Science

Chemical equation

A representation of a chemical reaction using chemical formulas, coefficients, and symbols.

Reactants

Substances that undergo chemical change, written on the left side of a chemical equation.

Products

New substances formed during a chemical reaction, written on the right side of a chemical equation.

Symbol $\rightarrow$

Indicates yields, produces, or forms in a chemical equation.

Coefficient

A whole number written before a chemical formula indicating the number of molecules or formula units.

Subscript

A small number written at the lower right of an element symbol indicating the number of atoms.

Symbol $(s)$

Denotes a substance in the solid state.

Symbol $(l)$

Denotes a substance in the liquid state.

Symbol $(g)$

Denotes a substance in the gas state.

Symbol $(aq)$

Denotes a substance that is aqueous or dissolved in water.

Symbol $\Delta$

Indicates that heat is supplied during the chemical reaction.

Catalyst

A substance that speeds up a chemical reaction without being permanently consumed.

Reason for balancing chemical equations

Atoms are neither created nor destroyed, so the number of atoms of each element must be equal on both sides.

Law of Conservation of Mass

Mass is neither created nor destroyed during a chemical reaction.

Fate of atoms during a chemical reaction

Atoms are rearranged to form new substances.

Immutable component when balancing chemical equations

Subscripts must never be changed when balancing.

Adjustable values when balancing chemical equations

Coefficients.

Inspection method

A trial-and-error method of adjusting coefficients until the number of atoms on both sides is equal.

Final step after balancing a chemical equation

Check that the number of atoms of every element is equal on both sides.

Standard coefficients used in a balanced equation

The smallest possible whole-number coefficients.

Balanced equation for $H_2 + O_2 \rightarrow H_2O$

$$2H_2 + O_2 \rightarrow 2H_2O$$

Balanced equation for $Fe + O_2 \rightarrow Fe_2O_3$

$$4Fe + 3O_2 \rightarrow 2Fe_2O_3$$

Balanced equation for $H_2 + Cl_2 \rightarrow HCl$

$$H_2 + Cl_2 \rightarrow 2HCl$$

Balanced equation for $Na + Cl_2 \rightarrow NaCl$

$$2Na + Cl_2 \rightarrow 2NaCl$$

Balanced equation for $Mg + O_2 \rightarrow MgO$

$$2Mg + O_2 \rightarrow 2MgO$$

Balanced equation for $N_2 + H_2 \rightarrow NH_3$

$$N_2 + 3H_2 \rightarrow 2NH_3$$

Balanced equation for $Al + O_2 \rightarrow Al_2O_3$

$$4Al + 3O_2 \rightarrow 2Al_2O_3$$

Rate of reaction

How fast reactants are converted into products.

Collision theory

The theory that reactant particles must collide with each other for a chemical reaction to occur.

Two requirements for a successful collision

Correct orientation and sufficient energy.

Correct orientation

Particles colliding in the proper position so chemical bonds can break and form.

Sufficient energy

Having enough kinetic energy during collision for a reaction to proceed.

Activation energy ($E_a$)

The minimum amount of energy needed for a reaction to occur.

Effect of low activation energy

The reaction can occur more easily and rapidly.

Effect of high activation energy

More energy is needed to start the reaction, making it slower or harder to occur.

Effect of increasing temperature on reaction rate

Particles move faster, leading to more frequent and energetic collisions, speeding up the reaction.

Effect of decreasing temperature on reaction rate

Particle motion slows down, resulting in fewer effective collisions and a slower reaction.

Concentration

The amount of particles of a substance present in a given volume.

Effect of increasing concentration on reaction rate

More collisions occur in a given volume, making the reaction faster.

Effect of increasing surface area on reaction rate

More particles are exposed for collisions, making the reaction faster.

Reason smaller particle size increases reaction rate

Smaller particles have a greater total surface area exposed for collisions.

Mechanism by which a catalyst speeds up a reaction

Provides an alternative reaction pathway with a lower activation energy.

Inhibitor

A substance that slows down or prevents a chemical reaction.

Corrosion

The gradual deterioration of metals due to chemical reactions with their surroundings.

Methods to prevent corrosion

Painting, coating, or protecting metal surfaces.

Heat of reaction

The overall energy absorbed or released during a chemical reaction.

Enthalpy change ($\Delta H$)

Another name for the heat of reaction at constant pressure.

Formula for enthalpy change ($\Delta H$)

$$\Delta H = E_{\text{products}} - E_{\text{reactants}}$$

Exothermic reaction

A reaction that releases heat energy into the surroundings.

Sign of $\Delta H$ in an exothermic reaction

Negative ($\Delta H < 0$).

Effect of an exothermic reaction on surroundings

Causes the surrounding temperature to increase (become warmer).

Potential energy of products in an exothermic reaction

Lower than the potential energy of the reactants.

Examples of exothermic reactions

Combustion, cellular respiration, hot packs, and bioluminescence.

Endothermic reaction

A reaction that absorbs heat energy from the surroundings.

Sign of $\Delta H$ in an endothermic reaction

Positive ($\Delta H > 0$).

Effect of an endothermic reaction on surroundings

Causes the surrounding temperature to decrease (become cooler).

Potential energy of products in an endothermic reaction

Higher than the potential energy of the reactants.

Examples