Introductory Chemistry: Study Guide for Matter, Energy, Gases, Liquids, Solids, and Solutions

Ch. 3: Matter and Energy

What Is Matter?

Matter is anything that has mass and occupies space. It is composed of atoms and molecules, which are in constant motion.

• Matter can exist in different physical forms and can be classified based on its state and composition.

• Atoms are the basic building blocks of matter.

Classifying Matter According to Its State: Solid, Liquid, and Gas

Matter exists in three primary states, each with distinct properties:

• Solids: Definite shape and volume; particles are closely packed and vibrate in place.

• Liquids: Definite volume but no definite shape; particles are close but can move past each other.

• Gases: No definite shape or volume; particles are far apart and move freely.

Classifying Matter According to Its Composition: Elements, Compounds, and Mixtures

Matter can also be classified by its chemical composition:

• Elements: Pure substances consisting of only one type of atom (e.g., O2, Fe).

• Compounds: Pure substances composed of two or more elements chemically combined in fixed ratios (e.g., H2O, CO2).

• Mixtures: Physical combinations of two or more substances; can be homogeneous (uniform, e.g., saltwater) or heterogeneous (non-uniform, e.g., salad).

Differences in Matter: Physical and Chemical Properties

Properties of matter can be classified as physical or chemical:

• Physical properties: Observed without changing the substance's identity (e.g., color, melting point, density).

• Chemical properties: Describe a substance's ability to undergo chemical changes (e.g., flammability, reactivity).

Changes in Matter: Physical and Chemical Changes

Changes in matter are categorized as:

• Physical changes: Do not alter the chemical composition (e.g., melting, boiling, dissolving).

• Chemical changes: Result in new substances with different properties (e.g., rusting, combustion).

Separating Mixtures Through Physical Changes: Mixtures can be separated by physical means such as filtration, distillation, or chromatography.

Conservation of Mass: There is No New Matter

The law of conservation of mass states that matter is neither created nor destroyed in a chemical reaction.

• Mass of reactants = Mass of products

Energy

Energy is the capacity to do work or produce heat. It exists in various forms, including kinetic and potential energy.

• Units of Energy: The SI unit is the joule (J); another common unit is the calorie (cal).

• 1 cal = 4.184 J

Energy and Chemical, and Physical Change

Energy changes accompany both physical and chemical changes. Exothermic processes release energy, while endothermic processes absorb energy.

Temperature: Random Motion of Molecules and Atoms

Temperature measures the average kinetic energy of particles in a substance. Higher temperature means faster particle motion.

Energy and Heat Capacity Calculations

Heat capacity is the amount of heat required to change a substance's temperature by 1°C.

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

• Where $q$ = heat (J), $m$ = mass (g), $c$ = specific heat capacity (J/g°C), $\Delta T$ = change in temperature (°C)

Ch. 8: Quantities in Chemical Reactions

Making Pancakes: Relationships between Ingredients

Recipes illustrate the concept of stoichiometry, where specific ratios of ingredients yield a certain amount of product.

Making Molecules: Mole-to-Mole Conversions

Stoichiometry uses balanced chemical equations to relate moles of reactants and products.

• Use coefficients from the balanced equation as conversion factors.

• Example: $2H_2 + O_2 \rightarrow 2H_2O$

Making Molecules: Mass-to-Mass Conversions

To convert between masses of reactants and products:

1. Convert mass to moles (using molar mass).

2. Use mole ratio from the balanced equation.

3. Convert moles back to mass.

More Pancakes: Limiting Reactants, Theoretical Yield, and Percent Yield

The limiting reactant is the reactant that is completely consumed first, limiting the amount of product formed.

• Theoretical yield: Maximum amount of product possible from given reactants.

• Percent yield: $\text{Percent yield} = \frac{\text{Actual yield}}{\text{Theoretical yield}} \times 100\%$

Limiting Reactants, Theoretical Yield, and Percent Yield from Initial Masses of Reactants

Calculate the amount of product formed by determining the limiting reactant and using stoichiometry.

Enthalpy: A Measure of the Heat Evolved or Absorbed in a Reaction

• Enthalpy change ($\Delta H_{rxn}$): The heat change at constant pressure during a reaction.

• Sign of $\Delta H_{rxn}$: Negative for exothermic, positive for endothermic reactions.

• Stoichiometry of $\Delta H_{rxn}$: Use the balanced equation to relate enthalpy change to moles of reactants/products.

Ch. 11: Gases

Kinetic Molecular Theory: A Model for Gases

The kinetic molecular theory explains the behavior of gases based on the motion of their particles.

• Gas particles are in constant, random motion.

• Collisions are elastic; no energy is lost.

• Volume of particles is negligible compared to the container.

Pressure: The Result of Constant Molecular Collisions

Gas pressure results from collisions of gas particles with the walls of their container.

• Pressure units: Atmosphere (atm), Pascal (Pa), torr, mmHg.

• 1 atm = 101,325 Pa = 760 mmHg = 760 torr

• Pressure unit conversion: Use the above relationships to convert between units.

Boyle’s Law: Pressure and Volume

At constant temperature, the pressure and volume of a gas are inversely related.

