Chapter 1: Chemistry and Measurement Vanessa Prasad-Permaul Valencia College CHM 1045

Matter: Anything that has both mass & volume Properties of Matter Chemistry: The study of composition, properties, and transformations of matter Matter: Anything that has both mass & volume Hypothesis: Interpretation of results Theory: Consistent explanation of observations

Conservation of Mass Law of Mass Conservation: Mass is neither created nor destroyed in chemical reactions.

Example 1: Conservation of Mass C(s) + O2(g)  CO2(g) 12.3g C reacts with 32.8g O2, ?g CO2 12.3g + 32.8g = 45.1g 0.238g C reacts with ?g O2 to make .873g CO2 0.238g + x = 0.873g = 0.873g-0.238g = 0.635g of O2 ?g C reacts with 1.63g O2 to make 2.24g CO2 x + 1.63g = 2.24g = 2.24g - 1.63g = 0.61g C

Example 1: Conservation of Mass Exercise 1.1 1.85g of wood is placed with 9.45g of air in a sealed vessel. It is heated and the wood burns to produce ash and gases. The ash is weighed to yield 0.28g. What is the mass of the gases in the vessel? 1.85g Wood + 9.45g Air heat 0.28g Ash + ? g gases 1.85 + 9.45 - 0.28 = 11.02g of gases What is the mass of wood that is converted to gas by the end of the experiment? 1.85g of Wood – 0.28g of ash = 1.57g

Matter is any substance that has mass and occupies volume. Matter exists in one of three physical states: solid liquid gas

In a solid, the particles of matter are tightly packed together. Solids have a definite, fixed shape. Solids cannot be compressed and have a definite volume. Solids have the least energy of the three states of matter.

Liquids cannot be compressed and have a definite volume. In a liquid, the particles of matter are loosely packed and are free to move past one another. Liquids have an indefinite shape and assume the shape of their container. Liquids cannot be compressed and have a definite volume. Liquids have less energy than gases but more energy than solids.

Gases can be compressed and have an indefinite volume. In a gas, the particles of matter are far apart and uniformly distributed throughout the container. Gases have an indefinite shape and assume the shape of their container. Gases can be compressed and have an indefinite volume. Gases have the most energy of the three states of matter.

Phases

Properties of Matter A physical change is a change in the form of matter but not in its chemical identity A chemical change or a chemical reaction is a change in which one of more kinds of matter are transformed into a new kind of matter or several new kinds of matter

Properties of Matter Physical Properties can be determined without changing the chemical makeup of the sample. Some typical physical properties are: Melting Point, Boiling Point, Density, Mass, Touch, Taste, Temperature, Size, Color, Hardness, Conductivity. Some typical physical changes are: Melting, Freezing, Boiling, Condensation, Evaporation, Dissolving, Stretching, Bending, Breaking.

Some typical chemical properties are: Properties of Matter Chemical Properties are those that do change the chemical makeup of the sample. Some typical chemical properties are: Burning, Cooking, Rusting, Color change, Souring of milk, Ripening of fruit, Browning of apples, Taking a photograph, Digesting food. Note: Chemical properties are actually chemical changes

Properties of Matter Exercise 1.2 Potassium (K) is a soft, silvery-colored metal that melts @ 64oC. It reacts vigorously with water (H2O), Oxygen (O2) and Chlorine (Cl2). Identify all physical properties: Soft Silvery-colored Melting point of 64oC Identify all chemical properties: Metal (its chemical identity) K reacts vigorously with H2O K reacts vigorously with O2 K reacts vigorously with Cl2

Classifications of Matter Matter can be divided into two classes: mixtures pure substances Mixtures are composed of more than one substance and can be physically separated into its component substances. Pure substances are composed of only one substance and cannot be physically separated.

There are two types of pure substances: Compounds Elements A compound is a substance composed of two or more elements chemically combined Compounds can be chemically separated into individual elements. Water is a compound that can be separated into hydrogen and oxygen. An element cannot be broken down further by chemical reactions.

