chem final

Topics Covered on the Final

1. Chemical Reactions & Equations

• Predict products of double replacement reactions

• Write full and net ionic equations

1. Mole Conversions

• Atoms ↔ Moles using Avogadro’s number 6.022 \times 10^{23}

• grams ↔ Moles using molar mass

• Percent composition

1. Empirical and Molecular Formulas

• Calculate from percent composition

• Determine molecular formula using molar mass

1. Hydrates

• Determine the formula and name based on mass loss

• Use mole ratios

1. Stoichiometry

• Identify limiting reactants

• Theoretical yield and excess reactant calculations

• Percent yield formula: \text{Percent Yield} = \left( \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \right) \times 100

1. Gas Laws

• Combined and Ideal Gas Law: PV = nRT

• \frac{V1}{T1} = \frac{V2}{T2}

• STP conditions (1 atm, 273.15 K)

1. Solutions

• Molarity: M = \frac{\text{mol solute}}{\text{L solution}}

• Dilution formula: M1V1 = M2V2

1. Acids & Bases

• Identify acid/base, conjugate acid/base pairs

• pH/pOH: \text{pH} = -\log[H^+], \text{pOH} = -\log[OH^-], \text{pH} + \text{pOH} = 14

1. Nomenclature

• Know names of common acids (HF, HNO₂, H₂CO₃, etc.)

1. Redox Reactions

• Assign oxidation numbers

• Identify oxidation and reduction

1. Intermolecular Forces

• Identify dispersion, dipole-dipole, and hydrogen bonding

• Ion-dipole interactions (especially in solutions like CaCl₂ in water)

1. Thermochemistry

• Review energy transfer, endo/exothermic, q = mc\Delta T, etc.

Important Formulas

• Avogadro’s Number: 6.022 \times 10^{23} \text{ particles/mol}

• Molar Mass: Sum of atomic masses (periodic table)

• Empirical Formula Steps:

  1. Convert % to grams

  2. Convert grams to moles

  3. Divide all by smallest number of moles

  4. Multiply to get whole numbers if needed

• Ideal Gas Law: PV = nRT, where R = 0.0821 L·atm/mol·K

• Percent Composition: % \text{Element} = \left( \frac{\text{mass of element in 1 mol of compound}}{\text{molar mass of compound}} \right) \times 100

Definitions to Know

• Limiting Reactant: The reactant that runs out first

• Empirical Formula: Simplest ratio of elements

• Molecular Formula: Actual number of atoms

• Hydrate: Compound with water molecules bound

• Oxidation: Loss of electrons

• Reduction: Gain of electrons

• Strong vs. Weak Acids/Bases: Complete vs. partial ionization

• Intermolecular Forces: Forces between molecules

Key Thermochemistry Formulas

1. Heat (q) Formula: q = mc\Delta T

• q = heat (J or kJ)

• m = mass (g)

• c = specific heat capacity (J/g·°C)

  • Water: 4.18 J/g·°C

• \Delta T = change in temperature (°C)

1. Phase Change Heat: q = n\Delta H

• n = moles

• \Delta H = enthalpy change (kJ/mol)

  • Use for phase changes like melting (\Delta H{\text{fusion}}) or boiling (\Delta H{\text{vaporization}})

1. Calorimetry: q{\text{lost}} = q{\text{gained}}

• Use to calculate heat transfer between substances in a system (e.g., metal in water)

1. Enthalpy of Reaction: \Delta H = \sum \Delta Hf^\circ \text{(products)} - \sum \Delta Hf^\circ \text{(reactants)}

• Use standard enthalpies of formation \Delta H_f^\circ (kJ/mol)

1. Hess’s Law:

• Add chemical equations and their enthalpy changes to find overall \Delta H

  • If a reaction is reversed, change the sign of \Delta H

  • If a reaction is multiplied, multiply \Delta H by the same factor

Important Thermochemistry Definitions

• Thermochemistry: Study of energy changes in chemical reactions.

• System: The part of the universe you’re focusing on (e.g., the chemical reaction).

• Surroundings: Everything outside the system.

• Endothermic: Absorbs heat (positive \Delta H); feels cold.

• Exothermic: Releases heat (negative \Delta H); feels hot.

• Enthalpy (H): Heat content of a system at constant pressure.

