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[Gen Chem]

Compounds and Bonding

Overview of Bonding

Introduction to Bonding

Put simply, chemical bonding join atoms together to form more complex structures (like molecules or crystals). Bonds can form between atoms of the same element, or between atoms of different elements. There are several types of chemical bonding which have different properties and give rise to different structures.

In many cases, atoms try to react to form valence shells containing eight electrons. The octet rule describes this, but it also has many exceptions

• Ionic bonding occurs between positive ions (cations) and negative ions (anions). In an ionic solid, the ions arrange themselves into a rigid crystal lattice. NaCl (common salt) is an example of an ionic substance. In ionic bonding there is an attractive force established between large numbers of positive cations and negative anions, such that a neutral lattice is formed. This attraction between oppositely-charged ions is collective in nature and called ionic bonding.

• Covalent bonds are formed when the orbitals of two non-metal atoms physically overlap and share electrons with each other. There are two types of structures to which this can give rise: molecules and covalent network solids. Methane (CH₄) and water (H₂O) are examples of covalently bonded molecules, and glass is a covalent network solid.

• Metallic bonding occur between atoms that have few electrons compared to the number of accessible orbitals. This is true for the vast majority of chemical elements. In a metallically bonded substance, the atoms' outer electrons are able to freely move around - they are delocalised to form an 'electron pool'. Iron is a metallically bonded substance.

Chemical bonding is one of the most crucial concepts in the study of chemistry. In fact, the properties of materials are basically defined by the type and number of atoms they contain and how they are bonded together.

So far, you have seen examples of intramolecular bonds. These bonds connect atoms into molecules or whole crystals. There are also intermolecular bonds that connect molecules into large substances. These are also called intermolecular forces, or IMF. IMF are weaker than intramolecular bonds, and as they do not permanently join two molecules or ions, it is generally considered incorrect to refer to them as bonds. Sometimes, a substance will not have both IMF and intramolecular bonds. In the case of ionic crystals (like salt) or covalent networks (like diamond), the solid is made out of a network of intramolecular bonds connecting all the component atoms or ions in a repeating pattern, with no separate units to be attracted to each other by IMF. In the case of metallic bonding, the atoms are all interconnected into one large piece of metal, but the electrons move freely rather than being confined to the static bonds of a crystal lattice or covalent network.

Electronegativity

What determines the type of bond formed between two elements? There are two ways of classifying elements to determine the bond formed: by electronegativity, or by metallic/non-metallic character.

Electronegativity

Electronegativity is a property of atoms that is reflected in the layout of the periodic table of the elements. Electronegativity is greatest in the elements in the upper right of the table (e.g., fluorine), and lowest in the lower left (e.g., francium).

Electronegativity is a relative measure of how strongly an atom will attract the electrons in a bond. Although bonds are the result of atoms sharing their electrons, the electrons can be shared unequally. The more electronegative atom in a bond will have a slight negative charge, and the less electronegative atom will have a slight positive charge. Overall, the molecule may have no charge, but the individual atoms will. This is a result of the electronegativity—by attracting the electrons in a bond, an atom gains a slight negative charge. Of course, if two elements have equal electronegativity, they will share the electrons equally.

Linus Pauling created a commonly-used measure of electronegativity.

Metallic elements have low electronegativity, and non-metallic elements have high electronegativity. If two elements are close to each other on the periodic table, they will have similar electronegativities.

Electronegativity is measured on a variety of scales, the most common being the Pauling scale. Created by chemist Linus Pauling, it assigns 4.0 to fluorine (the highest) and 0.7 to francium (the lowest).

Notation

When drawing diagrams of bonds, we indicate covalent bonds with a line. We may write the electronegativity using the symbols $\delta +$ and $\delta -$. Look at this example.

Hydrogen fluoride (HF): ${\begin{matrix}\delta +&&\delta -\\H&-&F\end{matrix}}$

The plus goes over the less electronegative atom. From the above diagram, we can see that the fluorine attracts the electrons in the covalent bond more than the hydrogen does. Fluorine will have a slight negative charge because of this, and hydrogen will have a slight positive charge. Overall, hydrogen fluoride is neutral.

Ionic Bonds

What are ions?

Ions are atoms or molecules which are electrically charged. Cations are positively charged and anions carry a negative charge. Ions form when atoms gain or lose electrons. Since electrons are negatively charged, an atom that loses one or more electrons will become positively charged; an atom that gains one or more electrons becomes negatively charged.

Description of Ionic Bonding

Ionic bonding is the attraction between positively- and negatively-charged ions. These oppositely charged ions attract each other to form ionic networks (or lattices). Electrostatics explains why this happens: opposite charges attract and like charges repel. When many ions attract each other, they form large, ordered, crystal lattices in which each ion is surrounded by ions of the opposite charge. Generally, when metals react with non-metals, electrons are transferred from the metals to the non-metals (this is because metals tend to have relatively low ionization energy and non-metals tend to have high electron affinity). The metals form positively-charged ions and the non-metals form negatively-charged ions. The smallest unit of an ionic compound is the formula unit, but this unit merely reflects that ratio of ions that leads to neutrality of the whole crystal, e.g. NaCl or MgCl₂. One cannot distinguish individual NaCl or MgCl₂ molecules in the structure.

It is however possible that the stacking consists of molecular ions like NH₄⁺ and NO₃^- in ammonium nitrate. In such structures, the ions are charged molecules rather than charged atoms.

The ions arrange themselves into a lattice where each ion is surrounded by ions of the opposite type.

An example of both atomic (Li+) and molecular (NO3-) ions

Ionic bonding may also be referred to as electrovalent bonding.

Characteristics

Example ionic compounds: Sodium chloride ($NaCl$), potassium nitrate ($KNO_{3}$).

Ionically bonded substances typically have the following characteristics.

• High melting point (solid at room temperature)
• Hard but brittle (can shatter)
• Many dissolve in water
• Conductors of electricity when dissolved or melted

In general the forces keeping the lattice together depend on the product of the charges of the ions it consists of. A comparison e.g. of NaCl (+1)*(-1) to MgO (+2)*(-2) shows that magnesium oxide is kept together much more strongly -roughly 4 times- than sodium chloride. This is why sodium chloride has a much lower melting point and also dissolves much more easily in a solvent like water than magnesium oxide does.

Solubility

Ionic compound dissolved in a polar solvent, in this case water. (not to scale)

Most ionic compounds are soluble in polar solvents. Polar solvents are simply solvents that have molecules with positive and negative poles. The most common example of a polar solvent is water (H₂O).

The reason that they are soluble in polar solvents is because the positive ions are attracted to the negative pole of the solvent and the negative ions are attracted to the positive pole of the solvent. Due to this, the ionic compound dissociates or splits in the solvent into its constituent ions.

For example, H₂O is composed of relatively more positive hydrogen atoms and relatively more negative oxygen atoms. When an ionic compound like NaCl is added to the water, the positive ions will be attracted to the more negative oxygen atoms and negative ions will be attracted to the more positive hydrogen atoms.

This is a very useful property of many ionic compounds, but not all, as there are exceptions where certain compounds cannot be dissolved in certain polar solvents.

Conduction

For a substance to conduct electricity, there must be some carrier of charge. For example, in metals, what transports the charge is the mobile electrons surrounding the positive ions.

However, in the solid state, ionic compounds lack any mobile particle capable of carrying charge as the ions cannot move within the lattice because the electrostatic force holding them in place is extremely strong. This changes when the compound is dissolved or molten. When dissolved, the ions dissociate and are capable of moving around. The same is true for molten ionic compounds. Which is why molten and dissolved ionic compounds can conduct electricity, while solid ionic compounds cannot.

Formation

The electron transfer between Na and Cl.

Ionic bonding occurs when metals and non-metals chemically react. As a result of its low ionization energy, a metal atom is not destabilized very much if it loses electrons to form a complete valence shell and becomes positively charged. As its affinity is rather large, a non-metal is stabilized rather strongly by gaining electrons to complete its valence shell and become negatively charged. When metals and non-metals react, the metals lose electrons by transferring them to the non-metals, which gain them. The total process -a small loss plus a large gain- leads to a net lowering of the energy. Consequently, ions are formed, which instantly attract each other leading to ionic bonding.

For instance, in the reaction of Na (sodium) and Cl (chlorine), each Cl atom takes one electron from a Na atom. Therefore each Na becomes a Na⁺ cation and each Cl atom becomes a Cl^- anion. Due to their opposite charges, they attract each other and are joined by millions and millions of other ions to form an ionic lattice. The lattice energy that results from this massive collective stacking further stabilizes the new compound. The formula (ratio of positive to negative ions) in the lattice is NaCl, i.e. there are equal numbers of positive and negative charges ensuring neutrality.

The charges must balance because otherwise the repulsion between the majority charges would become prohibitive. In the case of magnesium chloride, the magnesium atom gives up two electrons to become stable. Note that it is in the second group, so it has two valence electrons. The chlorine atom can only accept one electron, so there must be two chlorine ions for each magnesium ion. Therefore, the formula for magnesium chloride is MgCl₂. If magnesium oxide were forming, the formula would be MgO because oxygen can accept both of magnesium's electrons.

Exercise for the reader

Try figuring out what the formula for magnesium nitride would be. Use the periodic table to help.

It should also be noted that some atoms can form more than one ion. This usually happens with the transition metals. For instance Fe (iron) can become Fe²⁺ (called iron(II) or -by an older name- ferrous). Fe can also become Fe³⁺ (called iron(III) or -sometimes still- ferric).

