⚗ Chemistry
From atoms to reactions - everything you need for the High School Completion exam
Video Resources
🎥 Professor Dave Explains
Chemistry, physics, biology, and earth science. Short focused videos that skip the fluff.
Atomic Structure
►
🔭 Inside the Atom
All matter is made of atoms. Every atom has a central nucleus surrounded by a cloud of electrons.
1
Protons (+) - positively charged particles inside the nucleus. The number of protons defines what element it is.
2
Neutrons (neutral) - no charge, inside the nucleus alongside protons. Add mass but not charge.
3
Electrons (-) - negatively charged, orbit the nucleus in shells/orbitals. Tiny mass compared to protons and neutrons.
📊 Atomic Number and Mass Number
Atomic Number = number of protons (what makes the element unique)
Mass Number = protons + neutrons
Neutrons = Mass Number - Atomic Number
A neutral atom has equal protons and electrons. Carbon (C) has atomic number 6: 6 protons, 6 electrons in a neutral atom, and 6 neutrons (mass 12).
⚒ Isotopes and Electron Shells
Isotopes are atoms of the same element (same proton count) with different numbers of neutrons. Carbon-12 has 6 neutrons; Carbon-14 has 8 neutrons. Same element, slightly different mass.
Electrons orbit in energy shells. The first shell holds up to 2 electrons, the second up to 8, the third up to 18. The outermost shell electrons (valence electrons) are what determine how the atom bonds with others.
Electrons orbit in energy shells. The first shell holds up to 2 electrons, the second up to 8, the third up to 18. The outermost shell electrons (valence electrons) are what determine how the atom bonds with others.
The Periodic Table
►
📊 Layout and Organization
The periodic table organizes all known elements by atomic number (number of protons), increasing left to right, top to bottom.
1
Periods (rows) - horizontal rows. Elements in the same period have the same number of electron shells.
2
Groups (columns) - vertical columns. Elements in the same group have the same number of valence electrons and similar chemical properties.
🌎 Key Groups to Know
| Group | Name | Key Trait |
|---|---|---|
| Group 1 | Alkali Metals | Very reactive, 1 valence electron, soft metals (Li, Na, K) |
| Group 2 | Alkaline Earth Metals | Reactive, 2 valence electrons (Mg, Ca) |
| Group 17 | Halogens | Very reactive nonmetals, 7 valence electrons (F, Cl, Br, I) |
| Group 18 | Noble Gases | Unreactive, full valence shell (He, Ne, Ar) |
🧰 Metals, Nonmetals, and Metalloids
Metals (left side and middle): shiny, conduct electricity and heat, malleable, ductile. Most elements are metals.
Nonmetals (upper right): poor conductors, brittle in solid form, many are gases at room temperature (H, O, N, Cl).
Metalloids / Semimetals (staircase border): properties between metals and nonmetals. Silicon (Si) is a semiconductor used in electronics.
Nonmetals (upper right): poor conductors, brittle in solid form, many are gases at room temperature (H, O, N, Cl).
Metalloids / Semimetals (staircase border): properties between metals and nonmetals. Silicon (Si) is a semiconductor used in electronics.
🔬 Reading an Element Box
6
C
Carbon
12.011
Top number = Atomic Number (protons) = 6
Symbol = C
Name = Carbon
Bottom number = Atomic Mass (average mass of all isotopes) = 12.011
C
Carbon
12.011
Top number = Atomic Number (protons) = 6
Symbol = C
Name = Carbon
Bottom number = Atomic Mass (average mass of all isotopes) = 12.011
Atomic mass is not always a whole number because it is the weighted average of all naturally occurring isotopes of that element.
Chemical vs. Physical Changes
►
⚖ Physical Changes
A physical change alters the form or appearance of matter but does NOT create a new substance. The chemical composition stays the same.
Examples: cutting paper, melting ice, dissolving salt in water, crushing a can, tearing fabric, boiling water.
Physical changes are usually reversible. Melted ice can refreeze. Dissolved salt can be recovered by evaporating the water.
Examples: cutting paper, melting ice, dissolving salt in water, crushing a can, tearing fabric, boiling water.
Physical changes are usually reversible. Melted ice can refreeze. Dissolved salt can be recovered by evaporating the water.
🔥 Chemical Changes
A chemical change (chemical reaction) produces one or more new substances with different properties. The original substance is gone.
Signs of a chemical change:
Chemical changes are usually irreversible. Burned wood cannot become wood again.
