Chemical Bonding - Session 04 - PDF

CHEMICAL BONDING (Part 01) The combining of Atoms & Molecules Mr. DARSHANA SAMPATH BSc (Hons) in ...

CHEMICAL BONDING (Part 01) The combining of Atoms & Molecules Mr. DARSHANA SAMPATH BSc (Hons) in  1 LEARNING OBJECTIVES  To explain why chemical bonds are formed.  To explain the formation of ionic and covalent bonds.  To write the chemical formula and name of an ionic compound.  To describe properties of ionic and covalent compounds.  To draw Lewis structures for covalent compounds and polyatomic ions.  2  Do you think that the most stable form for an element is that of a neutral atom?  Can you find the elemental form of sodium atoms in nature?  Can atoms of different elements join together to form entirely different substances?  3 What Is a Chemical Bond?  A chemical bond is a force of attraction between atoms or ions.  Bonds form when atoms share or transfer valence electrons. (Valence electrons are the electrons in the outer energy level of an atom that may be involved in chemical interactions.)  Valence electrons are the basis of all chemical bonds.  4 Why Chemical Bonds are formed? To give atoms a more stable arrangement of electrons.  The noble gas elements are the least reactive of all the elements. They almost never form any type of compound. Their electron configuration is the most stable of all of the elements, having their s and p sublevels filled.  The noble gases have an “octet”, meaning they have eight valence electrons.  The other elements are typically more stable if they have an octet, too.  Therefore other atoms will gain electrons, lose electrons, or share electrons in order to obtain an octet.  5 Octet Rule : The tendency of atoms to prefer to have eight electrons in the valence shell.  Elements in the second period complete the octet when forming chemical bonds thereby achieving a greater stability.  The valance shell of elements in the third period and subsequent periods consist of d sub energy level in addition to s and p sub energy levels. Therefore, when forming chemical bonds, there could be instances where the number of electrons in the valance shell may exceed eight. The presence of d orbitals in the valence shell permits 18 electrons. 6  7 1s1 1s2 2s2 2p2 1s2 2s2 2p3 1s2 2s2 2p4 Valence #e 1 4 5 6 Valency 1 4 3 2  8 Types of Chemical Bonds There are 3 primary types of chemical bonds which are formed by atoms or molecules to yield compounds.  Ionic Bonds  Covalent bonds  Metallic Bonds https://youtu.be/NgD9yHSJ29I? si=6pwaE1eQHIPOSRrK  9 1. Ionic bonding  Ionic bonding is a type of chemical bonding that involves the electrostatic attraction between oppositely charged ions.  The electrons are transferred from one atom to another resulting in the formation of positive and negative ions. The electrostatic attractions between the positive and negative ions forms the ionic bond. Metal atoms that lose electrons make positively charged ions - Cation Non-Metal atoms that gains electrons make negatively charged ions - Anion  10 Formation of Cations Cations are positively charged ions; they are formed when an atom (usually a metal) loses valence electrons. Atoms that lose electrons have more protons than electrons and so have an overall positive charge. (Cations are smaller in size than the original atom). 1s2 2s2 2p6 3s1 1s2 2s2 2p6 = [Ne] 1s2 2s2 2p6 3s2 1s2 2s2 2p6 = [Ne]  11 The charge of cations is given by the number of electrons lost, followed by a “+” symbol (ex. 2+). The formula of cations is given by the chemical symbol followed by the charge. (ex. Mg2+ ). Mg  Mg2+ + 2e  12 Predicting Monatomic Cation Charges The alkali metals in group 1 are always +1 when they form cations. The alkaline earth metals in group 2 are always +2 when they form cations. Aluminum and the elements in group 3 are always +3 when they form cations. The metallic elements in groups other than 1, 2, or 3 also lose electrons to form cations, but they do so in less easily predicted ways.  13 Monatomic Cation Names  The names of monatomic cations always start with the name of the metal, sometimes followed by a Roman numeral to indicate the charge of the ion.  The Roman numeral in each name represents the charge on the ion and allows us to distinguish between more than one possible charge. Cu+ is copper(I) Cu2+ is copper(II)  If the atoms of an element always have the same charge, the Roman numeral is unnecessary Na+ is named sodium ion (without the Roman numeral for the charge) Polyatomic Cation Names  There is only one common polyatomic ion. Its formula is NH4+, and its name is ammonium.  