Abdullah Al-Salem University Chemistry 101 Lecture Notes PDF

Abdullah Al-Salem University Chemistry 101 Dr.Fatma Hussain 1  CHM 1 (Fall Semester) - 2024/2025 ❖Textbook: Chemistry 14th edition. By: By Jason overby and Raymond Chang, McGraw-Hill, Hiher education (2022) ❖ Course Contents: No....

Abdullah Al-Salem University Chemistry 101 Dr.Fatma Hussain 1  CHM 1 (Fall Semester) - 2024/2025 ❖Textbook: Chemistry 14th edition. By: By Jason overby and Raymond Chang, McGraw-Hill, Hiher education (2022) ❖ Course Contents: No. Subject Chapter Sections Excluded 1 Measurement and the properties of 1 1.1&1.9 matter 2 Atoms, Ions and Molecules 2 2.8 3 Mass Relationships in Chemical 3 3.4 reactions 4 Reactions in Aqueous Solutions 4 4.6 & 4.7 5 Thermochemistry 6 - 6 Quantum Theory and Electronic 7 7.9 structure of Atoms 7 Periodic Relationships Among the 8 - Elements 8 Compounds and Bonding 9 9.10 9 Structure and Bonding Theories 10 10.6, 10.7 & 10.8 10 Organic chemistry 24 24.3 ❖Grading: Midterm Exams: 30% (15% each, 2 Exams) Final Exam: 40% Quizzes: 15% Homework: 10% Classwork/Participitation: 5% ❖ Exams: 1st Midterm Exam : 15% (1, 2, 3 & 4) 2nd Midterm Exam : 15% (6, 7, 8 & 9) Attendance Policies ❖ 3 hours of absence results in a first warning. ❖ 6 hours of absence results in a second warning. ❖ Any hour of absence after 6 results in dismissal from the course (FA). Important: ❖ Coming to class after 15 minutes must be considered absent even if they attend the rest of the class. EXAMS DATES 1st Midterm Exam: - Chapters (1, 2, 3 & 4). 2nd Midterm Exam : - Chapters (6, 7, 8 & 9). Final Exam : - All Chapters. 5 Chapter 1 Introduction 6  Lecture Outline 1.2 The Scientific Method 1.3 Measurment 1.4 handling Numbers 1.5 Dimensional Analysis in Solving Problems 1.6 Real-World Problem Solving 1.7 Classifications of Matter 1.8 The Three Common States of Matter  The Scientific Method The Scientific method: Overall philosophy of approach to the study of nature. observation Law Hypotheses Theory Statements of fact tentative explanation tested explanation often in equation of set of of experiments form observation results (explains body of facts) qualitative quantitative no numbers numbers 2Na + Cl2 2NaCl concentration Note thattheory is not fact and it could be wrong but we need to proof it.  Example : 1. Classify each of the following statements as a hypothesis, a law, or a theory. (a) An autumn leaf gravitates toward the ground because there is an attractive force between the leaf and Earth. (b) All matter is composed of very small particles called atoms. The Basic Units of Measurement Units tell the standard quantity to which we are comparing the measured property. Without an associated unit, a measurement is without meaning. Scientists use a set of standard units for comparing all our measurements. So we can easily compare our results. Each of the units is defined as precisely as possible. 33  What Is a Measurement? The unit tells you to what standard you are comparing your object. The number tells you: 1.What multiple of the standard the object measures. 2. The uncertainty in the measurement. ❖More digits → more precision  Volume Derived unit. Any length unit cubed. Measure of the amount of space occupied. SI unit = cubic meter (m3) Commonly measure solid volume in cubic centimeters (cm3). Commonly measure liquid or gas volume in milliliters (mL).  Density Mass Density = Ratio of mass : volume. Volume Its value depends on the kind of material, not the amount. Solids = g/cm3 (1 cm3 = 1 mL) Liquids = g/mL Gases = g/L Density Density : solids > liquids > gases Except ice is less dense than liquid water! Density For equal volumes, the more dense (D) object has a larger mass (m). (↑ m ↑ D) For equal masses, the more dense (D) object has a smaller volume (v). (↓ v ↑ D) Heating objects causes objects to expand. This does not effect their mass! How would heating an object effect its density?  Using Density in Calculations Solution Maps: What Volume Does 100.0 g of Marble Occupy (D = 4.00 g/cm3)?  Home work Q1: The dimensions of a sample of iron are 1.20 cm, 1.30 cm, and 1.20 cm. Its mass 1.39 g. Calculate its density? Q2: A student is trying to find the density of a cube each side of the cube measure 4 cm and the mass is 503.68 g. Calculate the density?  