BASIC CONCEPTS of CHEMISTRY Please revise: names of element, Periodic Table (understanding the notation), acid, base and salt nomenclature

BASIC CONCEPTS of CHEMISTRY

Please revise: names of element, Periodic Table (understanding the notation), acid, base and salt nomenclature

1. Definitions of acid and base

There are several methods of defining acids and bases. While these definitions don't contradict each other, they do vary in how inclusive they are. Antoine Lavoisier, Humphry Davy, and Justus Liebig also made observations regarding acids and bases, but didn't formalize definitions.

Svante Arrhenius

 acids produce H ⁺ ions in aqueous solutions , e.g.

HBr → H ⁺ + Br ^– HNO 2 → H ⁺ + NO 2 -

H ₂ SO ₄ → H ⁺ + HSO ₄ ^- HSO 4 - ↔ H ⁺ + SO 4 2-

 bases produce OH ^- ions in aqueous solutions, e.g.

KOH → K ⁺ + OH ^-

Mg(OH) ₂ → Mg(OH) ⁺ + OH ^- Mg(OH) ⁺ ↔ Mg ²⁺ + OH ^- Zn(OH) 2 ↓ ↔ Zn(OH) ⁺ + OH ^- Zn(OH) ⁺ ↔ Zn ²⁺ + OH ^-

o water required, so only allows for aqueous solutions

o only protic acids are allowed; required to produce hydrogen ions

o only hydroxide bases are allowed

Johannes Nicolaus Brønsted - Thomas Martin Lowry

 acids are proton donors

 bases are proton acceptors Examples:

acid ₁ base 2 base ₁ acid ₂ HBr + H ₂ O ↔ Br ^- + H ₃ O ⁺ HCl + NH 3 ↔ Cl ^- + NH 4 +

H 2 O + NH 3 ↔ OH ^- + NH 4 +

NH ₄ ⁺ + H ₂ O ↔ NH 3 + H ₃ O ⁺ H 2 O + CN ^- ↔ OH ^- + HCN

o queous solutions are permissible

o bases besides hydroxides are permissible

o only protic acids are allowed

Gilbert Newton Lewis (last restrictive of acid-base definition)

 acids are electron pair acceptors

 bases are electron pair donors Examples:

acid base

H 3 O ⁺ + NH 3 ↔ H 2 O + NH 4 +

AlCl 3 + Cl ^- ↔ AlCl 4 -

Al ³⁺ + H ₂ O ↔ Al(H 2 O) ₆ ³⁺

BCl 3 + NH 3 ↔ BCl 3 NH 3

Properties of Acids

 taste sour (don't taste them!)... the word 'acid' comes from the Latin acere, which means 'sour'

 acids change litmus (a blue vegetable dye) from blue to red

 their aqueous (water) solutions conduct electric current (are electrolytes)

 react with bases to form salts and water

 evolve hydrogen gas (H 2 ) upon reaction with an active metal (such as alkali metals, alkaline earth metals, zinc, aluminum)

Properties of Bases

 taste bitter (don't taste them!)

 feel slippery or soapy (don't arbitrarily touch them!)

 bases don't change the color of litmus; they can turn red (acidified) litmus back to blue

 their aqueous (water) solutions conduct and electric current (are electrolytes)

 react with acids to form salts and water Question:

Which of them are acids or bases: NH 4 +

, OH ^- , CN ^- , ClO 4 -

, HCl, HCN, NH 3 ·H 2 O , BF 3 , Cl ^- , Fe ³⁺ , SO ₃ , H ₂ O, H ₃ O ⁺ , NH ₄ ⁺ , NH ₂ ^- , FeCl ₃ , Ag ⁺ , H ₂ PO ₄ ^- , S ^2- , Cu ²⁺ , F ^- , NH ₃ , CoCl ₂ ?

2. Formulas of common acids and bases

Here are the names and formulas of some of the common acids and bases.

 Binary Acids:

A binary compound consists of two elements. Binary acids have the prefix hydro in front of the full name of the nonmetallic element. They have the ending -ic. Examples include hydrochloric and hydrofluoric acid.

Hydrofluoric Acid - HF Hydrochloric Acid - HCl Hydrobromic Acid - HBr Hydroiodic Acid - HI Hydrosulfuric Acid - H 2 S

 Ternary Acids:

Ternary acids commonly contain hydrogen, a nonmetal, and oxygen. The name of the most common form of the acid consists of the nonmetal root name with the -ic ending, The acid containing one less oxygen atom than the most common form is designated by the -ous ending. An acid containing one less oxygen atom than the -ous acid has the prefix hypo- and the -ous ending. The acid containing one more oxygen than the most common acid has the per- prefix and the -ic ending.

