Chemistry - Chemical Reactions & Equations

Explain how chemical equations represent reactions by listing reactants on the left and products on the right, with coefficients balancing the equation under mass conservation and an arrow indicating direction.

Learn to write chemical equations by identifying reactants and products, specifying their formulas and states, and balancing the equation with coefficients per the conservation of mass.

Balance chemical equations using the law of conservation of mass, ensuring equal atom counts on reactants and products, and apply coefficients to balance.

Explore symbols in chemical equations, including arrows for direction, upper arrows for gaseous products, lower arrows for precipitates, reversible reactions, state symbols (s, l, g, aq), and triangle heat indicators.

Identify cations and anions and explain how losing or gaining electrons yields positive or negative charges in ionic substances. Define electro-positive and electro-negative elements and illustrate with examples.

Explore endothermic and exothermic reactions, energy changes in heating and burning, and the reactivity series that ranks metals by their tendency to lose electrons and form positive ions.

Explore the four main types of chemical reactions: combination (synthesis), decomposition, single replacement, and double replacement, and see how these patterns appear in daily life and laboratory experiments.

Explore combination or synthesis reactions where multiple reactants join to form a single product, such as Na plus Cl forming NaCl and carbon plus oxygen forming CO2.

A decomposition reaction splits one reactant into multiple products, often requires energy such as heat or electricity, and can be endothermic or exothermic, the opposite of a combination reaction.

Understand single replacement reactions where one element replaces another in a compound, forming cation or anion replacements based on reactivity, with zinc and hydrochloric acid, or sodium and water.

Explore double replacement reactions where cations and anions swap to form new compounds, such as NaOH with CaBr2 yielding NaBr and Ca(OH)2.

Learn how oxidation involves loss of electrons and reduction involves gain, driven by oxygen and hydrogen, with examples from iron with oxygen, copper oxide, carbon dioxide, and methanol.

Explore oxidation-reduction concepts through oxygen transfer and hydrogen transfer with examples like iron oxide reacting with carbon monoxide to form iron and carbon dioxide, and ethanol converting to ethanal.

Explore redox reactions, where oxidation and reduction occur together through electron transfer. See how metals like magnesium, zinc, and aluminum oxidize while oxygen, hydrogen, and copper ions reduce.

Burning magnesium ribbon in air shows a combination reaction with oxygen forming magnesium oxide, visible as a bright white flame and white ash, illustrating oxidation and signs of chemical reaction.

Observe the oxidation of copper powder to copper oxide as heated copper reacts with oxygen, turning brown to black and illustrating signs of chemical change and oxidation.

Explore the reaction between lead nitrate and potassium iodide, a double-displacement process that yields yellow lead iodide precipitate and potassium nitrate in solution.

Explore the double displacement reaction between sodium sulfate and barium chloride, yielding a barium sulfate precipitate and a sodium chloride solution, with signs of chemical change.

The experiment shows zinc reacting with dilute sulfuric acid to produce hydrogen gas, illustrating a single-displacement and redox reaction with safe, guided laboratory practice.

Students observe a single displacement reaction where iron displaces copper from copper sulfate, forming iron sulfate and depositing copper on the iron nail, with color changes from blue to green.

Explore the exothermic reaction between calcium oxide (quicklime) and water, forming calcium hydroxide, and observe signs like temperature rise and milky calcium carbonate when CO2 is introduced.

Heating ferrous sulfate crystals demonstrates decomposition and endothermic reaction, causing color change from light green to white to brown and releasing sulfur dioxide gas.