Reaction Predictions ! !

Types of Chemical Reactions

1) Single Displacement2) Double Displacement3) Decomposition4) Synthesis5) Combustion

Single Displacement

• A +BC AC + B• One element shoves the other element out!

Example 1:

• Ca + AlCl3 CaCl2 + Al

Example 2:

• AlCl3 + K

What do metals have to do with reaction predictions?

• Metals vary in reactivity – Lose electrons in a reaction – AND give the electrons to someone else

• Metal’s ability to give electrons depends on how reactive the metal is– How easily/quickly the metal wants to give up its

electrons

Activity Series of Metals

• Metals listed and arranged according to reactivity. • Metals will displace other metal ions in a

solution from any metal BELOW it• Generally,

– Metals > H2 on activity series will produce hydrogen gas (H2) when combined with an acid

– Metals < H2 cannot produce hydrogen gas (H2) from an acid

Double Displacement

• AB + CD AC + BD

• Elements switch partners !

Example 1:

• Na3PO4 + BaCl2 Ba3(PO4)2 + NaCl

Example 2:

K3PO4 + MgSO4

Decomposition

• AB A + B

• Breaking chemical compound up, going from BIG to SMALL !

Example 1:

• Au2O3 Au + O2

Example 2:

• H2O

Synthesis

• A + B AB

• Joining! Making new chemical compound !

Example 1:

• Mg + N2 Mg3N2

Example 2:

• Be + Cl2

Combustion

• HYDROCARBON (compound made up of just Cs and Hs) + O2 CO2 + H2O

• Chemical reactions involve a compound burning.

Example 1:

• C2H6 + O2 CO2 + H2O

Example 2:

• Propane (C3H8) burns

Thermodynamics

Energy

• Ability to do work • Units– Joules (J), we will use “kJ”• Can be converted to different types

• Energy change results from forming and breaking chemical bonds in reactions

Heat (q)

• Energy transfer between a system and the surroundings

• Transfer is instant from high----low temperature until equilibrium

• Temperature—– Measure of heat, “hot/cold”– the average kinetic energy of molecules

Heat (q) continued

• Kinetic theory of heat – Heat increase resulting in temperature change

causes an increase in the average motion of particles within the system.

• Increase in heat results in– Energy transfer– Increase in both potential and kinetic energies

Thermodynamics 101

• First Law of Thermodynamics– Energy is conserved in a reaction (it cannot be created

or destroyed)---sound familiar???

– Math representation: ΔEtotal = ΔEsys + ΔEsurr = 0• Δ= “change in” • ΔΕ= positive (+), energy gained by system• ΔΕ= negative (-), energy lost by system • Total energy = sum of the energy of each part in a chemical

reaction

Calorimetry

How do we find the change in energy/heat transfer that occurs in chemical reactions???

Calorimetry

• Experimentally “measuring” heat transfer for a chemical reaction or chemical compound

• Calorimeter– Instrument used to determine the heat transfer of a chemical reaction– Determines how much energy is in food – Observing temperature change within water around a reaction

container

** assume a closed system, isolated container– No matter, no heat/energy lost – Constant volume

Specific Heat• Amount of heat required to increase the temperature of 1g of a

chemical substance by 1°C

• Units: cal/g-K or J/g-K

• 4.184 J = 1 cal, K = 273 + °C• Allows us to calculate how much heat is released or absorbed by a

substance ! ! !

• Unique to each chemical substance – Al(s) = 0.901J/g°K

– H2O(l) = 4.18 J/g°K

Specific Heat Equations

• q = smΔΤ– s/Cp = specific heat (values found in reference

table)– m = mass in grams– ΔΤ= change in temperature

Example 1: How much energy is required to warm 420 g of water in a water bottle from 25C to

37C ?