Common Ions

Ionic bonding typically occurs in reactions between a metal and non-metal, but there are also certain molecules called polyatomic ions that undergo ionic bonding. Within these polyatomic ions, there can be covalent (or polar) bonding, but as a unit it undergoes ionic bonding. There are countless polyatomic ions, but you should be familiar with the most common ones. You would be well advised to memorize these ions.

Name	Formula	Name	Formula
Ammonium	NH4+	Hydronium	H3O+
Peroxide	O22-	Hydroxide	OH-
Nitrite	NO2-	Nitrate	NO3-
Sulfite	SO32-	Sulfate	SO42-
Hydrogen sulfite	HSO3-	Phosphate	PO43-
Hypochlorite	ClO-	Chlorite	ClO2-
Chlorate	ClO3-	Perchlorate	ClO4-
Carbonate	CO32-	Hydrogen carbonate	HCO3-

Covalent Bonds

Covalent bonds create molecules, which can be represented by a molecular formula. For chemicals such as a basic sugar (C₆H₁₂O₆), the ratios of atoms have a common multiple, and thus the empirical formula is CH₂O. Note that a molecule with a certain empirical formula is not necessarily the same as one with the same molecular formula.

Formation of Covalent Bonds

Covalent bonds form between two atoms which have incomplete octets — that is, their outermost shells have fewer than eight electrons. They can share their electrons in a covalent bond. The simplest example is water (H₂O). Oxygen has six valence electrons (and needs eight) and the hydrogens have one electron each (and need two). The oxygen shares two of its electrons with the hydrogens, and the hydrogens share their electrons with the oxygen. The result is a covalent bond between the oxygen and each hydrogen. The oxygen has a complete octet and the hydrogens have the two electrons they each need.

When atoms move closer, their orbitals change shape, letting off energy. However, there is a limit to how close the atoms get to each other—too close, and the nuclei repel each other.

One way to think of this is a ball rolling down into a valley. It will settle at the lowest point. As a result of this potential energy "valley", there is a specific bond length for each type of bond. Also, there is a specific amount of energy, measured in kilojoules per mole (kJ/mol) that is required to break the bonds in one mole of the substance. Stronger bonds have a shorter bond length and a greater bond energy.

The Valence Bond Model

One useful model of covalent bonding is called the Valence Bond model. It states that covalent bonds form when atoms share electrons with each other in order to complete their valence (outer) electron shells. They are mainly formed between non-metals.

An example of a covalently bonded substance is hydrogen gas (H₂). A hydrogen atom on its own has one electron—it needs two to complete its valence shell. When two hydrogen atoms bond, each one shares its electron with the other so that the electrons move about both atoms instead of just one. Both atoms now have access to two electrons: they become a stable H₂ molecule joined by a single covalent bond.

Double and Triple Bonds

Covalent bonds can also form between other non-metals, for example chlorine. A chlorine atom has 7 electrons in its valence shell—it needs 8 to complete it. Two chlorine atoms can share 1 electron each to form a single covalent bond. They become a Cl₂ molecule.

Oxygen can also form covalent bonds, however, it needs a further 2 electrons to complete its valence shell (it has 6). Two oxygen atoms must share 2 electrons each to complete each other's shells, making a total of 4 shared electrons. Because twice as many electrons are shared, this is called a 'double covalent bond'. Double bonds are much stronger than single bonds, so the bond length is shorter and the bond energy is higher.

Furthermore,, nitrogen has 5 valence electrons (it needs a further 3). Two nitrogen atoms can share 3 electrons each to make a N₂ molecule joined by a 'triple covalent bond'. Triple bonds are stronger than double bonds. They have the shortest bond lengths and highest bond energies.

Electron Sharing and Orbitals

Carbon, contrary to the trend, does not share four electrons to make a quadruple bond. The reason for this is that the fourth pair of electrons in carbon cannot physically move close enough to be shared. The valence bond model explains this by considering the orbitals involved.

Recall that electrons orbit the nucleus within a cloud of electron density (orbitals). The valence bond model works on the principle that orbitals on different atoms must overlap to form a bond. There are several different ways that the orbitals can overlap, forming several distinct kinds of covalent bonds.

The Sigma Bond

The first and simplest kind of overlap is when two s orbitals come together. It is called a sigma bond (sigma, or σ, is the Greek equivalent of 's'). Sigma bonds can also form between two p orbitals that lie pointing towards each other. Whenever you see a single covalent bond, it exists as a sigma bond. When two atoms are joined by a sigma bond, they are held close to each other, but they are free to rotate like beads on a string.

The electron density is in between the two atoms in an σ bond.

The Pi Bond

The second, and equally important kind of overlap is between two parallel p orbitals. Instead of overlapping head-to-head (as in the sigma bond), they join side-to-side, forming two areas of electron density above and below the molecule. This type of overlap is referred to as a pi (π, from the Greek equivalent of p) bond. Whenever you see a double or triple covalent bond, it exists as one sigma bond and one or two pi bonds. Due to the side-by-side overlap of a pi bond, there is no way the atoms can twist around each other as in a sigma bond. Pi bonds give the molecule a rigid shape.

Pi bonds are weaker than sigma bonds since there is less overlap. Thus, two single bonds are stronger than a double bond, and more energy is needed to break two single bonds than a single double bond.

The electron density lies above and below the atoms in a π bond.

Hybridization

Consider a molecule of methane: a carbon atom attached to four hydrogen atoms. Each atom is satisfying the octet rule, and each bond is a single covalent bond.

Now look at the electron configuration of carbon: 1s²2s²2p². In its valence shell, it has two s electrons and two p electrons. It would not be possible for the four electrons to make equal bonds with the four hydrogen atoms (each of which has one s electron). We know, by measuring bond length and bond energy, that the four bonds in methane are equal, yet carbon has electrons in two different orbitals, which should overlap with the hydrogen 1s orbital in different ways.

To solve the problem, hybridization occurs. Instead of a s orbital and three p orbital, the orbitals mix, to form four orbitals, each with 25% s character and 75% p character. These hybrid orbitals are called sp³ orbitals, and they are identical. Observe:

$C\quad {\frac {\uparrow \downarrow }{1s}}\;\;{\frac {\uparrow \downarrow }{2s}}\;{\frac {\uparrow \,}{2p_{x}}}\;{\frac {\uparrow \,}{2p_{y}}}\;{\frac {\,\,}{2p_{z}}}$

$C^{*}\quad {\frac {\uparrow \downarrow }{1s}}\;\;{\frac {\uparrow \,}{sp^{3}}}{\frac {\uparrow \,}{sp^{3}}}{\frac {\uparrow \,}{sp^{3}}}{\frac {\uparrow \,}{sp^{3}}}$

Now these orbitals can overlap with hydrogen 1s orbitals to form four equal bonds. Hybridization may involve d orbitals in the atoms that have them, allowing up to a sp³d² hybridization.

Exercise for the reader

Predict the hybridized electron configuration of carbon in ethene. How many sigma bonds are there? How many pi bonds?

Hint: Hybridized electrons form only sigma bonds. Pi bonds form only between p electrons.

Metallic Bonds

Review on structure

Atoms are particles composed of 3 subatomic particles. The protons, neutrons, and electrons. The protons and neutrons are in the nucleus and the electrons orbit this nucleus. These electrons can leave the atom and this is important to note especially for metals. The energy required to remove an electron is referred to as ionization energy and is very low for metals and so metals tend to lose their outermost (valence) electrons whenever possible. The attractive force between electrons and the nucleus is referred to as the electrostatic force of attraction. This electrostatic force is very strong and is capable of holding large structures together.

Overview

Metallic bonds occur among metal atoms. Whereas ionic bonds join metals to non-metals, metallic bonding joins a bulk of metal atoms. A sheet of aluminium foil and a copper wire are both places where you can see metallic bonding in action.

The "sea of electrons" is free to flow about the crystal of positive metal ions.

When metallic bonds form, the s and p electrons delocalize. Instead of orbiting their atoms, they form a "sea of electrons" surrounding the positive metal ions. The electrons are free to move throughout the resulting network. The electrostatic force of attraction between this "sea of electrons" and the positive ions acts as a form of "glue" holding together the metallic structure.

Properties

The delocalized nature of the electrons explains a number of unique characteristics of metals:

Metals are good conductors of electricity	The sea of electrons is free to flow, allowing electrical currents.
Metals are good conductors of heat	As heat energy is defined by the motion of particles, the sea of electrons provides an excellent medium for this kinetic energy to be transferred.
Metals are ductile (able to draw into wires) and malleable (able to be hammered into thin sheets)	As the metal is deformed, local bonds are broken but quickly reformed in a new position.
Metals are gray and shiny	Photons (particles of light) cannot penetrate the metal, so they bounce off the sea of electrons.
Gold is yellow and copper is reddish-brown	There is actually an upper limit to the frequency that is reflected. It is too high to be visible in most metals, but not gold and copper.
Metals have very high melting and boiling points	Metallic bonding is very strong, so the atoms are reluctant to break apart into a liquid or gas.

Reactivity and solubility

Nearly all metals are near completely insoluble in most solvents whether organic or inorganic. This is because the bonds between metals are exceedingly strong and solvents cannot dissolve the metal. However, if the metal undergoes a reaction and forms an ionic compound it can dissolve. Metals react when an electronegative substance attracts the electrons in the sea of electrons and the structure collapses without the force holding it together. Metals almost always form ionic compounds however on occasion they can participate in a form of covalent bonding. These compounds however, do not share any of the properties of metallic bonding. Furthermore, it is important to distinguish between a compound of a metal and an alloy.

