final review
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Title
HONORS CHEMISTRY Unit 1: Matter and Changes
NAME:
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Scientific Methodology
Definition: Method used by scientists for logical and systematic problem-solving.
Observation: Use of senses to obtain information.
Types of Observations:
Qualitative: Descriptive qualities from senses.
Quantitative: Numerical data from observations.
Hypothesis: Proposed explanation testable through prediction
Format: "If ____ then ____."
Experiment: Step-by-step procedure to test hypotheses.
Practice Questions
Qualitative observations: ______________________________
Quantitative observations: ____________________________
Variables:
Independent Variable: Changed by the experimenter.
Dependent Variable: Changes as a result of the independent variable.
Controlled Variable: Kept constant throughout the experiment.
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Variables
Experiments involve three types of variables:
Independent Variable: Intentionally manipulated.
Dependent Variable: Outcomes based on independent variable changes.
Controlled Variable: Variables kept constant to ensure valid results.
Practice Questions
Example 1: Changing ethanol volume in boiling point experiment.
Independent Variable: ____________
Dependent Variable: ____________
Controlled Variables: ____________
Example 2: Mouthwash effectiveness experiment.
Independent Variable: ____________
Dependent Variable: ____________
Controlled Variables: ____________
Practice Chemical Data
Mouthwash effectiveness measured by average bacterium per tooth
Variables in fish eggs hatching study setup with temperatures: 10℃, 20℃, 30℃, 40℃, and 50℃.
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Chemistry Fundamentals
Chemistry: Study of matter composition and changes.
Matter: Anything with mass and volume.
Chemical: Substance produced or used in chemistry.
Common chemicals: water (H2O), methane (CH4).
Important Concepts
Chemical Formula: Indicates number and types of atoms.
Example: H2O, CO2, NaCl.
Element: Simplest form of matter.
Listed in the Periodic Table; examples include Carbon (C), Helium (He).
Subscript: Number indicating atoms, e.g., H2O means 2 hydrogen atoms.
Coefficient: Indicates how many molecules in a formula, e.g., in 3H2O, 3 is the coefficient.
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Element Symbols and Chemical Formulas
Need for Naming System: Common names lack chemical composition details.
Element Symbols: One, two, or three-letter codes representing elements (Capital Letter + lower case).
Examples:
Copper: Cu, Argon: Ar, Lithium: Li.
Purpose of Element Symbols
Provides shorthand for elements in formulas.
Difference between element symbols and chemical formulas is that element symbols are single elements while formulas represent compounds.
Subscript Usage
Indicates quantity of each element in a compound.
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Composite Elements
Diatomic Elements: Composed of two atoms; found on the periodic table.
Examples include H2, O2, N2.
Tetratomic: Phosphorus as P4;
Octatomic: Sulfur as S8.
Practice: Write the chemical formula for the following elements.
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States of Matter
Matter: Anything with mass and volume.
Three Physical States:
Solids, liquids, gases.
Solids: Fixed shape and volume.
Example: Ice, iron.
Liquids: Fixed volume but take the shape of their container.
Example: Water, oil.
Gases: No fixed shape or volume; fill containers.
Example: Oxygen gas, air.
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Identifying States of Matter
Particle Level Drawings:
Solid: Particles close in pattern.
Liquid: Particles close but jumbled.
Gas: Particles far apart.
Periodic Table: Colors indicate state:
Black = solid, Blue = liquid, Red = gas.
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Atoms and Molecules
Atoms: Smallest units of elements.
Molecule: Two or more atoms bonded together.
Identifying Atoms and Molecules
Identifying: Use the periodic table to check if it has subscripts or multiple elements.
Practice distinguishing between atoms and molecules using examples.
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Diatomic and Polyatomic Elements
Diatomic Elements: H2, O2, N2, etc.
Polyatomic Elements: Phosphorus P4 and Sulfur S8.
Chemical Compound Formation:
Example: Covalent compounds form with specific ratios determined by element properties.
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Pure Substances and Mixtures
Matter Categories: Pure substances and mixtures.
Pure Substances: Consistent chemical composition, such as elements and compounds.
Mixtures: Combinations of two or more substances, can vary in composition.
Homogeneous Mixtures: Uniform compositions (saltwater).
Heterogeneous Mixtures: Distinct parts visible (salad).
