PREDICTING REACTION PRODUCTS AND BALANCING CHEMICAL EQUATIONSThe most common issue among students new to chemistry is writing equations; but more precisely, predicting the products of a reaction and then balancing equations.  There are a few key pieces of information which help this process which I am presenting here today.

The basicsPredicting chemical reactions starts first with knowing the basics of atoms, molecules, ions and compounds.  Get familiar with elements in the periodic table - their names, chemical symbols, their states at room temperature, and their potential charges as ions.

To get you started:

Most metals can exist in their elemental form, with an oxidation state of zero, in solid state, e.g. Al(s), Fe(s), Ag(s), etc.  The notable exception is Mercury, which exists at room temperature as a liquid, Hg(l) (some thermometers contain mercury to tell the temperature). Noble gases (Group VIII) exist as gases, but are generally not involved in chemical reactions, unless you are considering ionization energies or similar concepts. Non-metals which are not noble gases generally bond covalently with other non-metals to form molecules.  Small molecules like N2(g), O2(g), CO2(g) and NO(g) all exist as gases at room temperature.  As we observe the states of elements going down Group VII (halogens), we observe gas to liquid to solid: F2(g), Cl2(g), Br2(l), I2(s). Carbon reacts with hydrogen, oxygen, nitrogen, phosphorus, sulfur and some other elements to form organic molecules. I highly recommend you revise compounds of carbon if you are a little rusty.  The below tables should help you out:

Last, but definitely not least, memorise your common cations and anions - their elemental composition AND charge, which are shown in the below table. Know how to write chemical formulasfrom cations and anions by balancing the charges of cations and anions into a neutral compound.  I cannot stress this point enough - it is usually a large sticking point for most students and one you can't get around.  Spend the time memorising the below table and you will save yourself a lot of stress.  There are a few tips and tricks to memorising this table though, take note: The top four rows (plus sulfide) all pertain to main group elements which have a charge relevant to their position on the periodic table - Group II elements have a +2 charge; Group VII elements have a -1 charge, etc. Transition elements often have more than one oxidation state, however only the most common ones have been listed on the table.  Note that copper, iron and tin each have two common ions so be sure you know which one you are dealing with. Some of the polyatomic anions (sulfide, sulfite, sulfate, carbonate, phosphate) also have associated hydrogenated compounds.  If you can memorise the non-hydrogenated forms and their charges, then you can work out the chemical formula for every hydrogen added - add one hydrogen and add +1 to the charge of the compound.

Types of reactionsIt is super important to first identify the type of reaction you are looking at - this will give you vital cues as to the products of the reaction in question.  I have summarised the types of reactions pertinent to VCE below, excluding a couple related to the reactions of alkenes which is covered later in Year 12.

Combustion reactions:These reactions are essentially burning a compound - often a carbon/hydrogen/oxygen based compound or sometimes solid metal - in air. What does a fire need to burn?  Fuel, a source of ignition, and of course OXYGEN GAS.  Therefore, a chemical reaction representing the combustion of a compound is really a reaction with oxygen gas:

The reaction of a carbon compound with oxygen gas will ALWAYS produce carbon dioxide and water as products.   The reaction of a metal with oxygen is actually a redox reaction involving the oxidation of the metal, and will always produce a metal oxide.

Precipitation reactions:The formation of a precipitate in plain English is the formation of a solid compound in solution, and can be used as a quantitative determination which is termed gravimetric analysis.  Gravimetric analysis is useful because the solid formed as a result of a reaction can be separated from the liquid by filtration, and then can be dried and weighed.  Through stoichiometry you can determine the amount/concentration of the reactant(s).

To form a precipitate, two salts undergo a metathesis reaction - the swapping of cations/anions - which produces a salt which is insoluble in the solvent. It is

important to note that the formation of a solid is related to its solubility in that particular solvent.

There are some great solubility tables published in countless textbooks and on the internet - not only are there plenty of these about, but I believe they're not needed. At least for the purposes of VCE and basic chemistry courses, the products of a precipitation reaction are not going to be ambiguous.  The vital information you need to know is:All salts of Li+, Na+, K+, NH4+, CH3COO- and NO3- are ALWAYS going to be soluble in water. If you KNOW the reaction is definitely a precipitation reaction, then the other salt MUST be the solid.As you practice, you will come to know the common insoluble salts. e.g. PbI2, BaSO4, AgCl, etc.

Acid-base reactions:There are several different types of acid-base reactions, which can be summarised in the below graphic.

The most common type of acid-base reaction encountered in VCE chemistry is the one in light blue/aqua.  In acid-base titrations, usually it is not the amount of product measured, but the amount of one reactant to bring the reaction to the equivalence point - the point at which you have stoichiometrically equal amounts of the acid and the base.  This is determined by observation of the colour of an indicator; the colour change is termed the end point.  An indicator is chosen so that the equivalence point and end point are very close.Sometimes reactions involving carbonates or hydrogen carbonates are used quantitatively by measuring the loss of CO2 gas from the reaction system - the mass of CO2 lost can be weighed on a set of scales.

Dissolution reactions:Dissolution reactions are simply those where a solute is dissolving in a solvent to form a solution - or said another way, where compounds are breaking apart into their constituent anions and cations in the presence of a solvent, usually water:

They can also be used to depict an acid dissociating in water.  For a monoprotic

acid, this is a simple task:

For a polyprotic acid, there are several steps involved, where each reaction involves only the loss of one  hydrogen for the acid.  Note that the first dissociation reaction for a triprotic acid is shown as a strong acid and therefore complete dissociation is observed, denoted by the "normal" arrow pointing to the right. Whereas in the second and third reactions, dihydrogen species and hydrogen species are weak acids and thus incomplete dissociation is observed, denoted by the equilibrium arrows.

Balancing equationsThere are no hard and fast rules for balancing equations - everyone has their own way which works for them.  However, there are a few tips and tricks:

Keep a tally of all the elements on each side of the equation; change your tally when you add a stoichiometric coefficient (the number in front of each compound). I also recommend that if a polyatomic anion/cation does not break apart during the reaction, count it as one entity, rather than its constituent elemente.g. treat phosphate as one entity with a 3- charge rather than one phosphorus and four oxygens Start with elements other than carbon, hydrogen and oxygen (if applicable). Then proceed with C, H and O in the most appropriate order When balancing combustion reactions, I recommend you balance carbon first, then hydrogen, and finally oxygen. If there is one element by itself, e.g. O2, leave this one to last. Don't forget to add states at the end!

The most important factor in consolidating your understanding of predicting reaction products and balancing equations is PRACTICE

Trends in the Periodic Table

Periodic table is a devise that gives us the information of the elements in a systematic manner. The

periodic table consists of all known elements arranged in rows and columns in increasing

order of their atomic numbers. 