Nitric Acid - HNO 3

Nitrous Acid - HNO 2

Hypochlorous Acid - HClO Chlorous Acid - HClO ₂ Chloric Acid - HClO 3

Perchloric Acid - HClO 4

Sulfuric Acid - H ₂ SO ₄ Sulfurous Acid - H 2 SO 3

Phosphoric Acid - H 3 PO 4

Phosphorous Acid - H ₃ PO ₃ Carbonic Acid - H 2 CO 3

Acetic Acid - HC 2 H 3 O 2

Oxalic Acid - H ₂ C ₂ O ₄ Boric Acid - H 3 BO 3

Silicic Acid - H 2 SiO 3

 Bases:

Sodium Hydroxide - NaOH Potassium Hydroxide - KOH

Ammonium Hydroxide - NH ₃ ·H ₂ O (NH ₄ OH) Calcium Hydroxide - Ca(OH) 2

Magnesium Hydroxide - Mg(OH) 2

Barium Hydroxide - Ba(OH) ₂ Aluminum Hydroxide - Al(OH) 3

Ferrous Hydroxide or Iron (II) Hydroxide - Fe(OH) 2

Ferric Hydroxide or Iron (III) Hydroxide - Fe(OH) ₃ Zinc Hydroxide - Zn(OH) ₂

Lithium Hydroxide – LiOH

3. Strength of acids and bases

Strong electrolytes are completely dissociated into ions in water. The acid or base molecule does not exist in aqueous solution, only ions. Weak electrolytes are incompletely dissociated.

Strong acids

Strong acids completely dissociate in water, forming H+ and an anion, e.g. HCl + H 2 O → H ₃ O ⁺ + Cl ^-

Since it is a strong acid it dissociates completely, you will not have any HCl. In solution we have H 3 O ⁺ and Cl ^- ions which concentration is equal HCl concentration.

There are six strong acids. The others are considered to be weak acids. You should commit the strong acids to memory:

 HCl - hydrochloric acid

 HNO ₃ - nitric acid

 H ₂ SO ₄ - sulfuric acid

 HBr - hydrobromic acid

 HI - hydroiodic acid

 HClO ₄ - perchloric acid

100% dissociation isn't true as solutions become more concentrated. If the acid is 100%

dissociated in solutions of 1.0 M or less, it is called strong. Sulfuric acid is considered strong only in its first dissociation step.

H 2 SO 4 → H ⁺ + HSO 4 -

What Is the Strongest Acid?

None of the strong acids traditionally listed in a chemistry text holds the title of World's Strongest Acid. The record-holder used to be fluorosulfuric acid (HFSO 3 ), but the carborane superacids H(CHB11Cl11) are hundreds of times stronger than fluorosulfuric acid and over a million times stronger than concentrated sulfuric acid. The superacids readily release protons, which is a slightly different criterion for acid strength than the ability to dissociate to release a H ⁺ ion (a proton). The carborane acids are incredible proton donors, yet they are not highly corrosive. Corrosiveness is related to the negatively-charged part of the acid. Hydrofluoric acid (HF), for example, is so corrosve it dissolves glass. The fluoride ion attacks the silicon atom in silica glass while the proton is interacting with oxygen. Even though it is highly corrosive, hydrofluoric acid is not considered to be a strong acid because it does not completely dissociate in water.

Weak acids

A weak acid only partially dissociates in water to give H ⁺ and the anion. Examples of weak acids include hydrofluoric acid, HF, and acetic acid, CH 3 COOH.

e.g. HF + H 2 O ↔ H 3 O ⁺ + F ^-

In solution molecules HF and H ₃ O ⁺ and F ^- ions are in equilibrium with each other. We have to use the equilibrium expression to determine what concentration each species is at when the system is at equilibrium. The equilibrium expression is following:

Weak acids include:

 Molecules that contain an ionizable proton. A molecule with a formula starting with H usually is an acid.

 Organic acids containing one or more carboxyl group, -COOH. The H is ionizable.

