1.1 What is Chemistry? Chemistry can be defined as the science that deals with the materials of the universe and the changes that these materials undergo and the energy associated with those changes. Chemistry is often called the central science that is because most of the phenomena that occur around us involved chemical changes, in other words changes where one or more substance becomes different substance. Examples 1) Eggs, flour, sugar and baking powder are mixed together to yield a cake. 2) Wood burns in air, forming water, carbon dioxide and other substances. 3) The steel in a car rusts. 1.2 Measurements and Calculations In science; making observations is a key part of the scientific process, and the observation can be divided into two parts as in (Figure 1): Observations Qualitative Quantitative This mean how we see the substance is it yellow or red? and is it solid or liquied? This mean how much the substance weight? is it 1 gram or 2 grams? Quantitative observations is called Measurements Figure 1 1
* To be understandable; measurement always consists of two parts Number Unit a) The number part of a measurement tells us what is the value we have (numerical value). b) The unit part of the a measurement tells us what scale or standard is being used to represent the results of the measurement see (Figure 2). Examples: 10 cm 3 Unit 3 centimter (cm) Number Number Unit Figure 2 People always required common or standard units to make the practical life around the world possible. However, standard systems of units did arise but unfortunately; different systems were adopted in different parts of the world. The two most widely used systems are the English system used in the United State of America and the metric system used in most of the rest of the world; however the metric system is preferred for most scientific work. In 1960 an international agreement set up a system called the International System (abbreviated SI) [Le Système International in French], the SI unit system is based on the metric system so that the SI unit system considered as an expanded modernized version of the metric system. In chemistry, all measurements are made in the metric system or in the IS system. The International System of Units SI is the most widely used system of units. It is the most common system for everyday use in the world, and is almost universally used in sceince. 2
1.3 Fundamental Quantities & Fundamental Units. (Base) Fundamental quantities and the fundamenta units are such that every other quantities and unit can be generated from them. There are seven fundamental quantities and each one has a fundamental or base unit, they are listed in table (1) Fundamental Quantities Fundamental Unit Length [L] Mass [M] Time [T] Absolute temperature [θ] Electric current [A] Luminous intensity [I] Amount of substance [n] meter [m] kilogram [kg] second [s] kelvin [K] ampere [A] candela [cd] mole [mol] Table (1) Fundamental quantities and fundamental units 1.4 Derived Quantities & Derived Units The derived quantities and derived units are these quantities and units that we can generate them from any of the seven fundamental quantities and fundamental units respectively. For example; Volume is derived quantity. The volume is defined as the amount of three-dimensional space occupied by a substance (Figure 3). Volume = length x width x height fundamental quantity derived quantit m 3 = m x m x m fundamental unit derived unit h l w Figure 3 3
1.5 What are the Prefixes? Why we do need to use Prefixes? The fundamental units are not always a convenient size, the SI system uses prefixes to change the size of the unit from large size to small size or from the small size to large size that all depends on what we are dealing with. Some Prefix Used in the SI System Prefix giga mega kilo hekto Deka Deci Centi Milli Micro Nano Abbreviation Decimal Expression Exponential Expression G 1.000.000.000 10 9 M 1.000.000 10 6 k 1.000 10 3 h 100 10 2 da 10 10 1 d 0.1 10-1 c 0.01 10-2 m 0.001 10-3 μm (or mc) 0.000001 10-6 n 0.000000001 10-9 Table (2) 1.7 Exponential Numbers & Common Numbers 1.7.1 Exponential Numbers The numbers we are working with in the scientific measurements are often very large or very small; thus it is convenient to express them using powers of 10 this simplification known as exponential numbers. This method involves the use of a base number 10 raised to some power like base 10 X power or exponent 4