1. The modern or long form of Periodic table is based on the concept of atomic numbers and the

elements are arranged in different groups owing to their different physical properties and

chemical properties.

2. In this arrangement the elements that show similarities in their properties fall in the same

group.

3. Due to the difference in their electronic configurations however they are placed in different

periods.

4. Each period designates one one increase in the principle quantum number

5. There are different blocks in the periodic table depending on the electronic configuration.

6. 's' block elements are those in which the outermost electron/s is/are in the 's' orbital.

7. These have their own characteristic properties. Similarly 'p'. 'd'. and 'f' block elements are

separated in to respective blocks.

8. The set of properties that are unique for these elements are taken in to account

9. Zero group or group 18 elements are inert or noble gases. The elements in this group have

their outermost orbitals completely filled.

10. These are either non-reactive or feebly reactive. 's' block elements are called alkali and

alkaline earth metals. 'd' block elements are transition metals and 'f' block elements are inner

transition elements.

Some properties of elements both physical properties and chemical properties are periodic in nature.

This is greatly due to their electronic configurations. Since the modern periodic table is designed to

address the electronic configuration also this periodicity is maintained.

 

General Trends in the Periodic Table

This arrangement of elements in the modern periodic table is based on the modern periodic law,

which states that "The physical and chemical properties of elements are the periodic functions of their

atomic numbers", which, in simple words, means that when elements are arranged in increasing order

of their atomic numbers, we observe the repetition of certain properties of elements after regular

intervals in the periodic table.

These properties are called the Periodic properties. The periodic properties also show trends in the

periodic table, that is, they show gradation through out rows and columns.

1. Effective nuclear charge

The other electrons in same orbit the lower orbits screen the every electron to some degree so

that the effective nuclear charge on an electron is lesser than expected. This is

called screening effect. This gives an idea of the tendency of the elements towards the reactions.

The effective nuclear charge increases from left to right in a period and it increases slightly or remains

the same in a group from top to bottom.

2. Atomic size or radius

Covalent radius is generally taken as the atomic radius and the atomic radius shrinks with increasing

number of electrons in the outermost orbit till it reaches the maximum capacity of the orbit. Thus in a

period the atomic radius reduces from left to right and with increase in the orbit in a group from top to

down it increases. 

Inert gases, with completed outer orbit have larger radius than the neighboring halogen group atomic

radius. 

In d-block elements because the filling takes place in the n-1 level there is negligible change in the

atomic radius. 

3. Solvation energy or hydration energy

The energy changes involved in the salvation of substance in solvent is known as its salvation energy.

Larger the atomic size lesser the hydration energy or the energy decreases.

4. Melting points and Boiling points

In general melting points decrease down periodic groups containing non transition elements. In

transition elements the melting point usually increases from top to bottom but with a lot of exceptions.

Reactivity Trends in the Periodic TableAll the elements will not participate in the reactions. Noble gases seldom undergo reaction under

normal conditions. They are thus described as inert elements. Even the elements that undergo

reactions the types of reactions they undergo and the products formed are different. 

Generally the reactions occur between the electropositive metallic elements and electronegative non

metals or their compound ions. Basic reactions that are taken in to consideration are the formation of

hydrides and oxides. Sometimes the formation of halides is also considered in determining the

trends. 

1. Formation of oxides

Group 1 and group 2 elements form oxides which are having a general formula M2O and MO

respectively.

These are basic oxides and their hydroxides are alkaline. The elements in these groups tend

to form peroxides also.

Group 13 elements form oxides which are amphoteric in nature with general formula M2O3.

Group 14 elements form covalent oxides. These are mostly neutral and they form weak acids

when dissolved in water.

Group 15 elements form various oxides in different oxidation states.

Group 17 forms oxides which are acidic in nature.

Generally the reactivity of oxidation reduces as the size of the atom increases.

2. Formation of Hydrides

Group 17 elements form hydrides which are strongly acidic in nature.

Group 16 elements form hydrides which are neutral.

Group 15 hydrides are anhydrous compounds. Their hydroxides are basic in nature. Hydrides

of group 14 are covalent hydrides.

Group 13 hydrides are amphoteric in nature and Group 1 and 2 hydrides are acidic in nature.

The vigour of reaction reduces with the increase in atomic size.

Ionization Energy Trend on Periodic TableBack to Top

Ionization energy is the minimum energy required to remove the most loosely held electron

from the ground state of I mole of the isolated atom, molecule or ion in the gaseous state.

1. Successive removals of further ions are termed as 2nd, 3rd ionization energies.

2. The value of ionization energy depends on a number of factors.

3. Ionization energy decreases with the increase in the size of the atom.

4. Ionization energy is greater if the nuclear charge or effective nuclear charge.

5. More energy is required to remove the electron from the half filled or completely filled

shell and hence the ionization energy also is great.

In the periodic table Ionization energy decreases from top to bottom in a group and it will increase

from left to right in a period. However the trend is broken when s- orbital is completely filled (2nd

Group) or half filled p orbital (15th group).

Electronegativity Trends in the Periodic TableElectronegativity is defined as the power of an atom in a molecule to attract bonded or shared pair of

electrons to itself. It depends on both the ionization energy and the electron affinity. There are three

scales that gave values for the elements with regards to their electronegativities and they are

1. Paulings scale

2. Mullikan scale and

3. Alred Rachow scale

Although these three values are almost matching the general scale used is that of Paulings. 

According to this Hydrogen is given a value of 1.0 and Fluorine the maximum of 4.0 and the

electronegativity of other elements were calculated accordingly. In the periodic table electronegativity

increases across a particular period with increasing atomic number or from left to right. The

electronegativity decreases in a group from top to down.

Properties of Trends in Periodic Table1. Atomic size

"Distance from the center of an atom to its outermost shell."

In the periodic table, on moving from left to right in a row, the atomic size of elements decreases

because the number of electron shells in each atom remains the same in a row whereas the increase

in atomic number across the row causes an increase in the attractive forces between the valence

electrons and the positive protons in nucleus. This force of attraction pulls the valence shell closer to

the nucleus in atoms having greater atomic number.

The atomic size of elements increases as we move down a column in the periodic table. This is due to

the increase of the number of electron shells with each successive row.

2. Metallic character

"Tendency of an atom to form positive ions by losing electrons is called its metallic character."

The more readily an atom loses electrons, the more metallic it is.

