Chap 4 Forces Between Particles: Ions, Ionic Compounds, and Covalent Bonding

Forces Between Particles

Objectives

This section introduces the fundamental concepts of ionic and covalent bonding, the formation and naming of ions and compounds, and the prediction of molecular polarity. Mastery of these topics is essential for understanding chemical interactions in GOB Chemistry.

• Use electronic configurations to determine the number of electrons gained or lost by atoms to achieve an octet of electrons.

• Apply the octet rule to predict the ions formed during the formation of ionic compounds.

• Write correct formulas for binary ionic compounds containing a representative metal and a representative nonmetal.

• Name binary ionic compounds.

• Calculate formula weights for ionic compounds.

• Use molecular shapes to predict the polarity of a molecule.

• Use electronegativities to classify covalent bonds and determine whether molecules are polar or nonpolar.

• Write correct formulas for ionic compounds containing representative metals and polyatomic ions.

• Correctly name binary covalent compounds and compounds containing polyatomic ions.

Ions and Ion Formation

Formation of Ions

Atoms can gain or lose electrons to achieve a stable electron configuration, often resembling the nearest noble gas. This process results in the formation of ions, which are charged particles.

• Cation: An atom that loses one or more electrons, resulting in a positive charge. Example: Sodium (Na) loses one electron to form Na+.

• Anion: An atom that gains one or more electrons, resulting in a negative charge. Example: Chlorine (Cl) gains one electron to form Cl-.

• Valence electrons: The electrons in the outermost shell, which determine chemical reactivity. For main group elements, the group number indicates the number of valence electrons (e.g., Group 3A = 3 valence electrons).

• Octet rule: Atoms tend to gain, lose, or share electrons to achieve eight electrons in their valence shell.

Example: Sodium atom (Na) loses one electron to become Na+; Chlorine atom (Cl) gains one electron to become Cl-.

Predicting Ion Charges

The charge of an ion can be predicted based on its group in the periodic table:

• Group 1A: +1 charge (e.g., Na+)

• Group 2A: +2 charge (e.g., Mg2+)

• Group 3A: +3 charge (e.g., Al3+)

• Group 5A: -3 charge (e.g., N3-)

• Group 6A: -2 charge (e.g., O2-)

• Group 7A: -1 charge (e.g., Cl-)

• Group 8A: 0 charge (noble gases, already have a full octet)

Additional info: Transition metals can form multiple cations with different charges.

Naming Ions

Cations

Cations are named by stating the element name followed by the word "ion." For elements that form only one type of cation (Groups 1A, 2A, 3A):

Group 1A	Group 2A	Group 3A
Hydrogen ion (H+) Lithium ion (Li+) Sodium ion (Na+) Potassium ion (K+)	Magnesium ion (Mg2+) Calcium ion (Ca2+) Strontium ion (Sr2+) Barium ion (Ba2+)	Aluminum ion (Al3+)

Multivalent Cations (Transition Metals)

Some metals can form more than one type of cation. Their charge is indicated by a Roman numeral in parentheses.

Ion	Name
Cu+	Copper (I) ion
Cu2+	Copper (II) ion
Fe2+	Iron (II) ion
Fe3+	Iron (III) ion
Pd4+	Palladium (IV) ion
V5+	Vanadium (V) ion

Anions

Monatomic anions are named by taking the stem of the element name and adding the suffix "-ide."

Anion	Stem Name	Anion Name
F-	fluor	Fluoride
Cl-	chlor	Chloride
Br-	brom	Bromide
I-	iod	Iodide
O2-	ox	Oxide
S2-	sulf	Sulfide

Polyatomic Ions

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying an overall charge. They are treated as a single unit in chemical formulas.

Ion	Name
NH4+	Ammonium
OH-	Hydroxide
CN-	Cyanide
CH3COO-	Acetate
NO3-	Nitrate
NO2-	Nitrite
HCO3-	Hydrogen carbonate (bicarbonate)
CO32-	Carbonate
HSO4-	Hydrogen sulfate (bisulfate)
SO42-	Sulfate
PO43-	Phosphate

Additional info: Memorization of common polyatomic ions is recommended for success in GOB Chemistry.

