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Ionic bond

An ionic bond is a type of chemical bond that involves a metal and a nonmetalion (or polyatomic ions such as ammonium) through electrostatic attraction. In short, it is a bond formed by the attraction between two oppositely charged ions.

The metal donates one or more electrons, forming a positively charged ion or cation with a stable electron configuration. These electrons then enter the non metal, causing it to form a negatively charged ion or anion which also has a stable electron configuration. The electrostatic attraction between the oppositely charged ions causes them to come together and form a bond.

For example, common table salt is sodium chloride. When sodium (Na) and chlorine (Cl) are combined, the sodium atoms each lose an electron, forming cations (Na+), and the chlorine atoms each gain an electron to form anions (Cl−). These ions are then attracted to each other in a 1:1 ratio to form sodium chloride (NaCl).

Na + Cl → Na+ + Cl−→ NaCl

The removal of electrons from the atoms is endothermic and causes the ions to have a higher energy. There may also be energy changes associated with breaking of existing bonds or the addition of more than one electron to form anions. However, the attraction of the ions to each other lowers their energy. Ionic bonding will occur only if the overall energy change for the reaction is favourable – when the bonded atoms have a lower energy than the free ones. The larger the resulting energy change the stronger the bond. The low electronegativity of metals and high electronegativity of non-metals means that the energy change of the reaction is most favorable when metals lose electrons and non-metals gain electrons.

Pure ionic bonding is not known to exist. All ionic compounds have a degree of covalent bonding. The larger the difference in electronegativity between two atoms, the more ionic the bond. Ionic compounds conduct electricity when molten or in solution. They generally have a high melting point and tend to be soluble in water.

Ionic structure

Ionic compounds in the solid state form lattice structures. The two principal factors in determining the form of the lattice are the relative charges of the ions and their relative sizes. Some structures are adopted by a number of compounds; for example, the structure of the rock salt sodium chloride is also adopted by many alkali halides, and binary oxides such as MgO.

Strength of an ionic bond

For a solid crystalline ionic compound the enthalpy change in forming the solid from gaseous ions is termed the lattice energy. The experimental value for the lattice energy can be determined using the Born-Haber cycle. It can also be calculated using the Born-Landé equation as the sum of the electrostatic potential energy, calculated by summing interactions between cations and anions, and a short range repulsive potential energy term. The electrostatic potential can be expressed in terms of the inter-ionic separation and a constant (Madelung constant) that takes account of the geometry of the crystal. The Born-Landé equation gives a reasonable fit to the lattice energy of e.g. sodium chloride where the calculated value is −756 kJ/mol which compares to −787 kJ/mol using the Born-Haber cycle.

Polarization effects

Ions in crystal lattices of purely ionic compounds are spherical; however, if the positive ion is small and/or highly charged, it will distort the electron cloud of the negative ion, an effect summarised in Fajans' rules. This polarization of the negative ion leads to a build-up of extra charge density between the two nuclei, i.e., to partial covalency. Larger negative ions are more easily polarized, but the effect is usually only important when positive ions with charges of 3+ (e.g., Al3+) are involved. However, 2+ ions (Be2+) or even 1+ (Li+) show some polarizing power because their sizes are so small (e.g., LiI is ionic but has some covalent bonding present). Note that this is not the ionic polarization effect which refers to displacement of ions in the lattice due to the application of an electric field.

Ionic versus covalent bonds

In an ionic bond, the atoms are bound by attraction of opposite ions, whereas, in a covalent bond, atoms are bound by sharing electrons. In covalent bonding, the molecular geometry around each atom is determined by VSEPR rules, whereas, in ionic materials, the geometry follows maximum packing rules.

In reality, purely ionic bonds do no

Ionic radius

Ionic radius, rion, is a measure of the size of an atom's ion in a crystal lattice. It is measured in either picometres (pm) or Angstrom (Ã…), with 1 Ã… = 100 pm. Typical values range from 30 pm (0.3 Ã…) to over 200 pm (2 Ã…).

The concept of ionic radius was developed independently by Victor Goldschmidt and Linus Pauling in the 1920s to summarize the data being generated by the (at the time) new technique of X-ray crystallography: it is Pauling's approach which proved to be the more influential. X-ray crystallography can readily give the length of the side of the unit cell of a crystal, but it is much more difficult (in most cases impossible, even with more modern techniques) to distinguish a boundary between two ions. For example, it can be readily determined that each side of the unit cell of sodium chloride is 564.02 pm in length, and that this length is twice the distance between the centre of a sodium ion and the centre of a chloride ion:

2[rion(Na+) + rion(Cl−)] = 564.02 pm

However, it is not apparent what proportion of this distance is due to the size of the sodium ion and what proportion is due to the size of the chloride ion. By comparing many different compounds, and with a certain amount of chemical intuition, Pauling decided to assign a radius of 140 pm to the oxide ion O2−, at which point he was able to calculate the radii of the other ions by subtraction.

