Untitled Quiz

Chapter 1: Matter, Measurement, and Problem Solving

• Chemistry is a molecular approach, 2nd edition, by Nivaldo Tro.
• The book is authored by Roy Kennedy, from Massachusetts Bay Community College, Wellesley Hills, MA.
• Copyright is 2011 Pearson Education, Inc.
• The chapter focuses on matter, measurement, and problem solving.

Composition of Matter

• Matter is composed of atoms and molecules.
• The scientific method is used to study and understand matter.

Structure Determines Properties

• The properties of matter are determined by the atoms and molecules that compose it.
• Carbon monoxide is composed of one carbon atom and one oxygen atom; it's colorless, odorless, burns with a blue flame, and binds to hemoglobin.
• Carbon dioxide is composed of one carbon atom and two oxygen atoms; it's colorless, odorless, incombustible, and does not bind to hemoglobin.

Atoms and Molecules

• Atoms are submicroscopic particles.
• Molecules are two or more atoms attached together in a specific geometrical arrangement.
• The attachments between atoms are called bonds.
• Bonds have different strengths and come in different shapes and patterns.
• Chemistry is the study of matter.

The Scientific Approach to Knowledge

• Philosophers try to understand the universe through reasoning and thinking about ideal behavior.
• Scientists try to understand the universe through empirical knowledge gained through observation and experiment.

Gathering Empirical Knowledge - Observation

• Some observations describe the characteristics or behavior of nature (qualitative).
• Example: "The soda pop is a liquid with a brown color and a sweet taste. Bubbles are seen floating up through it."
• Some observations compare a characteristic to a standard numerical scale (quantitative).
• Example: "A 240 mL serving of soda pop contains 27 g of sugar."

From Observation to Understanding

• A hypothesis is a tentative interpretation or explanation for an observation.
• A good hypothesis is falsifiable.
• It's possible to prove a hypothesis wrong with one test.
• Example: "The sweet taste of soda pop is due to the presence of sugar."

Testing Ideas

• Ideas in science are tested with experiments.
• An experiment is a set of highly controlled procedures designed to test whether an idea about nature is valid.
• Experiments generate observations that will either validate or invalidate the idea.

From Specific to General Observations

• A scientific law summarizes all past observations and predicts future observations.
• Law of Conservation of Mass: "In a chemical reaction, matter is neither created nor destroyed."
• A scientific law allows prediction of future observations, allowing testing with experiments.
• Unlike state laws, scientific laws are not choices to be violated.

From Specific to General Understanding

• A hypothesis is a potential explanation for a single or small number of observations.
• A scientific theory is a general explanation for why things in nature are the way they are and behave the way they do.
• Theories use models, are the pinnacle of scientific knowledge, and are validated or invalidated by experiment and observation.

Scientific Method

• The scientific method involves observations, experiments, and the development of hypotheses, theories, and laws.
• Observations lead to a hypothesis which tests an idea by conducting experiments that can confirm or falsify that hypothesis.
• Hypotheses that are correct and repeatedly validated through experiments form basis for laws and theories.

Relationships Between Pieces of the Scientific Method

• Observations describe what happens.
• Laws summarize what happens.
• Hypotheses explain why things happen.
• Theories explain why things happen in general.

Classification of Matter

• Matter is anything that occupies space and has mass.
• Matter can be classified by its state (solid, liquid, gas) and its composition.

Classifying Matter by Physical State

• Matter can be classified as solid, liquid, or gas based on observable characteristics.
• Solids have a rigid shape and volume that do not change when placed in a new container; they are not compressible and do not flow.
• Liquids take the shape of their container and maintain their volume; they are not compressible but they flow.
• Gases take both the shape and volume of their containers; they are compressible and flow.

Solids

• The particles are packed closely together and are fixed in position, although they may vibrate.
• The close packing of the particles makes solids incompressible.
• The inability of the particles to move around makes solids retain their shape and volume.

Crystalline Solids

• Some solids have their particles arranged in repeating patterns – called crystalline solids.

Amorphous Solids

• Some solids have their particles arranged randomly distributed without any long-range pattern – called amorphous solids.

Liquids

• The particles in a liquid are closely packed, but they have some ability to move around.
• The close packing means that liquids are incompressible.
• The ability of the particles to move around allows liquids to take the shape of their container and flow.

