General Chemistry for Health Sciences Lecture Notes PDF

General Chemistry-for Health Sciences Dr. Montaha Anjass Almasri Assistant Professor Department of Chemistry University of Sharjah [email protected] www.anjassgroup.net 1  Content 1. Chapter 1. Chemistry: Methods and Measurements 2. Chapter 2. The Structure of Atoms and the Per...

General Chemistry-for Health Sciences Dr. Montaha Anjass Almasri Assistant Professor Department of Chemistry University of Sharjah [email protected] www.anjassgroup.net 1  Content 1. Chapter 1. Chemistry: Methods and Measurements 2. Chapter 2. The Structure of Atoms and the Periodic Table 3. Chapter 3. Structure and Properties of Ionic and Covalent Compounds 4. Chapter 4. Calculations and the Chemical Equation 5. Chapter 5. States of Matter: Gases, Liquids, and Solids 6. Chapter 6. Solutions 7. Chapter 7. Energy, Rate, and Equilibrium 8. Chapter 8. Acids and Bases and Oxidation Reduction 9. Chapter 9. The Nucleus, Radioactivity, and Nuclear Medicine General, Organic and Biochemistry, K. Denniston and D. Quirk 8th Edition 2  GENERAL CHEMISTRY General, Organic and Biochemistry, K. Denniston and D. Quirk 8th Edition Copyright© The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chemistry The study of matter, its chemical and physical properties, the chemical and physical changes it undergoes, and the energy changes that accompany these processes Matter: Anything that has mass and occupies space 1.2 The Categorization of Matter Properties - characteristics of matter scientists can use to categorize different types of matter. Ways to Categorize matter: 1. By State 2. By Composition 1.2 The Categorization of Three States of Matter 1.Gas - particles widely separated, no definite shape or volume 4 Matter 2. Liquid - particles closer together, definite volume but no definite shape 3. Solid - particles are very close together, define shape and definite volume  Three States of Water (a) Solid (b) Liquid (c) Gas  5 1.2 The Categorization of Physical property - is observed without changing the composition or identity of a substance Physical change - produces a Matter recognizable difference in the appearance of a substance without causing any change in its composition or identity - conversion from one physical state to another - melting an ice cube Separation by Physical Properties Magnetic iron is separated from other nonmagnetic substances, such as sand. This property is used as a large-scale process in the recycling industry. 1.2 The Categorization of Chemical property - result in a change in composition and can be 5 observed only through a chemical reaction Matter Chemical reaction (chemical change) - a chemical substance is converted into one or more different substances by rearranging, removing, replacing, or adding atoms. hydrogen + oxygen → water reactants products  Classify the following as either a 1.2 The Categorization of chemical or physical property: a. Color (Physical) b. Flammability (Chemical) Matter c. Hardness (Physical) d. Odor (Physical) e. Taste (Physical)  Classify the following as either a 1.2 The Categorization of chemical or physical change: a. Boiling water becomes steam Matter b. Butter turns rancid c. Burning of wood d. Mountain snow melting in spring e. Decay of leaves in winter 1.2 The Categorization of Intensive properties - a property of matter that is independent of the 6 quantity of the substance - Color Matter - Melting Point Extensive properties - a property of matter that depends on the quantity of the substance - Mass - Volume 1.2 The Categorization of Composition of Matter Matter Pure substance - a substance that has only one component 7 Mixture - a combination of two or more pure substances in which each substance retains its own identity, not undergoing a chemical reaction 1.2 The Categorization of Composition of Matter Matter Element - a pure substance that cannot be changed into a simpler form of matter by any chemical reaction 7 Compound - a pure substance resulting from the combination of two or more elements in a definite, reproducible way, in a fixed ratio 1.2 The Categorization of Composition of Matter Matter Mixture - a combination of two or more pure substances in which each substance retains its own identity 7 Homogeneous - uniform composition, particles well mixed, thoroughly intermingled Heterogeneous – nonuniform composition, random placement  1.3 The Units of Measurement Units - the basic quantity of mass, volume or whatever quantity is being measured – A measurement is useless without its units English system - a collection of functionally unrelated units – Difficult to convert from one unit to another – 1 foot = 12 inches = 0.33 yard = 1/5280 miles Metric System - composed of a set of units that are related to each other decimally, systematic – Units relate by powers of tens  Metric System Units 1.3 The Units of Mass - the quantity of matter in an object Measurement – not synonymous with weight – standard unit is the gram (g) – The pound (lb) is the common English unit. 1 lb = 454 g Mass