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Chemistry Unit Notes Viborg Hurley Science 2020**Intro to Chemistry Textbook McGraw Hill **HWK questions via Pearson Unit 1: Matter and Energy w/ Mathematics 1.1: Matter and its Classifications Matter is anything that takes up space and has mass. Mass is the quantity of matter. (Different from weight)Energy is not made of matter; it is the ability to do work. Composition of Matter Pure SubstanceAlways have the same chemical composition, no matter the origin. Cannot be separated by physical means. (aluminum, silicon)Types Elements: substance that cannot be broken down into simpler substances even by a chemical reaction. (Natural or Synthetic…83 of 118 are natural) Depicted in a periodic table sorted by similar properties into groups and periods. Metals or nonmetals. Distinguished by luster, thermal conductivity, electrical conductivity. Elements are represented by element symbols or a shorthand version of the longer name. (Represents the common name or Latin name pg. 6 McGraw book)Compounds: pure substances composed of two or more elements combined in definite proportions. Properties can differ from component elements. (Ex. Salt)Represented with chemical formulas that describe the composition of a compound using the symbols for the elements and subscripts to indicate the number of each element. Mixtures Consists of two or more pure substances that vary in composition. Can be separated by chemical means using physical and chemical properties. Types Homogeneous Mixtures Uniform composition throughout. Often called a solution. Most solutions are compounds dissolved in water and are clear. Ex: Saltwater, tea, acidsHeterogeneous Mixtures Not uniform throughout…different samples have components present in different proportions. Representing Matter Chemistry focuses on the microscopic level and to create models representing these particles on the macroscopic level. Microscopic structures of matter are called atoms. Smallest unit of an element that has the chemical properties of that element. Molecules are combinations of elemental units or two+ atoms bound together. States of Matter A physical state is the form that matter can take. Solid State Fixed shape, fixed volume, no volume change under pressure, particles are fixed and in a regular, crystalline shape. Liquid StateShape varies, fixed volume, slight volume change under pressure, particles are randomly arranged and free to move about. Gaseous StateShape varies, volume varies, large volume change under pressure, particles widely separated and move independently. Highest amount of energy. Representing states in chemical formulas Typically shown as subscripts behind the compound: (s), (g), or (l). The symbol, (aq), represents an aqueous solution in which a substance is dissolved in water. 1.2: Physical and Chemical Changes and Properties of Matter Various Classifications of Properties Quantitative or Extensive Properties Properties that are based on numerical values and depend on the amount of matter in a sample. Qualitative or Intensive Properties Properties that are based on observations and the type of matter, not how much. Ex: color, shape, texture, shininess, physical state. Physical Properties Characteristic that we can observe or measure without changing the composition of a substance. Ex: odor, taste, hardness, mass, volume, density, magnetism, conductivity, and temperature. Quantitative Physical Properties Mass Amount of matter. Typically measured in grams, milligrams, kilograms. (Metric conversions and metric ladder) **Sketch on board. Outline prefixes. (Dimension Analysis PP and Conversion pg. 41-45 McGraw Hill) Use dimensional analysis to convert between mass units. Ex: There are 50.0 mg of sodium in a soda, how many grams are present? How many pounds? (1 lb = 453.6 g)PP 1.5 pg. 16Volume Amount of space a substance occupies. Measured in liters, milliliters. (Metric conversion) 1 mL = 1 cm3Volume is measured using graduated cylinders, beakers, flasks, or by using the equation for solids. V = length x width x heightUse dimensional analysis to convert between volume units. Ex: What is the volume of a 12.0 oz can of soda in milliliters? What is the volume in liter? (1 oz = 29.57 mL)PP 1.6 pg. 18Density Ratio of its mass to its volume. That being said density is an intensive property as it is specific to the type of substance. (Demo density column) Gases have low densities (low volumes, gas particles spread out), while metals have high densities (packed together). What can we infer about the density of solid state of water compared to its liquid state?Density’s equation: Density = mass / volume and is measured in g/mL, g/cm3, or some type of mass/volume unit combination. Sample ProblemsWhat is the volume of 100.0 g of copper? The density is 8.9 g/cm3?PP 1.7 a and b pg. 20Temperature Measure of how hot or cold something is based on the kinetic energy of the particles. Three scales are used: Fahrenheit (United States), Celsius, and Kelvin. Conversions Tk = Tc + 273.15Tf = (1.8 x Tc) + 32Tc = (Tf – 32) / 1.8Intensive properties relating to temperature Boiling point and melting point. Absolute zero: lowest possible temperature. 0 on the Kelvin Scale and -273.15 degrees on the Celsius scale. Sample Problems The melting point of copper is 1,083 degrees C. Above what temperature, in Kelvin and Fahrenheit, is copper a liquid? What is the physical state of copper at 1,000 degrees C?PP 1.9 a-cPhysical Changes Process that changes the physical properties of a substance without changing its chemical composition. Separating mixtures utilizes physical changes via:Filtration (particle size)Magnetism Distillation (utilizes boiling points and condensation points to separate out mixtures into its components.) **Cherry Coke Lab: PBL??Chemical Change Process in which one or more substances are converted into one or more new substances. Ex: Rust, Tarnish Chemical changes are depicted in the form of a chemical reaction where reactants form products. Clues are transfers of energy, change in color, production of a gas, or formation of a precipitate. All physical and chemical changes follow the law of conservation of mass. (Law vs Theory)Chemical Properties Defined by what it is composed of and what chemical changes it can undergo. Ex: Reactivity and flammability Not all reaction observations can indicate physical or chemical changes. (Bubbles) 1.3 Energy and Energy Changes Energy is the capacity to do work or transfer heat, while work is the acting of a force over a given distance. Most chemical reactions require continuous energy inputs or are spontaneous. Do they do work?Types of Energy Kinetic energy: energy of motionThermal energy increases as kinetic energy of particles increase. Electrical energy increases as kinetic energy of particles increase. Potential energy: energy possessed by any object. Chemical potential energy is stored in bonds. Thermal and electric energy can also be stored or put into motion. Ex: Trinitrotoluene (TNT)Energy conversions follow the Law of Conservation of Energy. 1.4: Scientific Inquiry and Fields of Chemistry Review Inquiry through Penny Cleaning Problem Based Learning. (Use pg. 30-31 as a guide ) Fields of Chemistry Organic Chemistry Study of all chemicals containing carbon. Inorganic Chemistry Study of all chemicals without carbon. Biochemistry Study of processes that take place in living organisms. Analytical Chemistry Study that focuses on the composition of matter. Physical Chemistry Study that deals with the mechanism, rate, and energy transfer that occurs when matter changes. Green Chemistry Develop chemical processes that prevent pollution and reduce the amount of natural resources used to manufacture products. Types of Chemical Research Pure chemistry: pursuit of chemical knowledge. Applied chemistry: application of chemistry to a practical goal or application to a problem. PP 57 pg. 28 Pearson Chemistry 1.5: Math Toolbox for Chemistry Scientific Notation Tool to represent very large or very small numbers in an easier format. Based on rules of 10 with a coefficient and an exponent (positive or negative). Example Problem1.14 pg. 35 McGraw Hill a-ePP 1.14 a-e Rules of Scientific Notation Multiplication Multiply the coefficients and add the exponents. (pg. 36)Division Divide the coefficients and subtract the exponents. (pg. 36)Addition and Subtraction Express both numbers with identical exponents. Ex pg. 36Example Problem 1.15 pg. 36 a-dPP 1.15 a-d Accuracy, Precision, and Significant Figures Measurements are either exact (counting) or not exact. Precision and Accuracy Precision is the extent of agreement between repeated measurements of a given value. Accuracy is the difference between the value of a measured number and its expected or correct value. (Diagram bullseye model) Significant Figures All the digits in a number of which we are absolutely certain, plus one additional digit that is uncertain. Ex: Say we read on the side of a graduated cylinder and the line seems like it is at 18. What are we certain on and what are we uncertain on?Rules on Determining the Number of Sig Figs Nonzero numbers are always significant. All numbers are significant when counted. A zero alone in front of a decimal point is not significant. 0.2806. A zero to the right of the decimal point but before the first nonzero digit is not significant. Ex: 0.002806. A zero between nonzero numbers is significant. Ex. 2,806, 0.02806A zero at the end of a number and to the right of the decimal is significant. Ex: 0.0028060, 2,806.0A zero at the end of a number and to the left of the decimal point is not significant if the value was not counted. Ex: 28,060 (normally would write in scientific notation to be clear about 4 sig figs)Example Problems 1.16 a-b, PP 1.16 a-b pg. 38Sig Fig Rules in Calculations Multiplication and Division The product or quotient must have the same number of significant figures as the least precise number in the problem. Example Problem 1.17 a-d pg. 38PP 1.17 a-dAddition and Subtraction The sum or difference can only be as precise as the least precise number used in the calculation. Round to the first uncertain digit. Example Problem 1.18 a-d pg. 39PP 1.18 a-dMultistep Problems Keep track of sig figs throughout the steps (pg. 39)Ex 1.19 a-d pg. 40PP 1.19 a-dRounding Numbers Example 1.20 a-e Unit 1 Assessments: Distillation Lab, Penny Lab, Exam Unit 2: Atoms, Ions, Periodic Table, and Electron Structure 2.1: Dalton’s Atomic Theory Initial belief was that matter was continuous and could be divided continuously. (Aristotle) John Dalton challenged that notion based on multiple lines of evidence:Antoine Lavoisier’s experimental results and the law of conservation of mass. Open vs closed container. Joseph Proust’s experimental results and the law of definite proportions. States that all samples of the same compound always contain the same proportions by mass of the composed elements. (Ex: Pure water has a mass ratio of 8 g of oxygen for every 1 g of hydrogen) Dalton’s Atomic Theory All matter is composed of small, indivisible particles, called atoms. (Since been adjusted)All atoms of a given element are identical both in mass and chemical properties. Atoms of different elements have different masses and chemical properties. (Since been adjusted)Atoms are not created or destroyed in chemical reactions. Atoms combine in simple, fixed, whole-number ratios to form compounds. 