistry at Campbell University.


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Date Created: 08/31/16
Chapter 2 Reading Guide: Atoms and Elements By Madelyne Crawford August 31, 2016 Section 2.1: Brownian Motion (Atoms Confirmed)  Robert Brown (1827- Scottish botanist)- noticed pollen grains were always in motion continuously o Originally thought particles were alive (male sex cells of plants) because they were always moving o His hypothesis was incorrect because he observed continuous motion of stone particles, which definitely could not be alive o He never figured out what was causing their continuous motion  Answer: Albert Einstein (1905)- Brownian motion o His model quantitatively explained the continuous motion Brown had observed o Moving because of their constant encounter with thermal energy  Jean Perrin (1908)- French physicist who confirmed Einstein’s Brownian model o Won Nobel Prize- put an end to the curiosity of the presence of matter o Stopped doubt of particles and the nature of matter  Scanning tunneling microscope (STM)- can be used to pick up and move individual atoms- structures can be made one atom at a time  Nanotechnology- image, move, and build atoms  How big is an atom? Imagine a grain of sand- more atoms in the grain of sand than you could ever count in a lifetime  Atoms are KEY to relating/connecting macro and microscopic worlds Section 2.2: Early Ideas About the Building Blocks of Matter  First people to propose the idea of matter being composed of small, indestructible particles: o Leucippus (fifth century BC) and his student Democritus (460- 370 BC) o Greek philosophers o “atomos” o many different atoms and they all move randomly, according to them o Plato and Aristotle did not agree- believed that everything was composed of different amounts of fire, air, earth, and water- THIS VIEW PREVAILED, even though Leucippus is ultimately right  Nicholas Copernicus (beginning of modern science- 1543) o On the Revolution of Heavenly Orbs- said the sun was the center of the universe and not the earth (SCIENTIFIC REVOLUTION)  Advancement through the next 200 years (scientific approach became the modern way of learning about the physical world) o Francis Bacon, Johannes Kepler, Galileo Galilei, Robert Boyle, Isaac Newton o These advancements helped John Dalton come up with atomic theory and confirm Leucippus and Democritus’s perspectives and ideals on atoms Section 2.3: Modern Atomic Theory and the Laws That Led to It  During chemical reactions matter is neither created nor destroyed, but it does rearrange  Joseph Proust (1797)- French chemist, composition of compounds o Found that elements composing compounds always are in fixed proportions o Contrast: components of a certain mixture can be in any proportion of the mixture no matter what o Law of Definite Proportions: “All samples of a given compound, regardless of their source or how they were prepared, have the same properties of their constituent elements.” o Compounds have definite amounts of their elements because of their composition of atoms  Law of multiple proportions- “When two elements (A and B, for example) form two different compounds, the masses of element B that combine with 1g of element A can be expressed as a ratio of small whole numbers.” o John Dalton, 1804 o Had no fancy machine to detect individual atoms, so Dalton used masses of samples  Atomic Theory (in full) o 1. Elements are each made up of small, indestructible particles (ATOMS) o 2. All atoms of a certain element have same mass o 3. Atoms combine: simple, whole number ratios to make compounds o 4. Atoms are not interchangeable- cannot change into atoms of another element (atoms only change the way they are bound together with other atoms) 2.4: The Discovery of the Electron  JJ Thomson (late 1800s)- working at Cambridge o Cathode rays- beams of electrons that come out of the cathode ray tube (Thomson’s invention) o Applied high voltage and observed cathode rays moving from negatively charged electrode (cathode) to the positively charged one (anode)  Composition of particles in cathode ray: o Travel in straight lines o Do not depend on original material’s composition (cathode) o Negative electrical charge  Electrostatic forces- compose atoms, which results in attractive and repulsive forces  Electric field- area around a charged particle where the electrostatic forces exist  Measured charge-to-mass ratio by deflecting them using electric and magnetic fields o Cathode ray particle was 2000 times lighter than hydrogen (lightest know atom) o Discovery of the electron!!- negatively charged, low-mass particle that is in all atoms  Robert Millikan (1909- American physicist)- oil drop experiment o Sprayed oiled in small drops, fell through a small hole because of gravity, during fall the drops would gain electrons that Millikan had made by bombarding the air in the area with ionizing radiation o Drops acquired a negative charge o He could slow down or reverse the free fall of the negatively charged drops by creating an electric field between two metal plates o Millikan was able to calculate the charge of each drop by measuring the strength of the electric field that was needed to stop the free fall of the drops o Reasoned that if each drop has to contain a whole number of electrons, the charge must be a whole-number that is a multiple of the electron’s charge  Why did scientists want to know this? The magnitude of the charge of the electron determines the strength of an atom’s ability to hold electrons 2.5: The Structure of the Atom  Plum-pudding model: suggested by JJ Thomson, negatively charged electrons were tiny particles within a positively charged sphere  Radioactivity- emission of energetic particles from the center of unstable atoms o Discovered by Henri Becquerel and Marie Curie o Allowed scientists to probe the structure of an atom for the first time o Alpha particles- positively charged and most massive (in terms of gamma and beta)  Ernest Rutherford (1909) o Worked under Thomson and attempted to prove his plum- pudding model wrong by doing an experiment in which he directed positively charged alpha particles through thin sheet of gold foil- according to Thomson’s model, the probes should have passed right through the gold foil o Results were unexpected: some particles went through but were deflected and some bounced back o Proposed the nuclear theory to explain this!!  