Alloys

Metallic bonds can occur between different elements. A mixture of two or more metals is called an alloy. This is what differentiates alloys and compounds containing metals, as alloys are only mixtures. Depending on the size of the atoms being mixed, there are two different kinds of alloys that can form:

• 
  Substitutional alloy
• 
  Interstitial alloy

It is also worth note that substitutional/interstitial alloys can exist when there are more than 2 metals are involved in the formation of an alloy.

The resulting mixture will have a combination of the properties of both metals involved. Interstitial alloys tend to be (with exceptions) stronger than their component metals. This is because the particles of varying size tend to prevent the various layers in a metal from sliding over each other smoothly. An example of this would be the difference between bronze and its components copper and tin. Copper on its own is a softer metal unfit for use as a weapon, while bronze was widely used during the Bronze Age for swords and armour. Alloys can also have various other benefits such as corrosion resistance, torsional strength, and much more.

Molecular Shape

Covalent molecules are bonded to other atoms by electron pairs. Being mutually negatively charged, the electron pairs repel the other electron pairs and attempt to move as far apart as possible in order to stabilize the molecule. This repulsion causes covalent molecules to have distinctive shapes, known as the molecule's molecular geometry. There are several different methods of determining molecular geometry. A scientific model, called the VSEPR (valence shell electron pair repulsion) model can be used to qualitatively predict the shapes of molecules. Within this model, the AXE method is used in determining molecular geometry by counting the numbers of electrons and bonds related to the center atom(s) of the molecule.

The VSEPR model is by no means a perfect model of molecular shape! It is simply a system which explains the known shapes of molecular geometry as discovered by experiment. This can allow us to predict the geometry of similar molecules, making it a fairly useful model. Modern methods of quantitatively calculating the most stable (lowest energy) shapes of molecules can take several hours of supercomputer time, and is the domain of computational chemistry.

Table of Geometries

Orbital Hybridization
	sp	sp2	sp3	sp3d	sp3d2
2 Groups	Linear	Bent	Bent	Linear
3 Groups		Trigonal Planar	Trigonal Pyramidal	T-Shaped
4 Groups			Tetrahedral	See-saw	Square Planar
5 Groups				Trigonal Bipyramidal	Square Pyramidal
6 Groups					Octahedral

The hybridization is determined by how many "things" are attached to the central atom. Those "things" can be other atoms or non-bonding pairs of electrons. The number of groups is how many atoms or electron pairs are bonded to the central atom. For example, methane (CH₄) is tetrahedral-shaped because the carbon is attached to four hydrogens. Ammonia (NH₃) is not trigonal planar, however. It is trigonal pyramidal because it is attached to four "things": the three hydrogens and a non-bonding pair of electrons (to fulfill nitrogen's octet).

Tetrahedral Shape

Consider a simple covalent molecule, methane (CH₄). Four hydrogen atoms surround a carbon atom in three-dimensional space. Each CH bond consists of one pair of electrons, and these pairs try to move as far away from each other as possible (due to electrostatic repulsion). You might think this would lead to a flat shape, with each hydrogen atom 90° apart. However, in three dimensions, there is a more efficient arrangement of the hydrogen atoms. If each hydrogen atom is at a corner of a tetrahedron centered around the carbon atom, they are separated by about cos^-1(-1/3) ≈ 109.5°—the maximum possible.

Hybridization

To align four orbitals in this tetrahedral shape requires the reformation of one s and three p orbitals into an sp³ orbital.

CH₄ is a tetrahedral molecule.

Lone Electron Pairs

The VSEPR model treats lone electron pairs in a similar way to bonding electrons. In ammonia (NH₃) for example, there are three hydrogen atoms and one lone pair of electrons surrounding the central nitrogen atom. Because there are four groups, ammonia has a tetrahedral shape but unlike methane, the angle between the hydrogen atoms is slightly smaller, 107.3°. This can be explained by theorizing that lone electron pairs take up more space physically than bonding pairs. This is a reasonable theory: in a bond, the electron pair is distributed over two atoms whereas a lone pair is only located on one. Because it is bigger, the lone pair forces the other electron pairs together.

The lone pair occupies more space than a bonding pair, decreasing the angles.

Testing this assumption with water provides further evidence. In water (H₂O) there are two hydrogen atoms and two lone pairs, again making four groups in total. The electron pairs repel each other into a tetrahedral shape. The angle between the hydrogen atoms is 104.5°, which is what we expect from our model. The two lone pairs both push the bonds closer together, giving a smaller angle than in ammonia.

Linear and Planar Shapes

Electron-Poor Atoms

In some molecules, there are less than four pairs of valence electrons. This occurs in electron deficient atoms such as boron and beryllium, which don't conform to the octet rule (they can have 6 and 4 valence electrons respectively). In boron trifluoride (BF₃), there are only three electron pairs which repel each other to form a flat plane. Each fluorine atom is separated by cos^-1(-1/2) = 120°. A different set of hybrid orbitals is formed in this molecule: the 2s and two 2p orbitals combine to form three sp² hybrid orbitals. The remaining p orbital is empty and sits above and below the plane of the molecule.

Beryllium, on the other hand, forms only two pairs of valence electrons. These repel each other at cos^-1(-1) = 180°, forming a linear molecule. An example is beryllium chloride, which has two chlorine atoms situated on opposite sides of a beryllium atom. This time, one 2s and one 2p orbital combine to form two sp hybrid orbitals. The two remaining p orbitals sit above and to the side of the beryllium atom (they are empty).

Bent vs. Linear

The non-bonding pair causes sp² hybridization, leading to a bent shape.

No non-bonding pairs causes sp hybridization, leading to a linear shape.

Some elements will have a bent shape, others have a linear shape. Both are attached to two groups, so it depends on how many non-bonding pairs the central atom has.

Take a look at sulfur dioxide (SO₂) and carbon dioxide (CO₂). Both have two oxygen atoms attached with double covalent bonds. Carbon dioxide is linear, and sulfur dioxide is bent. The difference is in their valence shells. Carbon has four valence electrons, sulfur has six. When they bond, carbon has no non-bonding pairs, but sulfur has one.

Five and Six Groups

Recall that some elements, especially sulfur and phosphorus, can bond with five or six groups. The hybridization is sp³d or sp³d², with a trigonal bipyramidal or octahedral shape respectively. When there are non bonding pairs, other shape can arise (see the above chart).

How The Shapes Look

The yellow groups are non-bonding electron pairs. The white groups are bonded atoms, and the pink is the central atom. This is referred to as the AXE method; A is the central atom, X's are bonded atoms, and E's are non-bonding electron pairs.

Molecules are not static; their bonds are continually twisting, stretching and bending. According to quantum theory, the energies of these bond movements are quantized, and this fact forms the basis of infrared spectroscopy, an important chemical tool in analyzing organic molecules.

• 
  linear
• 
  trigonal planar
• 
  bent
• 
  tetrahedral
• 
  trigonal pyramidal
• 
  bent
• 
  trigonal bipyramidal
• 
  seesaw
• 
  t-shape
• 
  octahedral
• 
  square pyramid
• 
  square planar

Intermolecular Bonds

Dipoles

The polar bonds are symmetric, but they don't point in opposite directions. The result is a dipole (positive pointing down).

Covalent bonds can be polar or non-polar, and so can the overall compound depending on its shape. When a bond is polar, it creates a dipole, a pair of charges (one positive and one negative). If they are arranged in a symmetrical shape, so that they point in opposite directions, they will cancel each other. For example, since the four hydrogens in methane (CH₄) are facing away from each other, there is no overall dipole and the molecule is non-polar. In ammonia (NH₃), however, there is a negative dipole at the nitrogen, due to the asymmetry caused by the non-bonding electron pair. The polarity of a compound determines its intermolecular bonding abilities.

Polar and Non-Polar Shapes

When a molecule has a linear, trigonal planar, tetrahedral, trigonal bipyramidal, or octahedral shape, it will be non-polar. These are the shapes that do not have non-bonding lone pairs. (e.g. Methane, CH₄) But if some bonds are polar while others are not, there will be an overall dipole, and the molecule will be polar (e.g. Chloroform, CHCl₃).

The other shapes (with non-bonding pairs) will be polar. (e.g. Water, H₂O) Unless, of course, all the covalent bonds are non-polar, in which case there would be no dipoles to begin with.

Dipole-Dipole Bonds

When two polar molecules are near each other, they will arrange themselves so that the negative and positive sides line up. There will be an attractive force holding the two molecules together, but it is not nearly as strong a force as the intramolecular bonds. This is how many types of molecules bond together to form large solids or liquids.

Dipole-dipole forces hold these two HCl molecules together.

Hydrogen Bonding

Certain chemicals with hydrogen in their chemical formula have a special type of intermolecular bond, called hydrogen bonds. Hydrogen bonds will occur when a hydrogen atom is attached to an oxygen, nitrogen, or fluorine atom. This is because there is a large electronegativity difference between hydrogen and fluorine, oxygen, and nitrogen. Thus, molecules such as $HF$, $H_{2}O$, $NH_{3}$ are extremely polar molecules with very strong dipole-dipole forces. As a result of the high electronegativities of fluorine, oxygen, and nitrogen, these elements will pull the electrons almost completely away from the hydrogen. The hydrogen becomes a bare proton sticking out from the molecule, and it will be strongly attracted to the negative side of any other polar molecules. Hydrogen bonding is an extreme type of dipole-dipole bonding. These forces are weaker than intramolecular bonds, but are much stronger than other intermolecular forces, causing these compounds to have high boiling points.

The dotted line represents a hydrogen bond.

Covalent Networks

A covalent network

Silicon dioxide forms a covalent network. Unlike carbon dioxide (with double bonds), silicon dioxide forms only single covalent bonds. As a result, the individual molecules covalently bond into a large network. These bonds are very strong (being covalent) and there is no distinction between individual molecules and the overall network. Covalent networks hold diamonds together. Diamonds are made of nothing but carbon, and so is soot. Unlike soot, diamonds have covalent networks, making them very hard and crystalline.