2.2: Structure of the Atom Subatomic Particles Electron Negatively charged subatomic particle, demonstrated by J.J. Thomson. (1897)Cathode Ray Tube Experiment A voltage is applied to the tube and as electricity runs it forms a ray and can be observed by making glass coated materials glow. Thomson noticed that rays bent towards a positively charged plate and deflected away from negatively charged plates. (Repeated test with multiple materials)Millikan Oil Drop Experiment Experiment to determine the charge to mass ratio of these electrons. Exposed droplets to radiation to give them an electrical charge. Millikan that determined the size of the electric field needed to suspend oil droplets in air. (Charge was found to be -1.6022 x 10-19 coulboms) Thomson used charge-mass ratios to calculate a mass of 9.91094 x 10-28 g) For some time, this led to the misunderstanding that atoms contained thousands of electrons because an electron was 1,836x less than one hydrogen atom. Proton Positively charged subatomic particle that is equal, but opposite in charge to the electron. Discovery due to the fact that many atoms were neutral in charge. Nucleus Prior to its discovery, the big question was how are protons and electrons arranged within an atom. J.J. Thomson was the first to suggest a model called the “plum pudding” model. Assumed protons and electrons were evenly distributed in an atom. Ernest Rutherford Designed an experiment to test Thomson’s model, called the gold foil experiment. Rutherford’s goal was to bombard a thin gold foil with alpha particles, because alpha particles are massive and positively charged. Based on Thomson’s model the hypothesis was that these particles should zip through the gold foil unaffected. (Why?)Results Some were deflected slightly and some even zipped backwards. (Equivalent to shooting a military artillery shell at tissue paper and having it bounce back.)Conclusions Suggested that most of the mass of an atom is concentrated in a positively charged core, called the nucleus. Electrons were dispersed outside of the nucleus. Led to the development of the nuclear model of the atom. Biggest problem was that there was still a missing portion of the mass of the atom. Hypothesized the existence of a neutron or an uncharged particle in the nucleus of the atom. Later was tested and confirmed by Rutherford’s colleague James Chadwick. Found the mass to be 1.679 x 1024 g. Isotopes, Atomic Number, and Mass Number Atomic Number Number of protons (Z) of the element. Identifies the element. If the charge is neutral, this would also indicate number of electrons. Isotopes and Mass Number Isotopes of elements are atoms that contain a specific number of neutrons. Identical in chemical properties but differ in physical properties such as melting and boiling points. Mass number (A): sum of the number of protons and the number of neutrons (N) in the nucleus. (**Not an actual mass) A = Z + N Isotope Symbols (pg. 67) Example problem and Practice problem 2.2 2.3: Ions What happens when the number of protons and electrons is not equal?Atoms becomes charged, or simply, becomes an ion. Many elements exist naturally as an ion. Cations (positive charge) or Anions (negative charge) Example / Practice Problems 2.5 pg. 70-71Writing Isotope Symbols for Ions Pg. 71 Example 2.6 2.4: Atomic Mass Atoms were able to be measured using a technique called mass spectrometry. (Video) Due to masses of atoms being incredibly small, scientists devised a unit to make it more manageable. Atomic Mass Unit (amu) 1 amu = 1/12th the mass of 1 C-12 atom (1.66 x 10-24 g) In order to find the relative or average atomic mass of an element we must take the atomic mass of each isotope multiplied by its abundance in decimal form. The sum of all the isotopes equals an atom’s average atomic mass. (Value on most periodic tables) Ex: Ag-107 and Ag-109 with 51.82% and 48.18% relative abundance respectively. PP 2.7 and 2.8 on pg. 73 and 742.5: The Periodic Table Dmitri Mendeleev Developed and published the basic arrangement of the periodic table. (Show image)Ordered in order of increasing atomic mass (63 known at the time). Grouped together elements with similar properties into columns and rows so that they varied periodically. (Get it) He left spaces in his table based on predictions of elements to be discovered later. (Gallium, Scandium, and Germanium) While his original table was slightly incorrect, the basics led to the current periodic table in order of atomic number. Classification of Elements Groups or Families Represented with Roman numerals (I through VIII) with an A or B, or with a 1-18. Periods Horizontal rows with elements that vary in regular fashions. (1-7) Types of elements Metals, nonmetals, metalloids. Elements in groups with an A or 1-2, 13-18 are known as main-group or representative elements. Elements in groups with a B or 3-12 are transition metals. Inner transition metals Group of elements placed below the periodic table. Group Names (all elements in a group share similar properties)IA: Alkali metal (reactive, especially with water)IIA: Alkali earth metal (less reactive, more so than other metals)VIIA: Halogens (occur as diatomic molecules, form salt bonds)VIIIA: Noble gases (inert, do not react) Pg. 78 PP 2.9 Ions and the Periodic Table Position of an element on the PT helps predict the charge of its ion. Many atoms of the main-group elements gain or lose electrons to form ions with the same electron count as the nearest noble gas. Nonmetals typically gain electrons and metals lose electrons. Group IA elements for +1 charge, Group IIA form +2 charge, Group IIIA form +3 charge, Group IV varies, Group V forms -3 charge, and the pattern continues. Nobel gases are inert and are neutral. Pg. 79 and 80 PP 2.10 7.1: Electromagnetic Radiation and Energy Electromagnetic radiation or radiant energy Travel through space as oscillating waves and are composed of electric and magnetic fields. All move at the speed of light. (3.0 x 108 m/s)Spectrum consists of visible light, X-rays, UV light, infrared, microwaves, and radio waves. PropertiesWavelength (symbol is lambda) is the distance between two points on a wave. Shorter = more radiation. Frequency is a measure of the number of wave cycles moving through a point in space per 1 second. Measured in Hertz. (Wavelength increases, frequency decreases) Eyes only detect visible light (ROY G BIV)Atoms absorb light and reflect the color of light we perceive it to be. Also can alter the true color of objects. (Ex: red car and sodium vapor lights that emit only yellow light)Photons Unit of light energy that depends on the wavelength or frequency. (Increases with increasing frequency, decreases with increasing wavelength)Violet has the highest frequency / energy of visible light. PP 7.1 pg 263 and 264Calculating wavelength, frequency, and photon energy Ephoton = (hc / γ) C = speed of light (3.00 x 108 m/s)H = Planck’s constant = 6.626 x 10-34 J s Lamda = wavelength in meters PP 7.2 pg. 265Atomic Spectra Line spectrum Passing light from an element with electric current through a slit and a prism we separate the colored light into distinct colored lines representing single wavelengths of light. “Fingerprint” of an element. Must be heated or given a charge to produce energy. (PHET)7.2: Bohr Models Issues of the Rutherford model of the atomDidn’t follow the basic laws of physics known at the time. How could negative and positive charges remain separate?What prevents electrons from tumbling towards the nucleus?Niels Bohr (PHET Lab)Came to work with Rutherford a year after Rutherford published his model of the atom. Tasked with explaining the stability of the electrons in his arrangement. Proposed the planetary modelElectrons orbit the nucleusBased off the work of Max Planch who believed energy produced by atoms is quantized (can only have certain values) Bohr believed the energies of electrons in an atom are quantized. Proposed several fixed orbits, each with a specific radius and energyEach labeled with n. Electrons in orbits close to the nucleus have low energy. Electrons can absorb energy and move to a higher energy orbit. Electrons can release energy and drop to lower orbits. Change in energy = Final energy – Initial energy Negative = energy is releasedEphoton = Absolute value (Change in energy) Example 7.3 Using Hydrogen atom and line spectrum PP 7.3Limitations to the Bohr Model Explained only the hydrogen atom and did not work for atoms with more than one electron. 7.3: Modern Model of the Atom Edwin Schrodinger Developed a mathematical model that worked for all atoms. Similar to Bohr’s model Energies of electrons were quantized. Variation from Bohr’s modelDescribes electrons occupying orbitals, not orbits. Orbitals are 3D regions in space where electrons are likely to be found, not circular pathways. (Based on probabilities) Principle Energy Levels Similar to Bohr’s energy levels, Orbitals of similar size all reside in the same energy level. Smaller the energy level = lower energy = smaller the orbital. Orbital Shapes S, p, d, and f. (Diagram shapes on bored pg. 271) The s has one orientation and can hold two electrons, p has three orientations and can hold six electrons, d has five orientations and can hold 10 electrons, while the f orbital has seven orientations and can hold 14 electrons. Orbitals and Energy Levels Energy level 1 Consists of a single s orbital. Energy level 2 Consists of s and p orbitals. (sublevels are one type of orbital at a specific energy level) 2s contains 1 sublevel, while 2 p contains 3 sublevels. Energy level 3 Consists of s, p, and d orbitals. Contain 1, 3, and 5 sublevels respectively. Energy level 4Consists of s, p, d, and f orbitals with 1, 3, 5, and 7 sublevels respectively. Orbital Diagrams and Principles Aufbau Principle Electrons fill orbitals starting with the lowest energy orbitals first. Ex: Hydrogen’s electron is in the 1s orbital. Pauli Exclusion Principle A max of two electrons can occupy each orbital, and they must have opposite spins. (Represented with arrows pointed in opposite directions) Electrons spin on their axis. (Think of the Earth) Hund’s Rule Electrons are distributed into orbitals of identical energy (same sublevel) in such a way as to give the maximum number of unpaired electrons. Example Problems 7.4 pg. 275 a-c with Practice Problems Electron Configurations Shows the distribution of electrons among sublevels. Shorthand version of an orbital diagram. Number of the principal energy level written first followed by the symbol for the sublevel. A superscript is added to the sublevel to indicate the number of electrons in that level. Ex: Carbon’s electron configuration: 1s22s22p2 (Superscripts add to equal atomic number)Example and Practice Problems a-c 7.5 pg. 276 7.4: Periodicity of Electron Configurations What do you notice for the electron configurations of Li, Na, and K?Periodic Table Blocks Groups IA and IIA (including He) are in the s block. (The highest energy level electrons reside in an s orbital.)Groups IIIA through VIIIA are in the p block. The transition metals reside in the d block, while the inner transition metals reside in the f block. Energy Level Trends The principal energy level written matches the period number for elements in the s and p blocks. The principal energy level written for the d block is the period number -1 and for the f block, it is the period number -2. Ex: Electron Configuration for Mn. Example and Practice Problems 7.6 a-c on pg. 280-281Noble Gas Configurations (shorthand versions)Utilize the symbol for the previous noble gas in brackets followed by the remaining electron configuration. Example: Ca and Ar7.5: Valence Electrons for the Main-Group Elements Valence electrons are electrons located in the highest principal energy level. (Core electrons are all electrons in lower energy levels.)Electron configurations indicate the number of valence electrons:Main group elements have valence electrons that are equal to the number of electrons in the highest energy s and p sublevels. Valence electrons = group number. Ex: Group VIIIA has 8 valence electrons. Example 7.7 a-c pg. 2847.6: Electron Configurations for Ions How do we write the electron configuration for an ion?For cations we remove valence electrons from the configuration. Ex: Calcium vs Ca2+ ion. For anions we add valence electrons to the configuration. Ex: Cl vs Cl- ion. Notice that the calcium ion, the chlorine ion, and argon all have the same configuration. They are isoelectronic. Ex 7.8 pg. 285-286. 