CHEMISTRY 161 Chapter 4. CHEMICAL REACTIONS 2 HgO (s) → 2Hg (l) + O 2(g) aq 1. properties of aqueous solutions 2. reactions in aqueous solutions a) precipitation.

CHEMISTRY 161

Chapter 4

CHEMICAL REACTIONS

2 HgO(s) → 2Hg(l) + O2(g)

aq

1. properties of aqueous solutions

2. reactions in aqueous solutions

a) precipitation reactions

b) acid-base reactions (proton transfer)

c) redox reactions (electron transfer)

1.PROPERTIES OF AQUEOUS SOLUTIONS

homogeneous mixture of two or more substances

solvent solute

substance in a large amount substance in a small amount

N2 gas phase O2

(air)

Ag solid phase Au(alloys)

H2O liquid phase NaCl(sea water)

EXP1

iodine in ethyl alcohol (C2H5OH)

EXP2

table salt in water (H2O)

does not conduct electricity(molecular solid)

I2

does conduct electricity(ionic solid)

Na+Cl-

AQUEOUS SOLUTION

solutes

solute

water (H2O)

electrolytes non-electrolytes

solution conducts electricity

solution does not conduct electricity

EXP3

electrolytes non-electrolytes

solution conducts electricity

solution does not conduct electricity

non-electrolyte weak electrolyte strong electrolyte

methanol

sugar

ethanol

water

dark bright

ionic compounds

(NaCl, KF)

NaOH

HCl

H2SO4

CH3COOH

HCOOH

HF

medium

EXP5

SOLUTION

concentration

SOLUTION

percentage concentration

% = g [solute] / g solvent X 100

12 g of sodium chloride are solved in 150 g of water. Calculate the percentage concentration

8 %

solubility of a solute

number of grams of solute that can dissolve in 100 grams of solvent at a given temperature

SOLUTION

36.0 g NaCl can be dissolve in 100 g of water at 293 K

GAS PHASE SOLUTION

Saturn

solvent

H2/He

solute

CH4, PH3

LIQUID SOLUTION

Europa

solvent

H2O

solute

MgSO4

SOLID SOLUTION

Triton

solvent

N2

solute

CH4

methanol

sugar

ethanol

water

ionic compounds

(NaCl, KF)

NaOH

HCl

H2SO4

CH3COOH

HCOOH

HF

ELECTROLYTES

migrating negative and positive charges

Kohlrausch NaCl

DISSOCIATION

‘breaking apart’

NaCl (s) → Na+ (aq) + Cl- (aq)

NaOH (s) → Na+ (aq) + OH- (aq)

HCl (g) → H+ (aq) + Cl- (aq)

strong electrolytes are fully dissociated

Ca(NO3)2 (s) → Ca2+(aq) + 2 NO3- (aq)

EXP5

polyatomic ions do NOT dissociate

O

H H

δ-

δ+ δ+

SOLVATION

cations anions

SOLVATION

non-electrolyte

CH3COOH (aq) H+ (aq) + CH3COO- (aq)

weak electrolytes are not fully dissociated

reversible reaction

(chemical equilibrium)

→ ←

NaCl (s) → Na+ (aq) + Cl- (aq)

strong electrolytes are fully dissociated

CHEMICAL REACTIONS

1.properties of aqueous solutions





2.1. PRECIPITATION REACTIONS

solution 1 solution 2 solution 1 + solution 2

2.1. PRECIPITATION REACTIONS

formation of an insoluble product

(precipitate)

NaCl(aq) + AgNO3(aq)

AgCl(s) + NaNO3(aq) EXP 6

insoluble compounds

1.M+ compounds (M = H, Li, Na, K, Rb, Cs, NH4)

2. A- compounds (A = NO3, HCO3, ClO3, Cl, Br, I)(AgX, PbX2)

3. SO42-

(Ag, Ca, Sr, Ba, Hg, Pb)

4. CO32-, PO4

3-, CrO42-, S2-

(Ag, Ca, Sr, Ba, Hg, Pb)

NaCl(aq) + AgNO3(aq) → AgCl(s) + NaNO3(aq)

balanced molecular equation

(table to determine which compound precipitates)

balanced ionic equation

1. NaCl(s) → Na+(aq) + Cl-(aq)

2. AgNO3(s) → Ag+(aq) + NO3-(aq)

3. Na+(aq) + Cl-(aq) + Ag+(aq)+ NO3-(aq) →

AgCl(s) + Na+ (aq) + NO3-(aq)

spectator ions

Ba(NO3)2(aq) + Na3PO4(aq)