14 Formation of Anions Anions are negatively charged ions; they are formed when an atom (usually a nonmetal) gains valence electrons. Atoms that gain electrons have less protons than electrons and so have an overall negative charge (Anions are larger in size than the original atom). O : 1s2 2s2 2p4 1s2 2s2 2p6 = [Ne] Cl: 1s2 2s2 2p6 3s2 3p5 1s2 2s2 2p6 3s2 3p6 = [Ar]  15 The charge of anions is given by the number of electrons gained, followed by a “-” symbol (ex. 2-). The formula of anions is given by the chemical symbol followed by the charge (ex. S2-). S + 2e  S2-  16 Predicting Monatomic Anion Charges Group 15 : -3 negative charge Group 16 : -2 negative charge Group 17 : -1 negative charge  17 Monatomic Anion Names The monatomic anions are named by adding -ide to the root of the name of the nonmetal that forms the anion. H- : hydride ion N3-: nitride ion P3-: phosphide ion O2- : oxide ion S2- : sulfide ion F- : fluoride ion Cl- : chloride ion Br- : bromide ion Polyatomic Anion Names There is many polyatomic anions. The following anions are most common. OH- : hydroxide ion NO3- : nitrate ion CO32- : carbonate ion SO42- : sulfate ion PO43- : phosphate ion HCO3-: hydrogen carbonate ion HPO42-: hydrogen phosphate ion HS- : hydrogen sulfide ion  18 Ionic Compounds Compounds that contain ions are called ionic compounds and result when metals (cations) react with nonmetals (anions). The ions arrange themselves into organized patterns where each ion is surrounded by several ions of the opposite charge. The organized patterns of positive and negative ions are called lattice structures. The electrostatic attraction between the oppositely charge ions is quite strong (the strongest of all types of chemical bonds)  19 NaCl : Sodium Chloride The image shows the solid lattice structure of sodium chloride. Each sodium ion is touching six chloride ions; the four surrounding ones and one above and one below. Each chloride ion is touching six sodium ions in the same way.  20 Properties of ionic compounds  Generally hard, but brittle.  Very high melting and boiling points.  Ionic substances generally dissolve readily in water.  Solutions of ionic compounds and melted ionic compounds conduct electricity, but solid materials do not.  21 How to write the formula for an ionic compound 1. Write the symbol and charge of the cation (metal) first and the anion (nonmetal) second. 2. Transpose only the number of the positive charge to become the subscript of the anion and the number only of the negative charge to become the subscript of the cation. 3. Reduce to the lowest ratio. 4. Write the final formula. Leave out all subscripts that are 1.  22 Fill the following table by writing the chemical formula Na+ Ca2+ Fe3+ SO42- Na2SO4 Sodium sulphate NO3- CO32- Br- O2- OH-  23 Write down the chemical formula of following compounds. 1) Magnesium sulfide 10) Calcium Oxide 2) Lead(II) Nitrate 11) Copper(I) Bromide 3) Potassium Oxide 12) Aluminum sulfide 4) Potassium Carbonate 13) Hydrogen Carbonate 5) Aluminum Bromide 14) Iron(III) Chloride 6) Iron (III) nitrate 7) Iron(II) Chloride 8) Copper(II) Nitrate 9) Magnesium oxide  24 2. Covalent Bonds  Covalent bonds are formed when a pair of electrons is shared between two atoms of the same element or different elements.  The sharing pair of electrons contribute one electron from each atom to form the electron pair.  Consequently, stable electron configurations are often achieved by both atoms in respect to the total number of electrons in the valence shells.  25 When bonds are forming energy is released.  26 Types of Covalent bonds Although we defined covalent bonding as electron sharing, the electrons in a covalent bond are not always shared equally by the two bonded atoms. Unless the bond connects two atoms of the same element, there will always be one atom that attracts the electrons in the bond more strongly than the other atom does. (electronegativity difference)  27 Polar Covalent Bond In a polar covalent bond, two atoms share a pair of electrons unequally because of differences in their electro-negativities. The unequal distribution of electrons causes the bonded atoms to acquire partial positive and negative charges.  28 Non- Polar Covalent Bond  When atoms share an equal number of electrons, a non-polar covalent bond is formed.  There is no difference in the electronegativity of the two atoms.  It occurs when two atoms with the same electron affinities come together, as in diatomic elements.  29 How to determine whether polar bond or non polar bond?  The greater the difference in electronegativities, the greater the imbalance of electron sharing in the bond.  