A Chemist’s View of Explosives:

Chemical Bonding Notes

I. Chemical bond: a mutual electrical attraction between the nuclei and valence electrons of

different atoms that binds the atoms together. Another way to describe a chemical bond is to say the attractive forces between atoms or ions

in compounds. In ionic compounds it is an attractive force between positive and negative

ions.

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In ionic bonding valence electrons are actually TRANSFERRED between a nonmetal and metal. This happens because a non-metallic atom is much more electronegative and it can pull electrons away from the less electronegative metallic atom. In an ionic compound the positive and negative ions combine so that the overall charge is zero.

Sometimes the more electronegative atom is not “powerful” enough to completely take away the electrons from another atom so the atoms SHARE electrons. This sharing of electrons is called a covalent bond.

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Ionic Bonding occurs between metals and nonmetals.

Covalent Bonding occurs between nonmetals.

Bonds (and compounds) form in order to obtain

an electron configuration like

that of noble gases!

II. Formation of Ionic Bonds and Ionic Compounds

A. Electron Dot Structures: show the placement and transfer of valence electrons. Rules to remember when drawing electron dot structures:

1. Only valence electrons are shown. Valence electrons are the electrons in the outermost s and p sublevels. Transition metals could also have d sublevel valence electrons.

2. Valence electrons are shown as dots and are not drawn randomly! They are arranged around the element's symbol to correspond to the elements electron configuration. (Only 2 dots or electrons per side.)

3. Follow the Octet Rule which sates that atoms form bonds in order to obtain 0 or 8 valence electrons, because of this electron dot structures will show no more than 8 electrons for each atom or ion. Another way to think of the Octet rule: Atoms react by changing the number of their electrons so as to acquire the stable electron configuration of a noble gas.

B. Electron Dot Structures for Atoms:

Write the element's symbol and place the appropriate number of dots to represent the valence electrons around the symbol. (The electron configuration is given to help you understand the idea of valence electrons.)

a.) Ca [Ar]4s2 b.) Li [He]2s1 c.) Be [He]2s2 d.) O [He]2s22p4

e.) Br [Ar]4s23d104p5

C. Electron Dot Structures for Ions: Ions form when atoms lose or gain

valence electrons.

(1.) Cations - these form when atoms have LOST valence electrons. To draw the dot structures write the symbol, put [ ] around the symbol, and the charge of the ion outside the [ ]. There are NO dots because there are NO valence electrons! (You may want to write the electron configuration for the atom to help you see what happens when it ionizes.)

a.) Mg ion

b.) Li ion

c.) Al ion

d.) Ba ion

(2.) Anions-these form when atoms have GAINED valence electrons. To draw the dot structures write the symbol, draw 8 dots around the symbol, put [ ] around the symbol, and the charge of the ion outside the [ ]. (You may want to write the electron configuration for the atom to help you see what happens when it ionizes.)

a.) S ion

b.) Br ion

c.) N ion

d.) P ion

(3.) Transition and Inner Transition Elements-the number of valence electrons for these are harder to predict based on their position on the periodic table because some of these elements have valence electrons in the d sublevel. Example:

a.) How many valence electrons does an atom of

iron have? To answer this question write the

electron configuration for iron:

Are there any unstable electrons in the d level?

When iron ionizes what are the possible ions?

b.) How many valence electrons does an atom of titanium have?

Electron configuration for titanium:

Are there any unstable electrons in the d level?

When iron ionizes what are the possible ions?

D. Pseudo-noble gas electron configuration-elements that cannot acquire a noble gas electron configuration, but can become somewhat stable with 18 electrons in their outer shell. Examples are: Hg+2, Cd+2, Au+1, Cu+1

E. Electron Dot Structures for Ionic Compounds:

1. Write the electron dot structure for each of the elements involved.

2. Draw arrows from the electrons of the metallic atom to the non-metallic atom. This shows the transfer of electrons.

3. After that right the dot diagram for the new ionic compound, including charges.

A. Lewis Dot Structures for Ionic Compounds (compounds held together by ionic bonds – usually a M

Example: a.) Sodium and Chlorine

Na + Cl [ Na ]+ + [ • Cl• ]-

• •• ••••

•••

••••

http://www.beyondbooks.com/psc92/3b.asp

Examples:

b.) Magnesium and Oxygen

Examples:

c.) Aluminum and Oxygen

Examples:

d.) Calcium and Fluorine

Examples:

e.) Sodium and Nitrogen

F. Characteristics of ionic compounds (compared to molecular compounds)

-higher melting points

-higher boiling points

-generally hard, brittle solids

-when melted or dissolved in water they can conduct electricity

-shapes are crystalline in nature (page 177)

– square/cube

Copper sulfate has a triclinic crystal structure.