The metallic character of elements decreases as we move across a row from left to right in the

periodic table and increases as we move down a column in the periodic table. This is because on

moving across a row in the periodic table, the number of valence electrons increases which causes

elements to gain electrons rather than losing them. When we move down a column in the periodic

table, the metallic character increases because the number of valence electrons remains same and

the atomic size increases which makes the valence electrons more far away from the nuclear

attraction and thus easier to lose.

3. Ionization energy (Ionization Potential)

"The energy required for valence electrons to be lost by an isolated gaseous atom thus forming a

positively charged ion is called ionization energy for that element."

The ionization energy increases as we move across a row in the periodic table from left to right, since

it is difficult and a more energy consuming process to remove valence electrons from nonmetallic

elements than that for metallic elements, which are always ready to lose electrons.

The ionization energy decreases as we move down a column in the periodic table, because the lower

elements are able to lose electrons more easily than the upper elements in a group of the periodic

table.

4. Electron affinity

"It is the amount of energy gained by an atom when it gains one electron to form a negatively charged

ion."

Thus, non metallic compounds generally have larger electron affinities since they need electrons more

than metallic elements.On moving across a period from left to right, the electron affinity increases and

on moving down a group in the periodic table, the electron affinity decreases.

5. Electronegativity

"It is the tendency of an atom to gain an electron pair and thus form a chemical bond."

Thus, electronegativity of non metallic elements is more than that of metals.

Consequently, electronegativity increases across a period from left to right and decreases down a

group in the periodic table.

Pictorial Form of Trends in the Periodic Table

  

Electronegativity Trends

The long form of periodic table carries elements according to the increasing order of their

atomic numbers. This in turn, is actually the increasing number of outermost electrons. When we

look at the arrangements though, there are certain properties of these elements which are also

periodic in nature. 

These are called periodic properties and are discussed and understood before going over to chemical

bonding or the nature of compounds formed by any group of elements.

 

Define ElectronegativityConsider the formation of a covalent bond between two similar atoms of a molecule like Hydrogen. In

this molecule the electron pair participating in the formation of covalent bond is shared equally by both

hydrogen atoms, that is, the electron pair lies exactly in the center of the molecule. 

On the other hand, consider the formation of a covalent bond between two dissimilar atoms of a

molecule like HCl. In this molecule the electron pair participating in the formation of covalent bond is

not shared equally by the two atoms, the hydrogen and chlorine. The electron pair tends to lie nearer

the Cl atom than the Hydrogen atom. 

The reason for this unequal sharing of electron pair is given by the fact that Cl atom has a greater

tendency than the hydrogen atom to attract the shared electron pair between them towards itself.

Thus, we can also say that the Cl atom has more electronegativity than the Hydrogen atom.

From the above fact, electronegativity can be defined as

" The electronegativity of a bonded atom is defined as its relative tendency (or ability) to attract the

shared electron pair towards itself."

Electronegativity of an atom, A is generally expressed as χA.

The concept of electronegativity was proposed by the American chemist Linus Pauling in 1932.

Electronegativity is a relative quantity. It is basically written in comparing two atoms involved in the

covalent bond formation. It does not have any unit.

Some of the scales suggested for measuring the electronegativity are

1. Pauling's bond energy scale

Where bond energies or the energy required to break a bond, to get neutral atoms, is employed to

measure electronegativity of an element.

2. Mullikan's scale

According to this scale, electronegativity is taken as a mean difference between ionization potential of

an element and its electron affinity.

Factors Affecting ElectronegativityThe electronegativity of an element is influenced by following factors.

1. Nuclear charge

The higher the nuclear charge, more will be the electronegativity value of an element, since the

nucleus will be able to attract or pull more electrons towards itself.

2. Atomic size

Size of the atom is inversely proportional to electronegativity value. The smaller the element, higher is

its nucleus's reach towards the outer shell. Thus, more will be the electronegativity. An example of

this is Fluorine, the first member of the halogen series. 

Fluorine is the most highest electronegative element available and is the smallest one too.

3. Screening effect or shielding effect

If the outermost electrons shields the nucleus effectively, the electronegativity of the element

decreases.

Periodic Table Electronegativity Trend

In a period

Electronegativity increases on moving in a period of the periodic table from left to right. This is due to

the increase in nuclear charge as a result of which the added electrons can be held more tightly.

Thus, C-N bond should be shown as Cδ+ - N δ- or C→N, the arrow head being towards the more

electronegative element, N.

In the same period, on moving from left to right, the electronegativities increase with the increase in

the number of outer electrons. An example of this can be shown with the help of the second period.

Element Li Be B C N O F

Electronegativity value

1.0

1.5

2.0

2.5

3.0

3.5

4.0

In a group

In moving down a group, from top to bottom, a new shell is added from the top element to the bottom

one. Also, the nuclear charge increases from top to bottom.

The increase in the nuclear charge indicates that the electronegativity of the lower element should be

more than that of the top element. This is not the case though. 

The reason is because of the increase in the atomic radii as we move down the group. The increase

in atomic radii increases the electron shielding effect, and it is much more than the increase in the

nuclear charge. Consequently, the lower elements are less electronegative than the top elements.

Thus, as we move down a group. electronegativity decreases.

The most electronegative elements, say, Fluorine for example, are present at the top right hand

corner of the periodic table. while the most electropositive elements or the less electronegative

elements are present at the bottom left hand corner, for example, Cs.

Trend for ElectronegativityElectronegativity value of elements follows the below mentioned trend:

1. Electronegativity increases as we move down from left to right across a period. This is

because the atomic size decreases across the period and nuclear attraction over electrons

increases.

2. On moving down a group, the atomic size increases and the nuclear attraction over outer

electrons decreases. Consequently, electronegativity decreases as we move from top to

bottom in a group.

3. On comparing the elements based on their group, Halogens are the most electronegative

elements in their respective periods.

4. Alkali metals, have the least electronegativity value in their periods. The last group in the

periodic table, the 18th group consists of Noble gases. They have zero electronegativity value

because of their completely filled outer shells.

5. Fluorine, the smallest element has the highest electronegativity value among all elements and

Helium has the lowest electronegativity value.

6. The 15, 16, and 17th group of elements are comparatively more electronegative due to the

almost filled outermost shells.

7. The first and second group, alkali metals and alkaline earth metals are least electronegative.

These are electropositive in nature.

8. Electronegativity value is only applicable for bonded atoms, and not for the isolated gaseous

atoms.

9. It is usually calculated from experimentally determined values of bond energies compared to

a reference.

Properties of Metal-Nonmetals And Metalloids

(a) Properties of metals

Almost all metals which have low ionization energies and low electro negativities are electropositive in nature and have a tendency to lose electrons. Hence they are a good reducing agent. All metals show different reactivity towards various regents. Like alkali metals, alkaline earth metals are highly reactive and react vigorously with water and dilute acids. But transition elements are less reactive compared to alkali metals. 