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Elements and Compounds
Element: Pure substance consisting of one type of atom.
Compound: Pure substance with two or more elements.
Practice Distinction Between Element and Compounds
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Separation Techniques
Mixture Separation: Homogeneous mixtures are harder to separate.
Filtration: Separates solids from liquids.
Distillation: Utilizes different boiling points.
Chromatography: Separates based on affability for mobile phases.
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Continued Separation Techniques
Sublimation: Solid to gas separation process.
Crystallization: Produces pure solid from solution.
Decantation: Transfers a liquid while leaving solids behind.
Centrifugation: Spins mixture to separate components based on density.
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Reading the Periodic Table
Structure: Columns (groups) and rows (periods).
Information in Each Box: Atomic number, symbol, name, mass.
Identity through Atomic Number: Unique to each element.
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Atomic Mass and Isotopes
Atomic Mass: Weighted average reflecting the number of isotopes and their masses.
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Reaction Terminology
Chemical Reactions: Involve reorganization of atoms, represented in chemical formulas.
Reactants vs Products: Substance before and after a reaction.
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Yield and State Symbols
Yield Symbol: Indicates the production/result of a reaction.
State Symbols: Represent physical state (solid, liquid, gas, aqueous).
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Physical and Chemical Properties
Properties: Characteristics observed/ measured without altering substances.
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Further Physical Properties Practice
Identifying properties as chemical (C) or physical (P).
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Physical and Chemical Changes
Types of Changes: Physical changes do not alter identity; chemical changes create new substances.
Recognizing Changes: Changes can be identified through observable events like color change or gas formation.
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Equations for Changes and Identification
Physical vs. Chemical Reactants: Differentiated by characteristics seen in equations.
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UNIT 2: Number Handling
Measurements and conversions in chemical analysis.
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Fundamental Units
of Measurement
SI Units: Basis for each measurement in chemistry.
Examples: Meter (length), Kilogram (mass), Second (time).
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Derived Units
Derived Quantities: Calculated units from fundamental measurements.
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Scientific Notation
Purpose: Simplifies large/small numbers for calculations.
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Utilizing Scientific Notation on Calculators
Practical examples for performing operations with scientific notation.
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Measurement Basics
Components of Measurement: Magnitude, precision, unit.
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Measuring Liquids and Temperature
Techniques for ensuring accuracy in measurements.
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Significant Figures
Rules for Identifying: Practical methods for recognizing significant figures in numbers.
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More Significant Figures Practice
Further recognition of significant figures and rounding.
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Calculations with Significant Figures
Methodologies for determining results based on significant figures.
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Accuracy vs. Precision
Defining and distinguishing these two quality measures in experimentation.
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Density as a concept in Chemistry
Density Defined: Mass per volume ratio and related calculations.
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Determining Density Through Irregular Object Measurements
Methods of measurement: Techniques for calculating density using displacement.
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Dimensional Analysis
Conversion Factors: Significant for unit changes in measurements.
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Summary of Dimensional Analysis Procedures
Steps for effective unit conversion in calculations.
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Metric System Best Practices
Common units: Essential information for quantitative chemistry analysis.
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Basic Conversion Practice in Chemistry
Practical examples to reinforce understanding.
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Common English Conversion Factors
Relative measures for comparisons in units.
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UNIT 3: Nomenclature and the Mole
Introduction to essential naming conventions and mole calculations.
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Metals, Nonmetals, and Covalent Compounds
Properties and naming conventions in chemistry.
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Ionic Compounds and Naming Formulas
Understanding the balance of charges and compound formation.
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Acids and Their Formulations
Techniques for identifying and creating formulas transitionally.
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The Mole Concept and Molar Mass
Fundamental descriptions of moles and their relevance in calculations.
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Particle-Concentration Conversions with Moles
Using Avogadro's number in practical scenarios.
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Historical Context of Atomic Structure
Development through key historical figures and their contributions to modern atomic theory.
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Fundamental Particle Theory and Atomic Mass
Overview of quantum physics and subatomic particles in chemistry.
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Wave-Particle Duality of Light and Chemistry
In-depth analysis of light behavior in relation to atomic structures.
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Quantum Mechanical Model and Electron Configurations
Explanation of energy levels and electron distributions in atoms.