Bond

Ionic Bond“Electrons Taken not Shared”

Formation of Ionic Compounds• Ionic compounds are composed of cations (metals) and anions (nonmetals).

• Although they are composed of ions, ionic compounds are electrically neutral. The positive charge equals the negative.

• The positive charge and the negative charge attract one another. This is how the bond is formed.

Formula Units• Chemists represent the composition of substances

by writing a chemical formula. A chemical formula shows the kinds and number of atoms in a substance.

• Ionic compounds exists as collection of negative and positive charge arranged in repeated patterns.

• The chemical formula for an ionic compound, the formula unit, represents the smallest whole number ratio.

What is the overall charge on an ionic compound

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1. Positive

2. Negative

3. Neutral

How are ionic bonds formed

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1. Electrons are given away

2. Electrons are taken

3. Electrons are shared

4. 1 and 2

What two elements below will form an ionic bond

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1. Cl and Br

2. K and He

3. H and O

4. Li and Cl

What two elements below will form an ionic bond

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1. I and Na

2. Mg and Ne

3. P and O

4. 1 and 2

Properties of Ionic Compounds• Most ionic compounds are crystalline solids at room temperature

• Ionic compounds generally have high melting points.

• Ionic Compounds can conduct electricity when melted or dissolved. This is because ionic compounds break down into their IONS when melted or dissolved in water.

Ionic bonds are the strongest type of bond.

Which is NOT a property of ionic compounds?

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1. The conduct electricity in solutions

2. They will not break apart in water to form ions

3. They have low conductivity

4. 2 and 3

Metallic Bonds

What are Metals?

• A metal is an element that readily forms positive ions (cations) and has metallic bonds.

What are Metals?

• On the periodic table, a diagonal or stair step line drawn from boron (B) to polonium (Po) separates the metals from the nonmetals. Elements on this line are metalloids, sometimes called semi-metals; elements to the lower left are metals; elements to the upper right are nonmetals.

• Almost 80% of the elements on the periodic table are metals.

Metallic Bonds• The valence electrons of metal atoms can

drift freely from one part of the metal to another- this is sometimes

called a “sea of electrons” • Metallic bonds consist of the attraction between these free floating electrons and the positively charged metal ions (cations). This

attraction is the “bond” that holds metals together.

Metallic bonds only form between two “pure” metals – no metalloids.

Metallic bonds are t

the SECOND

strongest bond

type.

Physical Properties of Metals• Lusterous- they are shiny!

• High density- atoms are tightly packed.

• Good conductors of electricity and heat.– Reason- electrons can flow freely.

Physical Properties of Metals

• Ductile- they can be drawn into wires AND

• Malleable- they can be hammered into shapes– Reason- cations can slide easily past each

other because the sea of electrons insulates them and prevents strong repulsions.

…on the other hand

• Ionic compounds are brittle and break easily? Why

Physical Properties of Metals• Metal ions are arranged in very compact

orderly patterns.

–Similar to the way apples are stacked at the grocery store.

• Pure metals form the simplest kinds of crystals

Chemical Properties of Metals• Most metals are chemically unstable and

will react will oxygen in the air- that is they form oxides- over varying timescales (for example iron rusts over years and potassium burns in seconds).

Chemical Properties of Metals

-The alkali metals react quickest followed by the alkaline earth metals.

-The transition metals take much longer to oxidize (such as iron, copper, zinc, nickel). Others, like palladium, platinum and gold, do not react with the atmosphere at all.

-Some metals form a barrier layer of oxide on their surface which cannot be penetrated by further oxygen molecules and thus retain their shiny appearance and good conductivity for many decades (like aluminium, some steels, and titanium).

Alloys• Very few metals that you encounter daily are pure

metals. Most metals are alloys, a mixture of two or more elements of which at least one is a metal.– Examples:

• Brass is an alloy of copper and zinc• Sterling silver is an alloy of silver (92%) and copper (8%)• Stainless Steel is an alloy of iron (81%), chromium (18%), nickel

(1%), and trace amounts of carbon.

• Alloys are important because their properties are often superior to those of their component elements. – Examples:

• Sterling silver is harder and more durable than pure silver, but still soft enough to make jewelry and tableware.

• Brass is harder and easier to shape than either copper or zinc

Covalent Bonds – The nice bonds that share!

What do covalent bonds share?

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1. Electrons

2. Protons

3. Neutrons

4. Their cookies at lunch

Molecules and Molecular Compounds

• Compounds that are NOT held together by an electrical attraction, but instead by a sharing of electrons.

• Atoms held together by a sharing of electrons are joined by a covalent bond.

H• •H

Molecules and Molecular Compounds

• A molecule is a neutral group of atoms joined together covalent bonds. A compound composed of molecules is called a molecular compound.

• The chemical formula for a molecule is called the molecular formula.

Properties of Molecular Compounds

• Composed of two or more nonmetals.

• Usually gases or liquids at room temperature.

Properties of Molecular Compounds

• Molecular compounds tend to have a relatively lower melting and boiling point than ionic compounds.– Reason: There are no (or few and weak)

bonds holding the molecules together in molecular compounds.

Ionic CompoundMolecular Compound

Properties of Molecular Compounds

• Do not conduct electricity. They form nonelectrolytes in solution.– Reason: Molecular compounds do not break

apart into ions in solutions.

Polar covalent bonds are the most common bond type!

2 Types of Covalent Bonds – Polar and Nonpolar

1. Polar Covalent Bond – unequal sharing of electrons - When the electronegativity difference between the bonding atoms are greater than 0.3, a polar covalent bond results (one atom can attract electrons better, and pulls the electrons closer).

Electronegativity Table

• Example: HCl, H2O

• Even though the electrons are shared, the fact that the electrons are more attracted to the Cl atom results in a partial negative charge.

• The unequal sharing of electrons results in the formation of partial positive charges and partial negative charges.

Polar covalent bonds are weaker than both ionic and metallic, but stronger than non-polar covalent bonds.

• Non-Polar Covalent Bond – equal sharing of electrons

• Example: H2

• Electrons are equally shared by the 2 atoms

• The bonding for a compound can be nonpolar if the molecule shows symmetry.

Non-polar covalent bonds are the weakest type of bonds!

How do we know the type of bond that will be formed?

• Simplistically, Ionic bonds occur when electrons are given and taken, and covalent bonds occur when electrons are shared. However… there is a bit more to the story. There’s really a spectrum of bonding, from taking to partial sharing, to unequal sharing, to complete sharing of valence electrons.

Bond Polarity

• Most bonds are a blend of ionic and covalent characteristics.

• Difference in electronegativity determines bond type.

What determines what really happens to electrons in a

chemical bond?

Electronegativity of an atom determines what happens to electrons in a bond! (Remember electronegativity = the ability of an atom to attract electrons)

To determine the bond type, you have to find the electronegativity difference between the atoms.

How do I find Electronegativity difference?

Use your electronegativity chart to determine the difference of electronegativity between atoms (remember to subtract big-small)

Electronegativity Table

Bond Polarity

Examples:

• NaCl

• HCl

• Cl2

3.0-0.9=2.1Ionic

3.0-2.1=0.9Polar

3.0-3.0=0.0Non-Polar