must be measured on a balance (not a scale)  Length - the distance between two points – standard unit is the meter (m) 1.3 The Units of – The yard is the common English unit. Measurement 1 yd = 0.91 m Volume - the space occupied by an object – standard unit is the liter – The quart is the common English Unit 1 qt = 0.946 L 8  Volume = lxwxh 1.3 The Units of Measurement Volume = 10 cm x 10 cm x 10 cm = 1000 cm3 = 1dm3 1 dm3 = 1 L  Time 1.3 The Units of Measurement - metric unit is the second  Metric System Prefixes Basic units are the units of a quantity 1.3 The Units of without any metric prefix. Measurement 1.4 The Numbers of Measurement (start here) Information-bearing digits or figures in a number are significant figures The measuring devise used determines the number of significant figures in a measurement The degree of uncertainty associated with a measurement is indicated by the number of figures used to represent the information 1.4 The Numbers of Measurement Significant figures - all digits in a number representing data or results that are known with certainty plus one uncertain digit  Recognition of Significant Figures 1.4 The Numbers of All nonzero digits are significant Measurement 7.314 has four significant digits The number of significant digits is independent of the position of the decimal point 73.14 also has four significant digits Zeros located between nonzero digits are significant 60.052 has five significant digits  Use of Zeros in Significant Figures 1.4 The Numbers of Zeros at the end of a number (trailing zeros) are Measurement significant if the number contains a decimal point. 4.70 has three significant digits Trailing zeros are insignificant if the number does not contain a decimal point. 100 has one significant digit; 100. has three Zeros to the left of the first nonzero integer are not significant. 0.0032 has two significant digits  How many significant figures are in the following? 1.4 The Numbers of Measurement 1. 3.400 2. 3004 3. 300. 4. 0.003040  Scientific Notation 1.4 The Numbers of Used to express very large or very small Measurement numbers easily and with the correct number of significant figures Represents a number as a power of ten Example: 4,300 = 4.3  1,000 = 4.3  103  To convert a number greater than 1 to scientific notation, the original decimal point 1.4 The Numbers of is moved x places to the left, and the resulting Measurement number is multiplied by 10x The exponent x is a positive number equal to the number of places the decimal point moved 6200 = 6.2  103 What if you want to express the above number with three significant figures? = 6.20  103  To convert a number less than 1 to scientific 1.4 The Numbers of notation, the original decimal point is moved x places to the right, and the resulting number is Measurement multiplied by 10–x The exponent x is a negative number equal to the number of places the decimal point moved 0.0062 = 6.2  10–3  When a number is exceedingly large or small, 1.4 The Numbers of scientific notation must be used to input the number into a calculator: Measurement 0.000000000000000000000006692 g must be entered into calculator as: 6.692 x 10−24  Represent the following numbers in scientific notation: 1.4 The Numbers of Measurement 1. 0.00018 2. 3004 3. 300. 4. 0.00304  Types of Uncertainty 1.4 The Numbers of Error - the difference between the true value Measurement and our estimation – Random – Systematic Accuracy - the degree of agreement between the true value and the measured value Precision - a measure of the agreement of replicate measurements  Significant Figures in Calculation of Results 1.4 The Numbers of Rules for Addition and Subtraction The result in a calculation cannot have greater Measurement significance than any of the quantities that produced the result Consider: 9 37.68 liters 6.71864 liters 108.428 liters 152.82664 liters correct answer 152.83 liters  Report the result of each to the proper number of significant figures: 1.4 The Numbers of Measurement 1. 4.26 + 3.831 2. 8.321 − 2.4  Adding and Subtracting in Scientific Notation 1.4 The Numbers of There are two ways to solve the following: 9.47 x 10−6 + 9.3 x 10−5 Measurement SOLUTION 1: convert both numbers to standard form and add: 0.00000947 + 0.000093 0.00010247 correct answer 1.02 x 10−4  Adding and Subtracting in Scientific Notation 1.4 The Numbers of There are two ways to solve the following: 9.47 x 10−6 + 9.3 x 10−5 Measurement SOLUTION 2: change one of the exponents so that both have the same power of 10, then add. 9.47 x 10−6 changes to 0.947 x 10−5 0.947 x 10−5 + 9.3 x 10−5 10.247 x 10−5 correct answer 1.02 x 10−4  Rules for Multiplication and Division 1.4 The Numbers of The answer can be no more precise than the least precise number from which the answer is derived Measurement The least precise number is the one with the fewest significant figures 4.2 ´103 (15.94) -8 -4 = 2.9688692 ´10 (on calculator) 2.255 ´10 Which number has the fewest significant figures? 