Van der Waals forces

Van der Waals, or London dispersion forces are caused by temporary dipoles created when electron locations are lopsided. The electrons are constantly orbiting the nucleus, and by chance they could end up close together. The uneven concentration of electrons could make one side of the atom more negatively-charged than the other, creating a temporary dipole. As there are more electrons in an atom, and the shells are farther away from the nucleus, these forces become stronger.

Van der Waals forces explain how nitrogen can be liquified. Nitrogen gas is diatomic; its equation is N₂. Since both atoms have the same electronegativity, there is no dipole and the molecule is non-polar. If there are no dipoles, what would make the nitrogen atoms stick together to form a liquid? Van der Waals forces are the answer. They allow otherwise non-polar molecules to have attractive forces. These are by far the weakest forces that hold molecules together.

Melting and Boiling Points

When comparing two substances, their melting and boiling points may be questioned. To determine which substance has the higher melting or boiling point, you must decide which one has the strongest intermolecular force. Metallic bonds, ionic bonds, and covalent networks are very strong, as they are actually intramolecular forces. These substances have the highest melting and boiling points because they only separate into individual molecules when the powerful bonds have been broken. Breaking these intramolecular forces requires great amounts of heat energy.

Substances with hydrogen bonding, an intermolecular force, will have much higher melting and boiling points than those that have ordinary dipole-dipole intramolecular forces. Non-polar molecules have the lowest melting and boiling points, because they are held together by the weak van der Waals forces.

If you need to compare the boiling points of two metals, the metal with the larger atomic radius will have weaker bonding, due to the lower concentration of charge. When comparing boiling points of the non-polar gases, like the noble gases, the gas with the largest radius will have the highest points because it has the most potential for van der Waals forces.

Ionic compounds can be compared using Coulomb's Law. Look for substances with high ionic charges and low ionic radii.

Chemical Reactions

Naming Substances

Substances with carbon and hydrogen are organic compounds. They have special names that are beyond the scope of this book. For more information, see the Organic Chemistry Wikibook.

Some compounds have common names, like water for H₂O. However, there are thousands of other compounds that are uncommon or have multiple names. Also, the common name is usually not recognized internationally. What looks like water to you might look like agua or vatten to someone else. To allow chemists to communicate without confusion, there are naming conventions to determine the systematic name of a chemical.

Naming Ions and Ionic Compounds

Ions are atoms that have lost or gained electrons. Note that in a polyatomic ion, the ion itself is held together by covalent bonds. Monoatomic cations (positive) are named the same way as their element, and they come first when naming a compound. Monoatomic anions (negative) have the suffix -ide and come at the end of the compound's name.

Examples of ionic compounds

• NaCl - Sodium chloride
• MgCl₂ - Magnesium chloride
• Ca₃N₂ - Calcium nitride

Notice that there is no need to write how many ions there are. Between the periodic table and our knowledge of ionic bonding, we can determine the number of ions, based on which elements are used.

Polyatomic Ions

Polyatomic ions have special names. Many of them contain oxygen and are called oxyanions. When different oxyanions are made of the same element, but have a different number of oxygen atoms, then prefixes and suffixes are used to tell them apart. The chlorine family of ions is an excellent example.

Name	Formula
Chloride	Cl-
Hypochlorite	ClO-
Chlorite	ClO2-
Chlorate	ClO3-
Perchlorate	ClO4-

The -ate suffix is used on the most common oxyanion (like sulfate SO₄^2- or nitrate NO₃^-). The -ite suffix is used on the oxyanion with one oxygen atom fewer (like sulfite SO₃^2- or nitrite NO₂^-). Sometimes there can be a hypo- prefix, meaning one oxygen atom fewer than for -ite. There is also a per- prefix, meaning one more oxygen atom than an -ate molecule has.

One last prefix you may find is thio-. It means an oxygen has been replaced with a sulfur within the oxyanion. Cyanate is OCN^-, and thiocyanate is SCN^-.

Examples of polyatomic ions

• NH₄Cl - Ammonium chloride
• K(HCO₃) - Potassium hydrogen carbonate
• AgNO₃ - Silver nitrate
• CuSO₃ - Copper (II) sulfite

In the last example, copper had a roman numeral 2 after its name because most of the transition metals can have more than one charge. The charge on the ion must be known, so it is written out for ions that have more than one common charge. Silver always has a charge of 1+, so it isn't necessary (but not wrong) to name its charge. Zinc always has a charge of 2+, so you don't have to name its charge either. Aluminum will always have a charge of +3. All other metals (except the Group 1 and 2 elements) must have roman numerals to show their charge.

Common polyatomic ions that you should know are listed in the following table

Name	Formula
Acetate	C2H3O2-
Ammonium	NH4+
Carbonate	CO32-
Cyanide	CN-
Cyanate	CNO-
Thiocyanate	CNS-
Hypochlorite	ClO-
Chlorite	ClO2-
Chlorate	ClO3-
Perchlorate	ClO4-
Hypobromite	BrO-
Bromite	BrO2-
Bromate	BrO3-
Perbromate	BrO4-
Hypoiodite	IO-
Iodite	IO2-
Iodate	IO3-
Periodate	IO4-
Nitrite	NO2-
Nitrate	NO3-
Peroxide	O22-
Permanganate	MnO4-
Sulfite	SO32-
Sulfate	SO42-

In older texts, ions were assigned names based on their Latin root and a suffix. Common ions with this naming system include "plumbous/plumbic" for lead(II)/lead(IV) and "ferrous/ferric" for iron(II)/iron(III). These Latin-based names are outdated, so it's not important to learn them. We now use the Stock system instead.

Further explanation of the roman numerals is in order. Many atoms (especially the transition metals) are capable of ionizing in more than one way. The name of an ionic compound must make it very clear what the exact chemical formula is. If you wrote "copper chloride", it could be CuCl or CuCl₂ because copper can lose one or two electrons when it forms an ion. The charge must be balanced, so there would be one or two chloride ions to accept the electrons. To be correct, you must write "copper(II) chloride" if you want CuCl₂ and "copper (I) chloride" if you want CuCl. Keep in mind that the roman numerals refer to the charge of the cation, not how many anions are attached.

Common metal ions are listed below and should be learned:

Name	Formula
Iron(II)/Ferrous	Fe2+
Iron(III)/Ferric	Fe3+
Copper(I)/Cuprous	Cu+
Copper(II)/Cupric	Cu2+
Tin (II)/Stannous	Sn2+
Tin (IV)/Stannic	Sn4+
Lead (II)/Plumbous (most common)	Pb2+
Lead (IV)/Plumbic	Pb4+
Mercury (I) (Note: Mercury (I) is a polyatomic ion)	Hg22+
Mercury (II)	Hg2+

Naming molecules in chemistry

There are two systems of naming molecular compounds. The first uses prefixes to indicate the number of atoms of an element that are in the compound. If the substance is binary (containing only two elements), the suffix -ide is added to the second element. Thus water is dihydrogen monoxide. A prefix is not necessary for the first element if there is only one, so SF₆ is 'sulfur hexafluoride'. The prefix system is used when both elements are non-metallic.

Number	Prefix
1	Mono-
2	Di-
3	Tri-
4	Tetra-
5	Penta-
6	Hexa-
7	Hepta-
8	Octa-
9	Nona-
10	Deca-
11	Undeca-
12	Dodeca-

The second system, the stock system, uses oxidation numbers to represent how the electrons are distributed through the compound. This is essentially the roman numeral system that has already been explained, but it applies to non-ionic compounds as well. The most electronegative component of the molecule has a negative oxidation number that depends on the number of pairs of electrons it shares. The less electronegative part is assigned a positive number. In the stock system, only the cation's number is written, and in Roman numerals. The stock system is used when there is a metallic element in the compound. In the case of V₂O₅, it could also be called vanadium(V) oxide. Knowing that oxygen's charge is always -2, we could determine that there were five oxygens and two vanadiums, if we were given the name without the formula.

Naming acids

If an acid is a binary compound, it is named as hydro[element]ic acid. If it contains a polyatomic ion, then it is named [ion name]ic acid if the ion ends in -ate. If the ion ends in -ite then the acid will end in -ous. These examples should help.

Examples of acid names

• HCl - Hydrochloric acid
• HClO - Hypochlorous acid
• HClO₂ - Chlorous acid
• HClO₃ - Chloric acid
• HClO₄ - Perchloric acid

Formulas and Numbers

Calculating Formula Masses

In molecules but not ionic compounds, the formula mass is also known as the molecular mass.

The calculation of a compound's formula mass (the mass of its molecule or formula unit) is straightforward. Simply add the individual mass of each atom in the compound (found on the periodic table). For example, the formula mass of glucose (C₆H₁₂O₆) is 180 amu.

Molar masses are just as easy to calculate. The molar mass is equal to the formula mass, except that the unit is grams per mole instead of amu.

Calculating Percentage Composition

Percentage composition is the relative mass of one substance in a compound compared to the whole. For example, in methane (CH₄), the percentage mass of hydrogen is 25% because hydrogen makes up a total of 4 amu out of 16 amu overall.

Using Percentage Composition

Percentage composition can be used to find the empirical formula of a compound, which shows the ratios of elements in the compound. However, this is not the same as the molecular formula. For example, many sugars have the empirical formula CH₂O, which could correspond to a molecular formula of CH₂O, C₂H₄O₂, C₆H₁₂O₆, etc.

To find the empirical formula from percentage composition, follow these procedures for each element.