Lattice Energy: the energy released when one mole of an ionic crystalline compound is formed from gaseous ions. Negative values for lattice energy mean that energy was released when the ionic crystal is formed.

Formula unit: the simplest collection of atoms from which an ionic compound’s formula can be established. A formula unit is to an ionic compound as a molecule is to a covalent compound.

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Remember what happens when an ionic bond forms?

• One or more electrons from 1 atom are removed and attached to another atom, resulting in a cation and an ion which attract each other

Write this down!!Ionic Compounds - never exist as individual formula units, are solids

Molecular Compounds-can exist as an individual molecule, are usually liquids or gaseshttp://web.jjay.cuny.edu/~acarpi/NSC/6-react.htm

III. Formation of Covalent Bonds and Molecular Compounds

A. Covalent Bonds – a bond in which electrons are shared. Which compounds have covalent bonding?

1. Molecular (or covalent) compounds - these are two NON-METALS. These compounds always have covalent bonding

2. Polyatomic ions (PO-3, NO-1, CN-1). These ions are held together with covalent bonds.

B. Type of Covalent Bonds

l. Nonpolar covalent bond-a covalent bond in which the bonding electrons are shared equally by the bonded atoms, resulting in an evenly balanced charge. If the difference in electronegativity between two bonded atoms is less than 0.3 a nonpolar bond will exist.

Nonpolar Covalent Bond (equal sharing)

Oxygen AtomOxygen Atom Oxygen AtomOxygen Atom

Oxygen Molecule (O2)

2. Polar covalent bond-a bond is which the bonded atoms do NOT share the bonding electrons equally. A polar covalent bond is a bond in which the atoms have an unevenly balanced charge. If the difference in electronegativity between two bonded atoms is from 0.3 to 1.7, a polar bond will exist If the difference in EN is less than 0.3 then the bond is nonpolar covalent. The atom with the greater electronegativity will pull the electrons toward it, giving that atom a slightly negative charge. A partial negative charge is shown by - and the less electronegative atom will have a partial positive charge, designated +.

Bond PolarityBond Polarity

• ElectronegativityElectronegativity

–Attraction an atom has for a shared Attraction an atom has for a shared pair of electrons.pair of electrons.

–higher ehigher e--neg atom neg atom --

–lower elower e--neg atomneg atom ++

Polar Covalent Bonds: Unevenly matched, but willing to share.

The H-O bonds in water are polar covalent because oxygen is more electronegative than hydrogen, and therefore electrons are pulled closer to oxygen.

Practice: Find the differences in electronegativity (EN page 151 or on calculator) in the following pairs of atoms. Designate which, if any, atom is partially negative and partially positive.

a. H and Cl

b. F and Br

c. S and I

d. O and H

Another Example: cesium-fluorine bonding

Cs EN = 0.7F EN = 4

4 - 0.7 = 3.3

Ionic Bond

D. The Octet Rule and Dot Structures -chemical compounds tend to form so that each atom, by gaining, losing, or sharing electrons,

has an octet of valence electrons. Electron dot structures (also known as Lewis dot diagrams)

show valence electrons as dots around the element’s symbol. Dot structures for molecules

show atoms sharing dots (covalent bonds).

Ar

Covalent bonds are single, double, or triple

single bond-two atoms share one pair of electrons (1 sigma bond)

double bond-two atoms share two pair of electrons (1 sigma and 1 pi bond)

triple bond-two atoms share three pair of electrons (1 sigma and 2 pi bonds)

Rules for correctly illustrating the dot structure of a molecule:

1. Add up the TOTAL number of valence electrons in the substancea) be sure to subtract 1 electron if it is a positively-charged ion

(NH4+1)

b) be sure to add electrons for each negative charge on an ion (SO4

-2)2. Decide what is the central atom. The central atom is the one

that is least represented. (or the least electronegative)3. Hook the particles together using a short straight line (or 2 dots)

to indicate a covalent bond between atoms. Each of these "bonds" represents 2 shared electrons.

