JoyNotes

Stephen Beutel Mr. Littlejohn – Per. 7/8 Chemistry R Chemistry R Midterm Review Sheet Significant Figures o The #’s that are meaningful in a measurement or calculation o When measuring, measure to the level of uncertainty o That means whatever the device allows for, plus one extra space  Ex) A ruler measures to the 10ths of an inch, so you write your answer to the hundredths place. o Accuracy vs. Precision  Accuracy – how close you are to being correct  Precision – how close several measurements are to each other o Determining Sig Figs: Atlantic vs. Pacific Technique  Atlantic – • used when decimal is absent, so you start on the Atlantic side (right side) • work left until you hit the first non-zero • count first non-zero and everything to the left • Ex) 143,000 has 3 sig figs  Pacific – • used when decimal is present, so you start on the Pacific side (left side) • work right until you hit first non-zero • count first non-zero and everything to the right • Ex) 12,000. has 5 sig figs o Addition and Subtraction of Sig Figs  Answer cannot have more decimal places then the least # of decimal places you started with (ROUND!) o Multiplication and Division of Sig Figs  Answer must have same # of sig figs as the coefficient with the least # of sig figs (ROUND!) Percent Error o Shows how wrong or right you are o Tells how far your measured value is from correct answer o FORMULA: o o Always round to the tenths place!! Conversion Factors o A unique way of writing the # 1 o In the same system they are defined quantities so they have an unlimited # of sig figs  Numbers that you have directly (physically) counted have unlimited sig figs Scientific Notation o Makes it easier to express very large and very small numbers o Format:  32,800,000,000,000 = 3.28 x 1013  0.00000000073 = 7.3 x 10-10 Elements o The simplest form of matter o Cannot be separated into simper substances through chemical means Compounds o Substances that can be separated into simpler substances by chemical means Mixtures o A physical blend of 2 or more substances o 2 categories:  1) Heterogeneous (different parts are different)  2) Homogeneous (whole mixture is same) o Heterogeneous  Suspension: mixture separates into 2 or more layers o Homogeneous  1) Solution – has two parts • a) solute – the thing being dissolved • b) solvent – major component of solution, does the dissolving o water is the universal solvent  2) Colloid • opaque – particles are large but don’t settle o Separating Mixtures  1) Heterogeneous • physically separate compounds  2) Homogeneous • Fractional Distillation – boil liquid to produce a vapor, which is then condensed (returned to liquid states) and collected in a separate container What is Chemistry?? o The study of matter o Matter – anything that has a mass and occupies space o Matter can exist in 3 phases: Solid, Liquid, Gas  Solid Characteristics • Definite shape and volume  Liquid Characteristics • Definite volume • No definite shape • Particles touch each other but spread apart  Gas Characteristics • No shape or volume o Physical vs. Chemical Properties  Physical • Can be observed with senses • Can be determined without destroying object • Ex) density, color, odor • Changes in property just a change in phase  Chemical • Dictates how substance reacts with other substances • Original shape is always altered in observing chemical property • Ex) rust • Changes include color change, evolution of a gas, fire o Law of Conservation of Matter  Matter cannot be created or destroyed  Mass of reactants = mass of products Democritus – Greek philosopher who created the term “atom” Dalton’s Atomic Theory o 1) all atoms are composed of submicroscopic indivisible particles called atoms o 2) different elements are different atoms o 3) atoms chemically combine, simple whole-number ratios to form compounds o 4) elements cannot be changed chemically to another element o Dalton’s Atomic Diagram: Atoms o 3 particles make up an atom:  1) electron  2) proton  3) neutron o Electrons 0 -1 e o o o o  Negatively charged  No mass Protons  Positively charged1 1p  