Chapter 4: Forces Between Particles

1. Chapter 4:Forces Between Particles

2. Chapter 4 objectives  1. Draw correct Lewis structures for atoms of representative elements.  (Section 4.1; Exercise 4.2)   2. Use electronic configurations to determine the number of electrons gained or lost by atoms as they  achieve noble gas configurations.  (Section 4.2; Exercise 4.12)   3. Use the octet rule to correctly predict the ions formed during the formation of ionic compounds, and  write correct formulas for binary ionic compounds containing a representative metal and  representative nonmetal.  (Section 4.3; Exercises 4.20 and 4.22)   4. Correctly name binary ionic compounds.  (Section 4.4; Exercise 4.30)   5. Determine formula weights for ionic compounds.  (Section 4.5; Exercise 4.38)   6. Draw correct Lewis structures for covalent molecules.  (Section 4.6; Exercise 4.48)   7. Draw correct Lewis structures for polyatomic ions.  (Section 4.7; Exercise 4.50)   8. Use VSEPR theory to predict the shapes of molecules and polyatomic ions (Section 4.8;   Exercises 4.52 and 4.54)    9. Use electronegativities to classify covalent bonds of molecules, and determine whether covalent  molecules are polar or nonpolar.  (Section 4.9; Exercises 4.58 and 4.64)   10. Write correct formulas for ionic compounds containing representative metals and polyatomic ions,  and correctly name binary covalent compounds and compounds containing polyatomic ions.   (Section 4.10; Exercises 4.66, 4.70, and 4.72)   11. Relate melting and boiling points of pure substances to the strength and type of interparticle forces  present in the substances.  (Section 4.11; Exercises 4.78 and 4.80) 

3. Other ways of looking at electronic structure • Noble gas configurations (including those of transition metals) • Lewis structures • More on valence electrons

4. 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.

5. 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.

6. VALENCE ELECTRONS • 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.

7. LEWIS STRUCTUREEXAMPLE • 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:

8. Types of atom combinations and bonding Elements can combine in three basic ways. • Metals combining with metals • These can be called alloys if the elements are different. • Held together with metallic bonds – the core nuclei and electrons are held in a “sea” of valence electons. • Metals combining with non-metals. • Metal atoms lose electrons to form cations • Non-metal atoms gain electrons to form anions • The cations and anions are held together with electrostatic attractions – termed ionic bonds • Non-metals combining with non-metals • Atoms share electrons • Bonds are called covalent bonds

9. THE OCTET RULE • According to the octet rule, atoms will gain, lose or share 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 for representative elements. • SIMPLE ION • A simple ion is an atom that has acquired a net positive or negative charge by losing or gaining one or more electrons.

10. SIMPLE ION EXAMPLES • Magnesium, Mg, has two valence electrons which it loses to form a simple ion with a +2 electrical charge. The ion is written as Mg2+. • 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 O2-. • 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 -.

11. DETERMINING 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 Sr2+ ions and phosphorus, P, a group VA nonmetal forms P3- ions.

12. IONIC BONDFORMATION • 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 between simple ions when representative metal atoms lose valence electrons and 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.

13. ISOELECTRONIC • Isoelectronic is a term that literally means “same electronic,” used to describe atoms or ions that have identical electronic configurations. Examples include Na+, Mg2+, Ne and F- And K+, Ca2+, S2-, Cl-, Ar

14. IONIC COMPOUNDS • The substances that result 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. Cu2O CuO

15. BINARY IONIC COMPOUND FORMULAS • 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 will provide equal numbers of positive and negative electrical charges when combined together.

16. BINARY IONIC COMPOUND FORMULA EXAMPLES • Sodium and fluorine: • Sodium, a group IA metal, will form sodium ions with the symbol Na+. • Fluorine, a group VIIA nonmetal, will form fluoride 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.

17. BINARY IONIC COMPOUND FORMULA EXAMPLES (continued) • 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 sulfide ions with the symbol S2-. • The minimum number of ions required to give the same number of positive and negative charges is two Na+ ions and one S2- ion. • The two Na+ ions provide two positive charges and the one S2- ion provides two negative charges. • The resulting formula is Na2S.

18. BINARY IONIC COMPOUND FORMULA EXAMPLES (continued) • Aluminum and oxygen: • Aluminum is a group IIIA metal and will form ions with the symbol Al3+. • Oxygen is a group VIA nonmetal and will form ions with the symbol O2-. • The minimum number of ions required to give the same number of positive and negative charges is two Al3+ ions and three O2- ions. • The resulting formulas is Al2O3.

