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.

Explore the foundational concepts of Chemistry in Unit 1. This quiz covers key topics such as physical and chemical properties, changes, and the concept of density. Test your understanding of how substances interact and transform.