Comprehensive Study Notes: Properties, Matter, Measurements, Atomic Structure, Bonding, and Calculations
Types of properties
Physical properties — characteristics of a material that can be observed or measured without changing the substance
Examples: boiling point (BP), melting point (MP), color, density, odor, hardness, etc.
Chemical properties — describe how a material will interact with other substances
Examples: reactions with water, iron's behavior in water/air, sodium reacting with water, etc.
Changes in Matter
Physical changes — do not change the composition of a substance
Examples: change of state (melting, freezing, vaporization), ripping, breaking, tearing, etc.
Chemical changes — involve a change in the composition of the substance
Examples: rusting, cooking, digestion, burning, etc.
Matter, Substances, and Types of Pure Substances
Matter — anything with mass and volume
Substance — pure form of matter (one type of material present)
Element — a pure substance that cannot be broken down by chemical means into simpler substances
Compound — two or more elements chemically combined
Mixture — two or more substances not chemically combined
Homogeneous — particles are similar/identical and evenly distributed; looks the same throughout
Heterogeneous — uneven distribution of particles; looks different in different parts
Quick classifications (examples)
Substance, element or compound examples: H₂SO₄, Aluminum, CuSO₄, Chalk, Titanium
Mixture examples: Red Koolaid, Water w/ice, Milk (store), Milk (FFTC), Raw Hamburger, Pepperoni pizza
Mixture type notes:
Steel is typically treated as a homogeneous alloy (a homogeneous mixture)
Chunky Peanut Butter is a heterogeneous mixture due to chunky portions
Sugar, C₁₂H₂₂O₁₁ is a compound
Scientific Notation
Scientific notation defined: a number (coefficient) multiplied by ten to a power
Standard format rule: the coefficient has only one nonzero digit to the left of the decimal
Exponents indicate magnitude:
Positive exponent → number > 1
Negative exponent → number < 1
How to read the exponent position: count how many places the decimal point moves
Examples:
Change into/out of Scientific Notation
Practice conversions between standard form and scientific notation (examples listed in notes):
0.0135 →
2.73 × 10^{-3} →
17,250 →
1.50 × 10^{0} → 1.50
7.5 × 10^{2} → 750
Measurements and the Metric System
Measurements require both a number and a unit
The metric system is the preferred system in science
Mass — amount of matter in an object; unit: grams (g)
Volume — space an object takes up; unit: liters (L)
Length — distance between two points; unit: meter (m)
Metric Prefixes (prefix chart)
Mega — M — 1 × 10^{6}
Kilo — k — 1 × 10^{3}
Hecta — h — 1 × 10^{2}
Deci — d — 1 × 10^{-1}
Centi — c — 1 × 10^{-2}
Milli — m — 1 × 10^{-3}
Micro — u — 1 × 10^{-6}
Unit Analysis (Dimensional Analysis)
Definition: converting from one set of units to another using conversion factors
A conversion factor is a fraction formed from a known equivalence
Example: If 1 ft = 12 in, then use or
Factors are set up so that identical units appear on the diagonal to cancel appropriately
Typical problem: find minutes in a given number of weeks (convert weeks → days → hours → minutes)
Example conversion chain: 1 week = 7 days, 1 day = 24 hours, 1 hour = 60 min
Result example: 4.57 weeks = 46065.6 min
Metric Conversions (practice problems and answers)
Common conversions:
1) 150 mm → 15 cm
2) 2.46 mg → 0.00246 g → 0.00000246 kg
3) 2.75 mL → 0.00275 L
4) 635 Mm → 6.35 × 10^{11} mm
5) 5.79 GL → 5.79 × 10^{9} L → 5.79 × 10^{11} cL
6) 875 ng → 8.75 × 10^{-7} g
7) 418.2 m → 0.4182 km
8) 127 Mg → 1.27 × 10^{10} cg
Significant Figures (sig figs)
Sig figs definition: all known digits plus one estimated digit
Atlantic-Pacific Rule (trailing/leading zeros):
No decimal point: count from Atlantic (left) to Pacific (right); skip leading zeros
Decimal point present: count from Pacific (right) to Atlantic (left); skip trailing zeros
Defined or exact numbers have infinite sig figs
Rounding: determine which place you are rounding to; look at next digit
0–4 -> drop; 5–9 -> round up
Rounding rules for operations:
Addition/Subtraction: answer has the same number of digits to the right of the decimal as the measurement with the fewest digits to the right of the decimal
Multiplication/Division: answer has the same number of sig figs as the measurement with the fewest sig figs
Arithmetic Practice (sig figs emphasis)
Examples of mixed operations include:
255.35 + 4.25 + 1.0 + 17.2 + 17.312 + 3.275
-1.25 - 27.6 + 0.13 + 22.475 + 4.3
17.350 + 18.21 – 2.15 – 0.175
Note: Follow sig fig rules for both types of operations; refer to problems in notes for practice
More Arithmetic Practice
Additional problems include:
4.75 × 0.4593 × 6 × 2.16 × 140
15.40 ÷ 2.0
(23.3 + 18.25) ÷ 2.049
Unit Conversions and Calculations (Practice set)
Sample problems: adding, subtracting, multiplying and dividing with units; using conversion factors to obtain correct units
American Units (Common Equivalencies)
12 inches = 1 ft
3 ft = 1 yd
36 inches = 1 yd
5280 ft = 1 mile
16 oz = 1 lb
2000 lb = 1 ton
8 oz = 1 cup
2 cup = 1 pint
2 pints = 1 quart
4 quarts = 1 gallon
1 gallon = 128 oz
1 quart = 32 oz
1 pint = 16 oz
Basic time units: 60 s = 1 min, 60 min = 1 h, 24 h = 1 d, 7 d = 1 week, etc.
