Exam 1 Study Guide: Chemistry Fundamentals (Chapters 1, 2, 3, 4, 6, 7)

Course Introduction & Chapter 1: Chemistry in Our Lives

Scientific Method and Variables

The scientific method is a systematic approach to investigation in science. Understanding variables is essential for designing experiments.

• Independent Variable: The factor that is changed or controlled in an experiment.

• Dependent Variable: The factor that is measured or observed in response to changes in the independent variable.

• Example: In a study of plant growth, the amount of sunlight (independent) affects plant height (dependent).

Definition of Chemistry

Chemistry is the study of matter, its properties, how and why substances combine or separate, and how substances interact with energy.

• Example: Investigating how vinegar reacts with baking soda.

Accuracy vs. Precision

These terms describe the quality of measurements.

• Accuracy: How close a measurement is to the true or accepted value.

• Precision: How close repeated measurements are to each other.

• Example: If you weigh a sample three times and get 5.01g, 5.00g, and 5.02g, your measurements are precise.

Scientific Notation and Mathematical Calculations

Scientific notation expresses numbers as a product of a coefficient and a power of ten.

• Example: $3,500 = 3.5 \times 10^3$

• Rounding: Adjusting numbers to a specified number of digits.

• Percentages: $\text{Percent} = \frac{\text{Part}}{\text{Whole}} \times 100$

• Interpreting Graphs: Understanding data trends and relationships between variables.

Chapter 2: Chemistry and Measurements

Qualitative vs. Quantitative Measurements

Measurements can be descriptive or numerical.

• Qualitative: Describes qualities (e.g., color, texture).

• Quantitative: Expresses quantities (e.g., mass, volume).

Units of Measurement

Different systems are used for measuring physical quantities.

• English System: Inches, pounds, gallons, Fahrenheit, seconds.

• Metric System: Meters, grams, liters, Celsius, seconds.

• SI System: Standardized units: meter (m), kilogram (kg), second (s), kelvin (K), mole (mol).

• Common Units: Volume (L, mL), Length (m, cm), Mass (g, kg), Temperature (°C, K), Time (s).

Unit Conversions and Prefixes

Converting between units is essential in chemistry.

• Prefixes: kilo- ($10^3$), centi- ($10^{-2}$), milli- ($10^{-3}$), micro- ($10^{-6}$), deci- ($10^{-1}$).

• Example: $1 \text{km} = 1000 \text{m}$

• "cc" Unit: $1 \text{cc} = 1 \text{cm}^3 = 1 \text{mL} = 20 \text{drops}$

Measured vs. Exact Numbers

• Measured: Obtained by measurement; have uncertainty.

• Exact: Defined values (e.g., 1 dozen = 12).

Significant Figures

Significant figures reflect the precision of a measurement.

• Rules: All nonzero digits are significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros in a decimal are significant.

• Reporting: Measurements should be reported with the correct number of significant digits.

Dimensional Analysis (Factor-Label Method)

Used to convert units by multiplying by conversion factors.

• Example: $\text{Given} \times \frac{\text{Desired Unit}}{\text{Given Unit}}$

Density and Specific Gravity

• Density: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Specific Gravity: $\text{Specific Gravity} = \frac{\text{Density of Sample}}{\text{Density of Water}}$

• Application: Used in medicine to assess urine concentration; high SG indicates concentrated urine, low SG indicates dilute urine.

Chapter 3: Matter and Energy

Classification of Matter

Matter can be classified based on composition.

• Pure Substance: Has a fixed composition (element or compound).

• Mixture: Combination of two or more substances.

• Element: Simplest form of matter; cannot be broken down.

• Compound: Substance formed from two or more elements chemically combined.

Types of Mixtures

• Homogeneous: Uniform composition (e.g., salt water).

• Heterogeneous: Non-uniform composition (e.g., salad).

States of Matter and Properties

• Solid: Definite shape and volume; particles tightly packed.

• Liquid: Definite volume, no definite shape; particles less tightly packed.

• Gas: No definite shape or volume; particles far apart.

• Physical Properties: Observable without changing composition (e.g., color, melting point).

• Chemical Properties: Describe ability to undergo chemical change (e.g., flammability).

Physical and Chemical Changes

• Physical Change: Alters form but not composition (e.g., melting ice).

• Chemical Change: Produces new substances (e.g., rusting iron).

Temperature and Absolute Zero

• Absolute Zero: Lowest possible temperature; $0 \text{K}$ ($-273.15 \text{°C}$).

• Temperature Conversion: $\text{K} = \text{°C} + 273.15$

Energy Concepts

• Energy: Capacity to do work.

• Potential Energy: Stored energy.

• Kinetic Energy: Energy of motion.

• Unit Conversion: $1 \text{cal} = 4.184 \text{J}$

Specific Heat and Heat Calculations

• Specific Heat: Amount of heat required to raise the temperature of 1g of a substance by 1°C.