4. Subtract the number of electrons used in "hooking" the atoms together from the total valence electrons.

5. Use the "leftover" electrons, if any, to fill the octets of the peripheral atoms.

6. Place anymore "leftover" electrons on the central atom (in pairs).

http://chemsite.lsrhs.net/d_bonding/flashLewis.html

Practice: Draw the electron dot structures for the following:

Practice: Draw the electron dot structures for the following:

1. carbon tetrachloride (CCl4)

2. OF2

3. NF4+1

4. PCl35. CO2

6. N2

E. Exceptions to the Octet Rule

a. Hydrogen atoms have less then an octet.

Hydrogen only needs 2 valence electrons.

H2

HN3

F. Resonance – a concept in which two or more Lewis structures for the same arrangement of atoms (resonance structures) are used to describe the bonding in a molecule or ion. To show resonance, a double-headed arrow is placed between a molecule’s resonance structures.

Example: Ozone

G. Coordinate Covalent Bond is formed when one atom contributes BOTH bonding electrons in a covalent bond.

Examples:

carbon monoxide

SO42-

Chemical Bonding Notes Part 2

I. Metallic Bonds-a third type of bond. This is what holds pure metal atoms together.

Metallic bonding accounts for many physical properties of metals, such as strength, malleability, ductility, thermal and electrical conductivity, opacity, and luster.

What happens to form a metallic bond? 1. each metal donates its valence electron(s) to form

an electron cloud 2. this leaves positive particles which are "cemented"

together with the negative electron cloud, often called a “sea of electrons.”

Metallic Bonds: Mellow dogs with plenty of bones to go around.

A Sea of Electrons

Metals Form Alloys

Metals do not combine with metals. They form Alloys which is a solution of a metal in a metal.Examples are steel, brass, bronze and pewter.

II. Polarity - Polar and nonpolar molecules - if a

molecule contains a polar bond, is the molecule itself polar also?

It depends!!

1. Polar molecules - a polar molecule is positive at one point and negative at another point. For example, HBr contains a polar bond. As a result the hydrogen side of the molecule is partially positive and the bromine side of the molecule is partially negative. It "acts like a magnet".

Water is a molecule with two polar bonds. A molecule of water is also polar because there is an area of positive charge on the hydrogen atoms and an area of partially negative charge on the oxygen. Its bent shape allows it to act somewhat like a magnet. The presence of 2 unshared pair of electrons is a main factor for it being a polar molecule.

2. Nonpolar molecules - Carbon tetrachloride has four C-Cl bonds. Each bond is a polar covalent bond.

The molecule itself is nonpolar because of its

1.) shape. It is perfectly symmetrical, and

2.) the partially-positive carbon in the center which is covered by the 4 partially negative chlorine atoms. It cannot “act like a magnet”.

Cl

Cl

Cl

3. Helpful hints and practice: A. Hints to help you decide if a molecule is POLAR: 1. Does it have at least one polar bond? If so, it's probably polar. 2. Does it have any unshared pairs of electrons around the central

atom? If so, it is probably polar. 3. Can the molecule act like a magnet? If so, it is probably polar. B. Practice: Which of the following molecules are polar and which ones

are nonpolar molecules? If the molecule is polar, tell why it is polar.

1.) SO2

2.) H2S 3.) CO2

4.) BF3

5.) CH4

6.) ClO2-1

7.) CH3Cl8.) PO4

-

9.) MgCl2

III. Hydrates: Some compounds trap water inside their crystal structure and are known as hydrates. You will not be able to predict which compounds will form hydrates. CuSO4 5H20 is an example of a hydrate. This says that one formula unit of cupric sulfate will trap 5 molecules of water inside its crystal.

Hydrates are named by naming the ionic compound by the regular rules and then adding (as a second word) a prefix indicating the number of water molecules. You will use the word “hydrate” to indicate water. The above compound would be called cupric sulfate pentahydrate.

To find the formula mass of a hydrate, simply find the mass of the ionic compound by itself and then ADD the mass of water molecule(s) to that mass.

Practice: What is the formula mass of barium chloride dihydrate?

1. What is the formula mass of aluminum sulfate octahydrate?