Why Study Chemistry?? In this Chapter. States of Matter. Chemistry 141

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1 Why Study Chemistry?? Chemistry 141 In this Chapter What is Chemistry? Measurements The scientific notation Significant figures Unit conversions 2 States of Matter The 3 most common states of matter are: 1.Solids: Have a fixed shape and size. 2.Liquids: Occupy the shape of the container they are poured in. 3.Gases: No fixed shape or size and completely fill the container they are placed in. To distinguish between the different kinds of matter, we can determine and compare their properties. The two common types of properties are: 3 1

2 Physical Properties of Matter Physical properties: A characteristic of a substance that can be observed without changing the composition of the substance. (for example: the color of an object). Physical properties are subdivided into: Extensive: Depend on how much or the amount of the substance. (mass) Intensive: Independent of the amount of the substance. (temperature) 4 Chemical Properties of Matter These properties can cause a change in the chemical nature of the substance. Typically are independent of the amount of the substance. Formation of rust on a piece of iron left outside is an example. Matter can be further classified as: 5 Classification of Matter Matter (Has mass and takes up space) 3 states Solid Liquid Gas Pure substances (unique properties, fixed composition, can t be purified further) Mixtures (physical combinations of 2 or more pure substances) Compounds Homogeneous (uniform composition) Heterogeneous (nonuniform composition) Elements 6 2

3 Compounds Formed by the combination of 2 or more elements in a fixed proportion. The law of definite proportions states that the elemental composition of a pure compound is always the same regardless of its source. Can be broken down into simpler substances (elements) by chemical means. Water and carbon dioxide. 7 Elements 118 known elements present in a wide range of abundances. Considered to be chemically pure. Elements are usually represented by a chemical symbol derived from the element s name. The periodic table arranges all the known elements in a certain order. 8 Making measurements Science is based on making 2 types of observations: qualitative and quantitative. Qualitative observations are descriptive and normally do not use numbers. Quantitative observations are numerical and use a number and a unit. 9 3

4 Units 5000 years ago in ancient Egypt the cubit was the unit for measuring length. We need a standard, universally acceptable, set of units to describe our measurements. All acceptable measurements should have a number (85 in this case) associated with a unit ( F). Units make any measurement meaningful. 10 Systems of Measurements Here in the US we still use what is known as the English system of measurements. The rest of the world (including England) uses what is known as the SI system. The International System of Units is a system of standard units of measurement for science based on the metric system. 11 The Metric and the SI Systems Developed after the French revolution in about Based on the decimal system. Uses multiples of 10 to indicate numbers that differ by orders of magnitude from a basic unit. The metric unit has 2 parts: a base unit and a prefix. The prefix tells us how many times to divide or multiply the base unit by

5 Fundamental Units The following are considered as the fundamental or base units: 13 The Scientific Notation Used to express very large or very small numbers. There are two terms involved in any number expressed by scientific notation. Consider the number: 1.23 x Large Numbers by Scientific Notation The distance between the Earth and Mars is 55, meter. In scientific notation this number is written as: 5.5 x meter; this means that if we multiply 5.5 by 10, ten times we will get the original number. Here the decimal point has been moved to the left till we have one non-zero number to the left of the decimal point. The rest of the number is expressed as a multiple of ten with the exponent of 10 = number of places we move the decimal point. 15 5

6 Small Numbers by Scientific Notation Very small numbers are treated exactly opposite. The common cold virus has a length of meter. This number can be written as: 8.75 x 10-7, this is the same as dividing the number 8.75 by 10 seven times. Here the decimal point is moved to the right by 7 places and the exponent of 10 is expressed as negative seven. Useful for significant figures. 16 SI system and Prefixes PRINT THE SEPARATE PAGE FROM BB AND HOLD ON TO IT 17 Conversion Factors Do not make up your own conversion factors!! Most conversion factors are tabulated. Always remember to use the conversion factor so that the desired units are on top and are NOT cancelled. Always setup so that: New units Old units = New units Old units 18 6

7 Convert 3.58 x 10-8 pm to m Conversion factor to go from to. Is the answer going to be a large or a small number? Convert 19 Convert Convert x Yg to g Convert 9.56 x 10-5 TL to ml 20 Significant Figures (Sig figs) Significant figures tell you how precise a certain measurement is. Rules for significant figures: Read the digits from left to right. Start with the first non-zero number, count all digits x x x

