General and Organic Chemistry Chapter 1 PDF

Chapter 1 Chemistry and Measurement Modified by Dr. Mousa AlTarabeen 1.1 Chemistry ❖ The study of the composition of materials (water: H2O), structure (Geometry of water) of materials and of the changes that materials undergo (physical and chemical change). ❖ Application of chemistry:...

Chapter 1 Chemistry and Measurement Modified by Dr. Mousa AlTarabeen 1.1 Chemistry ❖ The study of the composition of materials (water: H2O), structure (Geometry of water) of materials and of the changes that materials undergo (physical and chemical change). ❖ Application of chemistry: Electronics; Communications; Biochemistry; Drug Discovery (Extraction, Separation, Structure Elucidation), Nano technology; …. Etc 1.2 Experiment and Explanation Experiment: is an observation of natural phenomena carried out in a controlled manner so that the results can be duplicated. Explanation: is an analysis of results and obtain rational conclusions. 1|2 What are the steps of the scientifc Scientific Method method? Observation Hypothesis Theory Law ❖ Observation: Experiments, Results, and Notes ❖ Hypothesis: is a temporary explanation of some regularity of nature ▪ Mental picture that explains observed laws ▪ Tentative explanation of data ▪ Make predictions ▪ Leads to further tests ▪ Go to laboratory and perform experiments 1|3  Scientific Method Observation Hypothesis Theory Law Experiments Results Hypothesis 1|4 Scientific Method Figure 1.7. an representative of the scientific method diagram experiment; General steps in the scientific method.  Scientific Method Observation Hypothesis Theory Law Theory: is a tested explanation of a basic natural phenomenon ▪ Depending on results, may have to modify theory ▪ Can never prove theory is absolutely correct Law is a brief statement or mathematical equation about a fundamental relationship or regularity of nature. ▪ Usually an equation ▪ Based on results of many Force = mass x acceleration experiments ▪ Only states what happens ▪ Does not explain why they happen 1|6  Matter and Its Classifications Mass: The quantity of matter in a material Matter: Whatever occupies space and can be perceived by our senses Weight: Force with which object is attracted by gravity ❖ Example: Mass vs. Weight Astronaut on moon and on earth – Weight on moon = 1/6 weight on earth – Same mass regardless of location 1.3 Law of conservation of mass ❖ Law of Conservation of Mass The total mass remains constant during a chemical change (chemical reaction). Example: Heating 2.53 grams of metallic mercury (M in air produces 2.73 grams of a red-orange residue. Assume that the chemical change is the reaction of the metal with O2 in the air. When the red- orange residue is heated strongly, it gives back the mercury and releases the oxygen. Determine the mass of oxygen released. According to the law of conversation of mass, Mass of mercury + mass of oxygen = mass of red-orange residue Substituting, you obtain, 2.53g + mass of O2 = 2.73 g Then, Mass of O2 = (2.73 – 2.53) = 0.20 g 1.4. Matter: Physical State and Chemical Composition States of Matter: Solids: ▪ Fixed shape and volume ▪ Particles are close together ▪ Have restricted motion Liquids: ▪ Fixed volume, but take container shape ▪ Particles are close together ▪ Are able to flow Gases: ▪ Expand to fill entire container ▪ Particles separated by lots of space e.g., Ice, water, steam  Physical Change and Chemical Change Physical Change is a change in the form of matter but not in its chemical identity For example, melting and dissolving No new substances formed Substance may change state or the ratios e.g., Ice melting Sugar or salt dissolving Stirring iron filings and sulfur together Separation of NaCl and water by distillation 10  Chemical Change = Chemical Reaction A change in which one or more kinds of matter are transformed into a new kind of matter or several new kinds of matter Example - Rusting – Formation of new substance or compound – Involves changing chemical makeup of substances – New substance has different physical properties – Can’t be separated by physical means e.g., – Compound containing sulfur and iron (FeS) No longer has same physical properties of free elements Can’t be separated using magnet  Learning Check: ▪ For each of the following, determine if it represents a chemical or physical change: Your Turn! Which one of the following represents a physical change? A. when treated with bleach, some dyed fabrics change color B. grape juice left in an open unrefrigerated container turns sour C. when heated strongly, sugar turns dark brown D. in cold weather, water condenses on the inside surface of single pane windows E. when ignited with a match in open air, paper burns  Physical Properties A characteristic that can be observed for a material without changing its chemical identity e.g., Color; Physical state; Electrical conductivity; Melting point and boiling point Chemical Property A characteristic of a material involving its chemical change For example; ▪ The ability of iron to react with oxygen to produce rust ▪ Compound containing sulfur and iron No longer has same physical properties of free elements Can’t be separated using magnet Example: Potassium is a soft, silvery-colored metal that melts ₒ at 64 C. It reacts vigorously with water, with oxygen, and with chlorine. Identify all of the physical properties and chemical properties given in this description. Physical Property Chemical Property  Elements Substance that can’t be decomposed into simpler materials by chemical reactions. Substances composed of only one type of atom (e-, proton, neutron). Simplest forms of matter that we can work with directly More complex substances composed of elements in various combinations Substance: A kind of matter that cannot be separated into other kinds of matter by any physical process For example, An NaCl solution can be separated into NaCl and water, but NaCl cannot be further physically separated into new materials. Compounds ▪ Formed from two or more atoms of different elements ▪ Always combined in same fixed ratios by mass ▪ Can be broken down into elements by some chemical changes e.g., Water decomposed to elemental hydrogen and oxygen 15  Pure Substance vs. Mixture Pure substances – Elements and compounds – Composition always same regardless of source Mixture – Can have variable compositions – Made up of two or more substances – Two broad categories of mixtures: Heterogeneous; oil in water, Ice and water (same composition but different composition state) Homogeneous; Air, soft drink, milk, sugar dissolved in water, metal alloy  Relationship of Elements, Compounds, and Substances Law of Definite Proportions ▪ A pure compound, regardless of its source, always contains constant proportions of the elements by mass. ‫ دائما‬،‫ بغض النظر عن مصدره‬،‫يحتوي املركب النقي‬.‫على نسب ثابتة من العناصر حسب الكتلة‬ For example, 1.0000 gram of sodium chloride will always contain 0.3934 gram of sodium and 0.6066 gram of chlorine.  Learning Check: Classification Your Turn! Coffee is a homogeneous mixture. When you add sugar to it and the sugar dissolves… A. the coffee becomes a heterogeneous mixture. B. the coffee is still a homogeneous mixture. C. the coffee, which used to be a pure compound, is no longer a pure compound. D. the coffee, which used to be an element, is no longer an element. E. the sugar reacts with water to form a new compound. 18  1.5 Measurement and Significant figures 1. Measurements involve comparison – Always measure relative to reference (Unit) e.g., Foot, meter, kilogram – Measurement = number + unit e.g., Distance between 2 points = 25 What unit? inches, feet, yards, miles Meaningless without units 2. Measurements are inexact – Measuring involves estimation – Always have uncertainty 1 | 19 Accuracy – how close a measurement is to the true value of the measured quantity Precision – how close a set of measurements are to each other accurate precise not accurate & but & precise not accurate not precise 1.8  true mass of 2.000 g Student A Student B Student C 1.964 g 1.972 g 2.000 g 1.978 g 1.968 g 2.002 g Average 1.971 g 1.970 g 2.001 g Student B results are more precise than those of student A Student C results are the most precise and the most accurate Generally highly accurate measurements are precise too, but highly precise measurements do not necessarily guarantee accurate results  Uncertainties in Measurements Measurements all inexact – Contain uncertainties or errors Sources of errors – Limitations of reading instrument Ways to minimize errors – Take series of measurements – Data clusters around central value – Calculate average or mean values – Report average value 22  Significant