• Specific Heat Capacity (c): Amount of energy needed to raise 1 g of a substance by 1°C.

• Calorimeter: Device used to measure heat changes.

• Heat of Fusion: Heat required to melt 1 mol of a substance at its melting point.

• Heat of Vaporization: Heat required to vaporize 1 mol of a substance at its boiling point.

• Standard Enthalpy of Formation (\Delta H_f^\circ): Change in enthalpy when 1 mole of a compound forms from its elements in their standard states.

Important Definitions for Reactions

1. Reactants: Substances present before a chemical reaction.

2. Products: Substances formed by a chemical reaction.

3. Chemical Equation: A representation of a chemical reaction using symbols and formulas.
   Example: 2H2 + O2 \rightarrow 2H_2O

4. Law of Conservation of Mass: Matter is neither created nor destroyed in a chemical reaction. Equations must be balanced.

5. Types of Chemical Reactions:

• Synthesis (Combination): A + B \rightarrow AB

• Decomposition: AB \rightarrow A + B

• Single Replacement: A + BC \rightarrow AC + B

• Double Replacement: AB + CD \rightarrow AD + CB

• Combustion: Hydrocarbon + O2 \rightarrow CO2 + H_2O

1. Aqueous (aq): Dissolved in water

2. Precipitate: A solid that forms from a solution in a chemical reaction

3. Spectator Ions: Ions that do not change during the reaction (appear on both sides)

4. Net Ionic Equation: A simplified chemical equation that only shows species that actually change

Formulas & Procedures for Reactions

1. Balancing Chemical Equations: Adjust coefficients to have the same number of each atom on both sides.

2. Mole Ratio (from balanced equations): Used in stoichiometry to relate amounts of reactants and products.

3. Writing Ionic & Net Ionic Equations:

   1. Write full balanced molecular equation.

   2. Split all aqueous compounds into ions (full ionic equation).

   3. Cancel spectator ions → Net ionic equation

4. Example:

• \text{NaOH (aq) + HCl (aq)} \rightarrow \text{NaCl (aq) + H}_2\text{O (l)}

• Full Ionic: \text{Na}^+ + \text{OH}^- + \text{H}^+ + \text{Cl}^- \rightarrow \text{Na}^+ + \text{Cl}^- + \text{H}_2\text{O}

• Net Ionic: \text{OH}^- + \text{H}^+ \rightarrow \text{H}_2\text{O}

1. Solubility Rules (for predicting products):

• Nitrates (NO3^⁻), alkali metals, and ammonium (NH4^⁺) are always soluble

• Use solubility rules to determine if a precipitate will form

Key Definitions for Moles

1. Mole (mol): A unit that represents 6.022 \times 10^{23} particles (Avogadro’s number), which could be atoms, molecules, ions, or formula units.

2. Avogadro’s Number: 1 mole = 6.022 \times 10^{23} particles

3. Molar Mass: The mass of one mole of a substance (g/mol), equal to the sum of the atomic masses from the periodic table.

4. Representative Particles:

• Atoms (elements like Na or He)

• Molecules (covalent compounds like H₂O)

• Formula Units (ionic compounds like NaCl)

Conversion Map for Moles

• particles (atoms, molecules)

  • ↑ ↓

  • (\times 6.022\times10^{23}) (\div 6.022\times10^{23})

• mass (g) ↔ moles ↔ particles

  • ↑ ↓

  • (\div mtext{molar mass}) (\times mtext{molar mass})

• g/mol

Core Mole Conversion Formulas

1. Mass ↔ Moles:

• \text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}

• \text{mass} = \text{moles} \times \text{molar mass}

1. Particles ↔ Moles:

• \text{moles} = \frac{\text{particles}}{6.022 \times 10^{23}}

• \text{particles} = \text{moles} \times 6.022 \times 10^{23}

1. Volume (at STP) ↔ Moles (for gases):

• \text{moles} = \frac{\text{volume (L)}}{22.4}

• \text{volume} = \text{moles} \times 22.4

  • (Only valid at Standard Temperature and Pressure: 0°C and 1 atm)

1. Moles ↔ Atoms in a Compound:

• Multiply moles of compound by the number of atoms of an element per formula unit.

  • Example: 2 mol of H₂O → 2 \times 2 = 4 mol H atoms

Helpful Tips

• Always label your units and cancel them as you go.

