Bonding and structure: Forces Between Particles

Bonding and structure: Forces Between Particles NOBLE GAS CONFIGURATIONS An electronic configuration that is characterized by two electrons in the valence shell of helium and eight electrons in the valence shell of all other group VIIIA noble gases. 1 2 LEWIS STRUCTURES A representation of an atom or ion in which the elemental symbol represents the atomic nucleus and all but the valence-shell electrons. The valence electrons are represented by dots arranged around the elemental symbol. DETERMINATION OF THE NUMBER OF VALENCE ELECTRONS IN AN ATOM Write an electronic configuration for the atom and identify the valence electrons as those having the largest n value in the configuration. A simpler alternative for representative elements is to refer to the periodic table and note the group to which the element belongs. The number of valence electrons is the same as the Roman numeral group number. Examples: Calcium, Ca, is in group IIA. The number of valence electrons is 2. Phosphorus, P, is in group VA. The number of valence electrons is 5. 3 4 1

EXAMPLE OF A LEWIS STRUCTURE Potassium, K, is in group IA and so it has one valence electron. The Lewis structure is K. Aluminum, Al, is in group IIIA and so it has three valence electrons. The Lewis structure is Al :.! THE OCTET RULE According to the octet rule, atoms will gain or lose sufficient electrons to achieve an outer electron arrangement identical to that of a noble gas. This arrangement usually consists of eight electrons in the valence shell. SIMPLE ION An atom that has acquired a net positive or negative charge by losing or gaining one or more electrons. 5 6 EXAMPLES OF SIMPLE IONS Magnesium, Mg, has two valence electrons which it loses to form a simple ion with a +2 electrical charge. The ion is written as Mg 2+. Oxygen, O, has six valence electrons. It tends to gain two electrons to form a simple ion with a -2 electrical charge. The ion is written as O 2-. Bromine, Br, has seven valence electrons. It tends to gain one electron to form a simple ion with a -1 electrical charge. The ion is written as Br -. A SIMPLE WAY TO OBTAIN IONIC CHARGES FOR REPRESENTATIVE ELEMENTS Representative metals will form ions having the same positive charge as the number (Roman numeral) of the group to which they belong. Representative nonmetals will form ions with a negative charge equal to 8 minus the number (Roman numeral) of the group to which they belong. For example, strontium, Sr, a group IIA metal forms Sr 2+ ions and phosphorus, P, a group VA nonmetal forms P 3- ions. 7 8 2

IONIC BOND FORMATION Ions with positive charges are attracted to ions with negative charges. The attractive force between such ions holds them together and is called an ionic bond. Ionic bonds form when representative metal atoms lose valence electrons. The electrons are gained by representative nonmetal atoms. Both atoms are changed into ions with noble gas configurations. The resulting ions are then attracted to each other. Learning Check 1. The number of valence electrons in aluminum is a) 1e - b) 2e - c) 3e - 2. The change in electrons for octet requires a a) loss of 3e - b) gain of 3e - c) gain of 5e - 3. The ionic charge of aluminum is a) 3 - b) 5 - c) 3 + 4. The symbol for the aluminum ion is a) Al 3+ b) Al 3- c) Al + https://www.youtube.com/watch?v=ftw7a5ccubs 9 10 10 Solution Learning Check 1. The number of valence electrons in aluminum is c) 3e - A. Number of valence electrons in sulfur is 1) 4e - 2) 6e - 3) 8e - 2. The change in electrons for octet requires a a) loss of 3e - 3. The ionic charge of aluminum is c) 3 + 4. The symbol for the aluminum ion is a) Al 3+ B. Change in electrons for octet requires a 1) loss of 2e - 2) gain of 2e - 3) gain of 4e - C. Ionic charge of sulfur is 1) 2+ 2) 2-3) 4-11 11 12 12 3

