Chapter 1 Introduction: Matter and Measurement Lecture Presentation Chapter 1 Introduction: Matter and Measurement James F. Kirby Quinnipiac University Hamden, CT

Matter and measurents Matter Numbers mixtures and pure substances homogenous and heterogeneous mixtures compounds and elements chemical and physical changes physical and chemical properties Numbers Accuracy and precision Significant figures Rounding off Scientific notation Metric units Dimensional analysis

Chemistry The study of the composition and structure of materials and of the changes that materials undergo.

Chemistry Chemistry is the study of the properties and behavior of matter. It is central to our fundamental understanding of many science-related fields.

Scientific Method Results Hypothesis Experiments

Scientific Method

Experiment Hypothesis Theory Law An observation of natural phenomena carried out in a controlled manner so that the results can be duplicated and rational conclusions obtained. Hypothesis A tentative explanation of some regularity of nature. Theory A tested explanation of a basic natural phenomenon. Law A concise statement or mathematical equation about a fundamental relationship or regularity of nature.

Basic laws about matter Law of conservation of mass In an ordinary chemical reaction, mass is neither lost nor created Law of constant proportions A chemical compound always contains exactly the same  proportion of elements by mass

Matter Matter is anything that has mass and takes up space.

Matter Atoms are the building blocks of matter. Each element is made of a unique kind of atom. A compound is made of two or more different kinds of elements. Note: Balls of different colors are used to represent atoms of different elements. Attached balls represent connections between atoms that are seen in nature. These groups of atoms are called molecules.

Methods of Classification State of Matter Composition of Matter

States of Matter

States of Matter The three states of matter are solid. liquid. gas. In this figure, those states are ice, liquid water, and water vapor.

Classification of Matter Based on Composition If you follow this scheme, you can determine how to classify any type of matter. Homogeneous mixture Heterogeneous mixture Element Compound Matter And Measurement

Classification of Matter—Substances A substance has distinct properties and a composition that does not vary from sample to sample. The two types of substances are elements and compounds. An element is a substance which can not be decomposed to simpler substances. A compound is a substance which can be decomposed to simpler substances.

Compounds and Composition Compounds have a definite composition. That means that the relative number of atoms of each element that makes up the compound is the same in any sample. This is The Law of Constant Composition (or The Law of Definite Proportions).

Classification of Matter—Mixtures Mixtures exhibit the properties of the substances that make them up. Mixtures can vary in composition throughout a sample (heterogeneous) or can have the same composition throughout the sample (homogeneous). Another name for a homogeneous mixture is solution.

*classify the following as elements, compounds, heterogenous mixtures or homogenous mixtures baking soda sand sugar chlorine iodized salt bronze copper brass air blood milk salad steel wool honey spring water

Types of Properties Intensive Properties are independent of the amount of the substance that is present. Examples include density, boiling point, or color. Extensive Properties depend upon the amount of the substance present. Examples include mass, volume, or energy.

Indicate which of the following are intensive or extensive properties Density Volume Mass Specific volume Temperature Moles Viscosity Thermal conductivity Color Electrical resistance Odor Luster Charge Pressure Malleability Ductility

Types of Properties Physical Properties can be observed without changing a substance into another substance. Some examples include boiling point, density, mass, or volume. Chemical Properties can only be observed when a substance is changed into another substance. Some examples include flammability, corrosiveness, or reactivity with acid.

which of the following are chemical or physical properties Melting point Flammability Viscosity Boiling Souring of milk Rusting of iron Food digestion Taking a bite of food Deployment of a car airbag Breaking of glass

Types of Changes Physical Changes are changes in matter that do not change the composition of a substance. Examples include changes of state, temperature, and volume. Chemical Changes result in new substances. Examples include combustion, oxidation, and decomposition.

Changes in State of Matter Converting between the three states of matter is a physical change. When ice melts or water evaporates, there are still 2 H atoms and 1 O atom in each molecule.

Chemical Reactions (Chemical Change) In the course of a chemical reaction, the reacting substances are converted to new substances. Here, the elements hydrogen and oxygen become water.

Separating Mixtures Mixtures can be separated based on physical properties of the components of the mixture. Some methods used are filtration. distillation. chromatography.