Unit Analysis (extended)
A method of converting units using conversion factors; example: minutes in 4.57 weeks: 4.57 week × (7 days / 1 week) × (24 h / 1 day) × (60 min / 1 h) = 46065.6 min
Temperature and Density
Temperature is a measure of average kinetic energy, not a direct measure of hot/cold
Celsius and Kelvin:
Density: ratio of mass to volume, units typically g/mL or g/cc
Pure substances can be identified by density
If two substances with different densities are mixed, the denser substance tends to sink (buoyancy concept)
Density of pure water at 4°C is 1.000 g/mL
Specific Gravity
Specific gravity = density of a substance / density of water
Since density of water is 1.000 g/mL at 4°C, SG ≈ density in g/mL
Atoms, Nuclei, Isotopes, and Ions (Basics)
Parts of the atom:
Protons: positive charge, located in nucleus, ~1 amu
Neutrons: neutral, located in nucleus, ~1 amu
Electrons: negative charge, outside nucleus, ~0 amu
1 amu = 1.661 × 10^{-24} g
Nucleus: center of the atom, positively charged
Atomic Number (Z): number of protons; defines identity of the element
Mass Number (A): protons + neutrons; always a whole number
Isotopes: atoms of the same element with different numbers of neutrons (different A)
Ions: atoms or groups with a net charge; charge = protons − electrons (positive if more protons, negative if more electrons)
Atomic Weight (Atomic Mass): weighted average of all isotopes; not necessarily a whole number
Isotopes and Notation
Example notation: Zn-65 indicates zinc with mass number 65 (an isotope)
Atomic weights are often not whole numbers due to multiple isotopes
Periodic Table Concepts
Groups (columns) → elements with similar properties; 18 groups
Periods (rows) → elements with varying properties; 7 periods
Classification: metals, nonmetals, and metalloids
Key families to know: alkali metals, alkaline earth metals, halogens, noble gases, main-group elements, transition elements, inner transition elements
Valence Electrons: electrons in the highest energy level; important for bonding
Octet Rule: atoms gain or lose electrons to achieve a full outer shell of 8 valence electrons
Ions and Bonding (ionic vs covalent)
Ions: positively charged (cations) or negatively charged (anions)
Cations: typically metals
Anions: typically nonmetals or metalloids
Ionic Bonds: electrostatic attraction between oppositely charged ions; usually metal + nonmetal; high melting points; conduct electricity when dissolved in water; typically solids at room temperature
Covalent Bonds: atoms share electrons; usually nonmetals or metalloids; lower melting points; do not conduct electricity in water; may be solid, liquid, or gas at room temperature
Ionic Compounds: Writing Formulas and Names
General rule: write the voltage (charge) of the cation first, then the anion
Cross charges to give subscripts (do not include the charges in the final formula)
If a subscript is 1, drop it
Use parentheses for polyatomic ions when needed, with a subscript outside the parentheses
Finally, reduce subscripts to the smallest whole-number ratio
Polyatomic Ions and Roman Numeral Nomenclature (examples)
Polyatomic ions listed: NH₄⁺, ClO₃⁻, C₂H₃O₂⁻, OH⁻, NO₃⁻, CO₃²⁻, SO₄²⁻, PO₄³⁻
Ions requiring Roman numerals in names: Copper (I or II), Chromium (II or III), Iron (II or III), Cobalt (II or III), Nickel (II or III), Manganese (II or III), Lead (II or IV), Tin (II or IV), Mercury (I or II)
Naming Ionic Compounds (General Rules)
Name the positive ion first (include Roman numeral if there are multiple possible charges)
Name the negative ion second
If the negative ion is monatomic, end the name with -ide; if polyatomic, end with -ate (usually)
Practice: Writing Formulas from Names and Naming from Formulas
Name to formula practice: Sodium sulfate, Cobalt (II) carbonate, Zinc oxide, Calcium acetate, Nickel (II) phosphide, Ammonium sulfide, Tin (II) bromide, Potassium nitride, Cadmium phosphate, Silver nitrate
Formula to name practice: including Fe(NO₃)₃, NaF, MnCl₂, Al(ClO₃)₃, CrS, MgCO₃, CuSO₄, SrBr₂, Ba₃(PO₄)₂, CdS, ZnI, NH₄NO₃
Lewis Dot Diagrams (LDD) and Electron Configurations
LDD shows only valence electrons
Steps to build LDD:
1) Count total valence electrons in the molecule
2) Draw LDDs for each atom
3) Determine central bonding scheme (CBA: completed octets, bond order, etc.)
