Chapter 1-8 Chemistry Review: Substances, Mixtures, States, and Energy
Classification: Pure Substances and Mixtures
Major categories of matter: Pure substance and Mixture.
Pure substance definitions:
A pure substance can be either an element or a compound.
If the substance is made of only one element throughout, it is a pure substance (an element).
Elements are the basic building blocks of the universe; the universe contains up to 118 elements.
A pure substance can also be a compound, which contains multiple elements combined in a definite form.
Examples:
Element example: A substance made of only one element (e.g., pure aluminum) is a pure substance.
Compound example: Carbon dioxide, CO$_2$.
CO$2$ composition: one carbon atom bound to two oxygen atoms forming CO$2$ molecules.
Throughout the substance, only CO$_2$ molecules are present (one type of molecule).
If a substance contains two elements but sometimes forms different molecules (e.g., CO and CO$_2$), and both molecule types coexist, it is a mixture, not a pure substance.
Mixtures:
A mixture consists of two or more substances physically mixed together, not chemically combined.
Key distinction: in a mixture, the components can be separated by physical means; there is no new substance formed.
Example of a mixture formed by simple mixing and later chemical reaction: carbon and oxygen gas mixed, then burned to form CO$_2$ (chemical change).
Types of mixtures:
Homogeneous mixture: uniform composition throughout; the different parts are not visually distinguishable.
Example: sugar dissolved in water -> sugar water solution; samples from anywhere are indistinguishable.
Heterogeneous mixture: non-uniform composition; different parts have different properties or amounts.
Example: water with sand; samples from different regions show different sand-to-water ratios.
Classification practice (from transcript):
Pasta and tomato sauce: mixture; heterogeneous (liquid + solid components).
Aluminum foil: pure substance (elemental aluminum, Al).
Helium: pure substance (elemental helium).
Air: mixture (composition includes O$2$, N$2$, CO$2$, SO$x$, H$_2$O vapor, etc.).
Additional notes on composition:
A pure substance can be composed of one element or multiple elements that form the same type of molecule throughout (e.g., CO$2$). If multiple different molecules (e.g., CO and CO$2$) are present in the same sample, it is a mixture.
States of Matter and Their Properties
Three major states: solid, liquid, gas.
What determines the state: strength of interactions between particles.
Solid: strong interactions; particles have little freedom; definite shape and definite volume; particles are close together in fixed arrangement.
Liquid: weaker interactions than in solids; definite volume but shape depends on container; particles can move past each other.
Gas: very weak interactions; no definite shape or volume; expands to fill container; particles move fast and are far apart.
Summary table (conceptually):
Solid: definite shape, definite volume, low particle freedom, strong intermolecular forces.
Liquid: shape follows container, definite volume, moderate particle freedom, intermediate forces.
Gas: shape and volume follow container, high particle freedom, very weak forces.
Physical properties (definition):
Properties observed or measured without changing the identity of a substance.
Examples: color, size, mass, density, boiling point, freezing point.
Measuring boiling/freezing points keeps the substance unchanged (physical properties).
Chemical properties (definition):
Properties observed when a substance reacts to form new substances; identity changes.
Example: carbon burning in oxygen to form CO$_2$ (chemical property).
Physical vs chemical changes (definition):
Physical change: no new substance formed; identity remains the same.
Chemical change: new substance formed; original substance transforms.
Examples:
Physical changes: breaking a pen, melting gold (still gold), cutting a pizza (shape/size changes but composition same).
Chemical changes: burning candle (wax + oxygen form CO$2$ and other products), iron rusting (Fe forms Fe$2$O$_3$), toasting a marshmallow (carbon-containing compounds oxidize).
Learning checks (examples from transcript):
Determine if a property is physical or chemical.
Determine if a change is physical or chemical based on whether a new substance forms.
Practical reminder: Many everyday processes involve both physical and chemical changes; use the criterion of whether a new substance is formed to classify.
Practice: Classification of Substances and Mixtures
Identify each of the following as a pure substance or mixture:
Pasta and tomato sauce: mixture.
If a mixture, classify as homogeneous or heterogeneous: heterogeneous (solid pieces dispersed in liquid).
Aluminum foil: pure substance (elemental aluminum, Al).
Helium: pure substance (elemental helium).
Air: mixture (multiple gases: O$2$, N$2$, CO$_2$, etc.).
Classification notes:
Homogeneous mixtures have uniform composition throughout (e.g., salt water, clear sugar solution).
Heterogeneous mixtures have non-uniform composition (e.g., sand-water mixture, salad).
Temperature Scales and Conversions
Major temperature scales used:
Fahrenheit (°F): commonly used in the United States for everyday weather/temperature.
Celsius (°C): common in most countries; defined by water's phase changes.
Kelvin (K): scientific scale; absolute scale with absolute zero at 0 K.
Definitions and reference points:
Celsius definition: 0 °C is the freezing point of water; 100 °C is the boiling point of water.
Kelvin definition: 0 K is absolute zero (the lowest possible temperature where particles stop moving).
Absolute zero is a theoretical limit; in the real universe the lowest observed temperature is about 4 K (background radiation) due to the Big Bang.
Key conversions (as given in transcript):
Fahrenheit to Celsius: TF = 1.8 imes TC + 32.
Celsius to Fahrenheit: TC = rac{TF - 32}{1.8}.
Kelvin to Celsius: TC = TK - 273.
