PRINCIPLES OF CHEMISTRY I CHEM 1211 CHAPTER 9 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university

CHAPTER 9 CHEMICAL BONDING

- The attractive force that holds atoms together - The result of interactions between electrons in the combining atoms Two types of chemical bonds - Covalent and Ionic (electrovalent) bonds CHEMICAL BOND

Covalent Bond - Formed through the sharing of one or more pairs of electrons between two atoms - Always involve two nonmetals - Electron sharing CHEMICAL BOND

Ionic Bond - Formed by attraction between two oppositely charged ions - Formed as a result of the transfer of electron(s) from atom(s) to another atom(s) - Often formed between metal and nonmetal ions through electrostatic attraction - Electron transfer CHEMICAL BOND

Two concepts - Valence Electrons - Octet Rule

VALENCE ELECTRONS - Not all electrons in a given atom participate in bonding - Only valence electrons are available for bonding (electrons in the outer most shell) - For representative and noble-gas elements these electrons are always found in the s or p subshells

VALENCE ELECTRONS - Using electron configuration to determine the number of valence electrons C: 1s 2 2s 2 2p 2 O: 1s 2 2s 2 2p 4 Na: 1s 2 2s 2 2p 6 3s 1 - Using electron-dot structure (Lewis symbol) to designate the number of valence electrons (place first 4 dots separately on four sides and pair up as needed) ∙C∙ :O∙ Na∙.....

VALENCE ELECTRONS Three important facts about valence electrons - Representative elements in the same group of the periodic table have the same number of valence electrons - The number of valence electrons for representative elements is the same as the group number (with A) in the periodic table - The maximum number of valence electrons for any given element is eight

OCTET RULE - Electrons arranged with 8 valence electrons are more stable than all others - The valence electron configuration of the noble gases are considered the most stable (all have 8 valence electrons; helium has 2) - All noble gases have the outermost s and p subshells completely filled

OCTET RULE - The noble gases are the most unreactive of all elements - Atoms of many elements tend to acquire the 8 valence electron configuration through chemical reactions - Atoms of elements tend to gain, lose, or share electrons to produce a noble-gas electron configuration - This results in the formation of compounds - This tendency is known as the OCTET RULE

IONIC BOND - Electron transfer - Metals donate electrons to form positive ions - Nonmetals accept electrons to form negative ions - The electrons lost by the metal are the same ones gained by the nonmetal

- The positive and negative ions attract one another to form ionic compounds - Ions combine in ratios to obtain charge neutrality (net charge = 0) - The symbol for positive ions is always written first IONIC BOND

Lewis Structures - Lewis structures involve compounds - Lewis symbols involve individual elements Na∙ + ∙Cl:[Na] + [:Cl:] - NaCl CaCl 2.. ∙Ca∙ +.. ∙Cl:.. [Ca] 2+ [:Cl:] -.. IONIC BOND

Energetics Removing an electron from Na(g) to form Na + (g) Na(g) → Na + (g) + e -  E = +496 kJ/mol Adding an electron to Cl(g) to form Cl - (g) Cl(g) + e - → Cl - (g)  E = -349 kJ/mol - Attraction between the unlike charges draws ions together causing energy to be released Heat of formation of ionic substances is quite exothermic Na(s) + 1/2Cl 2 (g) → NaCl(s)  H f o = kJ IONIC BOND

Energetics - Ionic compounds do not contain discrete molecules but ordered arrays of positive and negative ions (result of energy released) NaCl - Formula unit that indicates combining ratio - A given sodium ion has six immediate chloride ion neighbors - A given chloride ion has six immediate sodium ion neighbors IONIC BOND

Lattice Energy - The energy required to completely separate one mole of a solid ionic compound into its gaseous ions - Increases with increasing charge on the ions and decreasing distance between the radii of the ions (from electrostatic potential energy, E el ) NaCl(s) → Na + (g) + Cl - (g)  H lattice = +788 kJ/mol IONIC BOND

Lattice Energy - Highly endothermic indicating ions are strongly attracted to one another - Reason why ionic compounds are hard, brittle, and have high melting points Melting point of NaCl is 801 o C IONIC BOND

