PSMA (General Chemistry) pt. 2

Chemical Bonds

• are forces of attraction that exist between a positive ion and a negative ion or between molecules

Octet Rule

• an atom other than hydrogen tends to form bonds until it is surrounded by eight electrons

According to Gilbert Lewis

• atoms combine to achieve a more stable electron configuration.

• Maximum stability results when an atom is isoelectronic with a noble gas.

• Ionic bond

• Covalent bond

• Metallic bond

Three Types of lnteratomic Bonds

Ionic Bond

• Or electrovalent bond

• the electrostatic force that holds ions together in an ionic compound

• farmed by the transfer of electrons from an atom of low ionization energy (alkali and alkaline earth metal) to a more electronegative element (halogen and oxygen)

Covalent Bond

• bond in which two electrons are shared between atoms

• formed when electrons are shared equally between two atoms with the same or almost the same electronegativity

• Non-bonding pairs are called lone pairs.

• Lone pairs - pairs of valence electrons that are not involve in covalent bond formation

Polar Covalent Bond

• This is formed from sharing of electrons between two atoms with different electronegativity

ELECTRONEGATIVITY

• the ability of an atom to attract toward itself the electrons in a chemical bond.

• property that helps us distinguish a nonpolar covalent bond from a polar covalent bond

Multiple Bonds

• Atoms can form multiple covalent bonds

  • Single bond

  • Double bond

  • Triple bond

Single bond

two atoms are held together by one electron pair

Double Bond

two atoms share two pairs of electrons

Triple Bond

two atoms share three pairs of electrons

Coordinate Covalent Bond

A bond formed wherein one furnishes both the bonding pair of electrons

Metallic Bond

consists of group of cations held in a fixed position in the metal and the valence electrons which are free to move about among the different electron clouds

Intermolecular forces

attractive forces between molecules

• Van der Waals

• Hydrogen Bond

• Dipole Bond

• Resonance

4 Intermolecular forces

Van der Waals

Very weak forces of attraction between non-polar molecules

Hydrogen Bond

is an attractive force that exists between the hydrogen of one molecule and the more electronegative portion of another molecule (such as O, F, and N)

Dipole Bond

is an attractive force between polar molecules

Resonance

The use of two or more Lewis Structures to represent a particular molecule

Resonance Structure

is one of the two or more Lewis Structures for a single molecule that cannot be represented accurately by only one Lewis Structure

• Dipole-dipole (Keesom Forces)

• Dipole-induced dipole Forces

• London Dispersion Forces (Induced Dipole–Induced Dipole)

3 Types of Van der Waals Forces

Dipole-dipole (Keesom Forces)

Two polar molecules align so that OPPOSITE CHARGES are matched (electrostatic attraction)

Dipole-induced dipole Forces

• A dipole can induce (cause) a temporary dipole to farm in a nonpolar molecule

• The molecules then line up to match opposite charges

London Dispersion Forces (Induced Dipole–Induced Dipole)

• A temporary dipole farms in a nonpolar molecule which leads to a temporary dipole to farm in ANOTHER non-polar molecule

• Dispersion is the ONLY intermolecular attraction that occurs between non-polar molecules

MOLECULAR FORMULA

The molecular formula of an organic compound simply shows the number of each type of atom present. It tells you nothing about the bonding within the compound.

EMPIRICAL FORMULA

The empirical formula of an organic compound gives the simplest possible whole number ratio of the drfferent types of atom within the compound.

CONDENSED FORMULA

The condensed formula is also text-based; here, each carbon atom is listed separately, with atoms attached to it following. An exception is cyclic parts of molecules. e.g. benzene, where the carbons are grouped.