Celsius to Kelvin: TK = TC + 273. (Transcript uses 0 K = -273 °C and 4 K ≈ -269 °C.)
Examples from the transcript:
0 °C ≡ 273 K (as stated, room practice uses 0 °C = 273 K in the notes).
4 K ≡ -269 °C (since -273 °C + 4 °C = -269 °C).
25 °C ≡ 298 K (273 + 25).
Quick learning check (sample answers):
Temperature at which water freezes: 0 °C (definition of Celsius).
Temperature at which water boils: 100 °C (definition of Celsius).
Number of Celsius units between freezing and boiling points: 100 °C.
Convert -15 °F to Celsius using the relation; practice as described in lecture.
Practical note: Memorize the Fahrenheit-Celsius relation for exams (Tf = 1.8 Tc + 32). Memorization of conversions between Celsius, Kelvin, and Fahrenheit will be required.
Energy: Definitions, Types, and Examples
What is energy? The ability to do work.
Two main types of energy:
Kinetic energy (motion): energy due to movement.
Formula (conceptual): E_k = frac{1}{2} m v^2.
Key dependencies: mass (m) and velocity (v); larger mass or higher velocity increases kinetic energy.
Everyday intuition: a small fast bullet can have high kinetic energy; a very large slow object may also have notable kinetic energy.
Potential energy (stored energy): energy due to position or configuration; can be released to do work later.
Formula (gravitational): E_p = m g h.
$m$: mass, $g$: gravitational acceleration, $h$: height above reference level.
Examples:
A rock on a mountain has gravitational potential energy that increases with height $h$; if it falls, gravity does work and converts potential energy to kinetic energy.
The position of the rock relative to a reference ground determines its potential energy; higher $h$ means larger $E_p$.
Other forms of potential energy include elastic potential energy in springs (extension or compression stores energy).
Chemical energy as potential energy:
Bonds in molecules store potential energy; breaking bonds releases energy.
Example: in a carbon dioxide molecule, the C=O bonds store potential energy that can be released when bonds are broken during chemical reactions.
In biological systems, food molecules (e.g., carbohydrates) store chemical energy; digestion oxidizes these molecules and releases energy to form ATP, which drives cellular processes.
Examples of energy types in everyday items (as discussed):
Peanut butter and jelly sandwich: contains chemical energy stored in molecular bonds; energy can be released when metabolized (fuel for the body).
Gasoline in a tank: contains chemical energy stored in hydrocarbon bonds; burning releases energy to do work.
A spring: stores elastic potential energy when stretched or compressed; releasing it allows it to do work.
Summary of energy classification in the transcript:
Energy in motion: kinetic energy.
Energy stored due to position or configuration: potential energy (gravitational, elastic, chemical bonds).
In chemistry, chemical bonds store potential energy; energy can be released when bonds break and new substances form.
Energy Units and Reading Conversions
SI unit of energy: Joule (J).
Common larger unit: kilojoule (kJ).
1 ext{ kJ} = 1000 ext{ J}.
Calorie-based energy units (nutrition context):
Calorie (cal) and kilocalorie (kcal).
1 ext{ kcal} = 1000 ext{ cal} = 4184 ext{ J} = 4.184 ext{ kJ}.
In nutrition, sometimes Calorie with a capital C denotes a kilocalorie; this equals 1 kcal.
Quick conversion fact from transcript: 1 ext{ Cal} = 1 ext{ kcal} = 4184 ext{ J} = 4.184 ext{ kJ}.
Important note for exam prep: memorize the conversion between joules, kilojoules, calories, and kilocalories; specifically, 1 ext{ cal} = 4.184 ext{ J}.
Quick Reference: Key Formulas and Concepts
Pure substance concepts:
Pure substance: either an element or a compound.
Elements are pure substances made of a single type of atom.
Compounds are pure substances formed by chemical combinations of two or more elements in fixed ratios.
Mixtures are physical blends of two or more substances without chemical bonding.
States of matter:
Solid: definite shape and volume; strong interparticle forces; little particle movement.
Liquid: definite volume, variable shape; particles can flow past each other.
Gas: no definite shape or volume; particles far apart and move rapidly.
Physical vs chemical properties and changes:
Physical properties: observed without changing identity (color, density, boiling point, etc.).
Chemical properties: involve changing into new substances (e.g., burn in oxygen to form CO$_2$).
Physical changes: no new substance formed (e.g., melting, cutting).
Chemical changes: new substances formed (e.g., combustion, rust).
Temperature scales:
Fahrenheit: uses °F; formula to convert to Celsius: TF = 1.8 imes TC + 32.
Celsius: uses °C; freezing point 0 °C, boiling point 100 °C.
Kelvin: uses K; absolute zero at 0 K; relation to Celsius: TC = TK - 273. (Transcript uses 0 K = -273 °C; 4 K ≈ -269 °C; accurate value is 273.15 instead of 273.)
Energy concepts:
Kinetic energy: E_k = frac{1}{2} m v^2.
Potential energy: E_p = m g h.
Chemical energy and bond energy: energy stored in chemical bonds; energy released when bonds break and new substances form.
Real-world relevance:
Understanding purity vs mixture helps in predicting properties and separability.
State and energy concepts underpin material behavior in furnaces, engines, biology, and environmental science.
Temperature conversions are essential for lab measurements, international coursework, and thermodynamics.