Bonding Review

-Define Ionic, and Covalent Bonding-Discuss ionic and covalent properties-Learn to draw Lewis Structures for

Ionic and Covalent structures-Distinguish between Polar and Non-

polar Covalent compounds

• Ionic compounds are composed of both metals and non-metals. The bond that is formed is based on electrostatic forces between negatively(anion) and positively(cation) charged ions.

Ionic Compounds

• Ionic bonding occurs by the transfer of one or more electrons from one atom to another

Properties of Ionic Compounds• Usually form crystalline solids at

room temperature (made of ions)• Have high melting points (strong

bond formed by electrostatic force)• Usually do not conduct electricity as

a solid• Usually dissolve in water (water is

polar so attracts ions)• Usually conduct electricity when in

solution or molten state (ions)

• An ion is an atom or group of atoms that have a charge. Atoms normally have a neutral charge because most often they have the same number of electrons and protons. They become ions by the loss or addition of one or more electrons. This process is called ionization.

• An ion that has more electrons than protons is called an anion, and an ion that has fewer electrons than protons is called a cation.

Ion Review

THE OCTET RULE

• The interaction of ionic bonds is when atoms gain or lose electrons until the outer shell of electrons is full and stable with 8 electrons. This is part of the octet rule.

• Octet rule:

When atoms combine to form molecules they generally each lose, gain, or share valence electrons until they attain or share eight (full shells) and reach a noble gas electron configuration which makes the atom stable.

Ionic Bonds• Nonmetals usually have four or more electrons in

their outer shell. To make their outer shell full, it’s easier(it takes less ionization energy) for them to gain three or four electrons than to lose four or five electrons.

• When you look at the metals, they usually have three or less electrons in their outer shell. Opposite from nonmetals, it is easier (less energy) for metals to lose three or less electrons than to gain four or more.

• Therefore it makes sense that metals and nonmetals bond together easily.

Lewis Dot Structure

• In 1902 Gilbert Newton Lewis invented the valance bond theory.

• Lewis came up with an easy way to represent electrons in the outer shells of ions. His invention is called “Lewis Dot Symbols”.

• Lewis structures are used to visualize the valence electrons of elements. In the Lewis model, an element symbol is inside the valence electrons of the s and p subshells of the outer ring.

Lewis Dot Structures

See Video Below

Drawing Lewis Structures of Ionic Compounds

https://www.youtube.com/watch?v=Vk3uF7_Zldo

Covalent Compounds• Covalent compounds are made up of

two nonmetals. A single covalent bond is formed when a pair of electrons is shared between two atoms

• There are two types of covalent bonds: non-polar covalent and polar covalent.

Properties of Covalent Compounds

• Forms molecules• Low melting and boiling points

(weaker bond)• Larger more complex compounds

will have higher melting and boiling points

• Usually do not conduct electricity as a solid or when molten or in solution (no ions)

• Usually do not dissolve in water (depends on polarity)

Non-Polar Covalent Bonding

• Accounts for the bond that keeps two atoms of the same element together (diatomic molecules). (Cl2, H2)

• Atoms share electrons equally to achieve noble gas configurations

• Shared electrons are attracted to both nuclei, which keeps atoms together

• Electrons involved in bonding are called shared electron pairs, ones that are not are called lone electron pairs

Watch Video Lewis Structures and Covalent

Bonding https://www.youtube.com/watch?v=

NYFE5uslaNo

Polar Covalent Bonds• Account for the bonding found in HF• One electron from each. atom is shared but

not equally due to unequal attraction (electronegativity) for shared electrons

• The bond is referred to as polar because 2 poles are formed(+ and -)

• Electronegativity values allows us to determine which atom has a greater pull

• The atom with the greater electronegativity becomes the negative end of the polar bond.

• The atom with the lower electronegativity becomes the positive end of the polar bond

H F

Example-HF

FH

e- riche- poor

d+ d-

Electronegativity

• Electronegativity is the tendency of an atom to draw or attract the electrons in a bond toward itself

• Electronegativity is like a game of tug-of-war, atom's ability to pull determines what kind of bond it forms

• To form a covalent bond, two or more atoms with similar electronegativities will share electrons

• The greater the difference in electronegativity the more polar the bond.

Double Covalent Bonds• Compounds sometimes share two

pairs of electrons and form a double bond. This often occurs when two atoms of the same element bond, but also occur between different elements.

• This is called Double Covalent bonds

• Examples: O2, CO2

Triple Covalent Bond• Same idea as single and double• Two atoms of the same element or

two different elements share three pairs of electrons and form a triple bond

• Example: N2

Bond Length & Bond Energy• Typically the more bonds that are shared the

shorter the bond length (closer together the atoms are)

• Bond Energy: the amount of energy required to break the bond

• Bond energy increases as the number of bonds increases

• Bond energy increases when atoms are closer together.

Classifying Bond Type• Non-polar Covalent bonds = less than 0.3 difference in electronegativity

• Polar Covalent bonds = difference less than 1.7

• Ionic Bonds = difference of 1.7 or more.