Dalton’s Atomic Theory Law of Definite Proportions: Different samples of a pure chemical substance always contain the same proportion of elements by mass. Any sample of H2O contains 2 hydrogen atoms for every oxygen atom

There are two types of mixtures: homogeneous mixtures heterogeneous mixtures Homogeneous mixtures have uniform properties throughout. Salt water is a homogeneous mixture. Heterogeneous mixtures do not have uniform properties throughout. Sand and water is a heterogeneous mixture.

Which of the following represents a mixture? Example 2: Matter Which of the following represents a mixture?

Accuracy, Precision, and Significant Figures in Measurement Accuracy is how close to the true value a given measurement is. Precision is how well a number of independent measurements agree with one another.

Example 8: Accuracy & Precision Which of the following is precise but not accurate?

Accuracy, Precision, and Significant Figures in Measurement Significant Figures are the total number of digits in the measurement. The results of calculations are only as reliable as the least precise measurement!! Rules exist to govern the use of significant figures after the measurements have been made.

Accuracy, Precision, and Significant Figures in Measurement Rules for Significant Figures: Zeros in the middle of a number are significant Zeros at the beginning of a number are not significant Zeros at the end of a number and following a period are significant Zeros at the end of a number and before a period may or may not be significant.

Example 4: Significant Figures How many Significant Figures ? a) 0.000459 = 3 b) 12.36 = 4 c) 36,450 = 4 d) 8.005 = 4 e) 28.050 = 5

Accuracy, Precision, and Significant Figures in Measurement Rules for Calculating Numbers: During multiplication or division, the answer can’t have more significant figures than any of the original numbers.

Example 5: Significant Figures 218.2 x 79 = 17237.8 = 1.7 x 104 12.5 / 0.1272 = 94.33962264150943 = 94.3 0.2895 x 0.29 = 0.083955 = 0.084 32.567 / 22.98 = 1.417188859878155 = 1.417

Accuracy, Precision, and Significant Figures in Measurement -During addition or subtraction, the answer can’t have more digits to the right of the decimal point than any of the original numbers.

Example 6: Significant Figures 218.2 + 79 = 297.2 = 297 12.5 - 0.1272 = 12.3728 = 12.4 0.2895 + 0.29 = 0.5795 = 0.58 32.567 - 22.98 = 55.547 = 55.55 185.5+2.224 = 187.724 = 187.7

Accuracy, Precision, and Significant Figures in Measurement Rules for Rounding Numbers: If the first digit removed is less than 5 round down (leave # same) If the first digit removed is 5 or greater round up Only final answers are rounded off, do not round intermediate calculations

Example 7: Rounding and Significant Figures Round off each of the following measurements a) 3.774499 L to 4 sig. figs. = 3.774L b) 255.0974 K to 3 sig. figs. = 255K c) 55.265 kg to 4 sig. figs. = 55.27kg d) 1.2151ml to 3 sig. figs. = 1.22ml e) 1.2143g to 3 sig. figs. = 1.21g

Significant Figures Exercise 1.3 Give answers to the following arithmetic setups. Round to the correct number of significant figures: a) 5.61 x 7.891 = 4.864671 = 4.9 9.1 b) 8.91 - 6.435 = 2.475 = 2.48 c) 6.81 – 6.730 = 0.08 = 0.08 d) 38.91 x (6.81-6.730) = 38.91 x 0.08 = 3.1128 = 3

How to put into calculator Scientific Notation Changing numbers into scientific notation Large # to small # Moving decimal place to left, positive exponent 123,987 = 1.23987 x 105 Small # to large # Moving decimal place to right, negative exponent 0.000239 = 2.39 x 10-4 How to put into calculator

Example 3: Scientific Notation Put into or take out of scientific notation 1973 = 1.973 x 103 5.5423 x 10-4 = 0.00055423 0.775 = 7.75 x 10-1 3.55 x 107 = 35,500,000 8500 = 8.5 x 103

Measurement and Units SI Units

Some prefixes for multiples of SI units Measurement and Units Some prefixes for multiples of SI units * * * * * * * * Important

