most common types of ternary compounds consist of a metallic cation (positive ion) and a polyatomic anion (negative ion).
Example
Steps:
1.Write the symbols for the monoatomic and polyatomic ions in the compound.
ex. Potassium = K Hydroxide = OH
2. Look up the oxidation numbers of the ions involved and write them as superscripts to the right of the elemental symbols.
ex. Potassium = K+ Hydroxide = OH-
3.Use the correct combination of ions to produce a compound with a net charge of zero. Parenthesis must be used if you need more than one of a polyatomic ion.
ex. KOH
Ternary Ionic Compound:
usually contain only a kind of metal cation and a kind of oxyanion.
Steps in Naming: 1.First word: the complete name of the less electronegative element. 2. Second word: consisting of the stem of the more electronegative element with an -ate or an -ite suffix. Either the prefix per- or the prefix hypo- may also be used to help indicate the ON of the nonmetal X. Greek prefixes are not used with ionic compounds. 3. Some groups have special names; Cations are named like a metal. Anions like a nonmetal are named with an -ide on the ending
Examples:
1
Ba(ClO3)2
K2SO4
bariumchlorate
potassiumsulfate
*Upon learning a list of ions, it is frequently easier to name an ionic compound by simply listing the ion names in the order given in the formula.
In 1704, Isaac Newton famously outlined his atomic bonding theory, in "Query 31" of his Opticks, whereby atoms attach to each other by some "force". Specifically, after acknowledging the various popular theories in vogue at the time, of how atoms were reasoned to attach to each other, i.e. "hooked atoms", "glued together by rest", or "stuck together by conspiring motions", Newton states that he would rather infer from their cohesion, that "particles attract one another by some force, which in immediate contact is exceedingly strong, at small distances performs the chemical operations, and reaches not far from the particles with any sensible effect."
In 1819, Jöns Jakob Berzelius developed a theory of chemical combination stressing the electronegative and electropositive character of the combining atoms. Mid-19th century, Edward Frankland, F.A. Kekule, A.S. Couper, A.M. Butlerov, and Hermann Kolbe, building on the theory of radicals, developed the theory of valency, originally called "combining power", in which compounds were joined owing to an attraction of positive and negative poles. In 1916, chemist Gilbert N. Lewis developed the concept of the electron-pair bond, in which two atoms may share one to six electrons, thus forming the single electron bond, a single bond, a double bond, or a triple bond; in Lewis's own words, "An electron may form a part of the shell of two different atoms and cannot be said to belong to either one exclusively."
Polar Covalent Bond
The above figure is an illustrative representation of a water molecule having polar covalent bonds between the Oxygen atom and the Hydrogen atoms.
In a polar covalent bond, the electrons shared by the atoms spend a greater amount of time, on the average, closer to the Oxygen nucleus than the Hydrogen nucleus. This is because of the geometry of the molecule and the great electron negativity difference between the Hydrogen atom and the Oxygen atom.
The result of this pattern of unequal electron association is a charge separation in the molecule, where one part of the molecule, the Oxygen, has a partial negative charge and the Hydrogen have a partial positive charge.
A polar bond is formed whenelectronsare unequally shared between two atoms. Polar covalent bonding occurs because one atom has a stronger affinity for electrons than the other. The bonding electrons will spend a greater amount of time around the atom that has the stronger affinity for electrons.
Non Polar Covalent Bond
The H2moleculeis a good example of the first type of covalent bond, the non polar bond. Because bothatomsin the H2molecule have an equal attraction or affinity forelectrons, the bonding electrons are equally shared by the two atoms and whenever two atoms of the sameelementbond together, a non polar bond is formed.
Non-Polar bondingresults when twoidentical non-metals equally shareelectrons between them. One well known exception to the identical atom rule is the combination of carbon and hydrogen in all organic compounds.
In chemistry, an ionic compound is a chemical compound in which ions are held together in a lattice structure by ionic bonds. Usually, the positively charged portion consists of metal cations and the negatively charged portion is an anion or polyatomic ion. Ions in ionic compounds are held together by the electrostatic force between oppositely charged bodies. Ionic compounds have a high melting and boiling point, and they are hard and very brittle.