1. which compound falls out? 2. balanced molecular equation

3. balanced ionic equations4. identify spectator ions

Cs2CrO4(aq) + Pb(NO3)2(aq)

Ba(NO3)2 (aq) + Na2SO4 (aq)

CHEMICAL REACTIONS






ACIDS AND BASES

Arrhenius (1883)

ACIDS

BASES

NaOH (s) → Na+ (aq) + OH- (aq)

MOH → M+ (aq) + OH- (aq)

HCl (g) → H+ (aq) + Cl- (aq)

HAc → H+ (aq) + Ac- (aq)

ionization

IDENTIFICATION

Litmus Paper

acid

base

red

blue

Säure

Base

EXP7

ACIDS AND BASES

ACIDS BASES

HCl (aq) + NaOH (aq) → NaCl (aq) + H2O

HAc (aq) + MOH (aq) → MAc (aq) + H2O

and

NEUTRALIZE EACH OTHER

acid + base salt + water

H+ ≈ 10-15 m

Na+≈ 10-10 m

ACIDS AND BASES

ACIDS AND BASES

HCl (g) → H+ (aq) + Cl- (aq)

H+(aq) + H2O H3O+(aq)

HCl (g) + H2O → H3O+ (aq) + Cl- (aq)

one stephydronium ion

(aq) (l) (aq) (aq)

hydronium ion

acid base

cation hydronium ion

PROPERTIES OF ACIDS

1. acids have a sour taste

vinegar – acetic acidlemons – citric acid

2. acids react with some metals to form hydrogen

2 HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)

3. acids react with carbonates to water and carbon dioxide

2 HCl(aq) + CaCO3(s) → CaCl2(aq) + [H2CO3]H2CO3 → H2O(l) + CO2(g)

EXP8

EXP9

4. some acids are hygroscopic

H2SO4 (conc)

BASES

1. bases have a bitter taste

2. bases feel slippery

soap

3. aqueous bases and acids conduct electricity

EXAMPLES

KOH(aq) and HF(aq)

Mg(OH)2(aq) and HCl(aq)

Ba(OH)2(aq) and H2SO4(aq)

NaOH(aq) and H3PO4(aq)

(stepwise)

Bronsted (1932)

ACIDS

HAc → H+ (aq) + Ac- (aq)

proton donors

BASES

proton acceptor

B + H+ (aq) → BH+ (aq)

weak electrolyte

CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq)

NH3(aq) + H2O(l) NH4+ + OH-

strong electrolyte

HCl(aq) + H2O(l) → H3O+(aq) + Cl-(aq)

HNO3(aq) + H2O(l) → H3O+(aq) + NO3-(aq)

donor versus acceptor

CH3COOH(aq) + H2O(l) H3O+(aq) + CH3COO-(aq)

NH3(aq) + H2O(l) NH4+(aq)+ OH-(aq)

H2O(l) + H2O(l) H3O+(aq) + OH-(aq)

water can be either an acid or a base

AUTO DISSOCIATION

monoprotic acids

diprotic acid

HF, HCl, HBr, HNO3, CH3COOH

H2SO4 → H+(aq) + HSO4-(aq)

HSO4-(aq) H+(aq) + SO4

2-(aq)

triprotic acid

H3PO4 H+(aq) + H2PO4-(aq)

H2PO4-(aq) H+(aq) + HPO4

2-(aq)

HPO42-(aq) H+(aq) + PO4

3-(aq)

EXP10

CHEMICAL PROPOERTIES

1. Non-metal oxides react with water to form an acid

(acetic anhydrides)

3 2 2 4

2 5 2 3

2 2 2 3

SO ( ) H O H SO ( ) sulfuric acid

N O ( ) H O 2HNO ( ) nitric acid

CO ( ) H O H CO ( ) carbonic acid

g aq

g aq

g aq

Cl2O7, SO2, Br2O5

+ H2O

+ H2O

+ H2O

CHEMICAL PROPERTIES

2. Soluble metal oxides react with water to form a base

(base anhydrides)

