Guided Notes Matter

Chemistry I Guided Notes For Unit II What is Chemistry? Chemistry is the study of the composition, structure, and properties of matter and the changes it undergoes. 5 Major Branches of Chemistry  Organic Chemistry: The study of carbon containing substances. Common examples include pharmaceuticals and plastics.  Inorganic Chemistry: The study of substances not containing carbon. Inorganic chemists might be involved in ceramics, catalysis, or semiconductors.  Analytical Chemistry: This is concerned with the separation or determination of the composition of substances. Forensics, nutritional analysis, and assisting other scientists are among the research areas for analytic chemists.  Physical Chemistry: This deals with the laws, principles, and theories that describe the behavior of chemicals. Physical chemists study things like reaction rates, reaction mechanisms, and thermodynamics.  Biochemistry: The study of the chemistry of organisms. Examples include studying metabolism, fermentation, and medicine. Classification of Matter  Matter: anything that has volume and mass.  Pure substance: something made up of one kind of material with specific properties. Both elements and compounds are pure substances.  Mixture: combination of two or more pure substances. Some examples are granite, trees, and pencils.

Laliberte, 8/25/08

1

 Compound: substance that contains two or more elements combined in a fixed proportion. Compounds can be written as a formula. Common examples are water (H2O), carbon dioxide (CO2), and sodium chloride (NaCl).  Element: simplest type of pure substance; contains only one type of atom. Elements are the substances listed on the periodic table.  Heterogeneous mixture: mixture where the different phases (parts) can be seen. Dirt and a mixture of salt and pepper are examples.  Homogeneous mixture: mixture where the different phases (parts) cannot be seen. These can also called solutions. Salt water and brass are examples. Fundamental Building Blocks of Matter and Phases of Matter  Atom: the smallest particle of an element. Atoms are composed of a nucleus which is surrounded by a cloud of electrons. The nucleus houses protons and neutrons.  Molecule: the smallest particle of a covalent compound that still has the composition and properties of that substance. Substances can exist in three different states of matter depending on the strength of bonding and the amount of kinetic energy available. When bonding between particles is strong, the solid phase is more likely. Weak bonding favors the gaseous phase. The average kinetic energy of a substance is its temperature. At higher temperatures gases are favored, while solids are more likely at lower temperatures. Solid: In a solid, the particles (atoms or molecules) are in contact with each other and have a very limited ability to move. Characteristic properties include definite volume and shape, slight expansion when heated, and incompressible. Liquid: Like solids, the particles are in contact with each other; however, the

Laliberte, 8/25/08

2

particles can move around more freely. Characteristic properties include definite volume, takes the shape of its container, slight expansion when heated, incompressible, and has the ability to flow. Gases: In gases, the particles are not in contact with each other. Characteristic properties include takes the shape and volume of its container, large expansion when heated, compressible, and has the ability to flow. Distinguishing Between Physical and Chemical  Physical Change: process that alters a substance without changing its composition. ▫ Ex. Boiling water, dissolving sugar in water, carving wood.  Chemical Change: a chemical reaction ▫ Ex. Burning a candle, rusting.  Physical Property: characteristic that is determined without changing the composition. ▫ Ex. Color, boiling point  Chemical Property: characteristic that changes the composition when determined. ▫ Ex. Flammability. As can be seen above, the main distinction between chemical and physical is that there is a change in identity for chemical changes and properties. In a chemical change, new substances (elements or compounds) are created by breaking bonds and/or forming new bonds. There are five pieces of evidence that can help you identify if a chemical reaction has taken place. Evidence of a Chemical Reaction: 1. Precipitation–formation of a solid out of a liquid or gaseous phase. 2. Color change 3. Formation of a gas 4. Heat change–usually temperature increases but not always 5. Plating–formation of a thin metal coating Separation Techniques Laliberte, 8/25/08

3

Separation techniques are based upon differences in physical properties Many techniques involve solubility or phase changes. Evaporation  Used to recover a solid from solution  Solvent is often actually boiled off – not evaporated  Watch glass reduces loss of solid from spattering Filtration  Used to separate an insoluble solid from a liquid  The liquid that comes through is called the filtrate Distillation  Separates mixtures of liquids having different boiling points  Lower boiling liquid condenses in the condenser Fractional Distillation  Used for mixtures that are hard to separate  Typically used when there is more than one volatile substance Chromatography  Separates different substances dissolved in a liquid based upon the tendency to travel across the surface of another material Extraction  Separates different substances dissolved in a liquid based upon solubility