A major review of crystallographic data led to the publication of a revised set of ionic radii in 1976, and these are preferred to Pauling's original values. Some sources have retained Pauling's reference of rion(O2−) = 140 pm, while other sources prefer to list "effective" ionic radii based on rion(O2−) = 126 pm. The latter values are thought to be a more accurate approximation to the "true" relative sizes of anions and cations in ionic crystals.

The ionic radius is not a fixed property of a given ion, but varies with coordination number, spin state and other parameters. Nevertheless, ionic radius values are sufficiently transferable to allow periodic trends to be recognized. As with other types of atomic radius, ionic radii increase on descending a group. Ionic size (for the same ion) also increases with increasing coordination number, and an ion in a high-spin state will be larger than the same ion in a low-spin state. Anions (negatively charged) are almost invariably larger than cations (positively charged), although the fluorides of some alkali metals are rare exceptions. In general, ionic radius decreases with increasing positive charge and increases with increasing negative charge.

An "anomalous" ionic radius in a crystal is often a sign of significant covalent character in the bonding. No bond is completely ionic, and some supposedly "ionic" compounds, especially of the transition metals, are particularly covalent in character. This is illustrated by the unit cell parameters for sodium and silverhalides in the table. On the basis of the fluorides, one would say that Ag+ is larger than Na+, but on the basis of the chlorides and bromides the opposite appears to be true. This is because the greater covalent character of the bonds in AgCl and AgBr reduces the bond length and hence the apparent ionic radius of Ag+, an effect which is not present in the halides of the more electropositive sodium, nor in silver fluoride in which the fluoride ion is relatively unpolarizable.


The concept of ionic radii is based on the assumption of a spherical ion shape. However, from a group-theoretical point of view the assumption is only justified for ions that reside on high-symmetry crystal lattice sites like Na and Cl in halite or Zn and S in sphalerite. A clear distinction can be made, when the point symmetry group of the respective lattice site is considered, which are the cubic groups O6 and Td in NaCl and ZnS. For ions on lower-symmetry sites significant deviations of their electron density from a spherical shape may occur. This holds in particular for ions on lattice sites of polar symmetry, which are the crystallographic point groups C1, C1h, Cn or Cnv, n = 2, 3, 4 or 6. A thorough analysis of the bonding geometry was recently carried out for pyrite-type disulfides, where monovalent chemical reaction where the reactant entities are given on the left hand side and the product entities on the right hand side. The coefficients next to the symbols and formulae of entities are the absolute values of the stoichiometric numbers. The first chemical equation was diagrammed by Jean Beguin in 1615.


A chemical equation consists of the chemical formulas of the reactants (the starting substances) and the chemical formula of the products (substances formed in the chemical reaction). The two are separated by an arrow symbol (\rightarrow, usually read as "yields") and each individual substance's chemical formula is separated from others by a plus sign.

As an example, the formula for the burning of methane can be denoted:

CH|4| + 2 O|2| \rightarrow CO|2| + 2 H|2|O

This equation would be read as "CH four plus O two produces CO two and H two O." But for equations involving complex chemicals, rather than reading the letter and its subscript, the chemical formulas are read using IUPAC nomenclature. Using IUPAC nomenclature, this equation would be read as "methane plus oxygen yields carbon dioxide and water."

This equation indicates that oxygen and CH4 react to form H2O and CO2. It also indicates that two oxygen molecules are required for every methane molecule and the reaction will form two water molecules and one carbon dioxide molecule for every methane and two oxygen molecules that react. The stoichiometric coefficients (the numbers in front of the chemical formulas) result from the law of conservation of mass and the law of conservation of charge (see "Balancing Chemical Equation" section below for more information).

Common symbols

Symbols are used to differentiate between different types of reactions. To denote the type of reaction:

  • "=" symbol is used to denote a stoichiometric relation.
  • "\rightarrow" symbol is used to denote a net forward reaction.
  • "\rightleftarrows" symbol is used to denote a reaction in both directions.
  • "\rightleftharpoons" symbol is used to denote an equilibrium.

Physical state of chemicals is also very commonly stated in parentheses after the chemical symbol, especially for ionic reactions. When stating physical state, (s) denotes a solid, (l) denotes a liquid, (g) denotes a gas and (aq) denotes an aqueous solution.