• For compounds, calculate the total molar mass by adding up all atoms.

• Be extra careful with diatomic elements like H₂, O₂, N₂, etc.

Key Definitions for Formulas

1. Empirical Formula: The simplest whole-number ratio of atoms in a compound.

• Example: CH₂O is the empirical formula of glucose (C₆H₁₂O₆).

1. Molecular Formula: The actual number of atoms of each element in a molecule. It’s a multiple of the empirical formula.

2. Percent Composition: The percentage by mass of each element in a compound.

Formulas & Procedures for Formulas

1. Percent Composition

• % \text{Element} = \left( \frac{\text{mass of element in 1 mol of compound}}{\text{molar mass of compound}} \right) \times 100

1. Empirical Formula from Percent Composition Steps:

   1. Assume 100 g of compound (so % becomes g).

   2. Convert grams to moles for each element: \text{moles} = \frac{\text{grams}}{\text{atomic mass}}

   3. Divide each by the smallest number of moles.

   4. Multiply to get whole numbers, if needed: If you get something like 1.5, multiply all by 2. If you get 2.33, multiply all by 3, etc.

2. Molecular Formula from Empirical Formula

• Formula: \text{Molecular Formula} = (\text{Empirical Formula}) \times n

  • Where: n = \frac{\text{Molar Mass of Compound}}{\text{Molar Mass of Empirical Formula}}

Example Problem for Formulas

• Given: 40.00% C, 6.71% H, and 53.29% O

1. Assume 100 g → 40.00 g C, 6.71 g H, 53.29 g O

2. Convert to moles:

• C: \frac{40.00}{12.01} = 3.33 mol

• H: \frac{6.71}{1.008} = 6.66 mol

• O: \frac{53.29}{16.00} = 3.33 mol

1. Divide by smallest (3.33):

• C: 1, H: 2, O: 1 → Empirical formula = CH₂O

• If the molar mass is 180 g/mol, then: Molar mass of CH₂O = 30 g/mol

• \frac{180}{30} = 6 → Molecular formula = C₆H₁₂O₆

Key Definitions for Hydrate

1. Hydrate: A compound that includes water molecules chemically bound within its crystal structure.
   Example: \text{CuSO}4 \cdot 5\text{H}2\text{O} (copper(II) sulfate pentahydrate)

2. Anhydrate: The compound left behind after the water is removed (usually by heating).

• Example: CuSO₄ is the anhydrate of CuSO₄·5H₂O.

1. Water of Hydration: The water molecules in a hydrate’s structure.

Formulas & Calculations for Hydrate

1. Mass of Water Lost

• \text{Mass of Water} = \text{Mass of Hydrate} - \text{Mass of Anhydrate}

1. Moles of Water and Anhydrate

• \text{Moles of H}_2\text{O} = \frac{\text{Mass of Water Lost}}{18.02 \, \text{g/mol}}

• \text{Moles of Anhydrate} = \frac{\text{Mass of Anhydrate}}{\text{Molar Mass of Anhydrate}}

1. Mole Ratio of Water to Salt

• \text{Ratio} = \frac{\text{Moles of H}_2\text{O}}{\text{Moles of Anhydrate}} \Rightarrow \text{Round to nearest whole number}

• This gives the formula of the hydrate: \text{Salt} \cdot x\text{H}_2\text{O}

Example Problem of Hydrate

• Given: Mass of hydrate = 5.061 g

• Mass of anhydrate = 2.472 g → Mass of water = 5.061 - 2.472 = 2.589 g

1. Moles of water:

• \frac{2.589 \, \text{g}}{18.02 \, \text{g/mol}} \approx 0.144 \, \text{mol}

1. Moles of MgSO₄:

• \frac{2.472 \, \text{g}}{120.37 \, \text{g/mol}} \approx 0.0205 \, \text{mol}

1. Ratio:

• \frac{0.144}{0.0205} \approx 7 \Rightarrow \text{Formula} = \text{MgSO}4 \cdot 7\text{H}2\text{O}

• → Name: Magnesium sulfate heptahydrate

Key Definitions for Stoichiometry

1. Stoichiometry: The calculation of reactant and product quantities in chemical reactions using balanced equations.

2. Mole Ratio: The ratio of moles of one substance to another, derived from the coefficients in a balanced chemical equation.