1. Convert from percentage to grams (for simplicity, assume a 100 g sample).
2. Divide by the element's molar mass to find moles.
3. Simplify to lowest whole-number ratio.

For example, a compound is composed of 75% carbon and 25% hydrogen by mass. Find the empirical formula.

• 75g C / (12 g/mol C) = 6.25 mol C
• 25g H / (1 g/mol H) = 25 mol H
• 6.25 mol C / 6.25 = 1 mol C
• 25 mol H / 6.25 = 4 mol H

Thus the empirical formula is CH₄.

Calculating Molecular Formula

If you find the empirical formula of a compound and its molar/molecular mass, then you can find its exact molecular formula. Remember that the molecular formula is always a whole-number multiple of the empirical formula. For example, a compound with the empirical formula HO has a molecular mass of 34.0 amu. Since HO would only be 17.0 amu, which is half of 34.0, the molecular formula must be H₂O₂.

Exercise for the reader

An unknown substance must be identified. Lab analysis has found that the substance is composed of 80% Fluorine and 20% Nitrogen with a molecular mass of 71 amu. What is the empirical formula? What is the molecular formula?

Stoichiometry

The word stoichiometry derives from two Greek words: stoicheion (meaning "element") and metron (meaning "measure"). Stoichiometry deals with calculations about the masses (sometimes volumes) of reactants and products involved in a chemical reaction. It is a very mathematical part of chemistry, so be prepared for lots of calculator use.

Jeremias Benjaim Richter (1762-1807) was the first to lay down the principles of stoichiometry. In 1792 he wrote: "Die stöchyometrie (Stöchyometria) ist die Wissenschaft die quantitativen oder Massenverhältnisse zu messen, in welchen die chymischen Elemente gegen einander stehen." [Stoichiometry is the science of measuring the quantitative proportions or mass ratios in which chemical elements stand to one another.]

Molar Calculations

Your Tool: Dimensional Analysis

Luckily, almost all of stoichiometry can be solved relatively easily using dimensional analysis. Dimensional analysis is just using units, instead of numbers or variables, to do math, usually to see how they cancel out. For instance, it is easy to see that:

$grams\times {\dfrac {moles}{grams}}\times {\dfrac {atoms}{moles}}=atoms$

It is this principle that will guide you through solving most of the stoichiometry problems (chemical reaction problems) you will see in General Chemistry. Before you attempt to solve a problem, ask yourself: what do I have now? where am I going? As long as you know how many (units) per (other units), this will make stoichiometry significantly easier.

Moles to Mass

Where can you find the molar mass of these elements? The periodic table. You should always have one on hand—don't expect to get very far without one!

How heavy is 1.5 mol of lead? How many moles in 22.34g of water? Calculating the mass of a sample from the number of moles it contains is quite simple. We use the molar mass (mass of one mole) of the substance to convert between mass and moles. When writing calculations, we denote the molar mass of a substance by an upper case "M" (e.g. M(Ne) means "the molar mass of neon"). As always, "n" stands for the number of moles and "m" indicates the mass of a substance. To find the solutions to the two questions we just asked, let's apply some dimensional analysis:

$1.5\;mol\;Pb\;\times {\dfrac {207.2\;g\;Pb}{1\;mol\;Pb}}=310.8\;g\;Pb$

Can you see how the units cancel to give you the answer you want? All you needed to know was that you had 1.5 mol Pb (lead), and that 1 mol Pb weighs 207.2 grams. Thus, multiplying 1.5 mol Pb by 207.2 g Pb and dividing by 1 mol Pb gives you 310.8 g Pb, your answer.

Mass to Moles

But we had one more question: "How many moles in 22.34g of water?" This is just as easy:

$22.34\;g\;H_{2}O\;\times {\dfrac {1\;mol\;H_{2}O}{18\;g\;H_{2}O}}=1.24\;mol\;H_{2}O$

Where did the 18 g H₂O come from? We looked at the periodic table and simply added up the atomic masses of two hydrogens and an oxygen to get the molecular weight of water. This turned out to be 18, and since all the masses on the periodic table are given with respect to 1 mole, we knew that 1 mol of water weighed 18 grams. This gave us the relationship above, which is really just (again) watching units cancel out!

Calculating Molar Masses

Before we can do these types of calculations, we first have to know the molar mass. Fortunately, this is not difficult, as the molar mass is exactly the same as the atomic weight of an element. A table of atomic weights can be used to find the molar mass of elements (this information is often included in the periodic table). For example, the atomic weight of oxygen is 16.00 amu, so its molar mass is 16.00 g/mol.

For species with more than one element, we simply add up the atomic weights of each element to obtain the molar mass of the compound. For example, sulfur trioxide gas is made up of sulfur and oxygen, whose atomic weights are 32.06 and 16.00 respectively.

${\begin{matrix}{\hbox{M(SO}}_{3}{\hbox{)}}&=&32.06+3\times 16.00\\\ &=&80.06{\hbox{g}}/{\hbox{mol}}\\\end{matrix}}$

The procedure for more complex compounds is essentially the same. Aluminium carbonate, for example, contains aluminium, carbon, and oxygen. To find the molar mass, we have to be careful to find the total number of atoms of each element. Three carbonate ions each containing three oxygen atoms gives a total of nine oxygens. The atomic weights of aluminium and carbon are 26.98 and 12.01 respectively.

${\begin{matrix}{\hbox{M}}({\hbox{Al}}_{2}({\hbox{CO}}_{3})_{3})&=&2\times 26.98+3\times 12.01+9\times 16.00\\\ &=&233.99{\hbox{g}}/{\hbox{mol}}\\\end{matrix}}$

Empirical Formulae

The empirical formula of a substance is the simplest ratio of the number of moles of each element in a compound. The empirical formula is ambiguous, e.g. the formula CH could represent CH, C₂H₂, C₃H₃ etc. These latter formulae are called molecular formulae. It follows that the molecular formula is always a whole number multiple of the empirical formula for a compound.

Calculating the empirical formula is easy if the relative amounts of each element in the compound are known. For example, if a sample contains 1.37 mol oxygen and 2.74 mol hydrogen, we can calculate the empirical formula. A good strategy to use is to divide all amounts given by the smallest non-integer amount, then multiply by whole numbers until the simplest ratio is found. We can make a table showing the successive ratios.

Hydrogen	Oxygen
2.74	1.37	divide by 1.37
2	1	ANSWER

The empirical formula of the compound is H₂O.

Here's another example. A sample of piperonal contains 1.384 mol carbon, 1.033 mol hydrogen and 0.519 mol oxygen.

Carbon	Hydrogen	Oxygen
1.384	1.033	0.519	divide by 0.519
2.666	2	1	multiply by 3
8	6	3	ANSWER

The empirical formula of piperonal is C₈H₆O₃.

Converting from Masses

Often, we are given the relative composition by mass of a substance and asked to find the empirical formula. These masses must first be converted to moles using the techniques outlined above. For example, a sample of ethanol contains 52.1% carbon, 13.2% hydrogen, and 34.7% oxygen by mass. Hypothetically, 100g of this substance will contain 52.1 g carbon, 13.2 g hydrogen and 34.7 g oxygen. Dividing these by their respective molar masses gives the amount in moles of each element (as we learned above). These are 4.34 mol, 13.1 mol, and 2.17 mol respectively.

Carbon	Hydrogen	Oxygen
4.34	13.1	2.17	divide by 2.17
2	6	1	ANSWER

The empirical formula of ethanol is C₂H₆O.

Molecular Formula

Beware: In the case of H₂O, the whole number multiple is 1, so its empirical formula is the same as its molecular formula. This is not always the case!

As mentioned above, the molecular formula for a substance equals the count of atoms of each type in a molecule. This is always a whole number multiple of the empirical formula. To calculate the molecular formula from the empirical formula, we need to know the molar mass of the substance. For example, the empirical formula for benzene is CH, and its molar mass is 78.12 g/mol. Divide the actual molar mass by the mass of the empirical formula, 13.02 g/mol, to determine the multiple of the empirical formula, "n". The molecular formula equals the empirical formula multiplied by "n".

${\begin{matrix}{\hbox{M(CH)}}&=&13.02{\hbox{ g/mol}}\\{\hbox{M(benzene)}}&=&78.12{\hbox{ g/mol}}\\{\hbox{M(benzene)}}/{\hbox{M(CH)}}&=(78.12\ g/mol)/(13.02\ g/mol)&=6\\\end{matrix}}$

This shows that the molecular formula for benzene is 6 times the empirical formula of CH. The molecular formula for benzene is C₆H₆.

Solving Mass-Mass Equations

A typical mass-mass equation will give you an amount in grams and ask for another answer in grams.

To solve a mass-mass equation, follow these rules

1. Balance the equation if it is not already.
2. Convert the given quantity to moles.
3. Multiply by the molar ratio of the demanded substance over the given substance.
4. Convert the demanded substance into grams.

For example, given the equation $Cu^{2+}+2AgNO_{3}\to Cu(NO_{3})_{2}+2Ag^{+}$, find out how many grams of silver (Ag) will result from 43.0 grams of copper (Cu) reacting.

• Convert the given quantity to moles.

$43.0g~Cu\times {\frac {1~mol~Cu}{63.55g~Cu}}$

• Multiply by the molar ratio of the demanded substance and the given substance.

$43.0g~Cu\times {\frac {1~mol~Cu}{63.55g~Cu}}\times {\frac {2~mol~Ag}{1~mol~Cu}}$

• Convert the demanded substance to grams.

