Chapter 1: Matter, Measurement, and Problem Solving – Study Notes

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Chapter 1: Matter, Measurement, and Problem Solving

Introduction

This chapter introduces the foundational concepts of chemistry, focusing on the nature of matter, its classification, and the scientific methods used to study it. Students will learn about the properties of matter, measurement systems, and the importance of significant figures and uncertainty in scientific data.

Matter and Its Classification

What is Matter?

Matter is anything that occupies space and has mass.
All physical objects, including living organisms, are composed of matter.

Atoms and Molecules

Atoms are the smallest units of elements that retain the properties of the element.
Molecules are combinations of two or more atoms held together in a specific shape.
Chemistry is the science that studies the behavior of matter by examining atoms and molecules.
Example: Water (H2O) and hydrogen peroxide (H2O2) are both molecules but have different properties due to their different structures.

Classification of Matter

Matter can be classified by its state (solid, liquid, gas) and composition (element, compound, mixture).

States of Matter

Solid: Definite shape and volume, incompressible.
Liquid: Definite volume, no definite shape, incompressible.
Gas: No definite shape or volume, compressible.

Classification by Composition

Pure Substance: Made up of only one component; composition does not vary.
Element: Cannot be broken down into simpler substances.
Compound: Composed of two or more elements in a fixed proportion.
Mixture: Composed of two or more components; composition can vary.
Heterogeneous Mixture: Composition varies from one region to another.
Homogeneous Mixture (Solution): Uniform composition throughout.

Examples of Classification

Granite: Heterogeneous mixture
Mercury: Element
Brass: Homogeneous mixture (alloy)
Sweet tea: Homogeneous mixture

Physical and Chemical Properties and Changes

Physical vs. Chemical Properties

Physical properties: Observed without changing composition (e.g., color, density, melting point).
Chemical properties: Observed only by changing composition (e.g., flammability, reactivity).

Physical vs. Chemical Changes

Physical change: Alters state or appearance, not composition (e.g., melting ice).
Chemical change: Alters composition; atoms rearrange to form new substances (e.g., rusting iron).

Intensive and Extensive Properties

Intensive property: Independent of amount (e.g., temperature, density).
Extensive property: Dependent on amount (e.g., mass, volume).

The Scientific Approach to Knowledge

Observation: Gathering data about the world.
Hypothesis: Tentative explanation of observations.
Law: Describes how nature behaves (often mathematical).
Theory: Model explaining why nature behaves as it does.

Measurement and Units

Units of Measure

Standard quantities used to specify measurements.
Two main systems: Imperial System and Metric System.
SI Units (Système International d’Unités): Standard in science, based on the metric system.

Quantity	Unit	Symbol
Length	Meter	m
Mass	Kilogram	kg
Time	Second	s
Temperature	Kelvin	K
Amount of substance	Mole	mol
Electric current	Ampere	A
Luminous intensity	Candela	cd

Temperature Scales

Celsius (°C), Fahrenheit (°F), Kelvin (K)
Conversion equations:

SI Prefix Multipliers

Prefix	Symbol	Multiplier
kilo	k	1,000
centi	c	0.01
milli	m	0.001
micro	μ	0.000001
nano	n	0.000000001

Problem Solving and Dimensional Analysis

General Problem-Solving Strategy

Identify the starting point (given information).
Identify the end point (what you must find).
Devise a plan to get from start to end using known relationships (conceptual plan).
Dimensional analysis uses units as a guide to solving problems.

Conversion Factors

Ratios used to convert between units (e.g., 1 inch = 2.54 cm).
Exact numbers do not affect significant figures.

Units Raised to a Power

When converting squared or cubed units, raise both the number and the unit to the power.
Example:

Derived Units

Definition and Examples

Derived units are combinations of base units (e.g., speed = m/s, volume = m3).
1 cm3 = 1 mL; 1000 mL = 1 L (not SI).

Density

Density is the ratio of mass to volume.
For solids: typically g/cm3; for liquids: g/mL; for gases: g/L.
Formula:

Substance	Density (g/cm3)
Water	1.00
Ethanol	0.789
Copper	8.96

Significant Figures and Measurement Reliability

Significant Figures

Communicate the precision of a measurement.
Rules:
- All nonzero digits are significant.
- Interior zeroes are significant.
- Leading zeroes are not significant.
- Trailing zeroes after a decimal point are significant.
- Trailing zeroes before a decimal point are significant if a decimal is shown.
- Trailing zeroes before an implied decimal are ambiguous; use scientific notation.

Uncertainty in Measurements

Inexact numbers: Have uncertainty (from measurements).
Exact numbers: No uncertainty (from counting, definitions, or integral numbers in equations).

Significant Figures in Calculations

Multiplication/Division: Result has the same number of significant figures as the measurement with the fewest significant figures.
Addition/Subtraction: Result has the same number of decimal places as the measurement with the fewest decimal places.
Rounding: If the leftmost digit to be dropped is less than 5, leave the preceding number unchanged; if 5 or greater, increase the preceding number by 1.

Accuracy and Precision

Accuracy: How close a measured value is to the true value.
Precision: How close a series of measurements are to one another.

Practice Problems

Classify changes and properties as physical or chemical.
Determine significant figures in given values.
Apply dimensional analysis to solve unit conversion problems.
Calculate density, mass, or volume using the density formula.

Additional info: For more practice, refer to the end-of-chapter problems and laboratory exercises related to measurement and classification of matter.