- Generally, transition metals do not form ions that have the noble-gas configuration - Transition metals first lose valence-shell s electrons and then as many d electrons as required to form ions - Transition metals can form different cations Fe: Fe 2+ and Fe 3+ Sn: Sn 2+ and Sn 4+ Pb: Pb 2+ and Pb 4+ TRANSITION METAL IONS

COVALENT BONDING - Involve electron sharing - Usually occurs between two nonmetals - The basic structural unit in covalent bonding is a molecule - Forms molecular compounds

HH ∙∙ : HH Two hydrogen atoms H + H Hydrogen molecule H H 1s electronsShared electron pair COVALENT BONDING

- Two neclei attract the same shared electrons to form a covalent bond - Orbitals containing the valence electrons overlap to create a common orbital - The electrons move throughout the common orbital - The electrons are shared by both nuclei COVALENT BONDING

- The valence electrons help each atom achieve a noble-gas configuration H∙H∙∙H∙HH : HHH :F∙:F∙ :F : F::F:FF:F: H : F:F:F:H H2H2 ∙F:∙F: ∙F:∙F:.. F2F2 HFH∙H∙.. bonding electrons nonbonding electrons LEWIS STRUCTURES

Bonding Electrons - The pairs of valence electrons involved in the covalent bond formation Nonbonding Electrons (Lone Pairs of Electrons) - The pairs of valence electrons not involved in electron sharing LEWIS STRUCTURES

H2OH2O H ∙ O : O H H :ORO H H - Oxygen (O) has six valence electrons - Gains two more through electron sharing with H - Achieves a noble-gas configuration.. :.. : LEWIS STRUCTURES

NH 3 H ∙ N N H H :ORN H H H ∙ H H : : :..... LEWIS STRUCTURES - Nitrogen (N) has five valence electrons - Gains three more through electron sharing with H - Achieves a noble-gas configuration

CH 4 ∙ C ∙ C H H :ORC H H H H H ∙ : HH.... H ∙ LEWIS STRUCTURES - Carbon (C) has four valence electrons - Gains four more through electron sharing with H - Achieves a noble-gas configuration

SINGLE COVALENT BOND - Two atoms share one pair of valence electrons - Represented by one line - Bond order is one Bond Order - Number of electron pairs that are shared between two atoms Bond Length - The minimum energy distance between the nuclei of two bonded atoms in a molecule

DOUBLE COVALENT BOND - Two atoms share two pairs of valence electrons - Represented by two lines - Approximately twice as strong as a single covalent bond between the same two atoms - Bond order is two

DOUBLE COVALENT BOND CO 2 - C has four valence electrons and needs four more - Each O atom has six valence electrons and needs two more :O::C::O:orOCO - Possible for elements that need two electrons to complete their octet..

TRIPLE COVALENT BOND - Two atoms share three pairs of valence electrons - Represented by three lines - Approximately thrice as strong as a single covalent bond between the same two atoms - Bond order is three - Bond length decreases with increasing bond order

TRIPLE COVALENT BOND N2N2 - Nitrogen has five valence electrons and needs three more to complete its octet - Each nitrogen must share three of its electrons with the other :N:::N:or:N:NN:N: - Possible for elements that need three or more electrons to complete their octet

COORDINATE COVALENT BOND - Both electrons come from only one of the two bonding atoms - Oxygen often forms coordinate covalent bonds : +XY : XY filled orbitalvacant orbitalshared electron pair H : O : Cl : coordinate covalent bond Chlorous acid (HClO 2 )Hypochlorous acid (HOCl).. H : O : Cl : O :..