• Formula: $q = m \times c \times \Delta T$

• Water: High specific heat due to polar covalent bonds and hydrogen bonding.

Chapter 4: Atoms and Elements

Elements and the Periodic Table

• Element Names and Symbols: Refer to Table 4.2 (first 20 elements).

• Groups: Vertical columns; similar properties.

• Periods: Horizontal rows.

• Family Names: Alkali metals (1A), alkaline earth metals (2A), halogens (7A), noble gases (8A).

Classification of Elements

• Metals: Conduct electricity, malleable, shiny.

• Nonmetals: Poor conductors, brittle.

• Metalloids: Properties intermediate between metals and nonmetals.

• Representative Elements: Groups 1A-8A.

• Transition Metals: Groups in the center of the table.

Atomic Structure

• Protons: Positive charge, in nucleus.

• Neutrons: Neutral, in nucleus.

• Electrons: Negative charge, orbit nucleus.

• Role: Protons define element; neutrons affect mass; electrons affect chemical behavior.

Calculating Subatomic Particles

• Protons: Equal to atomic number.

• Neutrons: $\text{Neutrons} = \text{Mass Number} - \text{Atomic Number}$

• Electrons: Equal to protons in neutral atom.

Isotopes

• Definition: Atoms of the same element with different numbers of neutrons.

• Application: Used in medicine (e.g., radioactive iodine for thyroid).

Mass Number vs. Atomic Mass

• Mass Number: Total protons + neutrons.

• Atomic Mass: Weighted average of isotopes.

Energy Levels and Valence Electrons

• Energy Level: Regions where electrons are likely found.

• Valence Electrons: Electrons in the outermost energy level; determine chemical properties.

Periodic Trends

• Atomic Size: Increases down a group, decreases across a period.

• Ionization Energy: Energy to remove an electron; increases across a period.

• Electron Affinity: Tendency to gain electrons.

• Metallic Character: Increases down a group.

Electron Arrangement

• First 20 Elements: Write electron configuration based on atomic number.

• Example: Carbon (6): 2, 4

Chapter 6: Ionic and Molecular Compounds

Octet Rule and Ion Charges

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons.

• Cations: Positive ions (loss of electrons).

• Anions: Negative ions (gain of electrons).

Simple Ions and Formulas

• Representative Elements: Groups 1A-8A form predictable ions.

• Example: Sodium (Na) forms Na+; Chlorine (Cl) forms Cl-.

Writing Ionic Compound Formulas

• Combine cations and anions to balance charges.

• Example: Na+ + Cl- → NaCl

Bond Types

• Ionic: Transfer of electrons; metal + nonmetal.

• Covalent: Sharing of electrons; nonmetal + nonmetal.

• Polar Covalent: Unequal sharing of electrons.

Polyatomic Ions

Memorize the following polyatomic ions:

Name	Formula	Charge
Hydroxide	OH-	-1
Nitrate	NO3-	-1
Nitrite	NO2-	-1
Bicarbonate	HCO3-	-1
Carbonate	CO32-	-2
Cyanide	CN-	-1
Sulfate	SO42-	-2
Sulfite	SO32-	-2
Phosphate	PO43-	-3

Naming Compounds

• Ionic Compounds: Name cation first, then anion.

• Molecular Compounds: Use prefixes to indicate number of atoms.

• Example: CO2 is carbon dioxide.

Intermolecular Forces

• Ions: Strong electrostatic forces.

• Polar Covalent Molecules: Dipole-dipole interactions.

• Nonpolar Covalent Molecules: London dispersion forces.

Chapter 7: Chemical Quantities and Reactions

Avogadro's Number and Moles

• Avogadro's Number: $6.022 \times 10^{23}$ particles per mole.

• Use: Calculate number of particles in a given number of moles.

Molar Mass

• Definition: Mass of one mole of a substance (g/mol).

• Calculation: Add atomic masses from periodic table.

• Conversion: $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass (g/mol)}}$

Balanced Chemical Equations

• Reactants: Starting substances.

• Products: Substances formed.

• Balancing: Ensure same number of atoms on both sides.

Types of Chemical Reactions

• Combination: Two or more substances form one product.

• Decomposition: One substance breaks into two or more products.

• Single Replacement: One element replaces another.

• Double Replacement: Exchange of ions between two compounds.

• Combustion: Substance reacts with oxygen, producing energy.

Oxidation and Reduction

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Identify: Determine which reactant is oxidized or reduced.

Stoichiometry

• Mole-Mole Factor: Use balanced equation to relate moles of substances.

• Mass-Mass Calculations: Convert grams to moles, use mole ratio, convert back to grams.

Endothermic vs. Exothermic Reactions

• Endothermic: Absorbs heat; surroundings get colder.

• Exothermic: Releases heat; surroundings get warmer.