Positive and negative ions are attracted to each other to make ionic compounds. They easily form crystals. They tend to have high boiling and melting points. They conduct electricity when they dissolve in water.
Ions can be single atoms, as the sodium and chloride in common table salt sodium chloride, or more complex groups such as the carbonate in calcium carbonate. But to be considered an ion, they must carry a positive or negative charge. Thus, in an ionic bond, one 'bonder' must have a positive charge and the other a negative one. By sticking to each other, they resolve, or partially resolve, their separate charge imbalances. Positive to positive and negative to negative ionic bonds do not occur. (For an easily visible analogy, experiment with a pair of bar magnets.)
Ionic compounds are usually formed when metal cations bond with nonmetal anions. The only common exception I know to this is when ammonium is the cation - there's no metal in ammonium, but it forms ionic compounds anyhow.
Formation of Ionic Compounds
This is a tutorial on writing ionic compounds (by MarkResengarten on YouTube)
Whenever a chemical reaction occurs, the changes that people observe are caused by the creation or loss of certain types of materials. For instance, temperature can generate a chemical reaction. An easy way to recognize a chemical change is to compare the color of the original item with the new one. A number of chemical reactions cause color changes.
In chemistry, a colour reaction is a chemical reaction that is used to transform colourless chemical compounds into coloured derivatives which can be detected visually or with the aid of a colourimeter.
The concentration of a colourless solution cannot normally be determined with a colourimeter. The addition of a colour reagent leads to a colour reaction and the absorbance of the coloured product can then be measured with a colourimeter.
Colourimeter
A change in absorbance in the ultraviolet range cannot be detected by eye but can be measured by a suitably-equipped colourimeter. A special colourimeter is required because standard colourimeters cannot operate below a wavelength of 400 nanometes. It is also necessary to use fused quartz cuvettes because glass is opaque to ultraviolet.
Many different colour reagents have been developed for determining the concentrations of different substances. For example, Nessler's reagent can be used used to determine the concentration of a solution of ammonia.
purple (MnO4- ions) // blue (MnO4 3- ions) // green (MnO4 2- ions) // orange (Mn 3+ ions)
The secret of this color change is pH. Chemicals with a low pH (0-6) are acidic, while those with a high pH (8-14) are basic. (A pH of 7 is neutral: neither acidic nor basic.) Universal indicator is a chemical that changes color in the presence of acids and bases from a pH of 2 to 10. Acids turn the indicator red, pink, orange, and yellow, while bases turn it green, blue, and purple. Vinegar is an acid, so when you poured the indicator solution into the second flask, it turned red. Ammonia is a base, so when you mixed the acidic vinegar solution with ammonia, it raised the pH and the water turned blue. If you had enough vinegar in your last flask, the solution should have turned red again.
Chemistry Magic : Confounding Color (by homesciencetools.com) A video on Color Reaction
For thousands of years people have known that vinegar, lemon juice and many other foods taste sour. However, it was not until a few hundred years ago that it was discovered why these things taste sour - because they are all acids. The term acid, in fact, comes from the Latin term acere, which means "sour". While there are many slightly different definitions of acids and bases, in this lesson we will introduce the fundamentals of acid/base chemistry.
In the seventeenth century, the Irish writer and amateur chemist Robert Boyle first labeled substances as either acids or bases (he called bases alkalies) according to the following characteristics:
Acidstaste sour, are corrosive to metals, change litmus (a dye extracted from lichens) red, and become lessacidic when mixed with bases.
Basesfeel slippery, change litmus blue, and become less basic when mixed with acids.
While Boyle and others tried to explain why acids and bases behave the way they do, the first reasonable definition of acids and bases would not be proposed until 200 years later.
In the late 1800s, the Swedish scientist Svante Arrhenius proposed that water can dissolve many compounds by separating them into their individual ions. Arrhenius suggested that acids are compounds that contain hydrogen and can dissolve in water to release hydrogen ions into solution.
