CSET Chem Domain 2/3

Ionic, Covalent, and Metallic Bonds

• Ionic

  • Metal + nonmetal

  • Electrons transferred

  • Forms ions

  • Metal loses electrons → cation (+)

  • Nonmetal gains electrons → anion (−)

  • Opposite charges attract and hold the ions together.

    • Example: NaCl

    • Na → Na⁺

    • Cl → Cl⁻

• Covalent

  • Nonmetal + nonmetal

  • Electrons shared

  • Forms molecules

  • Can be polar or nonpolar

  • Example: H₂O

• Metallic

  • Metal + metal

  • Electrons delocalized/free-moving

  • Example: Cu (copper)

Bond & Molecular Polarity

• Bond polarity: unequal sharing of electrons due to different electronegativities

• Nonpolar bond: electrons shared equally → H–H

• Polar bond: electrons shared unequally → H–Cl

  • More electronegative atom = δ−

  • Less electronegative atom = δ+

• Molecular polarity: depends on bond polarity + molecular shape

  • Symmetrical → polarities can cancel → CO₂ = nonpolar

  • Asymmetrical → polarities don't cancel → H₂O = polar

Reaction Types

Basic reaction types

• Synthesis

  • A + B → AB

• Decomposition

  • AB → A + B

• Single replacement

  • A + BC → AC + B

• Double replacement

  • AB + CD → AD + CB

• Combustion

  • CH₄ + 2O₂ → CO₂ + 2H₂O

Endothermic Reactions 

• Reaction ABSORBS energy/heat from the surroundings.

• Energy goes IN

• Surroundings get colder

• Products have more energy than reactants

  • ΔH is positive (+)

Exothermic Reactions

• Reaction RELEASES energy/heat into the surroundings.

• Energy goes OUT

• Surroundings get warmer

• Products have less energy than reactants

  • ΔH is negative (−)

Catalysts

• A catalyst LOWERS the activation energy.

• This makes the reaction FASTER.

  • The catalyst is not consumed by the reaction.

• It does NOT change the overall energy difference between reactants and products (ΔH). 

  • This is because catalyst only changes the pathway the reaction takes — it doesn't change the starting or or ending points

Factors Affecting Reaction Rate 

• Temperature ↑ = reaction rate ↑

  • Particles move faster so more collisions have enough energy to react.

• Concentration ↑ = reaction rate ↑

  • More particles so there are more frequent collisions.

• Surface area ↑ = reaction rate ↑

  • More exposed particles so there are more collisions.

• Catalyst → reaction rate ↑

  • Lowers activation energy (Ea) so Provides an easier reaction pathway.

Gibbs free energy/spontaneity (ΔG)

• Tells you whether a reaction is thermodynamically favorable/spontaneous.

• Spontaneous = can occur without needing continuous outside energy.

• Formula: 

  • ΔG = ΔH - TΔS

• ΔH = change in enthalpy (heat/energy)

• T = temperature in Kelvin

• ΔS = change in entropy/disorder

• unit kJ/mol

What it tells us: 

• -ΔG = Spontaneous/favorable

• +ΔG = Nonspontaneous

• ΔG = 0 Equilibrium

Le Chatelier's Principle 

• When a system at equilibrium is disturbed, it shifts to oppose the change and re-establish equilibrium.

• What can cause a shift?

  • Add more reactant → shifts toward products

  • Remove reactant → shifts toward reactants

  
  

  • Add more product → shifts toward reactants

  • Remove product → shifts toward products

Equilibrium Constants (K)

Tells you the relative amounts of products and reactants at equilibrium.

It helps tell you whether equilibrium favors products or reactants.

• Formula: 

  • aA+bB ⇌ cC+dD

    • 
      

• Pure solids and pure liquids → DON'T include

• Gases and aqueous substances → include

• What does the size of K mean?

  • K > 1 → more products → equilibrium favors products

  • K < 1 → more reactants → equilibrium favors reactants

  • K ≈ 1 → significant amounts of both

Oxidation State Rules

• Free element: oxidation state = 0 → O₂, H₂, Fe

• Monatomic ion: equals its charge → Na⁺ = +1

• Neutral compound: oxidation states add to 0

• Polyatomic ion: oxidation states add to overall charge

• O: usually −2

• H: usually +1

Oxidation vs. Reduction

• Oxidation: loses electrons → oxidation number increases

• Reduction: gains electrons → oxidation number decreases

• OIL RIG: Oxidation Is Loss, Reduction Is Gain

• Redox: oxidation + reduction happen together

Oxidizing vs. Reducing Agents

• Oxidizing agent: gets reduced bc it steals electrons from another substance

• Reducing agent: gets oxidized bc it gives up its electrons to another substance

Example:

• Zn+Cu²⁺ → Zn²⁺ + Cu

  • Zn: 0 → +2 → loses 2e^-1 = oxidized

  • Cu: +2 → 0 → reduced gains 2e^-1 = reduced 

• Zn (Reducing Agent)

• Cu (Oxidizing Agent) 

Ochem and Biochem

• Functional Groups

  • Alcohol → –OH

  • Aldehyde → –CHO

  • Ketone → C=O within chain

  • Carboxylic/organic acid → –COOH

  • Ether → R–O–R

  • Amine → –NH₂

  • Ester → R–COO–R

• Basic Organic Reactions

  • Substitution: one atom/group is replaced by another

  • Addition: atoms/groups add across a double or triple bond

  • Esterification: carboxylic acid + alcohol → ester + water

Enthalpy

• The change in heat/energy of a reaction.

  • ΔH < 0 → exothermic → releases heat.

  • ΔH > 0 → endothermic → absorbs heat.

• ΔH = H(products) − H(reactants)

• Measured in kJ/mol for reactions.

Specific Heat

• The amount of energy needed to raise the temperature of 1 gram of a substance by 1°C.

• Different substances have different specific heats.

  • High specific heat → takes more energy to heat up.

  • Low specific heat → heats up more easily.

• Unit: J/g°C

• Water = 4.18 J/g°C

Heat Transfer Mechanisms

• Conduction

  • Heat transfer through direct contact.

  • Example: A metal spoon gets hot in hot soup.

• Convection

  • Heat transfer through the movement of a fluid (liquid or gas).

  • Example: Warm air rises and cool air sinks.

• Radiation

  • Heat transfer through electromagnetic waves; doesn't need direct contact or a medium.

  • Example: Heat from the Sun reaches Earth.