19. 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:

20. BINARY IONIC COMPOUNDNAME EXAMPLES • Name K2O: • name = metal name + nonmetal stem + -ide • name = potassium + ox- + -ide = potassium oxide • Name Mg3N2: • 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 AlBr3: • name = metal name + nonmetal stem + -ide • name = aluminum + brom- + -ide = aluminum bromide

21. NAMING BINARY IONIC COMPOUNDS(continued) • Some metal atoms, especially those of transition and inner-transition elements form more than one type of charged ion. (e.g. Iron forms both Fe2+ and Fe3+ 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 FeCl2 and FeCl3 contain iron ions with 2+ and 3+ charges, respectively. Their names are iron (II) chloride and iron (III) chloride. FeCl2 FeCl3

22. “Odd” element charges to know • Fe(III), Fe(II) • Cu(I), Cu(II) • Hg(I), Hg(II) • Sn(II), Sn(IV) • Pb(II), Pb(IV) • Ag+

23. IONIC COMPOUNDS CONTAINING 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.

24. COMMON POLYATOMIC IONS

25. EXAMPLES OF IONIC COMPOUNDS CONTAINING POLYATOMIC IONS • Compound containing K+ and ClO3- KClO3 • Compound containing Ca2+ and ClO3- Ca(ClO3)2 • Compound containing Ca2+ and PO43- Ca3(PO4)2

26. NAMING IONIC COMPOUNDS CONTAINING POLYATOMIC ANIONS • 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: • KClO3 is named potassium chlorate • Ca(ClO3)2 is named calcium chlorate • Ca3(PO4)2 is named calcium phosphate • CaHPO4 is named calcium hydrogen phosphate

27. 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.

28. IONIC COMPOUND FORMULAS & 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 crystal lattice. • 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 1023) of the simplest combining ratio of ions in the compounds.

29. COVALENT BONDS • A covalent bond is a type of bond 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, F2. • 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.

30. COVALENT BONDS(continued) • The sharing of electrons takes place when electron-containing orbitals of atoms overlap. This is shown below for the formation of the H2 molecule.

31. COVALENT BONDS(continued) • Electron sharing resulting in covalent bonding can occur between identical atoms or between different atoms. • Molecules such as F2, Cl2 Br2, I2 H2, O2 and N2 are formed when electron sharing occurs between identical atoms. • Molecules such as H2O, and CH4 are formed when electron sharing occurs between different atoms.

32. COVALENT BONDINGEXAMPLES

33. COVALENT BONDINGEXAMPLES (continued)

34. 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.

35. DRAWING LEWIS STRUCTURES FOR COVALENT MOLECULES (continued) • 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.

36. DRAWING LEWIS STRUCTURES FOR COVALENT MOLECULES EXAMPLE • Draw a Lewis structure for SO3. • 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 six valence electrons. The total number of electrons is 24 (six from the one S atom and 18 from the three O atoms).

37. DRAWING LEWIS STRUCTURES FOR COVALENT MOLECULES EXAMPLE (continued) • 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 six of the 24 available electrons. The remaining 18 are used to complete the octets of the atoms, beginning with the O atoms.

38. DRAWING LEWIS STRUCTURES FOR COVALENT MOLECULES EXAMPLE (continued) • 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.

39. 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.

40. 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.

41. NAMING BINARY COVALENT COMPOUNDS EXAMPLES • SO2: name = sulfur + di- + ox + -ide = sulfur dioxide • XeF6: name = xenon + hexa- + fluor + -ide = xenon hexafluoride • H2O: name = di- + hydrogen + mono- + ox + -ide = dihydrogen monoxide (also known as water) (Note, the final o of mono- was dropped for ease of pronunciation.)

42. COVALENT MOLECULE POLARITY • 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.

43. COVALENT MOLECULE POLARITY (continued) • 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+).

44. COVALENT MOLECULE POLARITY (continued) • 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.

45. THE POLARITY OF MOLECULES

46. THE SIGNIFICANCE OF POLARITY Molecules that are polar can attract/bond to other polar molecules. This attraction/bonding gives rise to higher melting/boiling points, is involved in things like dissolving and allows many kinds of biological interactions. Along with dispersion forces, polar interactions complete the difference kinds of interactions/forces that occur between atomic and molecular particles.

47. INTERPARTICLE FORCE SUMMARY • 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.

48. 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. H2O and CO) are held together by dipolar forces (also called dipole-dipole interactions), which are attractive forces that exist between the positive end of one polar molecule and the negative end of another.

49. TYPES OF MATERIALS (continued) • Some polar covalent molecules (e.g. H2O) experience hydrogen bonding, which is a special case of dipole-dipole interactions and the result of attractive dipolar forces between molecules in which hydrogen atoms are covalently bonded to very electronegative atoms (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. SiO2, diamond, carbon nanotubes).

50. TYPES OF MATERIALS (continued) • 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. O2 and CO2 – 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. The larger the molecules that greater the net attractions can be. E.g., Ar > Ne and C3H8 > CH4