Frederick A. Bettelheim William H. Brown Mary K. Campbell Shawn O. Farrell William H. Brown Beloit College Chapter 1 Matter, Energy, and Measurement

1-2 Chemistry Chemistry Chemistry is the study of matter. Matter Matter is anything that has mass and takes up space. Matter can change from one form to another. chemical change In a chemical change (chemical reaction), substances are used up and others formed in their place. Example: Example: When propane (bottled or LP gas) burns in air, propane and oxygen are converted to carbon dioxide and water. In a physical change, matter does not lose its identity. A common physical change is a change of state. Example: Example: Ice (solid water) melts to become liquid water; liquid water boils to become steam (gaseous water).

1-3 Scientific Method Fact Fact: A statement based on direct experience Hypothesis Hypothesis: A statement that is proposed, without actual proof, to explain a set of facts or their relationship. Theory Theory: The formulation of an apparent relationship among certain observed phenomena, which has been verified to some extent. In a sense, a theory is the same as a hypothesis except that we have a stronger belief in it because more evidence supports it. If, however, we find new evidence that conflicts with the theory, it must be altered or rejected.

1-4 Scientific Method

1-5 Exponential Notation Used to represent very large or very small numbers as powers of 10. Examples: is written as 2 x ,000,000 is written as 2 x 10 6

1-6 Metric System Table 1.1 Base units in the metric system

1-7 Metric System Table 1.2 The Most Common Metric Prefixes

1-8 Significant Figures 1. Nonzero digits are always significant. For example has four significant figures; 2.3g has two. 2. Zeros at the beginning of a number are never significant. For example L has two significant figures. 3. Zeros at the end of a number that contains a decimal point are always significant. For example 3.00L has three significant figures mm also has three. 4. Zeros at the end of of number that contains no decimal point may or may not be significant. For example $36,000 contains two significant figures and $36, contains seven significant figures. 5,000 mL contains one significant Figure and 5,000. mL contains four.

1-9 Atlantic-Pacific Rule for Sig Figs

1-10 Metric & English Systems Table 1.3 Some Conversion Factors

1-11 Mass and Weight Mass: Mass: The quantity of matter in an object. Mass is independent of location. Weight: Weight: The result of mass acted upon by gravity. Weight depends on location; depends on the force of gravity at the particular location.

1-12 Temperature Fahrenheit (F): Fahrenheit (F): Defined by setting the normal freezing point of water at 32°F and the normal boiling point of water at 212°F. Celsius (C): Celsius (C): Defined by setting the normal freezing point of water at 0°C and the normal boiling point of water at 100°C. Kelvin (K): Kelvin (K): Zero is the lowest possible temperature; also called the absolute scale. Kelvin degree is the same size as a Celsius degree K = °C + 273

1-13 Factor-Label Method Conversion factor Conversion factor: A ratio of two different units, used as a multiplier to change from one system or unit to another. For example, 1 lb = g Example: Example: Convert 381 grams to pounds. Example: Example: Convert gallons to milliliters.

1-14 The Three States of Matter Gas Has no definite shape or volume. Expands to fill whatever container it is put into. Is highly compressible.Liquid Has no definite shape but a definite volume. Is only slightly compressible.Solid Has a definite shape and volume. Is essentially incompressible.

1-15 Density Density: Density: The ratio of mass to volume. The most commonly used units are g/mL for liquids and solids, and g/L for gases. Example: Example: If 73.2 mL of a liquid has a mass of 61.5 g, what is its density in g/mL?

1-16 Specific Gravity Specific gravity: Specific gravity: The density of a substance compared to water as a standard. Because specific gravity is the ratio of two densities, it has no units (it is dimensionless). Example: Example: The density of copper at 20°C is 8.92 g/mL. The density of water at this temperature is 1.00 g/mL. What is the specific gravity of copper?

1-17 Energy Energy: Energy: The capacity to do work. May be either kinetic energy or potential energy The calorie (cal) is the base metric unit of energy. Kinetic energy (KE): Kinetic energy (KE): the energy of motion KE increases as the object’s velocity increases. At the same velocity, a heavier object has greater KE. Potential energy: Potential energy: The energy an object has because of its position; stored energy

1-18 Energy Examples of kinetic energy are mechanical energy, light, heat, and electrical energy. In chemistry, the most important forms of potential energy are chemical energy and nuclear energy. Chemical energy is stored in chemical substances as, for example, in foods such as carbohydrates and fats. It is given off when substances take part in chemical reactions. The law of conservation of energy Energy can neither be created nor destroyed. Energy can be converted from one form to another.

1-19 Heat and Temperature Heat is a form of energy. Heating refers to the energy transfer process when two objects of different temperature are brought into contact. Heat energy always flows from the hotter object to the cooler one until the two have the same temperature. Heat is commonly measured in calories (cal), which is the heat necessary to raise the temperature of 1 g of liquid water by 1°C.

1-20 Specific Heat Specific heat (SH): Specific heat (SH): The amount of heat necessary to raise the temperature of 1.00 g of a substance by 1.00°C. Table 1.4 Specific Heats of Common Substances

1-21 Specific Heat The following equation gives the relationship between specific heat, amount of heat, the mass of an object, and the change in temperature. Example: Example: How many calories are required to heat 352 g of water from 23°C to 95°C?

1-22 Chapter 1 Matter, Energy, and Measurement End Chapter 1