Matter and Energy

Introduction to Chemistry

Chemistry is the study of matter and the changes it undergoes. Matter is defined as anything that occupies space and can be perceived by our senses.

• Matter: Anything that has mass and takes up space.

• Chemical change: A process that alters the chemical identity of a substance.

• Physical change: A process that does not alter the chemical identity of a substance.

States of Matter

Matter exists in different physical forms, each with distinct properties.

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

• Liquid: Definite volume but indefinite shape; particles are close but can move past each other.

• Gas: Indefinite volume and shape; particles are far apart and move freely.

• Compressibility: Gases are highly compressible, liquids are slightly compressible, and solids are nearly incompressible.

Properties of Matter

Physical and Chemical Properties

Properties are characteristics used to describe matter. They can be classified as physical or chemical.

• Physical properties: Characteristics that can be observed without changing the chemical identity of the substance (e.g., color, melting point, density).

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

Examples of Properties

• Chalk (calcium carbonate) fizzes in dilute acid – Chemical property

• Ice melts when temperature reaches 32°F – Physical property

• Copper (II) oxide turns black when heated in air – Chemical property

• Bromine is a toxic, reddish-brown gas – Physical property

Intensive and Extensive Properties

Properties can also be classified based on their dependence on the amount of material present.

• Intensive properties: Do not depend on the amount of substance (e.g., color, melting point, density).

• Extensive properties: Depend on the amount of substance (e.g., mass, volume).

Changes in Matter

Physical and Chemical Changes

Changes in matter can be physical or chemical.

• Physical change: Alters the physical form but not the chemical identity (e.g., melting, boiling, dissolving).

• Chemical change: Produces one or more new substances with different chemical identities (e.g., burning, rusting).

Examples of Changes

• Baking a cookie – Chemical change

• Chewing a cookie – Physical change

• Digesting a cookie – Chemical change

Classification of Matter

Types of Matter

All matter is either a substance or a mixture of substances.

• Substance: A kind of matter that cannot be separated by physical processes.

• Mixture: Material that can be separated into two or more substances by physical means.

Elements, Compounds, and Mixtures

• Element: Comprised of one kind of atom; cannot be broken down by chemical means.

• Compound: Made of two or more elements chemically combined in fixed ratios; can be broken down by chemical means.

• Molecule: A type of substance where two or more atoms are chemically bonded in discrete units.

Types of Mixtures

• Heterogeneous mixture: Has physically distinct parts (e.g., iron filings/sawdust, sodium chloride and copper(II) sulfate pentahydrate).

• Homogeneous mixture: Uniform in its properties throughout (e.g., solutions such as salt water).

Law of Definite Proportions

The law states that a pure compound always contains the same proportions of elements by mass.

• For example, 100g of sodium chloride will always contain 39.3g of sodium and 60.7g of chlorine.

• 18g of water will always contain 16g of oxygen and 2g of hydrogen.

Matter Classification Table

Energy

Forms of Energy

Energy is the ability to do work. It exists in several forms relevant to chemistry.

• Kinetic energy: Energy due to motion; sometimes referred to as thermal energy.

• Potential energy: Energy stored in chemical bonds.

• Internal energy: The sum of kinetic and potential energies of molecules in a sample.

Law of Conservation of Energy

Energy can change forms but cannot be created or destroyed. Chemical reactions can release energy, and matter can be used to create energy (e.g., nuclear reactions).

Heat and Temperature

• Temperature: Measure of the average kinetic energy of a substance; measured in degrees Celsius (°C), Kelvin (K), or Fahrenheit (°F).

• Heat (q): Energy that flows due to a temperature difference between system and surroundings; measured in joules (J) or calories (cal).

Temperature Conversion Formulas

Energy and Its Units

• calorie (cal): Amount of heat needed to raise the temperature of 1 gram of water by 1°C.

• kilocalorie (kcal): 1000 cal.

• Food Calorie (Cal): 1000 cal.

• Joule (J): Metric unit of energy;

Energy Conversion Example

• A candy bar with 350 Calories contains joules.

Calorimetry

Calorimeters are devices used to measure heat flow in chemical reactions.

• Bomb calorimeter: Used for measuring energy released in combustion reactions.

• Coffee cup calorimeter: Used for reactions at constant pressure.

Specific Heat

Definition and Formula

Specific heat (SH) is the amount of heat needed to raise the temperature of 1 gram of material by one degree Celsius.

Example Calculation

• A 10.0 g sample of copper at 25°C is heated to 100°C with 289 J of heat added. The specific heat of copper can be calculated using the above formula.

Distribution of Energy

Kinetic Energy Distribution

In a sample of material, the kinetic energies of molecules follow a Boltzmann distribution.

• Average kinetic energy:

• Where m is mass and v is velocity.

States of Matter: Detailed Properties

Solids

• Condensed state; atoms or molecules are "touching".

• Strongest intermolecular forces; particles held rigidly in a 3D crystalline lattice.

Liquids

• Condensed state; atoms or molecules are "touching" but not rigidly.

• Intermolecular forces allow molecules to slide past each other.

Gases

• Virtually no intermolecular forces.

• Molecules are in constant random motion; velocity is related to temperature.

• Molecules collide with container walls and each other, bouncing off with no loss of energy.

Changes of State

Vaporization

• Most energetic molecules in a liquid escape into the gas phase.

• Once molecules are free as gases, they exert a pressure called vapor pressure.

• Vapor pressure depends on temperature; higher temperature increases vapor pressure.

Boiling Point

• Boiling point: Temperature where vapor pressure equals ambient pressure.

• Normal boiling point: Vapor pressure equals 760 torr.

• Boiling point depends on pressure; lower pressure lowers boiling point.

Freezing/Melting Point

• Melting point: Temperature at which a crystalline solid changes to a liquid.

• Freezing point: Temperature at which a liquid changes to a solid.

Energy Changes and Heating Curves

Heating Curve for Water

A heating curve shows the temperature change of a substance as heat is added.

• Plateaus on the curve represent phase changes (melting, boiling) where temperature remains constant as energy is used to change state.

Sample Heating Curve Table

Heat of Fusion

• Heat of fusion: Heat required to convert one mole of solid to liquid at its normal melting point.

• Represents energy needed to overcome intermolecular forces and allow molecules to slide around in the liquid phase.

Molar Heat of Vaporization

• Heat of vaporization: Heat required to convert one mole of liquid to gas at its normal boiling point.

• Represents energy needed to overcome intermolecular forces and allow molecules to escape into the gas phase.

Summary Table: Phase Changes and Energy

Practice Problem Example

• If a recipe says to bake a cake at 487 K, what temperature in °F should you set your oven? Use

Conclusion

This chapter covers the foundational concepts of matter and energy, including classification, properties, changes, and the relationship between energy and phase changes. Understanding these principles is essential for further study in chemistry.