Nuclear Theory: o 1. Nucleus- most of atom’s mass and all positive charge are in this small core o 2. Volume of atom is mostly empty space- negatively charged atoms are dispersed everywhere through the atom o 3. Lots of negative charge electrons outside nucleus and protons inside (electrically neutral atom)  Unaccounted for mass… WHY?? James Chadwick (one of Rutherford’s students) o Neutrons- neutral particles inside of the nucleus, no electric charge  Nucleus contains 99.9% of the atom’s mass, but not very much volume of the atom  Rutherford’s theory is still valid today- matter, at its core, is not as uniform as it appears 2.6: Subatomic Particles: Protons, Neutrons, and Electrons in Atoms  Atomic mass unit- (amu) common unit to express the mass of protons, neutrons, and electrons; ½ the mass of a carbon atom that contains six protons and six neutrons  The mass of a proton or neutron is about 1 amu, while electrons almost have an insignificant, or barely noticeable, mass  “The charge of the proton and the charge of the electron are equal in magnitude, but opposite in sign-” charges equal 0 when they are paired, neutron has no charge  matter is usually neutral (in terms of charge) because protons and electrons are present in equal #s  The number of protons in the nucleus of an element creates the element’s identity  Atomic number- # of protons in an atom’s nucleus (symbol Z) (range from 1-116)  Chemical symbol- one or two letter abbreviation (below atomic number on periodic table)  What determines the element’s name?- Latin names (Na for sodium because of Latin word natrium), reflective of the element’s properties (argon comes from the Greek word argos, which means inactive), where they are discovered (europium), or their discoverer’s name (curium)  Despite what Dalton stated in atomic theory, not all atoms have the same mass o Isotopes- atoms with the same number of protons, but different amount of neutrons o Relative amount of each isotope is roughly constant  Natural abundance- the natural amount of isotopes an element has  Mass number- the sum of the number of neutrons and protons in an atom (mass number is represented by A) o A= number of protons + number of neutrons o Therefore, the number if neutrons in an isotope is mass # - atomic # o A= mass number, Z= atomic number, and X= chemical symbol  ions- the result of the process in which during chemical changes, atoms lose and gain electrons, becoming charged particles  cations- positively charged ions (like Li+)  anions- negatively charged ions (like F-) 2.7: Finding Patterns: The Periodic Law and the Periodic Table  Dmitri Mendeleev- nineteenth century Russian chemistry professor (about 65 elements had been discovered by this time) o Noticed that certain element groups had similar properties (periodic pattern) o Periodic law- when the elements are arranged in order of increasing mass, certain sets of properties recur periodically o Left gaps because he predicted there would be more elements (silicon!!) o Clemens Winkler discovered silicon in 1886, just like Mendeleev predicted o Modern periodic table was created because of Mendeleev, except the elements are listed in terms of increasing atomic number instead of relative mass  Metals- lower left side and middle of the periodic table o Common properties: good conductors of heat and electricity, malleability, ductility, shiny, tendency to lose electrons during chemical reactions o Chromium, copper, strontium, lead  Nonmetals- upper right side of periodic table o Some are solids at room temp, while others are liquids or gases o Poor conductors of heat and electricity o Gain electrons during chemical reactions o Oxygen, carbon, sulfur, bromine, iodine  Metalloids- along zigzag line on periodic table o Semiconductors- electrical conductivity is highly temp dependent o Ability to change the conductivity of semiconductors (cell phones)  Main-group elements- properties are very predictable depending on their place on the periodic table (labeled with an A)  Transition elements- properties are less predictable (labeled with a B)  Family or group of elements- each column within the main-group regions of the periodic table  Noble gases- group 8A elements; mostly unreactive; neon, krypton, helium  Alkaline metals- group 1A elements; all reactive; sodium reacts when dropped into water, lithium, potassium  Alkaline earth metals- fairly reactive; magnesium, calcium, barium  Halogens- group 7A; very reactive nonmetals; chlorine, bromine, iodine  Main-group metals usually lose electrons, which form cations with the same # of electrons as the closest noble gas  Main-group nonmetals usually gain electrons, which form anions with same # of electrons as closest noble gas 2.8: Atomic Mass: The Average Mass of an Element’s Atoms  Atomic mass= average mass, we need this because isotopes cause atom’s to have different masses  “weighted according to the natural abundance of each isotope” o allows the assigning a characteristic mass to each element  Mass spectrometry- method that separates particles in terms of their mass o Measure masses of atoms and the percent of the abundance of elements o Can be used to identify an unknown molecule and the amount present o Tumors can now be analyzed by mass spectrometry to see whether it is cancerous 2.9: Molar Mass: Counting Atoms by Weighing Them  It is important to know the number of atoms in a given mass’ sample because chemical processes happen between particles  Mole- used to more efficiently count the number of atoms, amount of a substance o 1 mole= 6.02214 x 10^23 particles o this is called Avogadro’s number (named after Italian physicist Amedeo Avogadro)  The value of a mole is = to # of atoms in 12g of pure carbon-12 o 12gC = 1 mol C atoms = 6.022 x 10^23 C atoms  molar mass-mass on one mole of atoms of an element  an element’s molar mass (in g/mole) is equal to element’s atomic mass (in amu) Quizlet link!! cards/?new


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