Please note that a) Power or exponent indicates how many time the base number (10) is repeated. 10 5 = 10 x 10 x 10 x 10 x 10 = 100 000 [10 5 means 10 is repeated 5 times] 5 x10 2 = 5 x (10 x 10) = 500 [10 2 means 10 is repeated 2 times] 1 x 10 9 = 1 x 10 x 10 x 10 x 10 x 10 x 10 x10 x 10 x10 = 1000 000 000 10 1 = 10 [10 1 means 10 is repeated only 1 times]. b) Power or exponent (x) can be zero or positive or negative numbers. 10 0 & 10 5 & 10-5 c) If x is zero the product always equal 1 10 0 = 1 d) If x is positive number the product is always bigger than 1 10 2 = 100 which is bigger than 1 e) If x is negative number the product is always less than 1 10-1 = 1/10 = 0.1 which is less than 1 f) A negative exponent indicates the reciprocal of the same number with a positive exponent 10 3 10-3 BUT 10-3 = 1/10 3 = 1/10 x 1/10 x 1/10 = 1/1000 = 0.001 1.7.2 Common Numbers Common numbers are these numbers which has the form of numbers without any exponent as example: 2300, 4567, 84000, 129800, 2.3000, 8.230 1.7.3 How to Change Exponential Numbers to Common numbers If the power or exponent is POSITIVE all what you have to do is to move the decimal point to the right from the first number after the decimal point.. Example 1: Change the exponent number 3.4 x 10 3 to common number. Answer This is positive exponent with value 3, so to convert it to common number we have to move the decimal point to the right 3 places. So 3.4 x 10 3 = 3400 **************************** 5
But If the power or the exponent is NEGATIVE, all what you have to do is to move the decimal point to the left from the first number after the decimal point. Example 3 Change the exponent number 3.4 x 10-3 to common number Answer This is negative exponent with value -3, so to convert it to common number we have to move the decimal point to the left 3 places. So 3.4 x 10-3 =.0034 Example 3 Change the exponent number 2.45 x 10-4 to common number Answer This is negative exponent with value -4, so to convert it to common number we have to move the decimal point to the left 4 places. So 2.45 x 10-4 =.000245 **************************** 1.7.4 How to Change Common Numbers to Exponential Numbers To change from an exponential number to a common number, the decimal point you will need to put will only have one digit on its left side. Notice that if you move the decimal point to the left side the exponential number will be positive, and if you move the decimal point to the right exponential number will be negative. Example1 Change 5000 to an exponential number. Answer To have only one digit on the left of a decimal point we have to move the decimal point three places to the left side So 5000.0 = (5000) = 5 x 10 3 6
1.8 How to Convert from Large Unit to Small Unit and vies versus There are two simple methods to perform this easy operations. The first is direct one and the second are indirect method, however both of them are simple so feel free to use any one you think it is better for you. However to apply any of the methods you must know all the units and the relationship between them, that mean for example you must know that 1Kg = 1000g (10 3 g) 1Km = 1000m (10 3 m) 1 L = 1000ml (10 3 ml) 1.8.1 To convert from large unit to small unit 1.8.1.1 The direct method To convert from large unit to small unit all what you have to do is to multiply the amount you have (number) with the suitable prefix (Table 2, page 5). Examples: 1) Convert 5 kg to its g counter part. As we know kg = 1000 g So 5 kg = 5 x 1000 = 5000 g (5 x 10 3 )g 2) Convert 10 L to millilitres ml, As we know 1 L = 1000 ml So 10 L = 10 x 1000 = 10000 ml (1 x 10 4 )ml 3) Convert 3 Km to its m counter part As we know 1 km = 1000 m So 3 km = 3 x 1000 = 3000 m (3 x 10 3 )m 1.8.1.2 To convert from small unit to large unit To convert from large unit to small unit all what you have to do is to divide the amount you have (number) on the suitable prefix (Table 2, page 5). Example 1 Convert 5 g to its kg counter part Answer: Kg = 1000 g so 5g = 5/1000 = 0.005 OR you can multiply with the 10-3 that mean 5 x 10-3 = 0.005 1.8.1.3 The indirect method The indirect method is simple way as well and it depends on making an equation. 7
Example 1 Convert 5 kg to its g counter part. 1kg 5kg 1000g 5kg x 1000g = 1kg x Y 5kg x 1000g = 1kg x Y 1kg l kg 5kg x 1000g = 1kg x Y 5000 g = Y 1kg 1kg Example 2 Convert 10 L to millilitres ml. 1L 10 L Y 1000 ml This means 10 L x 1000 ml = 1L x Y 10 L x 1000ml = 1L x Y 1L l L Y = 10 x 1000 = 10000 ml (1 x 10 4 ) ml 1.9 Uncertainty in Measurement and Significant Figures Whenever a measurement is made, an estimate is always required that is because there is no 100% accuracy in any measurement. However there is no 100 % accurate measurement because the measurement process is affected by few factors like the type of the device we are using, the human who is doing the measurement (skill of the operator) and the scale we use. This can be illustrated by the following example. If I asked 5 students to measure the length of the black arrow using a ruler like in figure 4. Figure 4 8
The results might be as the following Student 1 2 3 4 5 8.67 8.66 8.65 8.64 8.63 Results of measurement Please notice that the first two digits in each measurement are the same regardless of who made the measurement; these are called the certain numbers of the measurement. However the third digit is estimated and it vary from student to another, this digit called uncertain number. 8.67 uncertain number certain numbers It is very important to realize that a measurement always has some degree of uncertainty. The numbers record in a measurement (all the certain numbers plus the first uncertain number) are called Significant Figures 8.67 certain numbers uncertain number Significant Figures 1.9.1 Accuracy and Precision The accuracy of the measurement refers to how close the measured value is to the true or accepted value. For example, if you used a balance to find the mass of a known standard 100.00 g mass, and you got a reading of 75.95 g, your measurement would not be very accurate. Precision refers to how close together a group of measurement actually are to each other. Precision has nothing to do with the true or accepted value of a measurement. 9