Exercise 1.4 Express the following quantities using an SI prefix and a Measurement and Units Exercise 1.4 Express the following quantities using an SI prefix and a base unit. For instance, 1.6 x 10-6m = 1.6mm. A quantity such as0.000168g could be written 0.168mg or 168mg. a) 1.84 x 10-9 m = 1.84 nm (nanometer) b) 5.67 x 10-12 s = 5.67 ps (picosecond) c) 7.85 x 10-3 g = 7.85 mg (milligram) d) 9.7 x 103 m = 9.7 km (kilometer) e) 0.000732 s = 0.732 ms (millisecond) = 732us (microsecond) f) 0.000000000154 m = 0.154nm (nanometer) = 154pm (picometer)

Changes in Physical State Most substances can exist as either a solid, liquid, or gas. Water exists as a solid below 0 °C; as a liquid between 0 °C and 100 °C; and as a gas above 100°C. A substance can change physical states as the temperature changes.

When a solid changes to a liquid, the phase change is called melting. Solid  Liquid When a solid changes to a liquid, the phase change is called melting. A substance melts as the temperature increases. When a liquid changes to a solid, the phase change is called freezing. A substance freezes as the temperature decreases.

A substance vaporizes as the temperature increases. Liquid  Gas When a liquid changes to a gas, the phase change is called vaporization. A substance vaporizes as the temperature increases. When a gas changes to a liquid, the phase change is called condensation. A substance condenses as the temperature decreases.

When a solid changes directly to a gas, Solid  Gas When a solid changes directly to a gas, the phase change is called sublimation. A substance sublimes as the temperature increases. When a gas changes directly to a solid, the phase change is called deposition. A substance undergoes deposition as the temperature decreases.

Diagram of the various phases of temperature change

Temperature Conversions: The Kelvin and Celsius scales have equal size units (a change of 1oC is equivalent to a change of 1K) 180oF 100 oC 100 K

Temperature Conversions: Celsius (°C) — Kelvin (K) temperature conversion: Kelvin (K) = t°C x 1K + 273.15K 1oC Fahrenheit (°F) — Celsius (C) temperature conversions: there are exactly 9oF for every 5oC. Knowing that 0oC = 32oF tF = tC x 9oF + 32 5oC tC = 5oC x (toF – 32) 9oF

Example 9: Temp. Conversions Carry out the indicated temperature conversions: a) –78°C = ? K = (-78oC x 1K/1oC) +273.15K = 195.15 = 195K b) 158°C = ? °F = (158oC x 9oF/5oC)+32oF = 316.4 = 316oF c) 373.15 K = ? °C = (373.15K x 1oC/1K)– 273.15K = 100K d) 98.6°F = ? °C = 5oC/9oF x (98.6oF – 32oF) = 37oC e) 98.6°F = ? K = (37oC x 1K/1oC) +273.15K = 310.15 = 310K

A person with a fever has temperature of 102.5oF. Measurement and Units Exercise 1.5 A person with a fever has temperature of 102.5oF. What is this temperature in oC? A cooling mixture of dry ice and isopropyl alcohol has a temperature of -78 oC. What is the temperature in kelvins? a) oC = 5oC x (oF – 32 ) = 0.555 x (102.5 – 32) = 39.2oC 9oF b) K = oC + 273.15 = -78 + 273.15 = 195 K

Volume Volume: how much three-dimensional space a substance (solid, liquid, gas) or shape occupies or contains often quantified numerically using the SI derived unit (m3) the cubic meter. The volume of a container is generally understood to be the capacity of the container, i. e. the amount of fluid (gas or liquid) that the container could hold, rather than the amount of space the container itself displaces.

Traditionally chemists use liter (L) Volume units of Volume: m3 or cm3 (cc) Traditionally chemists use liter (L) 1cm3 = 1cc = 1mL

Density: relates the mass of an object to its volume. Measurement and Units Density: relates the mass of an object to its volume. Density = mass / Volume D = m / V V = m / D m = V  D Density decreases as a substance is heated because the substance’s volume increases.

