Chemistry 111 Chem 111
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Date Created: 09/09/16
Chemistry Notes Lecture and Textbook Chapter 1 1.1 Vocabulary: ● Science: knowledge gained through experience or experiment (also known as empirical knowledge) ● Chemistry: study of matter and its interactions with other matter and energy ● Scientific method: investigations that are guided by theory and earlier experiments ● Hypothesis: possible explanation for an event ● Law: statement that summarizes a large number of observations ● Theory: an explanation of the laws of nature Important to Know: ● Experiments are the foundation of all sciences ● There isn’t one set scientific methodit adapts and changes depending on the individual experiment ● Difference between law, hypothesis and theory: ○ A law summarizes a large number of observations, but provides no explanation ■ Ex: Newton’s First Law An object at rest will stay at rest. An object in motion will stay in motion unless a force is acted upon it. ○ A hypothesis is a possible explanation for an observation ■ Ex: The rolling object eventually stops because gravity is acting upon it. ● Hypotheses must be testable and falsifiable ○ A theory is a known and accepted explanation of the laws of nature ■ Ex: Earth’s gravity, combined with friction, act on the rolling object to eventually stop its’ motion. ■ Ex: Heliocentric Theory: most scientists today accept the theory that the Earth and planets revolve around the Sun. 1.2 Vocabulary: ● Matter: anything that has mass and takes up space ● Mass: measures quantity of matter in an object ● Weight: force of attraction between an object and Earth ● Property: anything observed or measured about a sample ○ Extensive Property: dependent on the size of the sample ■ Ex: volume, mass, weight, ○ Intensive Property: independent on size of the sample ■ Ex: color, melting point, density ○ Physical Property: property that can be measured without changing the chemical composition of a substance ■ Ex: color, mass, density ○ Physical Change: a change that occurs that does not alter the chemical composition of a substance ■ Ex: melting ○ Chemical Property: describes the reactivity of the sample ■ Ex: explosiveness, flammability ○ Chemical Changes: a reaction the changes the chemical composition of a substance ■ Ex: rusting, burning ● Substance: material that is chemically consistent throughout ○ Elements: substances that cannot be broken down into smaller, simpler substances ○ Compound: substances that are composed of two or more elements that can be broken down ● Mixture: combination of two or more substances that can be broken down by physical methods into simpler substances ○ Heterogeneous mixture: composition of substance changes from one part to the other ○ Homogeneous mixture: composition of substance is consistent throughout ■ Also known as a solution ● Alloy: solid solution consisting of a metal and another substance Important to Know: ● Difference between mass and weight: ○ Mass is how much matter an object is composed of, and it stays constant. Whereas weight can change depending on location because it is the amount of attraction between an object and the Earth. ■ Ex: We weigh less on the moon because there is less gravitational pull, therefore less attraction, however we still have the same mass. ● Elements are essential to all matter ● A compound always contains the same ratio of the same elements ○ Ex: any sample of water will always be composed of 11.2% H and 88.8% O 1.3 Vocabulary ● Accuracy: agreement of the measured value with the true or accepted value ● Precision: agreement among repeated measurements Important to Know: ● Four aspects of an experiment that are crucial: ○ Measured object ○ The value ○ The units ○ The reliability of the measurement ● Accuracy and precision are important to the reliability of an experiment/measurement ○ Accurate data will be close to the true value while precise data will be close in value with a small range ● Significant Figures ○ All known digits of a measurement are presented in addition to an estimated digit ○ Sig Fig Rules: ■ All nonzero digits are significant ● Ex: 12378.98 has 7 sig figs ■ Zeros between two nonzero digits are significant ● Ex: 5609 has 4 sig figs ■ In a number with no decimal, the trailing zeros may or may not be significant. ● Ex: the number 643,000 has at least 3 sig figs ● Remove doubt by using scientific notation 5 ○ Ex: 643,000→6.43 x 10 (now it has only 3 sig figs) ■ Leading zeros are insignificant ● Ex: 0.00045 has 2 sig figs ■ The uncertainty of a measurement is at least +/1 unit on the last reported digit ○ There are numbers that are not limited to significant figures ■ Tallied numbers ● There are 400 chickens ■ Defined Numbers ● 1 dozen= 12 ■ Powers of 10 in exponential notation 4 ● 1.33 x 10 ○ In addition and subtraction, the difference/sum should have the same number of decimal places as the number with the LEAST amount of decimal places ■ Ex: 43.0120.5=22.5 ○ In multiplication and division, the product/quotient should have the same number of sig figs as the number with the LEAST amount of sig figs ■ Ex: 25 x 56.9= 1,422.5 →1.4 x 10 ○ Rounding Methods: ■ <5, round down ■ >5, round up ■ 5, round to even ● 23.5→23.6 1.4 Vocabulary ● Derived units: combinations of base units ● Conversion factor: fraction where the numerator is a quantity equal to the quantity in the denominator ○ Ex: 1kg/ 1000g ● Density: defined as mass per unit volume ○ d = m/v Important to Know: ● The SI system was adopted by scientists to have a universal science language; allows for easier conversions ○ Base Units in the SI System: Quantity Unit Abbreviation Length meter m Mass kilogram kg Time second s Temperature kelvin K Amount mole mol Electric current ampere A Luminous Intensity candela cd ● Prefixes in the SI System ○ Prefix Abbreviation Meaning mega M 10 3 kilo k 10 centi c 10 milli m 10 6 micro µ 10 nano n 109 12 pico p 10 ● Conversion among derived units 3 ○ Ex: volume = length ■ The base unit for length is meters, so the unit for volume would be m 3 ○ Two important conversion factors are liter and milliliter: ■ 1 L= 1000 mL= 1000cm 3 3 3 ● Density is an intensive property and is usually expressed as g/cm or kg/m ● Temperature Conversion factors: ○ Tf = Tc x (1.8℉/1.0℃) + 32℉ ○ Tc = (Tf32℉) x (1.0℃/1.8℉) ○ Tk = Tc +273.15 Chemistry Notes: Lecture and Textbook Chapter 2 2.1 Vocabulary ● Atoms: smallest unit of an element that still maintains all the properties of the element ● Element: composed one only one type of atom ● Compound: contains two or more types of atoms; relative number of atoms in the compound are the same Important to Know: ● Dalton’s Atomic Theory: ○ Atoms are the smallest units of an element that maintain the properties of that element ○ An element is composed of only one type of atom ○ A compound contains atoms of 2+ elements, where the relative number of atoms of each element is always the same ○ Atoms don’t change identity in a chemical reaction ● Law of Conservation of Mass: mass does not change in a chemical reaction illustrated by Dalton’s fourth law ● Law of Constant Composition all samples of a pure substance will have the same elements in the same proportions illustrates Dalton’s third law ● Law of Multiple Proportio in every compound formed by the same element, the masses of one element that combine with a fixed mass of a second element are always a ratio small numbers 2.2 Vocabulary ● Electrons: subatomic particles that are negatively charged ● Protons: subatomic particles that are positively charged ● Neutrons: subatomic particles that have no charge ● Nucleus: center of atom Important to Know: ● In the 1800s J.J. Thomson found that applying controlled magnetic and electric fields to cathode rays resulted in particles that are negatively charged. ○ He called this particle and electron ● Robert Milliken was the first to accurately measure the charge of an electron ● Atoms are charged neutral, meaning that there must be an equal amount of positive charges (protons) and negative charges (electrons). ● Ernest Rutherford discovered the alpha particle (????) ○ Has a charge of 2+ and a mass four times that of hydrogen ○ Rutherford’s experiments lead him to hypothesize that most of the atom’s weight is found in an extremely small central core (the nucleus) ○ Each element has a different positive charge on the nucleus ● Electrons have almost no effect on mass ● Protons account for about half of the mass of an atom, A neutral particle, called a neutron, (no charge) accounts for the rest of the mass. We rely on protons and neutrons for mass of elements. 