• Formula: $P_1V_1 = P_2V_2$

Charles’s Law: Volume and Temperature

At constant pressure, the volume of a gas is directly proportional to its absolute temperature (in Kelvin).

• Formula: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$

The Combined Gas Law: Pressure, Volume, and Temperature

Combines Boyle’s and Charles’s laws for a fixed amount of gas.

• Formula: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

Avogadro’s Law: Volume and Moles

At constant temperature and pressure, the volume of a gas is directly proportional to the number of moles.

• Formula: $\frac{V_1}{n_1} = \frac{V_2}{n_2}$

The Ideal Gas Law: Pressure, Volume, Temperature, and Moles

Relates all four variables for an ideal gas.

• Formula: $PV = nRT$

• Where $P$ = pressure, $V$ = volume, $n$ = moles, $R$ = gas constant, $T$ = temperature (K)

• Determining the Molar Mass of a Gas: $\text{Molar mass} = \frac{mRT}{PV}$, where $m$ = mass of gas

• Ideal and Nonideal Gas Behavior: Real gases deviate from ideal behavior at high pressures and low temperatures.

Mixtures of Gases

• Partial Pressure: The pressure exerted by each gas in a mixture.

• Dalton’s Law: $P_{total} = P_1 + P_2 + ... + P_n$

• Collecting Gases over Water: Account for water vapor pressure when measuring gas volume.

Gases in Chemical Reactions

• Molar Volume at STP: At standard temperature and pressure (0°C, 1 atm), 1 mole of gas occupies 22.4 L.

Ch. 12: Liquids, Solids, and Intermolecular Forces

Properties of Liquids and Solids

Liquids and solids have stronger intermolecular forces than gases, resulting in definite volume (and shape for solids).

Intermolecular Forces in Action: Surface Tension and Viscosity

• Surface Tension: The energy required to increase the surface area of a liquid; caused by cohesive forces between molecules.

• Viscosity: A liquid's resistance to flow; higher viscosity means slower flow.

Evaporation and Condensation

• Evaporation: Liquid to gas at the surface below boiling point.

• Condensation: Gas to liquid.

• Boiling: Rapid vaporization throughout the liquid at the boiling point.

• Energetics: Evaporation is endothermic; condensation is exothermic.

• Heat of Vaporization ($\Delta H_{vap}$): Energy required to vaporize 1 mole of liquid.

Melting, Freezing, and Sublimation

• Melting: Solid to liquid (endothermic).

• Freezing: Liquid to solid (exothermic).

• Sublimation: Solid to gas without passing through the liquid phase.

• Heat of Fusion ($\Delta H_{fus}$): Energy required to melt 1 mole of solid.

Types of Intermolecular Forces: Dispersion, Dipole-Dipole, Hydrogen Bonding, and Ion-Dipole

• Dispersion Forces: Present in all molecules; caused by temporary dipoles.

• Dipole-Dipole Forces: Occur between polar molecules.

• Hydrogen Bonding: Strong dipole-dipole interaction involving H bonded to N, O, or F.

• Ion-Dipole Forces: Between ions and polar molecules (important in solutions).

Ch. 13: Solutions

Solutions: Homogeneous Mixtures

A solution is a homogeneous mixture of two or more substances. The solute is dissolved in the solvent.

Solutions of Solids Dissolved in Water: How to Make Rock Candy

• Solubility: Maximum amount of solute that can dissolve in a solvent at a given temperature.

• Saturation: A saturated solution contains the maximum amount of dissolved solute.

• Electrolyte Solutions: Solutions containing dissolved ionic solids that conduct electricity.

• How Solubility Varies with Temperature: Generally, solubility of solids increases with temperature.

Solutions of Gases in Water: How Soda Pop Gets Its Fizz

• Solubility of gases in water decreases as temperature increases and increases as pressure increases.

Specifying Solution Concentration: Mass Percent

• Mass percent: $\text{Mass percent} = \frac{\text{mass of solute}}{\text{mass of solution}} \times 100\%$

• Used to describe concentration and for calculations.

Specifying Solution Concentration: Molarity

• Molarity (M): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Used for stoichiometric calculations and to determine ion concentrations.

Solution Dilution

• To dilute a solution: $M_1V_1 = M_2V_2$

• Where $M$ = molarity, $V$ = volume

Solution Stoichiometry

Use molarity and volume to calculate moles of solute for reactions in solution.

Freezing Point Depression and Boiling Point Elevation: Making Water Freeze Colder and Boil Hotter

• Freezing point depression: Adding solute lowers the freezing point of a solvent.

• Boiling point elevation: Adding solute raises the boiling point of a solvent.

• Formulas:

  • $\Delta T_f = iK_f m$

  • $\Delta T_b = iK_b m$

  • Where $i$ = van 't Hoff factor, $K_f$/$K_b$ = constants, $m$ = molality

Osmosis: Why Drinking Salt Water Causes Dehydration

Osmosis is the movement of solvent through a semipermeable membrane from low to high solute concentration. Drinking salt water causes water to leave cells, leading to dehydration.