Q = ? m = 420 gC(H2O (l)) = 4.18 J/g• C

ΔT = 37-25 = 12 C

Q = mc ΔT

Q = (420 g)(4.18 J/g• C)(12 C)Q = 21067 J or 21 kJ

“Coffee Cup” calorimeter

• Styrofoam cup with known water mass in calorimeter– Assume no heat loss on walls– Initial water temp and then chemical placed inside– Final temperature recorded

• Any temperature increase has to be from the heat lost by the substance SOOO– All the heat lost from the chemical reaction or

substance is transferred to H2O in calorimeter

• The specific heat of gold is 0.128 J/g°C. How much heat would be needed to warm 250.0 g of gold from 25°C to 100°C?

Example 3:

Heat of Fusion (Hf)

• Fusion means melting/freezing• amount of energy needed to melt/freeze 1g of a

substance

• Different for every substance – look on reference tables

• Q = mHf

Heat of Vaporization(Hv)

• Vaporization means boiling/condensing• amount of energy needed to boil/condense

1g of a substance• Different for every substance – look on

reference tables• Q = mHv

Examples:

• Calculate the mass of water that can be frozen by releasing 49370 J.

• Calculate the heat required to boil 8.65 g of alcohol (Hv = 855 J/g).

• Calculate the heat needed to raise the temperature of 100. g of water from 25 C to 63 C .

Phase Change Diagram

• The flat points represent a phase change – temperature does not change while a phase change is occurring even though heat is being added.

• Diagonal points represent the 3 phases

Enthalpy

Thermodynamics

Enthalpy (ΔH)

• 2 types of chemical reactions:

1)Exothermic—heat released to the surroundings, getting rid of heat, -ΔΗ

2)Endothermic—heat absorbed from surroundings, bringing heat in, +ΔΗ

**Enthalpy of reaction— amount of heat from a chemical reaction which is given off or absorbed, units = kJ/mol

Enthalpy of Reaction

DH = Hfinal - Hinitial Hinitial = reactants

Hfinal = products

If Hfinal > Hinitial then DH is positive

and the process is ENDOTHERMIC

If Hfinal < Hinitial then DH is negative

and the process is EXOTHERMIC

Enthalpy of ReactionHfinal < Hinitial and DH is negative

Enthalpy of ReactionHfinal > Hinitial and DH is positive

More Enthalpy

• The reverse of a chemical reaction will have an EQUAL but OPPOSITE enthalpy change

• HgO Hg + ½ O2 ΔH = + 90.83 kJ

• Hg + ½ O2 HgO ΔH = - 90.83 kJ

• SOOO-----total ΔH = 0

Methods for determining ΔH

1)Calorimetry

2)Application of Hess’ Law

3)Enthalpies of Formation

Guidelines for using Hess’ Law

– Use data and combine each step to give total reaction

– Chemical compounds not in the final reaction should cancel

– Reactions CAN be reversed but remember to reverse the SIGN on ΔH

Calculate DH for S(s) + 3/2O2(g) SO3(g)

knowing that

S(s) + O2(g) SO2(g) DH1 = -296.8 kJ

SO2(g) + 1/2O2(g) SO3(g) DH2 = -98.9 kJ

The two equations add up to give the desired equation, so

DHnet = DH1 + DH2 = -395.7 kJ

USING ENTHALPY

Example 4:

NO(g) + ½ O2 NO2 (g) ΔH° = ?

Based on the following: ½ N2(g) + ½ O2 NO (g) ΔH° = + 90.29 kJ

½ N2(g) + O2 NO2 (g) ΔH° = +33.2 kJ

Enthalpy of Formation (ΔHf°)

• Enthalpy for the reaction forming 1 mole of a chemical compound from its elements in a thermodynamically stable state.

– A chemical compound is formed from its basic elements present at a standard state (25°C, 1 atm)

– Enthalpy change for this reaction = ΔHf°

• ΔHf°= 0 for ALL elements in their standard/stable state.

Enthalpy of Formation cont.