• Example: 2H2 + O2 \rightarrow 2H2O \Rightarrow \frac{2 \, \text{mol } H2}{1 \, \text{mol } O2}, \frac{2 \, \text{mol } H2O}{2 \, \text{mol } H_2}, \text{etc.}

1. Limiting Reactant: The reactant that gets used up first, limiting the amount of product formed.

2. Excess Reactant: The reactant that is not completely used up in the reaction.

3. Theoretical Yield: The maximum amount of product that can be formed from the given reactants.

4. Actual Yield: The amount of product actually obtained from the reaction (usually given in a lab).

5. Percent Yield: \text{Percent Yield} = \left( \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \right) \times 100

Core Stoichiometry Formulas

1. Moles to Moles:

• \text{mol A} \times \frac{\text{mol B}}{\text{mol A}} = \text{mol B}

1. Mass to Mass:

   1. Convert grams of A to moles of A

   2. Use mole ratio to find moles of B

   3. Convert moles of B to grams of B

2. \text{mass A} \xrightarrow{\div \text{molar mass A}} \text{mol A} \xrightarrow{\times \text{mol B/mol A}} \text{mol B} \xrightarrow{\times \text{molar mass B}} \text{mass B}

3. Limiting Reactant Steps:

   1. Convert both reactants to moles of product.

   2. The one that makes less product is the limiting reactant.

4. Excess Reactant Left Over:

• Start - Used = Left Over

Example of Stoichiometry

• Given: 25.0 g N₂ and 25.0 g H₂

• Reaction: N2 + 3H2 \rightarrow 2NH_3

1. Convert grams to moles:

• N₂: \frac{25.0}{28.02} \approx 0.893 \, \text{mol}

• H₂: \frac{25.0}{2.02} \approx 12.38 \, \text{mol}

1. Mole ratio:

• N₂ needs 3 mol H₂ per 1 mol N₂ → Needs 0.893 \times 3 = 2.679 mol H₂

• Available H₂ = 12.38 mol → Excess

• N₂ is limiting reactant

1. Product moles:

• 0.893 \, \text{mol N}2 \times \frac{2 \, \text{mol NH}3}{1 \, \text{mol N}2} = 1.786 \, \text{mol NH}3

1. Convert to grams if needed:

• 1.786 \times 17.03 \approx 30.4 \, \text{g NH}_3

Key Definitions for Gas Laws

1. Gas Laws: Describe the relationships between pressure (P), volume (V), temperature (T), and amount (n) of a gas.

2. Standard Temperature and Pressure (STP):

• 0°C (273.15 K) and 1 atm

• At STP, 1 mole of any ideal gas = 22.4 L

1. Kelvin Temperature: Always convert Celsius to Kelvin: K = °C + 273.15

Core Gas Law Formulas

1. Boyle’s Law (Pressure-Volume)

• P1V1 = P2V2

  • Inverse relationship: as pressure increases, volume decreases (T constant)

1. Charles’s Law (Volume-Temperature)

• \frac{V1}{T1} = \frac{V2}{T2}

  • Direct relationship: as temperature increases, volume increases (P constant)

1. Gay-Lussac’s Law (Pressure-Temperature)

• \frac{P1}{T1} = \frac{P2}{T2}

  • Direct relationship: as temperature increases, pressure increases (V constant)

1. Combined Gas Law

• \frac{P1V1}{T1} = \frac{P2V2}{T2}

  • Use when P, V, and T all change

1. Ideal Gas Law

• PV = nRT

  • P = pressure (atm)

  • V = volume (L)

  • n = moles

  • R = 0.0821 L·atm/mol·K

  • T = temperature (K)

1. Gas Density & Molar Mass Density: d = \frac{PM}{RT}

   • Where M is molar mass

• Molar mass: M = \frac{dRT}{P}

1. Avogadro’s Law (Volume-Moles)

• \frac{V1}{n1} = \frac{V2}{n2}

  • More moles = more volume (at same T and P)

1. Dalton’s Law of Partial Pressures

• \text{P}{\text{total}} = P1 + P2 + P3 + \ldots

Quick Tips for Gas Laws

• Always use Kelvin for temperature in gas law calculations.

• Check units—especially for volume (L) and pressure (atm).