Laliberte, 8/25/08

4

Significant Digits Overview All measurements have uncertainty. In science, it is important to indicate the uncertainty when you are reporting numbers since this indicates the reliability of the data and the quality of instruments being used. When a scientist takes a measurement, she records all the digits she is certain of and includes one digit that has some uncertainty. For instance, if a standard metric ruler were used to find the length of a pen, the reading may fall between 15.4 cm and 15.5 cm. The scientist may estimate that the length is closer to the 15.5 mark and report the value as 15.47 cm. In this measurement, only the 7 is uncertain and is probably off by 1 or 2 at the most. Sometimes the uncertainty will be reported using the ± symbol. For example, 244 ± 3. When the ± symbol is not used, expect that the right-most digit is uncertain and could be off by 2. For example, a value of 34.23 mL means that the scientist is fairly certain the value is between 34.21 and 34.25 mL with the best guess being 34.23 mL. Besides measurements, scientists also report calculated values (like density). There is a two-part rule that indicates how to round calculated numbers so that they show the same degree of uncertainty as the measurements they were calculated from. To use the rule for multiplication and division you will need to know how many significant digits each factor has. The number of significant digits is the number of digits that indicate precision. It is the number of digits you are certain of and the one uncertain digit. The number of significant digits in 15.47 cm (the length of the pen) is 4. Determining the number of significant digits Atlantic-Pacific Rule: Draw a map of the US and label the Atlantic and Pacific oceans. When the decimal point is Present, count the digits starting with the first non-zero digit from the Pacific side (left side). When the decimal point is Absent, count the digits starting with the first non-zero digit from the Atlantic side (right side). Examples: 1. 34000 Decimal is absent. Starting from the right, the 1st non-zero digit is the 4. There are 2 significant digits. 2. 0.005030Decimal is present. Starting from the left, the

Laliberte, 8/25/08

5

1st non-zero digit is the 5. There are 4 significant digits. As you can see from the examples, some of the zeros get counted and some don’t. On the other hand, all non-zeros are counted all of the time. Reporting calculated numbers to the correct significant digits Many-Places Rule: When doing Multiplication or division, report your answer to as Many significant digits as the factor with the least significant digits. When doing Plus or minus (i.e., addition or subtraction), report your answer to the same number of Places as the factor with the least. Examples: 1. 0.4563 x 13 = 5.9 2. 28.432 + 37.2 = 65.6

(5.9319 before rounding) (65.632 before rounding)

Scientific Notation Very large and very small numbers are not uncommon in chemistry. Using scientific notation is the best way to work with such numbers. In scientific notation, the number is expressed as the product of two factors. The first factor is a number between 1 and 10 and the second factor is a power of 10. As an example, 250000 is written as 2.5 x 105 using scientific notation. When converting a number to scientific notation, the first step is to move the decimal point so that there is one nonzero digit to the left of the decimal point. Next, multiply by 10 raised to the power equal to the number of places that the decimal was moved to the left. If the decimal was used to the right, 10 will be raised to a negative exponent. Ex. 1 5,000,000 = 5 x 106

Ex. 2 0.00000036 = 3.6 x 10–7

When converting a number to standard notation, move the decimal the number of places equal to the exponent that 10 is raised to. Move to the right when the exponent is positive and to the left when the exponent is negative. Ex. 1 Laliberte, 8/25/08

Ex. 2 6

2.3 x 104 = 23,000

5.12 x 10–3 = 0.00512

Know the Right Word  Qualitative: A description that is descriptive. Ex. The rock is heavy.  Quantitative: A description that is numerical; measurement. Ex. The rock weighs 45 kg.  Intensive property: Property that doesn’t depend upon the amount of matter present. Example: Chalk is white.  Extensive property: Property that depends upon the amount of matter present. Example: The mass of a nickel is 5.0 grams.  Accuracy: Refers to how close a measurement is the accepted or correct value. A synonym for accurate is correct. Imagine 3 darts were thrown at a dart board. Where would they have to land to have good accuracy? Ans. If you were aiming for the bull’s eye, all the darts should be in the bull’s eye or at least very close.  Precision: Refers to how close a set of measurements are to each other. A synonym for precise is reproducible. Imagine 3 darts were thrown at a dart board. Where would they have to land to have good precision? Ans. The three darts should all be very close to each other. It doesn’t matter where on the dart board they land.  Mass: Mass is a measure of the amount of matter in a substance. Scientists prefer to use mass instead of weight because the mass of an object is not affected by changes in gravity. Mass is measured in grams or kilograms using a balance. A balance determines the mass by comparison with objects of known mass. Weight: Weight (w) is a measure of the force of gravitational attraction between the earth and an object. It is equal to the mass of the object (m) multiplied by the gravitational constant (g); w = mg. Since the gravity on the moon is about 1/6 of what it is on earth, a 180 lb person would only weigh 30 lbs on the moon. Weight is measured in pounds or newtons (the metric unit) using a scale. A scale measures force, often with the use of a spring.