Some general properties of metals are:

1. Metals react with air to form oxides:

2M + O2 → 2 MO

The reactivity varies from metal to metal. Some metals, like beryllium and magnesium, react vigorously with air and form oxides. Magnesium burns with a typical intense white flame, while calcium is quite reluctant to start burning, but then bursts dramatically into flame, burning with a white flame.

2Mg + O2 → 2MgO

Ca + O2 → CaO

Some metals like copper and silver do not react vigorously, but take time.

Cu + O2 → CuO(Black)

Ag + O2 → Ag2O (Grey)

Once the oxide layer coats the metal surface, it will prevent further attack of oxygen. That is the reason, some metals like aluminum are reluctant to reaction with air.

2. Metals react with water to form hydroxides and hydrogen gas

Alkali metals reacts vigorously with water to produce hydrogen gas and an alkaline solution.

2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g)

2Li(s) + 2H2O(l) → 2LiOH(aq) + H2(g)

Alkaline earth metals are less reactive than alkali metals. For example, calcium reacts quickly with water to form calcium hydroxide and the bubbles you see in the beaker are because of hydrogen gas. The solution turns milky white in color because of the precipitation of calcium hydroxide.

Ca + 2H2O → Ca (OH) 2 + H2

As we move down the reactivity series of metals, their reactivity towards water decreases. Like aluminium does not react with cold water but forms aluminium oxide and hydrogen gas with steam. Less reactive metals like iron, zinc show no reaction with cold water but form oxides with steam and form oxides and hydrogen gas but reaction will be much less vigorous.

3. Metals react with non-metals to form ionic compounds

Since metals are electropositive in nature hence they can easily loos electron and form a metal cation. On the contrary, non-metals are electronegative and form anions. The cation and anion combine together by electrostatic force of attraction known as ionic bond and such compounds are termed as ionic compounds. For example, sodium metal reacts with fluorine to form sodium fluoride.

4. Metals react with acid to form salt and hydrogen gas

Reaction of metal with acid depends on their reactivity. Highly reactive metals can react with dilute acids and release hydrogen gas while less reactive metals react with strong acid under drastic conditions only. For example, magnesium strip reacts readily in dilute sulfuric acid and forms magnesium sulphate and hydrogen gas.

Mg(s) + H2SO4 (l) → MgSO4(s) + H2 (g)

(b) Properties of non-metals

There are only 17 non-metals (excluded H) arranged at right in the periodic table. They are electronegative elements with high electronegativity and ionization energy. They have a tendency to accept electrona and form anions.

Non-metals react with oxygen to form non-metallic oxides, which are acidic in nature. Hence non-metallic oxides form an acidic solution when dissolved in water and turn litmus solution red. For example, carbon is oxidized to form carbon dioxide, which is acidic in nature.

C(s) + O2 (g) → CO2 (g)

Non-metals are good oxidizing agents and are oxidized in almost all of their reactions. Like aluminum is oxidized with bromine to form aluminum bromide.

Non-metals have a tendency to oxidize metals.

2 Mg(s)+O2(g) → 2 MgO(s)

They can easily oxidize those compounds with which they react.

2H2S (g) +3O2 (g) → 2SO2 (g) + 2 H2O (g)

Less electronegative non-metals like carbon & hydrogen can act as a reducing agent for some compounds like ferric oxide, copper (II) oxide. 

Fe2O3(s)+3 C(s) → 2 Fe(s)+3 CO(g)

CuO(s)+H2(g) → Cu(s)+H2O(g)

Generally, no reaction takes place between non-metals and acids. But non-metals react with bases to form salts. For example, chlorine reacts with calcium hydroxide to form bleaching powder.

(c) Properties of metalloids

1. Metalloids tend to show an intermediate property between metals and non-metals. Some

metalloids like arsenic and antimony are crystalline solids.

2. However, the chemical properties of metalloids are either as metals or non-metals. They form

amphoteric oxides as metals form basic oxides but non-metals are generally acidic oxides.

3. The most important feature of metalloids is their semi conductivity. Some metalloids like

boron, silicon and germanium behave as semiconductors.

4. The chemical reactivity of metalloids depends on the substance they react with. For example,

when boron reacts with fluorine, it acts as a metal. While in reaction with sodium, it acts as a

non-metal.

5. Metalloids are usually brittle in nature and behave as electrical insulators at room temperature

but act as a conductor at a certain temperature. They are used as do-pants in glasses in

semiconductor chips.

Chemistry 101Dr. A. J. Pribula

Some Common Types of Chemical Reactions

1. When two elements react, a combination reaction occurs (think: could any other type of reaction occur?), producing a binary compound (that is, one consisting of only two types of atoms). If a metal and a nonmetal react, the product is ionic with a formula determined by the charges on the ions the elements form. If two nonmetals react, the product is a molecule with polar covalent bonds, with a formula consistent with the normal valences of the atoms involved. Some pairs of elements may react only slowly and require heating for significant reaction to occur.

Examples: 

K + S8  K2S (ionic)Ca + O2  CaO (ionic)Al + I2   AlI3 (ionic)H2 + O2   H2O (covalent)I2 + Cl2   ICl, ICl3, or ICl5 (covalent)(exact product depends on relative amounts of I2 and Cl2)(NOTE: The above reactions are not balanced, nor were they intended to be. They, like the others in this handout, are meant only to show the correct formulae for the reactants and products. You may wish to balance the reactions in the handout as an exercise.)

2. Reaction of a metal oxide with water produces a metal hydroxide; that is, a strong base. Reaction of a nonmetal oxide with water produces an oxyacid in which the nonmetal is in the same oxidation state as in the oxide you started with. Both of these are combination reactions, and both can be reversed by heating the products. Metal hydroxides decompose on heating to give the metal oxide and water, and oxyacids decompose on heating to give water and the nonmetal oxide in the appropriate oxidation state.

Examples:

Na2O + H2O   NaOHMgO + H2O   Mg(OH)2

SO2 + H2O   H2SO3

Cl2O5 + H2O   HClO3

HNO3  N2O5 + H2O

Fe(OH)3   Fe2O3 + H2O

3. Reaction of a metal oxide with a nonmetal oxide gives an oxysalt; reaction of a metal hydroxide with a nonmetal oxide produces a "hydrogen" oxysalt. This is essentially a reaction of the O2- or OH- in the metal compound with the molecular nonmetal oxide. This combination reaction occurs only if no water is present; in the presence of water, the nonmetal and metal oxides react with the water to produce acid and hydroxide, respectively (as shown in (2) above), then these react as in (4) below.

