General Chemistry I&II Chapter Notes
General Chemistry I&II Chapter Notes CH 102
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GASES I. Pressure is the force exerted per unit area by gas molecules as they strike the surfaces around them. The total pressure exerted by a gas depends on several factors including the concentration of gas molecules. A. As volume increases, pressure goes down. This results in fewer molecular collisions, resulting in lower pressure. When volume decreases, pressure increases. B. Because of pressure, we can drink from straws, inflate basketballs, and breathe. Variation in pressure in earth’s atmosphere creates wind, and changes in pressure can help us to predict weather. 1. Pressure that a gas exerts is the force that results from the collisions of gas particles divided by the area of the surface with which they collide. 2. The pressure exerted by a gas sample depends on the number of gas particles in a given volume. The fewer particles, the lower the force per unit area and the lower the pressure. Pressure decreases with an increasing altitude. 3. Pressure is measured in several different units. A common unit of pressure, the millimeter of mercury (mmHg), also called a torr, originates from how pressure is measured with a barometer (and evacuated glass tube, the tip of which is submerged in mercury. The liquid mercury is forced upwards by atmospheric pressure. a. Another unit of pressure is the atmosphere (atm), the average pressure at sea level. 1 atm = 760 mmHg b. The SI unit of pressure is the pascal (Pa), defined as 1 newton. 1 Pa = 1 N/ m 2 . c. The pascal is a much smaller unit of pressure than the atmosphere. 1 atm = 101, 325 Pa d. 1 atm= 29.92 inHg; 1 atm= 14.7 psi C. The Manometer 1. Manometers can measure the pressure of a gas sample. It is a U shaped tube containing a dense liquid, usually mercury. II. Simple Gas Laws A. Boyle’s Law: Volume and Pressure 1. A sample of gas has four basic physical properties: pressure, volume, temperature, and amount in moles. 2. English scientist Robert Boyle and his assistant Robert Hooke used a J-tube to measure the volume of a sample of gas at different pressures. They trapped a sample of air in the J-tube and added mercury to increase the pressure on the gas. They found an inverse relationship between volume and pressure- an increase in one results in a decrease in the other. This is known as Boyle’s Law. a. Boyle’s law follows the idea that pressure results from collisions between gas particles with the walls of their containers. b. P 1 1P V2 2 B. Charle’s Law: Volume and Temperature 1. The first person to carefully examine the relationship between volume of a gas and its temperature was J.A.C. Charles. If the temperature increases, the volume increases in direct proportion so that the quotient is always equal to the same constant. V1 V2 2. = T1 T2 C. Avogadro’s Law 1. When the amount of gas in a sample is increased at constant temperature and pressure, its volume increase in direct proportion because the greater number of gas particles fill more space. V V 2. 1= 2 n1 n 2 D. The Ideal Gas Law 1. The volume of gas is directly proportional to the number of moles of gas to the temperature of the gas, but is inversely proportional to the pressure of the gas. RnT 2. V= P L∙atm 3. R, the ideal gas constant is .08206 mol∙K 4. PV= nRT 5. The relationship between pressure and temperature is known as Guy- Lussac’s law; the temperature of a fixed amount of gas in a fixed volume increases, the pressure increases. III. Applications of the Ideal Gas Law: molar volume, density, and molar mass of a gas A. Molar Volume at Standard Temperature and Pressure: The volume occupied by one mole of any gas at T= 0℃ (273 K) and P= 1.00 atm can be easily calculated using the ideal gas laws. These conditions are called standard temperature and pressure (STP). The volume occupied by one mole of any gas under these conditions is called the molar volume of an ideal gas. RnT B. Using the ideal gas law, molar volume is V= P = 22.4 L. Molar volume is useful because it gives the volume of a gas under standard conditions. C. Density of a Gas: Since one mole of an ideal gas occupies 22.4 L, the density of an ideal gas can be calculated under standard conditions. Since density is mass/volume, and since one mole of gas is simply the molar mass, the density of a gas under standard conditions is given by the following: molarmass 1.Density= molarvolume ; g/L 2. Notice density of a gas is directly proportional to its molar mass. The greater the molar mass of a gas, the more dense it is. PM 3. d= RT D. Molar Mass of a Gas 1. The ideal gas law can be used in combination with mass measurements to calculate the molar mass of an unknown gas IV. Mixtures of Gases and Partial Pressure A. The pressure due to any individual component in a gas mixture is called the partial pressure ( Pn¿ of that component and can be calculated from the ideal gas law by assuming that each gas component acts independently. For a multicomponent gas mixture, the partial pressure of each component can be computed from the ideal n gas law and the number of moles of that component (¿¿n) as ¿ follows: Pa CHEMISTRY 101 FINAL STUDY GUIDE CHAPTER 1: 1. Pure substances and mixtures Pure Substances: matter with a fixed composition and distinct properties; Can be elements or compounds o Elements: can be atomic or molecular o Cannot be chemically broken down o Compounds: cannot be separated by physical means o Ex- water, ammonia, diamond, helium Mixtures: matter with a variable composition and contains more than 1 pure substance o Can be separated by physical means o Homogeneous: Components are uniformly distributed o Ex-Air, sugar water, rain, vodka, detergent o Heterogeneous: Components are not uniform o Ex- Soup, pizza, blood, gravel, dressing 2. Physical and Chemical Properties/ Changes Physical properties: odor, taste, color, appearances, melting point, boiling point, density o Changes: Does not change composition o Boiling, melting, freezing, sublimating, dissolving Chemical properties: corrosiveness, flammability, acidity, and toxicity o Changes: Changes chemical composition o Rusting, Burning, colors fading, eggs cooking, bread rising, milk