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CHM 111 Midterm Review Guide

by: Madeline Kaufman

CHM 111 Midterm Review Guide CHM 111

Marketplace > University of Miami > Chemistry > CHM 111 > CHM 111 Midterm Review Guide
Madeline Kaufman
GPA 3.98

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These notes cover the topics on the midterm exam.
Prin of Chm I
Study Guide
scientific, notation, Vocabulary, atom, elements, Chemistry, Energy, levels, periodic, Table
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Date Created: 09/18/16
CHEM MIDTERM REVIEW GUIDE CHAPTER 2: DATA ANALYSIS VOCABULARY  Measurement:  Accuracy: how close a measured value is to an accepted value  Precision: how close a series of measurements are to one another  Density: how tightly packed molecules are in a substance CONCEPTS  Scientific Notation: a number between 1 and 10 multiplied by 10 to a power o To multiply: multiply the numbers and add the exponents o To divide: divide the numbers and subtract the exponents o To add or subtract: make exponents equal then add or subtract  Metric System: 7 base units, common measurement system o Time- seconds o Length- meters o Mass- kilograms  p-pico= 10 -12  n-nano= 10 -9 -6  -micro=10  m-milli=10 -3  c- centi=10-2  k- kilo=103 6  M-mega=10  G-giga=10 9  New:::::  Hecto (h)  Deka (da) 10 grams in a dekagram  Deci; (d) ten decigrams in a gram  Significant Digits (including math problems): the number of digits recorded in an experiment o All non-zero numbers o All zeros between non zeros o Underlined #s not significant: 0.03440 or 560 or 56.00 th th o Addition/subtraction: the final answer has the least 10 or 100 place o Multiplication/division: least number of significant digits  Dimensional Analysis: a method of problem solving using unit factors to describe a matter o Conversion factor: a ratio of equivalent values used to express the same quantity in different units that always equal one FORMULAS  Density: d= m/v measured in grams per cubic centimeters  Percent Error (Provided): the ratio of an error to an accepted value CHAPTER 3: MATTER- PROPERTIES AND CHANGES VOCABULARY  Matter: anything that takes up space or has mass  Phase: differing states of matter o Solid: has its own definite shape and volume, tightly packed molecules o Liquid: has indefinite shape but definite volume o Gas: has indefinite shape and indefinite volume  Vapor: the gaseous state of a substance that is a solid or liquid at room temperature  Mass: how much matter an object has  Volume: how much space an object occupies  Element: a pure substance that cannot be separated into simpler substances by chemical or physical means  Compounds: a combination of 2 or more different elements that are combined chemically (water, salt)  Pure substance: matter with uniform, unchanging composition (salt, water)  Physical vs Chemical o Physical Properties: a characteristic that can be observed or measured without changing the samples composition (density, color, odor, taste, melting and boiling point)  Extensive: depends upon the amount of substance present (length, mass)  Intensive: independent of amount present (density) o Chemical Properties: the ability of a substance to combine with and change into 1 or more other substances (and inability to do so) o Physical Changes: altering a substance without changing its composition (crumple, bend) o Chemical Changes: (chemical reaction) altering a substance such that the result has different compositions and properties  Mixtures: a combination of 2 or more pure substances in which each pure substance retains its individual chemical properties o Homogenous Mixtures: doesn’t blend smoothly, individual substances remain distinct o Heterogeneous Mixtures: constant composition throughout the mixture  Filtration: a mean of separation of a mixture using a porous barrier to separate a solid from a liquid  Law of Conservation of Mass: mass is neither created nor destroyed in a chemical reaction, it is conserved  Law of Definite Proportions: regardless of the amount, a compound is always composed of the same elements in the same proportion by mass (percent by mass= mass of element/mass of compound x 100)  Law of Multiple Proportions: when different compounds are formed by a combination of the same elements, different masses of one element combine with the same relative mass of the other in a ratio of small to whole numbers FORMULAS  Pecent Mass: mass of element/mass of compound x 100  Avg Atomic Mass: (abundance1 x atomic mass1) +(abundance2 x mass2) STRUCTURE OF ATOM AND ELEMENTS Vocabulary  Atom: the smallest particle of an element that retains the property of that element