SI-Derived Unit 1 cm = (110 m ) = 110−6 m3 3 −2 3 Area = m2 Volume = m3 1 dm = (110 m ) = 110−3 m3 3 −1 3 Speed = m/s 1 L = 1000 mL = 1000 cm3 = 1 dm3 Unit conversion Usually we use Factor-label method : units undergo the same kind of mathematical operation as numbers (Convert from one unit to another unit) Conversion Factor: a fraction that we use to change the units Converting from One Unit to Another Always write every number with its associated unit. Always include units in your calculations. You can do the same kind of operations on units as you can with numbers. cm × cm = cm2 cm + cm = cm cm ÷ cm = 1 Using units as a guide to problem solving is called dimensional analysis. Common Units and Their Equivalents Length 1 kilometer (km) = 0.6214 mile (mi) 1 meter (m) = 39.37 inches (in.) 1 meter (m) = 1.094 yards (yd) 1 foot (ft) = 30.48 centimeters (cm) 1 inch (in.) = 2.54 centimeters (cm) exactly Common Units and Their Equivalents, Continued Mass 1 kilogram (km) = 2.205 pounds (lb) 1 pound (lb) = 453.59 grams (g) 1 ounce (oz) = 28.35 (g) Volume 1 liter (L) = 1000 milliliters (mL) 1 liter (L) = 1000 cubic centimeters (cm3) 1 liter (L) = 1.057 quarts (qt) 1 U.S. gallon (gal) = 3.785 liters (L)  Example—Convert 7.8 km to Miles 1. Write down the Given quantity and Given: 7.8 km its unit. 2. Write down the quantity you want Find: ? miles to Find and unit. 3. Write down the appropriate Conversion 1 km = 0.6214 mi Conversion Factors.GIVEN Factor: 4. Write a Solution Map. Solution km mi Map: 0.6214 mi 1 km 5. Follow the solution map to Solve Solution: the problem. 0.6214 mi 7.8 km  = 4.84692 mi 1 km 6. Significant figures and round. Round: 4.84692 mi 7. Check. Check: Units and magnitude are correct. Example —Convert 2,659 cm2 into Square Meters 1. Write down the Given quantity and Given: 2,659 cm2 its unit. 2. Write down the quantity you want Find: ? m2 to Find and unit. 3. Write down the appropriate Conversion 1 cm = 0.01 m Conversion Factors. Factor: 4. Write a Solution Map. Solution cm2 m2 Map: 2  0.01 m     1 cm  5. Follow the solution map to Solve Solution: the problem. 110 −4 m 2 2,659 cm  2 = 0.2659 m 2 1 cm 2 6. Significant figures and round. Round: 0.2659 m2 7. Check. Check: Units and magnitude are correct.  Temperature Scale TC = TK – 273.15 TC = (TF – 32)  5 9 TF = (TC  9 ) + 32 TK = TC+ 273.15 5 e.g. Convert 172.9 0F to degrees Celsius. TC = 78.28 0C Sa. Ex. 1.12 Liquid nitrogen boil at 77 K convert to Fahrenheit. TF= -321 0F Scientific Notation Scientific Notation A way of writing large and small numbers. An atom’s The sun’s average diameter is diameter is 0.000 000 000 3 m. 1,392,000,000 m.  Scientific Notation Each decimal place in our number system represents a different power of 10. Scientific notation writes the numbers so they are easily comparable by looking at the power of 10. The sun’s The sun’s diameter is diameter is 1,392,000,000 m. 1.392 x 109 m. An atom’s average diameter is An atom’s average diameter is 0.000 000 000 3 m. 3 x 10-10 m.  Exponents(x 10 ) Exponent 1.23 x 10-8 Decimal part Exponent part When the exponent on 10 is positive, it means that the powers of 10 larger. When the exponent on 10 is negative, it means that the power of 10 smaller.  Conversion of Scientific Notation If you moved the decimal point If you moved the decimal point to the left, to the right, ✓ then n is + ✓ then n is − ✓ If the original number is ✓ If the original number is 1 or larger less than 1  Scientific Notation To compare numbers written in scientific notation: First compare exponents on 10. If exponents are equal, then compare decimal numbers 1.23 x 105 4.56 x 102 1.23 x 105 > 4.56 x 102 4.56 x 10-2 7.89 x 10-5 4.56 x 10-2 > 7.89 x 10-5 7.89 x 1010 1.23 x 1010 7.89 x 1010 > 1.23 x 1010 Writing a Number in Scientific Notation Example (1) : 12340 1. Locate the decimal point. 12340. 2. Move the decimal point to obtain a number between 0 and 10 (e.g. 1,2,3…..9). 1.234 3. Multiply the new number by 10n. Where n is the number of places you moved the decimal point. 1.234 x 104 4. If you moved the decimal point to the left, then n is +. 