2.3 Vocabulary ● Atomic Number (Z): number of protons of an atom (gives element it’s identity on the periodic table) ● Mass Number (A): sum of protons and neutrons of an atom ● Isotope: elemental atoms that have the same amount of protons and electrons, but different numbers of neutrons in the nucleus ● Ions: atoms that gain or lose electrons ○ Cation: an atom that has a positive charge because it LOST an electron ○ Anion: an atom that has a negative charge because it GAINED an electron Important to know: ● The number of protons determines the identity of the element ○ Ex: Hydrogen has 1 proton, so it’s atomic number is 1. ● About 75% of naturally occurring elements have two or more stable isotopes ○ Ex: Hydrogen has no neutrons, Deuterium has 2 neutrons, and Tritium has 3 neutrons. ● Notation for writing specific isotopes: ○ AX Z ○ A= atomic number ○ Z= mass number ○ X= element ○ Ex: Write the symbol for the isotopes with 11 protons and and 12 neutrons. 23 ■ 11a ● When writing symbols for ions, it is important to include the charge, unless the charge is +/ 1. The number may be left out. ○ Na+ ○ You determine charge based on the Group number from the periodic table. ○ Group 1 has a 1+ charge, Group 2 is 2+,Group 3 is 3+, etc. ○ The transition metals are excluded for now 2.4 Vocabulary ● Isotopic mass: mass (u) of an isotope ● Atomic mass: weighted average mass of the naturally occurring isotopes of that element Important to Know: ● Scientists have agreed upon a relative mass scale to express the masses of atoms [atomic mass unit (u)] 12 12 ○ u= 1/12 the mass of a C (because the mass of 1 C atom is exactly 12 u) or 1 27 u= 1.66054 x 10 kg ○ The masses of protons and neutrons in an atom are approximately 1 u ● Because electrons have a very small mass, the overall mass of an atom is not a whole number ○ When the atomic mass is rounded to a whole number, it equals the mass number ● A mass spectrometer measures the masses and relative abundance of isotopes in a sample ● About 75% of elements occur in nature as isotopes. Usually the relative abundance of each isotope is the same throughout nature as well. ○ Atomic mass = fractionA x isotopic massA + fractionB x isotopic massB 2.5 Vocabulary ● Periodic Table: arranges elements according to similar properties ○ Period: horizontal row ○ Group (family): vertical column ● Metal: found on the center and left side of periodic table; shiny, good conductor ● Nonmetal: top right of periodic table; nonconductor ● Metalloid: found on staircase separating metals and nonmetals on the periodic table; properties of metals and nonmetals Important to Know ● Periodic table proposed by Dimitri Mendelev and Lothar Meyer, both working independently. ● Atomic number and atomic mass are found on the periodic table ○ 8 represents the atomic number and 15.99 represents the atomic mass. ● Representative elements are those found in Groups 1A8A (12, 1318) ● Transition Metals: elements in the B Groups (312) ● Inner Transition Metals: two rows of metals at the bottom (also called lanthanides and actinides) ● Group 1A metals are called Alkali metals ○ Soft, reactivity decreases down the group ● Group 2A metals are called Alkali Earth Metals ● Group 7A Nonmetals are called Halogens ○ Reactivity decreases down the group, also called “saltformers” ● Group 8A gases are called Noble gases ○ Stable 2.6 Vocabulary ● Molecules: combination of atoms that act as one particle ○ Diatomic molecules: simplest molecules; contain only 2 atoms ○ Ex: H2 ● Molecular Compound: when two types of atoms form a molecule ● Molecular formula: shortened description of molecule that includes the number of each type of atom in the compound ● Molecular Mass: sum of the atomic masses of all atoms present in the molecular formula Important to Know: ● Molecular compounds are usually from nonmetals ● How to calculate atomic mass: ○ C4 3 ■ Take atomic mass from periodic table and multiply by how many atoms of the element are present ● C: 12.01 x 4 = 48.04 ● H: 1.008 x 3 = 3.024 ■ Add the two products