Unit 4 Ionic and Covalent Bonding Ionic and Covalent Compounds Molecular Structure

Polyatomic Ions you NEED to know Carbonate Sulfate Nitrate Hydroxide Phosphate Ammonium

Naming Ionic Compounds Name consists of metal cation first, followed by non-metal anion. Cation is always named “as is.” It is simply the name of the metal as it appears on the periodic table or if NH4, it’s ammonium. The anion name is as follows: If the anion is a monoatomic ion, the name consists of the root of the element’s name followed by the suffix -ide. If the anion is a polyatomic ion, the name is “as is.” It is simply the name of the polyatomic ion. If the metal in the compound is a transition metal (from Groups 3 thru 12), then the name of the metal must be followed by a roman numeral to indicate the charge number on the metal.

Naming Ionic Compounds When ionic compounds form, the positive metal cations attract to the negative non- metal ions in a ratio that ensures a neutral compound. In other words, the positive charge of the cation(s) has to be equal to the negative charge of the anion(s). The ionic formula shows the simplest whole number ratio of the metal and non-metal ions.

Criss-Cross method for writing formulas The criss-cross method can be used to determine the ionic formula. Using this method, write the cation with charge, then write the anion with charge next to it. Bring down the charge number  of the anion and make it the subscript  (lower number) of the cation.  Then, bring down the charge number of the cation and make  it the subscript (lower number) of the anion. In chemistry, a subscript of 1 is eliminated. No subscript indicates that there is only one ion in the formula. Also, if there is an equal number of cations and anions in a formula, it reduces to a 1:1 ratio and the subscripts are eliminated. Finally, when there is more than one polyatomic ion in a formula, you must write the polyatomic ion in parenthesis and follow it with a subscript to indicate the number of units in the formula.

Examples-Try These! Strontium phosphate Cu(OH)2 Calcium sulfide Al(NO3)3

CHARACTERISTICS OF COVALENT COMPOUNDS Formed between non-metals and non-metals. Atoms in covalent compounds share electrons to achieve a stable configuration of eight valence electrons (octet). The chemical bond that results from the sharing of valence electrons is a covalent bond. Covalent bonding generally occurs when elements are relatively close to each other on the periodic table. (non-metals on right side of PT) When 2 or more atoms bond covalently, a molecule is formed. Ex. Glucose: C6H12O6: Methane: CH4 ; Carbon Dioxide: CO2

Types of Covalent Bonds 1. Single-Occurs when each atom in the bond requires one additional electron to achieve an octet; thus, one pair of electrons is shared between the atoms, forming a single bond. Ex. H-H, F-F, Br-Br Bond length (distance between bonding nuclei) is longest in a single covalent bond and thus, a single bond is the weakest.

Types of Covalent Bonds 2. Double-Occurs when each atom in the bond requires two additional electrons to achieve an octet; thus, two pairs of electrons are shared between the atoms, forming a double bond. Ex. 0-O Bond length (distance between bonding nuclei) is shorter than in a single bond and thus, a double bond is stronger than a single bond.

Types of Covalent Bonds 3. Triple-Occurs when each atom in the bond requires three additional electrons to achieve an octet; thus, three pairs of electrons are shared between the atoms, forming a triple bond. Ex. N-N Bond length (distance between bonding nuclei) is shortest in a triple bond and thus, a triple bond is the strongest.

Characteristics of Covalent Compounds Molecular substances exist as gases, liquids that vaporize readily at room temperature (ex. gasoline or alcohols) or as relatively soft solids (ex. wax). Covalent solids are more suitable for insulating due to their softness and malleability Covalent bonds between atoms are very strong; however, attractions between individual molecules (intermolecular forces) are relatively weak when compared to ionic compounds; thus, covalent substances have lower melting and boiling points when compared to ionic compounds. Many molecular compounds dissolve in solvents but do not dissociate into charged ions since they are composed of neutral atoms that share electrons; thus, solutions containing molecular substances do not conduct electricity and are called non-electrolytes.

Naming Binary Covalent Compounds Molecular (covalent) compounds are formed when two or more nonmetals share electrons. The first element in the formula appears first in the name and is preceded by a prefix only if there is more than one atom present in the molecule. In other words, a covalent compound name will never begin with the prefix “mono-.” The second element in the formula appears second in the name and is always preceded with a prefix to indicate the number of atoms in the formula. The second part of the name consists of the root name of the element followed by the suffix “- ide.”

Important Prefixes (MUST know!) Number of Atoms Mono- 1 Hexa- 6 Di- 2 Hepta- 7 Tri- 3 Octa- 8 Tetra- 4 Nona- 9 Penta- 5 Deca- 10 Examples: SO2 sulfur dioxide P2O5 diphosphorous pentoxide NH3 nitrogen trihydride (common name ammonia) PCI3 phosphorous trichloride N2O dinitrogen monoxide

Electronegativity and Bond Polarity Electronegativity is a measure of an atom’s ability to attract another atom’s electrons in a chemical bond. Electronegativity increases across a period (row) on the periodic table as the number of protons in the nucleus of each atom increases, thus increasing the positive nuclear charge or “pull” on negatively charged electrons. Electronegativity decreases down a group (column) on the periodic table as the distance between the nucleus and the outermost energy levels increases, thus reducing the nuclear “pull” on electrons. NOTE: Electronegativity values are not assigned to the noble gases since they do not typically bond with other atoms. The most electronegative element is Fluorine and the least is Cesium.

