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Introduction to Matter and Measurement

Introduction to Matter and Measurement. Important Concepts used in the Study of Chemistry Ó. Written by JoAnne L. Swanson University of Central Florida 2002 - 2003. Chemistry is the study of the properties and behavior of matter. MATTER IS ANYTHING THAT HAS. All MATTER IS MADE UP OF THE

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Introduction to Matter and Measurement

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  1. Introduction to Matter and Measurement Important Concepts used in the Study of ChemistryÓ Written by JoAnne L. Swanson University of Central Florida 2002 - 2003

  2. Chemistry is the study of the properties and behavior of matter. MATTER IS ANYTHING THAT HAS All MATTER IS MADE UP OF THE . THE PROPERTIES OF MATTER ARE RELATED TO THE ELEMENTS THAT IT CONTAINS.

  3. THE ELEMENTS ARE COMPOSED OF ATOMS WHICH ARE ______________________________________________________________________________________________________________. MOLECULES ARE COMBINATIONS OF TWO OR MORE ______,_______________________________. THESE BONDS CAN ONLY BE _____________________________________________________.

  4. WHY LEARN ABOUT CHEMISTRY? You and everything around you is chemical. A series of chemical reactions is what causes your heart to beat. The air you breathe is a mixture of gases (Chemicals). The carpet on the floor, the paint on the walls, the materials your seat is made of, the medicines you take, were all originally put together by people with training in chemistry. Studying chemistry will also provide you with tools useful in other courses, advancement in all sciences, and also life outside of school. You will learn measuring techniques. You will have a better understanding of your environment, household products and more.

  5. The Classification of Matter MATTER Physically separable AKA SOLUTION Chemically separable

  6. Pure substances – are ___________________________. Elements are in the simplest form and cannot be broken down any further. Compounds are___________________ _______________________________________________ Heterogeneous Mixtures – mixtures of pure substances in which ______________________ between the different components of the mixture. Ex. Granola, or oil and vinegar. Homogeneous Mixtures – mixtures of pure substances in which _________________________. It looks the same throughout. Ex. Alcohol in water, or brass or air. Aka solutions. Mixtures ______________________________________ _______________________________________________

  7. Notes: Matter can be ___________________________________________- _______________________________________________________- Compounds can be broken down into elements by _______________ and elements can ________________________________________ ______________________________________________________ When a compound forms from a chemical reaction of two or more elements, the ____________________________________________ ___________________________________________________. For example: Sodium ____________________________________ ______________________________________________________ Chlorine is _______________________. When sodium reacts with chlorine, __________________________ _____________________________________________________.

  8. Pure substances that don’t react with one another can be mixed together __________________. A mixture can be separated back into pure substances by __________________________________ _____________. _________________________________________ ____________________________________________________ don’t involve changing the chemical composition of the components of the mixture. In a mixture, each substance keeps its own properties. For example, in a mixture of sucrose and water, _______________ ______________________________________________________ ______________________________________________________ ______________________________________________________.

  9. ALL MATTER CAN EXIST IN A SOLID, LIQUID OR GASEOUS STATE PROPERTIES OF DIFFERENT STATES OF MATTER: • SOLIDS – THEIR PARTICLES ARE ________________ and_________________________. • THEY ARE _____________________. • THEY HAVE _________________and________________. • THEY HAVE ALL ___________________________.

  10. LIQUIDS – THE PARTICLES _____________________________________________________ - THEY ARE _______________________ -THEY CAN ___________ -THEY HAVE ___________________ BUT NOT _______________________ -THEY HAVE _______________________ GASES – THE PARTICLES ARE _________________________ ____________________________________________. -THEY ARE ________________________. - THEY CAN __________. -THEY HAVE NO ________________ OR __________. -THEY HAVE __________________________.

  11. THE ELEMENTS There are about 112 known elements. We use symbols to represent the different elements. The symbols may have one or two letters. If two letters, the FIRST ONE MUST BE _____________________________________ ____________________________________. If there is one letter, it must be capitalized. This helps to differentiate between elements and compounds. For example: C is_________, Co is_______, CO is____________

  12. Some symbols come from foreign names: Iron is Fe from Copper is Cu from Lead is Pb from Potassium is K from Silver is Ag from Sodium is Na from Tin is Sn from Mercury is Hg from hydrargyrum NOW IS A VERY GOOD TIME TO BE SURE YOU KNOW ALL OF THE NAMES OF THE ELEMENTS AND THEIR SYMBOLS.