For example: if you used a balance to find the mass of a known standard 100.00 g mass, and you got a reading of 75.95 g, then you have repeated the measuring process five times and you obtained the following readings 75.96, 75.97, 75.98, 75.99 and 76.0; we say that your measurements are precise that because all the six measurement you have done are very closed to each other. In the same time we say your measurements are not accurate that because the values you obtained are far away from the true value. measured value the true value measured value the true value Not accurate Not precise Not accurate But precise the true value measured value accurate and precise It is quite possible to be very precise and totally inaccurate (as in the above example). In many cases, when precision is high and accuracy is low, the fault can lie with the instrument. If a balance is not working correctly, they might consistently give inaccurate answers, resulting in high precision and low accuracy. 1.10. Scientific Notation and Significant Figures 1.10.1 Scientific notation is a short way to write very large or very small numbers. For example, the distance from the earth to the sun is about 93000000 miles. Of course it is not exactly 93000000 miles. We say it has been measured as 93000 000 miles. Instead of writing 93000000 (common numbers), we use a shorter way which is (exponential numbers). This way is known as scientific notation. 93 000 000 is entered as 9.3 Exp 7. Sometimes you will see it written as 9.3e7 Scientific Calculators usually use either the EE or Exp key for scientific notation. 10
1.10.2 Significant Figures Is the number of meaningful digits in a measured or calculated quantity. How to Determining the Number of Significant Figures? There are 6 rules to determine the number of significant figures 1) All digits that are non zero are significant, regardless of the location of the decimal point. Examples: 1.59 contains three significant figures 15.9 contains three significant figures 3.2674357 contains eight significant figures 1456 contains four significant figures 276387 contains six significant figures 2) Zeros between nonzero digits are significant figures Examples: 100.24 contains five significant figures 101.0354 contains seven significant figures 20.4.1 contains three significant figures 3) Zeros to the left of the decimal point or to the left of a number are not Significant figures. Examples: 0.517 contains three significant figures 0.0039 contains two significant figures 4) Zeros to the right of a decimal point are significant if they are at the end of a number Example: 59.0 contains three significant figures 16.0000 contains six significant figures 760 (contains two significant figures) is written as 7.6 x 10 2 5) When a number ends in zero or zeros that are not to the right of a decimal point, the end zero or zeros may or may not be significant. Thus, 760 may or may not contain three significant figures, that depends on the accuracy of the measurement. However to make it clear, such as this numbers should be written by using of scientific notation (short way). 11
1.11 Density Density is an important property of matter, it gives information about the physical properties of any matter. The density can be defined the amount of matter present in a given volume of substance i.e. (density is mass per unit volume) and it can be expressed by the formula D = M / V Where D is the density, M is the mass and V is the volume. Remember that density is Derived Quantities and its unit is Derived Units The most common unit for density is g/ml which equal to g/cm 3 that is because 1mL = 1 cm 3 D = M / V Remember the most common unit for measuring the density is g/ml which equal to g/cm 3 that is because 1mL = 1 cm 3. M D V Example 1) What is the density of the mercury in a thermometer if 40 g of it occupies 2 ml? D = M/V 40/2 = 20 g/ml. 2) Alcohol has a density of 0.80 g/ml. How much will 100 ml of is weight? D = M/V M = D x V 0.80 g/ml x 100 ml = 80 g 3) Alcohol has density of 0.80 g/ml, and its weight is 80 g, what is its volume? D = M/V V = M/D 80g = 100mL 0.8g/mL. 12