Filtration In filtration, solid substances are separated from liquids and solutions.

Distillation Distillation uses differences in the boiling points of substances to separate a homogeneous mixture into its components.

Chromatography This technique separates substances on the basis of differences in the ability of substances to adhere to the solid surface, in this case, dyes to paper.

Numbers and Chemistry Numbers play a major role in chemistry. Many topics are quantitative (have a numerical value). Concepts of numbers in science Units of measurement Quantities that are measured and calculated Uncertainty in measurement Significant figures Dimensional analysis

Units of Measurements—SI Units Système International d’Unités (“The International System of Units”) A different base unit is used for each quantity.

Units of Measurement—Metric System The base units used in the metric system Mass: gram (g) Length: meter (m) Time: second (s or sec) Temperature: degrees Celsius (oC) or Kelvins (K) Amount of a substance: mole (mol) Volume: cubic centimeter (cc or cm3) or liter (l)

Units of Measurement— Metric System Prefixes Prefixes convert the base units into units that are appropriate for common usage or appropriate measure.

Mass and Length These are basic units we measure in science. Mass is a measure of the amount of material in an object. SI uses the kilogram as the base unit. The metric system uses the gram as the base unit. Length is a measure of distance. The meter is the base unit.

Volume Note that volume is not a base unit for SI; it is derived from length (m × m × m = m3). The most commonly used metric units for volume are the liter (L) and the milliliter (mL). A liter is a cube 1 decimeter (dm) long on each side. A milliliter is a cube 1 centimeter (cm) long on each side, also called 1 cubic centimeter (cm × cm × cm = cm3).

Temperature In general usage, temperature is considered the “hotness and coldness” of an object that determines the direction of heat flow. Heat flows spontaneously from an object with a higher temperature to an object with a lower temperature.

Temperature In scientific measurements, the Celsius and Kelvin scales are most often used. The Celsius scale is based on the properties of water. 0 C is the freezing point of water. 100 C is the boiling point of water. The kelvin is the SI unit of temperature. It is based on the properties of gases. There are no negative Kelvin temperatures. The lowest possible temperature is called absolute zero (0 K). K = C + 273.15

Temperature The Fahrenheit scale is not used in scientific measurements, but you hear about it in weather reports! The equations below allow for conversion between the Fahrenheit and Celsius scales: F = 9/5(C) + 32 C = 5/9(F − 32)

Density Density is a physical property of a substance. It has units that are derived from the units for mass and volume. The most common units are g/mL or g/cm3. D = m/V

Numbers Encountered in Science Exact numbers are counted or given by definition. For example, there are 12 eggs in 1 dozen. Inexact (or measured) numbers depend on how they were determined. Scientific instruments have limitations. Some balances measure to ±0.01 g; others measure to ±0.0001g.

Uncertainty in Measurements Different measuring devices have different uses and different degrees of accuracy. All measured numbers have some degree of inaccuracy.

Accuracy versus Precision Accuracy refers to the proximity of a measurement to the true value of a quantity. Precision refers to the proximity of several measurements to each other.

Significant Figures The term significant figures refers to digits that were measured. When rounding calculated numbers, we pay attention to significant figures so we do not overstate the accuracy of our answers.

Recording the proper number of significant figures **the digits in a measurement that are known with certainty plus one digit that is uncertain**

Graduated pipette Graduated cylinder Buret Thermometer

Significant Figures in reported data All nonzero digits are significant. Zeroes between two significant figures are themselves significant. Zeroes at the beginning of a number are never significant. Zeroes at the end of a number are significant if a decimal point is written in the number or they are under a bar (to show significance).