Examples:

CaO(s) + SO3(g)   CaSO4(s)NaOH(s) + CO2(g)   NaHCO3(s)

4. Reaction of an acid with a base gives a salt plus water. The cation in the salt comes from the base; the anion comes from the acid. The base may be a metal hydroxide, a metal oxide, or a weak base such as NH3. The acid and/or base may be pure solids, liquids, or gases, or in aqueous solution. The oxidation states of the anion of the acid and cation of the base normally remain unchanged.

Examples:

HCl(aq) + Ca(OH)2(aq)   CaCl2(aq) + H2O(l)H2SO4(aq) + Fe(OH)3(s)   Fe2(SO4)3(aq) + H2O(l)NH3(g) + HC2H3O2(l)   NH4C2H3O2(s)Al2O3(s) + HClO4(aq)   Al(ClO4)3(aq) + H2O(l)

5. Ammonium salts react with metal hydroxides and oxides in an acid-base reaction to produce ammonia. This is essentially the reverse of one of the reaction types mentioned in (4) above. Either or both of the reactants may be a pure material or in aqueous solution.

Examples:

NH4Cl(aq) + KOH(aq)   NH3(g) + H2O(l) + KCl(aq)NH4NO3(s) + CaO(s)   NH3(g) + H2O(l) + Ca(NO3)2(s)

6. Reaction of the salt of a weak acid (that is, a compound containing the anion of a weak acid) with a strong acid produces the weak acid and a salt. This is another example of an acid-base reaction, in addition to the ones given in (4) and (5) above. The original salt of the weak acid may be either a pure solid or in aqueous solution. The cation in the salt formed as the product comes from the weak acid salt; the anion in the product salt comes from the strong acid. In many cases, the weak acid produced is unstable and decomposes to give the oxide of a nonmetal and water (see (2) above). This is especially true if the nonmetal oxide is a compound of limited solubility in water such as SO2, CO2, or the nitrogen oxides. The best-known examples of this type of reaction involve carbonates, bicarbonates, sulfides, and sulfites, but many other examples are known as well. Normally, these reactions do not involve oxidation or reduction.

Examples:

BaCO3(s) + HBr(aq)   BaBr2(aq) + H2O(l) + CO2(g)NaHCO3(aq) + H2SO4(aq)   Na2SO4(aq) + CO2(g) + H2O(l) MgS(s) + HCl(aq)   H2S(g) + MgCl2(aq)K2SO3(aq) + HNO3(aq)   KNO3(aq) + SO2(g) + H2O(l)Ca3(PO4)2(s) + HCl(aq)   CaCl2(aq) + H3PO4(aq)Zn(C2H3O2)2(aq) + HBr(aq)   ZnBr2(aq) + HC2H3O2(aq)

7. Reaction of solutions of two soluble salts with one another can give a precipitate of an insoluble salt formed by a double replacement reaction (also called a metathesis). Whether or not a precipitate forms depends on the exact combination of salts used. To make a prediction

as to whether a reaction will take place or not, you must know the solubility rules for common salts (Ebbing 4/e, page 104; lab manual, Appendix 7). Some combinations of salts may give oxidation-reduction reactions (see (11) below), but most do not.

Examples:

CaCl2(aq) + K2CO3(aq)   CaCO3(s) + KCl(aq)AgNO3(aq) + FeCl3(aq)   AgCl(s) + Fe(NO3)3(aq)

but: NiSO4(aq) + MgI2(aq)   no reaction 

(NiI2 and MgSO4 are both soluble)Al(NO3)3(aq) + Pb(C2H3O2)2(aq)   no reaction (Al(C2H3O2)3 and Pb(NO3)2 are both soluble)

8. Heating an oxysalt produces a metal oxide plus a nonmetal oxide or a metal salt plus oxygen, or some combination of these two decomposition reactions.

Examples:

KClO3(s)   KCl(s) + O2(g)

CaCO3(s)   CaO(s) + CO2(g) 

Pb(NO3)2(s)   PbO(s) + NO(g) + NO2(g) + O2(g)

9. Heating a hydrated material initially causes a decomposition reaction to produce the anhydrous compound and water. Further heating may yield further decomposition, depending on the material. (See (2) and (8) above.) Most binary compounds are stable to heat.

Examples:

H2C2O4. 2H2O(s)   H2O(g) + H2C2O4(s); followed by

H2C2O4(s)   H2O(g) + CO(g) + CO2(g)

CaCl2. 6H2O(s)   H2O(g) + CaCl2(s); followed by

CaCl2(s)   no reaction

CuSO4. 5H2O(s)   H2O(g) + CuSO4(s); followed by

CuSO4(s)   CuO(s) + SO3(g) (requires strong heating)

10. Reaction of an element with a compound often gives a single replacement reaction in which a nonmetallic element can replace a combined nonmetal, and a metallic element can replace a combined metal, or hydrogen from an acid. As a general rule, a more active (reactive) element will replace a less active (reactive) element from its compounds. In general (but with many exceptions), the most reactive nonmetals are found to the upper right in the periodic table, and the most reactive metals are found to the lower left. The order of reactivity of the halogens is F2>Cl2>Br2>I2. For hydrogen and the more common metals, the order of reactivity (the activity series) is

Li>K>Ca>Na>Mg>Al>Zn>Cr>Fe>Ni>Sn>Pb>H2>Cu>Hg>Ag>Pt>Au

In these two series, one element can replace another one to its right in the series. Metals to the left of H2 can replace H+ from acids. The very reactive metals (Li, K, Na, Ca) can replace H+ from cold water; metals of intermediate reactivity (Mg, Al) can replace H+ from hot water or steam. Any single replacement reaction can also be categorized as an oxidation-reduction (redox) reaction.

Examples:

Al(s) + NiSO4(aq)   Al2(SO4)3(aq) + Ni(s)Fe(s) + HBr(aq)   FeBr3(aq) + H2(g)Cl2(g) + KI(aq)   KCl(aq) + I2(s)Na(s) + H2O(l)   NaOH(aq) + H2(g)Zn(s) + Cu(NO3)2(aq)   Cu(s) + Zn(NO3)2(aq)

but: Ag(s) + HClO4(aq)   no reaction

Br2(l) + ZnCl2(aq)   no reactionSn(s) + H2O(l)   no reactionPb(s) + CrF3(aq)   no reaction

11. Compounds containing one or more atoms in high oxidation states often act as oxidizing agents; compounds containing atoms in low oxidation states often act as reducing agents. For most elements, the (old) group number of the atom in the periodic table gives the highest oxidation state possible for that element. For nonmetals, the lowest oxidation state possible is given by the (old) group number minus eight. Elemental metals most often act as reducing agents (they are oxidized); nonmetals frequently act as oxidizing agents (they are reduced).