1.13 Temperature Scales There are three different temperature scales 1) Fahrenheit scale: This scale is widely used in the united states and Britain. Fahrenheit scale is employed in most of the engineering science. 2) Celsius scale: This scale is employed in the physical and life sciences. 3) Kelvin scale: Is another temperature scale used in the science. In chemistry, the Celsius (C) and Kelvin ( K) temperature scales are commonly used. How are the three temperature scales are compare with one another? To understand the relation between the three different temperature scales we need to run a very simple experiment and note what we will see (Figure 7). Put three different thermometers i.e. one Fahrenheit thermometer, one Celsius thermometer and one Kelvin thermometer inside a glass contain mixture of water and ice (Figure 7a). On the Fahrenheit thermometer you will see that the temperature reading is 32 ºF On the Celsius thermometer you will see that the temperature reading is 0 ºC On the Kelvin thermometer you will see that the temperature reading is 32 K. Fig. 7 All of the above temperatures readings are indicate the freezing point of water. They also indicate the melting point of ice. 13
Next if you have changed the ice-water mixture with boiling water; the thermometers reading will be as the following (figure 7b) On the Fahrenheit thermometer you will see that the temperature reading is 212 ºF On the Celsius thermometer you will see that the temperature reading is 100 ºC On the Kelvin thermometer you will see that the temperature reading is 373 K. These are temperatures indicate the boiling point of water. In your study of chemistry, you will some times need to convert from one temperature scale to another. How can we do such as these changing? To convert from Celsius temperature to corresponding Fahrenheit you need the following formula ºF = (9/5 ºC) + 32º Where ºF is the temperature in Fahrenheit ºC is the temperature in Celsius In the other hand if you need to convert from Fahrenheit temperature to Celsius you would need the following formula Where ºF is the temperature in Fahrenheit ºC is the temperature in Celsius ºC = 5/9 (ºF 32) The convert from Celsius to Kelvin and vies versus you need to use the following formula. K = ºC + 273 Where K is the temperature in Kelvin ºC is the temperature in Celsius Please note that K scale is given as symbol only without the degree sign. Please do not hesitate to ask if you have any question 14
1.14 Matter What is Matter? Matter is any thing that occupies space and has weight (mass). Every thing we see or feel is matter!!! As examples is stars in the sky, cars in the street, chairs we set on, gasoline we put in our cars, the food we eat so many other things. The above examples represent matter that we can see and/or feel. 1.14.1 States of Matter There are three states of matter 1) Solid 2) Liquid 3) Gas Fig 8. The three states of matter To understand the difference between the three states of matter, let us consider that matter is made out of particles. Let us see how these particles are arranged in the three different states of matter. 1.14.2 Solid In solids, strong forces hold the particles together and the particles are arranged into a definite, rigid shape. Because of the strong force that hold the particles in the solid materials; the motion of the particles are little and highly restricted. If heat were applied on the solid it caused the particles to move but very slightly, however the particles of a matter in the solid state is in continues vibration Figure 9.a. 15
1.14.3 Liquid Fig.9 Arrangement of particles in the solid (a) Liquid (b) and gas(c) In the liquid state of matter the forces which hold the particles are weaker than these in solid. That mean the particles are little far away from each other figure 9.b. In liquids the particles are moderately ordered, and the particles are not strongly hold each other so that liquids are flow and have no definite structure but they take the shape of the vassal they are in. Liquids are expanded when they are heated. 1.14.4 Gas In gas state; the force which holds the particles are so weak figure 9c, so that the particles are far away from each other. Gases take the shape and the volume of any container they exist in State Definition Examples Solid Has a definite volume, highly ordered, has fixed shape and volume, does not flow, expand slightly when heated. Ice cubs, diamond, iron bar Liquid Has a definite volume but it takes the shape of its container, flow, moderately ordered. Gasoline, water, alcohol, blood, mercury. Gas Has no definite volume or shape; takes the shape of and the volume of its container, flow, the particles are far apart of each other Air, helium, oxygen 16
1.15 Physical & Chemical Changes Any matter and at any of its three states (solid, liquid and gas) has properties by which we can classify and recognized these matters. There are two types of properties we need to learn which are; Physical and Chemical properties. Physical properties: are these properties that can be measured and or observed with out the need or use of chemical reactions, in the physical properties there is no production of a new material. For examples are colour, odor, density, hardness, solubility, melting point, and boiling point. Chemical properties are determined by the reaction of a substance with other substances i.e. a new materials are produced. Examples of chemical properties are combinations with acids and bases to form salt and water. Reaction of active metals, with oxygen to form metal oxides. Burning a piece of wood, to produce carbon dioxide and water. Properties of any matter can usually be changed. If the changing takes place in the physical properties we call this change as Physical change. If the changing takes place in the chemical properties of the matter we call this change as Chemical change. 