Solution A. Number of valence electrons in sulfur is 2) 6e - ISOELECTRONIC A term that literally means same electronic, used to describe atoms or ions that have identical electronic configurations B. Change in electrons for octet requires a 2) gain of 2e - C. Ionic charge of sulfur is 2) 2-13 13 14 IONIC COMPOUNDS The substances formed when ionic bonds form between positive and negative ions are called ionic compounds. When ionic compounds are formed by the reaction of only two elements the resulting ionic compound is called a binary ionic compound. FORMULAS FOR BINARY IONIC COMPOUNDS Binary ionic compounds typically form when a metal and a nonmetal react. The metal tends to lose one or more electrons and forms a positive ion. The nonmetal tends to gain one or more electrons and forms a negative ion. The symbol for the metal is given first in the formula. The formula for a binary ionic compound represents the minimum number of each ion that when combined together will provide equal numbers of positive and negative electrical charges. Cu 2 O CuO 15 NaCl 16 4

EXAMPLES OF FORMULAS FOR BINARY IONIC COMPOUNDS Sodium and fluorine: Sodium, a group IA metal, will form sodium ions with the symbol Na +. Fluorine, a group VIIA nonmetal, will form ions with the symbol F -. The minimum number of ions needed to give the same number of positive and negative charges is one of each. The one Na + provides one positive charge and the one F - provides one negative charge. The correct formula that results is NaF. Sodium and sulfur: Sodium is a group IA metal and will form sodium ions with the symbol Na +. Sulfur is a group VIA nonmetal and will form ions with the symbol S 2-. The minimum number of ions required to give the same number of positive and negative charges is two Na + ions and one S 2- ion. The two Na + ions provide two positive charges and the one S 2- ion provides two negative charges. The resulting formula is Na 2 S. 17 18 Aluminum and oxygen: Aluminum is a group IIIA metal and will form ions with the symbol Al 3+. Oxygen is a group VIA nonmetal and will form ions with the symbol O 2-. The minimum number of ions required to give the same number of positive and negative charges is two Al 3+ ions and three O 2- ions. The resulting formulas is Al 2 O 3. Select the correct formula for each of the following ionic compounds A. Na + and S 2-1) NaS 2) Na 2 S 3) NaS 2 B. Al 3+ and Cl - Learning Check 1) AlCl 3 2) AlCl 3) Al 3 Cl C. Mg 2+ and N 3-1) MgN 2) Mg 2 N 3 3) Mg 3 N 2 19 20 20 5

Solution A. Na + and S 2-2) Na 2 S check: 2Na + + S 2- = 2(1+) + 1(2-) = 0 NAMING BINARY IONIC COMPOUNDS Binary ionic compounds are named using the following pattern: name = metal name + stem of nonmetal name + -ide The stem names and ionic symbols for some common nonmetals are given in the following table: B. Al 3+ and Cl - 1) AlCl 3 check: Al 3+ + 3Cl - = (3+) + 3(1-) = 0 C. Mg 2+ and N 3-3) Mg 3 N 2 check: 3Mg 2+ +2N 3- = 3(2+) + 2(3-) = 0 21 21 22 EXAMPLES OF BINARY IONIC COMPOUND NAMES Name K 2 O: name = metal name + nonmetal stem + -ide name = potassium + ox- + -ide = potassium oxide Name Mg 3 N 2 : name = metal name + nonmetal stem + -ide name = magnesium + nitr- + -ide = magnesium nitride Name BeS: name = metal name + nonmetal stem + -ide name = beryllium + sulf- + -ide + beryllium sulfide Name AlBr 3 : name = metal name + nonmetal stem + -ide name = aluminum + brom- + -ide name = aluminum bromide NAMING BINARY IONIC COMPOUNDS IN WHICH METALS FORM IONS WITH MORE THAN ONE CHARGE Some metal atoms, especially those of transition and inner-transition elements form more than one type of charged ion. (e.g. Cobalt forms both Co 2+ and Co 3+ ions.) The binary compounds containing such ions are named following the pattern given earlier with one addition, the number of positive charges on the metal ion is indicated by a Roman numeral in parentheses following the metal name. For example, the compounds CoO and Co 2 O 3 contain cobalt ions with 2+ and 3+ charges respectively. Their names are cobalt (II) oxide and cobalt (III) oxide. 23 24 6

