Chemical Bonds

Force that holds atom together

Topics

• Stability in Bonding– Valence Electrons– Reactivity

• Types of Bonds– Ionic, Metallic, Covalent, Hydrogen bond

• Writing Formulas and Naming Compounds– Binary compound– Polyatomic ion– Hydrate

Reactivity

• Definition: Reactivity describes how likely an element is to form bonds with other elements.

• Reactivity of an element is determined by its valence electrons.

Valence Electrons

• Definition: Valence electrons are electrons in the outermost energy shell (level) that are available to be lost, gained or shared when elements form chemical bonds.

• Elements in the same group of the Periodic Table have similar chemical properties because of their valence electrons.

Element GroupsAnd Valence Electrons

Group Valence Electron

Category Reactivity

1 1 Alkali Metal Very High

2 2 Alkali Earth High

13 3

14 4

15 5

16 6 High

17 7 Halogen Very High

18 8 Noble Gas Stable

Stability in Bonding

• Some elements combine chemically and no longer have the same properties they did before forming a compound.

• A chemical formula is composed of symbols and subscripts indicating the number of atoms of an element in a compound.– E.g. NaCl, H2O, CaCO3, CO2

• Atoms form compounds when the compound is more stable than the separate atoms.– Noble gas are more chemically stable.– Elements that do not have full outer energy

shell are more stable in compounds.– Atoms can lose, gain, or share electrons to

get a stable outer energy level– A chemical bond is the force that holds

atoms together in a compound.

Types of Bonds

Chemical Bonds

Intramolecular

(Strong)

Covalent

Ionic

Metallic

Intermolecular

(Weak)

Hydrogen

Other noncovalent

Intramolecular Force

• Definition: force and its property within the structure of a single molecule.

• Ionic, Covalent, and Metallic bond

Intermolecular Force

• Definition: force that act between stable molecules or between functional groups of macromolecules.

• Example: dipole-dipole interaction, hydrogen bonds, di-sulphide bonds

Cystein-Cystein

Chemical bond

• Definition: The force that holds atoms together in a compound.

• Ions are charged particles because it has more or fewer electrons than protons.– When an atom loses an electron, it becomes

a positively charged ion; a superscript indicates the charge, e.g., H+, Ca2

+, Li+

– When an atom gains an electron, it becomes a negatively charged ion, e.g., HCO3

-, OH-

Ionic Bond

• Definition: Force of attraction between opposite charges of the ions, often forming (crystal) lattice.

• An ionic compound is held together by the ionic bond, often joining a metal to a nonmetal

• The result of this bond is neutral compound.

• The sum of the charges on the ions is zero.

Covalent Bond• Definition: the force of attraction between

atoms sharing electrons.

• Molecules are neutral particles formed as a result of sharing electrons.

• Atoms can form double or triple bonds depending on whether they share two or three pairs of electrons.

• Covalent bonds are usually involved when two nonmetals form a compound or when a nonmetal bond with a metalloid.

• Electrons shared in a molecule are held more closely to the atoms with the larger nucleus.

• A polar molecule has one end that is slightly negative and one end that is slightly positive although the overall molecule is neutral.

• In a nonpolar molecule electrons are shared equally.

Main Difference

• Ionic bonds form when atoms lose or gain electrons; covalent bonds form when atoms share electrons.

Metallic Bond

• electromagnetic interaction between delocalized electrons, called conduction electrons and gathered in an "electron sea", and the metallic nuclei within metals.

• Form metallic bond joins metal to metal.

• This type of bonding is collective in nature. There is no single metallic bond.

Electronegativityaka Electron Affinity

• Electronegativity measures the ability of an atom to attract electrons. Fluorine has the highest electronegativity.

Ionization Energy

• Definition: The amount of energy needed to remove an electron from a neutral atom is called ionization energy.

• Moving from left to right, ionization energy increases across a period.

• Moving from top to bottom in a group, it generally decreases.

Drawing Chemical Bond• Bohr’s Model

– Valence shells are drawn in circles for s suborbital (Group 1, 2) and p suborbital (Group 13 thru 18)

– Protons and Neutrons are indicated in nucleus. Electrons in shells travel in pairs.

