Reactions Unit Energy, Changes, Balancing equations, types of reactions, Solubility, Net ionic equations 

Enthalpy Almost all chemical reactions involve energy usually heat added or lost Enthalpy (∆ H) is the energy content of a system at constant pressure

Endothermic and Exothermic Reactions Exothermic rxns- Enthalpy ∆ H is always negative Energy released as new bonds formed Energy of products is < reactants Endothermic rxns- Enthalpy ∆H is always positive Energy absorbed as bond broken Energy of products is > reactants

Entropy (∆ S)  disorder Entropy is the degree of randomness in a substance Solids < Liquids< Gases Tendency in nature to move toward disorder +∆ S   more disorder - ∆ S   less disorder

Phase Changes Solid (high ordered, low entropy) Liquid (more random, higher entropy) Gas (max randomness, high entropy)

Physical Changes Substance divided into parts Entropy ↑ Large crystal broken into pieces Solid dissolves and dissociates

Chemical Changes If more products than reactants → entropy ↑ If products are simpler than reactants → entropy ↑ Compounds become less random 

Temperature Changes ↑ Temperature Example- Molecules move faster Disorder ↑ and entropy Example- 2KClO3 (s) ↔ 2KCl(s) +3O2 (g) +875kj Rxn goes to the right because ; Products simpler More products Reaction leads to a gas Exothermic

Balancing Chemical Reactions (Conservation of Mass Law) Write the unbalanced equation Formulas of reactants on left Formulas of products on right Arrow shows direction of reaction → ← ↔

Balancing (cont.) Treat → like an equal sign Balance and then rebalance each atom Balance formulas by ONLY PLACING COEFFICIENTS in front Subscripts will change the formula Balance the Equation Apply Law of Conservation of Mass Need same number of atoms on either side of equation Treat → like an equal sign List all types of atoms in center under →

More on Balancing Indicate the states of matter Use (g) for gases Use (s) for solids Use (l) for liquids Use (aq) for dissolved in water Use Pneumonic to Remember ! M-Metals-----balancing these first I –Ions---------Look for polyatomic ions . Do they stay intact? Non-metals—Look for Cl or S, these are common O—Oxygen is a “Brinclhof” 𝑂 2 diatomic molecule also 𝐵𝑟 2 𝐼 2 𝐻 2 𝐶𝑙 2 𝑁 2 𝐹 2

Types of Chemical Reactions 5 general and 2 special cases Decomposition Synthesis Double Replacement Single Replacement Combustion

Decomposition One reactant breaks into two or more products AB→ A + B product may be elements or simpler compounds Examples : HgO → Hg + O2 FeS → Fe + S CaC 𝑂 3 → CaO + CO2

Synthesis Two reactants  one product A +B  AB Elements combine to form compounds Mg + O2  MgO K + Cl2  KCl Two simple compounds combine to form a complex compound Na2O + CO2  Na2CO3 BaCl2 + O2  Ba(ClO3)2

Single Replacement/Displacement One element replaces another element in a compound Two possibilities for Single Replacement Cation switches AX + Y  YX + A Cu + AgNO3  Ag + Cu(NO3)2 Anion Switches A + XY  XA + Y Cl2 + NaBr  NaCl + Br2

Double Replacement/Displacement The cation and anion switch places Diatomic Molecules do NOT qualify for Double Replacement ONLY Single Replacement AB + XY  AY + XB K(OH) + H2(SO4)  K2(SO4) + H2O FeS + HCl  FeCl2 + H2S

Neutralization When and Acid and Base mix in the right proportions and neutralize each other Special case of double Replacement rxns Na(OH) + HCl  NaCl + H2O

Combustion Reaction of Oxygen gas with a hydrocarbon Hydrocarbon: a compound contain both Hydrogen and Carbon CxHy + O2  CO2 + H2O Two types of combustion Complete : all reactants for CO2 and H2O Incomplete: rxn dose not go to completion CO is produced as well Complete Combustion CH4 + O2  CO2 + H2O C2H5OH + O2  CO2 + H2O

Solubility The ability of a substance to dissolve for this class; in water (aqueous) Solute: what is being dissolved Solvent: the medium the substance is being dissolved into Solubility Rules Guidelines of which compounds will dissolve in water

On page 19 of the workbook

Am I Soluble? Ba(NO3)2 ZnS (NH4)2CO3 Fe2O3 PbCl2 Soluble Insoluble

Net ionic equations Recall-- Phase indicators Solids (s) Liquids (l) Gases (g) Aqueous (aq) Solutions which are dissolved in water determine if the compound breaks into ions

Process to solve Write a balanced equation with phase indicators Cross out spectator ions Ions that do not directly contribute to the reaction Balance for charge (subscripts) Re-write the equations with ions/compound that contribute to the reactions Balance for mass (coefficients) Determine if there is an insoluble product Separate all soluble compounds Do not separate: Solids Pure liquids Gases Insoluble compounds

Sample Problem Ni 2+(aq) +2(OH 1-)(aq)  Ni(OH)2(s) 2 2 ____Ni(NO3)2 + ____NaOH  _____ Ni(OH)2 + _____ NaNO3 Ni 2+(aq) +2(NO3)1-(aq)+2Na1+(aq)+2(OH 1-)(aq)  Ni(OH)2 + 2Na1+(aq)+ 2(NO3)1-(aq) Ni 2+(aq) +2(OH 1-)(aq)  Ni(OH)2(s) (aq) (aq) (aq) (s) (s)