Gases

• In the gas state, particles have freedom of motion and are not held together.
• The particles are constantly moving around, bumping into each other and the container.
• In a gas there is a lot of empty space between the particles on average.
• Gases are compressible because there is a lot of empty space, particles can be squeezed closer together.

Classifying Matter by Composition

• Another way to classify matter is by examining its composition.
• Composition includes the types of particles, the arrangement of the particles, and the attractions or attachments between them.
• Pure substances, such as elements and compounds, have a fixed composition.

Classification of Pure Substances - Elements

• Pure substances that cannot be decomposed into simpler substances by chemical reactions are called elements.

Classification of Pure Substances - Compounds

• Pure substances that CAN be decomposed into simpler substances are called compounds.
• Compounds are chemical combinations of elements.
• Molecules of a compound are identical.
• All samples of a compound behave the same way.

Classification of Mixtures

• Homogeneous mixtures have uniform composition throughout.
• Heterogeneous mixtures do not have uniform composition throughout.

Changes in Matter

• Changes that alter the state or appearance but not the composition are physical changes.
• Changes that alter the composition are chemical changes.

Physical Changes in Matter

• The boiling of water is a physical change.
• The water molecules are separated, but their structure and composition do not change.

Chemical Changes in Matter

• The rusting of iron is a chemical change.
• The atoms in the nail combine with oxygen to form a new substance called rust, a compound with a different composition.

Properties of Matter

• Physical properties are characteristics that can be changed without changing the composition of matter.
• Chemical properties are characteristics that determine how matter changes composition as a result of contact with other matter, or the influence of energy.

Common Physical Changes

• State changes (boiling, condensing, melting, freezing, subliming)
• Dissolving

Common Chemical Changes

• Rusting
• Burning
• Dyes fading/changing color

Energy

• Matter possesses energy.
• Energy is classified as kinetic or potential.
• Energy is converted between forms during chemical and physical changes.

Energy of Matter - Kinetic Energy

• Kinetic energy is energy of motion.
• Thermal energy (heat) is a form of kinetic energy.

Energy of Matter - Potential Energy

• Potential energy is stored energy.
• Chemical potential energy arises from electrostatic attractive forces between atoms, molecules, and subatomic particles.

Conversion of Energy

• You can interconvert kinetic and potential energy.
• The total amount of energy remains the same regardless of conversion.

Spontaneous Processes

• Materials with high potential energy tend to be less stable.
• Processes in nature tend to occur on their own when the result is material with lower total potential energy.

Changes in Energy

• When a process results in less potential energy, the lost potential energy is converted to kinetic energy and released to the environment.
• Released energy can be harnessed to do work.

Potential to Kinetic Energy

• Conversion of potential energy to kinetic energy can be observed in various mechanisms such as a car running on gasoline.

Standard Units of Measure

• The SI system is a set of international standard units for comparison of measurements.

The Standard Units

• Measurements are expressed using metric units.

Length

• Sl unit is meter.
• Commonly use centimeters.

Mass

• Sl unit is kilogram.
• Commonly measure in grams or milligrams.

Time

• Sl unit is second.
• Defined as period of time for a specific number of radiation events in cesium-133.

Temperature

• Measure of average kinetic energy.
• Heat flows from high to low thermal energy until they reach the same temperature.

Temperature Scales

• Fahrenheit, Celsius, and Kelvin are different temperature scales.

Fahrenheit vs. Celsius

• A Celsius degree is 1.8 times larger than a Fahrenheit degree.
• The standard for 0° on the Fahrenheit scale is a lower temperature than the standard used for 0° on the Celsius scale.

Kelvin vs. Celsius

• The degree size is the same as in Celsius.
• Kelvin begins at absolute zero (much lower than the Celsius scale).

Example: Converting Temperature

• Converting temperatures between Celsius, Kelvin, and Fahrenheit scales involves using specific equations.

Related Units in the SI System

• All SI units are related to a standard unit by a power of 10.
• Prefix multipliers are always the same, regardless of the standard unit.

Common Prefix Multipliers in the SI System

• Common prefix multipliers such as mega-, kilo-, deci-, centi-, milli-, micro-, nano-, and pico- are used in measurement.

Volume

• Derived unit—any length unit cubed.
• Measure of the amount of space occupied.
• Sl unit is cubic meter; commonly use cubic centimeters.

Common Units and Equivalents

• Provides equivalency statements for length, mass, and volume measurements using different units.