If the reaction requires energy, it is indicated above the arrow. A capital Greek letter delta (\Delta) is put on the reaction arrow to show that energy in the form of heat is added to the reaction. h\nu is used if the energy is added in the form of light.

Balancing chemical equations

The law of conservation of mass dictates the quantity of each element does not change in a chemical reaction. Thus, each side of the chemical equation must represent the same quantity of any particular element. Similarly, the charge is conserved in a chemical reaction. Therefore, the same charge must be present on both sides of the balanced equation.

One balances a chemical equation by changing the scalar number for each chemical formula. Simple chemical equations can be balanced by inspection, that is, by trial and error. Another technique involves solving a system of linear equations.

Ordinarily, balanced equations are written with smallest whole-number coefficients. If there is no coefficient before a chemical formula, the coefficient 1 is understood.

The method of inspection can be outlined as putting a coefficient of 1 in front of the most complex chemical formula and putting the other coefficients before everything else such that both sides of the arrows have the same number of each atom. If any fractional coefficient exist, multiply every coefficient with the smallest number required to make them whole, typically the denominator of the fractional coefficient for a reaction with a single fractional coefficient.

As an example, the burning of methane would be balanced by putting a coefficient of 1 before the CH4:

1 CH|4| + O|2| \rightarrow CO|2| + H|2|O

Since there is one carbon on each side of the arrow, the first atom (carbon) is balanced.

Looking at the next atom (hydrogen), the right hand side has two atoms, while the left hand side has four. To balance the hydrogens, 2 goes in front of the H2O, which yields:

1 CH|4| + O|2| \rightarrow CO|2| + 2 H|2|O

Inspection of the last atom to be balanced (oxygen) shows that the right hand side has four atoms, while the left hand side has two. It can be balanced by putting a 2 before O2, giving the balanced equation:

CH|4| + 2 O|2| \rightarrow CO|2| + 2 H|2|O

This equation does not have any coefficients in front of CH4 and CO2, since a coefficient of 1 is dropped.

Ionic equations

An ionic equation is a chemical equation in which electrolytes are written as dissociated ions. Ionic equations are used for single and double displacement reactions that occur in aqueoussolutions. For example in the following precipitation reaction:

CaCl2(aq) + 2AgNO3(aq) \rightarrow Ca(NO3)2(aq) + 2AgCl(s)

the full ionic equation would be:

Ca2+ + 2Cl− + 2Ag+ + 2NO3− \rightarrow Ca2+ + 2NO3− + 2AgCl(s)

and the net ionic equation would be:

2Cl−(aq) + 2Ag+(aq) \rightarrow 2AgCl(s)

or, in reduced balanced form,

Ag+ + Cl− \rightarrow AgCl(s)

In this aqueous reaction the Ca2+ an

From Yahoo Answers

Question:Is there any easy way to do this or is it just a ridiculously large equation I need to use? The question is...Calculate the ionic strength for an aqueous solution with 0.002 mol/kg NaCO3 and next question says to estimate the mean ionic activity coefficient of the aqueous solution at 25 degrees C. I really dont want to do this but I must! If you answer, please explain, this is just a sample question to prepare me for a test so straight up answers are of no use to me! Ta:)

Answers:if ci and zi are the valence and the normality of ions ionic strength is IS= sum of 1/2 ci*zi^2 your formula is wrong it is Na2CO3 Na2CO3--> 2Na+ + CO3- z Na+=1 z CO3=2 cNa+ =0.004 c CO3 =0.002 IS= 1/2*(0.004*1+0.002*4)=0.006 For the next question, i must know the equilibrium constant of dissociation , sorry

Question:a) 2AgNO3 (aq) + Na2SO4 (aq) ---> How would i write the ionic and net ionic equation for this without knowing the charge for Ag (silver)? I'm really lost =[ can someone give me an explaination as well? that would be really helpful!

Answers:Ag is +1 and Na is also +1 (remember HNO3 where H is +1)

Question:Write a balanced formula equation, complete ionic equation, and net ionic equation for the reaction between: (Be sure to include phases.) a. an alkaline earth salt and sulfuric acid, being sure to identify the precipitate. b a halogen with a less active halide, being sure to identify which is oxidized and reduced.