Signs of a chemical change:
1
Color change - iron turns orange when it rusts
2
Gas production - bubbles forming (baking soda + vinegar)
3
Precipitate forms - a solid appears from mixing two solutions
4
Temperature change - heat released (burning) or absorbed (cold packs)
5
Light or sound produced - fireworks, explosions
Chemical changes are usually irreversible. Burned wood cannot become wood again.
Examples Side by Side
Physical: Ice melting, sugar dissolving, glass breaking, water boiling, stretching rubber
Chemical: Iron rusting, wood burning, food digesting, milk souring, fireworks exploding, baking a cake
Chemical: Iron rusting, wood burning, food digesting, milk souring, fireworks exploding, baking a cake
Elements, Compounds, and Mixtures
►
🔭 Pure Substances
1
Element: a pure substance made of only one type of atom. Cannot be broken down into simpler substances by chemical means. Examples: gold (Au), oxygen (O), hydrogen (H). There are 118 known elements.
2
Compound: two or more elements chemically bonded together in a fixed ratio. Has different properties from its elements. Examples: water (H₂O), table salt (NaCl), sugar (C₆H₁₂O₆). Can only be separated by chemical reactions.
🫕 Mixtures
A mixture is two or more substances physically combined. Each substance keeps its own properties. Can be separated by physical methods.
Homogeneous mixture (solution): uniform throughout - every part looks the same. Examples: salt water, air, steel, lemonade.
Heterogeneous mixture: non-uniform - you can see different parts. Examples: salad, sand and gravel, pizza, oil and water.
Homogeneous mixture (solution): uniform throughout - every part looks the same. Examples: salt water, air, steel, lemonade.
Heterogeneous mixture: non-uniform - you can see different parts. Examples: salad, sand and gravel, pizza, oil and water.
🔬 Methods of Separation
1
Filtration: separates solids from liquids (sand from water through filter paper).
2
Evaporation: removes liquid to leave dissolved solid (recover salt from salt water).
3
Distillation: separates liquids with different boiling points (purifying water, making alcohol).
4
Magnetism: separates magnetic materials (iron filings from sand).
Chemical Bonding Basics
►
⚡ Why Atoms Bond
Atoms bond to achieve a full outer electron shell (usually 8 electrons for most atoms - the octet rule). A full outer shell is the most stable configuration. Noble gases (Group 18) already have full shells, which is why they don't react.
🔭 Ionic Bonds
Ionic bonds form between a metal and a nonmetal. One atom transfers electrons to the other.
Example: Sodium (Na) has 1 valence electron. Chlorine (Cl) has 7 valence electrons and wants 1 more. Na gives its electron to Cl.
Example: Sodium (Na) has 1 valence electron. Chlorine (Cl) has 7 valence electrons and wants 1 more. Na gives its electron to Cl.
Na → Na⁺ + e⁻ Cl + e⁻ → Cl⁻
Now Na⁺ and Cl⁻ attract each other (opposite charges). Ionic compounds form crystal lattice structures, have high melting points, and conduct electricity when dissolved in water.
🔭 Covalent Bonds
Covalent bonds form between two nonmetals. Instead of transferring electrons, atoms share electrons.
Example: Two hydrogen atoms each share their 1 electron to form H₂ - both atoms now effectively have a full shell.
Single bond: share 1 pair of electrons (H-H). Double bond: share 2 pairs (O=O). Triple bond: share 3 pairs (N≡N).
Electronegativity measures how strongly an atom attracts electrons. The bigger the difference between two atoms, the more ionic (polar) the bond.
Example: Two hydrogen atoms each share their 1 electron to form H₂ - both atoms now effectively have a full shell.
Single bond: share 1 pair of electrons (H-H). Double bond: share 2 pairs (O=O). Triple bond: share 3 pairs (N≡N).
Electronegativity measures how strongly an atom attracts electrons. The bigger the difference between two atoms, the more ionic (polar) the bond.
Quick Comparison
Ionic: Metal + Nonmetal, electron TRANSFER, forms ions, crystal solids, high melting point, conducts when dissolved. (NaCl, MgO)
Covalent: Nonmetal + Nonmetal, electron SHARING, forms molecules, lower melting points, often gas or liquid at room temp. (H2O, CO2, CH4)
Covalent: Nonmetal + Nonmetal, electron SHARING, forms molecules, lower melting points, often gas or liquid at room temp. (H2O, CO2, CH4)
Balancing Chemical Equations
►
⚖ Law of Conservation of Mass
Matter cannot be created or destroyed in a chemical reaction. The number of each type of atom must be the same on both sides of the equation.
To balance: add coefficients (numbers in front of formulas) - NEVER change subscripts inside a formula, as that changes the substance.