MgO, Al2O3

2 2

2 2

CaO( ) H O Ca(OH) ( ) calcium hydroxide

Na O( ) H O 2NaOH( ) sodium hydroxide

s aq

s aq

+ H2O

+ H2O

NAMING ACIDS AND BASES

2 2

HCl( ) hydrogen chloride HCl( ) chlor

H S( ) hydrogen sulfide H S( ) sulfur

g aq hydro ic acid

g aq hydro ic acid

prefix hydro- the suffix –ic to the stem of the nonmetal name followed by the word acid

binary acids

NAMING ACIDS AND BASES

oxo acids acids

contain hydrogen, oxygen, plus another element

main group 5

HNO3 nitric acidHNO2 nitrous acid

H3PO4 phosphoric acidH3PO3 phosphorous acid

H2SO4 sulfuric acidH2SO3 sulfurous acid

main group 6

main group 7

HClO4 perchloric acidHClO3 chloric acidHClO2 chlorous acidHClO hypochlorous acid

Acids in the Solar System

Venus

H2SO4(g)

Europa

H2SO4(s)

Acids in the Interstellar Medium

Orion

NH3, H2O, H2S

CH3COOH

HCOOH

HF, HCl

CHEMICAL REACTIONS






1. oxidation

KEY CONCEPTS

loss of electrons

2. reduction acceptance of electrons

NUMBER OF ELECTRONS MUST BE CONSERVED

1. oxidation

EXAMPLE

2. reduction

!!!balance electrons!!!

Na+Cl-

Na Na+ + e

Cl2 + 2 e 2 Cl-

CaO, Al2O3

substance that lost the electrons reduction agent

substance that gained the electrons oxidizing agent

oxidizing agent is reduced

reducing agent is oxidized

2 Na + Cl2 2 Na+Cl-

EXAMPLE 1

solid state reaction of potassium with sulfur

to form potassium sulfide

EXAMPLE 2

solid state reaction of iron with oxygen

to form iron(III)oxide

OXIDATION NUMBER

ionic compounds ↔ molecular compounds

NaCl HF, H2

Na+Cl- ?electrons are fully transferred covalent bond

charges an atom would have if electrons are transferred completely

HF H+ + F-

molecular compound ionic compound

F- oxidation state -1

H+ oxidation state +1

EXAMPLE 1

H2O


2 H+ + O2-

H+ oxidation state +1

O2- oxidation state -2

EXAMPLE 2

H2


H+ + H-

EXAMPLE 3

OXIDATION NUMBER OF FREE ELEMENTS IS ZERO

RULE 1

OXIDATION NUMBER OF FREE ELEMENTS IS ZERO

H2, O2, F2, Cl2, K, Ca, P4, S8

RULE 2

monoatomic ions

oxidation number equals the charge of the ion

group I M+

group II M2+

group III M3+ (Tl: also +1)

group VII (w/ metal) X-

RULE 3

oxidation number of hydrogen

+1 in most compounds

(H2O, HF, HCl, NH3)

-1 binary compounds with metals (hydrides)

(LiH, NaH, CaH2, AlH3)

RULE 4

oxidation number of oxygen

-2 in most compounds

(H2O, MgO, Al2O3)

-1 in peroxide ion (O22-) (H2O2, K2O2, CaO2)

-1/2 in superoxide ion (O2-) (LiO2)

RULE 5

oxidation numbers of halogens

F: -1 (KF)

Cl, Br, I: -1 (halides) (NaCl, KBr)

Cl, Br, I: positive oxidation numbers if combined with oxygen (ClO4

-)

RULE 6

charges of polyatomic molecules must be integers

(NO3-, SO4

2-)

oxidation numbers do not have to be integers

-1/2 in superoxide ion (O2-)

MENUE

1.oxidation states of group I – III metals

2.oxidation state of hydrogen (+1, -1)

3. oxidation states of oxygen (-2, -1, -1/2, +1)

4.oxidation state of halogens

5.remaining atoms

oxidizing agents

OCl- Cl-?????