7.7 Periodic Properties of Atoms Chemical Reactivity and Electron Configuration Valence electrons are the only ones that participate in chemical reactions because they are furthest from the nucleus. Alkali and Alkali Earth metals are the most reactive metals. (Group IA more than IIA) Halogens are the most reactive nonmetals. Most of these elements exist as ions in nature, which explains their reactivity. (Write configurations for each.) Reactivity trends down groups (Video of all alkali metals in Group 1 with water…students draw observations.)Reactivity increases as we go down the groups. Due to ionization energy. Ionization Energy The more easily metal atoms give up their valence electrons, the more reactive they are. Ionization energy is the measure of the energy required to remove a valence electron from a gaseous atom to form a gaseous ion. Atoms with low IE do not bind their valence electrons very tightly and therefore are very reactive. IE increases down a group due to an increase in electron shielding. (Decrease in attraction between a valence electron and the nucleus of any atom.)IE increases from left to right within a period. (Increase in protons makes it harder to remove valence electrons which requires more energy.) Example 7.9 pg. 290 Atomic Size and Ion Size Atomic Radius Radius increases down a group due to larger orbitals. (Discuss group IA)Decreases across a period due to the increased positive charge pulling valence electrons closer to the nucleus. Example 7.10 on pg. 293Ion Size Cations tend to be smaller than their neutral atoms due to loss of electrons. Anions are larger than neutral atoms due to increase in electrons. No single trend for the periodic tableTrend for isoelectric series: As the number of protons increases, the ion size decreases. Ex: S2-, Cl-, K+, Ca2+, and Sc3+ (Arranged from largest to smallest) Example 7.11 pg. 294 Unit 2 Assessment: Flame Lab, Atomic Mass Lab, Exam with Battleship Review Unit 3: Chemical Compounds and Chemical Bonding 3.1: Ionic and Molecular Compounds Electrolytes A substance that separates into ions when dissolved in water. (Some solutions create enough ions to conduct electricity.)Ex: SaltTypes Strong electrolyte: Dissociates extensively in water and conducts electricity well. (Salt, HCl)Weak electrolyte: Dissociates only partially and does not conduct electricity well, but more so than a nonelectrolyte. (Acetic acid)Nonelectrolytes Substances that have ions that stay intact and do not conduct electricity. Ex: Pure water, sugar solution Types of Compounds (PBL based Lab)Ionic CompoundConsists of oppositely charged cations and anions in proportions that give a net zero charge. Usually consist of a metal and a nonmetal such as salt. Properties Crystalline solid, hard, brittle solid, very high melting and boiling points, high density, strong electrolytes, good electrical conductivity. Molecular Compound Not composed of ions and typically contains two or more different nonmetals. Also referred to as covalent compounds. Examples include carbon dioxide and oxygen molecules. Properties Gas, solid, or liquid Soft solid, low melting and boiling points, low density, weak electrolyte, and poor conductors. Example 3.1 pg. 93 and 3.2 pg. 963.2: Monatomic and Polyatomic Ions Monatomic IonsIons of a single atom. For example, if we dissociate MgS we would get Mg2+ and S2- ions. Remember the charge is easily predictable for the main group elements based on their group number. Nomenclature for Ions (Intro by providing Table 3.2 on pg. 97 to see if students can get naming rules) Anions are named as the first part of the element name and –ide added as a suffix. Cations are just named as normal with the word ion added after. Example Problem 3.3 on pg. 97 Polyatomic Ions Ion that contains two or more atoms, usually of more than one element. Most common are anions that contain oxygen attached to some other element. (Oxoanion)Nomenclature for Polyatomic Ions (Same intro as for monatomic ions but use Table 3.4) If there are only two oxoanions of an element, the one with the greater number of oxygen atoms is named by combining the root word with –ate. The oxoanion with the lesser number of oxygen atoms uses the ending –ite. If they form more than two oxoanions, such as chlorine, they are named as before. However, the prefix per- is placed before the root of the element name and –ate is added to the end for the oxoanion with the most oxygen and the prefix hypo- is added with the ending –ite to the one with the least. Example Problems 3.4 on pg. 99 3.3: Formulas for Ionic Compounds Rules of Thumb for Writing Ionic Compound Formulas The sum of the positive charges must equal the sum of the negative charges for each ion. The bonded ions create the formula unit. Generally, the name of the element that is farther to the left or farther down on the periodic able is written first. In ionic bonds it typically is the metal. ExamplesSodium chloride Sodium sulfide Calcium carbonate (polyatomic ion example)Calcium nitrate (polyatomic ion example)Example Problems 3.6 pg. 103-104 3.4: Naming Ionic Compounds Rules Naming the cation followed by the anion ending in –ide. Ex: NaCl Naming ionic compounds with polyatomic ions is similar to the naming rules above with the exception of the –ide ending. Simply name the cation followed by the polyatomic ion name. Ex: Na2SO4Naming ionic compounds that contain metals capable of forming multiple ion chargesTypically occurs with transition metals. Name using the stock system. A roman numeral is written in parentheses behind the cation indicating the number of the charge followed by the anion name ending with –ide. Ex: Copper and Chloride (1+ and 2+)When writing the formula, the rules are straightforwardRoman numerals indicate the charge, so you need to figure out the charge of each ion and calculate the number of each ion needed to form a neutral charge. **Use CuCl2 as an example. When using the old naming system suffixes will indicate the charge number. When ending in the suffix –ic, it indicates the higher charge, while –ous indicates the lower charge. Ex: Copper and Oxygen examples. Example Problems 3.7 pg. 105, 3.8 on 107, 3.9 on 108, and 3.10 on 109 – 110. 3.5: Naming and Writing Formulas for Molecular Compounds (Introduce table 3.8 on pg. 110 followed by Example 3.11. See if students can come up with naming rules)Naming Rules Name elements in the order they appear in the formula. We use a numerical prefix added each element to indicate the number of each element present in the formula. We add the suffix –ide to the ending of the second element. Ignore the mono prefix in the first element. Writing Formulas Basically reverse the naming rules. Use prefixes to indicate the number of each element. Example 3.12 pg. 112. 3.6: Acids and Bases Acids are defined as a substance that when dissolved in water provides hydrogen ions. Usually contain a hydrogen that is removed when dissolved in water. Hydrogen ions end up binding with water to form H3O+ ions. Organic acids contain the combination of atoms –CO2H. (Carboxylic acids)Writing Formulas for Acids Place the hydrogen first and write (aq) after the formula to indicate that the compound is an acid when dissolved in water. Bases are defined as a substance that reacts with an acid in aqueous solution to form water. Most contain hydroxide (OH-) ions. When mixed with an acid they create water. Strong and Weak Acids and Bases Acids and bases are electrolytes. Strong indicates complete dissociation, while weak indicates little / partial dissociation. Rules for Naming Acids and Bases Binary acids are named with the prefix hydro followed by the stem of the name of the nonmetal with the suffix –ic. The word acid follows. Ex: HCl. Naming acids containing polyatomic ions: Hydro is not used. Remove the –ate ending from the name of the polyatomic ion and replace it with –ic with acid following. Ex: H2CO3Naming acids containing polyatomic ions with multiple ions: If the ion ends with –ate replace with –ic followed by acid. If the ion ends with –ite replace with –ous followed by acid. Ex: H2SO4 and H2SO3Example Problem 3.13 on pg. 115 8.1 Types of Bonds Chemical bonds are forces that hold atoms together in a molecule or compound. In each type of bond the goal is to form a more stable electron configuration for each participating atom. Ionic Bonding One or more electrons are transferred from a metal to a nonmetal, forming a positively charged metal cation and a negatively charged nonmetal anion. The ions are held together by electrostatic forces (between oppositely charged ions)Properties of ionic bonds include high melting points and hard, crystalline forms due to strength of the bond. Ex: LiF Covalent Bonding Electrons are shared between atoms in pairs. (Nonmetals and nonmetals)Properties include: gas, liquid, or soft solid due to low strength, lower melting and boiling points. Polar vs Nonpolar Covalent Bonds In some cases covalent bonds can have unequal sharing of electrons. Polarity of a bond is the degree of transfer of bonding electrons from one atom to another. Nonpolar covalent bonds are when electrons are equally shared and charge is evenly distributed. (Common in diatomic and triatomic molecules) Polar covalent bonds are when electrons are shared unequally. Electrons tend to be found closer to one of the atoms than the other resulting in partial charges. Electronegativity Idea proposed by Linus Pauling to describe how charges could be separated in a bond. Proposed a partial transfer of electrons that resulted in partial ionic characteristics within covalent bonds. Termed electronegativity as the ability of an atom to attract bonding electrons. If atoms are identical then electrons are shared equally and the bond is nonpolar covalent. Electronegativity Values Fluorine has the highest value of 4.0 while Francium as the lowest at 0.7. Increase from bottom to top within a group and from left to right across a period. (Noble gases do not have values as they don’t form bonds)The greater the different in electronegativity values, the more polar a bond is. Example problem 8.2 on pg. 3098.2: Ionic Bonding Lewis Symbols A symbol used to show the valence electrons present in an atom. Dots are placed singly on four sides and paired when necessary. Noble gases are surrounded by an octet. Example Problem 8.3 on pg. 310 Hydrogen and Helium are exceptions as they can only hold…at max….two electrons. Lewis symbols for ionic bonds require a multistep process:Show the Lewis dot for each individual ion present. Then show the two ions bonding to create an ionic bond. Ex: NaCl on pg. 311 followed by Example 8.4 pg. 311-312. 8.3: Covalent Bonding The Octet Rule Tendency of an atom to achieve an electron configuration having eight valence electrons. Basically, atoms tend to bond with atoms that help them achieve the octet rule. Ex: CF4 **Only share valence electrons. Lewis Structures for the Diatomic Elements Single Covalent Bonds Covalent bonds that consist of a pair of electrons shared by two atoms. **Greatest bond length. Lewis structures are essentially Lewis Dot Structures for bonds. Single bonds are represented with a single dash. Ex: H2Double and Triple Bonds Occurs when combinations of atoms do not have enough electrons to satisfy the octet rule. Share more than one pair of electrons. Double bonds are represented with two dashes and share two pairs of electrons. Ex: O2Triple bonds are represented with three dashes and share three pairs of electrons. Ex: N2 **Greatest bond strength. Example Problem 8.5 on pg. 316 Steps for Determining Lewis Structures Draw an atomic skeleton by placing the symbols of all elements in the compound in the correct locations. Element with the greater number of atoms usually surrounds the one with the lesser number of atoms or the atom able to form the most bonds goes in the center. Central atoms tends to be the one that is less electronegative and is present in the least quantity. Hydrogen generally is on the outside and the chemical formula can give clues about the arrangement. Sum the valence electrons from each atom to get the total number of valence electrons. Place two electrons, a single bond, between each pair of bonded atoms. If you have not placed all the valence electrons in the formula, add any remaining electrons as unshared electron pairs to satisfy the octet rule. If necessary, shift unshared electron pairs from nonbonded positions on atoms with completed octets to positions between atoms to make double or triple bonds. Example 8.6 on pg. 318-319 Lewis Structures for polyatomic ions Must take into account the charge when summing the total