CHEM210 PSU Oxidation Chemical Reaction Concept Research Paper

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Chapter 2. Elements and Compounds 2.1 The Structure of the Atom 2.2 Elements and the Periodic Table 2.3 Covalent Compounds 2.4 Ions and Ionic Compounds In this chapter, we discover that there are actually a number of different variations, or isotopes, of the atoms associated with each element. We explore the structure of the atom in further detail and learn about the composition of isotopes. We also discuss the different ways that molecules and compounds are represented and named. 2.1 The Structure of the Atom 2 Atomic Structure: Protons and Neutrons Mass of the atom is primarily in the nucleus. Charge of the proton is opposite in sign but equal to that of the electron. Scanning Tunneling Microscopy (STM) Scanning Tunneling Microscopy •Fe atoms arranged on Cu. Atom in Chinese Fe atoms arranged on Cu. Quantum Corral Ag atoms arranged on Si. Atomic Masses and the Mole The mass of 1 atom of carbon-12 is defined to be 12 amu. vb = #eZ = #p+ also vb A = #p+ + #n 2.1 Atomic Number, Mass Number, and Isotopes Most elements have two or more isotopes, -atoms that have the same atomic number (Z) but different mass numbers (A). 7 SAMPLE PROBLEM Determine the numbers of protons, neutrons, and electrons in each of the following species: Number of protons = Z number of neutrons = A – Z number of electrons = number of protons. Solution Atomic number is Z = 17 17 protons. Mass number is A = 35, neutrons is 35 – 17 = 18. Number of electrons = number of protons = 17 electrons. 8 2.1 The Atomic Mass Scale and Average Atomic Mass Atomic mass unit (amu) is defined as a mass exactly equal to one-twelfth the mass of one carbon-12 atom. 9 Percent Abundance • For each element’s isotopes, the percent abundance is found by: number of atoms of an individual isotope Percent abundance = total number of atoms of ×100 all isotopes of that element 10 Atomic Weight (Mass) • Atomic weight: Equals the average mass of all naturally occurring isotopes of the element – Accounts for relative abundance of isotopes average atomic weight =  ( exact mass )( fractional abundance ) all isotopes atomic mass = ( %abundance of isotope 1) × mass of isotope 1 100 %abundance of isotope 2 ) ( + × mass of isotope 2... 100 11 Isotope Abundance Element Symbol At. Wt. Mass number Isotope mass Natural abundance (%) Hydrogen H 1.00794 1 2 3 1.0078 2.1041 3.0161 99.85 0.015 0 Boron B 10.811 10 11 10.0129 11.0093 19.91 80.09 Magnesium Mg 24.305 24 25 26 23.985042 24.985837 25.982593 78.99 10.00 11.01 Gallium 69.723 69 71 68.926 70.925 60.40 39.60 Ga 1 SAMPLE PROBLEM 2.3 Oxygen is the most abundant element in both Earth’s crust and the human body. The atomic masses of its three stable isotopes, 168O (99.757 percent), 178O (0.038 percent), and 188O (0.205 percent), are 15.9949, 16.9991, and 17.9992 amu, respectively. Calculate the average atomic mass of oxygen using the relative abundances given in parentheses. Report the result to four significant figures. What should the answer be? What can you check to estimate the answer? 13 SAMPLE PROBLEM 2.3 Solution Atomic mass = Σ(fractional abundance)(isotope mass) 16 8O 17 (99.757 %) 18 8O 8O (0.038 %) (0.205 %) Check your answer, what is the value on the periodic table? 14 2.2 The Periodic Table Horizontal rows - periods Vertical columns - groups (families). 15 2.3 Molecules and Molecular Compounds Molecule is a combination of at least two non-metal atoms in a specific arrangement and ratio held together by electrostatic forces known as covalent chemical bonds. Homonuclear diatomic molecules, both atoms in each molecule are of the same element. A diatomic molecule can also contain atoms of different elements (Heteronuclear diatomic molecules). 7 elements exist as diatomic molecules are hydrogen (H2), nitrogen (N2), oxygen (O2), fluorine (F2), chlorine (Cl2), bromine (Br2), and iodine (I2). LEARN the 7 Diatomics! 