• For gas stoichiometry at STP: 1 mol gas = 22.4 L

Key Definitions for Solutions

1. Solution: A homogeneous mixture of a solute dissolved in a solvent.

2. Solute: The substance being dissolved (usually present in smaller amount).

3. Solvent: The substance doing the dissolving (usually water in aqueous solutions).

4. Aqueous (aq): A substance dissolved in water.

5. Molarity (M): A measure of concentration: M = \frac{\text{mol of solute}}{\text{L of solution}}

6. Dilution: The process of adding solvent to a solution to decrease its concentration.

7. Saturated Solution: Contains the maximum amount of solute that can dissolve at a given temperature.

8. Supersaturated Solution: Contains more solute than normally possible at that temperature (unstable).

9. Unsaturated Solution: More solute can still dissolve in the solvent.

Core Formulas for Solution

1. Molarity (Concentration)

• M = \frac{n}{V} = \frac{\text{mol of solute}}{\text{L of solution}}

• Convert grams to moles if needed: \text{mol} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}

1. Dilution Equation

• M1V1 = M2V2

  • M1, V1: concentration and volume of stock solution

  • M2, V2: concentration and volume of diluted solution

1. Percent by Mass

• % \text{mass} = \left( \frac{\text{mass of solute}}{\text{mass of solution}} \right) \times 100

1. Percent by Volume

• % \text{volume} = \left( \frac{\text{volume of solute}}{\text{volume of solution}} \right) \times 100

1. Molality (less common unless specified)

• m = \frac{\text{mol of solute}}{\text{kg of solvent}}

Other Useful Concepts for Solutions

• Solubility: How much solute dissolves in a given amount of solvent at a specific temperature.

• “Like dissolves like”: Polar solvents dissolve polar solutes; nonpolar dissolves nonpolar.

• Ion-Dipole Interactions: When ionic compounds dissolve in polar solvents like water.

Example Problem (Molarity) for Solutions:

• Q: What is the molarity of a solution containing 45 g of NaOH in 1.059 L of solution?

1. Find moles of NaOH:

• \text{mol} = \frac{45 \, \text{g}}{39.997 \, \text{g/mol}} \approx 1.125 \, \text{mol}

1. Use molarity formula:

• M = \frac{1.125}{1.059} \approx 1.06 \, \text{M}

Key Definitions for Acids and Bases

1. Acid:

• Arrhenius: Produces H⁺ (or H₃O⁺) in water

• Brønsted-Lowry: Proton (H⁺) donor

1. Base:

• Arrhenius: Produces OH⁻ in water

• Brønsted-Lowry: Proton (H⁺) acceptor

1. Conjugate Acid: Formed when a base gains a proton (H⁺)

2. Conjugate Base: Formed when an acid loses a proton (H⁺)

3. Strong Acids/Bases: Completely ionize in solution

• Example acids: HCl, HNO₃, H₂SO₄

• Example bases: NaOH, KOH, Ba(OH)₂

1. Weak Acids/Bases: Partially ionize in solution

• Example: HC₂H₃O₂ (acetic acid), NH₃ (ammonia)

Core Formulas for Acids and Bases

1. pH and pOH

• \text{pH} = -\log[H^+]

• \text{pOH} = -\log[OH^-]

1. pH + pOH Relationship

• pH + pOH = 14

1. Ion Concentration from pH or pOH

• [H^+] = 10^{-\text{pH}}

• [OH^-] = 10^{-\text{pOH}}

1. Kw — Ion Product of Water

• K_w = [H^+][OH^-] = 1.0 \times 10^{-14} \quad (\text{at 25°C})

Neutralization Reaction for Acids and Bases

• Acid + Base → Salt + Water

  • Example: HCl + NaOH → NaCl + H₂O

Naming acids

• Acid Formula Name Type

• No oxygen (HX) “Hydro- + -ic acid”

  • HCl → hydrochloric acid

• With oxygen ending in -ate “-ic acid”

  • HNO₃ → nitric acid

• With oxygen ending in -ite “-ous acid”

  • HNO₂ → nitrous acid

Example Problem: pH Calculation

• Q: What is the pH of a solution with [H^+] = 1.995 \times 10^{-7}\, \text{M}?