1.15.1 Physical Change: Is one in which no new substance is produced, although there may be a change of the state or density or colour or all of them but the chemical composition remains the same. For example; cube of ice has changed to water when it was heated up. The water still as its water (H 2 O), that means no new substance was produced but all what happened is the state of the matter has changed from solid to liquid. Examples on the physical changes of the matter. 1) Dissolving a sugar in water. 2) Breaking a piece of glass 3) Freezing water. 1.15.2 Chemical Change: In a chemical change, which is often called a chemical reaction, the atoms of a substance are rearranged. A chemical change requires that the new substance have a chemical composition that is different from the composition of the original substance. For example; if you pass an electric current in a water (do not try to do this experiment by your self), both H 2 and O 2 gases will be produced. These two gases are completely different from the starting material H 2 O. 17
Some matters can changed from its solid state to the gas state without becoming a liquid. As example of this case; is the dry ice (solid CO 2 ) which is changing to CO 2 gas without passing through the liquid state. This phenomenon is called Sublimation. 1.15.3 Comparison of Physical and Chemical Properties A physical property tells what a substance is is it white or it is red; it is odourless or it has strong odor; it is hard or it is soft i.e. (solid, liquid, gas). A chemical property tells what a substance does- it is burns or it does not burn; it reacts with another substance or it does not. 1.16 Intensive properties & Extensive properties 1.16.1 Intensive properties: A property that does not change when the amount of sample changes. Examples are density, pressure, temperature and colour. 1.16.1 Extensive properties A property that changes when the amount of matter in a sample changes. Examples are mass, volume, length, and charge. 1.17 Composition of Matter All matter can be divided into two major classes which are pure substance and mixture. The pure substance could be Element or could be compound. The mixture can be either Homogeneous mixture or heterogeneous as in figure 6. Matter solid, liquid gas Pure Substance constant composition Physical methods Separating mixture Preparing mixture Mixture variable composition Elements Compounds Homogeneous Heterogeneous Metals Non -metals Semimetals (metalloids) Fig.10 Composition of Matter 18
1.17.1 Element: is a substance that can not be broken down into simpler substance by chemical methods and consist of only one type of atom. Examples of elements are iron (Fe), Aluminium (Al), Oxygen (O), Hydrogen (H). There are about 117 confirmed elements, and they were gathered and classified in a chart called the Periodic Table. Elements can be classified into three types: Metals, Non-metals and Semimetals (metalloids). Each type has its own specific properties. (N.P. We will study the periodic table in other section). 1.17.2 Compounds Certain elements have special affinities for each other. They bind together in special ways to form Compounds. Compounds are a substance that can be broken down into simpler substance (pure components elements) only by chemical methods. Compounds are formed from at least two elements. Compounds can be organic compounds i.e. contain C and H for example CH 3 COOH, CH 3 COCH 3, or can be inorganic compounds like HCl, NH 3, H 2 O. Compounds have the following characteristics 1) They can be separated into their component substances by chemical means, and they are homogeneous in composition. 2) They have a definite proportion by weight of the substance from which they were made. 3) They have different properties from those of substance from which they were made. 1.17.3 Mixture: A mixture is a combination of two or more substances that are not chemically united and do not exist in fixed proportions to each other Mixtures: consist of various compounds and/or elements, with no specific formula have variable composition with either definite or varying properties depending on the sample can be broken down into individual components by physical methods example: any alloy like brass, steel, gold; Mixtures can be either homogeneous or heterogeneous. 19
1.17.4 Homogeneous Mixture: The prefixes "homo"- indicate sameness A homogeneous mixture has the same uniform appearance and composition throughout. Many homogeneous mixtures are commonly referred to as solutions as example dissolving sugar in H 2 O. 1.17.5 Heterogeneous Mixture: The prefixes: "hetero"- indicate difference. the individual components of a mixture remain physically separated and can be seen as separate components, for example mixing sand with H 2 O. Selected Problems: 8, 9, 11, 18, 20, 22, 24, 26, 28, 30, 39, 41, 46, 48, 58. 20