$43.0g~Cu\times {\frac {1~mol~Cu}{63.55g~Cu}}\times {\frac {2~mol~Ag}{1~mol~Cu}}\times {\frac {107.86g~Ag}{1~mol~Ag}}=1.46\times 10^{2}~g~Ag$

Summary

To solve a stoichiometric problem, you need to know what you already have and what you want to find. Everything in between is basic algebra.

Key Terms

• Molar mass: mass (in grams) of one mole of a substance.
• Empirical formula: the simplest ratio of the number of moles of each element in a compound
• Molecular formula: the actual ratio of the number of moles of each element in a compound

In general, all you have to do is keep track of the units and how they cancel, and you will be on your way!

Chemical Equations

Chemical equations are a convenient, standardized system for describing chemical reactions. They contain the following information.

• The type of reactants consumed and products formed
• The relative amounts of reactants and products
• The electrical charges on ions
• The physical state of each species (e.g. solid, liquid)
• The reaction conditions (e.g. temperature, catalysts)

The final two points are optional and sometimes omitted.

Anatomy of an Equation

${\hbox{H}}_{2(g)}+{\hbox{Cl}}_{2(g)}\to 2{\hbox{HCl}}_{(g)}$

Hydrogen gas and chlorine gas will react vigorously to produce hydrogen chloride gas. The equation above illustrates this reaction. The reactants, hydrogen and chlorine, are written on the left and the products (hydrogen chloride) on the right. The large number 2 in front of HCl indicates that two molecules of HCl are produced for each 1 molecule of hydrogen and chlorine gas consumed. The 2 in subscript below H indicates that there are two hydrogen atoms in each molecule of hydrogen gas. Finally, the (g) symbols subscript to each species indicates that they are gases.

Reacting Species

Species in a chemical reaction is a general term used to mean atoms, molecules or ions. A species can contain more than one chemical element (HCl, for example, contains hydrogen and chlorine). Each species in a chemical equation is written:

${\hbox{E}}_{x(s)}^{y}$

E is the chemical symbol for the element, x is the number of atoms of that element in the species, y is the charge (if it is an ion) and (s) is the physical state.

The symbols in parentheses (in subscript below each species) indicate the physical state of each reactant or product. For ACS Style the state is typeset at the baseline without size change.

• (s) means solid
• (l) means liquid
• (g) means gas
• (aq) means aqueous solution (i.e. dissolved in water)

For example, ethyl alcohol would be written ${\hbox{C}}_{2}{\hbox{H}}_{6}{\hbox{O}}_{(l)}$ because each molecule contains 2 carbon, 6 hydrogen and 1 oxygen atom. A magnesium ion would be written ${\hbox{Mg}}^{2+}$ because it has a double positive ("two plus") charge. Finally, an ammonium ion would be written ${\hbox{NH}}_{4}^{+}$ because each molecule contains 1 nitrogen and 4 hydrogen atoms and has a charge of 1+.

Coefficients

The numbers in front of each species have a very important meaning—they indicate the relative amounts of the atoms that react. The number in front of each species is called a coefficient. In the above equation, for example, one H₂ molecule reacts with one Cl₂ molecule to produce two molecules of HCl. This can also be interpreted as moles (i.e. 1 mol H₂ and 1 mol Cl₂ produces 2 mol HCl).

It is important that the Law of Conservation of Mass is not violated. There must be the same number of each type of atoms on either side of the equation. Coefficients are useful for keeping the same number of atoms on both sides:

$2{\hbox{H}}_{2}+{\hbox{O}}_{2}\to 2{\hbox{H}}_{2}{\hbox{O}}$

If you count the atoms, there are four hydrogens and two oxygens on each side. The coefficients allow us to balance the equation; without them the equation would have the wrong number of atoms. Balancing equations is the topic of the next chapter.

Other Information

Occasionally, other information about a chemical reaction will be supplied in an equation (such as temperature or other reaction conditions). This information is often written to the right of the equation or above the reaction arrow. A simple example would be the melting of ice.

${\hbox{H}}_{2}{\hbox{O}}_{(s)}+heat\to {\hbox{H}}_{2}{\hbox{O}}_{(l)}$, which could be written as ${\hbox{H}}_{2}{\hbox{O}}_{(s)}{\xrightarrow {heat}}{\hbox{H}}_{2}{\hbox{O}}_{(l)}$

Reactions commonly involve catalysts, which are substances that speed up a reaction without being consumed. Catalysts are often written over the arrow. A perfect example of a catalyzed reaction is photosynthesis. Inside plant cells, a substance called chlorophyll converts sunlight into food. The reaction is written:

$6{\hbox{CO}}_{2}+6{\hbox{H}}_{2}{\hbox{O}}+sunlight{\xrightarrow {chlorophyll}}{\hbox{C}}_{6}{\hbox{H}}_{12}{\hbox{O}}_{6}+6{\hbox{O}}_{2}$

Examples

${\hbox{CH}}_{4(g)}+2{\hbox{O}}_{2(g)}\to {\hbox{CO}}_{2(g)}+2{\hbox{H}}_{2}{\hbox{O}}_{(l)}$	This is the equation for burning methane gas (CH4) in the presence of oxygen (O2) to form carbon dioxide and water: CO2 and H2O respectively. Notice the use of coefficients to obey the Law of Conservation of Matter.
${\hbox{Pb}}_{(aq)}^{2+}+2{\hbox{I}}_{(aq)}^{-}\to {\hbox{PbI}}_{2(s)}$	This is a precipitation reaction in which dissolved lead cations and iodide anions combine to form a solid yellow precipitate of lead iodide (an ionic solid).
$2{\hbox{SO}}_{2(g)}+2{\hbox{V}}_{2}{\hbox{O}}_{5(s)}\to 2{\hbox{SO}}_{3(g)}+4{\hbox{V}}_{2}{\hbox{O}}_{(s)}$ $4{\hbox{V}}_{2}{\hbox{O}}_{(s)}+{\hbox{O}}_{2(g)}\to 2{\hbox{V}}_{2}{\hbox{O}}_{5(s)}$	These two equations involve a catalyst. They occur one after another, using divanadium pentoxide to convert sulfur dioxide into sulfur trioxide. If you look closely, you can see that the vanadium catalyst is involved in the reaction, but it does not get consumed. It is both a reactant and a product, but it is necessary for the reaction to occur, making it a catalyst.
$2{\hbox{SO}}_{2(g)}+{\hbox{O}}_{2(g)}{\xrightarrow {V_{2}O_{5}}}2{\hbox{SO}}_{3(g)}$	If we add both equations together, we can cancel out terms that appear on both sides. The resulting equation is much simpler and self-explanatory (although the original pair of equations is more accurate in describing how the reaction proceeds).

Balancing Equations

Balancing Equations

Chemical equations are useful because they give the relative amounts of the substances that react in a chemical equation.

${\hbox{N}}_{2}+3{{\hbox{H}}_{2}}\to 2{\hbox{NH}}_{3}$

In some cases, however, we may not know the relative amounts of each substance that reacts. Fortunately, we can always find the correct coefficients of an equation (the relative amounts of each reactant and product). The process of finding the coefficients is known as balancing the equation.

During a chemical reaction, atoms are neither created or destroyed. The same atoms are present before and after a reaction takes place; they are just rearranged. This is called the Law of Conservation of Matter, and we can use this law to help us find the right coefficients to balance an equation.

For example, assume in the above equation that we do not know how many moles of ammonia gas will be produced:

${\hbox{N}}_{2}+3{{\hbox{H}}_{2}}\to {\color {Red}?}{\hbox{NH}}_{3}$

From the left side of this equation, we see that there are 2 atoms of nitrogen gas in the molecule N₂ (2 atoms per molecule x 1 molecule), and 6 atoms of hydrogen gas in the 3 H₂ molecules (2 atoms per molecule x 3 molecules). Because of the Law of Conservation of Matter, there must also be 2 atoms nitrogen gas and 6 atoms of hydrogen gas on the right side. Since each molecule of the resultant ammonia gas (NH₃) contains 1 atom of nitrogen and 3 atoms of hydrogen, 2 molecules are needed to obtain 2 atoms of nitrogen and 6 atoms of hydrogen.

An Example

${\hbox{O}}_{2}+{{\hbox{H}}_{2}}\to {\hbox{H}}_{2}{\hbox{O}}$	This chemical equation shows the compounds being consumed and produced; however, it does not appropriately deal with the quantities of the compounds. There appear to be two oxygen atoms on the left and only one on the right. But we know that there should be the same number of atoms on both sides. This equation is said to be unbalanced, because the number of atoms are different.
${\hbox{O}}_{\color {Blue}2}+{\color {Blue}2}{{\hbox{H}}_{2}}\to {\color {Blue}2}{\hbox{H}}_{2}{\hbox{O}}$	To make the equation balanced, add coefficients in front of each molecule as needed. The 2 in front of hydrogen on the left indicates that twice as many atoms of hydrogen are needed to react with a certain number of oxygen atoms. The coefficient 1 is not written, since it assumed in the absence of any coefficient.
${\hbox{N}}_{2}+{{\hbox{H}}_{2}}\to {\hbox{NH}}_{3}$	Now, let's consider a similar reaction between hydrogen and nitrogen.
${\hbox{N}}_{\color {Blue}2}+{{\hbox{H}}_{2}}\to {\color {Blue}2}{\hbox{NH}}_{3}$	Typically, it is easiest to balance all pure elements last, especially hydrogen. First, by placing a two in front of ammonia, the nitrogens are balanced.
${\hbox{N}}_{\color {Blue}2}+{\color {Red}3}{{\hbox{H}}_{\color {Blue}2}}\to {\color {Blue}2}{\hbox{NH}}_{\color {Red}3}$	This leaves 6 moles of atomic hydrogen in the products and only two moles in the reactants. A coefficient of 3 is then placed in front of the hydrogen to give a fully balanced reaction.