Figures Those digits in a measured number that include all certain digits plus a final digit having some uncertainty; For example, in the measurement 9.12, the first two digits (9.1) are certain and the last digit (2) is estimated. ▪ Measure a volume by a graduated cylinder with a scale gives uncertainty of 1 ml. The measured volume is 6 ml; the actual volume is in the range 5 to 7, or 6 ± 1, in this case; the 5 is Uncertain; there is one significant figure. ▪ Measure the volume by cylinder with finer divisions and uncertainty of 0.1 ml. 6.0 ml or (6.0 ± 0.1) ml ; the 0 is Uncertain; the actual volume is 5.9 – 6.1 ml; two significant figures. ▪ Measure the mass with analytical balance … 3.7654 g; five significant figures; the 4 is uncertain  Significant Figures Any digit that is not zero is significant 1.234 kg 4 significant figures Zeros between nonzero digits are significant 606 m 3 significant figures Zeros to the left of the first nonzero digit are not significant 0.08, 0.008 and 0.0008 have 1 significant figure 0.08 x 10-2, or 0.8 x 10-3, or 8 x 10-4. 0.080 and 0.0080; 2 sig fig; 0.80 x 10-2, or 8.0 x 10-3 If a number is greater than 1, then all zeros to the right of the decimal point are significant 2.0 mg 2 significant figures 56.0707 ml 6 sig. figs Trailing zeros (the right most zeros) are significant when there is a decimal point in the number. 400. has three sig figures; 2.00 has three sign figures ; 0.050 has two sig. fig  Trailing zeros are not significant in numbers without decimal points. 47000 has two significant figures ; Ex; the number that attend a football match is approximately 47 thousands 400 or 4x102 indicates only one. Exact numbers have an infinite number of significant digits but they are generally not reported If you count 2 pencils, then the number of pencils is 2.000... ❖ The number of centimeters per inch (2.54) has an infinite number of significant digits, as does the speed of light (299792458 m/s). Example: 24 mL ----- 2 sig. no. 3001 g ------------- 4 sig. no. 0.0320 m3 ------------- 3 sig. no. 6.4 x 104 molecules ----------- 2 sig. no 3000 mL. Depend on the decimal, the number of significant figures may be four (3.000 × 103), three (3.00 × 103), two (3.0 ×103), or one (3 x 103). How many significant figures does each of the following numbers have? 1. 413.97 2. 0.0006 3. 5.120063 4. 161,000 5. 3600. Write the following number (2500) so that the number of significant figures are A. 2:---------------- B. 3:---------------- C. 4:---------------- 26  Your Turn! How many significant figures are in 19.0000? A. 2 B. 3 C. 4 D. 5 E. 6 How many significant figures are in 0.0005650850? A. 7 B. 8 C. 9 D. 10 E. 11 ▪ Could be rewritten as 5.650850  10-4 27 Rounding The procedure of dropping nonsignificant digits and adjusting the last digit reported in the final result of a calculation Rounding Procedure 1. Look at the leftmost digit to be dropped. 2. If this digit is greater than 5: Add 1 to the last digit to be retained. Drop all digits farther to the right. 3. If this digit is less than 5: Drop all digits farther to the right. 4. If this digit is = 5: If the digits farther to right is even then follow step three If the digits farther to right is off then follow step two Rounding result of a calculation For example, 1.2151 rounded to three significant figures is 1.22. 1.2143 rounded to three significant figures is 1.21. 1.2245:------------- 5.3275:-------------  Significant Figures Addition or Subtraction The answer cannot have more digits to the right of the decimal point than any of the original numbers. 89.332 +1.1 one significant figure after decimal point 90.432 round off to 90.4 3.70 two significant figures after decimal point -2.9133 0.7867 round off to 0.79  Significant Figures Multiplication or Division The number of significant figures in the result is set by the original number that has the smallest number of significant figures 4.51 x 3.6666 = 16.536366 = 16.5 3 sig figs round to 3 sig figs 6.8 ÷ 112.04 = 0.0606926 = 0.061 2 sig figs round to 2 sig figs  Significant Figures Exact Numbers Numbers from definitions or numbers of objects are considered to have an infinite number of significant figures The average of three measured lengths; 6.64, 6.68 and 6.70? 