For the representative elements (i.e., those in the first two and last six columns of the periodic table), oxidation states most often are two units apart. For example, Sn forms Sn(II) and Sn(IV); Br forms Br1-, Br(I), Br(III), Br(V), and Br(VII). For the transition elements, (i.e., those in the "center" ten columns of the periodic table), oxidation states are often one unit apart, but can be in almost any relationship to one another. For the transition elements, the common oxidation states (charges on their ions) must be memorized. For example, Fe forms Fe2+ and Fe3+; Cu forms Cu+ and Cu2+, etc. Some of the transition elements form oxyanions as well as cations. For example, Mn forms Mn2+, Mn3+, MnO4

2-, and MnO4-; Cr forms Cr2+, Cr3+,

CrO42-, and Cr2O7

2-.

Any atom in its highest possible oxidation state can only act as an oxidizing agent; any atom in its lowest possible oxidation state can only act as a reducing agent. Atoms in intermediate oxidation states can be either oxidized or reduced; that is, they can act as either reducing or oxidizing agents. Some of the oxidizing agents most commonly encountered are MnO4

-, CrO4

2-, Cr2O72-, HNO3, H2O2, and the halogens. Some of the more common reducing agents

are elemental H2, metals, carbon, and I-.

In predicting products of oxidation-reduction reactions, don't forget their name--oxidation and reduction   must   occur simultaneously ! It is impossible for oxidation to occur without reduction or vice versa.

Examples:

Sn2+(aq) + F2(g)   Sn4+(aq) + F-(aq)Mn2+(aq) + BiO3

-(aq)   Bi3+(aq) + MnO4-(aq)

(note that the Bi is in its highest possible oxidation state in BiO3-)

K(s) + P4O10(s)   K3PO3(s) (note that P is reduced from P(V) to P(III))MnO4

-(aq) + I-(aq)   Mn2+(aq) + I2(aq)CuS(s) + HNO3(aq)   Cu(NO3)2(aq) + S8(s) + NO2(g) (note S2-   S0 and N(V)   N(IV))Fe2O3(s) + C(s)   CO2(g) + Fe(s)

[REACT101.S94/AJP1]

Reaction of Metals with Oxygen

Metals like sodium (Na) and potassium (K) are some of the most reactive metals. Potassium, sodium,

lithium, calcium and magnesium react with oxygen and burn in air.

Metals from aluminium to copper in the activity series of metals, react slowly when heated in air to form the metal oxides. Aluminium is the fastest and copper is the slowest of them.

Sodium metal reacts with the oxygen of the air at room temperature to form sodium oxide.

Hence, sodium is stored under kerosene to prevent its reaction with oxygen, moisture and

carbon dioxide.

Sodium oxide is a basic oxide which reacts with water to form sodium hydroxide.

Mg does not react with oxygen at room temperature. On heating, Mg burns in air with intense

light and heat to form MgO.

Zinc metal burns in air only on strong heating to form zinc oxide.

Iron metal does not burn in dry air even on strong heating. In moist air, iron is oxidized to give

rust.

On heating in air it burns with a brilliant flame forming triferric

tetroxide. 

Copper is the least reactive metal and does not burn in air even on heating. However, on

prolonged strong heating copper reacts with oxygen and forms copper (II) oxide (CuO)

outside and copper (I) oxide (Cu2O) inside.

 

Gold and platinum do not react with oxygen in air.

Reaction of Metals with Water

Potassium, sodium, lithium and calcium react with cold water.

Sodium reacts vigorously with cold water forming sodium hydroxide and hydrogen.

Metals from magnesium to iron in the activity series of metals, react with steam (but not H2O)

to form the metal oxide and hydrogen gas.

Red hot iron reacts with steam to form Iron (II, III) oxide.

Note: The reaction between iron and steam is irreversible. Tin, lead, copper, silver, gold and platinum

do not react with water or steam.

Reaction of Metals with Acids

Potassium, sodium, lithium and calcium react violently with dilute H2SO4 and dilute HCl,

forming the metal salt (either sulphate or chloride) and hydrogen gas. The reaction is similar

to the reaction with water.

         

Magnesium, aluminium, zinc, iron, tin and lead react safely with dilute acid. Magnesium is the

fastest and lead is the slowest of the six.

Zinc with dilute sulphuric acid is often used for the laboratory preparation of hydrogen. The reaction is

slow at room temperature, but its rate can be increased by the addition of a little copper (II) sulphate.

Zinc displaces copper metal, which acts as a catalyst.

Metals below hydrogen (copper, silver, gold and platinum), will not react with dilute acid. They cannot

displace hydrogen from the non-metal anion.

Note:

Copper reacts with oxyacids like nitric acid and sulphuric acid because

these acids are strong oxidizing agents.

In general,

Hydrochloric acid makes a metal chloride.

Sulphuric acid makes a metal sulphate.

Reactions with nitric acid are more complex, the nitrate is formed but the gas is rarely

hydrogen, and more often, an oxide of nitrogen.

Reaction of Metals with Salt Solutions

Reactive metals can displace any metal less reactive than itself, from the oxide, chloride or sulphate

of the less reactive metal in solution or their molten state. If metal A displaces metal B from its

solution, it is more reactive than B.

Copper (II) sulphate solution is blue; iron sulphate solution is almost colourless when dilute. During

the displacement, the blue solution loses its colour and the iron metal is seen to turn pink-brown as

the displaced copper becomes deposited on it.

On heating the mixture of magnesium powder and black copper (II) oxide, white magnesium oxide is formed with brown bits of copper:

Adding magnesium to blue copper (II) sulphate solution, the blue colour fades as colourless magnesium sulphate is formed and brown bits of copper metal form a precipitate:

Summary of the Reaction of Metals with Air, Water and Acids

Alkaline Earth Metals Properties

The chemical properties of the alkaline earth metals are as follows

Formation of oxides

The alkaline earth metals form the normal oxides of MO type which are obtained by heating the metal

in O2 or by heating their carbonates at high temperature.

Ca + O2 → CaO

CaCO3 → CaO + CO2

These oxides are extremely stable white crystalline solids due to their high crystal lattice energy

obtained by packing doubly charged ions in a sodium chloride type of lattice.

Formation of hydroxides

The hydroxides are of the type M(OH)2. These are obtained by the reaction of the metals with water.

Ca + 2H2O → Ca(OH)2 + O2

Be does not react with water even at elevated temperature, Mg reacts only with steam while other

metals react with cold water as well.

The hydroxides can also be prepared by reaction of metal oxides with water.