together ● 48.04 + 3.024 = 51.06 2.7 Vocabulary: ● Empirical Formula: used for ionic compounds; uses the smallest whole number subscripts to show relative numbers of ions ● Ionic Compound: composed of cations and anions joined together ● Monatomic Ions: single atoms that have lost or gained electrons ● Polyatomic Ions: group of particles with a charge that act as one particle ● Formula Mass: sum of atomic masses of atoms in the empirical formula Important to Know: ● Ionic compounds form from joining of metals and nonmetals ● In an empirical formula, the charges of the anions and cations must balance each other out ○ Ex: BaCl 2 ○ Ba has a 2+ charge and Cl has a 1. You need two Cl atoms to balance the charged ○ Position of elements in the periodic table can help determine charges, as stated earlier ● Polyatomic Ions to Know: ○ Acetate: CH3O 2 ○ Carbonate: CO 3 ○ Hydroxide: OH ○ Nitrate: NO3 ○ Phosphate: PO 4 2 ○ Sulfate: SO4 2 ○ Sulfite: SO3 ○ Example of Polyatomic Ion Equation: ■ Write the equation for Sodium Hydroxide ● Na has a charge of 1+ ● We know hydroxide (OH) has a charge of 1 ● So the equation is NaOH ○ Always write the cation first ● Formula Mass ○ How to calculate formula mass: ■ NaOH ● Multiply the mass of each element by how many atoms of the element are present ○ Na: 1 x 22.99= 22.99 ○ O: 1 x 15.99= 15.99 ○ H: 1 x 1.008=1.008 ● Add all the products together ○ 22.99+15.99+1.008=39.99 2.8 Vocabulary ● Chemical nomenclature: system for naming compound ● Hydrocarbons: organic compounds that only contain hydrogen and carbon ● Acid: compound that when reacting with water, produces a salt Important to Know: ● Naming Ionic Compounds: ○ Write cation first ○ For monatomic ions the cation name stays the same, but the suffix on the anion changes to “ide” ■ Ex: KCl→ Potassium Chloride ■ Ex: NaBr→ Sodium Bromide ● Naming Transition Metals ○ Most metals can form more than one cation (excluding those in Groups 1A, 2A, 3B) ○ Those metals not in Groups 1A, 2A, and 3B must include a Roman Numeral in parenthesis equal to the charge of the ion ■ Ex: FeC3 → Iron (III) Chloride ■ Chromium (III) hydroxide → Cr(OH3 ● Naming Acids ○ Acids will always begin with an “H” (HCl= hydrochloric acid) ○ If the anion end with “ide” change the suffix to “ic” ○ Add the prefix “hydro” to the cation ○ Follow with the word acid ■ Ex: HCN→Hydrocyanic acid ○ For polyatomic ions: ■ If the anion ends in “ate” change the suffix to “ic” ■ If it ends in “ite” change the suffix to “ous” followed by acid and drop the prefix ● Ex: Sulfite will change into sulfuric acid ● Naming Molecular Compounds ○ Some molecular compounds have other names, like water ○ The way to name molecular compounds is similar to ionic compounds, with the suffix “ide” ○ There are some rules for the ordering of elements ■ The farthest left element on the periodic table is named first ■ Within groups, the element closer to the bottom is named first ■ If combining H with elements in Group 6A and 7A, H is named first. ■ Oxygen is second, unless being paired with Fluorine ○ Prefixes are used when naming molecular compounds to determine how many of each atom is present ■ Mono: one ■ Di: two ■ Tri: three ■ Tetra: four ■ Penta: five ■ Hexa: six ○ Ex: N3 4→ Dinitrogen tetrafluoride ○ If the first element being listed in the name has only one atom of that element present, you may exclude a prefix ■ Ex: CO2 arbon Dioxide ● Alkanes and Cycloalkanes are hydrocarbons ○ Alkanes have this general formula: n 2n+n ■ Ex: ● C4 10 ○ Cycloalkanes are the same as alkanes except they are in the shape of a ring with its carbon atoms ■ Formula: C H n 2n ■ To name alkanes and cycloalkanes you count the number of carbon atoms and add the corresponding prefix. The suffix with be “ane” ■ For cycloalkanes add the prefix “cyclo” ■ Prefixes: ● Meth: one ● Eth: two ● Prop: three ● But: four ● Pent: five ● Hex: six ■ Ex: ● From the picture above, the name would be Butane ● If they were arranged in a ring, it would be cyclobutane ○ Functional Groups: ■ Alcohols and ethers can be attached to hydrocarbons ● Alcohols will have an OH group and ethers will have the COC group ○ Alkanes and cycloalkanes sometimes have hydrogen atoms that are removed and replaced by other