1. In chemistry, activation energy is a term that means the minimum energy that must be input to a chemical system with
potential reactants to cause a chemical reaction.
2. Chemical changes happen between reactants (the initial substances that enter into the reaction) and products (the final
substances that are present at the end of the reaction). Chemical reactions involve a rearrangement of the atoms in
reactants to form products with new structures where atoms are conserved.
3.
Bohr’s model:
4. There are many types of chemical bonds and forces that bind molecules together. The two most basic types of bonds are
characterized as either ionic or covalent. In ionic bonding, atoms transfer electrons to each other. Ionic bonds require at
least one electron donor and one electron acceptor. In contrast, atoms with the same electronegativity share electrons in
covalent bonds, because neither atom preferentially attracts or repels the shared electrons.
5. Ionic bonding is the complete transfer of valence electron(s) between atoms. It is a type of chemical bond that generates
two oppositely charged ions. In ionic bonds, the metal loses electrons to become a positively charged cation, whereas the
nonmetal accepts those electrons to become a negatively charged anion. Ionic bonds require an electron donor, often a
metal, and an electron acceptor, a nonmetal.
6. Ionic bonding is observed because metals have few electrons in their outer-most orbitals. By losing those electrons, these
metals can achieve noble gas configuration and satisfy the octet rule. Similarly, nonmetals that have close to 8 electrons
in their valence shells tend to readily accept electrons to achieve noble gas configuration. In ionic bonding, more than 1
electron can be donated or received to satisfy the octet rule. The charges on the anion and cation correspond to the
number of electrons donated or received. In ionic bonds, the net charge of the compound must be zero.
7. A covalent bond is a chemical bond that involves the sharing of electron pairs between atoms. These electron pairs are
known as shared pairs or bonding pairs and the stable balance of attractive and repulsive forces between atoms when
they share electrons is known as covalent bonding.
8. Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. The Pauling scale is the
most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to
cesium and francium which are the least electronegative at 0.7.
If the atoms are equally electronegative, both have the same tendency to attract the bonding pair of electrons, and so
it will be found on average half way between the two atoms. To get a bond like this, A and B would usually have to be
the same atom. You will find this sort of bond in, for example, H2 or Cl2 molecules.
Note: It's important to realize that this is an average picture. The electrons are actually
in a molecular orbital, and are moving around all the time within that orbital.
This sort of bond could be thought of as being a "pure" covalent bond - where the electrons are shared evenly
between the two atoms.
What happens if B is slightly more electronegative than A?
B will attract the electron pair rather more than A does.
That means that the B end of the bond has more than its fair share of electron density and so becomes slightly
negative. At the same time, the A end (rather short of electrons) becomes slightly positive. In the diagram, " " (read
as "delta") means "slightly" - so + means "slightly positive".
Defining polar bonds
This is described as a polar bond. A polar bond is a covalent bond in which there is a separation of charge between
one end and the other - in other words in which one end is slightly positive and the other slightly negative. Examples
include most covalent bonds. The hydrogen-chlorine bond in HCl or the hydrogen-oxygen bonds in water are typical.
What happens if B is a lot more electronegative than A?
In this case, the electron pair is dragged right over to B's end of the bond. To all intents and purposes, A has lost
control of its electron, and B has complete control over both electrons. Ions have been formed.
9. Naming Chemical Compounds
Know what makes a compound ionic. Ionic compounds contain a metal and a nonmetal. Refer to the periodic table of
elements to see what categories the elements in the compound belong to.
Build the name. For a two element ionic compound, the naming is simple. The first part of the name is the name of the
metal element. The second part is the name of the nonmetal element, with the suffix “-ide.” Example: Al2O3. Al2 =
Aluminium; O3 = Oxygen. So the name would be “aluminium oxide.”
Recognize transition metals. Metals in the D and F blocks of the periodic table are known as transition metals. Their
charge is written with a Roman numeral when writing out the compound name. This is because they can have more than
one charge and make more than one compound.
Example: FeCl2 and FeCl3. Fe = Iron; Cl2 = Chloride -2; Cl3 = Chloride -3. The names would be iron(II) chloride and
iron(III) chloride.
Understand what a polyatomic compound is. These compounds are built off of groups of atoms that are bonded
together, and the entire group has a positive or negative charge. You can do three basic things to polyatomic compounds:
Add a hydrogen to the beginning of the compound. The word “hydrogen” is added to the beginning of the compound
name. This reduces the negative charge by one. For example, “carbonate” CO 32- becomes “hydrogen carbonate” HCO3-.
Remove an oxygen from the compound. The charge stays the same and the ending of the compound changes from “-ate”
to “–ite”. For example: NO3 to NO2goes from “nitrate” to “nitrite.”
Memorize the most common ion groups. These are the basis for forming most polyatomic compounds. Listed in order
of increasing negative charge, they are:

Hydroxide ion: OH-

Nitrate ion: NO3-

Hydrogen carbonate ion: HCO3-

Permanganate ion: MnO4-

Carbonate ion: CO32-

Chromate ion: CrO42-

Dichromate ion: Cr2O72-

Sulfate ion: SO42-

Sulphite ion: SO32-

Thiosulfate ion: S2O3<2-

Phosphate ion: PO43-

Ammonium ion: NH4+
10. Acid-Base Reactions. When an acid and abase are placed together, they react to neutralize
the acid and base properties, producing a salt. The H(+) cation of the acid combines with the OH(-) anion of the base to
form water. The compound formed by the cation of the base and the anion of the acid is called a salt.
11. Combustion reactions almost always involve oxygen in the form of O2, and are almost always exothermic, meaning
they produce heat. Chemical reactions that give off light and heat and light are colloquially referred to as "burning."
12. Redox reactions are comprised of two parts, a reduced half and an oxidized half, that always occur together. The
reduced half gains electrons and the oxidation number decreases, while the oxidized half loses electrons and the
oxidation number increases. There is no net change in the number of electrons in a redox reaction. Those given off in the
oxidation half reaction are taken up by another species in the reduction half reaction. The two species that exchange
electrons in a redox reaction are given special names. The ion or molecule that accepts electrons is called the oxidizing
agent; by accepting electrons it causes the oxidation of another species. Conversely, the species that donates electrons is
called the reducing agent; when the reaction occurs, it reduces the other species. In other words, what is oxidized is the
reducing agent and what is reduced is the oxidizing agent.
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A molecule is a group of atoms, frequently nonmetals, strongly linked by chemical bonds. Covalent
bonds generally can't be understood on the same basis as ionic bonds, which are bonded by the
attraction of oppositely charged ions. Consider the H2 molecule; the two atoms are held together
tightly although no ions are present. In 1916 G.N. Lewis proposed that the strong attractive force
between two atoms in a molecule results from a covalent bond, a chemical bond formed by the
sharing of a pair of electrons between atoms. A few years later the covalent bond in hydrogen was
quantitatively explained using the newly discovered theory of quantum mechanics. This provided
evidence in support Lewis' hypothesis. Lets look at the findings.
Describing Covalent Bonding
Consider the formation of a covalent bond between to H atoms. As the atoms approach one another,
their 1s orbitals begin to overlap. Each electron can then occupy the space around both atoms such
that the two electrons can be shared by the atoms. The electrons are attracted simultaneously by the
positive charges of the two hydrogen nuclei, which causes a covalent bond to form. Although ions do
not exist, the force that hold the atoms together can still be regarded as arising from the attraction of
oppositely charged particles, nuclei and electrons.
The potential-energy curve for H2 shows that the formation of the bond is energetically favorable.
The graph plots the potential energy against the distance between nuclei. If the nuclei are too close
the atoms repel (top left). Too far, and they do not attract one another (middle right). The bond
occurs when the H atoms achieve a lowest energy state observed on the potential energy diagram
(bottom middle). The distance between nuclei at this minimum energy is called the bond length of
H2. It is the normal distance between nuclei in the molecule. The bond length of H2 is 74 pm. The
bond energy is the average enthalpy change for the breaking of the bond in H2. Bond length is
equal to the difference between the minimum potential energy and PE = 0.
Recall that Lewis dot structures can be used to illustrate covalent bonds as follows.
Coordinate Covalent Bonds
Often covalent bonds are formed between atoms that both donate electrons as is the case of
hydrogen shown above. However, it is possible for both electron to come from the same atom. A
coordinate covalent bond (dative bond) is formed when both electrons of the bond are donated by
one atom.
This type of covalent bond is not essentially different from other covalent bonds. However, being
aware of and having an understanding of this type of bond can enhance your understanding of how
molecular compounds form.
Multiple Bonds
In the molecules we have described so far, each of the bonds has been a single bond; a covalent
bond in which a single pair of electrons is shared by two atoms. But it is possible for atoms to share
two or more electron pairs. A double bond is a covalent bond in which atoms share two pairs of
electrons. A triple bond is a covalent bond in which atoms share three pairs of electrons. A
comparison of the Lewis structures for methane, CH4, ethylene, C2H4, and acetylene, C2H2 shows
these bonds. Note the octet of electrons on each carbon atom. Double bonds are most often formed
by C, N, O, and S atoms. Triple bonds from mostly to C and N atoms.
Polar Covalent Bonds; Electronegativity
When the electronegativities of the atoms in a covalent bond are not equal, electrons are not shared
equally. In this case there may be partial transfer of electron density from one atom to the other
resulting in a bond or molecule with oppositely charged ends. The water molecule provides a familiar
example of this. The greater the electronegativity difference, the more ionic character the bond has.
Bonds that are partly ionic are called polar covalent bonds.
Nonpolar covalent bonds occur when electrons are shared equally between atoms. This type of
bond arises when the electronegativities of the two atoms are equal such as in the case of the H 2
molecule where the difference in electronegativity is 0.
Recall the electronegativity is a measure of the ability of an atom in a molecule to draw bonding
electrons to itself. Electronegativity is an experimentally derived quantity and several scales have
been proposed. The most widely used scale was developed by Linus Pauling. These values are
sometimes called "Pauling's electronegativity values" or more simply the "Pauling scale". Pauling's
values are often represented on the periodic table as show below.
The difference in electronegativity between atoms allows us to predict what type of bond the atoms
will form as shown below.
In general, if the electronegativity difference is between 0.0 and 0.4 the bond is nonpolar covalent. If
the difference is between 0.4 and 2.0 it is polar covalent while differences greater than two generally
result in an ionic bonds.
Example: Using Electronegativities to Obtain Relative Bond Polarities
Use electronegativity values to arrange the following bonds in order of increasing polarity: P-H, H-O,
and C-Cl.
Solution:
To solve this problem you would using the Pauling scale to determine the absolute value of the
electronegativity difference between the atoms in each bond. The bonds are then placed in order of
increasing polarity based on these differences. The greater the difference the more polar the bond.
The differences are: P-H, 0.0; H-O, 1.4; C-Cl, 0.5. The order is therefore P-H, C-Cl, H-O.