1.234 x 104 Example (2) : 0.00012340 1. Locate the decimal point. 0.00012340 2. Move the decimal point to obtain a number between 0 and 10 (e.g. 1,2,3…..9). 1.2340 3. Multiply the new number by 10n. Where n is the number of places you moved the decimal point. 1.2340 x 104 4. If you moved the decimal point to the right, then n is −. 1.2340 x 10-4  Writing a Number in Standard Form Positive Power = Big Number Negative Power = Small Number 1.234 x 106 1.234 x 10-6 ✓ Move the decimal point to the ✓Move the decimal point to the right 6 places. left 6 place. ✓ When you run out of digits to When you run out of digits to move around, add zeros. move around, add zeros. ✓ Add a zero in front of the Add a zero in front of the decimal point for decimal decimal point for decimal numbers. numbers. 1.234000 000 001.234 1234000..000001234000  More Example The U.S. population in 2007 was estimated to be 301,786,000 people. Express this number in scientific notation. 301,786,000 people = Write the Following in Scientific Notation 123.4 = 8.0012 = 145000 = 0.00234 = 25.25 = 0.0123 = 1.45 = 0.000 008706 = Write the Following in Standard Form 2.1 x 103 = 4.02 x 100 = 9.66 x 10-4 = 3.3 x 101 = 6.04 x 10-2 = 1.2 x 100 = Inputting Scientific Notation into a Calculator Input the decimal part of -1.23 x 10-3 the number. If negative press +/- key. Input 1.23 1.23 (–) on some. Press EXP. Press +/- -1.23 EE on some. Input exponent on 10. Press EXP -1.23 00 Press +/- key to change exponent to negative. Input 3 -1.23 03 Press +/- -1.23 -03 Calculator Exponents(x 10 ) Significant Figures Significant Figures Writing numbers to reflect precision.  Reporting Measurements The system of writing measurements we use is called significant figures. When writing measurements, all the digits written are known with certainty except the last one, which is an estimate. An example of estimating the last digit: 45.872 Estimated Certain  Significant Figures rules Non-zero digits are always significant e.g. 32 (2 sig.) 12.3 (3 sig.) Any zeros between two significant digits are significant e.g. 1005 (4 sig) 7.03 (3 sig) Zeros at the beginning of a number are never significant e.g. 000026 (2 sig)  Significant Figures rules A final zero or trailing zeros are significant if the number contains in a decimal point e.g. 0.0200 (3 sig.) 3.0 (2 sig.) 12.30 cm (4 sig.) Zeros at the end of a number without a written decimal point are ambiguous and should be avoided by using scientific notation. e.g. 150 (2 sig. or 3 sig.) 1.5 x 102 (2 sig.) 1.50 x 102 (3 sig.) What is the number of significant figures in each of the following measurements? (a) 4867 mi (b) 56 mL (c) 60,104 tons (d) 2900 g (e) 40.2 g/cm3 (f) 0.0000003 cm (g) 0.7 min (h) 4.6 × 1019 atoms EXAMPLE 2.4 1 Determining the Number of Significant Figures in a Number How many significant figures are in each of the following numbers? (a) 0.0035 (d) 2.97 × 105 (b) 1.080 (e) 1 dozen = 12 (c) 2371 (f) 100,000 Solution: The 3 and the 5 are significant. The leading zeros (a) 0.0035 two significant figures only mark the decimal place and are not significant. The interior zero and the trailing zero are (b) 1.080 four significant figures significant, as are the 1 and the 8. All digits are significant. (c) 2371 four significant figures All digits in the decimal part are significant. (d) 2.97 × 105 three significant figures Defined numbers have an unlimited number of (e) 1 dozen = 12 unlimited significant significant figures. figures This number is ambiguous. Write as 1 × 105 to (f) 100,000 ambiguous indicate one significant figure or as 1.00000 × 105 to indicate six significant figures.  Significant Figures in Calculations A. Multiplication and Division with Significant Figures For multiplication (×) and division (÷), the result contains the same number figures as the measurement with fewest significant figures. 