 Anions with an ionizable proton. (e.g., HSO 4 -

→ H ⁺ + SO 4 2-

)

 Cations

o transition metal cations

o heavy metal cations with high charge

o

NH 4 +

dissociates into NH 3 + H ⁺ Strong bases

Strong bases dissociate 100% into the cation and OH ^- (hydroxide ion). The hydroxides of the Group I and Group II metals usually are considered to be strong bases.

LiOH - lithium hydroxide NaOH - sodium hydroxide KOH - potassium hydroxide RbOH - rubidium hydroxide CsOH - cesium hydroxide *Ca(OH) ₂ - calcium hydroxide *Sr(OH) 2 - strontium hydroxide *Ba(OH) 2 - barium hydroxide

* These bases completely dissociate in solutions of 0.01 M or less. The other bases make

solutions of 1.0 M and are 100% dissociated at that concentration. There are other strong

bases than those listed, but they are not often encountered.

Weak Bases

Examples of weak bases include ammonia, NH 3 ·H 2 O, and diethylamine, (CH 3 CH 2 ) 2 NH

 Most weak bases are anions of weak acids.

 Weak bases do not furnish OH ^- ions by dissociation. Instead, they react with water to generate OH ^- ions..

 We use the equilibrium expression to determine what concentration each species are in solution (example is above)

3. Salt formation

When acids and bases react with each other, they can form a salt and (usually) water. This is called a neutralization reaction and takes the following form:

HA + BOH → BA + H 2 O

Depending on the solubility of the salt, it may remain in ionized form in the solution or it may precipitate out of solution. Neutralization reactions usually proceed to completion.

The reverse of the neutralization reaction is called hydrolysis. In a hydrolysis reaction a salt reacts with water to yield the acid or base:

BA + H ₂ O → HA + BOH

More specifically, there are four combinations of strong and weak acids and bases:

a) strong acid + strong base, e.g. HCl + NaOH → NaCl + H ₂ O

When strong acids and strong bases react, the products are salt and water. The acid and base neutralize each other, so the solution will be neutral (pH=7; neutral reaction) and the ions that are formed will not reaction with the water.

Hydrolysis of NaCl:

NaCl + H 2 O → HCl + NaOH

Na ⁺ + Cl ^- + H ₂ O→ H ⁺ + Cl ^- + Na ⁺ + OH ^- H 2 O→ H ⁺ + OH ^-

b) strong acid + weak base, e.g. HCl + NH ₃ ·H 2 O → NH 4 Cl + H ₂ O

The reaction between a strong acid and a weak base also produces a salt, but water is not usually formed because weak bases tend not to be hydroxides. In this case, the water solvent will react with the cation of the salt to reform the weak base (pH of the solution is below 7, acid reaction). For example:

NH 4 Cl + H 2 O ↔ HCl + NH 3 ·H 2 O

NH ₄ ⁺ + Cl ^- + H ₂ O ↔ H ⁺ + Cl ^- + NH ₃ ·H ₂ O NH 4 +

+ H 2 O ↔ H ⁺ + NH 3 ·H 2 O

c) weak acid + strong base, e.g., HClO + NaOH → NaClO + H 2 O

When a weak acid reacts with a strong base the resulting solution will be basic. The salt will be hydrolyzed to form the acid, together with the formation of the hydroxide ion from the hydrolyzed water molecules (pH of the solution is under 7, alkaline reaction).

NaClO + H 2 O ↔ HClO + NaOH

Na ⁺ + ClO ^- + H 2 O ↔ HClO + Na ⁺ + OH ^-

ClO ^- + H ₂ O ↔ HClO + OH ^-

d) weak acid + weak base, e.g., HClO + NH 3 ·H 2 O ↔ NH 4 ClO

The pH of the solution formed from the reaction of a weak acid with a weak base depends on the relative strengths of the reactants. For example, if the acid HClO has a K a of 3.4 x 10 ^-8 and the base NH ₃ ·H 2 O has a K _b = 1.6 x 10 ^-5 , then the aqueous solution of HClO and NH ₃ ·H 2 O will be basic because the K a of HClO is less than the K b of NH 3 ·H 2 O.

How to write salt formulas?

First you must know the formulas of common acids and bases!

