Unit 1 Study Guide and Review: Matter and Measurements

1a: Define chemistry.

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 liquid water 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.

 

1b: Distinguish between physical and chemical properties of matter.

We can describe matter by its physical and chemical properties.

    • How do the physical and chemical properties of matter differ?
    • List some examples of physical properties.
    • List some examples of chemical properties.

Physical properties describe our observations about the properties of a substance in which the substance itself does not change during or after our observation. Examples of physical properties we use to describe a substance include its boiling point, melting point, appearance, and density.

Chemical properties describe our observations about the properties of a substance during or after its chemical transformation. Examples of chemical properties we use to describe a substance include its acidity and the types of chemical reactions it can withstand or perform.

For this unit, be sure to review the table in section three which compares physical and chemical properties of the element sodium.

 

1c: Classify changes of matter as physical or chemical.

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

    • How do physical and chemical changes differ?
    • List some examples of physical changes.
    • List some examples of chemical changes.

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 comprised 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.

 

1d: Explain the solid, liquid, and gas states in terms of particles.

We can describe three states or phases of matter: solid, liquid and gas.

    • Describe and draw a picture of solid, liquid, and gas phase particles.
    • Which state or phase of matter usually has the lowest density?
    • Which state or phase of matter is considered a “condensed” state?

Molecules in solid state matter are arranged in an “ordered” fashion. The particles touch each other and are close together. Solid materials have a definite shape.

Molecules in liquid state matter are close together, but are not ordered like a solid. The particles can move or “slip” around each other. The liquids flow and move. Liquid materials take the shape of their container.

Molecules in gas state matter are far apart from one another and usually do not touch or interact. The molecules are completely disordered. Molecules in a gas take up the entire space of their container.

Review this simple diagram of the three states of matter.

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

We consider molecules in a liquid or solid phase condensed phase matter. This means the molecules are in direct contact with their neighboring molecules.

 

1e: Distinguish among a quantity, a unit, and a measurement standard.

In chemistry we use a defined set of units when we make measurements—to ensure consistency, accuracy, and make comparisons. We base units of measurement on a commonly-accepted, standard scale, so we can describe and communicate the results of our measurements with other researchers.

    • Define the following terms: quantity, unit, measurement standard.

Quantity describes the amount we measure in an experiment. For example, a quantity could describe the mass, volume, length, or another observable measure.

A unit relates to the standard measurement scale. A unit defines the amount of a quantity measured. For example, we use meters to measure length (50 meters), grams to measure mass (10 grams), and liters to measure volume (5 liters), according to the metric scale of measurement.

A measurement standard is a universally-agreed-upon object that defines a unit of measurement.

Measuring instruments are calibrated to match measurement standards scientists have agreed to follow. Think about the length of a ruler (how long is 12 inches or one meter?), a mutually-agreed-upon amount of water used to follow a recipe (how much is in a cup?), or a common temperature used to calibrate a thermometer (how cold is 20 degrees celsius or fahrenheit?).

 

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

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

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

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

Quantity

SI Unit

Length

meter (m)

Mass

kilogram (kg)

Time

second (s)

Volume (not actually SI)

Liter (l)


Review the table of SI decimal prefixes in “Section 2: The SI Units.”

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) = 103 g, or 1 kg = 1,000 g.

 

1g: Determine the number of significant figures in measurements.

You should use significant figures when reporting a measurement to convey your level of confidence in your measurement.

    • Explain why you need to use significant figures when reporting measured quantities.
    • Determine the number of significant figures in a given measurement.
    • Properly round numbers to a given number of significant figures.

Chemists consider the last digit in a measurement uncertain because it is an approximation. Significant figures tell us “how good” or “confident” your measurement is, according to the equipment you used to make the measurement.

For example, you usually have to make an approximation when you measure the distance between the smallest markings on a ruler. You might say the coin in this picture measures 2.7 centimeters, but you really have to make an educated guess about the last figure since it is not exact. Scientists recognize this last digit is often uncertain (in other words, you have less confidence in its accuracy).

Image of a coin and a measuring ruler: 2.7??? cm.

Review these general rules for determining the number of significant figures in a measured quantity:

  1. Leading zeros (single zeros that appear to the left of a decimal point) are not significant (for example, 0.35).
  2. Zeros that appear to the right of a nonzero number, before the decimal point, are significant (for example, 50.35).
  3. Trailing zeros (any zeros that appear to the right of a decimal point) are significant (for example, .350.
  4. Trailing zeros without a decimal point are ambiguous (for example, 350).

Here is an example.

For the number 0.003900270, the leading zero (the single zero to the left of the decimal point) is not significant. The trailing zero is significant because it follows the decimal point. Therefore, this number has seven significant figures (indicated in red in the diagram below).

(Note that in this diagram, the "placeholder" zeros mean they are showing the order of magnitude of the number rather than an actual measured value. So, the leading zeros in this number indicate the order of magnitude of the number and are not significant measured numbers.)

For the number 9024000.0, all of the zeros are significant because all zeros are between nonzero numbers, between nonzero numbers and a decimal point, or trailing after a decimal point. This number has eight significant figures. The last example, 9024000, has ambiguous trailing zeros.

Image with Significant, Ambiguous Numbers

When working with measured quantities, you should round the numbers properly to avoid confusion about the confidence you have in the measurement.

Review the rules for rounding significant figures in the yellow boxes in Section 2: Rules for Rounding.

When rounding, round down if the first insignificant digit is less than five. Round up if the first insignificant figure is greater than five.

For example, to round 45.556 to four significant figures, the number should become 45.56 because the first insignificant digit (the last digit, six, in this case) is greater than five.

When rounding the answer for a multi-step problem, it is important to keep track of significant figures, but you should not round your number until after you have completed all of your calculations.

 

1h: Perform mathematical operations involving 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.

    • Perform addition and subtraction using significant figures.
    • Perform multiplication and division using significant figures.
    • Perform logarithm calculations using 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.

Image of How to Line up Numbers by Decimal Point

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.

Image of How to Multiply and Divide Numbers

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 x 10b where a is a number greater than one, and less than 10.

Image of Number of Digits and Normalized Form of Logarithm

 

1i: Convert measurements 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.

    • Convert quantities into scientific notation.

Scientific notation uses multipliers of a x 10n, where “a” represents the part of the number that includes nonzeros. The decimal point is moved to follow the first nonzero number, and n represents the number of zeros that precede or follow the first nonzero number.

For example, 15,000 in scientific notation is 1.5 x 104. In this case, we move the decimal point to follow the first nonzero digit. Then count the number of digits that follow the first nonzero digit to get 104.

Similarly, 0.0007005 in scientific notation is 7.005 x 10-4. In this case, we move the decimal point to follow the first nonzero digit to get 7.005. Count back to the original decimal point to determine the number of zero digits before the first nonzero 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.

 

Unit 1 Vocabulary

      • Central science
      • Chemical change
      • Chemical reaction
      • Chemical transformation
      • Chemical property
      • Chemistry
      • Condensed state of matter
      • Density
      • Dimensional analysis
      • Gas
      • Liquid
      • Mass
      • Measurement standard
      • Physical property
      • Physical change
      • Physical transformation
      • Quantity
      • Scientific notation
      • Significant figures
      • Insignificant figures
      • SI units
      • Solid
      • Unit
Last modified: Tuesday, June 11, 2019, 8:50 AM