Introductory Chemistry — Chapter 3: Matter and Energy (Vocabulary Flashcards)
Matter and Energy
Matter is defined as anything that occupies space and has mass.
It can exist as solids, liquids, or gases, and some types (like steel, water, wood, plastic) are easily visible, while others (air, microscopic dust) require magnification to detect.
All matter is composed of atoms, the fundamental building blocks. In many substances atoms bond to form molecules (two or more atoms bonded in specific geometric arrangements).
Advances in microscopy allow imaging of atoms and molecules (e.g., Scanning Tunneling Microscope, STM).
Atoms and Molecules
All matter is ultimately composed of atoms.
In some substances (e.g., an aluminum can), atoms exist as independent particles.
In other substances (e.g., rubbing alcohol), atoms bond to form well-defined structures called molecules.
STM images cobalt atoms on a copper surface, showing individual atoms as raised bumps; STM works by scanning a surface with a tip of atomic dimensions to produce images.
DNA is the hereditary material with a double-stranded structure visible in STM imagery (DNA is yellow in the figure). The double helix is the hallmark of its structure.
States of Matter
The common states are solid, liquid, and gas.
Matter is classified by state: solid, liquid, or gas; these states differ in spacing and motion of particles.
Water exemplifies all three states: ice (solid), liquid water, and steam (gas).
The Three States of Matter in Water
Solid (Ice): molecules closely spaced; vibrate about fixed points.
Liquid: molecules closely spaced but free to move around and past each other.
Gas (Steam): molecules separated by large distances; free to move and interact negligibly.
Solids
In solids, atoms or molecules are packed tightly in fixed locations.
Neighboring particles may vibrate but do not move past one another.
Solids have fixed volume and a rigid shape.
Examples: ice, diamond, quartz, iron.
Types of Solids
Crystalline solids: atoms/molecules occupy specific positions forming a well-ordered, three-dimensional, long-range repeating structure.
Amorphous solids: lack long-range order; no well-defined repeating structure.
Examples:
Crystalline: salt (sodium chloride), diamond.
Amorphous: glass, rubber, plastic.
Crystalline Solids: Example
Sodium chloride (table salt) is crystalline; its well-ordered, cubic shape results from regular, cubic arrangement of its atoms.
Liquids
In liquids, particles are close but can move relative to one another.
Liquids have a fixed volume but take the shape of their container because particles can move past each other.
Examples: water, gasoline, alcohol, mercury.
Gases
In gases, particles are far apart and move freely.
Gases are compressible because particles are not in contact.
Gases take the shape and volume of their containers.
Examples: oxygen, helium, carbon dioxide.
Properties of Solids, Liquids, and Gases
Table (State, Motion, Spacing, Shape, Volume, Compressibility):
Solid: oscillation/vibration about fixed point; close together; definite shape; definite volume; incompressible.
Liquid: free to move relative to one another; close together; indefinite shape; definite volume; incompressible.
Gas: free to move relative to one another; far apart; indefinite shape; indefinite volume; compressible.
Pure Substances and Mixtures
Matter can be categorized into two broad types:
Pure substances: composed of only one kind of atoms or molecules.
Mixtures: composed of two or more kinds of atoms or molecules in variable proportions.
Elements
An element is a substance that cannot be broken down into simpler substances.
All known elements are listed in the periodic table; the smallest unit of an element is an atom.
Helium is an example of an element (a pure substance) composed of only helium atoms.
Compounds
A compound is a pure substance composed of two or more elements in fixed, definite proportions.
Compounds are more common than pure elements and can be decomposed into simpler substances.
Water (H₂O) is a compound made from two elements (hydrogen and oxygen) in fixed proportions.
Mixtures (Pure Substances vs Mixtures)
A compound is a pure substance with chemically bound elements in fixed proportions.
A mixture is made of different substances physically mixed in variable proportions (not chemically bonded).
Mixtures: Homogeneous vs Heterogeneous
Heterogeneous mixtures have variable composition and can be separated into distinct components (e.g., oil and water form two layers).
Homogeneous mixtures have uniform composition throughout (e.g., sweetened tea where sugar is dissolved uniformly).