IONIC COMPOUND STRUCTURE The stable form of an ionic compound is not a molecule, but a crystal in which many ions of opposite charge occupy lattice sites in a rigid three-dimensional arrangement called a crystal lattice. IONIC COMPOUND FORMULAS AND WEIGHTS Formulas for ionic compounds represent only the simplest combining ratio of the ions in the compounds, not the precise numbers of atoms of each element found in a molecule. Formula weight is the sum of the atomic weights of the atoms shown in the formula of an ionic compound. This is similar to molecular weight. One mole of an ionic compound contains Avogadro s number (6.022 x 10 23 ) of the simplest combining ratio of ions in the compounds. 25 26 COVALENT BONDING Covalent bonding is a type of bonding in which the octet rule is satisfied when atoms share valence electrons. The shared electrons are counted in the octet of each atom that shares them as illustrated below for fluorine gas, F 2. The sharing of electrons takes place when electroncontaining orbitals of atoms overlap. This is shown below for the formation of the H 2 molecule. The atoms sharing one or more pairs of electrons are each attracted to the shared electrons, and thus, are attracted to each other. The attraction to each other is called a covalent bond. The covalent bond may be represented by the shared pair or by a single line between the bonded atoms. 27 28 7

Electron sharing resulting in covalent bonding can occur between identical atoms or between different atoms. Molecules such as Cl 2, O 2 and N 2 are formed when electron sharing occurs between identical atoms. Molecules such as H 2 O, and CH 4 are formed when electron sharing occurs between different atoms. EXAMPLES OF COVALENT BONDING H 2 O CH 4 29 30 DRAWING LEWIS STRUCTURES FOR COVALENT MOLECULES Step 1: Use the molecular formula to determine how many atoms of each type are in the molecule. Step 2: Use the provided connecting pattern of atoms to draw an initial molecular structure with the atoms properly arranged. Step 3: Determine the total number of valence-shell electrons contained in the atoms of the molecule. Step 4: Put one pair of electrons between each bonded pair of atoms in the initial structure drawn in Step 2. Subtract the number of electrons used in this step from the total number determined in Step 3. Use the remaining electrons to complete the octets of all other atoms in the structure, beginning with the atoms that are present in greatest number in the molecule. Remember, hydrogen atoms only require one pair of electrons to achieve the electronic configuration of helium. Step 5: If all octets cannot be satisfied with the available electrons, move pairs that are not located between atoms to positions between atoms to complete octets. This will create double or triple bonds between some atoms. 31 32 8

EXAMPLE OF DRAWING A LEWIS STRUCTURE FOR A COVALENT MOLECULE Draw a Lewis structure for SO 3. Step 1: The formula indicates one S and three O atoms are in the molecule. Step 2: The connecting pattern is that each O is bonded only to the S. Thus, the following arrangement is drawn: O S O O Step 3: Sulfur and oxygen are both in group VIA, and so each atom has 6 valence electrons. The total number of electrons is 24 (6 from the one S atom and 18 from the three O atoms). Step 4: One pair of electrons is put between each O atom and the S atom of the arrangement drawn in step 2. This required 6 of the 24 available electrons. The remaining 18 are used to complete the octets of the atoms, beginning with the O atoms. 33 34 Step 5: After step 4, it is seen that the octet of S is not completed, even though all available electrons have been used. One nonbonding pair from any of the three O atoms will be moved to a location between the O and the S atoms. This pair will continue to count toward the octet of the O, but will also now count toward the octet of the S. The resulting correct Lewis structure contains one double bond (two shared pairs) between the S and one of the O atoms. POLYATOMIC IONS Polyatomic ions are covalently-bonded groups of atoms that carry a net electrical charge. Most common polyatomic ions are negatively charged. Lewis structures can be drawn for polyatomic ions using the same steps that were shown earlier for covalent molecules with one change. In Step 3, one electron is added to the total for each negative charge found on the polyatomic ion and one electron is subtracted from the total for each positive charge found on the polyatomic ion. All other steps are used unchanged. 35 36 9