4) Arrange LDDs around the central atom so that single electrons are adjacent
5) Form bonds by sharing electron pairs (single, double, triple bonds)
6) Try to achieve 8 electrons around each atom (octet)
7) Ensure the total number of electrons is accounted for
8) Avoid lone electrons in the final diagram; exceptions: H (2 e-), B (6 e-)Example: Draw LDDs for F₂, CO₂, SiCl₄, PBr₃, SF₂, SiN, COS, SiOCl₂, COCl₂, NH₄⁺, CO₃⁻², OH⁻
Sample Electron-Configuration Notes (Memorization suggestions)
Carbon example: Carbon: 1s² 2s² 2p⁴
Sulfur example: Sulfur: 1s² 2s² 2p⁶ 3s² 3p⁴
Argon example: Argon: 1s² 2s² 2p⁶ 3s² 3p⁶
Use the periodic table as a guide to fill orbitals in order of increasing energy
Orbital types: s (2 e⁻), p (6 e⁻), d (10 e⁻), f (14 e⁻)
Ground state = lowest-energy arrangement of electrons
Electron Configurations and Notation (Zn example)
Example electron configuration for Zn: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰
Last entry interpretation: 3d¹⁰ means energy level 3, d-orbital, 10 electrons in that orbital
Periods correspond to energy levels; s-block on the left, p-block on the right, d-block in the middle
Use the periodic table to assign electrons according to position
Writing Formulas for Ionic Compounds (Detailed Rules)
Know the names, symbols, and charges of ions
Write the positive ion with its charge first, followed by the negative ion
Cross the charges to determine subscripts; ignore the charges in the final formula
If a subscripts equals 1, drop the 1
If using a polyatomic ion more than once, enclose it in parentheses and add a subscript
Reduce the subscripts to the smallest whole-number ratio
Example Ionic Formulas (Practice Questions)
Examples: Al and S; Cl and Ca; Ba and O; NH₄ and SO₄; Fe(II) and P
Example Ionic Names and Formulas (Continued)
Names for compounds such as Fe(NO₃)₃, NaF, MnCl₂, Al(ClO₃)₃, CrS, MgCO₃, CuSO₄, SrBr₂, Ba₃(PO₄)₂, CdS, ZnI, NH₃, etc.
Covalent Compounds and Names
Prefix system for covalent compounds
Mono-, Di-, Tri-, Tetra-, Penta-, Hexa-, Hepta-, Octa-, Nona-, Deca-
Naming rule: use prefixes for both elements; second element ends with -ide
Common exception: do not use the prefix mono- for the first element
Example: CO₂ → carbon dioxide; P₄O₁₀ → tetraphosphorus decoxide
Covalent Naming Practice
Name or write the formula for:
Silicon dioxide, Selenium dichloride, Sulfur difluoride, Carbon monoxide, Disulfur trioxide, Dinitrogen monoxide, P₂O₅, P₄O₁₀
Additional practice: Cu₂SO₄, AgNO₃, Mg(C₂H₃O₂)₂, K₂CO₃, N₂O₄, SeF₂, AlCl₃, NF₃, Cd(ClO₃)₂
Additional Compound Formulas (Further Practice)
Write formulas for:
Tin (IV) sulfide, Chromium (II) hydroxide, Arsenic tribromide, Nitrogen dioxide, Zinc Phosphate, Silicon tetraiodide, Lead (II) acetate, Disulfur dichloride
Lewis Dot Diagrams (Expanded)
Key steps (reiterated):
Count total valence electrons
Draw LDDs for each atom
Determine the central atom and the connectivity
Arrange electrons to satisfy octets
Form bonds by sharing electrons
Avoid leaving any single electrons in the final diagram
End of Notes
Types of properties
Physical properties — These are characteristics of a material that you can observe or measure without changing what the substance actually is. Think about features you can see or test directly without altering its chemical makeup.