souring 3. Identify& Characterize Different States of Matter Shape Volum Compressibilit Flow e y Solid Fixed Fixed Non Vibrate Ex- ice, Compressible s aluminu tightly m, in diamon place d Liquid Variabl Fixed Non Flows Ex- e Compressible water, alcohol, gasolin e Gas Variabl Variabl Compressible Flows *Solid matter can e e Quickl be crystalline, like y table salt and diamond, meaning its atoms or molecules are in patterns and with long repeating order. They may be amorphous too, like glass and plastic; the atoms will not have any order. * 4. Scientific Method o Accuracy is an indication of how close a measurement comes to the actual value of the quantity o Precision is an indication of how close repeated measurements are to each other; how reproducible a measurement is 5. Dimensional Analysis& SI Units Conversion to other metric units commonly used in Chemimass kg 1 picometer (pm) = 10 lemgth m -9 1 nanometer (nm) = 10 t-6e sec 1 micrometer (m) = 10 m 1 millimeter (mm) = 10 mmperature K 1 centimeter (cm) = 10 m 3electric current Amp 1 kilometer (km) = 10 amount of mol substance luminous candela intensity Volume, Density 3 3 o Solids: g/ cm or Kg/ m o Liquids: g/mL, g/L 6. Solve problems with density and volume D= mass/volume 7. Convert between Celsius, Kelvin, and Fahrenheit (F32) C 1.8 K C 273.15 8. Significant Figures Mult/Div: Answer has same number as number with fewest sig figs Add/ Sub: Fewest decimal places 9. Formulas for Volume of Solid Objects Cube: s3 Sphere: 4 πr3 3 Cylinder: 2 πr h CHAPTER 2 1. Elements, Compounds, and Mixtures Atomic elements: exist in nature with single atoms as their basic units. o For example: He, Al, Fe, Mg. o Noble gases, metals, and metalloids Molecular Elements: do not normally exist in nature with single atoms as their basic units. Instead, these elements exist as molecules o For example: Non- metals O2, H2, N2, Cl2. o Others exist as polyatomic molecules, like phosphorous and sulfur. Molecular Compounds: composed of two or more covalently bonded nonmetals. o For example, water is composed of H2O molecules, dry ice is composed of CO2 molecules, and propane is composed of C3H8 molecules. 2. Mass Laws Law of Conservation of Mass: Mass is neither created nor destroyed during any chemical reaction. o The mass of substances produced (products) by a chemical reaction is always equal to the mass of the reacting substances (reactants). o Ex- Formation of water from its elements 2H 2 + O2 2H O 2 4 g + 32 g 36 g Law of Definite Composition: in a pure compound, the elements are always present in the same definite proportion by mass. (Mass of element/mass of compound is fixed for any compound) o Ex- H2O 2.0160g (2x1.0080) of hydrogen will combine with 15.9994g of oxygen to produce 18.0154g of H2O and 4.032 of hydrogen will combine with 31.988g of oxygen to produce 36.020g of H2O Ratio of H:O is always 2:16 or 1:8 Element AMass o Mass Ratio= ;Ratio=decimal:1 Element B Mas Law of Multiple Proportions: when two elements (call them A and B) form two different compounds, the masses of B that combine with 1 g of A can be expressed as a ratio of small, whole numbers o Ex- Two compounds of C and O are CO and 2O Ratio of the mass (or mass%) of O to the mass of C in the two compounds is 2:1 Mass % O in CO =272.7 % Mass % O in CO = 57.1 % Mass% C in CO =227.3 % Mass% C in CO = 42.9 % Mass of O/C in CO 2 2.66 Mass of O/C in CO = 1.33 MassOxygen o MassOxygen = 2 ¿ 2.67 ¿1gC∈CO2 1g¿∈CO¿= 1.33 3. Determine Mass % from Chemical Formula Mass of X= masselement of X∈1mol com×100 Mass1molComp 4. Relate Dalton’s Atomic Theory to Mass Law Dalton’s Atomic Theory Modern Atomic Theory All matter consists of extremely small All matter is composed of atoms that indivisible particles called atoms are divisible and made up of smaller subatomic particles, electrons, protons and neutrons. Atom is still the smallest particle that retains the unique identity of the element Atoms of an element are identical in All atoms of an element have the mass and other properties and are same number of protons and different from atoms of other elements electrons, which determine the chemical behaviour of the element Compounds result from the chemical Compounds are formed by the combination of two or more atoms of chemical combination of two or different elements in a specific ratio. more elements in specific ratios. Atoms of one element cannot be Atoms of one element cannot be converted into atoms of another element, converted to atoms of another atoms can neither be created not element in a chemical reaction destroyed, only rearranged in a chemical reaction to produce new compounds. 8. Average Atomic Mass from %Abundance and Isotopic mass Avg mass = (decimal abundance x mass) +(decimal abundance x mass) 9. Periodic Table 10. Grams to Moles to Atoms CHAPTER 3 1. Ionic and Covalent Compounds Ionic Compounds: When a metal interacts with a nonmetal o Transfers one or more electrons to the nonmetal. The metal becomes a cation and the nonmetal becomes an anion. The oppositely charged ions attract one another by electrostatic forces and make an ionic bond. o Polyatomic ions o Characterized by high melting points Covalent Compounds: When a nonmetal bonds with another nonmetal o Neither atom transfers its electron to the other. The bonding atoms share their electrons. o The shared electrons have lower potential energy than they would in isolated atoms because they interact with the nuclei of both atoms. This is called a covalent cond. o Covalently bonded compounds are called molecular compounds because the individual molecules are not covalently bound to one another. 