atoms. These other atoms are called substituents ■ Ex: an alkyl group is an alkane group with a base alkane name and an yl suffix ● Methyl ■ Halides are also substituents ● Cl would be chloro ■ You acknowledge the substituent by numbering it on the chain ● Ex: CH4H3Cl)CH ○ 2 Chloropropane 2.9 Vocabulary ● Electrolyte: substance that, when dissolved in water, forms ions ● Nonelectrolyte: substance that does not conduct electricity and forms a neutral molecule when dissolved in water ● Dissociate: when ionic compounds break down into smaller units Important to Know: ● Ionic compounds are usually composed of metals and nonmetals ● Molecular compounds are usually bonds between nonmetals ● Solid ionic compounds do not conduct electricity, therefore many ions are also electrolytes. When they dissolve in water they dissociate and can conduct electricity ● Most molecular compounds cannot do this Chemistry Lecture and Textbook Notes Chapter 3 3.1 Vocabulary ● Stoichiometry: study of quantitative relationships with substances in chemical reactions ● Chemical equation: describes identities and amounts of reactants/products in a reaction ● Reactants: substances that start a reaction (what is consumed) ● Products: substances formed ● Neutralization: reaction of an acid and a base to form a salt and water ○ Acid: dissolves in water that forms hydrogen ions ○ Base: dissolves in water that forms hydroxide ions ○ Salt: ionic compound with a base cation and an acid anion ● Combustion reaction: process of burning ● Oxidationreduction: reaction where electrons are transferred from one species to another ○ Oxidation number: can be assigned to each element in a substance based on rules ○ Oxidation: loss of electrons ○ Reduction: gain of electrons Important to Know: ● Chemical equations MUST be balanced ○ 2C + 2O2 CO2 ○ Add coefficients in from of reactants and products to make sure the same number of each element are on both sides of the equation ● Example of Neutralization ○ HCl + NaOH →NaCl + H O2 ■ HCl is the acid, and NaOH is the base ■ In a neutralization reaction water is always a product, so just take the remaining elements to form the salt ● Combustion ○ Hydrocarbons + oxygen ○ The products are always CO2 d H2 ○ CH + 2O→ O + 2H O 4 2 2 2 ● OxidationReduction ○ Rules ■ An atom by itself has an oxidation number of 0. ■ Monatomic ions have an oxidation number equal to its charge ■ F has an oxidation number of 1 when it’s with other elements. O is 2. H is +1 with nonmetals and 1 with metals. Other halogens are 1. ■ The sum of the oxidation numbers must equal the charge ○ Ex: ■ CO 2 ● We know the oxidation number of O is 2. However because there are two O molecules, the charge is 4 (2 x 2). ● Because the charges must equal, the oxidation number of C must be +4. 3.2 Vocabulary ● Mole: amount of substance that contains as many entities as the number of atoms ● Molar mass: mass of one mole of a substance Important to Know: 12 ● Avogadro's Number is the determined number of C atoms in 12g. 23 ○ Equal to 6.022 x 10 ○ Allows for conversion of moles and atoms 23 ■ Multiply the moles of a substance by 6.022 x 10 and that is how many atoms are present 23 ● One mole of anything contains 6.022 x 10 entities 23 ○ 1 mol H = 6.022 x 10 atoms of H ● Molar mass ○ The molar mass (in g/mol) of an element is equal to the atomic mass of the element ○ Used as the conversion factor between mass (g) to amount (mol) ○ Ar’s mass is 39.95 u. It’s molar mass is 39.95 g/mol 3.3 Important to Know ● Mass percentage ○ The composition of compounds is constant, so we can use this to determine the mass of each element of every sample ○ Simply find the atomic mass of the sample, then divide the individual element’s weights by the atomic mass and multiply by 100. ■ Ex: CO2 ■ Atomic mass: 12.01 + (15.999 x 2) = 44.01 ■ Carbon’s weight is 12.01 so, 12.01/44.01=0.273 ● 0.273 x 100= 27.3% Carbon ■ You can repeat that process for Oxygen, or simply subtract carbon’s percentage from 100. ● 10027.3= 88.7% Oxygen ● In combustion reactions the masses of C and H are determined by the masses of water and carbon dioxide formed ● Empirical formula ○ What is the empirical formula of a compound that is 70.57% O and 29.43% C? ■ Divide each element’s atomic mass by its percentage in the compound ● 70.57/15.999 and 29.43/12.01 ○ O: 4.41 ○ C: 2.45 ● Divide by the smallest number ○ O: 4.41/2.45= 1.8 ○ C: 2.45/2.45= 1 ● CO1.8does not work, so multiply the subscripts until they are closer to whole numbers ● C22 ● Molecular Formula: ○ The molecular formula must be a whole number multiple of the empirical formula ■ Take the sample’s actual weight given by the problem and divide it by the empirical formula’s molar mass ■ From above, the mass would be 56.018. 