The general rule is if the difference in electronegativities is  less than about 0.4, the bond is considered nonpolar  greater than 0.4, the bond is considered polar.  If the difference in electronegativities is large enough (generally greater than about 1.8), the resulting compound is considered ionic rather than covalent.  An electronegativity difference of zero, of course, indicates a nonpolar covalent bond.  30 Covalent compounds Compounds that are composed of only non-metals or semi-metals with non- metals will display covalent bonding. There are several types of structures for covalent substances; 1. Individual molecules ( HCl / SO2 / CO2) 2. Molecular structures ( ethanol / Solid I2) Solid I2 Ethanol  31 Covalent compounds 3. Macromolecular structures (polyethylene / proteins)  32 4. Giant covalent structures (graphite / diamond)  33 Properties of Covalent compounds  The boiling/melting points of covalent compounds are low. (Giant covalent substances have very high melting points and boiling points)  They are soft in nature and relatively flexible.  These compounds do not possess electrical conductivity.  Nonpolar covalent compounds dissolve poorly in water. Some Compounds Have Both Covalent and Ionic Bonds  Nonmetal atoms in polyatomic ions are joined by covalent bonds, but the ion as a whole participates in ionic bonding.  34 Lewis Structures The Lewis structure of a molecule show how the valence electrons are arranged among the atoms of the molecule. Carbon Chlorine Oxygen 1s2 2s2 2p2 1s2 2s2 2p6 3s2 3p 5 1s2 2s2 2p4 Nitrogen Sodium Magnesium 1s2 2s2 2p3 1s2 2s2 2p6 3s1 1s2 2s2 2p6 3s2  35 How to draw Lewis structures Step 1: Determine the total number of valence electrons. H2O : Oxygen atom contributes 6 electrons each Hydrogen atom contributes 1 electron ( 1 x2 =2) Total valence electron count (6 + 2 = 8)  If it is a negatively charged ion, then negative charges should be counted as well. OH- : Oxygen atom contributes 6 electrons Hydrogen atom contributes 1 electron Negative charge of the ion 1 Total valence electron count (6 + 1 + 1 = 8)  If the ion is positively charged, then an equivalent number is deducted from the total valance electron count. NH4+ : Nitrogen atom contributes 5 electrons 4 Hydrogen atom contributes 4 electrons Positive charge of the ion 1 Total valence electron count (5 + 4 - 1 = 8)  36 Step 2: Write the skeleton structure of the molecule. The element with the lower electronegativity is generally the central atom. (except H) H O H Step 3: Use two valence electrons to form each bond in the skeleton structure. A bond is denoted by a pair of dots between the central atom and a surrounding atom. H : O : H Step 4: Try to satisfy the octets of the atoms by distributing the remaining valence electrons as nonbonding electrons. The remaining electrons are distributed, starting from the most electronegative atom, to complete the octet. Total valence # e = 8 Bonding # e = 4 Remaining e = 4  37 Step 05 : If electron pairs are remaining after distributing electron pairs on the surrounding atoms (satisfying the octet rule), then left over pairs of electrons are marked on the central atom. Step 06 : After distributing all the electron pairs, the number of electrons on each atom should be compared with the number of electrons in the non-bonded state of the atom (free atom) to assign the formal charge and then check completion of the octet. In the case of a bond, one electron is counted for each atom and if lone pairs are present, both electrons are counted to the particular atom. Step 07 : Electron distribution shall be rearranged in order to minimize the formal charge on atoms and completion of octet by converting lone pair of electrons to bonding pairs of electrons.  38 Let’s draw the Lewis diagram of Carbon dioxide (CO2)  First calculate the total valance electrons. C:4 O:6x2 Total = (4+12) = 16 Let’s draw the Lewis diagram of Sulphate ion (SO4 2-) S :6 O :6x4 (-) : 2 Total = (6+24+2) = 32  39 Draw the Lewis structures of the following compounds. H2S : Hydrogen Sulphide SO3 : Sulphur trioxide HCN : Hydrogen cyanide CH2Br2 : Dibromomethane CCl4 : Carbon tetrachloride NH4+ : Ammonium ion CO32- : Carbonate ion  40 Now You Should be Able…….  To explain why chemical bonds are formed.  To explain the formation of ionic and covalent bonds.  To write the chemical formula and name of an ionic compound.  To describe properties of ionic and covalent compounds.  To draw Lewis structures for covalent compounds and polyatomic ions. CHEMICAL BONDING To be Continued………. THANK YOU.