Examples: The temperature at which something boils (boiling point, BP), the temperature at which it melts (melting point, MP), its color, how much mass it has in a given volume (density), its smell (odor), or how hard it is.
Chemical properties — These describe how a material will react or interact with other substances. To observe a chemical property, you usually have to change the substance into something new.
Examples: How iron rusts when exposed to water and air, how vigorously sodium metal reacts with water, or whether a substance is flammable.
Changes in Matter
Physical changes — These changes do not alter the fundamental identity or chemical composition of a substance. The substance is still the same, just in a different form.
Examples: Changing state (like ice melting into water, which is still H₂O; water freezing into ice; or water evaporating into steam), or physically altering its shape (like ripping paper, breaking glass, or tearing fabric).
Chemical changes — These changes do involve a change in the composition of the substance, forming one or more new substances. This is often an irreversible process.
Examples: Iron rusting (iron turns into iron oxide), cooking an egg (proteins denature), digestion of food (complex molecules break down), or burning wood (wood turns into ash, smoke, and gases).
Matter, Substances, and Types of Pure Substances
Matter — Anything that has mass (takes up space) and volume. Everything around us is made of matter.
Substance — A pure form of matter, meaning only one type of material is present with a consistent composition and properties throughout.
Element — The simplest type of pure substance. An element cannot be broken down into simpler substances by ordinary chemical processes. Each element is defined by its unique number of protons (atomic number).
Compound — A pure substance formed when two or more different elements are chemically combined in a fixed ratio. Compounds can be broken down into their constituent elements by chemical means, but not physical means.
Mixture — When two or more substances are combined but not chemically united. Each substance retains its individual properties and they can usually be separated by physical means.
Homogeneous mixture — Also known as a solution, this is a mixture where the particles are uniformly distributed and the mixture looks the same throughout. You cannot see individual components.
Example: Salt dissolved in water, air, or a brass alloy.
Heterogeneous mixture — A mixture where the particles are unevenly distributed, and you can usually see the different components or distinct phases.
Example: Sand and water, oil and vinegar, or a salad.
Quick classifications (examples)
Let's classify some common examples:
Substance, element or compound examples:
(Sulfuric Acid) — Compound
Aluminum (Al) — Element
(Copper Sulfate) — Compound
Chalk (mostly calcium carbonate, ) — Compound
Titanium (Ti) — Element
Mixture examples:
Red Koolaid — Homogeneous mixture (sugar, flavor, color dissolved in water)
Water with ice — Heterogeneous mixture (liquid water and solid water are distinct phases)
Milk (store-bought) — Homogeneous mixture (fats, proteins, water thoroughly mixed, appears uniform)
Milk (FFTC) — (Assuming this refers to 'full-fat, straight from the cow' which might separate) Heterogeneous mixture due to phase separation if not processed.
Raw Hamburger — Heterogeneous mixture (meat and fat are distinct, unevenly distributed)
Pepperoni pizza — Heterogeneous mixture (you can see discrete ingredients like crust, sauce, cheese, pepperoni)
Mixture type notes:
Steel is typically treated as a homogeneous alloy (a homogeneous mixture) because its components (iron and carbon) are uniformly combined at a microscopic level.
Chunky Peanut Butter is a heterogeneous mixture because you can see and feel different parts (smooth peanut butter and solid peanut chunks).
Sugar () is a compound, not a mixture, as its elements are chemically bonded together in a fixed ratio.
Scientific Notation
Scientific notation defined: A convenient way to express very large or very small numbers. It involves writing a number as a coefficient multiplied by ten raised to a power.
Standard format rule: The coefficient (the number before the '') must have only one non-zero digit to the left of the decimal point. For example, is correct, but is not in standard form.
Exponents indicate magnitude:
Positive exponent (e.g., ) → The original number is greater than 1. A larger positive exponent means a larger number.
Negative exponent (e.g., ) → The original number is less than 1 (a fraction or decimal). A larger negative exponent (further from zero) means a smaller number.
How to read the exponent position: The exponent tells you how many places the decimal point moved (and in which direction) to get to the standard scientific notation format.
Examples:
(Decimal moved 2 places to the left)
(Decimal moved 2 places to the right)
(Decimal moved 3 places to the left)
(Decimal moved 3 places to the right)
(Decimal moved 1 place to the left)
(Decimal moved 1 place to the right)
(Any number to the power of zero is 1, so )
Change into/out of Scientific Notation
Practice converting between standard form and scientific notation:
From standard to scientific notation:
(Move decimal 2 places right)
(Move decimal 4 places left)
From scientific to standard notation:
(Move decimal 3 places left)
(No decimal movement)
(Move decimal 2 places right)
Measurements and the Metric System
Measurements require both a number and a unit: A measurement is incomplete without both. For example, saying