DISPLAYED FORMULA

A formula that shows all of the atoms and all of the bonds present in an organic compound. The bonds are represented as lines

STRUCTURAL FORMULA

Similar to displayed formula • not all bonds are shown, although all atoms are still indicated using subscript numbers. Carbon hydrogen bonds are often simplified

SKELETAL FORMULA

in this formula, most hydrogen atoms are omitted, and line ends or vertices represent carbons. Functional groups and atoms other than carbon or hydrogen are still shown. Easiest to draw & commonly used

Binary Ionic Compounds

contain a positive ion (cation) always written first in the formula and a negative ion (anion)

Binary Ionic Compounds (Type I)

1. The cation is always named first and the anion second.

2. A monatomic (meaning "one-atom") cation takes its name f ram the name of the element. For example, Na is called sodium in the names of compounds containing this ion.

3. A monatomic anion is named by taking the root of the element name and adding -ide. Thus the Cl ion is called chloride.

Ions Present: Na+, Cl-

Name: Sodium chloride

Compound

NaCl

Ions Present: K+, I-

Name: Potassium iodide

Compound

KI

Ions Present: Ca2+, S2-

Name: Calcium sulfide

Compound

CaS

Ions Present: Li+, N3-

Name: Lithium nitride

Compound

Li3N

Ions Present: Cs+, Br-

Name: Cesium bromide

Compound

CsBr

Ions Present: Mg2+, O2-

Name: Magnesium oxide

Compound

MgO

Hydrogen

Common Monatomic Cations:

H+

Lithium

Common Monatomic Cations:

Li+

Sodium

Common Monatomic Cations:

Na+

Potassium

Common Monatomic Cations:

K+

Cesium

Common Monatomic Cations:

Cs+

Beryllium

Common Monatomic Cations:

Be2+

Magnesium

Common Monatomic:

Mg2+

Calcium

Common Monatomic Cations:

Ca2+

Barium

Common Monatomic Cations:

Ba2+

Aluminum

Common Monatomic Cations:

Al3+

Hydride

Common Monatomic Anions:

H-

Fluoride

Common Monatomic Anions:

F-

Chloride

Common Monatomic Anions:

Cl-

Bromide

Common Monatomic Anions:

Br-

Iodide

Common Monatomic Anions:

I-

Oxide

Common Monatomic Anions:

O2-

Sulfide

Common Monatomic Anions:

S2-

Nitride

Common Monatomic Anions:

N3-

Phosphide

Common Monatomic Anions:

P3-

Binary Ionic Compounds (Type II)

• There are many metals that form more than one type of positive ion and thus form more than one type of ionic compound with a given anion.

• For example, the compound FeCl2 contains Fe2+ ions, and the compound FeCl3 contains Fe3+ ions.

• In a case such as this, the charge on the metal ion must be specified.

• The systematic names for these two iron compounds are iron(II) chloride and iron(III) chloride, respectively, where the Roman numeral indicates the charge of the cation.

• Another system for naming these ionic compounds that is seen in the older literature was used for metals that form only two ions.

• The ion wjth the hjgher charge has a name ending in -ic, and the one with the lower charge has a name ending in -ous.

• In this system, for example, Fe3+ is called the ferric ion, and Fe2+ is called the ferrous ion. The names for FeCl3 and FeCl2 are then ferric chloride and ferrous chloride, respectively

Iron (III)

Binary Ionic Compounds (Type II)

Fe3+

Iron (II)

Binary Ionic Compounds (Type II)

Fe2+

Copper (II)

Binary Ionic Compounds (Type II)

Cu2+

Copper (I)

Binary Ionic Compounds (Type II)

Cu+

Cobalt (III)

Binary Ionic Compounds (Type II)

Co3+

Cobalt (II)

Binary Ionic Compounds (Type II)

Co2+

Tin (IV)

Binary Ionic Compounds (Type II)

Sn4+

Tin (II)

Binary Ionic Compounds (Type II)

Sn2+

Lead (IV)

Binary Ionic Compounds (Type II)

Pb4+

Lead (II)

Binary Ionic Compounds (Type II)

Pb2+

Mercury (II)

Binary Ionic Compounds (Type II)

Hg2+

Mercury (I)

Binary Ionic Compounds (Type II)

Hg2²+

Silver

Binary Ionic Compounds (Type II)

Ag+

Zinc

Binary Ionic Compounds (Type II)

Zn2+

Cadmium

Binary Ionic Compounds (Type II)

Cd2+

Binary Ionic Compounds (Type I)

Does the compound contain Type I or Type II cations?