5.02 × 89,665 × 0.10 = 45.0118 = 45 3 sig. figs. 5 sig. figs. 2 sig. figs. 2 sig. figs. 5.892 ÷ 6.10 = 0.96590 = 0.966 4 sig. figs. 3 sig. figs. 3 sig. figs. EXAMPLE 2.5 2 Significant Figures in Multiplication and Division Perform the following calculations to the correct number of significant figures. (a) 1.01 × 0.12 × 53.51 ÷ 96 (b) 56.55 × 0.920 ÷ 34.2585 Solution: Round the intermediate result (in blue) to (a) 1.01 × 0.12 × 53.51 ÷ 96 = 0.067556 two significant figures to reflect the two significant figures in the least precisely 3 sf 2 sf 4 sf 2 sf = 0.068 Result should known quantities (0.12 and 96). have 2 sf. Round the intermediate result (in blue) to (b) 56.55 × 0.920 ÷ 34.2585 = 1.51863 three significant figures to reflect the three = 1.52 4 sf 3 sf 6 sf significant figures in the least precisely Result should known quantity (0.920). have 3 sf. 3 SKILLBUILDER 2.5 Significant Figures in Multiplication and Division Perform the following calculations to the correct number of significant figures. (a) 1.10 × 0.512 × 1.301 × 0.005 ÷ 3.4 (b) 4.562 × 3.99870 ÷ 89.5 B. Addition and Subtraction with Significant Figures For addition and subtraction, the result has the same number of decimal places as the measurement with fewest decimal places. 5.74 + 0.823 + 2.651 = 9.214 = 9.21 2 dec. pl. 3 dec. pl. 3 dec. pl. 2 dec. pl. 4.8 - 3.965 = 0.835 = 0.8 1 dec. pl 3 dec. pl. 1 dec. pl. EXAMPLE 2.6 4 Significant Figures in Addition and Subtraction Perform the following calculations to the correct number of significant figures. 0.987 + 125.1 – 1.22 = 0.765 – 3.449 – 5.98 = Solution: Round the intermediate answer (in blue) to one (a) decimal place to reflect the quantity with the fewest decimal places (125.1). Notice that 125.1 is not the quantity with the fewest significant figures—it has four while the other quantities only have three—but because it has the fewest decimal places, it determines the number of decimal places in the answer. Round the intermediate answer (in blue) to two (b) decimal places to reflect the quantity with the fewest decimal places (5.98). C. Rounding When rounding to the correct number of significant figures, if the number after the place of the last significant figure is: 1. 0 to 4, round down. ✓ Drop all digits after the last significant figure and leave the last significant figure alone. ✓ Add insignificant zeros to keep the value, if necessary. ✓ Example: Rounding to 2 significant figures: e.g. 2.34 2.3 (2 sig. fig.) e.g. 2.349865 2.3 (2 sig. fig.)  Rounding, continued When rounding to the correct number of significant figures, if the number after the place of the last significant figure is: 2. From 5 to 9, round up. ✓ Drop all digits after the last significant figure and increase the last significant figure by one. ✓ Add insignificant zeros to keep the value, if necessary. ✓ Example: Rounding to 2 significant figures: e.g. 2.37 2.4 (2 sig. fig.) Rounding, Continued Example : Rounding to 2 significant figures e.g 0.0234 0.023 2.3 × 10-2 (Scientific notation) e.g. 0.0237 0.024 2.4 × 10-2 (Scientific notation) e.g. 0.02349865 0.023 2.3 × 10-2 (Scientific notation) Rounding, Continued Example : Rounding to 2 significant figures e.g 234 230 2.3 × 102 (Scientific notation) e.g. 237 240 2.4 × 102 (Scientific notation) e.g. 234.9865 230 2.3 × 102 (Scientific notation) Rounding, Continued Example : Rounding to 2 significant figures 5.82 × 96 = 558.72 3 sig. figs. 2 sig. figs. 2 sig. figs. = 56 (scientific notation) = 5.5872 × 102 = 5.6 × 102 (2 sig. fig.) √ D. Both Multiplication/Division & Addition/Subtraction with Significant Figures When doing different kinds of operations with measurements with significant figures. Follow the standard order of operations. Please Excuse My Dear Aunt Sally. ( )→ n → → + - 3.489 × (5.67 – 2.3) = 4 sf 2 dp 1 dp 3.489 × 3.37 = 12 4 sf 1 dp & 2 sf 2 sf Example 1: Perform the Following Calculations to the Correct Number of Significant Figures a) 1.10  0.5120  4.0015  3.4555 0.355 b) + 105.1 − 100.5820 c) 4.562  3.99870  (452.6755 − 452.33) d) (14.84  0.55) − 8.02 Example 1, continue a) 1.10  0.5120  4.0015  3.4555 = 0.65219 = 0.652 0.355 b) + 105.1 − 100.5820 4.8730 = 4.9 c) 4.562  3.99870  (452.6755 − 452.33) = 52.79904 = 53 d) (14.84  0.55) − 8.02 = 0.142 = 0.1  Accuracy, Precision and Significant Figure Accuracy – how close a measurement is to the true value : A -TRUE Precision – how close a set of measurements are to each other P - EACH accurate precise not accurate & but & precise not accurate not precise Example : A laboratory technician analyzed a sample three times for