  13. The periodic table is a tool used by all chemists. The elements on the periodic table are arranged by _____________ ___________ (to be discussed later), and are in vertical groups with elements that have ___________________________. The groups are numbered across the top of the table. There are three different numbering systems._______________________________ _______________________________________________________ ______________________________________________________. For example, group IA consists of elements _________________ IT IS VERY IMPORTANT TO REMEMBER THAT ELEMENTS IN THE SAME GROUP HAVE SIMILAR CHEMICAL PROPERTIES

  14. THE PROPERTIES AND CHANGES THAT MATTER UNDERGOES There are two classifications of properties: physical properties are those that ______________ ______________________________________________. Some examples are: boiling point, melting point, freezing point, density, color, odor, texture. Ex.____________________________________________.

  15. Chemical properties are those that _________________ _____________________________________________. Some examples are the ability to burn (combust), the ability to react with something, toxicity, the ability to tarnish. Ex. When Iron (Fe) rusts, _________________________ ______________________________________________ ______________________________________________ ______________________________________________. Demo: the ability of alcohol to burn and the property that water doesn’t burn.

  16. Physical properties are also categorized into INTENSIVE PHYSICAL PROPERTIES AND EXTENSIVE PHYSICAL PROPERTIES. ___________________________________ ARE INDEPENDENT OF QUANTITY. For example, water ______________________ _______________________________________________________ _______________________________________________________. Intensive properties can be used to identify a substance. For example, a lustrous metal with a density of 2.71 g / mL is Aluminum.

  17. EXTENSIVE PHYSICAL PROPERTIES DO DEPEND ONQUANTITY. Some examples of extensive properties are mass, volume, length, etc.

  18. Physical and Chemical changesrefer to an action, rather than an ability of a substance to act. Alcohol burning is a _______________. Flattening a piece of metal is a ___________________. Silver tarnishing is a ___________________. Boiling water is a ____________________. (Note, the ability of alcohol to burn is a property).

  19. Qualitative and Quantitative Observations Qualitative observations are those that ____________________ ________________________. For example, ________________________________________-. While we often make qualitative observations in chemistry, we make mostly quantitative measurements in the lab. Quantitative measurements involve ______________________. For example, _____________________________.

  20. Units of Measure in Science We use the metric system because every unit is a power of ten. System International, or SI units are the base units used to derive all other units. There are prefixes used to describe the relative size of the quantity with regard to the base unit. *

  21. Prefixes that indicate a quantity larger than the base unit: Mm = mega meter. There are 106 meters / 1 Mm Kg = kilogram. There are 103 grams / 1 kg * Note that while kg is the base unit because it is the standard used, the gram is the unit onto which the prefixes are added.

  22. Prefixes that indicate a quantity smaller than the base unit: Angstroms are used only for length and are represented by 1010 A / 1 m. They are used mostly in describing sizes of atoms.

  23. Examples: There are 10 dm / 1 m There are 102 cg / 1 g There are 103 mL / 1 L Notice that in both the bigger and smaller prefixes, if a 1 is always put in front of the LARGER unit, there is never a need for negative exponents. EXPLAINATION_______________________ ______________________________________________________.

  24. VOLUME IS DERIVED FROM LENGTH. VOLUME = L x W x H. A cm3 = 1cm x 1 cm x 1 cm. It turns out that 1cm3 = 1 mL And 103mL / 1 L 10cm = 1dm 10cm = 1dm 1cm3 = 1 mL 1cm 1cm 10cm = 1dm 1cm 1000cm3 = 1 dm3= 1 L

  25. TEMPERATURE Temperature is the _________________________________. Heat always flows from an area of ______________________________. Example, touching a hot pan. In science and in many other countries, the Celsius scale is used for everyday temperature measurement. The SI unit of Kelvin is used for many scientific calculations. Temperature conversions: To convert from oF to oC use, oC = (oF - 32) / 1.8 To convert from oC to oF use, oF = 1.8 oC + 32 To convert from oC to K use, K = oC + 273 To convert from oF to K, first convert to oC then to K.

  26. Examples: Convert 98.6 oF to oC Convert 40.0 oC to oF Convert 25.0 oF to K

  27. DENSITY Density is the _________________________________________. For solids and liquids the units for density are g / cm3 or g / mL (these are interchangeable since 1cm3 = 1mL) For gases the units used are g / L Water has a density = 1.00 g / mL (know this) Air has a density = 0.001 g / mL or 1 g / L

  28. Uncertainty in Measurement All measured quantities have some error or uncertainty. These are _________________. ________________________________ ____________________________________________________. Numbers that are known exactly are called exact numbers. They are counted numbers or ratios to 1. For example, 12 eggs / 1 dozen, 103g / 1 kg, 10 people / 1 table, 2.54cm / 1 inch

  29. When we make measurements in science we strive for ____________________________-. PRECISION – the ability to ______________________ __________________________________________. ACCURACY – ________________________________ ____________________________________________. For example, if we say the object had a mass of about 20 grams, __ ______________________________________________________. If the object is measured and found to have a mass of 20.0251 g, ___ _______________________________________________________ ______________________________________________________

  30. Significant Figures We use significant figures in measurement to indicate the uncertainty in a measurement. The last number in a measurement is always _________________ and also shows the _________________________________. Each measuring instrument has increments (divisions). We take a measurement to the smallest increment, and then ____________________________________________ _______________ , (_______________________.) If we believe that the measurement is exactly on the increment, then we __________________ to the number to indicate this.

  31. On this centimeter ruler, the smallest increment is ____ ______________________. The first arrow is pointing between 1.4 and 1.5 cm. We may estimate that it is ____or _____cm. The second arrow appears to be right on 4.4 cm and we indicate this by writing, 4.40 cm. Notice each measurement is out to the ___________________________________________ ______________________________________________.

  32. It turns out that if the smallest division (increment) is in the ones place, we write our measurement to the ___________. If the smallest division is in the tenths place, we write the measurement to the ___________________. If the smallest division is in the hundredths place, we write the measurement to the __________________. That extra place is our estimated value, even if it is a zero. Ex. 3.521 This measurement has 4 significant figures. The one is the estimated value, in the thousandths place. The two is the position of the smallest increment on the instrument, the hundredths place

  33. To determine the number of significant figures in a number there are some easy guidelines to follow. • Zeros before a non-zero digit are __________significant. 0.0000500 has 3 significant figures. • Zeros after a non-zero digit are shown to be significant iff ______________________ number. 5200 has 2 s.f., 5200. has 4 s.f., 5.200 has 4 s.f.

  34. When doing calculations with measured values, the correct number of significant figures must be maintained so as not to imply more or less accuracy in the final answer than the original measured values. There are two rules to follow: 1. For multiplication and division – the answer should contain the same number of significant figures as the operator with the ______________________________ in the problem. 2. For addition and subtraction – the answer should contain the same number of ______________________ _____ as the operator with the least.

  35. Ex. 3.20 x 4.0 = 12.80 = 13 (only two s.f. allowed in the answer). 3s.f 2 s.f. 2s.f. 3.21 + 4.2 + 5 = 12.41 = 12 (no digits right of the decimal). 2 1 none none

  36. DIMENSIONAL ANALYSIS This is a problem solving technique that is a valuable and easy way to solve many chemistry problems. It involves starting with a given value and unit, and algebraically canceling out units until you arrive at the desired unit. We do this by using conversion factors which are ratios such as: 10 mm / 1 cm or 103 g / 1 kg, etc.

  37. Examples: • Convert 2.0 x 105 nm to km • the given is • info that we are aware of is • Now we plug in these conversion factors so that the units cancel until we end up with the desired units of km. • 2.0 x 105 nm x = 2.0 x 10 –7 km

  38. There are 6.02 x 1023 atoms / 1 mole and 14.01 g N / 1 mole • How many atoms are there in 75.2 g of N? • (write the given) Remember you always want to cancel out the unit that you start with otherwise you wouldn’t be doing the problem.

  39. Conversion factors you are required to know • all of the metric conversions given earlier in this chapter • English / metric conversion factors such as: 2.54 cm / 1 inch, • 454 g / 1 lb, 1.61 km / 1 mile, 946 mL / 1 qt, • also some English conversions such as: 5280 ft / 1 mile, • 1 lb / 16 oz, 1 gal / 4 qts, 1 qt / 4 cups • 12 inches / 1 foot, 3 feet / 1 yard, 2000 lb / 1 ton

  40. Definitions – chemistry, matter, elements, compounds, atoms, molecules, temperature, intensive and extensive, heterogeneous, homogeneous, qualitative, quantitative, density, precision, accuracy, significant figures. • Symbols for the elements • Physical and chemical properties • Physical and chemical changes • States of matter and the characteristics of each • Classification of matter • Significant figures in calculations • Units of measure and conversion factors • Dimensional analysis problem solving technique Topics covered in this chapter

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