Has mass Neutrons  Neutral Counting  protons = electrons  atomic # = protons = electrons  rounded atomic mass = mass #  mass # = protons + neutrons  neutrons = mass # - protons  electrons = protons – charge Isotopes  Atoms of the same element that differ in atomic mass due to different # of neutrons (not protons) Theories o JJ Thomson  Discovered electron  Disproved Dalton’s indivisible atom  Plum Pudding: positive “dough” with negative “raisins” o Rutherford  Gold Foil  Small positively charged nucleus with all of the mass inside an empty space o Neils Bohr  Planetary Model  Electrons orbit nucleus in rings Quantized Energy o All electrons want to be at lowest possible energy level (ground state) o When electrons absorb energy, they jump to higher energy level (excited state) o When electrons go from excited to ground, it gives of quantized energy as light or heat Atomic Mass Unit (amu) o Atomic mass is not measured in grams, but rather in amu’s o 1 amu = 1/12 the mass of C-14 Average Atomic Mass and Weighted Atomic Mass o Atomic mass’s on the periodic table are weighted by the abundance of the naturally occurring isotopes o Avg amu = o Valence vs. Kernel Electrons o Valence electrons- outermost electrons (outer energy level)  Responsible for bonding with other elements o Kernel electrons – inside electrons, do nothing Rules for drawing Lewis Dot Diagrams o 1) start at either top, right, bottom, or left o 2) first two electrons should be paired, unless group 14 element-C or Si o 3) each electron that follows the first two, work around the symbol clockwise The Octet Rule o All elements want 8 valence electrons  1) lose electrons – become positively charged  2) gain electrons – become negatively charged Ions o Charged atoms  1) cation – positive ion  2) anion – negative ion o Lewis Dot Diagrams for Ions  No dots for cations How to write Ground State electron configuration o Copy it off the periodic table o Ex) Mg: 2-8-2 How to write Excited State electron configuration o Ex) N: 2-5 → 2-3-2, 1-6, 2-4-1 … Rules o Each energy level can only hold a certain # of electrons o Comes from equation 2n2 Energy # of Energy Levels Level electrons o Row of periodic table = 1 2 2 8  1) # energy levels 3 18  2) # sublevels 4 32 o an orbital is the most likely location of the electron o each orbital may hold a maximum of 2 electrons o Sublevels: Sublevels # Orbitals Max # of electrons S 1 2 P 3 6 D 5 10 F 7 14 Aufbau Principle o Electrons enter orbital of lowest energy first o Order to memorize:  1s22s22p63s23p64s2 o Applying it!  Ex) F: 1s22s22p5  Ex) Be: 1s22s2 Orbitals o An orbital can contain a max of 2 electrons o Filled orbitals contain electron with opposite spin: o Ex) Flourine: Nuclear Chemistry o Radiation – the emission of a particle or ray from an unstable nucleus o Radioisotopes – unstable nuclei that give off radiation o Instability is caused by an imbalance of protons and neutrons Nuclear Penetrating Power o Alpha < Beta < Gamma Types of Nuclear Equations o 1) Natural Transmutation  one element changes into another  always only one thing on left side of reaction o 2) Artificial Transmutation  occurs only in a lab  one nucleus bombarded by something else  always two things on left side of reaction o Balancing Nuclear Equations  The sum of mass on reactants side must equal the sum of the mass on products side o *Note: in all nuclear reactions, a very small amount of mass is converted into a very large amount of energy Half Lives o Amount of time it takes for ½ of a radioisotope to decay o Half lives do not change, and are constant o **Matter is not destroyed!! 