Chemistry 1 Notes: Phase, Mixtures, Compounds, Density, and Energy

Phase and Mixtures

Phase concept introduced via a glass of salt water: when properly prepared, you don’t see salt; it looks like a single phase.
Homogeneous mixtures are called solutions; synonymous term used: solution.
Heterogeneous mixtures have non-uniform composition (e.g., Italian salad dressing, oil and water).
Separation of mixtures relies on physical properties and does not change the chemical identities of the substances involved.
Example: salt water can be separated into salt and water by physical methods without altering the identities of salt or water.

Separation Techniques (based on physical properties)

Decantation: pouring off the top layer when substances differ in density (e.g., oil and water).
Centrifugation: rapid spinning (e.g., separating blood components in a centrifuge).
Distillation: separation based on differences in boiling points.
Filtration: separation by state of matter (solids from liquids) using filter paper.
Chromatography: separation based on differences in intermolecular forces.
Magnetism: using a magnet to remove magnetic components (e.g., iron filings from dirt).
All of these methods rely on physical properties and do not change the chemical identities of the substances.
Labs in Chemistry 1, Chemistry 2, and Organic Chemistry will involve these techniques.

Filtration Demonstration

Beaker with soil and water (heterogeneous): filter paper in a funnel captures soil; water passes through the filter.
Demonstrates filtration as a physical separation technique.

Types of Mixtures (Homogeneous, Heterogeneous) and Examples

Homogeneous mixtures (solutions): uniform composition; examples include solid-liquid (salt in water) and liquid-liquid (gasoline—mixture of liquids).
Heterogeneous mixtures: non-uniform composition.
Solid-solid mixtures can form alloys with uniform composition (e.g., brass, 14-karat gold).
Gas-gas mixtures: air (primarily nitrogen and oxygen with argon, CO₂, water vapor, etc.). At a given location, air is typically uniform, but altitude can change composition.
In a solution: solute is the minor component dissolved; solvent is the major component that dissolves the solute (often water is the solvent for liquid solutions).

Solutions: Key Vocabulary and Concepts

Solute: substance being dissolved (minor component).
Solvent: substance doing the dissolving (major component).
Concentration depends on the ratio of solute to solvent.
In most liquid solutions, water is the solvent.

Compounds vs Mixtures

Compounds: two or more elements chemically combined; definite composition and fixed ratio.
Mixtures: can be separated into pure substances by physical methods.
Compounds have properties very different from their constituent elements.
Examples of ionic vs molecular compounds:
- Ionic compounds: contain ions and typically include a metal; e.g., sodium chloride (NaCl). Presence of a metal indicates ionic character.
- Salt example: common name rust; chemical name: iron oxide (FeO or Fe2O3 depending on oxidation state). In many cases you’ll see iron oxide as rust with Fe2O3; balancing and naming follow rules learned later.
- Molecular compounds: contain only nonmetals (e.g., H₂O, CO₂, C₈H₁₈ for octane, C₆H₁₂O₆ for sugar, CH₄, NH₃).
Fixed ratio and definite composition: compounds have a specific proportion by mass (percent composition).
Example: iron pyrite FeS₂ has 46.55% Fe and 53.45% S; formula derived from fixed mass percentages.
Relationship to formulas: fixed composition ratio is tied to the chemical formula; different compounds with the same elements can have very different properties (e.g., water vs hydrogen peroxide).
Important distinction: compounds cannot be separated into their elemental components by physical methods alone; mixtures can be separated by physical methods.

Percent Composition and Formulas

Percent composition is the fixed proportion of each element by mass in a compound.
Example: FeS₂ has Fe mass fraction 46.55% and S mass fraction 53.45%.
From percent composition, you can infer the simplest whole-number ratio of atoms to write the empirical formula; further steps yield the molecular formula.

Physical vs Chemical Properties

Physical properties: observed without changing composition. Examples from lecture:
- Color, odor, density, melting point, thermal conductivity, volume, state of matter.
Chemical properties: observed only by changing the substance’s identity (chemical composition).
- Example: sodium reacts with water or oxygen; these reactions change the substance’s identity.
The properties help classify materials and predict behavior in reactions and separations.

Physical vs Chemical Changes

Physical change: a change in form/state without changing chemical identity (e.g., melting ice to liquid water).
Chemical change: a change in chemical identity, forming new substances (chemical reactions).
Chemical reactions are written with reactants on the left, products on the right, and an arrow indicating conversion.
Example reactions discussed:
- Hydrogen gas reacts with oxygen gas to form water: $2\,\mathrm{H2} + \mathrm{O2} \rightarrow 2\,\mathrm{H_2O}$
- Iron metal reacts with oxygen to form rust (iron(III) oxide): $4\,\mathrm{Fe} + 3\,\mathrm{O2} \rightarrow 2\,\mathrm{Fe2O_3}$
The diatomic nature of some elements in elemental form is noted with the mnemonic “Hofprinkles” (H, N, O, F, Cl, Br, I exist as diatomic molecules, e.g., H₂, O₂).
Coefficients in reactions balance the number of atoms of each element on both sides.