• Lewis (Electron) Dot Diagram– Easier way to illustrate electrons on outermost

energy shell and their reactivity.– Charged atoms and molecules are placed in

brackets with their charges indicated at top right corner.

Octet Rule

• Octet rule is a simple chemical rule of thumb that states atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas.

• This rule is applicable to main group elements.

Valence Shells

Bohr’s Model

Bond Using Bohr’s Model

Electron Dot (Lewis) Diagram

Chemical Nomenclature

• Nomenclature: The system of naming in arts or science

• Writing chemical formulas and naming compounds– Binary Metal-Nonmetal Compounds– Binary Nonmetal-Nonmetal Compounds– Polyatomic ions– Hydrate

Oxidation Number

• Oxidation Number – # of e- gained, lost, or shared to become stable

• Determine the oxidation number from the periodic table.

Balancing Oxidation Number

• Ba + F

Binary Metal-Nonmetal Compounds

• Composed of two elements

• When writing binary compounds:– Use the least common multiples of oxidation

number.– Compound is neutral.– Charges on the ions must be balanced.– Use first element, the root name of the second

element, and the suffix -ide.

• If the metal has more than one possible charge, indicate the ion with charge in roman numerals– FeCl2 - Iron (II) chloride.

• Alternatively, common names may be used. Add –ous for the lower charge and –ic for the higher charge:– FeCl2 – ferrous chloride

– FeCl3 – ferric chloride

More examples

Compound Systematic Name Common Name

FeF2 Iron (II) fluoride Ferrous fluoride

FeF3 Iron (III) fluoride Ferric fluoride

Hg2Br2 Mercury (I) bromide

Mercurous bromide

HgBr2 Mercury (II) bromide

Mercuric bromide

Binary Nonmetal-Nonmetal Compounds

• Add –ide to the second element• Use Greek prefixes for number of atoms:

• Examples:– CO – carbon monoxide

– CO2 – carbon dioxide

– N2O5 – dinitrogen pentoxide

1 – mono 2 – di 3 – tri 4 – tetra 5 – penta

6 – hexa 7 – hepta 8 – octa 9 – nona 10 - deca

• Common names: –ous and –ic (-ic has greater charge, OR has fewer atoms)

Formula Systematic Name Common Name

NO Nitrogen monoxide Nitric oxide

N2O Dinitrogen monoxide Nitrous oxide

NO2 Nitrogen dioxide Nitrogen peroxide

N2O5 Dinitrogen pentoxide Nitric anhydride

N2O3 Dinitrogen trioxide Nitrous anhydride

Polyatomic Ion

• The compound contains three or more elements.

• When writing names:– Positive ions first, followed by negative ion– Use the least common multiple of oxidation

number, and put parenthesis around the polyatomic ion before adding a subscript

Balancing Oxidation Number

• K + OH Ca + NO3

Polyatomic Compounds• Ammonium ion NH4

1+

• -ide ions: CN1- cyanide, OH1- hydroxide• Oxyanions

– -ate are more oxygen

• NO21- nitrite

• NO31- nitrate

– Some oxyanions have extra hydrogen

• SO42- sulfate

• HSO41- hydrogen sulfate (or bisulfate)

• Oxyanions (continued)

– If more than two possibilities

Formula Name

ClO1- Hypochlorite

ClO21- Chlorite

ClO31- Chlorate

ClO41- Perchlorate

Naming compounds with polyatomic ions

• Positive charge species on left• Negative charge species on right• Use parenthesis as needed

Formula Ions Name

BaSO4 Ba2+ and SO42- Barium sulfate

Ca(NO3)2 Ca2+ and NO31- Calcium nitrate

Ca(NO2)2 Ca2+ and NO21- Calcium nitrite

Fe(NO3)2 Fe2+ and NO31- Iron (II) nitrate or

ferrous nitrate

Hydrate• Compound with water chemically attached

to its ions

• Two types:

– Hydro Acids: Hydro + Halogen name + ic

• HCl – hydrochloric acid

• HF – hydrofluoric acid

– Oxoacids: polyatomic ion + acid

• HNO3 – nitric acid, from nitrate

• HNO2 – nitrous acid, from nitrite