1 element transmutates into another! o Determining half lives  Everything is on Reference Table N o Steps to solving half life problems:  1) determine # of half lives • # ½ lives = total time/ ½ life  2) draw # of arrows from step 1  3) cut initial sample in ½ at each arrow Fusion o Brings together 2 small nuclei to create a slightly larger nuclei, and a large amount of energy o Very difficult to do because positive nuclei repel each other Uses of Radioisotopes o C-14 for radioactive dating and organic reactions o Tracer – goes over the reaction pathway o Iodine-31 for thyroid diagnosis o Cobalt-60 to irradiate food o U-238 for geologic dating Medical Radioisotopes o Must have short half lives to be quickly eliminated from the body Nuclear Tidbits o Bismuth (element 83) is the last non-radioactive element o Elements 84+ are automatically radioactive o Uranium (92) is the last naturally occurring element Review o Fission reactions are always artificial, big to small o Fusion reactions always artificial, 2 small identical things added o Neutron cannot be sped up in particle accelerator because no charge History of Periodic Table o Dimitri Mendeleev was first person to do a good job of organizing the elements o His table was based on atomic mass o Henry Moseley made the modern table based on atomic # Periodic Table o Rows = Periods (there are 7) o Columns = Groups/Families (there are 18) o REASON: elements in same group behave similarly, same period act differently o Noble Gases = group 18 o Alkali Metals = group 1 o Alkaline Earth Metals = group 2 o Halogens = group 17 o Transition Metals = groups 3-12, form colorful salts o Inner Transition Metals = “f-block” elements Trends o Ionization Energy (I.E) – amount of energy required to remove an electron  high in non-metals  low in metals (metals are losers) o Eletronegativity (eneg) – desire for electrons  high in non-metals  low in metals o Atomic/Ionic Radius – size of an atom/ion o Ions:  Cation (+ ion) – loses electrons, smaller than their atoms  Anion (- ion) – gain electrons, larger than their atoms o Trend #1:  I.E increases across a period as elements get more metallic  I.E decreases as you move down a group as elements get more metallic o Trend #2:  Eneg increases across a period  Eneg decreases as you move down a group o Trends #3 and #4:  Atomic/ionic radii decrease across a period because # protons (nuclear charge) increases and draws electrons closer.  Atomic/ionic radii increase down a group because more energy shells o Elements  Metals • Shiny, malleable, ductile, conduct electricity  Metalloids • Have properties of both  Non-metals • Dull, don’t conduct electricity, brittle Oxidation #’s o A fancy term for charge (oxidation # = charge) o Rules:  1) hydrogen is always +1 in compounds unless in a metal hydride (then -1)  2) oxygen is always -2 in compounds unless in a peroxide (then -1)  3) the oxidation # of an ion is whatever the charge is for the ion  4) the sum of the oxidation #’s in a polyatomic ion is whatever the charge is on that ion  5) alkali metals are always +1 in a compound and alkaline earth metals are always +2 in a compound  6) substances that are on their own have an oxidation # of 0  7) the sum of the oxidation #’s in a compound is 0 Binary Ionic Compounds o All ionic compounds consist of a metal, and one or more nonmetals o Naming cations  Name the metal followed by word cation o Naming Anions  Change root ending to “-ide,” followed by word “anion” o Naming Binary Ionic Compounds  Name the metal (without the word “cation”)  Name the nonmetal with “-ide” but not “anion” Formula Writing o 1) write out ions; for nonmetals choose first charge o 2) see if sum of ions = 0; if yes, formula = 1:1 ratio o 3) if sum doesn’t = 0, criss cross charges and simplify to lowest whole # ratio Bivalent Metals o Metals that have more than one # in top right Stock Method o If metals have multiple charges, you must specify which charge you used with Roman Numerals (Roman Numeral = charge) o Naming with this:  Always check the metal  If multiple charges exist, must use Roman Numeral Polyatomic Ions o All are shown on Table E o Most polyatomic ions end in “-te” o Formulas:  If the polyatomic ion receives a subscript, you MUST place a parenthesis around the polyatomic ion Binary Covalent Compounds o 2 or more nonmetals = covalent Prefix # of Atoms o binary = 2 elements Mono 1 Di 2 Naming Tri 3 o They all end in “-ide” Tetra 4 o Never use “mono-“ on first Penta 5 element Hexa 6 o Covalents do not reduce Hepta 7 Reactions Octa 8 o Synthesis – 2+ reactants Nona 9 become one product (marriage) Deca 10 o Decomposition – 1 reactant becomes 2+ products (divorce) o Single Replacement – an element on its own switches with an element in a compound (cheating) o Double Replacement – two compounds and cations switch places (wife swap) Balancing Equations o Don’t ever touch the subscript o You can change the coefficients (must be in whole # ratio) o Use a “T-bar” Bonding o To achieve the octet rule (gain, lose, share) o Forming a bond releases energy o Breaking a bond, energy is absorbed Ionic Bonding o Transfer/Share Electrons  Transfer from metal to nonmetal  Metals lose, nonmetals gain o Eneg or exceptions  Eneg difference = greater than 1.7 o Types of atoms  Metal and nonmetal o Properties  Hard, crystal solids  High melting pts  Conduct electricity when melted or dissolved in water  No conductivity as a solid o Examples  NaCl, Ca(NO3)2, MgCls, Li2S o Ionic Character  The greater the difference in eneg, the more ionic you are Nonpolar Covalent o Transfer/Share of Electrons  Equal Sharing between 2 nonmetals o Eneg  Eneg diff = 0 – 0.4  Exception – compound automatically nonpolar if symmetrical o Types of atoms  Diatomic elements (H2, N2, O2 …) o Properties  Low melting pts  does not conduct electricity  solid, soft, brittle  will dissolve other nonpolar compounds o Examples  CO2, CH4 Polar Covalent o Transfer/Share of Electrons  Unequal share based on eneg o Eneg  Eneg diff = 0.4 – 1.7  Not symmetrical o Properties  Not symmetrical  Higher boiling pt than nonpolar covalent, but still low  No electricity  Contain a dipole (region of positive and negative charge) Metallic Bonding o Transfer/Share of Electrons  Neither o Eneg  Nothing o Types of Atoms  METALS! o Properties  Properties of metals o “sea of valence electrons” Network Covalent o Transfer/Share of Electrons  Sharing o Eneg  Nothing o Types of Atoms  Must contain carbon or silicon  Nonmetallic o Properties  Hardest substances known  Extremely high melting pts  Repeating crystal lattice accounts for properties o Examples  Diamond (pure carbon)  Sand (SiO2) Coordinate Covalent o Transfer/Share of Electrons  Sharing but one element donates both electrons to bond o Eneg  Nothing o Types of Atoms  Polyatomic ions • 1) Ammonium • 2) Hydronium Molar Mass (gram formula mass) o GFM = mass of a compound expressed in grams instead of amu’s o Calculating:  1) determine # of each type of atom in molecule  2) multiply the # of atoms by the atomic mass • rounded to nearest whole #, except Cu and Cl  3) Add up your products Moles o What is a mole?  A # = 6.02 x 1023 o Conversion factor  1 mole = gfm  unlimited sig figs (determined by what’s given) o Volume Conversion  1mol of any gas = 22.4 L at STP o Molecule Conversion  1 mol = 6.02 x 1023 molecules Percent Composition o Percent by mass of each element in a compound o % comp: o Steps:  1) determine gfm = whole  2) find mass of element in question = part  3) use formula! Empirical Formula o Tells you the lowest whole # ratio of elements within a compound o To convert from molecular to empirical, simply reduce o Calculating:  Percent Comp given  1) convert % to grams (x % = x g)  2) convert grams to moles (dimensional analysis)  3) divide answers to #2 by smallest whole # in step 2 If you get … Multiply by … .33 3 .25 4 .5 2 .67 3 Molecular Formula o 1) determine empirical formula o 2) get gfm of empirical formula o 3) solve for molecular formula  (Molecular mass / Empirical Mass) = whole # o 4) multiply subscripts of empirical formula by whole # from step 3 Stoichiometry o A combination of the mole and balancing equations o Similar to following a recipe o Steps:  1) check if equation is balanced  2) set unknown equal to what is known  3) use dimensional analysis to solve for unknown 

Document Info

Posted By:

Stephen Beutel

Date:

Thursday, January 24, 2008

School:

Clarkstown High School South

Class:

Chemistry 10R

Tags:

• Review Sheet
• Chemistry

About this Document

This is my review sheet for my Chemistry 10R Midterm that contains everything I have learned so far.