Energy in Physical and Chemical Changes

Energy is defined as the capacity to do work.
Kinetic energy: energy of motion; related to temperature, which measures average kinetic energy of particles.
Potential energy: energy of position; energy stored in systems, including chemical bonds and stored energy in chemicals (relevant to chemical energy and reactions, e.g., nuclear energy concepts).
All changes (physical or chemical) involve energy changes; energy is conserved in processes.
Everyday analogy: energy conservation illustrated by an object with potential energy converting to kinetic energy during motion, with energy ultimately transferred to surroundings.

Quick Tools for Problem-Solving in Density (Density, Mass, Volume)

Density formula: $D = \frac{M}{V}$
Units: mass in grams (g), volume in milliliters (mL) or cubic centimeters (cm³).
1 cm³ = 1 mL; density units are typically g/cm³.
Example for a rectangular bar: volume $V = l \times w \times h = (3\ {\rm cm}) \times (40\ {\rm cm}) \times (0.5\ {\rm cm}) = 60\ {\rm cm^3}$
- Given mass $M = 100\ {\rm g}$ , density $D = \frac{100}{60} \approx 1.666\ \text{g cm}^{-3}$
- Significant figures discussion (as taught): the least precise input determines the number of significant figures in the result; in the example, half a centimeter (0.5 cm) and 40 cm were treated as having limited precision, which affects the final rounding. The instructor notes this can lead to different rounding outcomes; the practical approach in ALEKS is to apply the rules consistently.
Rounding and significant figures quick notes:
- When dividing, count significant figures in each input; the result cannot have more significant figures than the input with the fewest.
- Round after the calculation, not during it.
- The rounding of trailing zeros and zeros after the decimal point affects how many significant figures they count:
- Example: 20.4558 (6 sig figs) ÷ 2.13 (3 sig figs) → result is limited to 3 sig figs, so the final value should reflect the least precise input.
Unit handling reminders:
- Mass in grams, volume in cm³ or mL depending on context.
- If converting between units (e.g., mg to g or cm³ to mL), use metric prefixes; 1 cm³ = 1 mL; g and g-related conversions follow standard metric conversions.
Practical density problem flow:
- If given density and mass, solve for volume: $V = \frac{M}{D}$
- If given density and volume, solve for mass: $M = D \times V$
- Always check units to ensure consistency (e.g., g, mL, cm³).
Triangle shortcut for density problems: cover the variable you’re solving for to see the rearranged formula quickly (Density = M/V; M = D × V; V = M/D).

Equations and Notation Highlight

Density: $D = \frac{M}{V}$
Volume from dimensions: $V = l \times w \times h$
Mass from density and volume: $M = D \times V$
Mass per mole and fixed composition concepts can be explored later in empirical vs molecular formulas (not detailed in this lecture but introduced here).
Chemical reaction balance examples:
- Hydrogen and oxygen to water: $2\,\mathrm{H2} + \mathrm{O2} \rightarrow 2\,\mathrm{H_2O}$
- Iron and oxygen to rust: $4\,\mathrm{Fe} + 3\,\mathrm{O2} \rightarrow 2\,\mathrm{Fe2O_3}$
Diatomic elements mnemonic: Hofprinkles (H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂).

Summary of Key Takeaways

Mixtures can be separated using physical properties without changing identities; solutions are homogeneous mixtures with a solute and solvent.
Compounds are chemically combined substances with fixed composition ratios and distinct properties from their elements; distinguish ionic vs molecular based on composition and bonding.
Density is a powerful, size-independent property that acts as a chemical fingerprint; calculate using $D = \frac{M}{V}$ and pay careful attention to units and significant figures.
Physical changes do not alter chemical identity; chemical changes form new substances via chemical reactions, which are represented by balanced chemical equations.
Energy considerations (kinetic and potential) play a central role in both physical and chemical processes, with the principle of energy conservation guiding all transformations.
Diatomic molecules in elemental form (H₂, N₂, O₂, F₂, Cl₂, Br₂, I₂) are common in reactions and must be accounted for in reaction equations.
Practice rounding and significant figures consistently, especially in division and when converting units, as required by ALEKS and course assessments.

If you’d like, I can convert these notes into a printable study sheet or tailor a quiz-style set of questions based on these sections to help you study for your exam.