2. Strength of an ionic compound 1 q1q2 Coulomb’s law: E= ; q1 is the charge on the cation& q2 is the charge 4πε r on the anion. r is the distance between the two. Trends in lattice based on charge: Compounds with higher total charge (q1q2) are higher in lattice energy. For example, CaO has a higher lattice energy than NaF. Trends based on ionic size: If charges are equal, we look at distance between the atoms which depends on size. The smaller ion will result in smaller distance, r, between the two atoms and therefore have a higher lattice energy, E. o 3. Polyatomic Ions 4. Binary Acids and Oxyacids 5. Functional Groups& Organic Naming Hydrocarbons o Alkanes: Saturated hydrocarbons- C H n 2n+2,H ,4C H2, 6 H3, 8 H 4 C10 5 12,C6H 14 o Alkenes: Unsaturated hydrocarbons - C H n 2n, 2 4 o Alkynes: Unsaturated hydrocarbons - C H n 2n-2,2 2 Functionalized Hydrocarbons o Alcohols: Functional group – OH (hydroxyl group) – ol ending CH 3H, C H 2H 5 o Carboxylic Acids: Functional Group – COOH (carboxyl group) – oic acid ending HCOOH, CH COOH 3 o Ether (-O-): CH OC3 3 - Dimethyl ether o Amine (-NH ): 2 CH 3H 2 o Aldehydes (-CHO) and Ketones (-CO-): CH 3HO and CH COCH 3 3 6. Molecular Mass vs. Molar Mass Molar Mass= grams per mole; mass of one mole of a substance Molecular Mass= mass of one atom or molecule of a substance; measured in amu (periodic table values) 7. Determine Mass Percent from Chemical Formula molarmass of element∈compound Mass % of element = molarmass of compound 8. Determine empirical formula from masses of elements Convert the given mass % to masses by assuming a mass of the compound to be 100 g. Convert the information in grams to moles using the molar mass of the element. Divide the resultant moles by the smallest number of moles If there are fractional subscripts convert these to whole numbers by multiplying through with the smallest integer that can convert the fractional subscripts to whole numbers. 9. Determine molecular formula knowing empirical formula and molecular mass Molar Mass Molecular Formula = Empirical Formula x n; n= EmpiricalFormulaMass 10. Empirical Formula from Combustion Analysis CHAPTER 6 Energy flow: Find change in internal energy: ∆ E=q+w Quantifying Heat: o Heat from specific heat capacity: q= mC ∆ T o Specific heat to determine temperature changes: q= mC ∆ T Heat Transfer −q =q o metal water Measuring Internal Energy of a Chemical Reaction: o Given grams, specific heat, temperature change: q =C ×∆T -> −q =q cal cal cal rxn -> ∆ E=qrxnJ)/moles o Given grams, delta E: Use the formula C H 8 18for gasoline to calculate n. Solve for number of moles, n, from the given mass. Substitute in the equation ∆ H rxn q rxnto solve for q rxnhen using the relationship given below, solve for T: -qrxn q cal CcalT Heats of Reaction and Enthalpy Change: o To solve for enthalpy change given heat, pressure, and volume change: use w=−P∆V to solve for w. Then convert L × atm to kilojoules using 101.3 J = 1 Latm. Use ∆ E=q+w to solve for internal energy change. Then use ∆ H=∆E+P ∆V to solve for enthalpy change in kilojoules. Delta H and Stoichiometric Calculations Given a chemical reaction, grams reactant, and heat, find q per mole reactant: convert grams to moles. Divide moles by heat. Given mol reactant, heat, write a reaction including physical states: Combustion= gas + oxygen = CO2 + H2O Find heat absorbed given reaction, grams reactant, and ∆ H : Convert grams to moles. Use ∆ H rxn= qrxn to solve for heat. Constant Pressure Calorimetry o ∆ H rxn Laboratory Method of measuring heat evolved or absorbed in a chemical reaction: o Finding q rxniven mL and M of two reactants (acid and base), and two temperatures: Write down chemical equation. Since density is 1.00 g/mL, add the two volumes to find total volume, and use that to find grams. Specific heat is 4.18 J/gC. Use q= mC ∆ T to solve for q in kJ, use −q rxn solno flip sign accordingly. Divide by the amount in moles to get q rxn . o Finding the enthalpy of neutralization given the density of one reactant and the specific heat capacity of the reactant: Relationships involving ∆ H rxn. o Use Hess’s Law to find ∆ H: o Examples: o Calculate H for the reaction, 2NOCl (g) N (g)2+ O (g) 2 Cl (g) f2om the following information 1/2N (2) + 1/2 O (g2 NO (g) H = 90.3 kJ NO(g) + ½ Cl (g2 NOCl (g) H = -38.6 kJ H = -103.4 kJ o Ammonia will burn in the presence of a platinum catalyst to produce nitric oxide, NO. 4NH (g3 + 5O (g)2 4NO(g) + 6H O(g) 2 o What is the heat of the reaction at constant pressure? Given N 2g) + O (g2 2NO (g) H = 180.6 kJ N 2g) + 3H (g2 2NH (g) 3 H = -91.8 kJ 2H (2) + O (g2 2H O(g)2 H = -483.7 kJ ∆ H = -906.3 kJ Determining ∆ Hrxn from ∆ H ° f given a reaction: o o o ∆ H rxn= n ∆ H f(products)m ∆ H f(reactants)ere, m is the molar coefficient of the reactant and n is the molar coefficient of the product. I. Elements and compounds A. Atomic elements 1. Atomic elements are those that exist in nature with single atoms as their basic units. For example: He, Al, Fe, Mg. 2. Noble gases, metals, and metalloids B. Molecular Elements 1. Molecular elements do not normally exist in nature with single atoms as their basic units. Instead, these elements exist as molecules, two or more atoms of the element bonded together. For example: Non- metals O2, H2, N2, Cl2. 2. Most molecular elements exist as diatomic molecules. Others exist as polyatomic molecules, like phosphorous and sulfur. 3. Do not confuse Molecular ELEMENTS with Molecular COMPOUNDS. Molecular compounds are usually composed of two or more covalently bonded nonmetals. For example: H2O and CO2. C. Molecular Compounds 1. Molecular compounds are usually composed of two or more covalently bonded nonmetals. For example, water is composed of H2O molecules, dry ice is composed of CO2 molecules, and propane is composed of C3H8 molecules. C. A chemical compound is a chemical substance consisting of two or more different chemically bonded chemical elements, with a fixed ratio determining the composition. 1. For example: H2O, NaCl D. Chemical Bonds 1. Compounds are made of atoms held by chemical bonds. Chemical bonds are forces of attraction between atoms. They result from opposing charges, and the electrostatic forces are responsible for chemical bonding. 