Chemistry Unit 1 Overview

Chemistry Unit 1

• Chemistry is the branch of science that identifies the substances matter is composed of. It also investigates their properties, interactions, combinations, changes, and uses these processes to create new substances.
• Physical properties include density, luster, malleability, elasticity, and melting point.
• Chemical properties include flammability, reactivity, nutritional information, and what a substance reacts with.
• Chemical changes involve chemical reactions, forming new bonds, moving electrons, accompanied by heat (fire), gas (smoke), or color changes.
• Physical changes involve a change of form without changing the chemical properties, like tearing paper or melting ice.
• Examples of chemical or physical changes include metal placed in water, tearing paper, digesting food, iron rusting, and water boiling.

Density

• Density is the physical property of matter, relating mass to volume (the space matter occupies).
• The density formula is D = Mass/Volume.
• Density is constant for a substance.
• Higher mass means higher volume.
• A direct proportion exists between mass and volume.
• Changes to the size of an object will affect the volume, not the density.
• Different substances have different densities due to their compositions.

Graphing and Algebraic Relationships

• Graphs are used to illustrate the relationship between mass and volume of objects in chemistry.
• The graph must plot mass against volume.
• Students need to choose appropriate scales for the axes (volume on the x-axis and mass on the y-axis).
• Calculate the density of each object.
• Describe any observations obtained from the graph about the objects.
• Identify the meaning of the slope of the line.

Density Problem

• Calculate the volume of a piece of aluminum foil with a given mass of 32.70 g.
• Follow the steps for doing calculations in chemistry:
  • write the equation
  • plug in numbers with units
  • report the answer with proper units and significant figures

Variables and Controls in Experiments

• Variables are aspects of an experiment that can change from one measurement to another.
• Controls are aspects of an experiment that remain constant.
• Students should identify potential variables in a density lab.

Measurements in Chemistry

• Measurements are accurate only as the measuring tool allows.
• Identify the units during measurements (e.g., grams, kilograms, milliliters).
• Know the graduations on the measuring tool (e.g., 1mL, 10mL, 100mL).
• Report one extra decimal point, representing an educated guess, based on the measurement tool.

Laboratory Tools

• Rulers are used to measure length in centimeter (cm).
• Gram scales are used to measure mass. The precise value displayed might vary slightly.

Accuracy and Precision

• Accuracy refers to how close a measurement is to the actual value.
• Precision refers to how close multiple measurements are to each other.
• Darts examples are used for illustrating accuracy and precision.
• Percentage error tells you how close your measured value is to the accepted value.

Observing and Collecting Data

• Observations involve using the senses (e.g., sight, hearing) to gather information about an object or event.
• Data can be qualitative (descriptive such as color, texture) or quantitative (numerical measurements).

Significant Figures

• Numbers used in scientific measurements provide information about measuring equipment and care taken.
• The more significant figures, the more precise the measurement; e.g., the instruments employed and estimations made impact the number of significant figures.
• Non-zero digits are significant.
• Zeros between non-zero digits are significant.
• Zeros in front of non-zero digits are not significant.
• Zeros at the end of a number after a decimal point are significant but zeros at the end of a non-decimal number may or may not be significant. Use scientific notation if needed.
• Exact numbers have no impact on significant figures in calculations

Significant Figures, Calculations

• Addition or subtraction should have the same number of decimal places (least decimal place) as the value with the fewest.
• Multiplication and division require the same number of significant figures (least number) as the value with the fewest significant figures.

Scientific Notation

• Scientific notation represents numbers as M × 10ⁿ, where M is a number between 1 and 10 and n is an exponent (whole number).
• A negative value for 'n' means a decimal fraction, as in 0.001 = 1 x 10^-3.
• A positive value of 'n' means a large number, such as 10000 = 1 x 10⁴.

Operations with Scientific Notation

• Calculations involving scientific notation use calculator functions (EE, EXP, x10ⁿ, etc).
• Numbers are entered first (e.g., 4.2), then the "x10n" button is pressed, followed by the exponent (e.g., 4).
• Avoid using the ^ or parenthesis when inputting.

Units of Measurement

• Metric system:
• conversion factors are ratios derived from equalities that convert one unit to another
• Dimensional analysis is a mathematical technique used to solve problems by using units.
• SI conversions use prefixes such as kilo, mega, giga, centi, milli, micro, nano, and pico. The prefixes are multiplied by factors of 10.
• Examples of SI base units: meter (m), kilogram (kg), second (s), Kelvin (K).
• Specific conversion factors convert one unit to another.

Mole

• A mole is a quantitative unit similar to a dozen or a ream.
• One mole is equal to Avogadro's number (6.02 x 10²³ particles).
• The periodic table shows the mass of a mole of each element, also called molar mass.

Avogadro's Number

• It is an extremely large numerical value that is related to a mole.
• Avogadro's number represents the number of atoms or molecules in one mole of a substance (6.022 x 10²³ entities per mole).
• Examples: use of Avogadro's number in comparison to the size of the Earth, etc.

Gram/Mole Conversions

• The molar mass of a substance is the mass of one mole of that substance.
• Molar masses are typically presented in g/mol units.
• A molar mass can be used as a conversion factor between grams and moles.
• Illustrate the relationship between moles and grams for specific examples.

Conversions, Continued

• Example problems to convert moles to grams.
• Example problems to express grams in milligrams and kilograms.
• Example problem to convert nanograms to milligrams.