o Atomic number: the number of protons in an atom (equal to the number of electrons) o Atomic mass: the average mass of isotopes of that element  Sum of (isotope mass x abundance) o Mass number: the number of protons plus the number of neutrons o Isotope: an atom with a different number of neutrons than protons  Period: horizontal row in the periodic table  Group: vertical column in the periodic table  Periodic Law:  Overall: Theories  Models of the atom: o Dalton: tiny, solid sphere o Thompson: plum pudding model. o Rutherford: electrons move around the nucleus  Famous Experiments o Thompson: shot a ray of particles that was bent by a magnet showing they had a charge and finds charge to mass ratio. o Rutherford: shot particles through gold foil, and it was repelled in some instances shown that they must have hit something dense- discovered the nucleus.  Mendeleev’s Periodic Table: Mendeleev was significant because he predicted the existence and properties of undiscovered elements.  Periodic table: o Representative vs Transition Metals o Group vs Period o Groups-  Halogens: group 17  Non-metals  Highly reactive  Noble gases: group 18  Highly unreactive since they have an octet of electrons  Alkali metals: all elements in group 1 except hydrogen  Very reactive  Alkaline earth metals: all elements in group 2  Very reactive  Transition: o Metals v Nonmetals v Metalloids (malleability/conductivity/luster):  Metals:  Shiny when smooth and clean  Solid at room temperature  Good conductors of heat and electricity  Nonmetals:  Generally gases or dull solids  Poor conductors of heat and electricity  Metalloids:  Have chemical and physical properties of both metals and nonmetals  Conduct well under certain conditions- good for use with computers o Solids v Liquids v Gases (in the periodic table): CHAPTERS 5 AND 6 ELECTRONS IN ATOMS AND Vocabulary  Quantum: small measured amount of energy which matter gains energy in (j x s)  Photon: massless packet of energy (light)- a particle that carries a quanta of energy  Atomic emission spectra: atoms only emit specific frequencies of light o When an atom loses energy, it emits light.  Bohr’s Model of the Hydrogen Atom: atoms have only certain allowable energy levels which is accurate. Wrongly believes that the electrons travel in orbitals around the nucleus.  Schrodinger: discovered atomic orbitals (3d region around the nucleus to predict the location of electrons) and developed an equation to do so  Configuration: o Energy levels: (n) distance from nucleus- 1,2, or 3  Sublevels: ( ) shape of region of possibilities  S: spheres, one orbital possible  P: dumbells, 3 orbitals possible  D: 5 orbitals possible  F: 7 orbitals possible  Amplitude: wave height from origin to crest or origin to trough  Wavelength (y) : the shortest distance between equivalent points on a wave (m)  Frequency (v): the number of waves that pass a given amount of points per second (measured in Hz or s )-1  Electromagnetic spectrum: encompasses all forms of electromagnetic radiation; o radio waves, micro waves, infrared, visible light, ultraviolet, xrays, gammarays o Electromagnetic radiation: energy exhibiting wave-like behavior  Ground state: the lowest energy state of an atom  Excited state: when an atom gains energy  Trends o Atomic radius: the radius of an atom  Decreases as you move across a period because the nucleus executes a higher positive force pulling the electrons in  Increases as you move down a group as the electrons fill in bigger energy levels o Ionization energy: the energy it takes to remove an electron from an atom  Increases as you move across a period because with closer and closer electrons to an octet, the more they do not want to lose their electrons  Decreases as you move down a group because with wider energy levels and more electrons, they are more difficult to hold on to o Electronegativity: how tightly the nucleus holds its electrons  Noble gases are not electronegative  Increases as you move across a period because the more valence electrons it has the more it wants to hold on to them to reach an octet  Decreases as you move down a group because with more electrons it does not want to hold on to so many CHAPTERS 8 AND 9 IONIC AND COVALENT BONDING Miscellaneous - Valence electrons: electrons in the outermost orbital of an atom that determine its chemical properties (s and p electrons beyond the last noble gas) - Law of Constant Composition: - - - Remember: the stronger the bond, the more bonds present, the shorter the bond Ionic Bonding - Ion: a charged atom o Cation: positively charged ion (metals, as they lose electrons to form octets) o Anion: negatively