2. Ionic Compounds: a. When a metal interacts with a nonmetal, it can transfer one or more electrons to the nonmetal. The metal becomes a cation and the nonmetal becomes an anion. The oppositely charged ions attract one another by electrostatic forces and make an ionic bond. b. The basic unit of an ionic compound is the formula unit, the smallest, electrically neutral collection of ions. c. A polyatomic ion is an ion composed of two or more atoms. d. Characterized by high melting points 3. Covalent Bonds 1. When a nonmetal bonds with another nonmetal, neither atom transfers its electron to the other. The bonding atoms share their electrons. The shared electrons have lower potential energy than they would in isolated atoms because they interact with the nuclei of both atoms. This is called a covalent cond. Covalently bonded compounds are called molecular compounds because the individual molecules are not covalently bound to one another. II. Chemical Formulas and Molecular Models A. The easiest way to represent a compound is with its chemical formula (Ex- H2O). Chemical formulas usually list the more metallic, or the positively charged elements first followed by the negatively charged element. B. Chemical formulas can either be empirical, molecular, or structural. 1. An empirical formula gives the relative number of atoms of each element in a compound. A molecular formula gives the actual number of atoms of each element in a molecule of compound. a. The empirical formula for hydrogen peroxide is HO, but the molecular formula is H2O2. b. The molecular formula is always a whole number multiple of the empirical formula. Sometimes they are identical. 2. A structural formula uses lines to represent covalent bonds and shows how atoms in a molecule are bonded to each other. a. The structural formula for H2O2 is H—O—O—H. This can be written in a different way to give a sense of the molecule’s geometry. C. Molecular Models 1. A molecular model is a more accurate and complete way to specify a compound. A ball and stick molecular model represents atoms as ball and chemical bonds as sticks; how the two connect reflects the molecule’s shape. 2. In a space fitting molecular model, atoms fill the space between each other to more accurately represent how a molecule might appear if scaled to size. III. Ionic Compounds: Formulas and Names A. Ionic compounds occur throughout earth’s crust as metals. Ionic compounds are generally very stable because the attractions between cations and anions within ionic compounds are strong and because each ion interacts with several oppositely charged ions in the crystalline lattice. B. Writing formulas for ionic compounds 1. Since ionic compounds are charge neutral, and since many elements form only one type of ion with a predictable charge, we can get the formulas from their constituent elements. NAMING The most metallic element comes first when naming. IV. Formula Mass and the Mole Concept for Compounds A. The average atomic mass of an element is atomic mass. The average mass of a molecule of a compound is the formula mass. The formula mass is the sum of the atomic masses of all the atoms in its chemical formula. B. Molar Mass of a Compound 1. An element’s molar mass= mass in grams of one mole of compound. Using Avogadro’s number as a conversion factor will give you number of atoms in a given mass. The molar mass is a conversion factor between mass in grams and amount in moles. V. Organic Naming A. Hydrocarbons 1) Alkanes: Saturated hydrocarbons- C H n 2n+2 CH ,4C H2, 6 H 3 8 H 4 C10 5 12, 6 14 2) Alkenes: Unsaturated hydrocarbons - C H n 2n C 2 4 3) Alkynes: Unsaturated hydrocarbons - C H n 2n-2 C H 2 2 B. Functionalized Hydrocarbons 1) Alcohols: Functional group – OH (hydroxyl group) – ol ending CH O3, C H O2 5 2) Carboxylic Acids: Functional Group – COOH (carboxyl group) – oic acid ending HCOOH, CH COOH3 3) Ether (-O-): CH OC3 3 - Dimethyl ether Amine (-NH ): 2 CH N3 2 4) Aldehydes (-CHO) and Ketones (-CO-): CH 3HO and CH COC3 3 VI. Mass Percent A. Mass % of element = molar mass of element in comp x 100 Molar mass of comp B. Using mass percent to determine empirical formula Empirical Formula: Get the smallest whole number ratio from the given information. 1. Convert the given mass % to masses by assuming a mass of the compound to be 100 g. 2. Convert the information in grams to moles using the molar mass of the element. 3. Divide the resultant moles by the smallest number of moles 4. If there are fractional subscripts convert these to whole numbers by multiplying through with the smallest integer that can convert the fractional subscripts to whole numbers. C. Combustion Analysis g CO 2mol CO m2l C g C % C g H 2 mol H O 2 mol H g H % H % O = 100 % - (% C + % H) VII. Balancing Chemical Equations A. Balancing Chemical Equations 1. Translate statement to a skeletal equation by putting down respective formulas on the reactant and product sides. 2. Balance the numbers of each type of atom on both sides of the equation 3. Adjust the coefficients so that all are whole numbers 4. Check if the equation is balanced by making sure the number and type of atoms are the same on both sides of the equation 5. Specify the states of matter for each entity in parenthesis after its chemical formula B. Points to remember while balancing equations 1. A coefficient (the number in front of a formula) operates on all atoms in the formula including the subscript. 2. An equation remains balanced when you multiply all the coefficients by the same factor 3. Chemical formulas cannot be altered (do not change subscripts on an element) 4. Other reactants or products cannot be added to balance the equation as this would alter the chemical reaction ATOMS AND ELEMENTS THE HISTORY OF THE PERIODIC TABLE I. History A. Dmitri Mendeleev, a Russian chemistry professor, found that when he listed elements in order of increasing mass, similar properties recurred in a periodic pattern. Mendeleev summarized these observations in the periodic law. 1. He organized the elements in a table consisting of a series of rows in which mass increases from left to right. Elements with similar properties fall in the same vertical columns. 2. Mendeleev’s table contained some gaps, which allowed him to predict the existence of some undiscovered elements, like silicon. 