6.64 + 6.68 + 6.70 = 6.67333 = 6.67 =7 3 Because 3 is an exact number and does not determine the number of significant figures Example: Carry out the following arithmetic operations to the correct number of significant figures: (a) 11,254.1 g + 0.1983 g (a) 66.59 L − 3.113 L  (c) 8.16 m × 5.1355 (d) 0.0154 kg  88.3 mL (e) 2.641 × 103 cm + 3.27 × 102 cm First we change 3.27 × 102 cm to 0.327 × 103 cm and then carry out the addition (2.641 cm + 0.327 cm) ×103. The answer is 2.97 × 103 cm.  Learning Check Round each of the following to three significant figures. Use scientific notation where needed. 1. 37.459 2. 5431978 3. 132.7789003 4. 0.00087564 5. 7.665 For each calculation, give the answer to the correct number of significant figures. 1. 10.0 g + 1.03 g + 0.243 g = 2. 19.556 °C – 19.552 °C = 3. 327.5 m × 4.52 m = 4. 15.985 g ÷ 24.12 mL = 35 Your Turn! Round 0.00564458 to four significant figures and express using scientific notation. A. 5.64  10-2 B. 5.000  10-3 C. 5.645  10-4 D. 0.56446 E. 5.645  10-3 Give the value of the following calculation to the correct number of significant figures. 1.199 + 0.001 =------------------ A. 1.21213 B. 1.212132774 D. 1.2 E. 1 When the expression, 412.272 + 0.00031 – 1.00797 + 0.000024 + 12.8 is evaluated, the result should be expressed as: A. 424.06 B. 424.064364 C. 424.1 D. 424.064 E. 424 36  Learning Check For the following calculation, give the answer to the correct number of significant figures. 1. (71.359 m − 71.357 m) = (0.002 m) 0.00200 m = 0.0017 m/s2 (3.2 s × 3.67 s) (11.744 s2 ) 12 s2 2. (13.674 cm × 4.35 cm × 0.35 cm ) = (20.818665 cm )= 21 = 0.0088 cm3/s 3 (856 s + 1531.1 s) (2387.1 s) 2387 3. 2.568×5.8 = 15 = 3.6 4.168 4.168 Using a calculator, the answer is 3.573512476; should be rounded to two significant figures: 3.6 4. 4.18–58.16  (3.38 – 3.01)= 4.18- 58.16* (0.37)= 4.18 – 22 = -18  Decimal Multipliers ❖ Angstrom (Å) : Non-SI unit traditionally used in measuring small lengths Is equal to 10–10 m  Laboratory Measurements Four common: 1. Length 2. Volume 3. Mass 4.Temperature 1. Length ▪ SI Unit is meter (m) ▪ Meter too large for most laboratory measurements ▪ Centimeter (cm) ▪ 1 cm = 10–2 m = 0.01 m ▪ Millimeter (mm) ▪ 1 mm = 10–3 m = 0.001 m 2. Volume Dimensions of (length)3 SI unit for Volume = m3 Most laboratory use V in liters (L) – 1 L = 1 dm3 (exactly) Chemistry glassware marked in L or mL – 1 L = 1000 mL What is a mL? 1 mL = 1 cm3 3. mass; Kg or g= 10-3Kg  4. Temperature Measured with thermometer with three common scales A. Fahrenheit scale Water freezes at 32 °F and boils at 212 °F 180 degree units between melting and boiling points of water B. Celsius scale Most common for use in science Water freezes at 0 °C and boils at 100 °C 100 degree units between melting and boiling points of water C. Kelvin scale SI unit of temperature is kelvin (K) ▪ Note: No degree symbol in front of K Water freezes at 273.15 K and boils at 373.15 K ▪ 100 degree units between melting and boiling points Absolute Zero ▪ Zero point on Kelvin scale corresponds to nature’s lowest possible temperature Temperature Conversions How to convert between˚F and ˚C?  9 F  t F =   t C + 32F 5 C Example: 100 °C = ? ˚F  9 F  t F =   100 C + 32F 5 C tF = 212 ˚F 41 Temperature Conversions Common laboratory thermometers are marked in Celsius scale Must convert to Kelvin scale ‫يجب التحويل إلى مقياس كلفن‬ 1K TK = (t C + 273.15 C)   1 C Amounts to adding 273.15 to Celsius temperature Example: What is the Kelvin temperature of a solution at 25 °C? 1K TK = (25 C + 273.15 C)    = 298 K 1 C 42  Learning Check: T Conversions 1. Convert 121 ˚F to the Celsius scale. æ 9 °F ö æ 5 °C ö t F = çç ÷÷ t C + 32 °F è 5 °C ø ( ) t C = t F - 32 °F çç ÷÷ è 9 °F ø ( tC = 121oF − 32°F  5 °C    9 °F  ) = 49 °C 2. Convert 121 ˚F to the Kelvin scale. – We already have in °C so… 1K 1K TK = ( tC + 273.15°C) = (49+ 273.15°C) TK = 322.15 K 1°C 1°C 43  Your Turn! In a recent accident some drums of uranium hexafluoride were lost in the English Channel. The melting point of uranium hexafluoride is 64.53 °C. What is the melting point of uranium hexafluoride on the Fahrenheit scale? A. 67.85 °F B. 96.53 °F C. 116.2 °F D. 337.5 °F E. 148.2 °F On an absolute temperature scale, 100 °F is not double 50 °F (i.e., not ‘twice as hot’). What temperature, in °F, would really be double 50 °F (hint: convert 50 °F to K, double the Kelvin temperature, then convert back to °F)? A. 560 °F B. 25 °F C. 200 °F D. 283.15 °F E. 566 °F 44