16 2.3 Molecules and Molecular Compounds Chemical Formula denotes the composition of the substance, metal-nonmetal, non-metal-non-metal, or metal-metal. Molecular Formula shows the exact number of atoms of each element in a molecule of non-metal atoms. Allotrope is one of two or more distinct forms of an element. Two of the allotropic forms of the element carbon—diamond and graphite—have dramatically different properties. 17 2.3 Molecules and Molecular Compounds Condensed structural formula Bonding formula Structural formula Ball and Stick model Space-Filling model 18 2.3 Molecules and Molecular Compounds Molecular substances can also be represented using empirical formulas. empirical means “from experience” in the context of chemical formulas, “from experiment.” Empirical formula tells what elements are present in a molecule and in smallest whole-number ratio 19 2.3 Molecules and Molecular Compounds 20 2.3 Molecules and Molecular Compounds Binary compounds consist of just two different elements. 1st name the element that appears first in the formula (usually the metal atom). Name the second element (usually the nonmetal), changing the ending of its name to –ide. metal non-metal HCl hydrogen chloride NO nitrogen oxide SiC silicon carbide 21 Prefixes used in Naming Binary Nonmetal Compounds (Molecular Compounds) ONLY Number Prefix 1 Name mono Formula 2 di Name Formula Sulfur dioxide tri H2O2 4 Hydrogen peroxidetetra SO2 Sulfur trioxide SO3 5 Ammonia Carbon monoxide CO Carbon dioxide CO2 Chlorine monoxide ClO Water 3 6 7 Hydrazine 8 Nitric oxide H2O penta NH3 hexa N2H4 hepta octaNO 9 Nitrogen dioxide nona NO2 10 deca 12 Disulfur decafluoride S2F10 dodeca Prefix mono– is generally omitted for the first element except carbon monoxide (CO). 22 2.3 Oxoanions Two forms exist: –ate and –ite suffix endings are used. Increase # O ▪ More oxygen = “-ate” ▪ Less oxygen = “-ite” SO42- sulfate ion NO3- nitrate ion SO32- sulfite ion NO2- nitrite ion If they contain H, add a prefix “hydrogen” HSO4- hydrogen sulfate ion (common name=bisulfate ion) HCO3- hydrogen carbonate ion (common name=bicarbonate ion) 2.3 Oxoanions When four forms exist, add prefix Increase O –Add “per_____ate” and “hypo____ite” names FO4FO3FO2FO- perfluorate fluorate fluorite hypofluorite ClO4ClO3ClO2ClO- perchlorate chlorate chlorite hypochlorite 2.3 Ions and Ionic Compounds Ions that consist of a combination of two or more atoms are called polyatomic ions. LEARN THEM! Quiz in Lab! 25 2.3 Molecules and Molecular Compounds Acid is a substance that produces hydrogen ions (H+) when dissolved in water. Rules for naming simple acids of this type are as follows: 1. remove the –gen ending from hydrogen (leaving hydro–) 2. change the –ide ending on the second element to –ic 3. combine the two words 4. add the word acid. 26 2.3 Molecules and Molecular Compounds Binary compounds consist of just two different elements. 1st name the element that appears first in the formula (usually the metal atom). Name the second element (usually the nonmetal), changing the ending of its name to –ide. metal non-metal HCl hydrogen chloride SiC silicon carbide NaBr sodium bromide 27 Interactive Figure 2.4.2 - Explore Monoatomic ion Formulas Transition Metal have different oxidation states (charges) Charge of cation is determined by the non-metal anion paired with 28 2.4 Ions and Ionic Compounds Cations that can have more than 1 charge (oxidation state) are use Roman numerals. Cations with only two possible charges – use ending suffix – ous to the cation with the smaller positive charge – ic to the cation with the greater positive charge: Cu1+ : Cuprous Cu2+ : Cupric ▪ exceptions: lead (Pb2+, Pb4+), tin (Sn2+, Sn4+)… ▪ exceptions: silver (Ag+), zinc (Zn2+), cadmium (Cd2+) 29 2.4 Naming Ionic Compounds Ionic Compound: A neutral compound in which the total number of positive charges must equal the total number of negative charges. Binary Ionic Compounds aluminum sulfide: Al3+ +6 Al23+ S2+ -6 S32- =0 =0 Al2S3 2.4 Ions and Ionic Compounds Formulas of Ionic Compounds Aluminum Oxide Calcium Phosphate Check to make sure charges cancel out and equal 0! 31 2.4 Ions and Ionic Compounds Naming Ionic Compounds 32 2.4 Naming Ionic Compounds Metal of fixed charge with a polyatomic ion Cation Anion Formula Name K+ OH− KOH potassium hydroxide Ca2+ OH− Ca(OH)2 calcium hydroxide Al3+ SO24- Al2(SO4)3 aluminum sulfate Pb(SO4)2 lead(IV) sulfate Fe(NO2)2 iron(II) nitrite ferrus nitrite 24 Pb4+ SO Fe2+ NO2- 33 2.4 Ions and Ionic Compounds Familiar Inorganic Compounds 34 CHAPTER 3.1 3.2 3.3 3.4 3.5 Chapter 3. Stoichiometry The Mole and Molar Mass Stoichiometry and Compound Formulas Stoichiometry and Chemical Reactions Stoichiometry and Limiting Reactants Chemical Analysis In this chapter we will explore the chemical counting unit that links the atomic and macroscopic scales, the mole. The mole will allow us to study in greater detail chemical formulas and chemical reactions. Specifically, we will investigate stoichiometry, the relationship between quantities of materials in chemical reactions. 3.1 The Mole and Molar Masses The Mole People use a variety of counting groups to conveniently indicate the number of objects in some set: 1 “pair” objects 2 objects 1 “dozen” objects 12 objects 1 “gross” objects 144 objects 1 “million” objects 1,000,000 objects 1 “trillion” objects 1012 objects 1 “mole” objects 6.023 × 1023 objects 2 3.1 The Mole and Molar Masses The Mole A “mole” is a counting group, defined as the number of atoms in exactly 12 g of carbon-12, number of atoms in 12 g of carbon-12 is known as Avogadro’s Number (NA): NA = 6.0221418 × 1023 objects 1 mol objects 6.022  1023 objects 6.022  10 objects 1 mol objects 23 If you order one dozen doughnut, you are asking for 12 doughnuts. If you order one mole of doughnuts, you are asking for 6.022 × 1023 doughnuts! Fun Fact: “Avogadro’s Number” is named for Amedeo Avogadro, 1776-1856 3 Interactive Figure 3.1.1 - Recognize How the Mole Connects Macroscopic and Atomic Scales One mole quantities of (from left to right) the elements copper, aluminum, sulfur and the compounds potassium dichromate, water and copper (II) chloride dihydrate 4 SAMPLE PROBLEM 3.5 A typical human body contains roughly 30 moles of calcium. Determine (a) the number of Ca atoms in 30.00 moles of calcium and (b) the number of moles of calcium in a sample containing 1.00 × 1020 Ca atoms. Setup 1 mol Ca atoms 6.022  1023 Ca atoms 6.022  10 Ca atoms 23 1 mol Ca atoms 5 3.1 Molar Mass • Molar mass (MM): Molecular weight in grams of one mole of a substance • Molecular weight – Formula weight for substances that exist as individual molecules • Formula weight – Sum of the atomic weights of the elements that make up a substance multiplied by the number of atoms of each element in the formula 6 3.1 The Mole and Molar Masses Interconverting Mass, Moles, and Numbers of Particles 7 3.1 The Mole and Molar Masses Formula MM, Molar mass (g/mol) m, Mass of sample (g) SO2 n, Moles of NA, Number of Atoms, Sample (mol) Molecules or Formula Units 4.25 2.65 x 1024 molecules PCl3 OCl2 0.0186 6.38 x 1022 molecules NaBr Calculate the mass of carbon contained in 3.45 x 1021 molecules of aspirin, C9H8O4. [MM C9H8O4 = 180.158] 21 3.45 x 10 molecules x 1 mol 9 mol C x 12.01 1g = 0.619 g C x 23 1 mol 1 mol C 6.022 x 10 molecules atoms 8 3.2 Element Composition • Stoichiometry: Study of the relationship between relative amounts of substances – Formula of a compound provides information on the relative amount of each element in either one molecule or one mole of the compound • One molecule of acetic acid, CH3CO2H, contains: – Two atoms of oxygen and one mole of acetic acid – 2 mol of oxygen atoms 9 3.2 Percent Composition of Compounds (2 * Composition 1.008 amu) Percent - Mass of an element present in exactly 100 g of a compound composition by mass. % element = (number of atoms of element in formula) (molar mass of element) x 100% mass of 1 mol of compound Formula mass is: (2 * 1.008 (2 * 1.008 amu) amu) + (2 * 16.00 amu) = (2 * 1.008 34.02 