8 Special Rules for Zeroes Zeroes at the beginning of a number (called leading zeroes) are NOT significant has 3 significant figures. Zeroes between non-zero digits are ALWAYS significant has 5 significant figures. Zeroes at the end of a number (trailing zeroes) containing a decimal point are ALWAYS significant has 5 significant figures, while has only 3 significant figures. 22 Exceptions to Rules for Significant Figures Exceptions to the previous rules include Conversion factors Physical constants Special numbers like π and e Significant figures will change with any mathematical operations being performed. In multi-step calculations carry all your sig figs do not round off till the very end. 23 Exact Numbers A number that arises when items are counted or units/conversion factors are defined. 45 students 12 eggs 1 inch = 2.54 cm In a calculation ignore the number of significant figures in an exact number. 24 8

9 Addition/Subtraction Number of decimal places in the final answer stays the same as the number with the lowest number of decimal places. Consider: 25 Multiplication/Division Number of sig figs in the answer = number of sig figs in the term with the lowest number of sig figs. Consider: 26 Rounding off Round off 1.46, 1.44, 1.45 and 1.35 to 2 sig figs: 1.46 is rounded off to The last digit that is retained is increased by 1, if the digit to the right of it is greater than is rounded off to Left unchanged, if the digit to the right of it is less than is rounded off to If the digit to its right is 5, then leave the retained digit unchanged if it is even is rounded off to If the retained digit is odd, increase it by one. 27 9

10 Strategy for Problem Solving Identify where you want to go. Choose the right starting point. List the correct form of the conversion factors. Multiply the starting measurement through by the conversion factors. DOES THE ANSWER MAKE SENSE? 28 Conversion Factors Do not make up your own conversion factors!! Most conversion factors are tabulated. Always remember to use the conversion factor so that the desired units are on top and are NOT cancelled. Always setup so that: 29 Intersystem Conversions How many km is the Indy 500? The Indy 500 has exactly miles

11 Cell phone problem In the month of July you used your cell phone for 4050 minutes. Your bill was $ How much did you pay per minute? How many hours did you spend each day using your cell phone? 31 Swimming Pool Problem We need to find the volume required here by multiplying the dimensions of the pool. Thus: Volume = 20 ft x 30ft x 6ft = 3600 ft 3 Since we are multiplying three quantities in feet the answer here has the units ft 3. Now we have our starting point: 32 Convert the cubic feet to gallons by using the conversion factors: 1 foot = 12 inches, 1 inch = 2.54cm, 1000 ml = 1L and 3.78 L = 1 gallon. This is a slightly tricky setup and keep track of how the units change here. Also notice that I have cubed some conversion factors so they cancel and it is perfectly acceptable to do this. Remember 1cm 3 = 1mL

12 Common Experimental Quantities Mass or weight: Describes the quantity of matter in an object. Mass and weight are used interchangeably even though there is a slight difference between the 2. Weight = Mass x gravitational acceleration. The weight of an object will be different on the Earth and the Moon while the mass will stay the same. 1 gram = 10-3 kg = 1/454 pound (lb). 34 Common Experimental Quantities Length: The distance between 2 points. 1 meter (m) = 100 cm = 39.4 inch Volume: The space occupied by an object. 1 liter = 1000 milliliters (ml) = 1.06 quarts. A graduated cylinder and a volumetric flask are examples of common laboratory glassware used for measuring volume. Temperature: Measures the amount of heat contained in an object or the hotness of an object. Three common units that are based on the boiling and freezing temperatures of water are:

13 Temperature Conversion o From F to C o o F 32 F C = 1.8 Convert 80 F to C o o 80 F 32 F o C= = 26.6 F 1.8 o = 27 C (rounding off) From C to F: F = 1.8 C + 32 From C to Kelvin: K = C Density The density of an object is defined as the ratio between its mass and its volume and is given by the formula: density density volume mass (m) (d) = The density of an volume (v) object also volume = mass decides if an mass object will float or = density sink in a liquid. 38 Density Density is an intensive property (independent of the amount of the substance) and is a distinguishing characteristic of every substance. The most common units for density are: g/ml or g/cm 3 or g/cc where 1 ml = 1 cm 3 = 1 cc (cubic centimeter) Read as gram per milliliter. Might see this written as g.ml

14 How Good are My Numbers? Errors: Every measurement has an error that accompanies it. Precision and accuracy indicate the error. Precision: Measures how well same measurements agree with each other. Accuracy: How close is the number to the actual value? (Need to know the actual value of the number). ACCURACY AND PRECISION ARE NOT THE SAME!!! 40 Precision vs. Accuracy 41 Precision vs. Accuracy 42 14

15 Precision vs. Accuracy 43 Chapter Summary Chemistry is the study of matter and energy. Substances have unique physical and chemical properties such as density, temperature etc. Expressing measurements properly, needs the use of the SI system, units and significant figures. Must be able to convert between systems of units and handle scientific notation