3. In the modern table, the elements are arranged by atomic number rather than atomic mass. Elements arranged according to increasing atomic number from left to right in boxes Boxes are arranged in a grid of periods 1-7 (horizontal rows) and vertical columns 1-8 (groups) with letter A or B. The eight A groups contain the main group elements, and the ten B groups contain the transition elements. The lanthanides and actinides fit in horizontal rows in between the transition metals The diagonal from B to At as well as Ge and Sb are metalloids. Elements to the left of these are metals and the elements to the right of these are non-metals Elements within a group have similar chemical properties and those across a period have different chemical properties Main group metals lose electrons to form cations so that they have the same number of electrons as the nearest noble gases while the main group non-metals gain electrons to form anions with the same number of electrons as the nearest noble gas. The Group 8A elements are called noble gases as they are chemically inert WEIGHING ATOMS BY COUNTING THEM I. The masses of atoms and percent abundances of isotopes of elements are measured by using mass spectrometry, a technique that separates particles according to their mass. II. A mole is a chemist’s dozen: a mole is the amount of material containing 6.022 x 10 23 23 A. 1 23le of anything= 6.022 x 10 (Ex- 1 mol of marbles= 6.022 x 10 marbles. 1. A mole (mol) of any substance is a quantity that contains an Avogadro’s number of items where Avogadro’s Number = 6.022 x 10 23 2. 6.022 x 10 23atoms of 12C = 1 mol of 12C = 12 g of 12C 3. Since atomic masses are relative masses, if a mol of C- 12 atoms weighs 12 g, a mol of O-16 weighs 16 g and 1 mol of Na atoms is 22.99 g etc. B. The mass of 1 mole of atoms of an element is the molar mass. An element’s molar mass in grams per mole is numerically equal to the element’s atomic mass in amu. 1. Ex- Copper has an atomic mass of 63.55 amu and a molar mass of 63.55 g/mol. One mole of copper atoms therefore has a mass of 63.55 g. 2. Number of Particles (atoms, molecules, ions) = Number of moles = Mass in g OBSERVATIONS THAT LED TO AN ATOMIC VIEW OF MATTER I. The Law of Conservation of Mass (Lavoisier 1743-1749) A. The total mass of substances remains the same during a chemical reaction. Mass is neither created nor destroyed during any chemical reaction. The mass of substances produced (products) by a chemical reaction is always equal to the mass of the reacting substances (reactants). Ex- Formation of water from its elements 2H 2 + O 2 2H O2 4 g + 32 g 36 g Metabolism of glucose C 6 12 6 + 6O 2 6CO +26H O 2 180 g + 192 g = 372 g 264 g + 108 g =372 g II. Law of Constant Composition/ Law of Definite Proportions (Proust 1754-1826) A. States that in a pure compound, the elements are always present in the same definite proportion by mass. (mass of element/mass of compound is fixed for any compound) Ex- H2O 2.0160g (2x1.0080) of hydrogen will combine with 15.9994g of oxygen to produce 18.0154g of H2O and 4.032 of hydrogen will combine with 31.988g of oxygen to produce 36.020g of H2O Ratio of H:O is always 2:16 or 1:8 Ex- % mass of an element in a compound is fixed. Example mass % of C in CO = 212 g/ 44g ) x 100 = 27 % mass % of O in CO =2(32 g/ 44g) x 100 = 73 % Total mass % = 100 % III. Law of Multiple Proportions (Dalton 1766-1844) A. When two elements (call them A and B) form two different compounds, the masses of B that combine with 1 g of A can be expressed as a ratio of small, whole numbers Ex- A = Carbon, C B = Oxygen, O Two compounds of C and O are CO an2 CO Ratio of the mass (or mass%) of O to the mass of C in the two compounds is 2:1 Mass % O in CO =272.7 % Mass % O in CO = 57.1 % Mass% C in CO =227.3 % Mass% C in CO = 42.9 % Mass of O/C in CO = 2.66 Mass of O/C in CO 2 = 1.33 DISCOVERY OF THE ATOM I. J.J Thompson and the Discovery of Electrons A. Thompson used a cathode ray tube with a magnet and discovered that the green beam it produced was made up of negatively charged material. He performed many experiments and found that the mass of one of these particles was almost 2,000 times lighter than a hydrogen atom. From this he decided that these particles must have come from somewhere within the atom and that Dalton was incorrect in stating that atoms cannot be divided into smaller pieces. Thomson went one step further and determined that these negatively charged electrons needed something positive to balance them out. So, he determined that they were surrounded by positively charged material. This became known as the 'plum pudding' model of the atom. The negatively charged plums were surrounded by positively charged pudding. 1. Key discoveries Charge to mass ratio (c/m) of the negatively charged particles The particles released were negatively charged II. Millikan’s Oil Drop Experiment A. When the x-rays hit an atom in the oil droplet, electrons were ejected which then makes the oil droplet positively charged. This then attracts other electrons. So each oil droplet could have one or more electrons making them negatively charged. Applying a voltage across the plates such that the force of attraction of the negative charged droplet to the positive plate is equal to the force of attraction due to gravity (qE = mg) allowed Millikan to calculate the charge on the droplets which were multiples of 1.6 x 10^-19 coulombs. He therefore, concluded that was the charge on one electron. III. Rutherford and the Nucleus A. Ernest Rutherford, one of Thomson's students, did some tests on Thomson's plum pudding model. He fired a beam of positively charged particles called alpha particles ( 2+¿ ) at a very thin sheet of He gold foil. Because these alpha particles had so much mass, he fully expected that all of the alpha particles would go right through the gold foil. This is because, if Thomson were correct about the plum pudding model of the atom, the alpha particles would just go through the positively charged matter and hit the detecting screen on the other side. However, when they hit the gold nuclei they were repelled almost 180 degrees because of the mass and charge of the gold nuclei. This led to the nuclear model of the atom. Dalton’s Atomic Theory Modern Atomic Theory All matter consists of extremely All matter is composed of small indivisible particles calledatoms that are divisible and atoms made up of smaller subatomic particles, electrons, protons and neutrons. Atom is still the smallest particle that retains the unique identity of the element Atoms of an element are All atoms of an element have identical in mass and other the same number of protons properties and are different and electrons, which determine from atoms of other elements