MO + H2O → M(OH)2

Be(OH)2 is amphoteric in nature where as hydroxides of other metals are basic in nature. Their basic

character increases on moving down the group.

Formation of halides

The halides of these metals are of the type MX2. They can be prepared by

By heating the metals with halogens at appropriate temperatures

By reacting metal carbonates with dilute halogen acids

Halides of Be are covalent in nature. Therefore, do not conduct electricity in fused state. They are

hygroscopic and fume in air due to hydrolysis. They have low melting points. The covalent character

of Be is due to high polarizing power of small Be2+ cation.

The chlorides and fluorides of other metals are ionic in nature and hence good conductor of electricity

in the fused as well as in the dissolved state. The melting point and the conductivity increases on

moving down the group due to the increase in the ionic character of the halides.

Formation of hydrides

All the elements of this group form hydrides of MH2 type. BeH2 can't be prepared by direct

combination of Be and H2. It is however prepared by reducing BeCl2 by LiH or LiAlH4.

BeCl2 + 2LiH → BeH2 + 2LiCl

2BeCl2 + LiAlH4 → 2BeH2 + LiCl + AlCl3

Mg, Ca, Sr and Ba form their hydrides when heated directly with H2.

Formation of carbonates

All the carbonates decompose on heating at appropriate temperature to give CO2.

MCO3 → CO2 + MO

The stability of the carbonates of these metals increases on moving down the group. The marked

instability of BeCO3 is due to the strong polarizing effect of small Be2+ ion on the large polarizable

CO32- ion.

Carbonates are insoluble in water and therefore occur as solid rock minerals in nature. However, they

dissolve in H2O containing CO2 due to the formation of bicarbonates.

CaCO3 + CO2 +H2 → Ca(HCO3)2

Formation of nitrides

All the elements burn in nitrogen to form nitrides, M3N2

3Ca + N2 → Ca3N2

These react with H2O to liberate NH3,

Ca3N2 + 6H2O → 3Ca(OH)2 + 2NH3

Be3N2 is volatile while other nitrides are not so.

Formation of carbides

When the elements from Mg to Ba or their oxides are heated with carbon, then carbides are formed

(MC2).

CaO + 3C → CaC2 + CO

All these carbides are ionic in nature and have NaCl type structure.

Chemical Properties of Boron

Action of air and O2

Amorphous Boron is very reactive. When it is heated at 700 C in the air or O2, it burns with a reddish

flame forming a mixture of oxide and nitrate.

4B + 3O2 → 2B2O3

2B + N2 → 2BN

Action of acids

Halogen acids have no action on boron but it is attacked by oxidizing acids like HNO3 and H2SO4.

B + 3HNO3 → H3BO3 + 3NO2

2B + 3H2SO4 → 2H3BO3 + 3SO2

As reducing agentsBoron is a powerful reducing agent, since it can replace carbon from CO2, silicon from SiO2, and nitrogen from NO.

4B + 3CO2 → 2B2O3 + 3C

4B + 3SiO2 → 2B2O3 + 3Si

4B + 6NO → 2B2O3 + 3N2

Action of Alkalies

It is vigorously attacked by fused alkalies or by fusion with oxidizing fusion mixture.

2B + 6NaOH → 2Na3BO3 + 3H2

Action of metals

Except for Cu, Ag and Au, Boron combines with other metals at high temperatures to form borides.

These borides are extremely hard substances. Mg3B2 and Be3B2 are attacked by acids to form boron

hydrides.

Action of non-metals

Boron also reacts with non-metals. For example: Boron combines with N2 at temperatures greater

than 900 C to form boron nitride, BN. Boron forms boron sulphide B2S3 when heated with S. On

heating with Carbon in an electric furnace an extremely hard substance namely boron carbide, B4C is

formed.

Action of steam

Boron reacts with steam at red heat liberating H2.

2B + 3H2O → B2O3 + 3H2

Boron Compounds

Boron forms various compounds with different elements. 

The main compounds of boron are,

1. Boric acid

2. Borax

3. Borazine

4. Boron nitride

5. Boron trihalides

6. Ffluoroboric acid

Boric Acid

Boron forms several boric acids. Boric acids like orthoboric acid, metaboric acid, pyroboric acid, and

tetraboric acid are known. Out of these acids, orthoboric acid which is generally referred to as boric

acid is the most important and is stable both in the solid state and as solution. The remaining acids

are stable in the solid state and gradually change into orthoboric acid in solution. The formula of

orthoboric acid or boric acid is H3BO3.

HBO2 + H2O → H3BO3

H2B4O7 + 5H2O → 4H3BO3

Small quantities of boric acid are present in the jets of steam called saffioni which is mixed with other

gases in the volcanic district of Tuscany. Besides this it is also found in many mineral waters. It is also

present in traces of hops, in batteries and fruits and often in wines.

Borax

1. The formula of borax is Na2B4O7.10H2O. It is the sodium salt of tetraboric acid, H2B4O7.

2. It occurs naturally as tincal or suhaga in the dried up lakes of Tibet, Ceylon, California and

India. In India, it is found in Ladakh and the Puga valley of Kashmir.

3. Tincal contains about 55% Borax. Borax can be obtained from colemanite and by the action

of Na2CO3 on H3BO3.

4. It is sparingly soluble in cold water and fairly soluble in hot water. The solution is alkaline in

nature.

Borazine

1. Borazine is also known as borazole.

2. It has the formula B3N3H6. This compound is isoelectronic with benzene and hence has been

called inorganic benzene by Weberg. Borazine is a colorless, mobile volatile liquid. It freezes

at -58 C.

3. It's boiling points and melting points are: 63 C and -58 C respectively. Borazine gets slowly

hydrolyzed by water to produce H2, boric acid and NH3.

4. Hydrolysis is favored by increase in temperature.

Boron Nitride

1. The formula is BN. It is a white powder with a density of 2.34. It melts under pressure at 3000

C.

2. It is a very stable and unreactive substance. It remains unaffected by mineral acids, solutions

of alkalies and Cl2 at red heat.

3. It gets decomposed when heated in steam to give NH3.

4. It also decomposes when heated with fused KOH.

BN + 3H2O → H2BO3 + NH3

BN + KOH → K3BO3 + NH3

BN possesses the same hardness as a diamond and can withstand temperatures of more than 300 C.

Due to this property it is used for coating crucible linings.

Boron TrihalideBoron reacts with halogens to form trihalides BX3, where X= F, Cl, Br, I. All the trihalides of boron

except BF3 can be prepared by the treatment of a mixture of B2O3 and carbon with the appropriate

halogen at higher temperature.

