CHEM101: General Chemistry I

1a. Classify properties of matter and changes of matter as physical or chemical

• How do physical and chemical changes differ?
• What are some examples of physical changes?
• What are some examples of chemical changes?

Matter describes everything around us that has mass; chemistry is the academic discipline that studies matter. This includes solid objects, such as the table where we sit, the water (a liquid) we drink, and the air (a gas) we breathe. Chemistry is part of everything we touch, feel, see and smell.

Chemistry studies the properties and structure of matter, including chemical reactions, which describe the transformation of matter. Many people call chemistry the central science because we use it in most science and technology fields.

Scientists classify properties of matter as physical or chemical. Physical properties are those we can observe without altering the identity of the substance. For example, the melting or boiling point of a substance is a physical property because they do not alter the identity of the substance.

We can only observe chemical properties when we alter the identity of the substance. Rusting is an example of a chemical property because it is a chemical reaction that changes the composition of the substance.

Matter can undergo two types of changes: physical and chemical.

Physical changes do not alter or change the identity of a substance. Examples of physical transformations include freezing, melting, or boiling.

For example, when a solid ice cube melts to become liquid water, the water's chemical identity does not change. The transformation is physical because the identity of the initial and final substance remains the same, in this case water. Likewise, water vapor (gas) that results from boiling is still composed of water molecules.

Chemical changes, on the other hand, which we also call chemical transformations and chemical reactions, do alter the identity of the substance.

For example, a nail that rusts represents a chemical change because the rusting process creates a new substance that has a different chemical composition than the original nail. The transformation is chemical because the identity of the substance has changed. Similarly, burning represents a chemical change because the chemical composition of the substance changes during the burning process.

Review this material in Physical and Chemical Properties and Physical and Chemical Changes.

1b. Name and use SI units for length, mass, time, and volume

• Can you list the base SI units for length, mass, time, and volume?
• Can you recognize and use the standard SI decimal prefixes to convert among different units of measure?

Systeme Internationale (SI) units describe the internationally-recognized set of units of measurement that are standard in all scientific fields.

The base SI units for length, mass, time, and volume are as follows:

We use SI prefixes to convert among units that have different orders of magnitude. For example, you should use millimeters to measure lengths that are extremely short, and centimeters, meters, or kilometers, to measure longer units of measurement or distances.

We use dimensional analysis to convert among units since it makes it easier to compare quantities in different units. For example, from the SI decimal prefixes table we see that one kilogram (1 kg) = $10^{3}$ g, or 1 kg = 1,000 g.

Review this material in Units of Measure and Units and Dimensions in Chemistry.

1c. Perform mathematical operations involving significant figures

• Can you add and subtract using significant figures?
• Can you multiply and divide using significant figures?
• Can you perform logarithm calculations using significant figures?

Chemists often need to perform calculations on the quantities they have measured. They use significant figures to convey their level of confidence (or level of accuracy), in their measurements and follow specific rules for adding and subtracting, multiplying, and dividing quantities with significant figures.

For addition and subtraction, we determine the answer's number of significant figures by decimal places. Look at your input quantities and identify the quantity that has the fewest number of decimal places.

Line your addition or subtraction up vertically, according to the decimal point, to make this more clear. Your answer should have the same number of decimal places as the input quantity that had the fewest number of decimal places.

For multiplication and division, your answer should have the same number of significant figures as the input quantity that had the fewest number of significant figures. Notice how the leading zeros in front of the nonzero digits are insignificant figures. This is shown in the 0.000029 example below.

For base 10 logarithms, the answer will have the same number of significant figures as the normalized form of the logarithm. Normalized means the logarithm is given in scientific notation $a\times 10^{b}$ where a is a number greater than one, and less than 10.

Review this material in Significant Figures, Using Significant Figures and Significant Figure Practice.

1d. Convert measurements into scientific notation

• How do you convert quantities into scientific notation?

Scientific notation allows scientists and mathematicians to express small and large numbers more succinctly because they do not include all of the zeros in their notations and conversions. For example, scientists frequently use scientific notation when making a dimensional analysis, to convert measurements from one unit to another.

Scientific notation uses multipliers of $a\times 10^{n}$, where $a$ represents the part of the number that includes non-zero numbers. The decimal point is moved to follow the first non-zero number, and $n$ represents the number of zeros that precede or follow the first non-zero number.

For example, 15,000 in scientific notation is $1.5\times 10^{4}$. In this case, we move the decimal point to follow the first non-zero digit. Then count the number of digits that follow the first non-zero digit to get $10^{4}$.

Similarly, 0.0007005 in scientific notation is a\times $7.005\times 10^{-4}$. In this case, we move the decimal point to follow the first non-zero digit to get 7.005. Count back to the original decimal point to determine the number of zero digits before the first non-zero number. Since you can count four digits until you hit the original decimal point, the multiplier is $10^{-4}$. Be sure to use a negative exponent when the original number is less than one.

Review this material in Introduction to Scientific Notation.

1e. Perform dimensional analysis conversions between different units of measure

• Can you convert feet to inches?
• How many minutes are in one day?

We use dimensional analysis to convert among units since it makes it easier to compare quantities in different units. For example, from the SI decimal prefixes table above we see that one kilogram (1 kg) = $10^{3}$ g, or 1 kg = 1,000 g.

Dimensional analysis allows chemists to convert from a given unit of measurement to a desired unit of measurement. For example, if you are given 12 inches and want to convert inches to meters, you would use the following dimensional analysis conversion:

Review this material in Convert Units.

1f. Perform calculations involving density

• What is the density of an object having a volume of 100 mL and a mass of 10 g?
• What is the density of water at 4°C?

Density measures how tightly packed the particles in a substance are.

We define density ($d$) as the mass or volume of a substance at a given temperature. We write $d = m/v$ where $d$ is density, $m$ is mass, and $v$ is volume. If we know two of the variables in this equation, we can solve for the third algebraically. The units for density are a mass unit divided by a volume unit. The units used to describe density often differ for the phases of matter: solids (g/cm³), liquids (g/mL), and gases (g/L).

Density measures the mass of an object per unit volume. In other words, a denser object has a higher mass than a second object that shares the same volume. Most liquids and solids have significantly higher densities than gases.

Review this material in Density and Its Uses and Calculations Using Density.

Unit 1 Vocabulary

• Central science
• Chemical
• Chemical change
• Chemical reaction
• Chemical property
• Chemical transformation
• Chemistry
• Density
• Dimensional analysis
• Gas
• Insignificant figures
• Liquid
• Mass
• Matter
• Measurement standard
• Physical
• Physical change
• Physical property
• Physical transformation
• Scientific notation
• Significant figures
• SI units
• Solid
• Unit