4.2  103 has only 2 The answer is therefore, 3.0  10-8  Exact and Inexact Numbers 1.4 The Numbers of Inexact numbers have uncertainty by definition Measurement Exact numbers are a consequence of counting A set of counted items (beakers on a shelf) has no uncertainty Exact numbers by definition have an infinite number of significant figures  Rules for Rounding Off Numbers 1.4 The Numbers of When the number to be dropped is less than 5 the preceding number is not Measurement changed When the number to be dropped is 5 or larger, the preceding number is increased by one unit Round the following number to 3 significant figures: 3.34966  104 =3.35  104  Round off each number to three 1.4 The Numbers of significant figures: Measurement 1. 61.40 2. 6.171 3. 0.066494  1.5 Problem Solving Problem Solving is much like planning a road trip. First, students must form a plan and then follow that plan through to answer the problem. Do not try to get any answer; try to get an answer that makes sense in the context of the problem.  1.7 Additional Experimental Quantities Temperature - the degree of “hotness” 1.3 The Units of Measurement of an object 1.7 Additional Experimental Conversions Between Fahrenheit and Celsius ToC = ToF − 32 11 Quantities 1.8 ToF = 1.8 x ToC + 32 1. Convert 75oC to oF 2. Convert -10oF to oC 1. Ans. 167 oF 2. Ans. -23oC 1.7 Additional Experimental Kelvin Temperature Scale The Kelvin scale is another temperature Quantities scale. It is of particular importance because it is directly related to molecular motion. As molecular speed increases, the Kelvin temperature proportionately increases. TK = ToC + 273.15 1.7 Additional Experimental Energy Energy - the ability to do work kinetic energy - the energy of motion Quantities potential energy - the energy of position (stored energy) Energy is also categorized by form: light heat electrical mechanical chemical 1.7 Additional Experimental Characteristics of Energy (start here) Energy cannot be created or destroyed Energy may be converted from one form to Quantities another Energy conversion always occurs with less than 100% efficiency All chemical reactions involve either a “gain” or “loss” of energy 1.7 Additional Experimental Units of Energy Basic Units: calorie or joule 1 calorie (cal) = 4.184 joules (J) Quantities A kilocalorie (kcal) also known as the large Calorie. This is the same Calorie as food Calories. 1 kcal = 1 Calorie = 1000 calories 1 calorie = the amount of heat energy required to increase the temperature of 1 gram of water 1oC. 1.7 Additional Experimental Concentration Concentration: – the number of particles of a substance – the mass of those particles Quantities – that are contained in a specified volume Often used to represent the mixtures of different substances – Concentration of oxygen in the air – Pollen counts – Proper dose of an antibiotic 1.7 Additional Experimental Density and Specific Gravity Density – the ratio of mass to volume 12 Quantities – an extensive property – use to characterize a substance as each substance has a unique density – Units for density include: g/mL g/cm3 mass m d= = g/cc volume V 1.7 Additional Experimental Quantities water liquid mercury cork brass nut 1.7 Additional Experimental Calculating Density A 2.00 cm3 sample of aluminum is found to weigh 5.40 g. Calculate the density in g/cm3 and g/mL. Quantities Plan your trip: Starting point is 2.00 cm3 and 5.40 g Destination is density The Road we need is d = m/V 5.40 g, 2.00 cm3 Density Formula g/cm3 1.7 Additional Experimental Calculations Using Density A 2.00 cm3 sample of aluminum is found to weigh 5.40 g. Calculate the density in g/cm3 and g/mL. Quantities Solution: Substitute the information given in the problem into the density formula. d = 5.40 g = 2.70 g/cm3 2.00 cm3 Since 1 cm3 = 1 mL, = 2.70 g/mL 1.7 Additional Experimental Calculations using Density Calculate the volume, in mL, of a liquid that has a density of 1.20 g/mL and a mass of 5.00 g. Quantities Plan your trip: Starting point is 5.00 g Destination is mL The Road we need is density = 1.20 g/mL 5.00 g 1.20 g/mL mL 1.7 Additional Experimental Calculating Density Calculate the volume, in mL, of a liquid that has a density of 1.20 g/mL and a mass of 5.00 g. Quantities Solution: The density can be written as a conversion factor 1.20 g 1 mL or 1 mL 1.20 g 5.00 g x 1 mL = 4.17 mL 1.20 g 1.7 Additional Experimental Density Calculations Air has a density of 0.0013 g/mL. What is Quantities the mass of 6.0-L sample of air? Calculate the mass in grams of 10.0 mL of mercury (Hg) if the density of Hg is 13.6 g/mL. Calculate the volume in milliliters, of a liquid that has a density of 1.20 g/mL and a mass of 5.00 grams. 1.7 Additional Experimental Specific Gravity Values of density are often related to a standard Specific gravity - the ratio of the density of the Quantities object in question to the density of pure water at 4oC Specific gravity is a unitless term because the 2 units cancel Often the health industry uses specific gravity to test urine and blood samples density of object (g/mL) specific gravity = density of water (g/mL)