amu amu) (2 * 1.008 amu) (2 * 1.008 amu) 10 3.2 Empirical Formula • Whole-number ratio of elements present in the compound • Determined by converting mass of each element to an amount in moles • Reverse of the process that is used to determine percent composition 11 3.2 Molecular Formula (2 * 1.008 formula amu) • Molecular is a multiple scaled by a factor “n”, the molecular and empirical molar masses must scale by the same ratio  g molar formula mass  mol =n g empirical formula mass mol   empirical formula molecular formula n x C XH Y O Z = CnXHnYOnZ n = 2,3,4… 12 The molecular formula can be determined if the molecular mass is known and vice versa. Example: Vitamin C has the empirical formula C3H4O3 and molecular mass = 176.12 g/mol. Empirical mass: 3(12.01 g/mol) + 4(1.008 g/mol) + 3(15.99 g/mol) = 88.03 g/mol C3H4O3 Multiple repeat units = Molecular mass Empirical formula mass = 176.12 g/mol 88.03 g/mol Molecular formula = 2(C3H4O3) = C6H8O6 =2 3.2 Percent Composition & Empirical Formulas Check your WORK! What is Empirical Formula Mass? How many Repeat Units? What is the Molecular Formula Mass? Do the masses and the repeat units make sense? 14 Hydrated Compounds • Hydrated ionic compound: Ionic compound that has a well-defined amount of water trapped within the crystalline solid • Water of hydration: Water associated with the compound • Hydrated compound formula includes the term nH2O – Where n is the number of moles of water incorporated into the solid per mole of ionic compound 15 3.3 Chemical Equations Interpreting and Writing Chemical Equations Chemical Reaction, as described in the third hypothesis of Dalton’s atomic theory, is the rearrangement of atoms in a sample of matter. Examples include the rusting of iron and the explosive combination of hydrogen and oxygen gases to produce water. Chemical Equation uses chemical symbols to denote what occurs in a chemical reaction. 16 3.3 Chemical Equations Chemical Equations = Chemical Sentence “Ammonia and hydrogen chloride react to produce ammonium chloride.” “Calcium carbonate reacts to produce calcium oxide and carbon dioxide.” 17 3.3 Chemical Equations Each chemical species that appears to the left of the arrow is called a reactant. Reactants are those substances that are consumed in the course of a chemical reaction. Each species that appears to the right of the arrow is called a product. Products are those substances that are produced in the course of a chemical reaction Physical state of compounds are indicated using symbols 18 3.3 The Atomic Theory Dalton’s Atomic Theory 1. Elements are composed of extremely small particles called atoms. All atoms of a given element are identical, having the same size, mass, and chemical properties. The atoms of one element are different from the atoms of all other elements. 1. Compounds are composed of atoms of more than one element. In any given compound, the same types of atoms are always present in the same relative numbers. 2. Chemical Reaction rearranges atoms in chemical compounds; it does not create or destroy them. 19 3.3 Law of Conservation of Mass Hg(NO3)2(aq) + 2KI(aq) 3.25 g 3.32 g 3.25 g + 3.32 g = 6.57 g HgI2(s) + 2KNO3(aq) 4.55 g 2.02 g 4.55 g + 2.02 g = 6.57 g Balancing Chemical Equations • Relative amounts of the reactants and products involved in the reaction must be determined by balancing the equation 21 Balancing Chemical Equations 1. Write an unbalanced equation with correct formulas for all substances and states (). 2. Balance the atoms of one element. a. Start with the most complex molecule (keep polyatomics together) b. Change the coefficients in front of the molecules c. DO NOT alter the chemical formulas (i.e. do not change the subscripts!!!!) 3. Balance the remaining elements. 4. Check the atoms are all balanced. 