• pH = -log(1.995 \times 10^{-7}) ≈ 6.70 → Slightly acidic

Basic Definitions for Nomenclature

1. Element: A substance made up of atoms that all have the same number of protons.

2. Compound: A substance formed when two or more elements chemically bond.

3. Molecule: The smallest unit of a compound that retains its chemical properties.

4. Ions: Atoms or molecules that have gained or lost electrons, resulting in a charged species.

Ionic Compounds Nomenclature

• Ionic compounds are made from metal cations (positive ions) and nonmetal anions (negative ions).

• Naming Ionic Compounds: Name the cation (metal) first, followed by the anion (nonmetal).

  • For monatomic cations (e.g., Na⁺), use the element’s name.

  • For monatomic anions (e.g., Cl⁻), use the element’s root and add “-ide.”

  • If the metal can have multiple oxidation states (transition metals), include the oxidation state in Roman numerals in parentheses (e.g., Fe²⁺ → Iron(II)).

• Example: NaCl is Sodium Chloride.

• Formula for Ionic Compounds: The formula reflects the balance of charges to form a neutral compound. For Na⁺ and Cl⁻, one of each is needed for neutrality, giving the formula NaCl.

Covalent (Molecular) Compounds Nomenclature

• Covalent compounds form when two nonmetals share electrons.

• Naming Molecular Compounds:

  • The first element is named first, with its full element name.

  • The second element is named as if it were an anion, ending in “-ide.”

  • Prefixes (mono-, di-, tri-, etc.) are used to indicate the number of atoms of each element.

• Examples:

  • CO₂ = Carbon Dioxide.

  • N₂O₄ = Dinitrogen Tetroxide.

• Formula for Molecular Compounds: The formula indicates the number of atoms of each element.

Acids Nomenclature

• Acids are compounds that release hydrogen ions (H⁺) when dissolved in water.

• Naming Acids:

  • Binary Acids (two elements, often hydrogen + a halogen): Start with “hydro-” and end with “-ic acid.”

  • Example: HCl = Hydrochloric Acid.

  • Oxyacids (hydrogen + a polyatomic ion with oxygen): The name depends on the polyatomic ion.

  • If the ion ends in “-ate”, the acid name ends in “-ic acid.”

  • If the ion ends in “-ite”, the acid name ends in “-ous acid.”

  • Example: H₂SO₄ (sulphate ion) = Sulfuric Acid, H₂SO₃ (sulfite ion) = Sulfurous Acid.

Polyatomic Ions

• Polyatomic ions are ions made up of more than one atom.

• Common Polyatomic Ions:

  • Ammonium: NH₄⁺

  • Hydroxide: OH⁻

  • Sulfate: SO₄²⁻

  • Nitrate: NO₃⁻

  • Phosphate: PO₄³⁻

Oxidation States Nomenclature

• Oxidation states (or numbers) represent the charge on an atom in a molecule or ion.

  • Some elements have fixed oxidation states (e.g., alkali metals always have an oxidation state of +1), while others (like transition metals) can have multiple oxidation states.

Percent Composition and Empirical Formula Nomenclature

• Percent Composition: The percentage by mass of each element in a compound.

  • \text{Percent Composition} = \left( \frac{\text{Mass of element in compound}}{\text{Total mass of compound}} \right) \times 100

• Empirical Formula: The simplest whole number ratio of atoms in a compound.

  • Find the ratio of moles of each element in the compound and simplify to the smallest whole numbers.

Molar Mass Nomenclature

• The molar mass is the mass of one mole of a substance (in g/mol).

  • Add up the atomic masses of all atoms in the formula.

Stoichiometry Nomenclature

• Stoichiometry is the calculation of reactants and products in chemical reactions.

  • Use the mole ratio from the balanced chemical equation to find quantities of reactants or products.

Basic Definitions for Redox Reactions

1. Redox Reactions: A redox (reduction-oxidation) reaction involves the transfer of electrons between two species.

• One species undergoes oxidation (loses electrons), and the other undergoes reduction (gains electrons).

1. Oxidation: The process of losing electrons.

2. Reduction: The process of gaining electrons.

3. Oxidizing Agent (Oxidant): The substance that gains electrons and gets reduced.

4. Reducing Agent (Reductant): The substance that loses electrons and gets oxidized.

Oxidation States (Numbers) for Redox Reactions

1. Oxidation State: A measure of the degree of oxidation of an atom in a compound.

• The oxidation state of free