Laliberte, 8/25/08

7

Metric System Base Units in the Metric System Quantity Name length meter mass grams time seconds amount of substance mole temperature kelvin electric current ampere light intensity candela

Symbol m g s mol K A cd

Useful Derived Units in the Metric System Quantity Name

Symbol kg g g 3 or 3 or m cm mL m ⋅ kg N or 2 s kg Pa or m ⋅ s2 m2 ⋅ kg J or s2

Density Force

Newton

Pressure

Pascal

Energy or Quantity of Heat Volume Metric prefixes Prefix Abbre Exp. v Factor tera T 1012 giga G 109 mega M 106 kilo k 103 hecto h 102 deka da 101 100 deci d 10–1 Laliberte, 8/25/08

Joule

L or dm3

Liter Meaning

1,000,000,000,000 1,000,000,000 1,000,000 1,000 100 10 1 0.1

Example 1 Tm = 1x1012 m 1 Gm = 1x109 m 1 Mm = 1x106 m 1 km = 1000 m 1 hm = 100 m 1 dam = 10 m 1 meter 1 dm = 0.1 m 8

centi milli micro nano pico femto atto

c m µ n p f a

10–2 10–3 10–6 10–9 10–12 10–15 10–18

0.01 0.001 0.000 001 0.000 000 001 0.000 000 000 001 0.000 000 000 000 001 .000 000 000 000 000 001

1 cm = 0.01 m 1 mm = 0.001 m 1 µm = 1x10–6 m 1 nm = 1x10–9 m 1 pm = 1x10–12 m 1 fm = 1x10–15 m 1 am = 1x10–18 m

Metric Conversions There are two methods commonly used to perform conversions within the metric system. While the first method is certainly easier, the second method introduces you to a method that is much more powerful and will be very useful for solving difficult problems.

Method 1: Moving the decimal Because the metric system is designed around powers of 10, you can convert between different units by moving the decimal point. Move the decimal to the left when you are changing to a larger unit and move to the right when changing to a smaller unit. The number of places to move is equal to the difference in the exponents from the Exp. Factor column in the above table. Ex. 1 Convert 450 mL to L Move to the left 3 places; .45 L

Ex. 2 Convert 0.25 km to cm Move to the right 5 places; 25000 cm

Method 2: Conversion Factors This method takes more time for simple problems like the previous two examples; however, it does make more difficult problems easier to do. The basic idea is to multiply by a fraction that contains the ending units you are trying to get to in the numerator and the starting units in the denominator. Write the number 1 in front of the larger unit. For the smaller unit, write in the number of smaller units it takes to make up one of the larger units. Finally, perform the calculation. The example below should make this more clear.

Laliberte, 8/25/08

9

Ex. Convert 450 mL to L Step 1: Write the given number, the multiplication symbol, a fraction bar, and an equals sign. 450 mL x ––––––––––– = Step 2:

Write the units you are starting with in the denominator and the units you are ending with in the numerator. L 450 mL x ––––––––––– = mL

Step 3:

Put a “1” in front of the larger unit (in this case, L). Next, make the numerator and denominator equivalent by writing in the number of smaller units it takes to make up one of the larger units. 1L 450 mL x ––––––––––– = 1000 mL

Step 4:

Calculate (multiply by things on top and divide by things on the bottom). Notice that the units of mL cancel out. 1L 450 mL x ––––––––––– = 0.45 L 1000 mL

Unit Conversions The method of moving the decimal place cannot be used for conversions within the U.S. customary system of measurements and between the U.S. system and the metric system. For these conversions we will use the conversion factor method. Step 3 from this method will be modified. You have been given a list of conversions in the Reference Packet. Use these values for the numbers that you put into the conversion factor. Ex. 1 Convert from 175 lb to kg. Steps 1 & 2:

Laliberte, 8/25/08

kg 175 lb x ––––––––––––– = lbs 10

0.45359237 kg 175 lb x ––––––––––––– = 79.4 kg 1 lb

Steps 3 & 4:

Note: The answer has been rounded to the proper significant digits. Ex. 2 Convert from 75

mi km to . hr hr

in Steps 1 & 2:

75 mi

Note: “mi” was placed

–––––– x –––––––––– = hr mi 75 mi

Steps 3 & 4:

km

denominator so it would cancel out.