the chemical behaviour of the element Compounds result from the Compounds are formed by the chemical combination of two or chemical combination of two or more atoms of different more elements in specific elements in a specific ratio. ratios. Atoms of one element cannot Atoms of one element cannot be converted into atoms of be converted to atoms of another element, atoms can another element in a chemical neither be created not reaction destroyed, only rearranged in a chemical reaction to produce CHEMICAL QUANITIES AND AQUEOUS RELATIONSHIPS I. Chemical Quantities A. Greenhouse gases allow sunlight to enter the atmosphere and warm Earth’s surface. But it also may prevent some of the heart generated by the sunlight from escaping. The balance between incoming and outgoing energy from the sun determines Earth’s average temperature. B. Reaction Stoichiometry: 1. The amount of carbon dioxide emitted by fossil fuel combustion is related to the amount burned. Balanced chemical equations give these relationships. The coefficients in a chemical reaction specify the relative amounts of moles in each of the substances in the reaction. The numerical relationships between chemical amounts in a balanced equation are called reaction stoichiometry. Multiply by the mole ratio to convert between moles. C. Limiting Reactant, Theoretical Yield, and Percent Yield 1. The limiting reactant limits the amount of product in a chemical reaction. It’s the reactant that makes the least amount of product, and is completely consumed in a chemical reaction. The reactants that do not limit the product are said to be in excess. The reactant in excess it the reactant that occurs in a quantity greater than is required. a. Finding limiting reactant: Given grams or moles, convert reactant amounts to product amount in moles and decide which reactant produces less. The outcome is the theoretical yield. b. Finding excess amount: after deciding the limiting reactant, convert the amount of limiting reactant into moles of excess reactant. Subtract this number from the original moles of excess reactant. 2. The theoretical yield is the amount of product that can be made in a chemical reaction based on the amount of the limiting reactant. The actual yield is the amount of product actually produced by a chemical reaction. 3. The percent yield is the percentage of the theoretical yield that was actually attained: actual yield % Yield = theoreticalyieldx 100 II. Solution Concentration and Solution Stoichiometry A. A homogenous mixture of two or more substances is called a solution. The majority component of the mixture is the solvent, and the minority is the solute. An aqueous solution is a solution in which water acts as the solvent. B. Solute Concentration 1. The amount of solute in a concentration is variable. You can add a little bit of solute to make a dilute solution, one that contains a small amount of solute relative to the solvent. Or you can add a lot of solute to make a concentrated solution, one that contains a large amount of solute relative to the solvent. 2. Solute concentration can be expressed in molarity, the amount of solute (in moles) divided by the volume of the solution (in liters). Molarity is a ratio of the amount of solute per liter of solution (mol/L). To make an aqueous solution of a specified molarity, you put the solute into a flask and add water until you have the desired volume. a. Ex- to make 1 L of a 1 M NaCl solution, add one mole of NaCl to a flask then ad water to make 1 L of solution. b. Preparing Molar solutions: use M= mol/ L c. Finding amount of solute in grams from M and V: use M and V to find moles, then convert to grams d. Molarity from mass percent and density: Assume 100 g, so turn percent into grams. Use D= g/mL to solve for volume in liters. Convert grams to moles. Use M= mol/ L to find molarity. 3. The mass of a solution is equal to the mass of the solute plus the mass of the solvent. 4. To dilute a solution, use the formula M1V1=M2V2. For example, a procedure calls for 3.00 L of a .500 M CaCal2 solution. To prepare the solution from a 10.0 M stock solution, use the formula to solve for the amount of liters of solution. a. Diluting molar solutions: Use formula M1V1=M2V2 b. mL solution from Molarity and grams: Convert grams to moles using equation, then use M= mol/ L to solve for volume. III. Types of Aqueous Solutions and Solubility A. Electrolyte and Non- Electrolyte Solutions: The difference in the way that a salt (an ionic compound) and sugar (a molecular compound) dissolve in water illustrates a fundamental difference between types of solution. A salt conducts electricity while a sugar does not. 1. Electrolytes are substances that dissolve in water to form solutions that conduct electricity. Strong electrolytes are substances that dissociate completely into ions when dissolved, such as NaCl. 2. Most molecular compounds – with the exception of acids- dissolve in water as intact molecules. Sugar dissolves because the attraction between sugar molecules and water molecules overcomes the attraction of sugar molecules to each other. Compounds like this that don’t dissociate into ions are called nonelectrolytes, and do not conduct electricity. B. Strong Acids 1. Hydrochloric acid is an example of a strong acid, one that completely ionizes in a solution. Strong acids are strong electrolytes. C. Weak Acids 1. Many acids are weak acids that do not completely ionize in water. Weak acids are weak electrolytes. a. Moles of ions in strong acids and bases: if ions dissociate, use molar subscripts for number of moles b. Molarity of ions in strong acids and bases: Using ionic equation, take given grams of compound and convert them to moles of the ion. Use M=mol/L to find the molarity. Multiply by appropriate subscript to find molarity of the ion. IV. The Solubility of Ionic Compounds Solubility Rules V. Reactions A. Precipitation Reactions 1. Precipitation reactions are ones in which a solid or precipitate forms upon mixing two solutions. They do not always occur when two aqueous solutions are mixed. They key is knowing only insoluble compounds form precipitates. The reactants can be soluble, but you must identify if the product is insoluble. If one of the possible products is insoluble, a precipitation reaction occurs. B. Molecular, Ionic, and Complete Ionic Equations A. Molecular equations are equations showing the complete neutral formulas for each compound in the reaction as if they existed as molecules. However in actual solutions of soluble ionic compounds dissociated substances are present as ions. B. Complete ionic equations list individually all of the ions present as either reactants or products in a chemical reaction. 