charged ion (nonmetals, as they gain electrons to form octets) - Ionic bonds: bonds between a metal and nonmetal of opposite charges formed by their attraction to one another o Type 1: between a metal and nonmetal  Name the metal then the nonmetal with –ide at the end o Type 2: between a transition metal and a nonmetal  Name the metal (with number in parentheses) and the nonmetal with –ide at the end  Transition metals that do not need number: o Zn is always +2 o Ag is always +1 o Al and Ga are always +3 o Type 3: between a metal and a polyatomic ion  Name the metal and then the ion with no ending change  Polyatomic ion: covalently bonded atoms with a charge  Nitrate:  Sulfate:  Carbonate:  Phosphate:  Hydroxide:  Acetate:  Chromate:  Ammonium: - Covalent Bonds: a bond between atoms in which the atoms share valence electrons o Molecules: covalently bonded compounds o Naming- name the first element with the prefix (except for mono on the first element) and then the second element with prefix and ending with ide  Prefixes: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca o Acids: covalent molecules that behave like ions in solution because they have a hydrogen that disassociates in solutions.  Binary: consist of hydrogen and an element  Naming: hydro+elementname+ic+acid  Oxyacid: consist of hydrogen and a polyatomic ion  Naming: polyatomic ion name, if it ends in ate, end with –ic if it ends in ite, end with –ous+ acid - Lewis Dott molecular structures o Terminal Element: outside elements (hydrogen always terminal_ o Central element: central element, least electronegative element o Exceptions to the octet rule  Boron is happy with 6 electrons  Hydrogen is happy with 2 electrons  Xenon is happy with more than 8 electrons o Resonance: a double bond can exist in any valid area of the structure - VSEPR Models: Valence shell electron pair models (3D) o Linear: 2 areas of electrons on central atom o Trigonal: 3 areas of electrons on central atom  Trigonal planar: 3 bonded pairs, 0 lone pairs on central atom  Bent: 2 bonded pairs, 1 lone pair on central atom o Tetrahedral: 4 areas of electrons on central atom  Tetrahedral: 4 bonded pairs on central atom, 0 lone pairs  Trigonal pyramidal: 3 bonded pairs on central atom, 1 lone pair  Bent: 2 bonded pairs on central atom, 2 lone pairs - Polarity: bonding between unlike atoms resulting in an unequal sharing of electrons o Ionic: if the difference in electronegativity is greater than or equal to 2 o Polar: if the difference in electronegativity is between .4 and 1.9 o Nonpolar: if the difference in electronegativity is between .1 and . 3 o Purely Covalent: if the difference in electronegativity is 0 o When a molecule has polar bonds, it is not necessarily a polar molecule. The more electronegative atom pulls the electrons to it but if the charges are symmetrical (trigonal planar, linear, tetrahedral), they cancel each other out and the compound is not charged and therefore not polar SYMBOLS OF ELEMENTS Symbols - H: hydrogen - Be: beryllium - N: nitrogen - He: helium - B: boron - O: oxygen - Li: lithium - C: carbon - F: fluorine - Ne: neon - Sc: scandium - Ag: silver - Na: sodium - Ti: titanium - Au: gold - Mg: magnesium - V: vanadium - Hg: mercury - Al: aluminum - Cr: chromium - Sn: tin - Si: silicon - Mn: manganese - Pb: led - P: phosphorus - Fe: iron - Br: bromine - S: sulfur - Co: cobalt - I: iodine - Cl: chlorine - Ni: nickel - Kr: krypton - Ar: argon - Cu: copper - Xe: xenon - K: potassium - Zn: zinc - Ca: calcium - Ba: barium - - MEMORIZE - Equations (including temp!) - Transition metal - Metric system prefixes - Transition metal exceptions - Emission spectra - - - QUESTIONS - 1. As an atom loses energy does it emit light? Or as an atom gains energy? When it loses! 2. Do we need to know Aufbau’s principle; each electron occupies the lowest orbital available, Pauli; 2 electrons may occupy a single atomic orbital only if they spin in opposite directions, Hund’s rule; single electrons with the same spin must occupy each equal energy level orbital before opposite spins occupy one. 3. SPDF SECTIONS on the periodic table? - - - 4. Solid v Liquid v Gases in periodic table? 5. Do we only need to know the polyatomics on the review guide? Or the ones we knew for the test?- - - ones from test! 6. What is ternary?- an compound with 3 elements 7. How do you show ionic bonding using electron dot structures? 8. What are the P and S exceptions to the octet rule? What about xenon? Don’t need to know 9. What is malleable? And why are metals malleable? They can be moldable- some are more malleable than others -


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