Examples: Air (mostly nitrogen and oxygen) and seawater (salt and water).
Classification of Matter (Figure 3.8 Concept)
Pure substances include Elements and Compounds.
Mixtures include Homogeneous and Heterogeneous.
Summary of Matter Classification
Matter can be a pure substance or a mixture.
A pure substance is either an element or a compound.
A mixture can be homogeneous or heterogeneous and can contain two or more elements, two or more compounds, or a combination of both.
Physical vs Chemical Properties
Physical properties: observed without changing the substance's composition (e.g., odor, boiling point, color, density).
Chemical properties: observed only through chemical change in composition (e.g., rust susceptibility of iron, flammability of gasoline).
Physical vs Chemical Changes
Physical changes alter appearance but not composition (e.g., cutting or crushing; state changes like melting).
Chemical changes involve a chemical reaction that changes composition (e.g., burning butane with oxygen to form CO₂ and H₂O).
State changes (melting, boiling) are always physical changes.
Changes involving chemical reactions are always chemical changes and may involve color or temperature changes; result in a new substance.
Examples of Physical and Chemical Changes
Physical change example: ice melting to water; appearance changes but composition remains H₂O.
Chemical change example: copper turning green when exposed to air due to formation of new compounds; reaction with gases in air.
Distillation and Filtration (Separation of Mixtures) - Physical Changes
Distillation: Separates a mixture of liquids with different boiling points by heating; the most volatile component boils first, vapor is condensed and collected as a pure liquid.
Setup components: distilling flask, condenser, collection of pure liquid, cooling water in/out.
Filtration: Separates a mixture of a liquid and a solid by pouring through filter paper; solid is trapped by the paper, liquid passes through.
Setup components: funnel, filter paper, stirring rod.
Law of Conservation of Mass
Matter is neither created nor destroyed in ordinary chemical reactions.
During physical and chemical changes, the total amount of matter remains constant; mass is conserved.
In a chemical reaction, the sum of the masses of the reactants equals the sum of the masses of the products.
Example: burning 58 g of butane with 208 g of O₂ forms 176 g of CO₂ and 90 g of H₂O; total mass before = 266 g, total mass after = 266 g.
Energy and Its Role in Chemistry
Energy is the capacity to do work; work is force applied over a distance.
The behavior of matter is driven by energy; understanding energy is essential to chemistry.
Like matter, energy is conserved; it can be transformed from one form to another and transferred between objects.
Energy cannot be created from nothing or vanish into nothing.
Potential and Kinetic Energy
Total energy of a system = kinetic energy (energy of motion) + potential energy (energy of position or composition).
Forms of Energy
Electrical energy: energy of moving electric charges.
Thermal energy: energy due to the random motions of atoms and molecules; hotter objects have more thermal energy.
Chemical energy: potential energy associated with positions of particles in a chemical system.
Units of Energy
SI unit: the joule (J).
Calorie (cal): energy to raise 1 g of water by 1°C; 1 cal = 4.184 J.
Nutritional Calorie (Cal): 1 Cal = 1000 cal.
Kilowatt-hour (kWh): common unit for electricity; 1 kWh = 3.60 × 10^6 J.
Typical residential electricity cost around $0.12 per kWh.
Energy Conversion Factors (Table 3.2)
1 cal = 4.184 J
1 Cal = 1000 cal
1 kWh = 3.60 × 10^6 J
Unit Conversion Example
Example: A candy bar contains 225 Cal of nutritional energy. To convert to joules:
225 Cal × (1000 cal/Cal) × (4.184 J/cal) = 9.41 imes 10^{5} ext{ J}
Energy: Calculations Involving Heat and Temperature Change
Relationship between heat added, mass, specific heat, and temperature change:
q = m \, C \, \Delta T
where: q = heat (J), m = mass (g), C = specific heat capacity (J g^{-1} °C^{-1}), \Delta T = Tf - Ti.
The temperature change is defined as the difference between final and initial temperatures (in °C when using SI units for C).