Arrhenius defined bases as substances that dissolve in water to release hydroxide ions (OH-) into solutions.
The Arrhenius definition of acids and bases explains a number of things. Arrhenius's theory explains why all acids have similar properties to each other (and, conversely, why all bases are similar): because all acids release H+ into solution (and all bases release OH-). The Arrhenius definition also explains Boyle's observation that acids and bases counteract each other. This idea, that a base can make an acid weaker, and vice versa, is called neutralization.
Neutralization
As you can see from the equations, acids release H+ into solution and bases release OH-. If we were to mix an acid and base together, the H+ ion would combine with the OH- ion to make the molecule H2O, or plain water:
H+(aq)
+
OH-(aq)
=
H2O
The neutralization reaction of an acid with a base will always produce water and a salt, as shown below:
Acid
Base
Water
Salt
HCl
+
NaOH
=
H2O
+
NaCl
HBr
+
KOH
=
H2O
+
KBr
Though Arrhenius helped explain the fundamentals of acid/base chemistry, unfortunately his theories have limits. For example, the Arrhenius definition does not explain why some substances, such as common baking soda (NaHCO3), can act like a base even though they do not contain hydroxide ions.
In 1923, the Danish scientist Johannes Brønsted and the Englishman Thomas Lowry published independent yet similar papers that refined Arrhenius' theory. In Brønsted's words, "... acids and bases are substances thatare capable of splitting off or taking up hydrogen ions, respectively." The Brønsted-Lowry definition broadened the Arrhenius concept of acids and bases.
The Brønsted-Lowry definition of acid is very similar to the Arrhenius definition, any substance that can donate a hydrogen ion is an acid (under the Brønsted definition, acids are often referred to as proton donors because an H+ ion, hydrogen minus its electron, is simply a proton).
The Brønsted definition of bases is, however, quite different from the Arrhenius definition. The Brønsted bases is defined as any substance that can accept a hydrogen ion. In essence, a base is the opposite of an acid. NaOH and KOH, as we saw above, would still be considered bases because they can accept an H+ from an acid to form water. However, the Brønsted-Lowry definition also explains why substances that do not contain OH- can act like bases. Baking soda (NaHCO3), for example, acts like a base by accepting a hydrogen ion from an acid as illustrated below:
Acid
Base
Salt
HCl
+
NaHCO3
=
H2CO3
+
NaCl
In this example, the carbonic acid formed (H2CO3) undergoes rapid decomposition to water and gaseous carbon dioxide, and so the solution bubbles as CO2 gas is released.
pH
Under the Brønsted-Lowry definition, both acids and bases are related to the concentration of hydrogen ionspresent. Acids increase the concentration of hydrogen ions, while bases decrease the concentration of hydrogen ions (by accepting them). The acidity or basicity of something, therefore, can be measured by its hydrogen ion concentration.
In 1909, the Danish biochemist Sören Sörensen invented the pH scale for measuring acidity. The pH scale is described by the formula:
pH = -log [H+]
Note: concentration is commonly abbreviated by using square brackets, thus [H+] = hydrogen ion concentration. When measuring pH, [H+] is in units of moles of H+ per liter of solution.
For example, a solution with [H+] = 1 x 10-7 moles/liter has a pH equal to 7 (a simpler way to think about pH is that it equals the exponent on the H+ concentration, ignoring the minus sign). The pH scale ranges from 0 to 14. Substances with a pH between 0 and less than 7 are acids (pH and [H+] are inversely related - lower p H means higher [H+]). Substances with a pH greater than 7 and up to 14 are bases (higher pH means lower [H+]). Right in the middle, at pH = 7, are neutral substances, for example, pure water. The relationship between [H+] and pH is shown in the table below alongside some common examples of acids and bases in everyday life.
