General Chemistry 1, ch. 1-4 Study Guide
General Chemistry 1, ch. 1-4 Study Guide CHEM 1110 - 09
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This 5 page Study Guide was uploaded by Kristen Manda on Friday September 16, 2016. The Study Guide belongs to CHEM 1110 - 09 at University of Tennessee - Chattanooga taught by Rebecca E. Stimson in Fall 2016. Since its upload, it has received 61 views. For similar materials see General Chemistry 1 in Science at University of Tennessee - Chattanooga.
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Date Created: 09/16/16
Chemistry Chapters 14 Study Guide CHAPTER 1 Chemistry: the study of matter and changes that matter undergoes. Scientific Method: observations, hypothesis, experiment, theory Hypothesis: a testable explanation which explains the observations Theory: explains data from accumulated experiments, used to predict other results Matter: anything that occupies space and has mass Substance: form of matter that has a constant composition throughout and has distinct properties. Mixture: combination of 2 or more substances in which their identities remain distinct. Homogenous mixtures: composition is the same throughout, sugar, water, iced tea Heterogeneous mixtures: composition is different throughout. Sand and iron fill, trail mix Physical Properties: can be observed and measured without changing its identity Chemical Properties: exhibited when interacting with other substances. EX) corrosiveness, flammability Physical Changes: state changes but identity doesn’t. EX) melting, boiling, freezing Chemical Change: change in composition; original no longer there. EX) combustion, digestion, oxidation Significant Figures: 1) Any nonzero digit is significant 2) 0s between non zero digits is significant 3) 0s to the left of a nonzero digit aren’t 4) zeros to the right of the last nonzero digit are significant if decimal is present **When multiplying and dividing, the number of significant figures in the final product or quotient is determined by the original number that has the smallest number of significant figures. **When adding and subtracting, the answer cannot have more digits to the right of the decimal point than any of the original numbers. SI Units/Prefixes Length: meter (m) Amount of Substance: mole (mol) Mass: kilogram (kg) luminous intensity: candela (cd) Temperature: kelvin (K) electric current: ampere (A) CHAPTER 2 Atom: smallest quantity of matter that still retains the properties of matter Element: substance that cannot be broken down into two or more simpler substances. Alpha Rays: positively charged particles Beta Rays: skews away from negative field Gamma Rays: highest intensity Atomic Number: (z) the number or protons and also the # of electrons. Mass Number: (A) protons + neutrons Isotopes: atoms that have the same atomic number but different mass number. A different number of neutrons and isotopes have similar chemical properties. Atomic Mass: mass of an atom (amu). 1 amu= ½ the mass of a carbon 12 atom Average Atomic Mass: average mass of the naturally occurring mixture of isotopes. (isotopic mass)(natural abundance) + (isotopic mass)(natural abundance) **change all percent’s into decimals Mole: the amount of a substance that contains as many elementary entities as there are atoms in exactly 12g of carbon 12. Avogadro’s Number: 6.022e23. There are 6.022e23 atoms in 1 mol. Molar Mass: the mass in grams of one mole of the substance (g/mol). People: John Dalton: said that atoms are tiny indivisible particles JJ Thompson: discovery of the electron; used the cathode ray tube RA Millikan: determined the charge of an electron; examined the motion of tiny oil drops suspended in an electric field Wilhelm Rontgen: discovered xrays Antoine Becquerel: discovered radioactivity Ernest Rutherford: gold foil experiment: the nucleus is positively concentrated James Chadwick: discovered the neutron Einstein: a beam of light is a stream of particles Schrodinger: equation for the probability of finding an electron in a certain area of space Louis de Broglie: electrons can act like waves Max Plank: radiant energy is only emitted or absorbed in discrete quantities Chapter 3 Energy: the capacity to do work or transfer heat Kinetic Energy: the energy that results from motion Potential Energy: energy possessed by an object by virtue of its position. Thermal Energy: form of kinetic energy, energy associated with the random motion of atoms and molecules Chemical Energy: the energy stored within the structural units of chemical substances Electrostatic Energy: is potential energy that results from the interaction of charged particles Law of Conservation of Energy: energy cannot be created or destroyed Joule: unit of energy, amount of energy possessed by a 2kg mass moving at a speed of 1m/s. Wavelength: peak to peak of a wave Frequency: cycles per second Amplitude: height of wave, center to peak or troph Node: amplitude is zero Constructive interference: adding waves in a phase Destructive interference: adding waves out of phase Photoelectric Effect: a phenomenon in which electrons are ejected from the surface of a metal exposed to light at minimum frequency Threshold Frequency: the minimum frequency of light required to eject an electron from the surface of a metal Bohr’s Theory of the Hydrogen Atom: Bohr made the atomic orbitals. Heisenberg’s uncertainty principle: it is impossible to know simultaneously both the momentum and the position of a particle with certainty. Schrodinger: described electron density ( ) and atomic orbitals. Electron density is the probability where the electron is most likely to be Principle Quantum Number (n): designates size, collection of orbitals, only in integral numbers Angular Momentum Number (l): designates shape, depends on n, can be anything from 0 to (n1) l 0 1 2 3 Orbita s1 p3 d5 f7 l desig. Magnetic Quantum Number (Ml): specifies orientation Ms: is used to specify an electrons spin. Either positive or negative ½ **(n), (l), and (ml) all required to describe an atomic orbital Pauli Exclusion Principle: no two electrons in an atom can have the same four quantum numbers. Aufbau Principole: states that electrons are added to the lowest energy orbitals first. Hund’s Rule: the most stable arrangement of electrons is the one in which the number of electrons with the same spin is maximized. Noble Gas Core: shortcut for electron configurations. Only use if asked. **Exceptions on periodic table for electron configurations are ones ending in the d subshell like chromium and copper. ** Equationsyoushouldknow: ConstantsIwillprovide: K=C+273.15 SpeedoflightC=3.00x10 m/s D=M/V Planck’sconstanth=6.63x10 J s orKgm/s 2 23 c=λ NA=6.022x10 E=h 1”=2.54cm;1Km=0.6215mi;1lb=453.6g;1gal= 3.785L Chapter 4 John Newland: law of octaves, when elements arranged in order, every 8 element had similar properties Mendeleev: first periodic table Main Group Elements: s block and p block, 1a through 7a Noble Gases: group 8a and have a full p subshell Transition Metals: group 1B and 3B through 8B Lanthanides and Actinides: Make up the f block elements Valance Electrons: the outermost electrons of an atom that are involved in bonds and predict chemical properties. Metals: low ionization energies, cations, good conductors of heat and electricity, shiny, lustrous, malleable, and ductile Nonmetals: vary in color, not shiny, brittle, poor conductors of electricity and heat, have high electron affinities, anions Metalloids: elements with properties intermediate between those of metals and nonmetals Ionic Radius: the radius of a cation or anion, when an atom loses an electron, radius decreases. Isoelectronic Series: a series of two or more species that have identical electron configurations but different nuclear charges. Periodic Table Trends: **Study Periodic Table, Know first 18 Elements** Effective nuclear Charge: is the actual magnitude of positive charge that is “experience” by an electron in the atom Shielding: where an electron partially shielded from the positive charge of the nucleus by the other electrons Zeff increases from left to right Atomic Radius: the distance between the nucleus of an atom and its valence shell, or half the distance between adjacent, identical nuclei Increases from top to bottom and decreases from left to right Ionization Energy: is the minimum energy required to remove an electron from an atom in the gas phase Increases from left to right Takes more energy to remove core electrons Electron Affinity: is the energy released when an atom in the gas phase accepts an electron Increases from left to right meaning that it becomes more favorable to accept an electron
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