300 g/mol was measured. ■ 300/ 56.018= 5.3 ● (C22 5 ● C H 1010 3.4 Important to Know: ● Guideline for Stoichiometry ○ Write balanced equation ○ Calculate # of moles of species for which the mass is given ○ Use coefficients in equation to convert moles of given substance into moles of substance desired ○ Calculate mass of desired species ● Theoretical yield ○ Maximum amount of product that can be obtained from a chemical reaction based on amount of reactants 3.5 Vocabulary ● Limiting reactant: reactant that it totally used up when a chemical reaction takes place ● Actual yield: mass of product isolated in a reaction Important to Know: ● From the given amount of grams per element in the problem, divide by the specific elements atomic mass from the periodic table (if there are 2 oxygen molecules, account for that) ● The element with the smallest number is the limiting reactant, and is therefore how many moles of a compound that can be produced ● In most reactions not all of the product can be isolated ● To calculate percent yield, ○ Percent yield= actual yield/theoretical yield x 100 Chemistry Lecture and Textbook Notes Chapter 4 4.1 Vocabulary ● Solvent: compound that has the same physical state as the solution ● Solute: substance being dissolved ● Aqueous solution: water is the solvent ● Strong electrolyte: compound that separates into ions in water ● Weak electrolyte: only partially ionizes when dissolved in water ● Ionization: separation of molecular compound into cations and anions when dissolved in water Important to Know: ● Solubility Rules: + ○ Group 1A cations and NH are 4luble with no exceptions ○ Nitrates (NO )3re soluble without exception ○ Perchlorates (ClO )a4 soluble without exception ○ Acetates are soluble without exception (CH COO) + 2+ ○ Chlorides, bromides and Iodides (Cl, Br, I) are soluble, except with Ag , Hg 2 , Pb + ○ Sulfates (SO )4re soluble except Sr , Hg2+ 2, Pb , Ba 2+ ● Insolubility Rules: ○ Carbonates (CO ) except Group 1A cations and NH + 33 4+ ○ Phosphates (PO ) ex4pt Group 1A cations and NH 4 ○ Hydroxides (OH) except Group 1A cations, NH , Sr , and Ba 2+ 4 ● A precipitation reaction involves the formation of an insoluble product or products from the reaction table of soluble reactants. ○ Ex: mixing AgNO and LiCl (both soluble) form insoluble AgCl 3 ○ It helps to lay out all the reactants and possible products, labeling each as soluble or insoluble ● Complete Ionic Equation: shows all strong electrolytes as ions in solution ● Net Ionic Equation: only shows species in the solution that actually undergo chemical change ● Spectator Ions:do not participate in chemical reaction 4.2 Vocabulary ● Molarity: number of moles of solute in one liter of solution Important to know: ● Molarity = moles of solute/liters of solution ● Solutions of lower concentration can be prepared by diluting more concentrated solutions of known molarity ○ In dilution problem: ■ Moles of solute in dilute solution = moles of solute delivered from the concentrated solution ● molarity(conc) x volume(conc)= molarity(dil)x volume(dil) 4.3 Important to Know: ● Molarity is used to calculate moles from volume of solution analogous to using molar mass to calculate moles from mass of a solid 4.4 Vocabulary ● Equivalence point: the point (volume) in a titration where equivalent amounts of the two reactants have been added ● Indicator:compound that changes color at the end point Important to Know: ● In a titration the concentration and volume of a solution of known concentration is used to determine the concentration of an unknown solution ● The end point volume should be close to the equivalence point volume
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