Tricks in balancing certain reactions

Combustion

A combustion reaction is a reaction between a carbon chain (basically, a molecule consisting of carbons, hydrogen, and perhaps oxygen) with oxygen to form carbon dioxide and water, plus heat. Combustion reactions could get very complex:

$2{\hbox{C}}_{6}{\hbox{H}}_{6}+15{\hbox{O}}_{2}\to 12{\hbox{CO}}_{2}+6{\hbox{H}}_{2}{\hbox{O}}$

Fortunately, there is an easy way to balance these reactions.

First, note that the carbon in C₆H₆ can only appear on the product side in CO₂. Thus, we can write a coefficient of 6 in front of CO₂.

Next, note that the hydrogen in C₆H₆ can only go to H₂O. Thus, we put a 3 in front of H₂O.

We have 15 oxygen atoms on the product side, so there are ${\frac {15}{2}}$ O₂ molecules on the reactant side. To make this an integer, we multiply all coefficients by 2.

Another Example

Helpful Hint!
Note: Fractions are technically allowed as coefficients, but they are generally avoided. Multiply all coefficients by the denominator to remove a fraction.

${\hbox{C}}_{4}{\hbox{H}}_{10}+{{\hbox{O}}_{2}}\to {\hbox{CO}}_{2}+{\hbox{H}}_{2}{\hbox{O}}$	As reactions become more complex, they become more difficult to balance. For example, the combustion of butane (lighter fluid).
${\hbox{C}}_{\color {Red}4}{\hbox{H}}_{\color {Blue}10}+{{\hbox{O}}_{2}}\to {\color {Red}4}{\hbox{CO}}_{2}+{\color {Blue}5}{\hbox{H}}_{\color {Blue}2}{\hbox{O}}$	Once again, it is better to leave pure elements until the end, so first we'll balance carbon and hydrogen. Oxygen can then be balanced after. It is easy to see that one mole of butane will produce four moles of carbon dioxide and five moles of water.
${\color {OliveGreen}2}{\hbox{C}}_{4}{\hbox{H}}_{10}+{{\hbox{O}}_{\color {OliveGreen}2}}\to {\color {OliveGreen}8}{\hbox{CO}}_{2}+{\color {OliveGreen}10}{\hbox{H}}_{2}{\hbox{O}}$	Now there are 13 oxygen atoms on the right and two on the left. The odd number of oxygens prevents balancing with elemental oxygen. Because elemental oxygen is diatomic, this problem comes up in nearly every combustion reaction. Simply double every species except for oxygen to get an even number of oxygen atoms in the product.
$2{\hbox{C}}_{4}{\hbox{H}}_{10}+{\color {Blue}13}{{\hbox{O}}_{\color {Blue}2}}\to {\color {Blue}8}{\hbox{CO}}_{\color {Blue}2}+{\color {Blue}10}{\hbox{H}}_{2}{\hbox{O}}$	The carbon and hydrogens are still balanced, and now there are an even number of oxygens in the product. Finally, the reaction can be balanced.

Limiting Reactants and Percent Yield

Limiting Reactants

When chemical reactions occur, the reactants undergo change to create the products. The coefficients of the chemical equation show the relative amounts of substance needed for the reaction to occur. Consider the combustion of propane:

${\hbox{C}}_{3}{\hbox{H}}_{8}+5{\hbox{O}}_{2}\to 3{\hbox{CO}}_{2}+4{\hbox{H}}_{2}{\hbox{O}}$

For every one mole of propane, there must be five moles of oxygen. For every one mole of propane combusted, there will be three moles of carbon dioxide and four moles of water produced (along with much heat). If a propane grill is burning, there will be a very large amount of oxygen available to react with the propane gas. In this case, oxygen is the excess reactant. There is so much oxygen that the exact amount doesn't matter—it will not run out.

On the other hand, there is not an unlimited amount of propane. It will run out far before the oxygen runs out, making it a limiting reactant. The amount of propane available will decide how far the reaction will go.

Example

2H₂ + O₂ → 2H₂O

If there are three moles of hydrogen, and one mole of oxygen, which is the limiting reactant? How much product is created?

Twice as much hydrogen than oxygen is required. However, there is more than twice as much hydrogen. Thus hydrogen is the excess reactant and oxygen is the limiting reactant. If the reaction proceeds to completion, all of the oxygen will be used up, and one mole of hydrogen will remain. You can imagine this situation like this:

3H₂ + O₂ → 2H₂O + H₂

The reactant that is left over after the reaction is complete is called the "excess reactant". Often, you will want to figure out how much of the excess reactant is left after the reaction is complete. to do this, first use mole ratios to determine how much excess reactant is used up in the reaction.

Here are the ratios that need to be used:

$\left({\frac {\text{moles of limiting reactant}}{1}}\right)*\left({\frac {\text{coefficient of product}}{\text{coefficient of limiting reactant}}}\right)={\text{moles of excess remaining}}$

Percent Yield

Usually, less product is made than theoretically possible. The actual yield is lower than the theoretical yield. To compare the two, one can calculate percent yield, which is ${\frac {actual~yield}{theoretical~yield}}\times 100$.

The percent yield tells us how far the reaction actually went.

Types of Chemical Reactions

Helpful Hint!
Synthesis reactions always yield one product. Reversing a synthesis reaction will give you a decomposition reaction.

The general form of a synthesis reaction is A + B → AB. Synthesis reactions "put things together".

$2{\hbox{H}}_{2(g)}+{\hbox{O}}_{2(g)}\to 2{\hbox{H}}_{2}{\hbox{O}}_{(l)}$	This is the most well-known example of a synthesis reaction—the formation of water via the combustion of hydrogen gas and oxygen gas.
$2{\hbox{Na}}_{(s)}+{\hbox{Cl}}_{2(g)}\to 2{\hbox{NaCl}}_{(s)}$	Another example of a synthesis reaction is the formation of sodium chloride (table salt).

Because of the very high reactivities of sodium metal and chlorine gas, this reaction releases a tremendous amount of heat and light energy. Recall that atoms release energy as they become stable, and consider the octet rule when determining why this reaction has such favorable features.

Decomposition Reactions

These are the same number of synthesis reactions, with the format AB → A + B. Decomposition reactions "take things apart". Just as synthesis reactions can only form one product, decomposition reactions can only start with one reactant. Compounds that are unstable decompose quickly without outside assistance. Also they help other atoms to decompose better.

$2{\hbox{H}}_{2}{\hbox{O}}_{(l)}{\xrightarrow {electricity}}2{\hbox{H}}_{2(g)}+{\hbox{O}}_{2(g)}$	One example is the electrolysis of water (passing water through electrical current) to form hydrogen gas and oxygen gas.
$2{\hbox{H}}_{2}{\hbox{O}}_{2(l)}\to 2{\hbox{H}}_{2}{\hbox{O}}_{(l)}+{\hbox{O}}_{2(g)}$	Hydrogen peroxide slowly decomposes into water and oxygen because it is somewhat unstable. The process is sped up by the energy from light, so hydrogen peroxide is often stored in dark containers to slow down the decomposition.
${\hbox{H}}_{2}{\hbox{CO}}_{3(aq)}\to {\hbox{H}}_{2}{\hbox{O}}_{(l)}+{\hbox{C}}{\hbox{O}}_{2(g)}$	Carbonic acid is the carbonation that is dissolved in soda. It decomposes into carbon dioxide and water, which is why an opened drink loses its fizz.

Single Displacement Reactions

Single displacement reaction, also called single replacement, is a reaction in which 2 elements are substituted for another element in a compound. The starting materials are always pure elements, such as a pure zinc metal or hydrogen gas plus an aqueous compound. When a displacement reaction occurs, a new aqueous compound and a different pure element are generated as products. Its format is AB + C → AC + B. Single Diplacement Adding hydrochloric acid to zinc will cause a gas to bubble out:

${\hbox{Zn}}_{(s)}+2{\hbox{HCl}}_{(aq)}\to {\hbox{ZnCl}}_{2(aq)}+{\hbox{H}}_{2(g)}$

Double Displacement Reactions

In these reactions, two compounds swap components, in the format AB + CD → AD + CB

This is also called an "exchange". Here are the examples below:

1.) HCl + NaOH ----> NaCl + H2O

Precipitation

A precipitation reaction occurs when an ionic substance comes out of solution and forms an insoluble (or slightly soluble) solid. The solid which comes out of solution is called a precipitate. This can occur when two soluble salts (ionic compounds) are mixed and form an insoluble one—the precipitate.

${\hbox{2Pb}}({\hbox{NO}}_{3})_{2(aq)}{\xrightarrow[{}]{\Delta }}{\hbox{2PbO}}_{(s)}+4{\hbox{NO}}_{2(aq)}+{O}_{2(g)}$	An example is lead nitrate mixed with potassium iodide, which forms a bright yellow precipitate of lead iodide.
${\hbox{Pb}}_{(aq)}^{2+}+2{\hbox{NO}}_{3(aq)}^{-}+2{\hbox{K}}_{(aq)}^{+}+2{\hbox{I}}_{(aq)}^{-}\to {\hbox{PbI}}_{2(s)}+2{\hbox{K}}_{(aq)}^{+}+2{\hbox{NO}}_{3(aq)}^{-}$	Note that the lead iodide is formed as a solid. The previous equation is written in molecular form, which is not the best way of describing the reaction. Each of the elements really exist in solution as individual ions, not bonded to each other (as in potassium iodide crystals). If we write the above as an ionic equation, we get a much better idea of what is actually happening.
${\hbox{Pb}}_{(aq)}^{2+}+2{\hbox{I}}_{(aq)}^{-}\to {\hbox{PbI}}_{2(s)}$	Notice the like terms on both sides of the equation. These are called spectator ions because they do not participate in the reaction. They can be ignored, and the net ionic equation is written.

