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Chemistry Notes 1

by: Isabell Notetaker

Chemistry Notes 1

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Isabell Notetaker
DMACC - Ankeny Campus
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Notes for Unit 1
General Chemistry 1
Class Notes
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This 10 page Class Notes was uploaded by Isabell Notetaker on Sunday September 25, 2016. The Class Notes belongs to at DMACC - Ankeny Campus taught by in Spring 2015. Since its upload, it has received 3 views. For similar materials see General Chemistry 1 in Science at DMACC - Ankeny Campus.

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Date Created: 09/25/16
Chemistry Chemistry- the study of matter and it’s properties - Changers of the matter and the energy’s associated with the changes Properties of Matter - Matter = mass+ volume - Properties of matter = characteristics that make matter unique - Physical properties o No change o No interaction with other substance o EX: melting point of metal, electric conductivity, density o When there is a physical change > effects physical properties – but not composition EX: Ice melts (state of matter changes) -= hardness, density, buoyancy, ability to flow. Ice (solid) > water (liquid) - Chemical properties- Changes when the substance interacts with another substance o Convert a substance into another substance EX: water electric current > hydrogen + oxygen States of Matter - Solid - Liquid - Gas Changes & Energy - Changes (both physical and chemical) involve energy changes - Energy- ability to do work - Work- ability to move something - An object doing work transfers energy to another object where work is done - Total energy = Potential energy + Kinetic energy - Potential energy- energy due to the position of an object relative to another - Kinetic energy- motion of the object - 2 Key Concepts= o 1 When energy is converted from one form to another it is conserved not destroyed o 2 Situations of lower energy are more stable & more favored in nature than higher energy states which are less stable Units and Conversion Factors - Measured quantities = number and a unit - Conversion factors are ratios to express a quantity in different units Measurement in a Scientific Study SI units Physical Quantities Unit Name Unit Abbreviation Mass kilogram kg Length meter m. Time Seconds s. Temperature Kelvin K Electric Current Amps A Amount of Substance Mole mol. Luminosity Candle cd. Length - Biological cells = micrometers o 1 biological cell = 10^-6 m - Atomic Scale= nanometers (10^-9 m) o = pedometers (10^-12 m) - Proteins = diometers =nm - Atomic diameter = 200 pm o Angstron 1(A with a circle above) = 10^-10 m Volume - Liters ( L) - Milliliter (mL) - ** Physicians use cubic decimeters to measure body fluid - 1 L= 1 dm^3 = 10^-3 m^3 Mass - Mass is a constant - Weight is a variable due to gravity - Example is in NOTES Density D=m/v m= dxv V=m/d - SI unit for density is kg/m^3 - Unit in chemistry = g/L or g/dm^3 or g/mL or g/cm^3 - Example in NOTES Temperature- how hot or cold something is relative to another. - Heat- energy that flows from hot to cold - Thermometer- heat from substance flows to the thermometer and the fluid expands. - 3 Temp scales; Celsius (C) , Kelvins (K), Fahrenheit (F) - Example in NOTES Celsius Kelvin Fahrenheit Boiling Point 100 C 373.15 K 212 F Freezing Point 0 C 273.15 K 32 F C > K C > F F > C EXAMPLE IN NOTES 0 C > 273.15 K 9/5 x C + 32 F – 32 x 5/9 The Atomic Theory - Structure of an Atom - Atoms are neutrally charged - The nucleus contains both protons and neutrons and has a positive charge - In order for an atom to be neutral the protons have to equal the electrons o Atomic #- (Z) the number of protons in the element. o Mass # (A) – Protons and neutrons - Proton mass- 1.67262 x10^-27 kg - Neutron mass – 1.67493 c 10^-27 kg - Electron mass- 0.00091 x10^-27 kg - Atomic Mass Unit (AMU) o Proton = 1 AMU o Neutron = 1 AMU o Electron = 0.00055 AMU - Atomic System = 117 elements, 90 in nature, rest (27) made in lab - Arranged by similar properties, same colum Isotopes - All atoms of an element have the same atomic # but different masses (different # of neutrons) - LOOK IN NOTES FOR EXAMPLES Periodic Table - Mendeleev - Organized by similar properties - Periodic