1. To simplify the equation, spectator ions can be omitted. Spectator ions are the ones that appear unchanged on both sides of the reaction and don’t participate in the reaction. C. Acid- base and Gas Evolution Reactions 1. In an acid base reaction, or a neutralization reaction, an acid reacts with a base and the two neutralize each other, producing water, or in some cases a weak electrolyte. In a gas evolution reaction, a gas forms, resulting in bubbling. Many gas evolution reactions are also acid-base. Antacids acts as neutralizing agents. a. Acids are substances that produce H+ ions in an aqueous solution (Arrhenius definition). Protons associate with water molecules to produce hydronium ions. b. Bases are substances that produce OH- ions in aqueous solutions (Arrhenius definition) c. Some acids, called polyprotic acids contain more than one ionizable proton and release them sequentially. For example, H2SO4 is a diprotic acid, strong in its first ionizable proton but weak in its second. d. Likewise, some bases such as Sr(OH)2 produce two moles of OH- per mole of base. 2. When an acid and base are mixed, whether weak or strong, the H+ combines with the OH- from the base to form H2O. a. 3. Acid base reactions generally form water and an ionic compound called a salt, which usually remains dissolved in the solution. a. The net ionic equation for many acid- base −¿→H O 2 ¿ reactions is: +¿+O¿H H b. Another example is: H SO + 2KOH → 2 4 2 H2O + K 2 O 4 c. Acid+ Base = Salt + Water 4. Acid Base Titrations: a. Given Molarity and Liters, find moles. Make any necessary molar conversions to find the moles of the other reactant. Use moles and volume to find molarity. b. Use M1V1=M2V2 4. Gas Evolution Reactions a. Aqueous reactions that form a gas when two solutions are mixed. Some gas evolution reactions form a gaseous product directly when the cation of one reactant combines with the anion of the other. i. For example, when sulfuric acid reacts with lithium sulfide, di-hydrogen sulfide is formed. ii. Other gas evolution reactions form an intermediate product that then decomposes into a gas. For example, when aqueous hydrochloric acid is mixed with aqueous sodium bicarbonate, it forms the unstable intermediate carbonic acid, which breaks down to form water and carbon dioxide. iii. Other important gas- evolution reactions form H2SO3 or NH4OH as intermediates. D. Oxidation- Reduction Reactions 1. Oxidation- Reduction or redox reactions are reactions in which electrons are transferred from one reactant to another. The rusting of iron, the bleaching of hair, and the production of electricity in batteries involve redox reactions. Many redox reactions involve the reaction of a substance with oxygen. a. They do not always have to involve oxygen. For example, sodium and chlorine combine to make NaCl. Sodium also reacts with oxygen to form sodium oxide. In both cases, a metal with a tendency to lose electrons reacts with a nonmetal with a tendency to gain electrons. b. The fundamental definition of oxidation is the loss of electrons, a fundamental definition of reduction is the gain of electrons. c. The transfer of electrons need not be a complete transfer (as occurs in the formation of an ionic compound) for the reaction to qualify as a redox. For example, the reaction between hydrogen and chlorine gas to form HCl is a redox. Hydrogen loses some electron density and it has partially transferred its electron to chlorine. 2. The oxidation number is given to each atom based on the electron assignment. It is the “charge” it would have if all shared electrons were assigned to the atom with greater attraction for those electrons. 3. Oxidation numbers can be used to identify redox reactions. Something is oxidized if there is an increase in the oxidation number. Something is reduced if there is a decrease in the oxidation number. 4. The oxidizing agent oxidizes another substance and IS reduced. The reducing agent reduces another substance and IS oxidized. E. Combustion Reactions 1. Combustion reactions are a type of redox reaction by which we get most of our energy. They are characterized by the reaction of a substance with O2 to form one or more oxygen containing compounds, often including water. 2. These reactions emit heat. Usually the reacting carbon compound is oxidized and the oxygen is reduced. 1.3-1.4: Basis of Chemistry I. Chemistry is the science that seeks to understand the behavior of matter by studying the behavior of atoms and molecules. 1. Matter- anything that has mass and occupies space. 2. Atoms are submicroscopic particles that constitute the fundamental building blocks of matter. Free atoms are rare; instead they bind to form molecules. Classification of Matter: 1. Based on physical state Shape Volume Compressibilit Flow y Solid Fixed Fixed Non Vibrates Ex- ice, Compressible tightly in aluminum, place diamond Liquid Variable Fixed Non Flows Ex- water, Compressible alcohol, gasoline Gas Variable Variable Compressible Flows Quickly 3. Solid matter can be crystalline, like table salt and diamond, meaning its atoms or molecules are in patterns and with long repeating order. They may be amorphous too, like glass and plastic; the atoms will not have any order. 