B2O3 + 3C + 3X2 → 2BX3 + 3CO

BF3 and BCl3 are gases, BBr3 is a liquid and BI3 is a solid at room temperature. Boron trihalides are

covalent in nature due to small size and the high charge density on B3+ ion. Some other points which

favor the covalent character of boron trihalides are: as liquids they do not conduct electricity, their

boiling points are very low as compared to the halides of the elements of groups I A and II A, they

exist as discrete molecular species.

Fluoroboric Acid

1. The formula is HBF4. It is a strong acid, stronger than HF. It is decomposed when heated.

2. The salts given by this acid are fluoroborides or fluoroborates.

3. These salts are prepared by dissolving the corresponding metal borides, hydroxides or

carbonates in aqueous hydrofluoric acid.

4. Fluoroborates are also obtained in solution by treating alkali acid fluorides with H3BO3.

Oxygen FamilyReaction with Metals

Dioxygen forms alkaline oxides with active metals. But less reactive metals like gold and platinum. Active alkali metal & alkaline earth metal form oxides, peroxide & super oxide with dioxygen.

4K + O2(g) → 2K2O

2K + O2(g) → K2O2

K + O2(g) → KO2

With other metals, dioxygen forms metal oxides.

2 Zn(s) + O2 (g) → 2 ZnO(s)

Reaction with Non-metal

Just like reaction with metal, dioxygen reacts with non-metals also like hydrogen, carbon, sulfur and phosphorus & form oxides. Reaction occurs at high temperature or in electric discharge.

For example, dioxygen forms water with hydrogen and with solid carbon; it forms carbon monoxide or carbon dioxide. Similarly, with solid phosphorus, it forms tetraphorphorus heptoxide or tetraphorphorus decoxide. When dioxygen reacted with solid sulfur, sulfur dioxide gas forms.

2 H2 (g) + O2(g) → 2 H2O(g)

C(s) + O2 (g) → CO(g) or CO2(g)

P4(s) + O2(g) → P4O6(g) or P4O10(g)

S8(s) + 8 O2 (g) → 8 SO2(g)

Reaction with other Compounds

Dioxygen can react with different organic and inorganic compounds to form various products. 

For example,

With sulfur dioxide it form sulfur trioxide at 723 K temperature and 2 atm pressure in the presence of platinum of V2O5 as a catalyst. This reaction used to prepare sulfuric acid in contact process. Reaction is reversible and exothermic in nature. 

              723k,2atm

2SO2(g)+O2(g)       ⇌         2SO3(g)

        V205

With ammonia, dioxygen forms nitric oxide at 500K temperature and in the presence of platinum. Reaction is used in Ostwald process for the preparation of nitric acid.

4NH3(g) + 5O2(g) −→−−−−500K,Pt 4NO(g) + 6H2O(g)

In Decon's process, dioxygen oxidized hydrochloric acid to form water and chlorine gas. Reaction talks place at 700 K temperature and in the presence of CuCl2 catalyst.

4HCl (g) + O2 (g) −→−−−−−−700K,CuCl2 2H2O(g) + Cl2(g)

Electric discharge

Dioxygen forms ozone under the action of silent electric discharge.

3O2(g) → 2O3(g)

Respiration

It's a very important chemical property of dioxygen. It involves in respiration of all living bodies. Basically respiration is a combustion process of carbohydrates to produce carbon dioxide and water with a large amount of energy.

C6H12O6 (aq) + 6 O2 (g) → 6 CO2 (g) + 6 H2O(l)

Chemical Properties of Halogens1. Oxidizing Power

Since all halogens have a strong tendency to accept electrons, they act as good oxidizing agents. Out

of all the halogens, fluorine is the strongest oxidizing agent and can oxidize all other halide ions to

halogen in a solution. As we move down the group from F to I , oxidizing power decreases. Hence

chlorine can oxidize bromide ion to bromine as well as iodide ions to iodine.

Cl2 + 2Br¯ → Br2 + 2Cl¯

Cl2 + 2I¯ → I2 + 2Cl¯

In the same way, bromine can oxidize iodide ion to iodine.

Br2 + 2I¯ → I2 + 2Cl¯

Cl2(aq) Br2(aq) I2(aq)

Cl–

(aq)

Stays yellow solution (no

reaction)

Stays brown solution (no

reaction)

Br–

(aq)

Yellow solution forms

(Br2 forms)

Cl2 + 2 Br- → 2 Cl- + Br2

Stays brown solution (no

reaction)

I–(aq) Brown solution forms (I2 Brown solution forms(I2 forms)

forms)

Cl2 + 2 I- → 2 Cl- + I2Br2 + 2 I- →2 Br- + I2

  On the contrary, halide ions behave as reducing agents. Their reducing ability decreases from

fluoride ion to iodide ion.

least powerful F¯ < Cl¯ < Br¯ < I¯ most powerful reducing agent

2. Reaction with hydrogen

All halogens react with hydrogen to form hydrogen halides which are acidic in nature. The acidity of

hydrogen halides decrease from HF to HI. But the reactivity of halogens towards hydrogen decreases

from fluorine to iodine. Fluorine reacts violently in the dark; chlorine requires sunlight, while bromine

combines with hydrogen only on heating. Iodine reacts with hydrogen on heating in the presence of a

catalyst.

In dark

H2 + F2 → 2HF

In sunlight

H2 + Cl2 → 2HCl

Δ

H2 + Br2 → 2HBr

Δ

H2 + I2 → 2HI

3. Reaction with oxygen

Like other elements, halogens also form oxides with oxygen. But most of the oxides of halogen are

unstable. Apart from oxides, halogens also form halogen oxoacids and oxoanions. The general

formula for oxides is in the range from X2O to X2O7, while the general formula for oxoacids is from

HOX to HOXO3 (there is only HOF with fluorine) and for oxoanions are formed in the range XO - to

XO4-.

4. Reaction with metals

Due to the high reactivity of halogens, they readily react with most of the metals to form the

corresponding metal halides. For example, sodium reacts with chlorine gas to form sodium chloride.

Formation of sodium chloride is an exothermic reaction and produces a bright yellow light with a large

amount of heat energy.

2Na(s) + Cl2(g) → 2NaCl(s)

Metal halide is ionic in nature due to the high electro-negativity of halogen and high electro positivity

of metals. The ionic character of metal halides decreases from fluorine to iodine.

5. Reaction with other Halogens

Halogens react with each other to form Inter halogen compounds. The general formula of these

compounds is XYn, where n = 1, 3, 5 or 7. In a given formula, ‘X’ must be the less electronegative

halogen compared to ‘Y’.

XY XY3 XY5 XY7

ClF, BrF, BrCl, ICI, IBr,

IFCIF3, BrF3,IF3, ICI3

Br

F5

IF7