• DHrxn = Hfinal – Hinitial Really,

– ΔHf (products) - ΔHf (reactants)

• Calculate ΔHrxn based on enthalpy of formation (ΔHf)– aA + bB cC + dD

ΔH° =[c (ΔHf°)C + d(ΔHf°)D] - [a (ΔHf°)A + b (ΔHf°)B ]

Entropy

Spontaneous vs. Nonspontaneous

1)Spontaneous Process – Occurs WITHOUT help outside of the system, natural– Many are exothermic—favors energy release to

create an energy reduction after a chemical reaction • Ex. Rusting iron with O2 and H2O, cold coffee in a mug

– Some are endothermic• Ex. Evaporation of water/boiling, NaCl dissolving in water

Spontaneous vs. Nonspontaneous

2) Nonspontaneous Process– REQUIRES help outside system to perform

chemical reaction, gets aid from environment

– Ex. Water cannot freeze at standard conditions (25°C, 1atm), cannot boil at 25°C

**Chemical processes that are spontaneous have a nonspontaneous process in reverse **

Entropy (S)

• Measure of a system’s disorder – The degree of randomness associated with particles (molecules, etc.)

• Disorder is more favorable than order • ΔS = S(products) - S(reactants)

– ΔS is (+) with increased disorder – ΔS is (-) with decreased disorder

• State function– Only dependent on initial and final states of a reaction

• Ex. Evaporation, dissolving, dirty house

When does a system become MORE disordered from a chemical reaction?

(ΔS > 0)1)Melting 2)Vaporization 3)More particles present in the products than

the reactants– 4C3H5N3O9 (l) 6N2 (g) + 12CO2 (g) + 10H2O (g) + O2 (g)

4)Solution formation with liquids and solids 5)Addition of heat, increasing temperature

Thermodynamic Laws

1st Law of Thermodynamics – Energy cannot be created or destroyed

2nd Law of Thermodynamics – The entropy of the universe is always increasing. – Naturally favors a disordered state

Gibbs Free Energy

DG is the change in Gibbs free energy.

DG can be calculated asDGo = DHo - TDSo

The term DH represents enthalpy or heat energy which is available to do work.

The term DS represents entropy or random motion which is not available to do work.

Gibbs Free EnergyGibbs Free Energy

Le’Chatlier’s Principle

Basic Concept

• Main concept in chemical equilibrium

• Change in a variable that alters the equilibrium of a system (chemical reaction) products a shift in the OPPOSITE direction– Reaction shifts to counteract the variable’s influence

• Reaction tries to get back to equilibrium state-----SO reaction shifts

Basic Concept (cont.)

• What happens to a chemical reaction when equilibrium shifts

• When one side of a chemical reaction is stressed, the reaction shifts to the side of LEAST stress !

Example:

A + B AB

A) Heat AB) Cool BC) Add ABD) Take out B

Gas and Pressure

Kinetic Molecular Theory

1) Gases are made up of TONS of particles, constantly moving, and spread out.

2) Gas particles drive straight until they hit/collide with something (ex. Wall, particles).

3) Small particles, HUGE space for them to roam! Gas volume mainly empty space

Kinetic Molecular Theory (cont.)

4) No force of attraction! –gas particles randomly move around

5) When gas particles collide with each other or a container wall, no kinetic energy is lost when they collide.

6) Temperature determines average kinetic energy of gas particles.

-not all gas particles are moving at the same kinetic energy

Acids and Bases: an Introduction

An acid H+ in waterAn acid H+ in water

ACIDS

HCl hydrochloricHBr hydrobromicHI hydroiodicHNO3 nitric

H2SO4 sulfuric

HNO3

The following are examples of strong acids.

Base OH- in water Base OH- in water

BASES

NaOH(aq) Na+(aq) + OH-(aq)

NaOH is a strong base

Strong bases are strong electrolytes and soluble in waterStrong bases are strong electrolytes and soluble in water

04m08an104m08an1

• Based on the concentration of H+ or OH- ions in a solution.

• Strong Acids/Bases: completely dissociate into ions in a solution.

• Weak Acids/Bases: do NOT completely dissociate into ions in a solution.

Strength of Acids and Bases

Conjugate Acids/Bases

• Acids and bases are related to each other through the addition/loss of hydrogen ions – Conjugate acid-base pairs

• Acids produce conjugate bases• Bases produce conjugate acids

Conjugate Examples

• HA + H2O H3O+ + A-

• HNO3 + NH3 NH4+ + NO3

-

Strength of Acids and Bases

• Based on the concentration of H+ or OH- ions in a solution.