To balance: add coefficients (numbers in front of formulas) - NEVER change subscripts inside a formula, as that changes the substance.
📊 Worked Example 1: Water
1
Unbalanced: H₂ + O₂ → H₂O
2
Count atoms: Left side: 2H, 2O. Right side: 2H, 1O. Oxygen is unbalanced.
3
Put a 2 in front of H₂O to get 2 oxygen on the right: H₂ + O₂ → 2H₂O
4
Now hydrogen: Left: 2H. Right: 4H (2 x H₂O). Unbalanced.
5
Put a 2 in front of H₂: 2H₂ + O₂ → 2H₂O
6
Check: Left: 4H, 2O. Right: 4H, 2O. Balanced!
2H₂ + O₂ → 2H₂O
📊 Worked Example 2: Combustion of Methane
1
Unbalanced: CH₄ + O₂ → CO₂ + H₂O
2
Count: Left: 1C, 4H, 2O. Right: 1C, 2H, 3O.
3
Balance H: 4H on left needs 2H₂O: CH₄ + O₂ → CO₂ + 2H₂O
4
Count O now: Left: 2O. Right: 2+2 = 4O. Need 2O₂ on left.
5
Balanced: CH₄ + 2O₂ → CO₂ + 2H₂O
CH₄ + 2O₂ → CO₂ + 2H₂O
📊 Worked Example 3: Iron and Oxygen
Unbalanced: Fe + O₂ → Fe₂O₃
Balanced: 4Fe + 3O₂ → 2Fe₂O₃
Check: Left: 4Fe, 6O. Right: 4Fe (2x2), 6O (2x3). Balanced! This is the reaction that produces iron oxide (rust).
Acids, Bases, and the pH Scale
►
💉 The pH Scale
pH Scale: 0 -------- 7 -------- 14
Acidic Neutral Basic (Alkaline)
pH 0-6 = Acid (more H⁺ ions)
pH 7 = Neutral (pure water)
pH 8-14 = Base (more OH⁻ ions)
The pH scale is logarithmic: pH 4 is 10 times more acidic than pH 5, and 100 times more acidic than pH 6.
Acidic Neutral Basic (Alkaline)
pH 0-6 = Acid (more H⁺ ions)
pH 7 = Neutral (pure water)
pH 8-14 = Base (more OH⁻ ions)
☀️ Acids
Acids donate H⁺ (hydrogen) ions when dissolved in water. pH below 7.
Properties: sour taste (do not taste acids in lab!), react with metals to produce hydrogen gas, turn blue litmus paper pink.
Examples: HCl (hydrochloric acid, in stomach), H₂SO₄ (sulfuric acid, car batteries), CH₃COOH (acetic acid, vinegar), citric acid (lemons).
Properties: sour taste (do not taste acids in lab!), react with metals to produce hydrogen gas, turn blue litmus paper pink.
Examples: HCl (hydrochloric acid, in stomach), H₂SO₄ (sulfuric acid, car batteries), CH₃COOH (acetic acid, vinegar), citric acid (lemons).
◎ Bases
Bases accept H⁺ ions (or donate OH⁻ ions) when dissolved in water. pH above 7.
Properties: bitter taste, feel slippery/soapy, turn pink litmus paper blue.
Examples: NaOH (sodium hydroxide, drain cleaner), NH₃ (ammonia, cleaning products), NaHCO₃ (baking soda, pH 8.3), Mg(OH)₂ (milk of magnesia, antacid).
Properties: bitter taste, feel slippery/soapy, turn pink litmus paper blue.
Examples: NaOH (sodium hydroxide, drain cleaner), NH₃ (ammonia, cleaning products), NaHCO₃ (baking soda, pH 8.3), Mg(OH)₂ (milk of magnesia, antacid).
⚖ Neutralization
When an acid and a base react, they neutralize each other, producing water and a salt.
Acid + Base → Salt + Water
HCl + NaOH → NaCl + H₂O
Indicators are substances that change color based on pH. Litmus paper turns pink in acids and blue in bases. Universal indicator shows a range of colors. Phenolphthalein is colorless in acid and pink in base.
States of Matter
►
🧠 The Four States
1
Solid: definite shape and definite volume. Particles are tightly packed in fixed positions; they can only vibrate in place. Very high density. Examples: ice, rock, steel.
2
Liquid: definite volume but NO definite shape (takes the shape of its container). Particles can slide past each other but stay close together. Examples: water, mercury, molten lava.
3
Gas: NO definite shape or volume. Particles move rapidly and randomly, spread to fill any container. Much lower density than solids or liquids. Examples: oxygen, steam, carbon dioxide.