EXP10

reducing agent

2 Na + 2 H2O H2 + 2 NaOH

EXP11/12

NONO2

NO+

NO-NO2-

NO3-

PO43- SO4

2-

SO3

SO2

KO2

K2O

BrO-KClO4

1.redox reactions

2. oxidation versus reduction

3. oxidation numbers versus charges

4. calculation of oxidation numbers

REVISION

TYPES OF REDOX REACTIONS

1.combination reactions

A + B → C

2. decomposition reactions

C → A + B

3. displacement reactions

A + BC → AC + B

4. disproportionation reactions


A + B → C

two or more compounds combine to form a single product

S8(s) + O2(g) → SO2(g)

1. oxidation numbers

2. balancing charges

MENUE

1.oxidation states of group I – III metals

2.oxidation state of hydrogen (+1, -1)

3. oxidation states of oxygen (-2, -1, -1/2, +1)

4.oxidation state of halogens

5.remaining atoms


C → A + B

breakdown of one compound into two or more compounds

HgO(s) → Hg(l) + O2(g)

1. oxidation numbers

2. balancing charges

KClO3(s) → KCl(s) + O2(g)


A + BC → AC + B

an ion or atom in a compound is replaced by an ion or atom of another element

3.1. Hydrogen displacement

3.2. Metal displacement

3.3. Halogen displacement


group I and some group II metals (Ca, Sr, Ba)

react with water to form hydrogen

Na(s) + H2O(l) → NaOH + H2(g)

less reactive metals form hydrogen and the oxide in water (group III, transition metals)

Al(s) + H2O(l) → Al2O3(s) + H2(g)


even less reactive metals form hydrogen in acids

Zn(s) + HCl(aq) → ZnCl2(aq) + H2(g)

EXP12

Li K Ba Ca Na Mg Al Zn Cr Fe Cd Co Ni Sn Pb H Cu Hg Ag Pt Au

activity series of metals

displace H from water

displace H from steam

displace H from acids

Li K Ba Ca Na Mg Al Zn Cr Fe Cd Co Ni Sn Pb H Cu Hg Ag Pt Au

likes to donate electrons does not like so much to donate electrons

EXP13

3.2. Metal displacement

V2O5(s) + 5 Ca(s) → 2 V(s) + 5 CaO(s)

TiCl4(g) + 2 Mg (l) → Ti(s) + 2 MgCl2(l)

3.3. Halogen displacement

F2 > Cl2 > Br2 > I2

reactivity (‘likes’ electrons)

Cl2(g) + 2 KBr(aq) → 2 KCl(aq) + Br2(l)0 0+1+1 -1 -1

Br2(g) + 2 KI(aq) → 2 KBr(aq) + I2(s)


an element in one oxidation state is oxidized and reduced

at the same time

H2O2(aq) → 2 H2O(l) + O2(g)

Cl2(g) + 2 OH-(aq) → ClO-(aq) + Cl-(aq) + H2O(l)

SUMMARY


A + B → C


C → A + B


A + BC → AC + B


STOCHIOMETRY(CONCENTRATION)

molar concentration

Molarity

(M)

solution of literssolute of moles (M)molarity =

How many grams of AgNO3 are needed to prepare250 mL of 0.0125 M AgNO3 solution?

30.531 g AgNO

How many mL of 0.124 M NaOH are required

to react completely with 15.4 mL of 0.108 M H2SO4?

26.8 mL NaOH

2 NaOH + H2SO4 Na2SO4 + 2H2O

How many mL of 0.124 M NaOH are required

to react completely with 20.1 mL of 0.2 M HCl?

NaOH + HCl NaCl + H2O

How many grams of iron(II)sulfide have to react with hydrochloric acid to generate 12 g of hydrogen sulfide?

How many moles of BaSO4 will form if 20.0 mL of

0.600 M BaCl2 is mixed with 30.0 mL of 0.500 M MgSO4?

BaCl2 + MgSO4 BaSO4 + MgCl2

This is a limiting reagent problem

40.0120 mol BaSO

How many ml of a 1.5 M HCl will be used to neutralize

a 0.2 M Ba(OH)2 solution?

How many ml of a 1.5 M HCl will be used to prepare

500 ml of a 0.1 M HCl?

dil dil concd concdV M V MX X=

LIMITING REACTANT

C2H4 + H2O C2H5OHEXP14

excess reactantlimiting reactant

How many grams of NO can form when 30.0 g NH3 and 40.0 g O2 react according to

4 NH3 + 5 O2 4 NO + 6 H2O

CHEMISTRY 161 Chapter 4. CHEMICAL REACTIONS 2 HgO (s) → 2Hg (l) + O 2(g) aq 1. properties of aqueous solutions 2. reactions in aqueous solutions a) precipitation.

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