number of valence electrons. Utilize brackets to show that the charge is spread out over an entire ion. Polyatomic ions with the same number of valence electrons have similar Lewis structures. Resonance Resonance occurs when electron arrangement in molecules can be represented by two or more different Lewis structures. Ex: Ozone (O3)Drawing a resonance structure is quite simple. There are two methods:Drawing each structure and separating them with a double arrow. (pg. 321)Drawing one structure and representing the location of the bond that varies with a dashed line. (pg. 321)Ex: Pg. 321-322 Example 8.7 Exceptions to the Octet Rule Three categories: Odd-electron molecules Molecules with an odd sum of valence electrons. Typically lower in electronegativity, but highly reactive. Color is a good indicator of odd-electron molecules. Ex: NO Incomplete Octets Some atoms that participate in covalent bonding but do not have enough valence electrons to form octets. Most notable is Boron. BH3 and BF3. Both are very reactive. Expanded Valence Levels Compounds that have more than eight electrons around the central atom. Ex: SF6 and XeF48.4: Bonding in Carbon Compounds Importance of Carbon Part of more chemical compounds than any other element except Hydrogen. Backbone of all living things. Versatile in bonding forming single, double, and triple bonds. C-C bonds are very strong and stable. Hydrocarbons Compounds that contain C-H bonds. Classes Aliphatic Hydrocarbons (class in which the bonds are all localized single, double, or triple bonds.)Alkanes: Hydrocarbons that contain carbon-carbon single bonds. Alkenes: Hydrocarbons that contain carbon-carbon double bonds. Alkynes: Hydrocarbons that contain carbon-carbon triple bonds. **Many large alkanes have similar compositions but different arrangements. (Isomers)Aromatic Hydrocarbons (class in which bonds are arranged in six-atom ring structures with alternating single and double bonds)Benzene and compounds with one ring. Compounds with more than one ring. Functional Groups Other groups of atoms are substituted for one or more hydrogen atoms on the hydrocarbon framework. Classes (Chart pg. 325)Alcohols contain a hydroxyl group –OHEthers –O-Alehyde KetoneCarboxylic acid Ester Amine Example Problem 8.8 pg. 325-3268.5 Shapes of Molecules Shapes of molecules are immensely important to everyday life. Your body relies on certain molecular shapes for taste and smell. Valence-Shell Electron-Pair Repulsion Theory (VSEPR) Tendency of electron pairs to adjust orientation of their orbitals to maximize the distance between them. In order to predict the geometric shape we need to know the number of unshared electron pairs and single, double, or triple bonds surrounding the central atom. Utilize Lewis structures for this step. Shape is characterized by the bond angle between the central atom and the atoms bonded to it. Parent Structures LinearNumber of electron domains (Atoms or electron pairs) is two. Bond angle is 180 degrees. Trigonal planar Number of domains is three. Bond angle is 120 degrees. Tetrahedral Number of domains is four. Bond angle is 109.5 degrees. Example Problem 8.9 pg. 329Molecular Shapes (Derived from Parent Structures) **All bond angles are assumed to be the same as the parent structure. Linear parent shapes have linear molecular shapes. There is no variation. Derivatives of Trigonal Planar If there are no unshared pairs of electrons, the molecular shape is trigonal planar. If there is 1 unshared pair of electrons, it is bent. Derivatives of Tetrahedral IF there are no unshared pairs of electrons, the molecular shape is tetrahedral. If there is 1 unshared pair of electrons, it is trigonal pyramidal. If there are two unshared pairs of electrons, it is bent. Example 8.10 pg. 331Determining shapes for molecules with more than one central atomTo explain let’s examine the molecule: H2NCH2CO2H (pg. 332)Sketch the Lewis structure first. Identify central atoms: N, both C, and one O. Predictions (pg. 332) Have students determine shape. Can include multiple shapes. Assessment: Ionic vs Covalent lab, Assessment Unit 4: Chemical Composition 4.1: Percent Composition (Chewing Gum Lab) Percent Composition by mass Expression of the portion of the total mass contributed by each element. %E = (mass of E / mass of sample) x 100 Consistent with law of definite proportions, so the percent composition will be constant no matter the sample size. Example problems 4.1 and 4.2 on pg. 131-1324.2: Mole Quantities (Math Toolbox 4.1 Sample Problems and HWK Practice) Moles and Particles A mole (similar to the concept of a dozen) contains 6.022 x 1023 atoms, molecules, ions, or formula units (compounds). The number is referred to Avogadro’s number. We can use this conversion to convert between moles and number of formula units or atoms. Example Problem 4.3 on pg. 135 Molar Mass Term used to describe the mass of 1 mol of a substance. Labeled in grams per mole. Same value as the relative atomic mass so we can use the mass given on the PT. Molar mass allows us to connect moles, mass, and particle units. Molar mass for any bond is found by adding the molar masses of each individual element in the bond. Example Problem 4.4 on pg. 136. Dimensional Analysis allows us to use molar mass to convert between the mass in grams and the number of moles. Example 4.5 pg. 137 and 4.6 on pg. 138-139. We can also go a step further and calculate the number of molecules from the mass using dimensional analysis. Example 4.7 pg. 139-1404.3: Determining Empirical and Molecular Formulas Empirical and Molecular Formulas Empirical formulas express the simplest ratios of atoms in a compound using the simplest whole-number subscripts. The molecular formula expresses the actual number of atoms in a molecule. Typically the molecular formula is the same as the empirical formula or a multiple of it. Examples: H2O and H2O2Example Problem 4.8 on pg. 142 Determining Empirical Formulas We can calculate the empirical formula using data from its chemical composition. If we know the moles of each element in a compound we simply divide each by the smallest number of moles. Ex: Chalcopyrite consists of 0.0200 mol Cu, 0.0201 mol Fe, and 0.0399 mol S. Divide by the smallest number of moles and round to the nearest whole number. Answer: CuFeS2Determining Empirical Formulas via Percent Composition Use the percent mass to calculate a starting mass of each element. Since the percent mass is constant for a compound we can assume 100 g of each substance (45% would mean 45 g out of 100 g). We can then use the molar mass of each element and the starting mass to calculate the moles and finish using the steps we used earlier. Example 4.9 pg. 144-145 Empirical Formulas for Compounds with More Than Two Elements Process is the exact same, there are just more conversions to be completed. Example Problem 4.10 on pg. 145-146 Empirical Formulas with Fractional Mole Ratios In most cases the mole ratios we calculate come out to whole numbers or very close (1.98, 1.99). Some formulas end up with fractional values such as 1.25, 1.5, etc. In a fractional mole ratio you must multiply all subscripts by a number that allows all subscripts to be whole numbers. Example Problem 4.11 pg. 146-147Molecular Formulas from Empirical Formulas In order to determine if the formulas are the same or a multiple we need to compare the molar mass. If the experimental molar mass is the same as the calculated one, the molecular formula is the same as the empirical formula. If the experimental molar mass is greater than the calculated empirical formulas, then the molecular formula is some multiple based on the ratios of the masses. Example: Propene’s empirical formula is CH2 and a molar mass is 42.12 g/mol. However, the molar mass from the empirical formula is only 14.04 g/mol. The ratio of the two masses indicates a multiple of 3. The molecular formula would be C3H6. Example Problem 4.12 on pg. 148 Determining Percent Composition without masses We can concert a chemical formula using the number of moles and the molar mass. Example Problem 4.13 pg. 149 4.4: Chemical Composition of Solutions Chemical reactions between solids are slow, so we carry out reactions by dissolving the compounds in a solution. Solutions are mixtures that are homogeneous and consist of two parts:Solute: substance that is dissolvedSolvent: substance doing the dissolvingConcentration Relative amounts of solute and solvent in it. Dilute (small amount of solute) or Concentrated (large amount of solute)Percent by Mass %mass = (mass of solute / mass of solution) x 100 Molarity Number of moles of solute dissolved in 1 L of solution. Represented by an M. Molarity = moles of solute / liters of solution (not solvent)Example 4.15 pg. 153 Moles from Volume and Molarity Example 4.16 pg. 154 Mass from Volume and Molarity Molarity to moles of solute using volume, then mole of solute to grams using molar mass. Example 4.17 pg. 155Volume of solution Required to make a Solution Grams of solute to moles of solute using molar mass, followed by moles of solute to volume of solution using molarity. Example Problem 4.18 pg. 155-156Dilution Dilution is the process of adding more solvent to a solution thus increasing the number of solvent particles and increasing the volume. Stock solution is the original solution. Equation to calculate between stock and dilution solutions. MdilVdil = MstockVstockExample problem 4.19 pg. 158 Assessment: Percent Comp Lab, Dilution and Molarity PHET Labs,Unit 5: Chemical Reactions and Equations 5.1: What is a chemical reaction?Conversion of one substance or set of substances into another. Reactants forms into products. All reactions follow the Law of Conservation of Mass. Atoms are not destroyed or created, just rearranged. Identifying reactants and products: Example 5.1 pg. 174 5.2: How do we know a reaction occurs? (PBL lesson?)Change in color, production of light, formation of a solid (precipitate, smoke, metal coating), formation of a gas, or absorption / release of heat. (Demo each) 5.3: Writing Chemical Equations Chemical equations are symbolic representations of a chemical reaction. Reactants yield products. The physical state (s, l, g, or aq) is written as a subscript by the atomic symbols of the reactants and products. Any special conditions needed for a reaction to occur (heat, electricity) are written above the arrow. Equation must be balanced meaning that the number of atoms of each element is the same for the products side and the reactants side. Coefficients are added to balance chemical equations. **Cannot change subscripts! To balance start by writing the skeleton equation. Example 5.2 on pg. 179-180Example 5.4 on pg. 181-182 (Dealing with ionic compounds)Balancing Chemical Equations with Fractional Coefficients Example Problem 5.3 on pg. 180-1815.4: Predicting Chemical Reactions Periodicity can be used to predict the products of a chemical reaction. Principle of periodicity suggests that alkali metals will react with water to form some sort of metal hydroxide and hydrogen gas. Ex: If we place lithium metal in water, what would the products be?Classes of Chemical Reactions Decomposition: AB yields A + B Combination: A + B yields AB Single-displacement: A + BC yields AB + C Double-displacement: AB + CD yields AC + BD Example 5.5 on pg. 185 Decomposition Reactions Breaks down into the elements of which it is composed or into simpler compounds. Predictable results, usually stable, small molecules form. (especially gases) Example 5.6 on pg. 187-188 Combination Reactions Occurs when two substances react to produce a single compound. When a metal element reacts with a nonmetal element, we would get an ionic compound. Again, can be predictable based on the ions they form. When both reactants are nonmetal we would produce a molecular compound. Example 5.7 on pg. 189-190 Single Displacement Reactions A free element displaces another element from a compound to produce a different compound and a different free element. The displaced element is often a metalPredicting which metals react with water or acids relies on the activity series. A more active element displaces a less active element from its compounds. Activity series pg. 191. Example 5.8 on pg. 191 Predicting which metal displaces another metal. The more active metal displaces the less active metal from its ionic compound, but a less active metal cannot displace a more active metal. (Use the activity series) Example 5.9 on pg. 193 Double Displacement Reactions Two compounds exchange ions or elements to form new compounds. Precipitation reaction: one product compound separates from the reaction mixture because it is insoluble and forms a solid. (Demo) Prediction of whether such a compound will form follows a set of rules:Most compounds of alkali metals and ammonium ions are soluble. All nitrates and acetates are soluble. Most sulfates are soluble. (Exceptions include sulfates with Ba, Sr, Pb, Ca, Hg, and Ag) Most chlorides, bromides, and iodides are soluble. (Exceptions include Ag, Hg, Pb)Silver compounds are insoluble unless paired with a nitrate or chlorate. Oxides and hydroxides are insoluble unless paired with an alkali metal. Sulfides are insoluble unless paired with a soluble group above. Chromates are insoluble unless paired with a group above. Carbonates, phosphates, sulfites, and silicates are insoluble unless paired with alkali metal or ammonium ions. Ex: What results when we mix BaSO4 and NaCl? Example Problem 5.10 on pg. 195-196Gas formation reaction: One separates the reaction as an insoluble gas. Gases that commonly form are H2S, CO2, and SO2 gases. Example 5.11 on pg. 197-198Acid-Base neutralization reaction: Results in the formation of water. A neutralization reaction occurs when an acid and a base react to form an ionic compound and water. Fairly simple to predict. Example 5.12 on pg. 198 Combustion Reactions Any reaction that involves oxygen molecules as a reactant and that rapidly produces heat and flame. Combustion reactions can vary as both metallic and nonmetallic elements can undergo combustion reactions but also be combination reactions. (Ex: S + O2) Hydrocarbons are the most common compound involved in combustion reactions and produce CO2 and water. Example 5.13 on pg. 200Oxidation-Reduction Reactions A reaction in which electrons are transferred. Two separate processes that occur simultaneously:Oxidation: process of losing one or more electrons. Reduction: process of gaining one or more electrons. Example: Zn(s) + CuCl2(aq) yields ZnCl2(aq) + Cu(s) First step is to analyze the charges of each element/ion involved. (pg. 573) Identify the oxidation step and the reduction step. Write and balance the equations for each step separately. (pg. 574) You will have to add electrons to the reactant or product side depending on the step you are writing. Identify the following:Reducing agent: the agent causing the reduction of an element in another reactant because it provides the electrons for reduction. (Zn in the example)Oxidizing agent: the reactant that contains the element that is reduced by accepting the electrons. (CuCl2 in the example) Oxidation Numbers (14.2) Oxidation numbers are a charge assigned to the atoms in any compound. For ionic compounds it is the charge, but for molecular compounds we use a set of rules. Rules on pg. 577 (Handout) An oxidation-reduction reaction occurs if one or more elements changes its oxidation number. If it increases, electrons have been lost. If it decreases, electrons have been gained. Total positive oxidation numbers + total negative oxidation numbers = net chargeEx: C + O2 yields CO2 In this case, based on our rules we know the oxidation numbers for C, O2, and the O atoms involved in CO2. What is our oxidation number for the C in CO2? In addition, we should be able to ID the oxidation and reduction steps as well as the respective agents. (pg. 580) Example Problem 14.2 on pg. 578-579Example 14.3 on pg. 580-581 Balancing Oxidation-Reduction Equations Example 14.6 on pg. 591-592 5.5: Representing Reactions in Aqueous Solutions To be accurate in representing reactions that occur in solution, we should represent soluble ionic compounds by the formulas of their component ions. Molecular equation: Pb(NO3)2(aq) + K2CrO4 (aq) yields PbCrO4(s) + 2KNO3(aq) Ionic Equations break up the molecular equation into its soluble ionic substances. Spectator ions are the ions that occur on both sides of the ionic equation. Net ionic equations include only those substances that are involved in the reaction. In the case of our example it would be: Pb2+(aq) + CrO42- yields PbCrO4(s)Example 5.14 on pg. 201-202Assessment: Chemical reaction clues lab, Chemical reaction types lab, Net ionic practice, Assessment Unit 6: Quantities in Chemical Reactions 6.1: The Meaning of a Balanced Equation Coefficients only represent relative numbers of reactants and products. So if 1 molecule of compound A reacts with 2 molecules of compound B to yield 3 molecules of compound C and 4 molecules of compound D. So if I have 2 molecules of compound A, how much of the others would I need / produce? We can also use the coefficient to determine the relative number of mols. 1.000 mols. Because 1 mole = 6.022 x 1023 molecules we can estimate actual number of molecules as well. Stoichiometry Process of determining the amounts of substances in a chemical reaction. 6.2: Mole-Mole Conversions A convenient way to relate the moles of a reactant or a product to the other reactants or products using mole ratios. Mole ratios are obtained by using the coefficients of the balanced equation. Example: C3H8 + 5O2 yields 3CO2 + 4H2O. The mole ratio of C3H8 to the other reactants/products would be 1/5, 1/3, 1/4 respectively. A mole to mole conversion uses the following conversion Moles of known substance to moles of unknown substance using the mole ratio as a conversion factor. Example 6.1 on pg. 218-2196.3: Mass-Mass Conversion Typically mass is used to measure reactants and products because we do not have a device to measure moles. In this case we rely on two conversion factors: molar mass and mole ratios. Grams of substance A is converted to moles using molar mass of A. The moles of A is converted to moles of B using mole ratios from the balanced equation. Moles of B is converted back to mass using B’s molar mass. You can then use the Law of Conservation of Mass to calculate the mass of a substance that is unknown as Mass of reactants must equal Mass of Products. Ex: 2Na + Cl2 yields 2NaCl We have 9.20 g of sodium. We want to know how many grams of chlorine gas should react with this amount of sodium in this reaction.Ex: 6.2 on pg. 222-2236.4: Limiting Reactants **Smore Lab Introduce this concept with HCl and magnesium reaction!! Ask students: Which is in excess at the end of the reaction? Which component limits the reaction?Limiting reactants are the reactants that react completely. The limit the amount of the other reactant that can react and limits the product that can form. In excess reactants are the reactants left over. Example 6.3 on pg. 224 and Example 6.4 on pg. 225-226. Limiting Reactants on the Molecular Level 2H2 + O2 yields 2H2O *In this reaction the reactants are the parts that combine in a ratio to make the products. For this to occur the H2 and O2 must be present in a molecular 2:1 ratio. If not, then one reactant would become the limiting reactant. In the reaction above, we mix 8 molecules of H2 with 5 molecules of O2. How many molecules of H2O would form (keep in mind the 2:1 ratio) and which reactant is the limiting reactant? (pg. 226) Can utilize a chart similar to that on pg. 227 Rules are simple:If calculated B = actual B there is no limiting reactant. If calculated B is greater than actual B, B is the limiting reactant because it will react completely. If calculated B is less than actual B, A is the limiting reactant. Example 6.5 on pg. 228-229 **Count particles! Limiting Reactants on the Mole Scale Similar process that utilizes mole scales instead of molecular scales. Mole and molecular ratios are the same essentially. Example: 2Na + Cl2 yields 2NaCl **Suppose we mix 0.50 mol Na with 0.20 mol Cl2. What is the limiting reactant? Use the chart method on pg. 230! Example 6.6 on pg. 230-231 Example 6.7 on pg. 232 **Limiting Reactants when given the mass of reactants. Simply use the molar mass to convert masses into moles to complete the limiting reactant process and then convert back to grams once you have your answer. 6.5: Percent Yield (Lab)Describes how much of a product is actually formed in comparison to how much should have been formed. Theoretical yields are the max amount of product that can be obtained from the given amounts of reactants. Actual yields are the amounts of products measured in a lab. Percent yield = actual yield / theoretical yield x 100 Example 6.8 on pg. 234 6.6: Energy Changes Law of Conservation of Energy Energy can be converted or transferred, but it cannot be created or destroyed. Ex: Burning fuel in your car turns chemical energy into mechanical and heat energy. Most energy is lost to the environment as heat. Efficiency: measure of the amount of useful work that is achieved from an energy conversion, but it is a scientific law that no conversion is totally efficient. From power plant to desk lamp the overall efficiency is 1.6% (Diagram pg. 236)Energy Changes that Accompany Chemical Reactions Energy is often stored as potential energy in chemical bonds. *Bond Dissociation Energy Activity! (Lab)A reaction that releases energy is called an exothermic reaction. The reaction profile (diagram it, pg. 237) shows that the reactants contain more energy than products. A reaction that absorbs energy is call an endothermic reaction. The reaction profile shows the products have more energy than the reactants. Example 6.9 on pg. 237Quantities of Heat Heat is energy that is transferred between two objects due to a difference in temperatures. Easiest to measure in a chemical reaction. Always transfers from hotter to colder objects until it reaches thermal equilibrium. Units of Energy Chemists measure energy in joules (J) or calories (cal). 4.184 J = 1 cal. Nutritionists use Calories (C). 1 C = 1,000 J (kilocalorie or kilojoule)Example 6.10 on pg. 238-239Specific Heat and Calorimetry (Lab) Substances have a unique property called specific heat. Specific heat is the amount of heat needed for 1 g of a substance to raise its temperature by one degree Celsius. Units: joules per gram per degree Celsius or calories per gram per degree Celsius. (Specific Heat chart in student books) Amount of heat transferred to or from a substance is related to the mass, specific heat, and the change in temperature. Example 6.11 on pg. 240 Calculating heat using specific heat Q (heat) = mass x C (specific heat) x temperature change (Tf-Ti) If q is positive, heat is absorbed and it gets hotter, but when q is negative heat is released and substance cools. Example 6.12 on pg. 241 Analyzing the energy of a systemWhat happens when we want to determine the heat change for an object for which we do not know the specific heat or if the initial temperature of the object is difficult to measure?Typically an insulated container called a calorimeter is used to prevent heat loss. The measure of heat transfer is called calorimetry. A good example would be heating a pipe and then cooling it. (similar to what welders do)If we use a calorimeter we assume no heat is lost so qsystem + qsurroundings = 0 based on the Law of Conservation of Energy. In the example above, the pipe is our system because that is what we are interested in, while the water (which is more easily measurable) is the surroundings. Example 6.13 on pg. 2436.7: Heat Changes in Chemical Reactions Bomb Calorimeters are similar to calorimeters for objects, but are used primarily for chemical reactions and to determine the energy content of food. (Zoos can use these)A specific amount of reactant is placed in a compartment with an ignition wire. Electricity triggers the ignition wire to start the combustion reaction. The water bath is kept in a separate compartment, but the temperature change of the water is measured. We assume qreaction + qwater = 0. Heat changes are reported in kilojoules per gram for chemical reactions. We can also report in kilojoules per mole. The conversion from moles of a substance to the heat change is simple and uses the heat of reaction value as the conversion. (Example pg. 245) Example 6.14 on pg. 245Energy released per mole generally increases with the number of carbon atoms in the molecular formula. (pg. 246) Hydrogen provides the most fuel per gram. (Hydrogen fuel) In relation to food, fats have the most energy per gram compared to proteins and carbs. Assessments: Specific heat lab, calorimetry lab, exothermic / endothermic lab, percent yield lab, limiting reagent lab / activity, Exam Unit 7: Reaction Rates and Equilibrium 12.1: Reaction Rates A measure of how fast or slow a reaction will occur. (Some are almost instant, while others can take millions of years.) Conditions that affect the rate of reaction includes: temperature, reactant concentration, surface area, and the presence of a catalyst / inhibitor. Higher temps, Greater concentration, greater surface area, and the presence of a catalyst all result in faster reactions. Rate Laws An expression for the rate of a reaction in terms of the concentration of the reactants. Rate = k (specific rate constant) x [A] *Large k value indicates products form quickly. Orders of Reactions Just means the power to which the concentration of a reactant must be raised to match the experimental data on concentration and rate. First-order reactions are reactions where the rate is directly proportional to the concentration of only one reactant. Ex: A yields B…if the concentration of A is cut in half, so will the reaction rate. Higher order reactions are reactions that rely on two substances to react to make products. Rate = k[A]x[B]y In order to calculation the overall order of the reaction (recall in first order, the overall order is 1 due to only one reactant) we must calculate the sum of the exponents. If, for example, both exponents were 1, the overall order of the reaction would be 2. Example 18.1 on pg. 606 in Pearson book 12.2: Collision Theory Collision theory states that in order for a reaction to occur, reactant molecules must collide in the proper orientation and with sufficient energy. Usually a small fraction because molecules have to hit in the exact orientation to allow for atoms to be transferred to create new bonds. Molecules also have to have a high enough kinetic energy to allow reactions to occur. (Think darts)Energy and Collision Theory Energy barriers must be overcome by the reactants before they can change to products. Minimum amount of energy needed to reach this is called activation energy (Ea). Once the energy requirement is met, the reactant molecules that collide with the proper orientation are called the activated complex. Short lived, unstable, high-energy chemical species. Example 12.1 pg. 490-491Every reaction has its own specific energy diagram and activation energy value. The larger the activation energy, the slower the reaction. 12.3: Conditions that Affect Reaction Rates (Lab…possibly PBL)ConcentrationAn increase in concentration increases the rate of the reaction due to the increased number of reactants per unit volume. Leads to a greater number of collisions per unit of time. (Fraction of effective collisions remains the same.)Surface Area Increasing surface area also increases rate of reaction. It also leads to a greater number of collisions per unit of time by increasing the number of atoms exposed for collisions. (Think ground sugar vs sugar cube)Temperature As temperature increases, so does the average kinetic energy of a substance. Causes an increase in reaction rates in two ways:Increases the collision rate (molecules move faster and collide more frequently)Increases the fraction of collisions that are effective. (KE increases so more of the reactant molecules can attain the required activation energy)Reaction rates double approximately every 10 degrees Celsius. Example Problems 12.2 and 12.3 on pg. 493-494Catalysts Substance that alters the pathway in which a reaction occurs without itself being consumed in the reaction. New reaction pathway is a lower-energy pathway with a lower activation energy. Increases the reaction rate, because a lower amount of energy is needed for the reaction to occur. Examples (in a reaction, catalysts are written above the arrow as they take no part in the reaction) Enzymes are molecules that catalyze specific reactions within living organisms. (many life-sustaining chemical processes would be so slow that they would not occur at our body temperatures)Each enzyme has a unique depression or hole called an active site that interacts with specific reactant molecules called substrates. (Lock and key method)Catalysts remain unchanged Example: Chlorine acts as a catalyst in the destruction of the ozone layer. Step 1: O3 + Cl yields ClO + O2Step 2: ClO + O3 yields Cl + 2O2A couple things to note here:Notice that Cl is used in step 1 and regenerated in step 2, which indicates that Cl is the catalyst. Notice that ClO is formed in step 1 and used in step 2, indicating it is an intermediate because it only forms temporarily. If we combine the two steps into one reaction our net reaction becomes 2O3 yields 3O2Example 12.5 on pg. 498 12.4: Chemical Equilibrium Many reactions occur until they reach a state where there is no change in the concentrations of reactants and products. Established when a single reaction occurs in which reactants are converted to products and the products back into reactants at an equal rate. Typically we represent these reactions with a double arrow. N2O4 (double arrow) 2NO2 **Rates are equal, concentrations are equal. Diagram figure 12.15 on pg. 500 12.5: Equilibrium Constant At equilibrium the amounts of reactants and products may be equal. Most often the amounts will favor the reactants or products. (Position of Equilibrium) Equilibrium Constant Expression Keq = [C]c[D]d / [A]a [B]b **[x] indicates molar concentrations while the superscripts indicate the coefficients. If the equilibrium constant value is greater than 1, we say equilibrium lies to the right and favors products. If less than one, it lies to the left and favors reactants. If it is equivalent to 1, it lies in the middle and reactants and products are similar. Example 12.6 pg. 503-504 Heterogeneous Equilibrium So far, all the reactions we have examined have been in the same state. (gas) An equilibrium in which reactants and products are in the same state is homogeneous. Different states are in heterogeneous equilibrium. Example: Br2(l) yields Br2(g) **Chemists omit pure liquids and solids from equilibrium constant expressions. Therefore, we would write Keq = [Br2(g)] / 1 Example: BaF2 (s) (double arrow) Ba2+(aq) + 2F-(aq) Keq = 1.7 x 10-6Keq = [Ba2+] [F-]2Example 12.8 on pg. 50912.6: Le Chatelier’s Principle This principle states that if a system at equilibrium is disrupted, it shifts to establish a new equilibrium. Changes that can disrupt include: Change in concentration, Change in the volume of a gas reaction container, and temperature change. Reactant or Product Concentration When the concentration of a reactant or product is increased, the equilibrium will shift to consume the added substance. If they are reduced, it will shift to produce more of the removed substance. Add reactant, shift rightAdd product shift leftRemove reactant, shift leftRemove product shift right Example 12.9on pg. 511-512Volume of the Reaction Container Before understanding how equilibrium shifts we must understand the relationship between volume and concentration. **Only occurs with reactants and products in gaseous state. A decrease in volume, increases the concentration due to increased pressure. The opposite occurs if you increase volume. Equilibrium Shifts Due to Volume Changes If the reactants are less than the products Increasing volume shifts right, decreasing volume shifts left. If the reactants are greater than the products Increasing volume shifts left, decreasing volume shifts right. If the reactants and products equal, there is no shift due to changing volume. Example 12.10 on pg. 514 Temperature (Demo video of N2O4 and 2NO2)Temperature and Equilibrium Shifts depend on the type of reaction:Endothermic reactionsIncreasing temperature shifts the equilibrium to the right and Keq increases. Decreasing temperature shifts it left, and Keq decreases. Exothermic reactions Increasing temperature shifts it left and Keq decreases. Decreasing temperature shifts it right and Keq increases. Example Problem 12.11 on pg. 516 **Catalysts don’t shift equilibrium as they are not a part of the reaction. Increasing Product Yield (Lab idea?) Useful for predicting the effects of changes on a chemical equilibrium. Chemical companies want to maximize their yield of product and do so by imposing conditions that shift the equilibrium toward the product side of a reaction. Example 12.12 on pg. 517-518 with partnersAssessments: Chemical Reaction Lab, Product Yield Lab, Le Chatelier Lab, Exam Unit 8: Solutions and Their Properties 11.1: Composition of Solutions Solutions are homogeneous mixtures with uniform composition throughout. Recall that the solute is the substance being dissolved in the solvent. Solubility Rules pg. 445…should have from previous chapter) Solutions in which the solvent is water are aqueous solutions. Electrolyte solutions contain a solute that dissociates or ionizes in a solvent producing ions. (most sport drinks accomplish this)Electrolytes include soluble ionic compounds, strong acids, and strong bases. (Recall strong vs weak vs nonelectrolytes) Writing dissociation equationsThe solvent is not typically part of the reaction so when writing the chemical formula for something that dissociates or not, we write H2O (l) above the arrow. Using the solubility rules, we should be able to predict the chemical equation for a given set of reactants. Example 11.1 on pg. 446-447Hydrates Compounds that contain water. Typically lose water to form anhydrous compounds. Typically hydrates are named using the numerical prefixes. Bond is not very strong. Anhydrous if it does not contain water. Can determine the percent mass Percent by mass H2O = (mass of water/mass of hydrate) x 100PP 15.1 (Pearson)**Hydrate lab! 11.2: The Solution Process How do substances form a solution?Typically the first step is for ionic bonds in the solute to break. Hydrogen bonds between water molecules break as well. Ion-dipole forces form between ions and water molecules. Attraction between an ion (from the ionic bond) and a polar molecule (water molecule). Heat of solution The difference between the energy required for the separation of solute and solvent and the energy released upon formation of ion-dipole interactions. If the solvation process provides more energy than is needed to separate the pure solute and solvent particles, the heat of solution is negative and the process is exothermic. If the separation process requires more energy than the solvation process, the heat of solution is positive and it is endothermic. Diagrams pg. 449 Example 11.2 on pg. 450 Entropy Measure of the tendency for the energy of a system’s particles to become more dispersed, resulting in an increase in randomness of the system. Basically it is a measure of disorder. Think about NaCl before and after it is in a solution. 11.3: Factors that Affect Solubility Structure Polar solvents dissolve polar or ionic solutes. Nonpolar solvents dissolve nonpolar solutes. Exception: Soaps and detergents are made so that they have a polar and nonpolar end. The nonpolar end is attracted to oil stains and the polar end is attracted to water. Allows use of soapy water to clean food or dishes.Temperature Solubility of most solids increases in water as the temperature of the solution increases. KE increases and breaks forces holding particles together. (Think of soda left open in the refrigerator vs on the counter) Relationship is often depicted using solubility curves (pg. 254) Pressure Solubility of gases is strongly affected by pressure. Solubility of a gas in a liquid is directly proportional to the pressure of the gas above the liquid. (Henry’s Law)Increase in pressure increases the number of gas particles above the solution, increasing the chance that some will dissolve. (Increased pressure = increased solubility) Examples: Pop, Oxygen masks, and scuba divers. 11.4: Measuring Concentrations of Solutions Concentration, if you recall, is the relative amount of solute and solvent that make up a solution. Dilute vs concentrated. Solubility (lab)Defined as a ratio that identifies the maximum amount of a solute that will dissolve in a particular solvent to form a stable solution under specified conditions. (AKA the max concentration) Types of Solubility Saturated solution: Solution that contains the max concentration of dissolved solute.Unsaturated solution: A solution that contains less than the maximum possible amount of solute. Supersaturated solution: Usually obtained by preparing a saturated solution at a high temperature. Unstable. (adding one crystal could result in excess solute being precipitated) **Video Measurements of Solubility and Units Percent by Mass Grams of solute / grams of solution x 100 Example 11.3 on pg. 458 Percent by volume Volume of solute / volume of solution x 100 Example 11.4 on pg. 459 Mass/volume percent Grams of solute / volume of solution x 100 Parts per million and Parts per billion Common when referring to water or atmospheric amounts that are too small to share as a percentage. Parts per million = grams of solute / grams of solution x 106 (ppm) Parts per billion = grams of solute / grams of solution x 109 (ppb)Example 11.5 on pg. 460-461) Molarity One of the most common expressions in chemistry. Molarity = moles of solute / liters of solution (labeled as mol/L or M) Example 11.6 on pg. 461-462 Molality The problem with molarity is that it will not stay the same as temperature changes because the volume changes. Molality is used instead, as it does not change with changing temperature. Molality = moles of solute / kilograms of solvent (labeled as mol/kg)Example 11.7 on pg. 462 and 463 11.5: Quantities for Reactions that Occur in Aqueous SolutionsPrecipitation Reactions Recall that in a precipitation reaction the cation of one ionic compound combines with eh anion of another ionic compound to form an insoluble, solid substance known as a precipitate. We often need to use molarity with precipitation reactions in order to complete calculations. Example 11.8 on pg. 464-466 (Complex problems using molarity, molar mass, mole ratios, and volume) Acid-Base Titrations (cover in Acid-Base unit) 11.6: Colligative Properties Colligative Properties of Solutions Property of solutions that depend only upon the number of solute particles, not upon the identity. Three important propertiesVapor-Pressure Lowering Vapor pressure (pressure exerted by a vapor on its liquid) is lower in solutions with nonvolatile (don’t form a gas) solutes than the pure solvent. Freezing Point Depression Difference between the freezing point temperatures of a solution and of the pure solvent. Addition of 1 mol of solute to 1,000 g of water lowers the freezing point by 1.86 degrees C. Ex: salt on the roads, antifreezeBoiling Point Elevation Difference between the boiling point temperature of a solution and a pure solvent. Depends on concentration and takes more KE to pull apart solutes from the liquid thus raising the boiling point of a solution compared to the pure solvent. Ex: Antifreeze Freezing Point Depression and Boiling Point Elevation Calculations Magnitude of freezing point depression and the boiling point elevation of a solution are directly proportional to the molal concentration assuming the solute is molecular and not ionic. ΔTf = Kf x m Kf is the molal freezing point depression constant with the units being degrees C/m. ΔTb = Kb x m Example 11.10 on pg. 473 **Water has a Kb value of 0.52 degrees C/m and a Kf value of -1.86 C/m. Colligative Properties and Strong Electrolytes Colligative properties are proportional to the number of solute particles present in the solution. Determine the degree of dissociation. Often expressed as the average number of particles produced by one formula unit of the solute when it dissolves. The molal concentration of ions is the number of ions produced. Ex: 1 m NaCl solution would have a concentration of ions that is 2m. Nonelectrolytes don’t dissociate so a 1 molal concentration of say sugar would have the same molal concentration. Based on this we can assume that a strong electrolyte (such as NaCl) would have 2x the effect on freezing point depression and boiling point elevation as a nonelectrolyte. (even if both are 1 m solutions)Example 11.11 on pg. 474Assessments: Solubility factors lab, calculation practice, colligative properties lab, hydrate labUnit 9: Acids and Bases 13.1 What are Acids and Bases?Acid-Base Theories Arrhenius Definition of Acids and Bases (Svante Arrhenius and electrolytes)Acids produce hydrogen ions (H+) when placed in an aqueous solution. Ex: HClBases produce hydroxide ions (OH-) when placed in aqueous solutions. Explains that acids and bases neutralize each other via a reaction that produces water and a salt. HCl and NaOHClassifies Acids Monoprotic: Ionizes one hydrogen ion. (HCl)Diprotic: Ionizes two hydrogen ions. (H2SO4)Triprotic: Ionizes three hydrogen ions. (H3PO4)Limitations Assumes the release of free H+ ions which are unlikely to exist in aqueous solution because they bond with water molecules. Today, we say that hydronium ions (H3O+) represent H+ ions. Assumes all bases contain OH- ions. (ammonia)Bronsted-Lowry Definition of Acids and Bases (N. Bronsted and T.M. Lowry)Acids are any substances that can donate an H+ ion to another substance. HCl + H2O = H3O+ + Cl- Bases accept an H+ ion. NH3 + H2O = NH4+ + OH-Water is amphoteric…..acting as an acid and a base. Amines can at as bases in water (compounds that contain the –NH2 group)Ex: CH3NH2 + Water Bases can also include anions, such as the carbonate ion. Ex:Na2CO3 = 2NA+ + CO32- Ex: CO32- + H2O = HCO3- + OH-Example 13.1: ID the Bronsted-Lowry acid and base reactants pg. 532HNO3 + H2OSO42- + H2OH2CO3 + CH3NH2Conjugate Acid-Base Pairs When an acid or base donates or gains an H+ ion, it changes the classification of the products formed. Conjugate acids are the products form as a result of gaining an H+ ion, and a conjugate base are the products that forms as a result of losing an H+. Example HCl + H2O = H3O+ + Cl-Acid Base C.A CBExample 13.2 pg. 533: ID the conjugate base for each acid and explain its charge. HOClH2PO4-H2OLewis Definition of Acids and Bases Acids are any substances that can accept a pair of nonbonding electrons, or simply an electron pair acceptor. Bases are any substances that can donate a pair of nonbonding electrons or simply an electron pair donor. ExampleBF3 + NH3 = BF3NH3 (Draw the Lewis Dot)Boron trifluoride accepts an electron pair so it is an acid, while ammonia donates a pair and is the base. Practice problems (other Chemistry book)***Practice Worksheet 13.2 Strong and Weak Acids and Bases Day 2Strong acids and bases are acids and bases that are strong electrolytes and completely ionize in water. (Always use one arrow as it dissociates completely)Strong acids consist of hydrochloric acid, hydrobromic acid, hydroiodic acid, nitric acid, chloric acid, perchloric acid, and sulfuric acid. Strong bases include lithium hydroxide, sodium hydroxide, potassium hydroxide, magnesium hydroxide, calcium hydroxide, and barium hydroxide. Includes most ionic hydroxides of groups 1 and 2. Includes many cleaners. Weak acids and bases are acids and bases that are weak electrolytes and only partially ionize in water. (Always use double arrows in equations due to the fact it forms at equilibrium)Weak acids do not ionize completely when dissolved in water. An equilibrium forms between a weak acid and its conjugate base. Includes acetic acid, carbonic acid, citric acid, lactic acid, oxalic acid, phosphoric acid, tartaric acid, hydrofluoric acid, hypochlorous acid, and malic acid. Fewer than 10% of the molecules transfer an H+ to water, leaving the conjugate base. Ex: CH3CO2H + H2O (double arrows) CH3CO2- + H3O+ Equilibrium lies far to the left of the equation. Weak bases do not completely ionize when dissolved in water. The equilibrium forms between a weak base and its conjugate acid. Most common weak bases are organic compounds that contain the amine group. (-NH2) Others include ammonia, calcium carbonate, and calcium hypochlorite. Ex: CH3NH2 + H2O (double arrow) CH3NH3+ + OH-Weak bases can also be determined by the anions. If the anion is the conjugate base of a weak acid, it will behave as a base when added to water. (Therefore it is a weak base)Example: Sodium acetate (NaCH3CO2) dissolves in water for form Na+ and CH3CO2-. The latter ion is a conjugate base of acetic acid (weak acid) so it indicates Sodium acetate is a weak base. If it is a conjugate base of a strong acid, it will not react. Example 13.4 ID each as a strong or weak acid or base. Write an equation to describe its reaction with water. Pg. 539CH3CH2NH2 (weak base)HF (weak acid)NaF (weak base)KOH (strong base)On your Own Practice HINaCH3CO2NH4+NH313.3 Relative Strengths of Weak Acids Day 3Acid strength depends on the degree of ionization. Stronger the acid, larger its equilibrium constant value for ionization in water and the greater percentage of H3O+ and conjugate base ions produced. Example Ionization of acetic acid: Keq = 1.8 x 10^-5Ionization of hydrocyanic acid: Keq = 6.2 x 10^-10Which is stronger?The stronger acid produces more products at equilibrium and the further to the right equilibrium is. Therefore, the larger the equilibrium constant, the stronger the acid. Acid Ionization Constants Describes the equilibrium that forms when an acid reacts with water. Value is represented by Ka. Larger the Ka value, the stronger the acid. Keq values equal Ka values. Stronger the acid, the weaker the conjugate base. Examples 13.6 Significance of KaWhich acid ionizes in water to the greater extent, hypochlorous acid or hydrocyanic acid, HCN? Which would have the stronger base?Which solution has the greatest concentration of H+, a 0.10 M solution of HOCl or a 0.10 M solution of HCN?Calculating Acid and Base dissociation constants for Weak Acids Ka = [H+] [A-] / [HA]Calculations via ICE Method ((Physical_and_Theoretical_Chemistry)/Equilibria/Le_Chatelier%27s_Principle/Ice_Tables )Example: In a 0.100 M solution of ethanoic acid, [H+] = 1.34 x 10-3 M. Calculate the Ka of theis acid. We assume it ionizes to equilibrium so the ion concentration should equal the H+ oncentration. The amount of acid left can be determined by subtracting the starting concentration by the H+ concentration assuming we are at equilibrium. Ka = 1.82 x 10-5Can also use Ka values and the ICE methods to find H+ concentrations and ion concentrations. % Dissociation = [H3O+] / [HA] x 100Can also calculate pH using later equation. (See worksheet) Kb = [conjugate acid] [OH-] / [base]Polyprotic Acids Day 4 with Practice In-ClassAcids that contain more than one acidic hydrogen atom. They lose their hydrogens as a stepwise process. Example: Carbonic acid ionizing in waterH2CO3 + H2O (double arrow) H3O+ + HCO3- Ka = 4.5 x 10^-7HCO3- + H2O (double arrow) H3O+ + CO32- Ka = 4.7 x 10^-11The lower the Ka value indicates that the equilibrium lies towards the reactants and only a small portion ionizes. Small ionization indicates a weak acid. Example Problem 13.7The tartaric acid, H2C4H4O6, in grapes promotes a crisp flavor and graceful aging in wine. Its Ka values are 1.0 x 10^-3 and 4.3 x 10^-5. Write equations that show the ionization of tartaric acid in water. Besides water, which ion or molecules has the highest concentration in solution when tartaric acid is added to water?Do the same problem but for Oxalic acid. Its Ka values are 5.6 x 10^-2 and 1.5 x 10^-413.4 Acidic, Basic, and Neutral Solutions Day 5 with HWKWe base our concentrations and equilibrium constants off of the self-ionization of water. Kw = [H3O+] [OH-] = 1.0 x 10-14 at 25 C. **Constant Pure water (neutral): [H3O+] = [OH-] = 1.0 x 10-7 M Acidic: Greater H3O+ than OH-Basic: Greater OH- than H3O+Example 13.8Given the concentration of H3O+ ion in water, calculate the concentration of OH- ion in the solution. Identify each solutions as acidic, basic, or neutral:[H3O+] = 1.0 x 10-9 M[H3O+] = 0.0010 M Without doing any calculations, determine if a solution with an OH- concentration of 1.0 x 10-10 M would be acidic or basic? Given the concentration of OH- ion in each solution, calculate the concentration of H3O+ in the solution. ID each as acidic, basic, or neutral. [OH-] = 1.0 x 10-8 M[OH-] = 0.010 M Example 13.9 Stomach acid is about 1.0 M HCl. What is the H3O+ concentration in stomach acid? What is the OH- concentration?Strong acids and bases ionize completely, so the concentration of H3O+ or OH- concentrations will match the concentration on the bottle. In this case, the [H3O+] = 1.0 M. What is the concentration of OH-?Sodium hydroxide is sometimes called lye or caustic soda. What is the OH- concentration in a 0.85 M lye solution? What is the H3O+ concentration? 