Some Names of Polyatomic Ions The names of common polyatomic anions End in ate. NO 3 nitrate PO 3 4 phosphate With one oxygen less end in ite. NO 2 nitrite PO 3 3 phosphite With hydrogen attached use prefix hydrogen (or bi). HCO 3 hydrogen carbonate (bicarbonate) HSO 3 hydrogen sulfite (bisulfite) Names and Formulas of Common Polyatomic Ions 37 37 Copyright 2007 by Pearson Education, Inc. Publishing as Benjamin Cummings 38 38 Names and Formulas of Common Polyatomic Ions Naming Compounds with Polyatomic Ions The positive ion is named first followed by the name of the polyatomic ion. NaNO 3 sodium nitrate K 2 SO 4 potassium sulfate Fe(HCO 3 ) 3 iron(iii) bicarbonate or iron(iii) hydrogen carbonate (NH 4 ) 3 PO 3 ammonium phosphite Copyright 2007 by Pearson Education, Inc. Publishing as Benjamin Cummings 39 39 40 40 10

TABLE 4.8 Some Compounds with Polyatomic Ions Copyright 2007 by Pearson Education, Inc. Publishing as Benjamin Cummings 41 41 SHAPES OF MOLECULES AND POLYATOMIC IONS Most molecules and polyatomic ions are not flat twodimensional objects. Most have distinct three-dimensional shapes. The shapes of molecules or polyatomic ions can be predicted using a theory called the valence-shell electron repulsion theory, or VSEPR theory (sometimes pronounced "vesper" theory). According to the VSEPR theory, electron pairs in the valence shell of an atom will repel each other and get as far away from each other as possible. When the theory is applied to the valence-shell electron pairs of the central atom in a molecule or ion, the shape of the molecule or ion can be predicted. A central atom is an atom that is bonded to other surrounding atoms such as the S atom in the SO 3 molecule whose Lewis structure was drawn in the earlier example. 42 When the VSEPR theory is used, two rules are followed: Rule 1. All valence-shell electron pairs around the central atom are considered to behave the same regardless of whether they are bonding or nonbonding pairs. Rule 2. Double or triple bonds between surrounding atoms and the central atom are treated like a single pair of electrons when shapes are predicted. ARRANGEMENTS OF ELECTRON PAIRS The VSEPR theory can be applied to molecules or ions that have up to six pairs of electrons around the central atom. This discussion will be limited to molecules with no more than four pairs. 43 44 11

According to the VSEPR theory, the arrangement of electron pairs around the central atom (represented by E) depends on the number of electron pairs. Two pairs locate opposite each other. Three pairs arrange themselves in a flat triangle around the central atom. Four pairs become located at the four corners of a pyramidlike shape called a tetrahedron. EXAMPLES OF USING THE VSEPR THEORY Draw a Lewis structure and determine the shape of the molecule for CO 2. Solution: The Lewis structure drawn according to the rules given earlier is O=C=O. The central C atom is surrounded by two double bonds. Each double bond counts as a single electron pair, thus the central atom has two pairs of electrons around it. They will take up positions on opposite sides of the C. The O, C, and O atoms are arranged in a line and the molecule is linear. 45 46 Draw a Lewis structure and determine the shape of the molecule for NH 3. Solution: The Lewis structure drawn according to the rules given earlier is: The shape of the molecule is determined only by the positions of the atoms, not by the positions of the unshared electron pair. Thus, the NH 3 molecule has the shape of a pyramid with a triangular base. The N atom is at the peak of the pyramid and an H atom is at each corner of the base as shown below: The central atom is N. It has four electron pairs surrounding it. The four pairs will be located at the corners of a tetrahedron, with the N in the middle: 47 48 12