1) When a binary acid reacts with a base it forms a halide salt plus water. The salt contains the halogen of the acid and the cation of the base. Hydrochloric acid reacts with bases to form chloride salts. The specific cation will depend upon what base was used to begin with. For example:

HCl + NaOH → NaCl + H 2 O

 hydrochloric acid reacts with a base to produce chloride salts

 hydrobromic acid reacts with bases to produce bromide salts

 hydroiodic acid reacts with bases to produce iodide salts

 hydroflouric acid reacts with bases to produce flouride salts

Notice that the ending on the salt ends in "ide" whereas the ending of the acid is "ic"

 hydrobromic acid reacts with calcium hydroxide to form water and calcium bromide salt

 hydroiodic acid reacts with lithium hydroxide to form water and lithium iodide salt

 hydrochloric acid reacts with aluminum hydroxide to form water and aluminum chloride salt

Let's do some practice. Please answer the following questions:

1. What is the name of HCl?

2. If HCl were reacted with calcium hydroxide what would be the name of the salt produced? What would be its formula?

3. What is the name of HBr?

4. If HBr were reacted with sodium hydroxide what would be the name of the salt produced? What would be its formula?

2) When a ternary acid reacts with a base it forms a ternary salt plus water.

 HClO (hypochlorous acid) reacts with magnesium hydroxide base to form water and magnesium hypochlorite

HClO + Mg(OH) 2 → Mg(ClO) 2 + 2H 2 O

 HClO ₂ (chlorous acid) reacts with bases to form water and chlorite salts. For example, when chlorous acid reacts with sodium hydroxide base it forms water and sodium chlorite.

HClO ₂ + NaOH → NaClO ₂ + H ₂ O

 HClO 3 (chloric acid) reacts with a base to form water and chlorate salts. For example, chloric acid will react with lithium hydroxide base to form water and lithium chlorate.

HClO ₃ + LiOH → LiClO ₃ + H ₂ O

 HClO 4 (perchloric acid) reacts with a base to produce water and perchlorate salts. For example, perchloric acid will react with berylium hydroxide to form water and berylium perchlorate

HClO 4 + Be(OH) 2 → Be(ClO 4 ) 2 + 2H 2 O

We can form other series of oxy-halogen acids by using a bromine or an iodine instead of a chlorine. The same pattern of prefixes and endings will prevail however. HBrO is called hypobromous acid, HBrO 2 is called bromous acid, etc. What would HBrO 3 be called? HBrO 4 ? How about HIO? HIO ₂ ?

The salts that form when these acids react with a base also follow the same pattern.

Hypohalous acids form hypohalite salts, halous acids form halite salts, halic acids form halate salts, and perhalic acids form perhalate salts. Just replace the hal letters with the respective halogen letters.

 HNO ₂ (nitrous acid) reacts with a base to form water and nitrite salts. For example nitrous acid will react with magnesium hydroxide to form water and magnesium nitrite.

HNO ₂ + Mg(OH) ₂ → Mg(NO ₂ ) ₂ + 2H ₂ O

 HNO 3 (nitric acid) reacts with bases to form nitrate salts. So for example, when nitric acid reacts with sodium hydroxide then it forms water and sodium nitrate.

HNO 3 + NaOH → NaNO 3 + H 2 O

Another observation should be noted here. Salts of weak inorganic acids always end in "ite"

and salts of strong inorganic acids always end in "ate". Check it out above.

 H 2 SO 3 (sulfurous acid) reacts with a base to form sulfite salts. For example sulfurous acid will react with potassium hydroxide to form water and potassium sulfite:

H 2 SO 3 + 2KOH → K 2 SO 3 + 2H 2 O

 H ₂ SO ₄ (sulfuric acid) reacts with a base and produce water and the sulfate salt. For example:

H 2 SO 4 + 2LiOH → Li 2 SO 4 + 2H 2 O

 H ₂ CO ₃ (carbonic acid) reacts with a base to produce salt plus a carbonate salt. For example, carbonic acid will react with calcium hydroxide to form water and calcium carbonate.

H ₂ CO ₃ + Ca(OH) ₂ → CaCO ₃ + 2H ₂ O

Organic acids do not follow the rule that says that weak acids should end in "ous". All organic carboxylic acids are weak including carbonic acid yet they all end in "ic", but the rule was made for non-carbon containing acids. Carbonic acid like sulfurous acid is unstable and decomposes to produce water and carbon dioxide gas.

H 2 CO 3 → H 2 O + CO 2

 HC 2 H 3 O 2 (acetic acid) another carbon containing acids that are weak yet end in "ic".

acetic acid reacts with bases to form water and acetate salts. For example, acetic acid will react with potassium hydroxide to form water and potassium acetate

HC 2 H 3 O 2 + KOH → KC 2 H 3 O 2 + H 2 O