1.609344 km

–––––– x –––––––––– = 120 hr

km hr

1 mi

Ex. 3 (Note: This problem requires 2 conversions.) Convert from 13.6

mi km to gal . L

13.6 km Change km to mi

––––––– x ––––––––––– = 8.45064821 L

Change from L to gal

1 mi mi L

1.609344 km

8.45064821 mi 3.785412 L –––––––––––– x –––––––––– = 32.0

mi gal

L

1 gal

It would be more efficient to combine these two steps as shown below:

Laliberte, 8/25/08

11

13.6 km

1 mi

3.785412 L mi

––––––– x ––––––––––– x –––––––––– = 32.0 gal L 1.609344 km 1 gal

Density If you had a block of aluminum and a block of lead of equal size, the lead block would be heavier. This is because lead has a higher density. Density is a measure of mass per unit volume. The formula is Density, D =

mass m = volume V

When trying to determine density of an object, you will probably have to determine its volume. The volume of a rectangular solid can be found using V = lwh. The volume of a cylinder can be found using V = Πr2h. The volumes for irregular solids are generally determined by displacement of water using a graduated cylinder. The first example illustrates this. Ex. 1 A jeweler examines a 5.00 carat gemstone to determine if it is a diamond. She fills a 10 mL graduated cylinder with 6.00 mL of water. When she adds the stone, the volume reading increases to 6.18 mL. Is the gemstone likely to be a diamond, which has a density of 3.51

g ? cm3

Ans. Since 1 carat = 0.2 grams, the mass of the diamond is 5.00 carat x

0.2 g = 1.00 g 1 carat

The volume of the diamond is the difference between the final volume, 6.18 mL, and the initial volume, 6.00 mL. V = Vf – Vi = 6.18 mL – 6.00 mL = 0.18 mL Finally, density is found by dividing the mass by the volume. Laliberte, 8/25/08

12

Density = D =

m 1.00 g g = = 5.6 V 0.18 mL mL

Since 1 mL = 1 cm3, the density can be written as 5.6

g . The stone cm3

could be cubic zirconia, which has a density that ranges between 5.5 and 5.9

g . We do know, however, that the stone is not diamond. cm3

Ex. 2 A plastic rectangular block has a length of 24 cm, a width of 3.5 cm, and a height of 3.5 cm. The block has a mass of 278.4 g. Would it float in water? Ans. The volume is calculated as follows: V = lwh = (24 cm) x (3.5 cm) x (3.5 cm) = 294 cm3 Next, find the density. Density = D =

m 278.4 g g = . 3 = 0.95 V 294 cm cm3

An object will float in water if its density is less than that of water. Since water has a density of 1.00

g , the block will float. mL

Solving Chemistry Word Problems Solving chemistry word problems is considered one of the most difficult aspects of learning chemistry. The strategy below will help you organize your thoughts so you will become a better problem solver. List: Plan:

Laliberte, 8/25/08

Read the problem. Determine what information is given in the problem and what you are supposed to find. Write it down. The more familiar you are with units the easier this will be. Write down an equation that contains both the known and unknown information. Some problems require more than one step, and more than one equation. Map out a pathway to the solution. Write everything down.

13

Isolate: Solve: Answer:

Rearrange the equation to isolate the quantity you want to find. Write it down. Substitute the numerical values for the quantities in the equation, including units. Solve using a calculator. Write down you answer including units. Round off to the appropriate number of significant figures. Circle the answer. Note: Do not round off until you get to the final answer.

Here’s an example: What is the mass of 2.00 L of soda that has a density of 1.098 List:

m=?

Plan:

D=

m V

V = 2.00 L

D = 1.098

g ? mL

g mL

Volume of soda is in L but density uses mL. The

volume units need to be the same before using the density equation. 1000 mL = 1 L 1000 mL = 2000 L 1L

Isolate:

2.00 L x

Solve:

m = 2000 L) x (1.092

Answer:

2180 g

Laliberte, 8/25/08

m = VD

g ) = 2184 g mL

14