13.5 / 19.2 The pH Scale (Should be familiar with this at this point)pH allows us to connect the pH value to the concentration of H3O+. Log scale based on a power of 10. pH = -log[H3O+] Example 113.10 Calculating pHWhat is the pH of each of the following solutions?0.0010 M HBr0.035 M HNO30.035 M KOH0.00085 M HCl 0.10 M NaOH1.0 M HNO3Calculating pOHpOH = -log[OH-]pH + pOH = 14Application QuestionOlivia has blue flowers in her garden, but wants some of her plants to show flowers that are pink. After some research, she discovers that the flower color, in part, depends on the pH of the soil. Blue flowers occur when soil is in the pH range of 5.5 to 6.5. For pink flowers, the pH needs to be slightly more alkaline, in the range of 7.0 to 7.5. When she measures the pH of a soil solution from her garden, she learns that it is slightly acidic with a pH of 6.20. She learns that making the soil more alkaline to produce pink flowers can be achieved by adding calcium oxide to the soil. What is the [H3O+] of her soil?10-pH = [H3O+]10-pOH = [OH-]Example 13.11 Jake measured the pH of water in a swimming pool as 8.10. What is the OH- concentration in the pool water?Measuring pH pH meters (very accurate) pH indicators organic dyes that indicate if a solution is acidic or basic based on the color it forms. HIn is an acid or a base that dissociates in a known pH range. HIn (double arrow) H+ + In-HIn is dominant at low pH while In- is dominant at high pH. Differ based on color…see pg. 660 in Pearson book. 13.5a (19.4 via Pearson) Day 6 with in class practice + Lab Neutralization Reactions Acids and bases react to produce a salt and a water via a neutralization reaction. Ex: HCl + NaOH yields Water and NaClFinding the Moles Needed for Neutralization The term neutralization is used to describe both the reaction and the point at which a neutralization reaction is complete. How many moles of sulfuric acid are required to neutralize 0.50 mol of sodium hydroxide? H2SO4 + 2NaOH yields Na2SO4 + 2H2O Must rely on mole ratios to solve0.50 mol NaOH x (1 mol H2SO4 / 2 mol NaOH) = 0.25 mol H2SO4. Titration Neutralization is useful to determine the concentration of an unknown solution via titration. 3 steps for acid-base titrations A measured volume of an acid solution of unknown concentration is added to a flask. Several drops of an indicator (phenolphthalein) are added to the solution while the flask is gently swirled. Measured volume of a base of known concentration (standard solution) is mixed into the acid until the indicator just barely changes color (end point). (faint pink)Equivalence point: point at which neutralization occurs and the moles of hydrogen ions and hydroxide ions are equal. Very near the end point. Produce a titration graph on pg. 674 of Pearson book. Determining Concentration by Titration A 25 mL solution of H2SO4 is neutralized by 18 mL of 1.0 M NaOH. What is the concentration of the acid solution? H2SO4 + 2NaOH yields Na2SO4 + 2H2OSimple mole ratios and some conversion factors 0.018 L NaOH x (1.0 mol NaOH / 1 L NaOH) = 0.018 mol NaOH0.018 mol NaOH x (1 mol H2SO4 / 2 mol NaOH) = 0.0090 mol H2SO4Molarity = 0.0090 mol H2SO4 / 0.025 L = 0.36 M H2SO413.6 / 19.5 Buffer and Salts Day 7 with HWK and Salt Lab Salt Hydrolysis Remember that salts form during a neutralization reaction. A salt consists of an anion from an acid and a cation from a base. Salt hydrolysis occurs when the cations or anions of a dissociated salt remove hydrogen ions from, or donate hydrogen ions to, water. Salts that produce acidic solutions have positive ions that release hydrogen ions to water. Salts that produce basic solutions have negative ions that attract hydrogen ions from water. ExampleCH3COONa yields CH3COO- + Na+CH3COO- + H2O yields CH3COOH + OH-Because the salt’s ion attracts hydrogen, it is basic. What about the dissociation of NH4Cl?Rules Strong acid + strong base yields neutral solutionStrong acid + weak base yields acidic solution Weak acid + strong base yields basic solution. Buffers Buffer systems are combinations of a weak acid and its conjugate base or a weak base and its conjugate acid in equal concentrations. Strong acids don’t make good buffers as they ionize completely. Regulates pH by resisting changes in pH. Ex: carbonic acid – hydrogen carbonate buffer system that maintains blood pH. How it works?Contains one component that can react with hydrogen ions and another component that can react with hydroxide ions. Buffer capacity is the amount of acid or base that can be added to a buffer solution before a significant change in pH occurs. Writing Buffer Systems using H2CO3 / HCO3- Buffer System H2CO3 + H2O (double arrow) HCO3- + H3O+When a base is added the OH- reacts with the acid ro produce water. When an acid is added, the hydrogen ions react with the conjugate base to form water. (pg. 556 McGraw)Practice Problems Another buffer system in our blood is the H2PO4-/HPO42- system. Describe, using a balanced equation, how this system prevents the pH from falling too low when an acid enters the bloodstream. What about when a base enters the bloodstream?Which of the following, when added to water, can act as a buffer system? For each buffer system, write a balanced equation. HF and NaF (Buffer system)HNO3 and KNO3 (Not a buffer)NH3 and NH4Cl (Buffer system)Henderson-Hasselbalch Equation (Video) Useful expression for buffer calculations. pH = pKa + log ([A-] / [HA])Assessments: pH calculation lab, Titration lab, Buffer lab Unit 10: Gas Laws 9.1: The Behavior of Gases Gases are often described in terms of pressure, volume, temperature, density, and the amount of the gas. However more specific properties that explain the properties above include:Gases consist of particles that are relatively far apart. (explains the low densities of gases)Gas particles move about rapidlyGas particles have little effect on one another unless they collide. Gases expand to fill their containers. Temperature and Density Higher the temperature = lower the density. (Think gases in the atmosphere, hot air balloonsPressure Pressure is the amount of force applied per unit are. P = force / areaFor gases it typically means Pressure = force of gas particles / area of container. (Demo difference in pressure using a nail board and lay on it.) Atmospheric pressure measured with a barometer. Used to be filled with mercury and the height of the mercury would change to match atmospheric pressure. 1 atm = 76 cm Hg = 760 mm Hg = 760 torr1 atm = 101,325 Pa (Pascal is the SI unit of pressure)Example 9.1 on pg. 350 9.2: Factors that affect the properties of gases (PBL)Volume and Pressure (Lab)Boyle’s Law: For a given mass of gas at constant temperature, volume varies inversely with pressure. Simply, as pressure increases, volume decreases. (Graph on pg. 352)P1V1 = P2V2A gas that follows this law and does not vary is known as an ideal gas. Example 9.3 on pg. 354 Volume and Temperature (Lab)Charles’s Law: For a given mass of gas at constant pressure, volume is directly proportional to temperature on an absolute (Kelvin) scale. Simply, if temperature increases, volume increases. (Graph on pg. 355)V1/T1 = V2/T2 Example 9.5 on pg. 357 Volume, Pressure, and Temperature Combined Gas Law (Lab)States for a constant amount of gas, volume is proportional to absolute temperature divided by pressure. (P1V1/T1) = (P2V2/T2) Example 9.6 on pg. 359-360Gay-Lussac’s Law States that the pressure of a gas is directly proportional to the Kelvin temperature if the volume remains constant. P1/T1 = P2/T2Example 14.3 on pg. 461 in Pearson book Avogadro’s Hypothesis Hypothesized that the volume occupied by a gas at a given temperature and pressure is proportional to the number of gas particles and thus to the moles of the gas. (V1/n1) = (V2/n2) Example 9.7 on pg. 361-362The official hypothesis states that at a given pressure and temperature, equal volumes of all gases contain equal numbers of moles. Molar volume is the volume of any ideal gas. 22.414 L. Occurs at a temperature of 0 degrees Celsius and 1 atm. (Standard temperature and pressure: STP) Example 9.8 and 9.9 on pg. 362-3639.3: Ideal Gas Law The earlier relationships all assume a certain constant occurs in order for the relationship to occur. They can all be combined to form a single relationship known as the ideal gas law. PV = nRT **R represents the ideal gas constant. Under STP the value of R was found to be 0.08206 L atm / mol K Calculations with Ideal Gas Law Moles of a Gas Example 9.10 on pg. 365 Mass of a Gas Example 9.11 on pg. 366 Density of a Gas Density = mass / volume Apply to the gas in Examples 9.19 and 9.11 Determining molar mass from gas densityUsed in situation where an unknown gas’s density is given. You can use the density and the ideal gas law to calculate the molar mass and potentially its identity. Example problem on pg. 367-368 Dalton’s Law of Partial Pressures Useful for mixtures of gases where we want to know the properties of one or more of the individual substances. Dalton’s Law Gases in a mixture behave independently and exert the same pressure they would exert if they were in a container alone. Ptotal = PA + PB + PC............When water is involved we often need to take the total pressure and subtract the vapor pressure at the given temperature in order to the find the partial pressure of our gas. (Table 9.2 on pg. 369) Example Problem 9.12 on pg. 369-370Graham’s Law Two terms used to describe the movement of gases:Diffusion: the tendency of molecules to move toward areas of lower concentration. (Think Body spray) Effusion: A gas escapes through a tiny hole in its container. Gases of lower molar mass diffuse and effuse faster than gases at higher molar mass. Graham’s law of effusion States that the rate of effusion of a gas is inversely proportional to the square root of the gas’s molar mass. (RateA / RateB) = sqr root (molar massB / molar massA ) Example problem 14.8 on pg. 474 in Pearson book 9.4: Kinetic-Molecular Theory of Gases Model that explains experimental observations about gases under normal temperature and pressure conditions we encounter. Five postulates Gases are composed of small and widely separated particles. (explains volume, density)Particles of a gas behave independently of one another (explains Dalton’s Law)Each particle in a gas is rapid, straight-line motion until it collides with another molecule or with its container. (explains why gas fills its container)The pressure of a gas arises from the sum of the collisions of the particles with the walls of the container. (Boyle’s law)The average kinetic energy of gas particles depends only on the absolute temperature. (explains why gases with smaller molar mases have greater energy) 9.5: Gases and Chemical Reactions Product Volume from Reactant Volume The conversion below is used to obtain the product volume from the reactant volume in a chemical equation:Volume A is converted to mol A using the ideal gas law. Mol A is converted to mol B using mole ratios. Mol B is converted to volume of B using ideal gas law. Example 9.13 on pg. 373-374Moles and Mass From Volume Conversion allows us to find the moles and mass of a product from the reactants using the following conversion:Volume A converted to moles of A using ideal gas law. Moles of A converted to moles of B using mole ratios. Moles of B converted to mass of B using molar mass. Example 9.14 on pg. 375 Assessments: Gas Law Labs, Calculation Practices, Exam ................
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