3.3 Balancing Chemical Equations Balance : Al + Fe2O3 step 1 Al Al2O3 + Fe + Fe2O3 1 Al Al2O3 + Fe (2Fe + 3O) (2Al + 3O) 1Fe not balanced not balanced step 2 Al 1 Al + Fe2O3 Al2O3 + 2 Fe (2Fe + 3O) (2Al + 3O) balanced not balanced 2Fe step 3 2 Al + Fe2O3 2Al (2Fe + 3O) Al2O3 + 2 Fe (2Al + 3O) 2Fe step 4 – balanced 3.3 Balancing Chemical Equations Polyatomic ion on both sides of an equation? Balance Polyatomic ion as a collective species. NaNO3(s) + H2SO4(aq) Na + NO3 + 2H + SO4 Na2SO4(aq) + HNO3(aq) 2Na + SO4 + H + NO3 not balanced Balance Na in Na2SO4 2 NaNO3(s) + H2SO4(aq) 2Na + 2NO3 + 2H + SO4 2Na + SO4 + H + NO3 not balanced 2 NaNO3(s) + H2SO4(aq) 2Na + 2NO3 + 2H + SO4 Na2SO4(aq) + HNO3(aq) Na2SO4(aq) + 2 HNO3(aq) 2Na + SO4 + 2H + 2NO3 balanced 3.3 Chemical Arithmetic: Stoichiometry Stoichiometry: Relative proportions in which elements form compounds or in which substances react. aA + bB Grams of A Molar Mass of A Moles of A cC + dD Moles of B Mole Ratio Between A and B (Coefficients) Grams of B Molar Mass of B Stoichiometric Factor or Ratio 2A + 3B 2 As combine with 3Bs A2B3 3B 2A or Conversion factors 3B 2A 26 Stepwise method for solving, need to determine what is being asked and what you know. 1. Write correct formulas for reactants and products 2. BALANCE the chemical equation! 3. Decide what is known and what is unknown from the question • write it down under the chemical species 4. Map out a strategy for answering the question • What insights/hints are in the question • What are you solving for? 5. Convert everything to moles! 6. Use mole-mole stoichiometric ratio to convert from 1 chemical species to another chemical species 7. Check the answer! Make sure it is reasonable, think about the answer, does it Make Sense? Aqueous solutions of sodium hypochlorite (NaOCl), best known as household bleach, are prepared by reaction of sodium hydroxide with chlorine gas. How many grams of NaOH are needed to react with 25.0 g Cl2? 2 NaOH(aq) + Cl2(g) ____ g NaOCl(aq) + NaCl(aq) + H2O(l) 25.0 g Grams of Cl2 Molar Mass Moles of Cl2 Mole Ratio Moles of NaOH Grams of NaOH Molar Mass Aqueous solutions of sodium hypochlorite (NaOCl), best known as household bleach, are prepared by reaction of sodium hydroxide with chlorine gas. How many grams of NaOH are needed to react with 25.0 g Cl2? 2NaOH(aq) + Cl2(g) 25.0 g Cl2 x 1 mol Cl2 70.9 g Cl2 NaOCl(aq) + NaCl(aq) + H2O(l) 2 mol NaOH x 1 mol Cl2 = 28.2 g NaOH Significant figures!!! 40.0 g NaOH x 1 mol NaOH Sec 3.3 Exercise: Strategy Map Data Information: Mass of reactant Step 1: Write the balanced chemical equation Eq. gives mole ratios (stoichiometry) Step 2: Convert mass to moles amount of reactant in moles Step 3: Convert moles reactant to moles product amount of products in moles Step 4: Convert moles of products to mass mass of products in grams 30 3.4 Limiting Reactants The reactant used up first is called the limiting reactant, , the amount of this reactant limits the amount of product that can form. When all the limiting reactant has been consumed, NO MORE PRODUCT can be formed! Excess reactants are those present in quantities greater than necessary to react with the quantity of the limiting reactant. Initial amount – amount Used = amount Remaining (excess) 31 3.4 Reactions with Reactant in Limited Supply Given 10 slices of cheese and 14 slices of bread and a sandwich is 1 cheese and 2 breads, how many sandwiches can you make? Balanced equation 1 cheese + 2 bread 1 sandwich 1 cheese ≡ 2 bread 1 cheese ≡ 1 ...
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Article Writing
Institutional Affiliation
What is Oxidation? In the chemical reactions chemistry concept, the two major kinds of
biochemical response of reactions are reduction and oxidation. Oxidation is indeed not
associated with oxygen. Oxidation is the occurrence where there is electrons loss by ion,
molecule or an atom at the time of a reaction. Oxidation takes place when a molecule, ion or
atom oxidation condition goes up. The adverse process is known as reduction. An ancient
definition of oxidation was during oxygen addition to a compound as oxygen gas was initially
known as an oxidation fac...

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