Example problem (Gallium):
Gallium melts at 29.9 °C and has specific heat capacity C = 0.372 \, \text{J g}^{-1} {°C}^{-1}.
If 2.5 g of gallium is heated from 25.0 °C to 29.9 °C, the heat required is:
\Delta T = 29.9 - 25.0 = 4.9° C
q = m C \Delta T = (2.5 \, \text{g}) (0.372 \, \text{J g}^{-1} {°C}^{-1}) (4.9° C) \approx 4.56 \, \text{J}
Note: This is the heat required to raise the temperature to the melting point; melting itself would involve the heat of fusion, which is not included here.
Temperature Scales and Conversions
Temperature reflects the average kinetic energy of particles; heat is the transfer of thermal energy due to a temperature difference.
Fahrenheit (°F): Water freezes at 32 °F; boils at 212 °F; room temperature ~ 72 °F.
Celsius (°C): Water freezes at 0 °C; boils at 100 °C; room temperature ~ 22 °C.
Kelvin (K): Absolute scale; 0 K = absolute zero; water freezes at 273 K; boils at 373 K; room temperature ~ 295 K.
Relationship between scales:
^{ ext{°C}} = \dfrac{5}{9} (^{\text{°F}} - 32)
^{\text{°F}} = \dfrac{9}{5} (^{\text{°C}}) + 32
K = ^{\circ C} + 273.15
Degree symbol usage: The degree symbol is used with Fahrenheit and Celsius scales but not with Kelvin.
Temperature Changes: Specific Heat Capacity in Practice
Specific heat capacity is the amount of heat required to change the temperature of 1 g of a substance by 1 °C.
Substances vary widely in C; water has the highest value among common substances listed, meaning it heats up slowly and stores a lot of energy for a small temperature change.
Data table (typical values):
Lead: C = 0.128 \, \text{J g}^{-1} {°C}^{-1}
Gold: C = 0.128 \, \text{J g}^{-1} {°C}^{-1}
Silver: C = 0.235 \, \text{J g}^{-1} {°C}^{-1}
Copper: C = 0.385 \, \text{J g}^{-1} {°C}^{-1}
Iron: C = 0.449 \, \text{J g}^{-1} {°C}^{-1}
Aluminum: C = 0.903 \, \text{J g}^{-1} {°C}^{-1}
Ethanol: C = 2.42 \, \text{J g}^{-1} {°C}^{-1}
Water: C = 4.184 \, \text{J g}^{-1} {°C}^{-1}
The high specific heat of water explains why coastal regions experience moderated temperatures due to large energy absorption without large temperature changes.
Energy and Specific Heat: Worked Problem Outline
Given: mass m, specific heat C, temperature change ΔT; find q using q = m C \Delta T.
Steps:
Identify initial and final temperatures to compute \Delta T.
Multiply by the mass and the substance's specific heat capacity.
Keep track of units; ensure correct significant figures.
Review: Core Concepts (condensed)
Matter definition and states: solid, liquid, gas; crystalline vs amorphous.
Classification: Elements, Compounds, Mixtures (homogeneous vs heterogeneous).
Physical vs Chemical properties and changes.
Law of Conservation of Mass; energy conservation; energy forms.
Temperature scales: Fahrenheit, Celsius, Kelvin; conversions between them.
Specific heat capacity and energy transfer calculations.
Separation of mixtures by physical means: distillation and filtration.
Exothermic vs Endothermic reactions; TNT example illustrating exothermic release of energy.
What You Should Learn in Chapter 3 (Learning Objectives)
LO: Define matter, atoms, and molecules.
LO: Classify matter as solid, liquid, or gas.
LO: Classify matter as element, compound, or mixture.
LO: Distinguish between physical and chemical properties.
LO: Distinguish between physical and chemical changes.
LO: Apply the law of conservation of mass.
LO: Recognize the different forms of energy.
LO: Identify and convert between energy units.
LO: Distinguish between exothermic and endothermic reactions.
LO: Convert between Fahrenheit, Celsius, and Kelvin temperature scales.
LO: Relate energy, temperature change, and heat capacity.
LO: Perform calculations involving transfer of heat and changes in temperature.