Answers:(A) Among the alkaline earth metals you can chose the soluble salts of calcium, strontium or barium as one of the reactants, because the sulfates of these salts are slightly soluble. The least soluble one is barium sulfate and suppose we choose it as the product. All nitrates of metals are soluble, therefore we can choose barium nitrate as the reactant. Formula equation; Ba(NO3)2(aq) + H2SO4(aq) -------> BaSO4(s) + 2HNO3(aq) Ionic equation: Ba^2+(aq) + 2NO3^-(aq) + 2H^+(aq) + SO4^2-(aq) -------> BaSO4(s) + 2H^+(aq) + 2NO3^-(aq) Net ionic equation: (obtained by eliminating the spectator ions from both sides) Ba^2+(aq) + SO4^2-(aq) -------> BaSO4(aq) (B) Activity of halogens decreases from top to bottom within the group ( F > Cl > Br > I ) In the elemental state all halogens are diatomic molecules. F2 and Cl2 are gases, Br2 is liquid and I2 is solid. F2 replaces all other halogens. Cl2 replaces Br2 and I2. Br2 can only replace I2. Since I2 is the least active one it cannot replace any halogen. Formula equation; Cl2(g) + 2NaBr(aq) -------> 2NaCl(aq) + Br2(l) (note: all salts of sodium, potassium and ammonium are soluble) Ionic equation: Cl2(g) + 2Na^+(aq) + 2Br^-(aq) ------> 2Na^+(aq) + 2Cl^- (aq) + Br2(l) Net ionic equation: Cl2(g) + 2Br^-(aq) ------>2Cl^- (aq) + Br2(l) As it is clearly seen from the net ionic equation, Cl2 is reduced from 0 to -1 and Br^- is oxidized from -1 to 0.

Question:Write overall, total ionic and net ionic equations when two beakers containing the following solutions are mixed. Include phases for each species. A. Acid Base Reaction: hydrochloric acid and potassium hydroxide. B. Acid-Base Reaction: ammonia and nitric acid. C. Activity Series: Mn(s) and hydrochloric acid(aq). I'm pretty clueless as to where to begin with these... can you explain to me? Thank you!

Answers:A. Hydrochloric acid is a strong acid and potassium hydroxide is a strong base.. In a solution; - A strong acid completely ionizes. HCl(aq) + H2O(l) -----> H3O+(aq) + Cl-(aq) or HCl(aq) -----> H+(aq) + Cl-(aq) - A strong base completely dissociates. KOH(aq) --------> K+(aq) + OH-(aq) Strong acid - strong base reactions produces a slightly ionizable water. Overall: HCl(aq) + KOH(aq) -------> KCl(aq) + H2O(l) Total ionic: H+(aq) + Cl-(aq) + K+(aq) + OH-(aq) ----> K+(aq) + Cl-(aq) + H2O(l) Net ionic: H+(aq) + OH-(aq) ---->H2O(l) B. Ammonia is a weak base and nitric acid is a strong acid. In a solution; - A strong acid completely ionizes. HNO3(aq) -----> H+(aq) + NO3-(aq) - A weak base cannot ionize completely. NH3(aq) + H2O(l) <-----> NH4+(aq) + OH-(aq) Although this equation looks similar to the ionization of HCl, the remarkable difference is the shapes of the arrows representing the equations. -----> represents the complete ionization <----> represents the partial ionization (equilibrium reaction) The other difference is not visible, but it is a fact that the extent of the ionization cannot exceed 5%. Therefore, OH-(aq) cannot represent a weak base. Overall: NH3(aq) + HNO3(aq) ----> NH4NO3(aq) + H2O(l) Total ionic: NH3(aq) + H+(aq) + NO3-(aq) ---> NH4+(aq) + NO3-(aq) + H2O(l) Net ionic: NH3(aq) + H+(aq) + ---> NH4+(aq) + H2O(l) C. To complete such oxidation - reduction reactions, the activities of metals should be known. Manganese (Mn) is more active than H2. In other words, Mn(s) can reduce H+ to H2(g) (oxidation number = 0) or H+ can oxidize Mn(s) (oxidation number = 0) to Mn^2+(aq). Overall: Mn(s) + 2HCl(aq) -------> MnCl2(aq) + H2(g) Total ionic: Mn(s) + 2H^+(aq) + 2Cl^-(aq) -------> Mn^2+(aq) + 2Cl^-(aq) + H2(g) Net ionic: Mn(s) + 2H^+(aq) -------> Mn^2+(aq) + H2(g) I hope I am not late this time.

From Youtube

Net Ionic Equation :Free Science Help at Brightstorm! brightstorm.com How to write a net ionic equation.

Net Ionic Equations :You can see copies of the math at www.coolstudyguides.com We got a double replacement- We got the equation, it says find the net ionic of the precipitate, son. First we convert the whole thing into our ions. This way of writing, it's called complete ionic equation. Next we find the precipitate, we look on our 'lil sheet. It says that this one's insoluble, even in heat. We get rid of all the ions watching it occur, not participating or even precipitating, just being spectators. You've got it now. Word.