[H+]
pH
Example
Acids
1 X 100
0
HCl
1 x 10-1
1
Stomach acid
1 x 10-2
2
Lemon juice
1 x 10-3
3
Vinegar
1 x 10-4
4
Soda
1 x 10-5
5
Rainwater
1 x 10-6
6
Milk
Neutral
1 x 10-7
7
Pure water
Bases
1 x 10-8
8
Egg whites
1 x 10-9
9
Baking soda
1 x 10-10
10
Tums® antacid
1 x 10-11
11
Ammonia
1 x 10-12
12
Mineral lime - Ca(OH)2
1 x 10-13
13
Drano®
1 x 10-14
14
NaOH
Acids: acid, is a substance which reacts with a base. Commonly, acids can be identified as tasting sour, reacting withmetals such as calcium, and bases like sodium carbonate. Aqueous acids have pHs of less than 7, and turn blue litmus paper red. Chemicals or substances having the property of an acid are said to be acidic.
There are three common definitions for acids: the Arrhenius definition, the Brønsted-Lowry definition, and the Lewis definition. The Arrhenius definition states that acids are substances which increase the concentration of hydronium ions (H3O+) in solution. The Brønsted-Lowry definition is an expansion: an acid is a substance which can act as a proton donor. Most acids encountered in everyday life are aqueous solutions, or can be dissolved in water, and these two definitions are most relevant.
PROPERTIES OF ACIDS
Corrosive ('burns' your skin)
Sour taste (e.g. lemons, vinegar)
Contains hydrogen ions (H+) when dissolved in water
Has a pH less than 7
Turns blue litmus paper to a red color
Reacts with bases to form salt and water
Reacts with metals to form hydrogen gas
Reacts with carbonates to form carbon dioxide, water and a salt
EXAMPLES OF ACIDS
Hydrochloric acid (HCl) in gastric juice
Sulphuric acid (H2SO4)
Nitric acid (HNO3)
Carbonic acid in softdrink (H2CO3)
Uric acid in urine
Ascorbic acid (Vitamin C) in fruit
Citric acid in oranges and lemons
Acetic acid in vinegar
Tannic acid (in tea and wine)
Tartaric acid (in grapes)
Base: A base in chemistry is a substance that can accept hydrogen ions or more generally, donate electron pairs. The Brønsted-Lowry theory defines bases as proton (hydrogen ion) acceptors, while the more general Lewis theory defines bases as electron pair donors, allowing other Lewis acids than protons to be included
Bases can be thought of as the chemical opposite of acids. A reaction between an acid and base is called neutralization. Bases and acids are seen as opposites because the effect of an acid is to increase the hydronium ion (H3O+) concentration in water, whereas bases reduce this concentration. Bases and acids are typically found in aqueous solution forms. Aqueous solutions of bases react with aqueous solutions of acids to produce water and salts in aqueous solutions in which the salts separate into their component ions. If the aqueous solution is a saturated solution with respect to a given salt solute any additional such salt present in the solution will result in formation of a precipitate of the salt.
PROPERTIES OF BASES AND ALKALIS
Corrosive ('burns' your skin)
Soapy feel
Has a pH more than 7
Turns red litmus paper to a blue colour
Many alkalis (soluble bases) contain hydroxyl ions (OH-)
Reacts with acids to form salt and water
EXAMPLES OF BASES AND ALKALIS
Sodium hydroxide (NaOH) or caustic soda
Calcium hydroxide ( Ca(OH)2 ) or limewater
Ammonium hydroxide (NH4OH) or ammonia water
Magnesium hydroxide ( Mg(OH)2 ) or milk of magnesia
Many bleaches, soaps, toothpastes and cleaning agents
Common Acids
pH
Common Bases
pH
Hydrochloric acid
Sulphuric acid
Stomach juice
Lemons
Vinegar
Apples
Oranges
Grapes
Sour milk
White bread
Fresh milk
0.1
0.3
1-3
2.3
2.9
3.1
3.5
4
4.4
5.5
6.5
Human saliva
Distilled water
Blood plasma
Eggs
Seawater
Borax
Milk of magnesia
Ammonia water
Limewater
Caustic soda
6-8
7
7.4
7.8
7.9
9.2
10.5
11.6
12.4
14
INDICATORS
An indicator, when added to an acid, a neutral substance or a base, will change different colours.
