Chemistry Lecture: Fundamentals and Sample Problems PDF

CHEMISTRY ENGR. ZOREN P. MABUNGA, ECE, ECT, MSc.  Introduction Chemistry is the branch if science that deals with the study of matter. It also deals with the processes and the changes that matter undergoes.  Branches of Chemistry 1. Organic Chemistry – i...

CHEMISTRY ENGR. ZOREN P. MABUNGA, ECE, ECT, MSc.  Introduction Chemistry is the branch if science that deals with the study of matter. It also deals with the processes and the changes that matter undergoes.  Branches of Chemistry 1. Organic Chemistry – it deals with the study of chemical substances that contains carbon-carbon bonds. Examples: plastics, pharmaceuticals 2. Inorganic Chemistry - it deals with the study of chemical substances that do not contain carbon-carbon bonds. Examples: metals, minerals, semiconductors 3. Physical Chemistry – it deals with the behavior and changes in matter and the related energy changes. Examples: reaction rates, reaction mechanism 4. Analytical Chemistry – deals with the study of different components and composition of substances. Examples: Food nutrients 5. Biochemistry – Deals with the study of matter and processes in living organisms.  SIGNIFICANT FIGURES Rules for Counting Significant Figures 1. Nonzero integers are always significant 2. Zeros Leading zeros never count as significant figures. Captive (embedded) zeros are always significant. Trailing zeros are significant if the number has a decimal point. Examples: 2.5000 – 5 significant figures 1400 – 2 significant figures 0.005201 – 4 significant figures 1.500 x 10-3 – 4 significant figures  Addition and Subtraction with Significant Figures Result is limited by the number with the smallest number of significant decimal places Examples: 13.214 + 234.6 + 7.0350 + 6.38 = 1247 + 134.4 + 450 + 78 = 13.7 + 1.3 = 15600 + 172.49 =  Multiplication/Division with Significant Figures Result has the same number of significant figures as the measurement with the smallest number of significant figures Round the result so it has the same number of significant figures as the measurement with the smallest number of significant figures Examples: 2.5 x 3.42 = 3.10 x 4.520 = (4.52 x 10^-4) ÷ (3.980 x 10^-6) =  Matter Matter is anything that has mass and occupies space.  States of Matter 1.Solid 2.Liquid 3.Gas Two Special States: 1.Plasma – similar to gases except that it is made up of free electrons and ions from the element 2.Bose –Einstein Condensate (BEC) – newest state of matter as it was discovered only in 1995 by Cornell and Weiman Phase Changes  Properties of Matter 1. Physical Properties – can be measured without changing the composition of the original substance 2. Chemical Properties – can be observed or measured after changing chemically the substance under test. Physical Properties Chemical Properties Mass Conductivity Color Malleability Odor Reactivity Size Flammability Freezing Point Toxicity Boiling Point Radioactivity Density Oxidation  Classification of Measurable Properties 1.Extensive Properties – depend on the amount 2.Intensive Properties – independent on the amount Extensive Properties Intensive Properties Mass Boiling Point Volume Freezing Point Length Density Weight Temperature Area Pressure Width Melting Point Energy Color  Fundamental Chemical Laws The Law of Conservation of Mass – “Mass can neither be created nor destroyed.” The Law of Definite Proportion – “A given compound always contains exactly the same proportion of elements by mass.” The Law of Multiple Proportions – “When elements A and B form a series of compounds, the ratio of masses of B that combine with 1 gram of A can be reduced to small whole numbers.”  