2. Based on Composition Pure Substances- Matter that Impure Substances (Mixtures) has a fixed composition and – Matter that has a variable distinct properties. composition and contains more than one pure substance. Substances are not chemically bound and can be separated by physical means. Elements- Compounds- Homogenous- Heterogeneous One type of Cannot be Saltwater- - atom or atoms; separated by H2O , NaCl, Seawater- Can be atomic physical means; sweetened tea H 2 , NaCl, or molecular Contains more MgCl , sand ( Ex- Copper than one 2 (Cu) element- Only iO2 ) Oxygen ( molecular O 2 ) Ex- CuO, Water CANNOT BE More than one CHEMICALLY type of atom in a fixed ratio BROKEN DOWN All metals are CAN BE BROKEN atomic. DOWN Diatomic elements are pure substances. Question 1: What is the difference between elements and compounds? Compounds can be broken down; elements cannot. Elements can be molecular and atomic whereas compounds can only be molecular. II. Separation of Mixtures: Physical methods of separation Evaporation- heating the mixture to separate two substances that dissolve into each other (ex- separate salt water) Filtration- used to separate two substances that don't dissolve into each other; mixture is poured through filter paper in a funnel (ex- sand and water) Distillation- used to separate two liquids that boil at two different temps; mixture is heated to boil off the more volatile liquid, which is then condensed and collected. Decantation- physically separating two substances that don't dissolve into each other (ex- pour water out of a beaker while sand is settled at bottom) A. Changes 1. Physical Change – When a substance alters its physical form without changing its Composition. Ex- boiling, melting, freezing, sublimating, dissolving a. When kinetic energy is great enough to overcome attractions binding substance, a physical change will occur. Example: H O(s) → H O(l2 2 2. Chemical Change – The composition of the substance is altered and it becomes a different substance. a. Ex- rusting, burning, colors or dye fading, etc. Example: Formation of Rust 4Fe(s) + 3O (g)2→ 2Fe O (s)2 3 B. Properties 1. Physical properties: odor, taste, color, appearances, melting point, boiling point, density 2. Chemical properties: corrosiveness, flammability, acidity, and toxicity C. How do you identify matter? Properties of Matter 1. Physical Properties (qualitative) – Color, Odor, Hardness, Solubility Electrical Conductivity 2. Physical Properties (quantitative)- Density, Melting point, Boiling point 3. Chemical Properties – Flammability, Corrosiveness Extend your thinking (EYT1): What are stalagmites or stalactites made up of? Would you consider them to be compounds or mixtures? http://www.goodearthgraphics.com/virtcave/staltite/staltite.html III. 1.5: ENERGY & MATTER 1. Changes in matter accompany changes in energy. A. Changes in matter, both physical and chemical, result in the matter either gaining or releasing energy B. Energy is the capacity to do work C. Work is the action of a force applied across a distance a. A force is a push or a pull on an object b. Electrostatic force is the push or pull on objects that have an electrical charge 2. High energy= less stable A. Processes resulting in higher energy take more time and energy B. Forms of energy: a. Kinetic (energy of motion) ex- thermal heat i. Chemical Kinetic Energy – Motion of particles in a substance b. Potential (stored energy of position possessed by an object) ex- chemical potential energy from attractive forces between atoms i. Chemical Potential Energy – Results from positions and interactions of particles in a substance C. Ball rolling down hill: Kinetic energy to Potential energy D. Total Energy = Kinetic Energy + Potential Energy 3. The Law of Conservation of Energy states whatever you do to convert between energies, the total amount of energy remains the same. If a process results in the system having less potential energy at the end than it had at the beginning, the “lost” potential energy was converted into kinetic energy, which is released to the environment 4. Burning of gasoline: 2C H8(18 + 25O (g) 2 16CO (g) +12 H O (l) + 2nergy IV. 1.2: THE SCIENTIFIC APPROACH Observations- When observations can be consistently repeated, they become laws Ex- Cars develop rust faster in coastal areas Hypothesis – Ex- Atmospheric Conditions in the coastal area hasten rusting. Experiment: Ex- Measure the time in which rust develops on iron nails under various conditions, such as presence of water, presence of salt and water, presence of salt alone etc. 1. Theory- Explains why it happens. Ex- Salt and humidity hasten the rusting process Applies to single or Applies to all small number of events events 2. The Describes what Observation Law scientific happens approach is Explains why things Hypothesis Theory in fact the happen exact opposite of Plato’s “sensible” learning. Scientific knowledge is empirical (based on observation and experiment). Some observations are qualitative, noting or describing a process, or quantitative, measuring or quantifying something about the process. A. Antoine Lavoisier made the observation that combustion does not change the mass of objects. B. Observation leads to hypothesis. A good hypothesis is falsifiable; predictions can be refuted or proved by observations. Hypotheses are tested by experiments. Some observations lead to the development of a scientific law. a. Ex- Lavoisier’s Law of Conservation of Mass- in a chemical reaction, matter is neither created nor destroyed. Laws are also subject to experiments. C. One or more well established hypotheses might form the basis for a scientific theory. A scientific theory is a model for the way nature is and why. For example, John Dalton proposed the Atomic Theory, which states that matter is composed of small indestructible particles called atoms. Since these particles are rearranged, the total mass is the same. The Atomic Theory is a model for the physical world, and explains our laws and observations. 1.6: UNITS OF MEASUREMENT http://www.straightdope.com/columns/read/637/whats-the-origin-of-miles-and-yards http://en.wikipedia.org/wiki/Ancient_Roman_units_of_measurement#Weight System of Unit Mass Length English lb in Metric kg cm SI kg m I. Scientific Measurement – Expressed in SI units ; 7 fundamental units for the fundamental quantities of mass kg length m time sec temperature K electric current Amp amount of substance mol luminous intensity candela A. Heat flows from the matter that has high thermal energy into matter that has low thermal energy until they reach the same temperature i. heat flows from hot object to cold ii. heat is exchanged through molecular collisions between the two materials
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