Significant Figures 1. All non-zero digits count 1,234,456,789 2. All leading zeroes do not count 0.000 123 456 3. Confined zeroes do count 0.123 024 4. Trailing zeroes: -With a decimal, all trailing zeroes count 0.125 000 -In the presence of bars, zeroes under the bar count 36,020,100 -in the absence of bars and decimal, zeroes do not count 36,020,100

Rules for mathematical operations Rounding off 1. If the 1st digit to be dropped is less than 5, then that digit and all digits that follow are dropped 15.123 rounded to 3 sig figs becomes….. 2. If the first digit to be dropped is a 5 or greater, followed by at least one non-zero digit, the excess digits are dropped and the last retained digit increases by 1 62.782 rounded to 3 sig figs becomes… 3.If the 1st digit to be dropped is a 5 not followed by any other digits or followed by zeros, drop the 5 and any zeroes AND a. increase the last retained digit by 1 if it is odd b. Do not change the last retained digit if it is even 62.650 rounded off to 3 sig figs becomes 62.6…(even rule) 62.350 rounded off to 3 sig figs becomes 62.4….(odd rule)

Rules for mathematical operations Multiplication and division: -The calculated value has the same number of significant figures as the number with the fewest significant figures Addition and subtraction: -the uncertainty in the calculated value must be the same as the uncertainty in the input with the greatest uncertainty (last significant place) We are treating only the simplest of operations

Perform the following calculation and round your answer to the correct number of significant figures: Calculator answer: 3.23225000 The answer should be rounded to two significant figures because 0.456 – 0.421 = 0.035: 3.2

Calculate the following: 1. (2.05 X103) + (3.11 X 103) 2. (8.66 X 105) + (1.20 X 105) 3. (2.0 X 104) + (0.02 X104) 4. (5.401 X 103) + (2.101 X 10-8) 5. (6.0 X 106) + (7.75 X 10-8) 6. (1.01 X 103) - (9.952 X 10-2) 7. (4.4 X 102) - (9.56 X102) 8. (1.53 X 107) - (1.12 X 106) 9. (5.99 X105) - (8.65 X 102) 10. (4.5 X107) – (5.567 X10-5) 11. (7.65 X 104)- (9.9 X 103) 12. (1.9901 X 105) – (2.01 X104) 13. (2.0 X 102) X(2 X 102) 14. (1.0521 X 105) X (5.2 X 10-5) 15. (9.91 X 103) X (1.123 X 10-3) 16. (7.666 X 107) X (5.89 X108) 17. (1.0 X 10-6) X (2.25 X 106) 18. (6.857 X 10-5) ÷ (1.20 X 104) 19. (1.250 X 105) X (8.000 X 102) 20. (2.5 X 10-3) X (4.000 X 105) 21. (5.00 X 102) ÷ (5.000 X10-3) 22. (8.0 X 103) ÷ (2.000 X 103) 23. (5.65 X 107) ÷ (5.62 X 105) 24. (7.000 X 102) ÷ (3.5 X 104)

Dimensional analysis *this is a general problem solving approach in which the relationships between quantities (or factors) are used as a guide in setting up the calculation *this applies to one step as well as complex problems * in multiple step problems, an answer from a preceding step serves as a given for the next step Desired quantity= Desired quantity unit Given quantity unit X given quantity Conversion factor

*convert 12 grams into kilograms -given: 12 g; desired: kilograms kilograms = (1 kg/1000g) X 12 g = 0.012 kg -always check sig figs **convert 11 days into seconds -given: 11 days, desired: kilograms -multistep process -strategy: days to hours to minutes to seconds Seconds = 11 days X (24 hrs/1 day) X (60 mins/1 hr) X (60 secs/1 min) = 950,400 s

*other dimensional analysis problems involve do not merely changing the units of the same quantity Q:You're throwing a pizza party for 15 and figure each person might eat 4 slices. How much is the pizza going to cost you? You call up the pizza place and learn that each pizza will cost you $14.78 and will be cut into 12 slices A: $ 73.90

1. Express an acceleration of 9.81 m/s2 in ft/s2 2. How many eggs are in 10 and a half dozens of eggs? 3. Express the speed of 85 km/hr in inches per second. 4. Express 0.025 m in mm, km and miles 5. The mass of the earth is estimated at 6.6 X1021 metric tons. Express this mass in grams. 6. Convert a volume of 56L to mL, µL, nL and gallons 7. Express 2.5 L in mm3, in3 and ft3 Conversion factors: 1 metric ton = 1000 kg 1 ft = 12 inches 1 inch = 2.54 cm 1 gallon = 3.78 L