Density What is the density of glass (in mL) if a sample weighing 26.43 g has a volume of 12.40 cm3? d = ? m = 26.43 g V = 12.40 cm3 = 12.40 mL d = m = 26.43 g = 2.13145 = 2.131 g/mL V 12.40 mL

Density What is the volume of an unknown solution if the mass is 12.567 g and the density is 14.621 g/mL ? d = m/V V x d = m V = m/d V = 12.567 g / 14.621 g/mL = 0.85952 mL 12.567g x 1mL = 0.85952 mL 14.621g

Density What is the mass of an unknown solution if the volume is 20.2 mL and the density is 2.613 g/mL? d = m/V m = d x V m = 2.613g x 20.2 mL = 52.7826 = 52.8 g mL

Exercise 1.6 A piece of metal wire has a volume of 20.2 cm3 and a Density Exercise 1.6 A piece of metal wire has a volume of 20.2 cm3 and a mass of 159 g. What is the density of the metal? D = m = 159 g = 7.87128712 = 7.87 g /cm3 V 20.2cm3 We know that the following metals have the following densities. Which metal is the wire made of? Mn = 7.21 g/cm3 Fe = 7.87 g/cm3 Ni = 8.90 g/cm3

Ethanol (grain alcohol) has a density of 0.789 g/cm3. Exercise 1.7 Ethanol (grain alcohol) has a density of 0.789 g/cm3. What volume (mL) of ethanol must be poured into a graduated cylinder to equal 30.3 g? d = m/V V x d = m V = m / d V = 30.3 g x 1 cm3 = 38.4cm3 0.789 g

Dimensional Analysis & Units Dimensional-Analysis method uses a conversion factor to express the relationship between units. Original quantity x conversion factor = equivalent quantity Example: express 2.50 kg  lb. Conversion factor: 1.00 kg = 2.205 lb 2.50 kg x 2.205 lb = 6.00 lb 1.00 kg

Dimensional Analysis & Units

The oxygen molecule (O2) consists of two oxygen Dimensional Analysis & Units Exercise 1.8 The oxygen molecule (O2) consists of two oxygen atoms a distance of 121 pm apart. How many millimeters (mm) is this distance? 121 pm x 10-12 m x 1mm = 1.21 x 10-7 mm 1 pm 10-3

Dimensional Analysis & Units Exercise 1.9 A large crystal is constructed by stacking small identical pieces of crystal. A unit cell is the smallest piece from which a crystal can be made. A unit cell of a crystal of gold metal has a volume of 67.6 A3. What is the volume in dm3? 67.6 A3 x 10-10 m x 10 dm = 6.76 x 10 -26 dm3 1 A3 1 m 3

1 in = 2.54cm (exactly) 1 yd = 36in (exactly) Dimensional Analysis & Units Exercise 1.10 Using the following definitions, obtain the conversion factor for yards to meters. How many meters are there in 3.54 yd? 1 in = 2.54cm (exactly) 1 yd = 36in (exactly) 1 yd x 1 in x 1 cm = 1.093613298 = 1.094 yd/m 36 in 2.54 cm 10-2 m 3.54 yd x 1 m = 3.24 m 1.094 yd

a) 1.267 km  m  cm b) 0.784 L  mL c) 3.67 x 105 cm  in Conversions 1.267km x 1000m x 100cm = 126700cm = 1.267 x 105 1km 1m b) 0.784 L  mL 0.784L x 1000mL = 784L 1L c) 3.67 x 105 cm  in 3.67 x 105cm x 1in = 144488.1889in = 1.44 x 105in 2.54cm

d) 79 oz  g e) 9.63 x 10-3 yd  ft f) 23.5 cm2  m2 Conversions 79oz x 28.35g = 2239.65g = 2.2 x 103g 1oz e) 9.63 x 10-3 yd  ft 9.63 x 10-3yd x 1m x 1km x 0.62137mile x 5280ft 1.0936yd 1000m 1km 1mile = 0.0289ft f) 23.5 cm2  m2 23.5cm2 x 1m2 = 0.235m2 100cm2

g) 1.34 x 1012 pm  m 10-12pm h) 4.67 x 10-7 nm  pm 10-9nm 1m Conversions g) 1.34 x 1012 pm  m 1.34 x 1012pm x 1m = 1.34 x 1024m 10-12pm h) 4.67 x 10-7 nm  pm 4.67 x10-7nm x 1m x 10-12pm = 4.67 x 10-12pm 10-9nm 1m