Name the cation using the element name

Binary Ionic Compounds (Type II)

• Using the principle of chargebalance, determine the cation charge.

• Include in the cation name a Roman numeral indicating the charge

Binary Ionic Compounds

• This case most commonly occurs for compounds containing transition metals, which often farm more than one cation.

• Elements that farm only one cation do not need to be identified by a Roman numeral.

• Common metals that do not require Roman numerals are the Group 1 A elements, which form only +1 ions; the Group 2A elements, which form only +2 ions; and aluminum, which forms only Al3 +

• In virtually all its compounds silver is found as the Ag+ 1 ion. Therefore, although silver is a transition metal (and can potentially form ions other than Ag), silver compounds are usually named without a Roman numeral. Thus AgCl is typically called silver chloride rather than silver(I) chloride, although the latter name is technically correct.

• Also, a Roman numeral is not used for zinc compounds, since zinc forms only the Zn+2 ion

Binary Ionic Compounds

• This case most commonly occurs for compounds containing __, which often farm more than one cation.

• Elements that farm only __ do not need to be identified by a Roman numeral.

• Common metals that do not require Roman numerals are the , which form only +1 ions; the , which form only +2 ions; and aluminum, which forms only Al3 +

• In virtually all its compounds silver is found as the Ag+ 1 ion. Therefore, although silver is a transition metal (and can potentially form ions other than Ag), __ are usually named without a Roman numeral. Thus AgCl is typically called silver chloride rather than silver(I) chloride, although the latter name is technically correct.

• Also, a Roman numeral is not used for __, since zinc forms only the Zn+2 ion

Ionic Compounds with Polyatomic Ions

• We have not yet considered ionic compounds that contain polyatomic ions. For example, the compound ammonium nitrate, NH4NO3, contains the polyatomic ions NH4+ and NO3-

• Polyatomic ions are assigned special names that must be memorized to name the compounds containing them

• Several series of anions contain an atom of a given element and different numbers of oxyger) atoms. These anions are called oxyanions.

• When there are two members in such a series, the name of the one with the smaller number of oxygen atoms ends in -ite and the name of tfle one with the larger number ends in -ate.

• For example, sulfite (SO3²-) and sulfate (SO4²-)

• When more than two oxyanions make up a series, hypo- (less than) and per(more than) are used as prefixes to name the members of the series with the fewest and the most oxygen atoms, respectively

Mercury (I)

Common Polyatomic Ions

Hg2²+

Ammonium

Common Polyatomic Ions

NH4+

Nitrite

Common Polyatomic Ions

NO2-

Nitrate

Common Polyatomic Ions

NO3-

Sulfite

Common Polyatomic Ions

SO3²-

Sulfate

Common Polyatomic Ions

SO4²-

Hydrogen sulfate (bisulfate)

Common Polyatomic Ions

HSO4-

Hydroxide

Common Polyatomic Ions

OH-

Cyanide

Common Polyatomic Ions

CN-

Phosphate

Common Polyatomic Ions

PO4³-

Hydrogen phosphate

Common Polyatomic Ions

HPO4²-

Dihydrogen phosphate

Common Polyatomic Ions

H2PO4-

Thiocyanate

Common Polyatomic Ions

NCS- or SCN-

Carbonate

Common Polyatomic Ions

CO3²-

Hydrogen carbonate (bicarbonate)

Common Polyatomic Ions

HCO3-

Hypochlorite

Common Polyatomic Ions

ClO- or OCl-

Chlorite

Common Polyatomic Ions

ClO2-

Chlorate

Common Polyatomic Ions

ClO3-

Perchlorate

Common Polyatomic Ions

ClO4-

Acetate

Common Polyatomic Ions

C2H3O2-

Permanganate

Common Polyatomic Ions

MnO4-

Dichromate

Common Polyatomic Ions

Cr2O7²-