Electronegativity and Bond Polarity A non-polar covalent bond occurs between identical atoms (same electronegativity values) or between atoms with very similar electronegativity values (difference 0- 0.4). This results in equal sharing of electrons. Ex.  H-H    O=O    N-N    H-P

Electronegativity and Bond Polarity A polar covalent bond occurs between atoms with different electronegativity values (difference 0.5 to 1.9). This results in unequal sharing of electrons. The more electronegative atom in the bond attracts the shared electrons more, resulting in a partial negative charge on that atom (6”) and a partial positive charge on the less electronegative atom (6’). NOTE: Do not confuse polar covalent bonds with ionic bonds. Ionic bonds occur between metal and non-metal atoms with a large difference in electronegativity values (2.0 or higher). Metal atoms completely transfer their electrons to non- metal atoms to form fully charged ions which then attract in an ionic bond.

Molecular Shapes Shapes are determined by the VSEPR theory (Valence Shell Electron Pair Repulsion). Shapes are determined by the number of atoms that surround a central atom in a molecule. Shapes are determined by the presence or absence of unshared pairs of electrons on the central atom in a molecule (lone pairs). Unshared electrons on a central atom repel each other (due to their like charges) and cause the molecule to bend downwards.

Molecular Shapes-MUST know! Linear: 2 atoms bonded together. 2 atoms bonded to a central atom with no lone pairs on the central atom. Bent: 2 atoms surrounding a central atom with lone pairs on the central atom. Trigonal Planar: 3 atoms surrounding a central atom with no lone pairs on the central atom. Trigonal Pyramidal: 3 atoms surrounding a central atom with lone pairs of electrons on central atom. Tetrahedral: 4 atoms surrounding a central atom with no lone poirs of electrons on central atom.

Examples Formula BeCl2 BCl3 CH4 NH3 H2O Name Beryllium chloride Boron trichloride methane ammonia water Bonding Pairs 2 3 4 Valence Electrons 5 6 Lone Pairs 1 Angles between bonding pairs 180 120 109.5 107 105 Name of shape Linear Trigonal Planar Tetrahedral Bent

Polarity in Molecules/Solubility Polarity in molecules is determined by the types of bonds between atoms in a molecule and the overall symmetry of the molecule. A non-polar molecule has a symmetrical shape and charge distribution, resulting in a molecule that contains “even” attraction of electrons on all sides. A  polar molecule has an asymmetrical shape with an uneven charge distribution, resulting in a molecule that contains “uneven” attraction of electrons on its sides. These molecules have a partial positive side (6") and partial negative side (6 ).

Polarity in Molecules/Solubility In general, “like dissolves like.” Polar molecules dissolve in polar solvents. Non-polar molecules dissolve in non-polar solvents. Ionic compounds dissolve in polar solvents. Alcohols (can be polar or non-polar) tend to dissolve both polar and non-polar compounds but do not dissolve ionic compounds.

Intermolecular Forces (IMFs) Ionic and covalent bonds within a compound are strong but the attraction between molecules (intermolecular forces) is weaker.  The type of attractive force between molecules determines whether a molecular compound is a gas, liquid or solid at a given temperature

Types of IMFs London Dispersion Forces: WEAKEST IMF between non-polar molecules Do not have partial charges, so they are not attracted to one another due to differentials Attraction is due to momentary shifts in the position of shared electrons Ex. F-F   Br-B         0=C=O     CCI4

Types of IMFs Dipole Interactions: Medium strength intermolecular force between polar molecules. These molecules have partial positive and partial negative sides and the partial positive side of one molecule is attracted to the partial negative side of the neighboring molecule. Ex. HCl, PBr3 CH3OH

Types of IMFs Hydrogen Bonds: STRONGEST IMF between "super" polar molecules Defined as those containing Hydrogen bound to the most electronegative atoms on the periodic table (Nitrogen, Oxygen, and Fluorine).  The partial charges in these bonds are more dramatic and molecules containing these bonds have stronger partial positive and partial negative sides, resulting in the greatest attraction between molecules. Ex. NH3 H2O  H-F

Characteristics of Acids Acids are a special class of covalent compounds that contain hydrogen (non-metal) bound to other non-metals. Acids differ from other covalent compounds in that the hydrogen ionizes in solution to produce the H* cation. Acids can be distinguished from other covalent compounds in that their formulas almost always start with hydrogen. The hydrogen is combined with other non-metal anions, with a formula resembling that of an ionic compound.

Examples of dissociation equations Bromic acid: HBr Dissociation in Solution: HBr -->  H+ + Br-  Nitric acid: HNO3 Dissociation in Solution:  HNO-->  H+ + NO3-  Sulfuric acid: H2SO4 Dissociation in Solution: H2SO4 --> 2H+ + SO4-2 Acids produce charged ions in solution; thus, they conduct electricity and are classified as electrolytes.

Bonding Energy and Stability Bonds form between atoms in a molecule so atoms can gain stability by achieving the desired octet of valence electrons. Energy is always released when bonds form. Energy is always required/needed to break bonds. The amount of energy required to break a specific covalent bond is called bond dissociation energy. Because breaking bonds requires the addition of energy, bond dissociation energy is always positive.