Chemical Bonds

Ionic Bonds

An ionic bond is formed by the electrostatic attraction between a cation and an anion. Typically, a metal transfers electrons to a nonmetal, resulting in the formation of oppositely charged ions.

• Written as cation first, followed by the anion (e.g., NaCl).

• The total positive and negative charges must balance to give an overall charge of zero.

Example: Magnesium loses two electrons to form Mg2+; Oxygen gains two electrons to form O2-; the formula is MgO.

Writing Formulas for Ionic Compounds

To write the formula for an ionic compound:

1. Write the symbol and charge for the cation and anion.

2. Criss-cross the charges to become subscripts for the opposite ion.

3. Reduce subscripts to the simplest whole-number ratio.

4. For polyatomic ions, use parentheses if more than one is needed (e.g., Ca(NO3)2).

Example: Calcium ion (Ca2+) and nitrate ion (NO3-): $\text{Ca}^{2+} + 2\text{NO}_3^- \rightarrow \text{Ca(NO}_3)_2$

Naming Ionic Compounds

• Name the cation first (drop the word "ion"), then the anion.

• For transition metals, indicate the charge with a Roman numeral.

• For polyatomic ions, use the ion's name as is.

Examples:

• NaCl: sodium chloride

• CuCl2: copper (II) chloride

• FeSO4: iron (II) sulfate

• (NH4)2SO4: ammonium sulfate

Covalent (Molecular) Compounds

Covalent Bonds

A covalent bond is formed when two nonmetals share one or more pairs of electrons to achieve a full valence shell.

• Each shared pair of electrons is represented by a single line (e.g., H—H).

• Compounds formed are called molecular compounds.

Example: Water (H2O) is formed by sharing electrons between hydrogen and oxygen.

Naming Binary Covalent Compounds

• Name the less electronegative element first, then the more electronegative element with the suffix "-ide."

• Use prefixes to indicate the number of each atom:

Prefix	Number
mono-	1
di-	2
tri-	3
tetra-	4
penta-	5
hexa-	6
hepta-	7
octa-	8
nona-	9
deca-	10

Examples:

• CO2: carbon dioxide

• SF6: sulfur hexafluoride

• PBr3: phosphorus tribromide

• ICl: iodine monochloride

Bond Polarity and Molecular Polarity

Electronegativity and Bond Polarity

Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond.

• Nonpolar covalent bond: Electrons are shared equally (electronegativity difference < 0.5).

• Polar covalent bond: Electrons are shared unequally (electronegativity difference between 0.5 and 1.9).

• Ionic bond: Electrons are transferred (electronegativity difference > 1.9).

Example: In H—Cl, chlorine is more electronegative and pulls the shared electrons closer, making the bond polar.

Molecular Polarity

The shape of a molecule and the distribution of polar bonds determine whether a molecule is polar or nonpolar.

• Nonpolar molecule: Symmetrical distribution of electron density (e.g., CO2).

• Polar molecule: Asymmetrical distribution of electron density (e.g., H2O).

• Polarity is indicated by partial charges: $\delta^+$ (less electronegative) and $\delta^-$ (more electronegative).

Example: Ammonia (NH3) is a polar molecule due to its trigonal pyramidal shape and the presence of a lone pair on nitrogen.

Molecular Shapes

Common Molecular Geometries

The shape of a molecule affects its physical and chemical properties, including polarity.

• Linear: Atoms arranged in a straight line (e.g., CO2).

• Trigonal planar: Three atoms arranged around a central atom in a plane (e.g., BF3).

• Tetrahedral: Four atoms arranged around a central atom (e.g., CH4).

• Trigonal pyramidal: Three atoms and one lone pair around a central atom (e.g., NH3).

• Bent: Two atoms and two lone pairs around a central atom (e.g., H2O).

Additional info: The VSEPR (Valence Shell Electron Pair Repulsion) theory is used to predict molecular shapes.