The Atom Atom is the basic building block of matter. It is the smallest particle of an element.  Subatomic Particles Particle Charge in Unit Mass in Location Coulomb kg Proton +1.6 x 10-19 1.6726 x 10-27 Nucleus Neutron 0 1.6749 x 10-27 Nucleus Electron -1.6 x 10-19 9.1094 x 10-31 Outside  The Atomic Models 1.The Thompson Model 2.The Rutherford Model  The Atomic Models 3. The Bohr Model  Atomic Number and Atomic Mass The atomic number is just the number of protons inside the nucleus of an element. The atomic mass/mass number/atomic weight is the sum of the number of protons and neutrons found inside the nucleus of an atom. Atomic Mass = No. of Protons + No. of Neutrons Isotopes are atoms with the same number of protons but that have a different number of neutrons.  Atomic Weight Atomic Weight is the average of an element, calculated using the relative abundance of naturally occurring isotopes. 𝐴𝑡𝑜𝑚𝑖𝑐 𝑊𝑒𝑖𝑔ℎ𝑡 = 𝑚1 𝑝1 + 𝑚2 𝑝2 + 𝑚3 𝑝3 + ⋯ Where: mn = mass of isotopes 1, 2, 3…. pn = percent abundance of isotopes 1,2,3, …  Mole Concept and Molar Mass Mole is the amount of pure substance containing the same number of chemical units, as there are atoms in exactly 12 grams of carbon-12. Avogadro’s Number refers to the number of particles in a mole of a substance. (6.02 x 1023 particles/mole)  Mole – Mass Conversions The formula for calculating among mass, gram-formula mass (also known as molar mass), and the number of moles: 𝑚 𝑛= 𝑀 Where: n = number of moles m = mass of the substance in grams M = molar mass in grams per mole Designation of Molecules and Compounds 1. Molecular Formula – shows the exact number of atoms from each element that is present in the molecule or compound. 2. Empirical Formula – simplified ratio of the atoms from each element that is present in the molecule or compound.  Steps in Finding the Empirical Formula 1. Start with the number of grams of each element. If percentage are given, assume that total mass of the substance to be 100 grams. 2. Convert the mass of each element to moles. 3. Divide each mole value by the smallest number of moles calculated. 4. Round-off the values to the nearest whole number to get mole ratio. This is represented as subscripts in the empirical formula.  Steps in Finding the Molecular Formula 1. Determine the empirical formula. 2. Calculate the mass of empirical unit. 3. Divide the molecular mass by the mass of empirical formula. 4. Multiply the subscripts of the empirical formula by the calculated ratio to find the molecular formula. SAMPLE PROBLEMS/PAST BOARD EXAM QUESTIONS 1.Calculate the moles of Magnesium (Mg) present in 93.5 g of Mg? (Mg atomic mass = 24.31g/mol) a.3.85 moles b.4.15 moles c.5.38 moles d.3.35 moles SAMPLE PROBLEMS/PAST BOARD EXAM QUESTIONS 2.How many number of atoms are there in 1.32 x 103 g of Lead (Pb)? Pb atomic mass is 207.2 g/mole. a.3.84 x 1023 atoms b.4.38 x 1023 atoms c.3.84 x 1024 atoms d.4.38 x 1024 atoms SAMPLE PROBLEMS/PAST BOARD EXAM QUESTIONS 3. How many grams are there in 4.57 x 1021 amu? a.6.95 x 10-3 g b.5.45 x 10-3 g c.5.96 x 10-3 g d.7.59 x 10-3 g SAMPLE PROBLEMS/PAST BOARD EXAM QUESTIONS 4. How many moles of chloroform (CHCl3) are there in 210.45 g of chloroform? C = 12.01 amu, H = 1.008 amu, Cl = 35.45 amu a.2.51 moles b.4.34 moles c.2.12 moles d.1.76 moles SAMPLE PROBLEMS/PAST BOARD EXAM QUESTIONS 5. How many moles of sodium chlorate (NaClO3) are in 284 grams of 12.0% sodium chlorate solution? a. 0.52 moles b. 2.64 moles c. 1.17 moles d. 0.32 moles SAMPLE PROBLEMS/PAST BOARD EXAM QUESTIONS 6. How many protons (P) and neutrons are there in the nucleus are present in a Pb nucleus of atomic mass of 206? a. P = 92, N = 156 b. P = 85, N = 160 c. P = 82, N = 124 d. P = 90, N = 150 SAMPLE PROBLEMS/PAST BOARD EXAM QUESTIONS 7. How many moles are there in one atom? a. 3.6 x 10-23 b. 1.66 x 10-24 c. 2.86 x 10-4 d. 6.86 x 10-9 SAMPLE PROBLEMS/PAST BOARD EXAM QUESTIONS 8. What is the molecular formula of a sample of a compound containing 6.444 g of boron (B) and 1.803 g of hydrogen (H)? The compound has a molar mass of about 30 g. Given boron has 10.81 amu and hydrogen has 1.008 amu. a. B2H3 b. BH c. B2H6 d. BH3  Chemical Bonding 1. Ionic Bonding Formed when one atom accepts or donates one or more of its valence electrons to another atom. It involves a metal and a non-metal 2. Covalent Bonding Formed when two atoms share valence electrons between them. It involves two or more non-metals. 3. Metallic Bonding Force of attraction that exists within elemental metals.  Chemical Reactions Chemical Reaction is a process in which a substance or a combination of substances undergoes a change in appearance or properties and further transform into a different substance or a combination of new substances.  Classification of Chemical Reactions 1. Direct Combination or Synthesis involves the combination of two or more reactants to form one product. The reactants can be element or compound. 𝐴 + 𝐵 → 𝐴𝐵 2. Decomposition reaction involves the breakdown of a single reactant into two or more products. 𝐴𝐵 → 𝐴 + 𝐵 3. Single-Replacement Reactions occurs when an uncombined element replaces another element that is part of a compound. As a result, the replaced element becomes uncombined. 𝐴 + 𝐵𝐶 → 𝐴𝐶 + 𝐵 4. Double-Replacement Reactions occurs when two elements in different compounds replace each other. 𝐴𝐵 + 𝐶𝐷 → 𝐴𝐷 + 𝐶𝐵  Solutions Solutions is a homogenous mixture composed of only one phase. Solubility - the maximum amount of a substance that can be dissolved in a given amount of solvent. Solute – the substance being dissolved in another substance. Solvent – a substance in which another substance is dissolved.  Types of Solutions 1.Saturated Solution – has a maximum amount of solute that are dissolved. 2.Unsaturated Solution – has less amount of solute than the solvent can possibly dissolve. 3.Supersaturated Solution – has an amount of solute greater than the amount of solute in saturated solution. In this type there is a possibility of formation of crystals. Methods of Expressing Concentrations of Solutions Mole Fraction – The number of moles of solute is divided by the number of moles of solvent and all solutes. 𝑛𝐴 𝑋𝐴 = 𝑛𝐴 + 𝑛𝐵 – Where: n = number of moles of each component Methods of Expressing Concentrations of Solutions Normality – The number of gram equivalent weights of solute per liter. A solution is “normal” if there is exactly one-gram equivalent weight per liter. 