ELECTRONEGATIVITY - The ability of an atom to attract to itself the electrons in a chemical bond - Electronegativity depends on atom size nuclear charge number of inner shell electrons - Increases from left to right across periods on the periodic table

- Increases from bottom to top within groups on the periodic table - Flourine is the most electronegative of all the elements - Nonmetals are more electronegative than metals - Indicative of the fact that nonmetals gain electrons and metals lose electrons ELECTRONEGATIVITY

LEWIS STRUCTURES - Calculate the total number of valence electrons in the molecule (use group numbers in the periodic table) HClO 2 H (group IA) has 1 valence electron Cl (group VIIA) has 7 valence electrons O (group VIA) has 6 valence electrons Total electron count = (6) = 20

- Determine the central atom The central atom - mostly appears only once (SO 3, SO 2, CH 4 ) - is usually any additional element other than H and O (HNO 3, H 2 SO 4 ) - is C in almost all carbon-containing compounds - is neither H nor F (can make only one covalent bond) - for O and H containing compounds O is bonded to the central atom and H to O HClO 2 (Cl is the central atom) LEWIS STRUCTURES

- Write the atoms in the order in which they are bonded together - Place a pair of electrons between each pair of atoms H : O : Cl : O HClO 2 LEWIS STRUCTURES

- Add nonbonding electron pairs to all atoms except the central atom - Each atom should have eight electrons - H needs only 2 electrons HClO 2 H : O : Cl : O : 16 out of the 20 electrons have been used up.. LEWIS STRUCTURES

HClO 2 H : O : Cl : O : 20 out of the 20 electrons have been used up.. LEWIS STRUCTURES - Place any remaining electrons on the central atom of the structure..

HClO 2 H : O : Cl : O :.. LEWIS STRUCTURES.. - This step is not needed in this case since Cl has completed its octet - If the central atom has less than eight move nonbonding electron pairs to form double or triple bonds

HClO 2 H : O : Cl : O :.. LEWIS STRUCTURES.. - Count the total number of electrons in the Lewis structure (must equal the initial number) 20 electrons equal to the intial 20

HCN H (group IA) has 1 valence electron C (group IVA) has 4 valence electrons N (group VA) has 5 valence electrons Total electron count = = 10 LEWIS STRUCTURES - Calculate the total number of valence electrons in the molecule (use group numbers in the periodic table)

HCN (C is the cental atom) LEWIS STRUCTURES - Determine the central atom The central atom - mostly appears only once (SO 3, SO 2, CH 4 ) - is usually any additional element other than H and O (HNO 3, H 2 SO 4 ) - is C in almost all carbon-containing compounds - is neither H nor F (can make only one covalent bond) - for O and H containing compounds O is bonded to the central atom and H to O

HCN H : C : N LEWIS STRUCTURES - Write the atoms in the order in which they are bonded together - Place a pair of electrons between each pair of atoms

HCN H : C : N : LEWIS STRUCTURES.. 10 out of the 10 electrons have been used up - Add nonbonding electron pairs to all atoms except the central atom - Each atom should have eght electrons - H needs only 2 electrons

HCN H : C : N : LEWIS STRUCTURES.. 10 out of the 10 electrons have been used up - Nothing left to be placed on the central atom - Place any remaining electrons on the central atom of the structure

HCN H : C : N : LEWIS STRUCTURES.. H : C ::: N : - If the central atom has less than eight move nonbonding electron pairs to form double or triple bonds

HCN H : C : N : LEWIS STRUCTURES.. H : C ::: N : - Count the total number of electrons in the Lewis structure (must equal the initial number) 10 electrons equal to the initial 10

POLYATOMIC IONS The total number of electrons for negative charges - increase the number of electrons by the magnitude of the charge SO 4 2- S (group VIA) has 6 valence electrons O (group VIA) has 6 valence electrons Charge of -2 Total number of electrons = 6 + 4(6) + 2 = 32

NH 4 + N (group VA) has 5 valence electrons H (group IA) has 1 valence electron Charge of +1 Total number of electrons = 5 + 4(1) - 1 = 8 The total number of electrons for positive charges - decrease the number of electrons by the magnitude of the charge POLYATOMIC IONS

Ionic compound containing polyatomic ion - The cation and anion are treated separately Na 2 SO 4 [Na] + S :O: :O: :O: O:O::O:O 2-.. POLYATOMIC IONS

BOND POLARITY Nonpolar Covalent Bond - Two atoms involved in electron sharing have equal or similar electronegativity - Typically less than Equal sharing of electrons F 2, H 2, O 2

BOND POLARITY Polar Covalent Bond - There exists unequal sharing of electrons - One atom is more electronegative than the other - One atom attracts electrons more strongly than the other - Electronegativity difference is between 0.4 and 1.5 HCl, CO

BOND POLARITY - Increasing bond polarity renders a bond more ionic - Ionic bonds have electronegativity difference greater than Most bonds are a mixture of pure ionic and pure covalent - No natural boundary between ionic and covalent bonding For electronegativity difference between 1.5 and ionic bond if metal and a nonmetal are involved - polar covalent bond if two nonmetals are involved

- Polar covalent bonds create partial positive and negative charges on the atoms involved - Delta (δ) is used to designate these partial charges δ+ for less electronegative atom δ- for more electronegative atom HCl: BOND POLARITY.. δ+δ+δ-δ-

HCl: - An arrow with a cross can also be used - The arrowhead is near the more electronegative end of the bond BOND POLARITY..

DIPOLE MOMENTS - A dipole establishes whenever two electrical charges of equal magnitude but opposite sign are separated by a distance - The quantitative measure of the magnitude of the dipole is known as the dipole moment µ = Qr µ = dipole moment Q = electrical charge (two equal and opposite charges Q+ and Q-) r = distance between the centers of Q+ and Q- Units: debyes (D) 1 D = 3.34 x coulomb-meters (C-m)

FORMAL CHARGE - Used to predict stability and connectivity To Calculate the Formal Charge - All nonbonding (unshared electrons) are assigned to the atom on which they are found - Half of the number of bonding electrons are assigned to each atom in the bond Formal Charge = Number of electrons assigned to the atom Number of valence electrons in the isolated atom - - Sum of formal charges equals the overall charge - Sum of formal charges in neutral atoms equals zero

FORMAL CHARGE Formal Charge = Number of electrons assigned to the atom Number of valence electrons in the isolated atom - [:C N:] - Six electrons in the triple bond C: 2 nonbonding electrons + 3 bonding electrons = 5 Number of valence electrons = 4 N: 2 nonbonding electrons + 3 bonding electrons = 5 Number of valence electrons = 5 Formal Charge of C = = -1 Formal Charge of N = = 0 [:C N:] -

RESONANCE STRUCTURES Ozone (O 3 ) SO 3 O OO OO O : : : : : : : : : : : : SSS OO O O O OO O O : :: :: ::: : : : : : : : : : : : :: : : :

EXCEPTIONS TO THE OCTET RULE Odd Number of Electrons (NO, ClO 2, NO 2 ) N O :: :. :: :. and - Called radicals and are very reactive For example The immune system uses NO to fight bacteria

Less Than an Octet of Valence Electrons (Electron Defficient) - Usually in compounds of boron, beryllium, and aluminum - BF 3 (only six valence electrons around boron) - BeH 2 - BeF 2 - BH 3 - AlH 3 EXCEPTIONS TO THE OCTET RULE

More Than an Octet of Valence Electrons (Expanded) - Occurs in elements of period 3 and beyond - No d orbitals in periods 1 and 2 to hold extra electrons - PCl 5 (10 valence electrons around phosphorus) - SF 6 (12 valence electrons around sulfur) - XeF 4 EXCEPTIONS TO THE OCTET RULE

STRENGTH OF COVALENT BONDS - Determined by the energy required to break the bonds - Bond enthalpy is the enthalpy change for breaking the bond in one mole of a gaseous substance - D(Cl — Cl) denotes bond enthalpy in Cl 2 - Bond enthalpies are always positive (energy is consumed) - To decompose CH 4 into C and 4H,  H = 1660 kJ There are 4 equivalent C — H bonds Average C — H bond enthalpy = D(C — H) = (1660/4) kJ/mol = 415 kJ/mol

BOND ENTHALPIES - Bond breaking is an endothermic process - Bond formation is an exothermic process - Bond enthalpy increases with increasing number of bonds - Bond length decreases with increasing number of bonds