In the solution, there exists both lead and iodide ions. Because lead iodide is insoluble, they spontaneously crystallise and form the precipitate.

Acid-Base Neutralization

In simple terms, an acid is a substance which can lose a H⁺ ion (i.e. a proton) and a base is a substance which can accept a proton. When equal amounts of an acid and base react, they neutralize each other, forming species which aren't as acidic or basic.

${\hbox{HCl}}_{(aq)}+{\hbox{NaOH}}_{(aq)}\to {\hbox{H}}_{2}{\hbox{O}}_{(l)}+{\hbox{NaCl}}_{(aq)}$	For example, when hydrochloric acid and sodium hydroxide react, they form water and sodium chloride (table salt).
${\hbox{H}}_{(aq)}^{+}+{\hbox{OH}}_{(aq)}^{-}\to {\hbox{H}}_{2}{\hbox{O}}_{(l)}$	Again, we get a clearer picture of what's happening if we write a net ionic equation.

Acid base reactions often happen in aqueous solution, but they can also occur in the gaseous state. Acids and bases will be discussed in much greater detail in the acids and bases section.

Combustion

The combustion of methane (releasing heat and light)

Combustion, better known as burning, is the combination of a substance with oxygen. The products are carbon dioxide, water, and possible other waste products. Combustion reactions release large amounts of heat. C₃H₈, better known as propane, undergoes combustion. The balanced equation is:

${\hbox{C}}_{3}{\hbox{H}}_{8}+5{\hbox{O}}_{2}\to 3{\hbox{CO}}_{2}+4{\hbox{H}}_{2}{\hbox{O}}$

Combustion is similar to a decomposition reaction, except that oxygen and heat are required for it to occur. If there is not enough oxygen, the reaction may not occur. Sometimes, with limited oxygen, the reaction will occur, but it produces carbon monoxide (CO) or even soot. In that case, it is called incomplete combustion. If the substances being burned contain atoms other than hydrogen and oxygen, then waste products will also form. Coal is burned for heating and energy purposes, and it contains sulfur. As a result, sulfur dioxide is released, which is a pollutant. Coal with lower sulfur content is more desirable, but more expensive, because it will release less of the sulfur-based pollutants.

Organic Reactions

This is carboxylic acid. All functional groups end with an "R"—a placeholder for the rest of the molecule.

Organic reactions occur between organic molecules (molecules containing carbon and hydrogen). Since there is a virtually unlimited number of organic molecules, the scope of organic reactions is very large. However, many of the characteristics of organic molecules are determined by functional groups—small groups of atoms that react in predictable ways.

Another key concept in organic reactions is Lewis basicity. Parts of organic molecules can be electrophillic (electron-loving) or nucleophillic (nucleus, or positive loving). Nucleophillic regions have an excess of electrons—they act as Lewis bases—whereas electrophillic areas are electron deficient and act as Lewis acids. The nucleophillic and electrophillic regions attract and react with each other. Organic reactions are beyond the scope of this book, and are covered in more detail in Organic Chemistry. However, most organic substances can undergo replacement reactions and combustion reactions, as you have already learned.

Redox

The formation of hydrogen fluoride from the elements requires reduction of fluorine and oxidation of hydrogen.

Redox is an abbreviation of reduction/oxidation reactions. This is exactly what happens in a redox reaction, one species is reduced and another is oxidized. Reduction involves a gain of electrons and oxidation involves a loss, so a redox reaction is one in which electrons are transferred between species. Reactions where something is "burnt" (burning means being oxidised) are examples of redox reactions, however, oxidation reactions also occur in solution, which is very useful and forms the basis of electrochemistry.

Redox reactions are often written as two half-reactions showing the reduction and oxidation processes separately. These half-reactions are balanced (by multiplying each by a coefficient) and added together to form the full equation. When magnesium is burnt in oxygen, it loses electrons (it is oxidised). Conversely, the oxygen gains electrons from the magnesium (it is reduced).

${\begin{matrix}{\hbox{Mg}}&\to &{\hbox{Mg}}^{2+}+2e^{-}&\times 2\\{\hbox{O}}_{2}+4e^{-}&\to &2{\hbox{O}}^{2-}&\times 1\\2{\hbox{Mg}}+{\hbox{O}}_{2}+4e^{-}&\to &2{\hbox{MgO}}+4e^{-}&\ \\\end{matrix}}$

Redox reactions will be discussed in greater detail in the redox section.

Energy Changes in Chemical Reactions

Exothermic and Endothermic Reactions

The release of energy in chemical reactions occurs when the reactants have higher chemical energy than the products. The chemical energy in a substance is a type of potential energy stored within the substance. This stored chemical potential energy is the heat content or enthalpy of the substance.

If the enthalpy decreases during a chemical reaction, a corresponding amount of energy must be released to the surroundings. Conversely, if the enthalpy increases during a reaction, a corresponding amount of energy must be absorbed from the surroundings. This is simply the Law of Conservation of Energy.

absorbtion reactions is when a chemical reaction releases more energy than it absorbs and you can also see them die on the inside.

reletion reactions is when a chemical reaction absorbs more energy than it releases.

You are already familiar with enthalpy: melting ice is endothermic and freezing water is exothermic.

Examples

When methane burns in air the heat given off equals the decrease in enthalpy that occurs as the reactants are converted to products. ${\hbox{CH}}_{4(g)}+2{\hbox{O}}_{2(g)}\to {\hbox{CO}}_{2(g)}+2{\hbox{H}}_{2}{\hbox{O}}_{(g)}+energy$ The enthal difference between the reactants and the products is equal to the amount of energy released to the surroundings. A reaction in which energy is released to the surroundings is called an exothermic reaction. In this type of reaction the enthalpy, or stored chemical energy, is lower for the products than the reactants.

When ammonium nitrate is dissolved in water, energy is absorbed and the water cools. This concept is used in "cold packs". ${\hbox{NH}}_{4}{\hbox{NO}}_{3(s)}+water+energy\to {\hbox{NH}}_{4(aq)}^{+}+{\hbox{NO}}_{3(aq)}^{-}$ The enthalpy difference between the reactants and the products is equal to the amount of energy absorbed from the surroundings. A reaction in which energy is absorbed from the surroundings is called an endothermic reaction. In endothermic reactions the enthalpy of the products is greater than the enthalpy of the reactants.

Because reactions release or absorb energy, they affect the temperature of their surroundings. Exothermic reactions heat up their surroundings while endothermic reactions cool them down. The study of enthalpy, along with many other energy-related topics, is covered in the Thermodynamics Unit.

Activation Energy

Think about the combustion of methane. It releases enough heat energy to cause a fire. However, the reaction does not occur automatically. When methane and oxygen are mixed, an explosion does not instantly occur. First, the methane must be ignited, usually with a lighter or matchstick. This reveals something about reactions: they will not occur unless a certain amount of activation energy is added first. In this sense, all reactions absorb energy before they begin, but the exothermic reactions release even more energy. This can be explained with a graph of potential energy:

This graph shows an exothermic reaction because the products are at a lower energy than the reactants (so heat has been released). Before that can happen, the energy must actually increase. The amount of energy added before the reaction can complete is the activation energy, symbolized E_a.

Predicting Chemical Reactions

Types of Reactions

There are several guidelines that can help you predict what kind of chemical reaction will occur between a mixture of chemicals.

• Several pure elements mixed together may undergo a synthesis reaction.
• A single compound may undergo a decomposition reaction. It often forms water or hydrogen gas.
• A pure element mixed with an ionic compound may undergo a single replacement reaction.
• Two different ionic compounds are very likely to undergo a double replacement reaction.
• An organic compound (containing carbon and hydrogen) can usually react with oxygen in a combustion reaction.

However, not all elements will react with each other. To better predict a chemical reaction, knowledge of the reactivity series is needed.

Reactivity

When combining two chemicals, a single- or double-replacement reaction doesn't always happen. This can be explained by a list known as the reactivity series, which lists elements in order of reactivity. The higher on the list an element is, the more elements it can replace in a single- or double-replacement reaction. When deciding if a replacement reaction will occur, look up the two elements in question. The higher one will replace the lower one.

Elements at the very top of the series are so reactive that they can replace hydrogen from water. This explains the explosive reaction between sodium and water:

$2{\hbox{Na}}_{(s)}+2{\hbox{H}}_{2}{\hbox{O}}_{(l)}\to 2{\hbox{NaOH}}_{(aq)}+{\hbox{H}}_{2(g)}$

Elements in the middle of the list will react with acids (but not water) to produce a salt and hydrogen gas. Elements at the bottom of the list are mostly nonreactive.

Elements near the top of the list will corrode (rust, tarnish, etc.) in oxygen much faster than those at the bottom of the list.

The Reactivity Series

• Red: elements that react with water and acids to form hydrogen gas, and with oxygen.
• Orange: elements that react very slowly with water but strongly with acids.
• Yellow: elements that react with acid to form hydrogen gas, and with oxygen.
• Grey: elements that react with oxygen (tarnish).
• White: elements that are often found pure; relatively nonreactive.

Most Reactive

Cs

K

Na

Li

Sr

Ca

Rb

Ba

Mg

Al

(C)

Mn

Zn

Cr

Fe

Cd

Co

Ni

Sn

Pb

(H2)

Sb

Bi

Cu

Hg

Ag

Pt

Au

Least Reactive