Density

• Density is the ratio of mass to volume.
• It is an intensive property.
• For solids = g/cm³.
• For liquids = g/mL.
• For gases = g/L.

Density as a Conversion Factor

• Density can be used to convert between mass and volume.

Example: Calculating Density

• Calculating density for known masses and volumes.

Significant Figures

• Significant figures in measurements tell the range of values to expect for repeated measurements.
• Non-zero digits are always significant.
• Interior zeros (zeros between other digits) are always significant.
• Leading zeros (zeros at the beginning of a number) are never significant.
• Trailing zeros (zeros at the end of a number) are significant only when a decimal point is explicitly shown.

Counting Significant Figures

• Rules for determining the number of significant figures in a measurement, whether exact or estimated.

Significant Figures and Exact Numbers

• Numbers whose values are known with complete certainty are called exact numbers.
• Exact numbers have an unlimited number of significant figures.

Example: Determining Significant Figures

• Demonstrating how to determine the number of significant figures in different measurements.

Practice: Determine Significant Figure

• Exercise involving determining the significant figures in a set of measurements, along with the range of precision.

Multiplication and Division with Significant Figures

• When multiplying or dividing, the result will have the same number of significant figures as the measurement with the fewest significant figures.

Addition and Subtraction with Significant Figures

• When adding or subtracting, the result will have the same number of decimal places as the measurement with the fewest decimal places.

Rounding

• Rules and principles for rounding numbers to the correct number of significant figures.

Precision and Accuracy

• Precision refers to how close repeated measurements are to each other.
• Accuracy refers to how close a measurement is to the actual value.

Uncertainty in Measured Numbers

• Measurements have limitations, leading to uncertainties in the values measured.

Precision

• Imprecision is caused by random errors.
• Random errors result from random fluctuations, so cannot be corrected.
• We evaluate precision of measurements by evaluating how far they are from each other and the actual value.

Accuracy

• Inaccuracy in measurements is caused by systematic errors.
• Systematic errors are caused by limitations in instruments or techniques.
• These errors can be reduced by using more accurate instruments or more refined procedures.

Accuracy vs. Precision

• Examination of how to determine accuracy and precision, as well as how to avoid systematic error.

Solving Chemical Problems; Equations & Dimensional Analysis

• Solving problems using equations and dimensional analysis.
• Units are important when doing calculations as well as the numbers.
• Dimensional analysis provides a guide to problem solving that involves the units in the relationships between different measurements or parameters in calculations.

Problem Solving and Dimensional Analysis

• Many chemistry problems need relationships to convert measurement units.
• Conversion factors can be used to convert between units.

Conceptual Plans

• A visual outline illustrating the methods to a problem.
• Helps understand the steps needed to obtain desired quantities from given quantities.

Conceptual Plans and Conversion Factors

• Procedures to convert units such as inches to centimeters.
• Conversion factors are derived from equivalency statements.

Systematic Approach to Problem Solving

• Steps for solving problems involves sorting information, strategizing (conceptual plan), applying steps in the conceptual plan, and checking the answer with proper units.

Example: Unit Conversion

• Shows the way of performing a unit conversion step-by-step.
• Illustrates examples of problems solved in the conceptual plan.

Practice: Unit Conversion

• Practice converting mL to quarts.

Conceptual Plans for Units Raised to Powers

• Methods for converting a cube unit example: cubic inches to cubic centimeters.

Example: Unit Conversion with Powers

• Shows how to perform a unit conversion step-by-step.
• Illustrates examples of problems solved in the conceptual plan.

Practice: Unit Conversion with Powers

• Demonstrating how to perform unit conversions with raised powers.

Density as a Conversion Factor

• Density is used as a conversion factor between mass and volume.

Example: Density as a Conversion Factor

• How to use density to calculate mass from volume, or volume from mass.

Example: Calculation Using Density

• How to use density as a conversion factor given known mass and volume.

Practice: Calculation Using Density

Order of Magnitude Estimations

• Method for quickly approximating the magnitude of an answer by rounding numbers.
• Emphasizes identifying the proper exponent in scientific notation.

Estimate of Answers

• Approximating the magnitude of answers by rounding during the calculation.

Problem Solving with Equations

• Solving problems that require using equation-based conceptual plans.
• The plan will give all variables except one, then substitute and compute.

Example: Calculating Density using Equations

• Procedure for calculating density when given mass, length and radius.

Practice: Calculating Mass

• Example problems to calculate mass when given density and volume, following an organized approach.