4
Plasma: ionized gas with freely moving charged particles (ions and electrons). The most common state of matter in the universe. Found in stars, lightning bolts, neon signs, and fusion reactors.
📊 Comparing the States
| State | Shape | Volume | Particle Motion |
|---|---|---|---|
| Solid | Definite | Definite | Vibrate in place |
| Liquid | Indefinite | Definite | Slide past each other |
| Gas | Indefinite | Indefinite | Move freely and rapidly |
| Plasma | Indefinite | Indefinite | Ions and electrons move freely |
Phase Changes
►
🔥 The Six Phase Changes
1
Melting: solid → liquid. Absorbs heat energy. Melting point of water = 0°C (32°F).
2
Freezing: liquid → solid. Releases heat energy. Freezing point of water = 0°C.
3
Vaporization/Evaporation: liquid → gas. Absorbs heat. Boiling point of water = 100°C (212°F).
4
Condensation: gas → liquid. Releases heat. Dew forming on grass in the morning.
5
Sublimation: solid → gas (skips liquid phase). Dry ice (solid CO₂) sublimes at room temperature.
6
Deposition: gas → solid (skips liquid phase). Frost forming on a cold window on a humid morning.
📈 Heating Curve
A heating curve shows temperature vs. heat added for a pure substance. Key features:
Temperature
^
| ___Gas
| ____Liquid/
|___Solid/ Vaporizing
| Melting
+------------------------> Heat Added
Flat sections = phase changes (all energy goes
into breaking bonds, not raising temperature)
^
| ___Gas
| ____Liquid/
|___Solid/ Vaporizing
| Melting
+------------------------> Heat Added
Flat sections = phase changes (all energy goes
into breaking bonds, not raising temperature)
The heat of fusion is the energy needed to melt a solid (or released when it freezes). The heat of vaporization is the energy needed to vaporize a liquid (much larger than heat of fusion for most substances).
The Mole Concept
►
🧸 What Is a Mole?
A mole is a counting unit for atoms and molecules - just like a "dozen" means 12, a "mole" means a very specific (enormous) number.
1 mole = 6.022 x 10²³ particles (Avogadro's Number)
6.022 x 10²³ = 602,200,000,000,000,000,000,000. That's why we need this unit - atoms are unimaginably small, so we count them in moles.
📊 Molar Mass
Molar mass = the mass of 1 mole of a substance, in grams. It equals the atomic mass from the periodic table.
1 mole of Carbon (C) = 12.0 grams (atomic mass = 12)
1 mole of Oxygen (O) = 16.0 grams (atomic mass = 16)
1 mole of H₂O = 2(1.0) + 16.0 = 18.0 grams/mol
For a compound, add up the molar masses of all atoms in the formula.
🔄 Converting Between Moles and Grams
grams = moles x molar mass
moles = grams / molar mass
Example
How many grams are in 3 moles of water (H₂O)?
Molar mass of H₂O = 18.0 g/mol
Mass = 3 mol x 18.0 g/mol = 54.0 grams
Molar mass of H₂O = 18.0 g/mol
Mass = 3 mol x 18.0 g/mol = 54.0 grams
Stoichiometry Basics
►
📊 What Is Stoichiometry?
Stoichiometry uses a balanced equation to calculate how much of each reactant is needed and how much product is formed. The coefficients in a balanced equation give the mole ratios.
🔄 Step-by-Step Process
1
Write the balanced chemical equation.
2
Identify the given substance and the unknown substance.
3
Convert given grams to moles (divide by molar mass).
4
Use the mole ratio from the balanced equation to find moles of unknown.
5
Convert moles of unknown to grams if needed (multiply by molar mass).
📊 Worked Example
Balanced equation: 2H₂ + O₂ → 2H₂O
Question: How many moles of H₂O are produced from 4 moles of H₂?
Solution
Mole ratio from equation: 2 mol H₂ produces 2 mol H₂O (ratio = 1:1)
So: 4 mol H₂ x (2 mol H₂O / 2 mol H₂) = 4 mol H₂O
So: 4 mol H₂ x (2 mol H₂O / 2 mol H₂) = 4 mol H₂O
Another Example: N2 + 3H2 → 2NH3
How many moles of NH₃ from 6 moles of H₂?
Mole ratio: 3 mol H₂ : 2 mol NH₃
6 mol H₂ x (2 mol NH₃ / 3 mol H₂) = 4 mol NH₃
Mole ratio: 3 mol H₂ : 2 mol NH₃
6 mol H₂ x (2 mol NH₃ / 3 mol H₂) = 4 mol NH₃