Law- arranged by elements in order of increasing masses - Properties that would occur periodically - In the table each element has a box - Atomic # increases from left to right - The boxes are arranged in periods – horizontal rows (7) and vertical groups called columns (8 main) (8A) 2 left and 6 right B Group Elements (Between 2A and 3A) - Contain transition elements - 2 horizontal series of inner transition elements = lanthonides and actinides  In periodic system- o Metals – lower left side and middle system / Good conductors of heat and electricity o Non Metals- upper right side, generally gasses, brittle solids at room temp (r.t.) , Bromine is only solid at r.t., nonmetals (poor conductivity) o Metalloids- (semi metals) along staircase line, properties are inbetween metals and nonmetals. - Group 8A- Nobel Gasses – mostly unreactive - Group 1A- Alkali Metals- very reactive - Group 2A- Alkaline Earth Metals- fairly reactive - Group 7A- Halogens- very reactive nonmetals Ions and Periodic System  In chemical reactions o Metals > loose electrons >positive charge and are now cations o Nonmetals > accept electrons > negative charge anions  Main Group Elements o Metals > form ions = same # of electrons as the nearest nobel gas o Nonmetals > gain electrons > anions > same number of electrons as nearest nobel gas Moles and Molar Mass  Mole (mol) amount of material containing 6.02214 x 10^23 particals  Avagadros Number  One mole of atom, or ions, or molecules etc. corresponds to 6.02214 x 10^23 Bonding - Transferring of electrons from one element to another = ionic compound - Sharing of electrons = covalent compounds - Chemical bonds = forces that keep atoms together in a compound Ionic Compounds- Ions = charged particles made by gaining or losing electrons  Binary ionic compounds = 2 elements  Metal > lost e > cation  Nonmetal > gain e > anion o BOTH are transferring of electrons  EX: NaCl o Na (metal) [11e] > looses 1 e > Na+ [10e] (cation) o Cl (nonmetal) [17e] > gain 1 e > Cl- [18e] (anion)  -Sodium gives one electron to chlorine and stays together with Sodium. Oppositely charged ions held together through ionic bond and create crystallizing lattice (crystal appearance) * blue is Cl, yellow is Na Coulomb’s Law- energy of the attraction or repulsion between two particles. (energy is directly proportional) Energy [directly proportional] (charge 1 x charge 2) /(distance) Predicting # of electrons lost or gained - In ionic compounds – neutral  +=- cations = anions - Group A elements  Metals > loose electrons > [ions with same # of electrons as nearest noble gas]  Group 1A- loose 1 e  Group 2A – loose 2 e  Group 3A- loose 3 e  Nonmetals > gain electrons >[ions with same # of electrons as nearest noble gas]  Group 7A- gain 1e  Group 6A- gain 2e  Group 5A- gain 3 e  Group 4A- mostly shares in covalent bonds, a cation or anion from a single atom (monoatomic ions) Covalent Bonds= two nonmetals that share e EX: H 2 2 H atoms In covalent bonds - Each e no longer belongs to a particular atom, so we say the 2 e are shared by the 2 nuclei - THEY ARE NOT SEPERAT ATOMS - Atoms of different elements that share electrons Distinguish Ionic vs Covalent bonds - Covalent compounds consist of molecules (H O) 2 - Ionic compounds, no molecules, (NaCl) Polyatomic Ions - Many of ionic compounds= polyatomic ions - Two or more atoms bonded covalently = net charge + or – - EXAMPLE IN NOTES How to write a chemical formula 1. Show type and number of each atom in a substance a. Rules- for all ionic compounds names and formulas, positive ion first, then negative ion b. For binary ionic compounds; name metal, followed by name of anion and suffix –ide Compounds of Elements which form one Ion - Monoatomic ions = main-group elements o Alkali Metals- 1A- Li, K, Na, Rb, Cs, Fr [charge + 1] o Halogens- 7A- F, Cl, Br, I [ charge -1] o Group 2A- charge 2+ o Group 4A- exceptions – Sn+2 Sn+4 Pb+2 Pb+4 - For anions= ions charge = group number (A) – 8 o Ex: For S (sulfur) 6A-8=-2 Sn-2 - For Ionic compounds= +=- - The charge without the sign of one ions becomes the subscript of the other. (EXAMPLE IN NOTES) Naming Binary ionic compounds containing a metal that forms more than one cation Name of cation + Charge of cation Base of anion Suffix -ide + name + Systematic Name Metals- (positive charge) Chromium > Chromium (II) & (III) Cobalt > Cobalt (II) & (III) Iron > Iron (II) & (III) Copper > Copper (I) & (II) Common Name 2+ 3+ Chromium Cr - Chromous Cr - Chromic Iron Fe - Ferrous Fe 3+ -Ferric Compounds with Polyatomic Ions- Polyatomic ions- stay together as a charged unit Most Common Polyatomic Ions Cations NH 4charge +1) Ammonium H o (charge +1) Hydromium 3 Anions CH 3OO (charge -1) Acetate OH (charge -1) Hydroxide ClO 3 (charge -1) Chlorate NO 3charge -1) Nitrate SO (charge -2) Sulfate 4 MnO (4harge -1) Permanganate CO 3charge -2) Carbonate HCO (charge -1) Hydrogen carbonate 3 (bicarbonate) Cr2O 7charge -2) Dichromate PO (charge -3) phosphate 4 Oxoanions – a non-metal is bonded to one or more oxygen atoms - More oxygen atoms > nonmetal + _______ate - Less oxygen atoms > nonmetal + ______ite - EX: SO4 (charge 2-) = sulfate ; SO3 (charge 2-) = sulfite - EX: NO3 (charge 1-)= nitrate ; NO2 (charge 1- ) = nitrite  Halogens bonded to oxygen: o Least oxygen atoms > root + ____ite  EX: ClO (charge 1-) = hypochlorite o Most oxygen atoms > prefix per- +nonmetal +suffix –ate  EX: ClO4 (charge 1-) = perchlorate o One Fewer Oxygen > root + -ate  Ex: ClO3 (charge 1-) = chlorate o Two fewer Oxygen atoms > root + -ite  EX: ClO2 (charge 1-) = chlorite Hydrated Ionic Compounds 1 Mono- 2 di- - Hydrates = 3 Tri- specific number of water molecules in 4 Tetra- formula o Ex: MgSO4 = 5 Penta- Magnesium Sulfate o MgSO4 * (dot) 7 6 Hexa- H2O = Magnesium Sulfate 7 Hepta- heptahydrate o H2O is shown 8 Octa- after a dot + greek numeral before 9 Nona- the word hydrate o EX: CaSO4 *(dot) 10 Deca- 2 H2O = Calcium Sulfate dehydrate Acid Naming - Acids- compounds that release H+ (Hydrogen protons) - EX: HCl = hydrochloric acid - HCl > water > ionize (release H+) - HCl > gas state = molecular compound  Binary Acids o Hydro + nonmetal root + suffix –ic +acid o EX: HCl = hydrochloric acid  Oxoacids o Oxoanions  - ate in anion > -ic in acid  -ite in anion > -ous in acid  EX: sulfate > sulfuric ; sulfite > sulfurous o The oxoanions with prefixes  Hypo- & per- (they are retained in the acid)  Ex: BrO4 = perbromate > HBrO4 = perbromic acid Binary Covalent Compounds  First – lower group number +Next – higher group number+ suffix –ide o EX: NF3 = Nitrogen fluoride  If Both elements are in the same group > First – element with high period # + Next- Lower period # + suffix –ide o EX: Group 6A sulfer + oxygen o SO2 > sulfer dioxide o CoF2 > cobalt (II) fluoride o SF4 > Sulfer tetraflouride Ionic Bonds • metal + nonmetal Covalent Bonds • nonmetal +nonmetal Molecular Masses & Chemical Formula - Molecular masses – sum of atomic masses o EX: H2O = 2(atomic mass of H) + 1(atomic mass of oxygen) - Ionic Compounds – polyatomic ions o Molecular mass = # of atoms of each element inside parenthesis, multiplied by subscript outside parenthesis  EX: Ba(NO3)2 = Barium Nitrate (WORK SHOWN IN NOTES) The Mole Mol- Si unit for an amount of substance - Amount of substance that contain same # of entities of atoms  Converting o Mass (g) = amount in moles X (Molar mass / 1 mole) o Amount in Moles = Mass (g) X (1 mole/ molar mass) o EXAMPLES IN NOTES Determine mass percent from a chemical formula Mass percent of an element X = atoms of X x atomic mass of X Molecular mass of compound EXAMPLE IN NOTES Find mass of an element = mass of compound x mass of element Mass of mol compound Determine chemical formula from experimental data Empirical formula > diverted from mass analysis & lowest number of moles Pseudo Formula - Divide all subscripts by smallest subscript number - Multiply by smallest integer > to turn all subscripts into integers - EXAMPLE IN NOTESc Obtain chemical formula from mass percent composition Example in notes Elemental analysis of latic acid Example in notes Molecular Formula - Show us the actual number of atoms of each element in a molecule - Situates where empirical formula is same as molecular formula  For others, molecular formula will be a whole number of multiples of empirical formula - 1 part mass (H) : 16 parts oxygen (O) - Empirical formula= HO (17.01 g) - Whole # multiple =  Molar mass = 34.02 g = 2 (multiply empirical formula)  Empirical formula mass 17.01 How to Write and Balance chemical equations - The same number of atoms must appear on each side - EXAMPLE IN NOTES


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