PRINCIPLES OF CHEMISTRY I CHEM 1211 CHAPTER 9

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

CHAPTER 9 CHEMICAL BONDING

CHEMICAL BOND - 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: 1s22s22p2 O: 1s22s22p4 Na: 1s22s22p63s1 - 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

IONIC BOND - 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 .. .. ∙Cl: [Na]+ Na∙ + [:Cl:]- NaCl .. .. .. ∙Cl: .. .. [:Cl:]- .. CaCl2 ∙Ca∙ + [Ca]2+ .. .. [:Cl:]- .. ∙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) + 1/2Cl2(g) → NaCl(s) Hfo = -410.9 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, Eel) NaCl(s) → Na+(g) + Cl-(g) Hlattice = +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 oC

TRANSITION METAL IONS - Generally, transitionmetals 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: Fe2+ and Fe3+ Sn: Sn2+ and Sn4+ Pb: Pb2+ and Pb4+

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

COVALENT BONDING 1s electrons Shared electron pair : H H ∙ ∙ H H Two hydrogen atoms H + H Hydrogen molecule H H

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

LEWIS STRUCTURES - The valence electrons help each atom achieve a noble-gas configuration H∙ ∙H H : H H H H2 .. .. .. .. .. .. F2 :F∙ ∙F: :F : F: :F F: .. .. .. .. .. .. .. .. .. HF H∙ ∙F: H : F: H F: .. .. .. 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 H2O H H H ∙ . .. . H : : OR H O : O O : .. .. .. H ∙ - Oxygen (O) has six valence electrons - Gains two more through electron sharing with H - Achieves a noble-gas configuration

LEWIS STRUCTURES NH3 H ∙ H H . .. . H : : OR H N : N H ∙ N : .. . H H H ∙ - Nitrogen (N) has five valence electrons - Gains three more through electron sharing with H - Achieves a noble-gas configuration

LEWIS STRUCTURES CH4 H ∙ H H H ∙ . .. H : : H OR H C H C ∙ C ∙ .. . H ∙ H H H ∙ - 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 CO2 - C has four valence electrons and needs four more - Each O atom has six valence electrons and needs two more .. :O::C::O: or O C O .. - 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 N2 - 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 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 : : X + Y X Y filled orbital vacant orbital shared electron pair Hypochlorous acid (HOCl) Chlorous acid (HClO2) .. .. .. .. .. H : O : Cl : H : O : Cl : O : .. .. .. .. .. coordinate covalent bond

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

ELECTRONEGATIVITY - 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

LEWIS STRUCTURES - Calculate the total number of valence electrons in the molecule (use group numbers in the periodic table) HClO2 H (group IA) has 1 valence electron Cl (group VIIA) has 7 valence electrons O (group VIA) has 6 valence electrons Total electron count = 1 + 7 + 2(6) = 20

LEWIS STRUCTURES - Determine the central atom The central atom - mostly appears only once (SO3, SO2, CH4) - is usually any additional element other than H and O (HNO3, H2SO4) - 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 HClO2 (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 HClO2 H : O : Cl : O

LEWIS STRUCTURES - Add nonbonding electron pairs to all atoms except the central atom - Each atom should have eight electrons - H needs only 2 electrons HClO2 .. .. H : O : Cl : O : .. .. 16 out of the 20 electrons have been used up

LEWIS STRUCTURES - Place any remaining electrons on the central atom of the structure HClO2 .. .. .. H : O : Cl : O : .. .. .. 20 out of the 20 electrons have been used up

LEWIS STRUCTURES - If the central atom has less than eight move nonbonding electron pairs to form double or triple bonds HClO2 .. .. .. H : O : Cl : O : .. .. .. - This step is not needed in this case since Cl has completed its octet

LEWIS STRUCTURES - Count the total number of electrons in the Lewis structure (must equal the initial number) HClO2 .. .. .. H : O : Cl : O : .. .. .. 20 electrons equal to the intial 20

LEWIS STRUCTURES - Calculate the total number of valence electrons in the molecule (use group numbers in the periodic table) 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 = 1 + 4 + 5 = 10

LEWIS STRUCTURES - Determine the central atom The central atom - mostly appears only once (SO3, SO2, CH4) - is usually any additional element other than H and O (HNO3, H2SO4) - 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 (C is the cental 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 HCN H : C : N

LEWIS STRUCTURES - 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 : .. 10 out of the 10 electrons have been used up

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

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

LEWIS STRUCTURES - Count the total number of electrons in the Lewis structure (must equal the initial number) HCN .. H : C : N : .. H : C ::: N : 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 SO42- 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

PRINCIPLES OF CHEMISTRY I CHEM 1211 CHAPTER 9

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