percent iron and got the following results: 22.43% Fe, 24.98% Fe, and 21.02% Fe. The actual percent iron in the sample was 22.81%. The analyst's A.Precision was poor but the average result was accurate. B.Accuracy was poor but the precision was good. C.Work was only qualitative. D.Work was precise. E.C and D. Problem Solving and Dimensional Analysis Many problems in chemistry involve using relationships to convert one unit of measurement to another. Conversion factors are relationships between two units. May be exact or measured. Both parts of the conversion factor have the same number of significant figures. Conversion factors generated from equivalence statements. e.g., 1 inch = 2.54 cm can give or 2.54cm 1in 1in 2.54cm Matter, Mixture and Separation methods : A Matter any thing occupies a space and has a mass gas Three States liquid solid  In Your Room Everything you can see, touch, smell or taste in your room is made of matter. Chemists study the differences in matter and how that relates to the structure of matter.  What Is Matter? Matter is defined as anything that occupies space and has mass. Even though it appears to be smooth and continuous, matter is actually composed of a lot of tiny little pieces we call atoms and molecules. Atoms are the tiny particles that make up all matter. In most substances, the atoms are joined together in units called molecules. The atoms are joined in specific geometric arrangements. Classifying Matter by Physical State Matter can be classified as solid, liquid, or gas based on what properties it exhibits. State Shape Volume Compress Flow Solid Fixed Fixed No No Liquid Indefinite Fixed No Yes Gas Indefinite Indefinite Yes Yes The atoms or molecules have different structures in solids, liquids, and gases − leading to different properties. Fixed = Property doesn’t change when placed in a container. Indefinite = Takes the property of the container.  Anything occupies space and has mass 2 or more substances each retain their distinct identities Form of matter has definite composition substance with constant substance that cannot mixture composition mixture composition composition that can be be decompose into simpler is the same not uniform broken down into elements form by chemical or by chemical process physical change OR Two or more different atoms bind together by a chemical bond What is the different between element and atom? A mixture is a combination of two or more substances in which the substances retain their distinct identities. 1. Homogenous mixture – composition of the mixture is the same throughout. soft drink, milk, solder 2. Heterogeneous mixture – composition is not uniform throughout. cement, iron filings in sand Example: Classify each of the following as an element, a compound, a homogeneous mixture, or a heterogeneous mixture: (a)water from a well (b) Argon gas (c) Sucrose (d) chicken noodle soup  Physical and Chemical Change A physical change does not alter the composition or identity of a substance. sugar dissolving ice melting in water A chemical change alters the composition or identity of the substance(s) involved. hydrogen burns in air to form water Some Physical Properties Mass Volume Density Solid Liquid Gas Melting point Boiling point Volatility Taste Odor Color Texture Shape Solubility Electrical Thermal Magnetism conductance conductance Malleability Ductility Specific heat capacity Some Chemical Properties Acidity Basicity (aka alkalinity) Causticity Corrosiveness Reactivity Stability Inertness Explosiveness (In)Flammability Combustibility Oxidizing ability Reducing ability Example: Which of the following represents a chemical change? A.Boiling water to form steam B.Burning a piece of coal C.Heating lead until it melts D.Mixing iron powder and sand at room temperature E.Breaking glass  Extensive and Intensive properties An extensive property of a material depends upon how much matter is is being considered. mass length volume An intensive property of a material does not depend upon how much matter is is being considered. density temperature m.p, f.p, b.p etc color Example: Which of the following is an extensive property? A.Density. B.Specific heat. C.Mass. D.Color. E. Melting point.