𝐸𝑞𝑢𝑖𝑣𝑎𝑙𝑒𝑛𝑡 𝑤𝑒𝑖𝑔ℎ𝑡 𝑖𝑛 𝑔𝑟𝑎𝑚𝑠 𝑁𝑜𝑟𝑚𝑎𝑙𝑖𝑡𝑦 = 𝑉𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑖𝑛 𝑙𝑖𝑡𝑒𝑟𝑠  Methods of Expressing Concentrations of Solutions Molarity – Molarity (M) is defined as the number of moles of solute dissolved in 1 liter of solution. 𝑛𝑠𝑜𝑙𝑢𝑡𝑒 𝑀= 𝑉𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑖𝑛 𝑙𝑖𝑡𝑒𝑟𝑠  Methods of Expressing Concentrations of Solutions Molality – Molality (m) is defined as the number of moles of solute dissolved in 1 kg of solvent. 𝑛𝑠𝑜𝑙𝑢𝑡𝑒 𝑚= 𝑘𝑔𝑠𝑜𝑙𝑣𝑒𝑛𝑡  Methods of Expressing Concentrations of Solutions Percent Volume refers to the number of millimeters of solute dissolved in 100ml of solution. 𝑣𝑜𝑙𝑢𝑚𝑒 𝑠𝑜𝑙𝑢𝑡𝑒 %𝑣𝑜𝑙𝑢𝑚𝑒 = 𝑥100 𝑣𝑜𝑙𝑢𝑚𝑒 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛  Dillution Dilution is the process of adding solvent to a concentrated solution to achieve a solution of the desired concentration. 𝑛𝑎𝑓𝑡𝑒𝑟 = 𝑛𝑏𝑒𝑓𝑜𝑟𝑒 𝑀𝑉𝑎𝑓𝑡𝑒𝑟 = 𝑀𝑉𝑏𝑒𝑓𝑜𝑟𝑒 SAMPLE PROBLEMS/PAST BOARD EXAM QUESTIONS 9. What is the percent composition by mass of oxygen (O) element in sulfuric acid (H2SO4)? H = 1.008 amu, S = 32.07 amu, O = 16 amu a.62.25 % b.63.34 % c.64.45 % d.65.25 % SAMPLE PROBLEMS/PAST BOARD EXAM QUESTIONS 10. How many grams of sugar ( C12H22O11) can be crystallized to form 250 ml of a 0.11 molar solution? a.8.5 g b.9.4 g c.8.8 g d.9.9 g SAMPLE PROBLEMS/PAST BOARD EXAM QUESTIONS 11. What is the molality of a solution containing 75.5g of sucrose (molar mass = 342 g/mol) in 400.0 g of water? a.0.553 mol/kg b.0.421 mol/kg c.0.386 mol/kg d.0.215 mol/kg SAMPLE PROBLEMS/PAST BOARD EXAM QUESTIONS 12. What is the mole fraction of the solute in an 40 % by mass ethanol solution in water? 1 mol of ethanol has a mass of 46 grams while 1 mol of water has a mass of 18 grams. a.0.36 mol b.0.29 mol c.0.21 mol d.0.15 mol SAMPLE PROBLEMS/PAST BOARD EXAM QUESTIONS 13. Hydrogen peroxide solution for hair bleaching is usually prepared mixing 5 grams of hydrogen peroxide (H2O2), Molecular weight = 34 g/mole per 100 mL of solution. What is the molarity of this solution? a. 1.0 M b. 1.5 M c. 1.95 M d. 1.8 M  Acids Acids is any compound that dissociates in water into H + ions. Properties of Acids 1. Acids conducts electricity in aqueous solutions. 2. Acids have a sour taste. 3. Acids turn blue litmus paper to red. 4. Acids have pH between 0 and 7. 5. Acids neutralizes bases 6. Acids react with active metals to form hydrogen. 7. Acids react with oxides and hydroxides to form salts and water.  pH Equation 1 𝑝𝐻 = log + 𝐻 Where: H+ = hydrogen ion concentration in moles per liter pOH Equation 1 𝑝𝑂𝐻 = log 𝑂𝐻 − Where: OH- = hydroxide ion concentration in moles per liter pH and OH relationship 𝑝𝐻 + 𝑝𝑂𝐻 = 14  Base Bases is any compound that dissociates in water into OH- ions. Properties of Base 1. Bitter taste 2. Turn red litmus paper to blue 3. Have pH between 7 to 14 4. Bases react with acids to form salt and water Sample MCQs 14. In a graphic representation of the energy contents of reactants and resulting products, which would have a higher energy content in exothermic reaction? a. the reactants b. the products c. Both reactants and products d. Neither one Sample MCQs 15. What principle states about the fundamental limitation that, for a particle as small as the electron, one cannot know exactly where it is and at the same time know its energy and how it is moving? a. Aufbau principle b. Uncertainty principle c. Pauli exclusion principle d. Kinetic molecular theory