• Strong Acids/Bases: completely dissociate into ions in a solution.

• Weak Acids/Bases: do NOT completely dissociate into ions in a solution.

“BIG 6”---Strong Acids(Know them!!)

• HClO4

• HI• HCl• HNO3

• HBr• H2SO4

Strong Bases (Know them!!)

• Group I metal hydroxides (NaOH, KOH, etc.)

• Soluble/Slightly soluble Group II metal hydroxides ( Ca(OH)2, Sr(OH)2, Ba(OH)2 )

Conjugate Acid/Base Strength

• Stronger the acid, the weaker the conjugate base

• Stronger the base, the weaker the conjugate acid

• Weak acids/bases have strong conjugate bases/acids

• Binary acids– containing only 2 elements (one is hydrogen)– 1) Prefix “hydro—” with binary acid– 2) root name for second element– 3) End the name with “IC acid”

• Oxyacids– acids containing hydrogen, oxygen, and a nonmetal– 1) use the given polyatomic ion name from anion– 2) add “IC acid”

Acid Nomenclature

• Write name

Example 1: HBr

• Write the name.

Example 3: H2SO4

Let’s take a look at water

• 2H2O H3O+ + OH-

• [H3O+][OH-]= 1x10-14M2

–[H3O+]= 1x10-7M–[OH-]= 1x10-7M–Kw= 1x10-14M2

Example 1

• An acid is added to water and gives a hydroxide ion concentration [OH-] of 1.0x10-12M. What is the hydrogen ion concentration [H+] ?

What is pH?

• Pouvoir hydrogene: “hydrogen power” • pH = measure of [H3O+]------- Acidity

• [H3O+] expressed in powers of 10– Ex. 10-14 to 10-1

pH Scale

• Range from 0-14.• NEUTRAL, pH=7. (pure water)• BASE, pH > 7. (ocean water, milk of magnesia, baking

soda)• ACID, pH < 7. (stomach acid/HCl, vinegar, soft drinks)

pH Scale

Equations

• pH = -log[H+]

• pOH = -log[OH-]

• pH + pOH = 14

• [H+][OH-] = 1x10 -14 M2

• Example 2: What is the pH of a solution with a hydrogen ion concentration of 1x10-4M ? Is this solution acidic or basic ?

• Example 4: Calculate the [H+] and [OH-] of a vinegar solution with a pH of 2.5.

Net Ionic Equations

1) Molecular– Form you are most familiar with– Reactants and products written as neutral

compounds

– Ex. KCl + NaF NaCl + KF

3 Forms of Chemical Equations

2) Full Ionic– All chemical compounds are written as ions if they

can be (ex. Strong electrolytes, strong acids, strong bases, etc.)

– Soluble compounds—dissociate into ions– Insoluble compounds—solid precipitate– USE SOLUBILITY RULES!!!

– Ex. Ba+2 + 2Cl- + 2Na+ + SO4-2 BaSO4 (s) + 2Na+ +

2Cl-

Chemical Equations (cont. )

• Insoluble compounds are written as SOLIDS

• Pure liquids and gases are written as is, not broken up into ions, electrically neutral.

• Which compounds are broken up into ions????– Soluble ionic compounds– Strong acids (HCl, HBr, HI, HNO3 , HClO4 , H2SO4)

– Strong Bases ( Group IA bases and Ca(OH)2, Sr(OH)2 , Ba(OH)2

How do we write a FULL IONIC equation?

• Write the full ionic equation

Ex. 2: 2CH3CO2H + Sr(OH)2 Sr(CH3CO2)2 + 2H2O

3) Net Ionic Equations– Next step after full ionic equation. – Elimination of Spectator Ions (ions found on

both sides of the equation, not changed with reaction)

– Includes chemical compounds and ions DIRECTLY involved in chemical reaction.

Forms of Chemical Equations (cont.)

Na2CO3 + Ca(NO3)2 2NaNO3 + CaCO3 (s)

Write net ionic equation

Example 3: