Chemistry Foundations of Chemistry Test. This is due:

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1 Chemistry Foundations of Chemistry Test This is due: Directions: Answer the following questions on a separate sheet of paper (or on this paper if you have room), staple to this paper (if you used a separate piece of paper) and turn in on the date due. You need to be able to write an answer showing that you understand all of the following information: 1. Provide a definition for the following terms: a) Sum: this term means to add two or more number together; this is what the answer to an addition problem is called b) Difference: this term means to subtract to or more numbers; this is what the answer to a subtraction problem is called c) Product: this term means to multiply two or more numbers together; this is what the answer to a multiplication problem is called d) Quotient: this term means to divide two or more numbers; this is what the answer to a division problem is called 2. What are the most important lab safety rules for the Chemistry classroom? a) Follow directions, no horseplay, wear PPEs, etc. 3. Why is lab safety important to us? a) To keep everyone safe during the lab experiments and to not waste or destroy equipment/chemicals 4. What is a qualitative observation? Provide an example. What is a quantitative observation? Provide an example. 5. You need to be able to interpret data that show a variety of possible relationships between variables a) Draw a line or curve showing the following relationships: 1. Positive 3. No relationship 2. Negative 6. What is the definition of chemistry? a) Chemistry is the study of matter b) Matter is anything with mass and volume c) Chemistry is the study of anything with mass and volume 7. What is the definition of matter? a) Matter is anything with mass and volume 8. What is the definition of mass? Volume? a) Mass is the amount of matter in a substance, similar to weight b) Volume is the amount of space an object takes up 9. What are the three phases of matter? What is different and similar between the three phases? Draw pictures of the molecules of each of these phases of matter. a) Solid 1. Defined volume and shape 2. Molecules are rigidly locked together b) Liquid 1. Defined volume, but variable shape 2. Takes on the shape of its container 3. Molecules are on average the same distance away from each other 4. Molecules are able to move around and are able to be poured 5. Molecules clump together c) Gas 1. No set volume or shape 2. Molecules act as individual molecules even if they collide

2 10. How can matter be classified? What are the two major categories? Provide an example of both categories. a) Pure substance and mixture 1. Pure substance = compound or element 1. carbon 2. Mixture = homogeneous and heterogeneous 1. Air 3. A mixture is made up of two or more different molecules/atoms that are not chemically bonded 4. A pure substance is just one type of molecule or compound 11. How can pure substances be further classified? What are the two major categories? Provide an example of both categories. a) Pure substance = compound or element 1. A compound is composed of two or more different elements bonded together 1. water 2. An element is just one type of atom Aluminum 12. How can mixtures be further classified? What are the two major categories? Provide an example of both categories. a) Mixture = homogeneous and heterogeneous 1. Homogenous the same throughout; two substances that are mixed together in a uniform manner 2. Heterogeneous different throughout; two substances that are mixed together but not in a uniform manner 13. Define pure substance, mixture, element, compound, homogenous, and heterogeneous. a) Pure substance 1. Pure substances are defined as substances that are made of only one type of atom or only one type of molecule (a group of atoms bonded together) b) Mixture 1. a substance made by mixing other substances together. Two or more pure substances mixed together c) Element 1. Elements are chemically the simplest substances and hence cannot be broken down using chemical methods. d) Compound 1. A compound is a substance formed when two or more chemical elements are chemically bonded together e) Homogeneous 1. A mixture which has uniform composition and properties throughout. For example, air is a homogeneous mixture of gases. A teaspoonful of table salt stirred into a glass of water also makes a homogeneous mixture. f) Heterogeneous 1. A mixture is a combination of two or more pure substances in which the original substances retain their chemical properties. In some mixtures, the initial substances cannot be detected after they have been mixed

3 14. Determine whether the following is an element, compound or mixture. a) a. Oxygen gas element b) b. Kool-Aid mixture (homogeneous mixture) c) c. Carbon Dioxide (CO 2) compound d) d. Vinegar (CH 3COOH) compound 15. Classify Matter (based on arrangement) a) Identify each picture below with the type of matter from the list 1. Heterogeneous Mixture 2. Homogeneous Mixture (Solution) 3. Compound 4. Element 5. Solid 6. Gas Gas gas gas and solid gas solid Compound element mixture mixture element Pure substance pure substance heterogeneous homogeneous pure substance 16. What type of matter can be separated by physical means? What type of matter can only be broken down by chemical means? a) A mixture can be separated by physical means b) A compound can only be broken down by chemical means c) NOTE an element can never be broken down 17. Define physical properties/reactions. Provide an example for each. a) A physical property is a property that does not change that chemical nature of matter. 1. Color, shape, texture, volume, mass 1. A physical reaction is a reaction that changes the appearance/shape/look of the substance, but does not change the chemical composition of the substance. It is still the same substance even though it looks different. 1. Cutting in half, melting, boiling, changing phase b) Physical can be observed without changing the chemical composition (color change by using your pencil on your notebook paper); changes the appearance of the item, but it is still the same substance before and after the change 1. Color, texture, size, etc. 18. Define chemical properties/reactions. Provide an example for each. a) A chemical property is a property that does change that chemical nature of matter. 1. Flammability, ionization, combustibility 1. A chemical reaction is a reaction that changes the appearance/shape/look of the substance, and also changes the chemical composition of the substance. It is a new substance at the end of the reaction. 1. Reacting with an acid or base, producing bubbles without changing the temperature, producing a precipitate/solid without changing the temperature, producing heat and light, exploding, burning, spoiling, rusting, rotting 2. Chemical changes the chemical composition of the substance (ie. fire); what you end with is different than what you start with 1. Flammability, reactivity, etc. 19. Determine whether the following are chemical or physical changes. a. Breaking glass physical b. Two substances mixed generate heat chemical c. Ripping paper into two physical d. Liquid water freezing physical 20. Define density. What is the formula? a) Density is the ratio of mass to volume for an object 1. D = m/v 1. m = D * v 2. V = m/d

4 21. How can density be used in chemistry? a) Can be used to identify an unknown substance because every pure substance has its own unique density b) b. Can be used to determine if an object will float or sink (must be less dense that the substance you want to object to float in) 22. Show an example density problem. a) Density is mass divided by volume, so that the density is 45 g divided by 15cm 3, which is 3.0 g/cm3. 1. Density = mass/volume 2. Density = 45g / 15cm 3 = 3.0 g/cm An object has a density of 1.38 g/ml. What is the volume of 35.4 g of this object? a) D = m/v 1. m = D * v 2. V = m/d b) V = m/d g / 1.38 g/ml 2. = ml g of an object is found to have a volume of 4.35 ml. What is the density of the object? a) D = m/v 1. m = D * v 2. V = m/d b) D = m/v g / 4.35 ml 2. = 1.49 g/ml 25. Using the graph below, calculate the slope of Line B. Show all work and units a. Since the graph has mass on the y and volume on the x, this means that the slope (y/x) is Density b. This means that you need to pick a data point on the line where you can determine the mass and volume exactly and then divide y/x or mass/volume i. I picked the data point 80 g / 7 ml ii. The slope also known as the density is g/cm If the mass of a sample of Line B was 160 grams, what is the volume? a) D = m/v 1. m = D * v 2. V = m/d b) Since we know that the density is g/cm 3 and the mass is 160 grams V = m/d = 160 g / g/cm3 = cm Determine the density of an object from a data table or from a graph of mass vs volume. Volume (cm 3 ) Mass (g)

5 Plot the data above see graph Determine the density of each sample of the substance. o 7.8 g/cm 3 o 8.0 g/cm 3 o 7.8 g/cm 3 o 8.0 g/cm 3 o 7.8 g/cm 3 o 7.8 g/cm 3 o o 7.9 g/cm 3 = average density Determine the average density o 7.9 g/cm 3 = average density Solve for the volume 150g of the substance occupy. Show work; use labels. 150 g x 1 cm = 19 cm 3 use the density equation to solve for volume 7.9 g 28. Define accuracy and precision. Draw a dart board with high accuracy and low precision. Draw a dart board with high accuracy and high precision. Draw a dart board with high precision and low accuracy. a) Accuracy tells me how correct you were; how close were you to the number you were supposed to get; getting 100% on a test b) Precision requires multiple trials how close together several trials are; getting an 80% on every test in the class 29. Define percent error. What is it used for? What is the formula? a) Percent error or percentage error expresses as a percentage the difference between an approximate or measured value and an exact or known value. It is used in chemistry and other sciences to report the difference between a measured or experimental value and a true or exact value. b) Formula = (accepted experimental) / accepted * Calculate the % error for the following: Show all work and units! True value: g/ml Student s lab result: g/ml (11.85 g/ml g/ml) / g/ml x 100 = 3.97 % (Remember no negative values) 31. Explain how scientific notation can be used in science. Write a large number and small number in scientific notation. Show the conversion from standard to scientific notation. Show the conversion from scientific notation to standard. a) Used for large or small numbers to make expressing and using them easier 1. What are the rules for writing a number in scientific notation? 1. Example 1: = 3.6 x If you move the decimal to the right, then n is a negative number 2. If the original number is smaller than 1, n is a negative number 2. Example 2: 25,000 = 2.5 x If you move the decimal to the left, then n is a positive number

6 1. If the original number is more than 1, then n is a positive number 2. What are the rules for changing a number in scientific notation to regular notation? 1. If the exponent is negative then you need to move the decimal to the left 1. Your final answer should be less than 1 2. If the exponent is positive then you need to move the decimal to the right 1. Your final answer should be greater than Convert to either scientific notation or decimal form: a. 58,200, x 10 7 b x 10-4 c x d x Explain how to multiply and divide numbers that are in scientific notation. Provide an example. a) Multiplication (remember that you just need to multiply the main numbers and add the exponents). Then adjust to fit proper scientific notation 1. (2.0 x 10 2 )(3.0 x 10 3 ) = 6.0 x ii. (2 x 10 2 ) x (6 x 10 3 ) = 12 x 10 5 = 1.2 x 10 6 b) Divide the coefficients; subtract the exponents 1. (6.0 x 10 7 ) / (3.0 x 10 4 ) = 2.0 x (3.0 x 10 7 ) / (6.0 x 10 4 ) = 0.5 x 10 3 =5 x What are the rules for significant figures? Why do we use them? a) Any number that is not a 0 is significant b) A zero bookended by two non-zero digits is significant c) A zero at the beginning of a number (meaning a decimal number) is never significant d) A zero at the end of a number is only significant is there is a decimal point e) The M part of scientific notation is significant (what comes before the x10) f) We use significant figures because: 1. Measurements are used in chemistry to communicate information 2. We use the measurements taken during the lab and experiments to support our findings 3. We need a universal way to discuss these measurements 4. We use them as a universal way to express and evaluate precision and accuracy 35. The best measurement, using significant figures, at the arrow above would be 2.20 cm cm When reading a ruler, we generally report a measurement by recording all of the certain digits plus uncertain digits. a) one 37. What are the rules for adding and subtracting with significant figures? a) Round to the number with the lower number of decimal places 38. What are the rules for multiplying and dividing with significant figures? a) Round to the number with the lower number of sig. figs. 39. Express the answer in the correct number of significant figures. Label with appropriate units. a g = g/cm cm 3 b cm 2 x 1.2 cm = cm cm c m x 61.5 m = m 2 d cm = cm/in 8.50 in

7 40. How is the metric system used in chemistry? a) The metric system is used with measurements b) We use the metric system for conversions to allow us to complete calculations 41. Write the metric system showing how each unit is related to the base unit. 42. Convert measurements into various sizes. a) What is the size of the following metric measurements: 1. 1 km = 1000 meters 2. 1 meter = 100 cm 3. 1 liter = 1000 ml

8 43. Show a sample factor label problem including units. 44. Complete the indicated conversions: a. 37 g x = mg 37 g x 1000 mg = mg = mg 1 g b. 138 m x = km 138 m x 1 km = km = km 1000 m c. 4.7 kg x = g 4.7 kg x 1000 g = 4700 g = 4700 g 1 kg d mm x = _ m 4021 mm x _1 m = m = m 1000 mm 45. What property of mass does not change even though the appearance of the matter may? a) Its mass does not change b) Law of conservation of mass 46. What is the law of conservation of mass? a) The law of conservation of mass states that the mass before a reaction must equal the mass after.

9 47. Identify the law of conservation of mass using graphs. a) Demonstrate knowledge that matter is conserved in ordinary chemical reactions. Example: If 3.6 g of water is broken down, 3.2 g of oxygen and what mass of hydrogen will be produced? 3.2 g 3.6 g Mass Mass? g Water Oxygen Hydrogen The law of conservation of mass states that the mass before a reaction must equal the mass after. So if we know that we started with 3.6 g of water, and that 3.2 g of oxygen is produced, what must be the mass of hydrogen to keep the mass before and after the reaction the same? The answer is that the mass of hydrogen is 0.4 g. This is because you start with 3.6 grams and know that part of what you end with is 3.2 g. If you subtract these, you will see that the hydrogen makes up the remaining 0.4 grams to get us to 3.6 grams on both sides of the arrow. 48. The mass of a flask filled with water was recorded as 252 g. A chunk of alka seltzer was massed with the flask and the combination was recorded as g. The chunk of Alka-Seltzer was then placed in the flask. After some bubbling occurred and a few minutes passed, the flask was massed again and recorded as g. a. According to the above recorded-masses, what was the mass of the alka seltzer chunk? a. The flask mass before the alka seltzer was added was 252 g b. The mass of the flask combined with the alka seltzer is g c. If you subtract these two masses, you will get the mass of the alka selter d. = g 252 g = 3.72 g b. How much Carbon Dioxide gas was released in this reaction? a. We know that our starting mass is g for the experiment and that the ending mass is g b. If we subtract these two masses, then we will know how much carbon dioxide (bubbles) were given off in the reaction c. = g g = 0.16 g 49. Define solute, solvent, saturated, unsaturated, and supersaturated. 50. Using Kool-Aid as the solution, what is the solute? What is the solvent? 51. Explain how to read a solubility curve. 52. Looking at the curve below, if you have dissolved 10 grams of KNO 3 at 40 o C, is your solution saturated, unsaturated, or supersaturated?

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