POLARITY OF COVALENT MOLECULES The shared electrons of covalent bonds are not always shared equally by the bonded atoms. Electrons of a covalent bond are attracted toward atoms of highest electronegativity. Unequal sharing of the bonding electrons of a covalent bond cause the bond to become a polar covalent bond. For atoms bonded by a polar covalent bond, the more electronegative atom acquires a partial negative charge (δ-) and the less electronegative atom acquires a partial positive charge (δ+). When the resulting partial charges are distributed symmetrically in a molecule, the molecule is nonpolar. When the partial charges are distributed nonsymmetrically, the molecule is polar. 49 50 THE POLARITY OF MOLECULES NAMING BINARY COVALENT COMPOUNDS The pattern used to name binary covalent compounds is similar to that used to name binary ionic compounds: name = name of least electronegative element + stem of more electronegative element + -ide In addition to the pattern, the number of each type of atom in the molecule is indicated by means of the following Greek prefixes: Note: The prefix mono is not used when it appears at the beginning of the name. 51 52 13

EXAMPLES OF NAMING BINARY COVALENT COMPOUNDS SO 2 : name = sulfur + di- + ox + -ide = sulfur dioxide XeF 6 : name = xenon + hexa- + fluor + -ide = xenon hexafluoride H 2 O: name = di- + hydrogen + mono- + ox + -ide = dihydrogen monoxide (also known as water) (Note, the final o of mono- was dropped for ease of pronunciation.) WRITING FORMULAS OF IONIC COMPOUNDS THAT CONTAIN POLYATOMIC IONS The rules for writing formulas for ionic compounds containing polyatomic ions are essentially the same as those used for writing formulas for binary ionic compounds. The symbol for the metal is written first, followed by the formula for the negative polyatomic ion. Equal numbers of positive and negative charges must be represented by the formula. When more than one polyatomic ion is required in the formula, parentheses are placed around the polyatomic ion before the subscript is inserted. ( ) ( ) Na PO Mg PO NH PO 3 4 3 4 2 4 3 4 53 54 NAMING IONIC COMPOUNDS THAT CONTAIN A POLYATOMIC ANION The names of ionic compounds that contain a polyatomic anion are obtained using the following pattern: name = name of metal + name of polyatomic anion Examples: KClO 3 is named potassium chlorate Ca(ClO 3 ) 2 is named calcium chlorate Ca 3 (PO 4 ) 2 is named calcium phosphate CaHPO 4 is named calcium hydrogen phosphate EXAMPLES OF FORMULAS FOR IONIC COMPOUNDS CONTAINING POLYATOMIC IONS Compound containing K + and ClO 3- : KClO 3 Compound containing Ca 2+ and ClO 3- : Ca(ClO 3 ) 2 Compound containing Ca 2+ and PO 4 3- : Ca 3 (PO 4 ) 2 55 56 14

A SUMMARY OF INTERPARTICLE FORCES Ionic and covalent bonds represent two of the forces that occur between atomic-sized particles and hold the particles together to form the matter familiar to us. Other forces also exist that hold the particles of some types of matter together. These include: metallic bonding, dipolar forces, hydrogen bonding, dispersion forces. TYPES OF MATERIALS Ionic compounds (e.g. NaCl) are held together by ionic bonds, which are attractive forces that hold together ions of opposite charge. Polar covalent compounds (e.g. H 2 O and CO) are held together by dipolar forces, which are attractive forces that exist between the positive end of one polar molecule and the negative end of another. 57 58 Some polar covalent molecules (e.g. H 2 O) experience hydrogen bonding, which is the result of attractive dipolar forces between molecules in which hydrogen atoms are covalently bonded to very electronegative elements (O, N, or F). Network solids are solids in which the lattice sites are occupied by atoms that are covalently bonded to each other (e.g. SiO 2 and diamond pictured below). Metals (e.g. Cu) are held together by metallic bonds, which originate from the attraction between positively charged atomic kernels that occupy lattice sites and mobile electrons that move freely through the lattice. Nonpolar covalent molecules (e.g. O 2 and CO 2 shown below) are only held together by dispersion forces, which are very weak attractive forces acting between the particles of all matter that result from momentary nonsymmetric electron distributions in molecules or atoms. 59 60 15

RELATIVE STRENGTHS OF INTERPARTICLE FORCES 61 62 16