Chapter 10 Chemical Bonding Tro 2 nd ed

Chapter 10 Chemical Bonding Tro, 2 nd ed.

Nonmetals foundtotothe theleft right metalloids. Metals arearefound of of thethe metalloids

LEWIS STRUCTURES OF ATOMS Metals form cations and nonmetals form anions to attain a stable valence electron structure. These rearrangements occur by losing, gaining or sharing electrons. The Lewis structure of an atom is a representation that shows the valence electrons for that atom. Valence electrons: the electrons that occupy the outermost energy level of an atom. Valence electrons are responsible for the electron activity that occurs to form chemical bonds.

The Lewis structure of an atom uses dots to show the valence electrons of atoms. B 2 1 2 s 2 p The number of dots equals the number of s and p electrons in the atom’s outermost shell.

Lewis Structures of the first 20 elements Notice that C’s e- config is 2 s 22 p 2.

CHEMICAL BONDING Atoms will do one of three “things” to get to a noble gas electron configuration: 1. Take electrons from another atom 2. Give electrons to another atom 3. Share electrons with atom(s) In choices 1 & 2: cause ions to form, then ionic bonds In choice 3: sharing electrons results in covalent bonds With the exception of hydrogen & helium, this structure consists of eight electrons in the outermost energy level (The Octet Rule)

The Ionic Bond Cations (+) have given up e-s and anions (-) have gained e-s, and now have opposite electrical charges Results in strong electrostatic force of attraction All cations and anions exist in crystal lattices, defined geometric structures with repeating 3 -D pattern

The 3 s electron of sodium transfers to the 3 p is orbital of chlorine. The force holding Na+ and Cl- together an ionic bond. A sodium ion (Na+) and a chloride ion (Cl-) are formed. Lewis representation of sodium chloride formation.

2+ and two Cl- together are 2+ -) are The forces holding Mg Two 3 s electrons of magnesium the A magnesium ion (Mg ) and twotransfer chloride to ions (Cl 3 p ionic bonds. orbitals of two chlorine atoms. formed.

In Na. Cl theiscrystal made up each of sodium cubic crystals. ion is surrounded by six chloride ions.

In the crystal each chloride ion is surrounded by six sodium ions.

The ratio of Na+ to Cl- is 1: 1 There is no molecule of Na. Cl

The Ionic Bond Using your Periodic Table, determine the cation and anion each atom is likely to form, then write the Lewis structures and made the compounds. Finish by writing the compound’s chemical formula. Practice: Al & F, Mg & O, Na & N, Al & O

A sodium ion is smaller than a sodium atom because: (1) the sodium atom has lost its outermost electron. (2) the 10 remaining electrons are now attracted by 11 protons and are drawn closer to the nucleus.

A chloride ion is larger than a chloride atom because: (1) the chlorine atom has gained an electron and now has 18 electrons and 17 protons. (2) The nuclear attraction on each electron has decreased, allowing the chlorine to expand.

Transition Metals form cations – a little different Transition metals lose their “s” electrons first, because they are in the highest principle energy level, then they lose their “d” electrons. Zn 2+ + 2 e- Cu Cu+ + 1 e[Ar]4 s 23 d 10 [Ar]4 s 13 d 10 [Ar]3 d 10

COVALENT BONDING A covalent bond consists of a pair of electrons shared between two atoms. In the millions of chemical compounds that exist, the covalent bond is the predominant chemical bond. Substances which covalently bond exist as molecules.

Carbon dioxide bonds covalently. It exists as individually bonded covalent molecules containing one carbon and two oxygen atoms.

The term molecule is not used when referring to ionic substances. Instead they are called Formula Units. Sodium chloride bonds ionically. It consists of a large aggregate of positive and negative ions. No molecules of Na. Cl exist.

COVALENT BONDING Nonmetal Atoms have deficiency of electrons in outermost shell and want to gain electrons to get full shell Since two nonmetal atoms both want more electrons, they will share electrons to get full shell H has 1 e- and wants 2 Cl has 7 e-s and wants 8 Both satisfied if they share a pair of electrons between them Each contributes 1 e- to the pair and each gets to share the 2 e-s in the pair H has 2 e-s and Cl has 8 e-s and they are “HAPPY”

COVALENT BONDING IMPORTANT: in giving e-s to be shared, atom actually gains e-s The number of e-s an atom contributes to be shared is equal to the number of e-s it needs to have an octet! (or a full shell) A pair of shared e-s is called a covalent bond 1 pair of e-s between two atoms = single bond 2 pairs of e-s betwn two atoms = double bond 3 pairs “ “ “ = triple bond

LEWIS STRUCTURES OF COMPOUNDS In writing Lewis structures, the most important consideration forming a stable compound is that the atoms attain a noble gas configuration. The most difficult part of writing Lewis structures is determining the arrangement of the atoms in a molecule or an ion. In simple molecules with more than two atoms, one atom will be the central atom surrounded by the other atoms.

Cl 2 O has two possible arrangements. The two chlorines can be bonded to each other. Cl-Cl-O The two chlorines can be bonded to oxygen. Cl-O-Cl Usually the single atom will be the central atom. (also usually the “leftist or lowest” on the Periodic Table)

Valence Electrons of Group A Elements Atom Group Valence Electrons Cl VIIA 7 H IA 1 C IVA 4 N VA 5 S VIA 6 P VA 5 I VIIA 7

Procedures for Writing Lewis Structures Step 1. Obtain the total number of valence electrons to be used in the structure by adding the number of valence electrons in all the atoms in the molecule or ion. If you are writing the structure of an ion, add one electron for each negative charge or subtract one electron for each positive charge on the ion. Step 2. Write the skeletal arrangement of the atoms and connect them with a single covalent bond (two dots or one dash). Choose the “leftist or lowest” element as the central atom. Arrange terminal atoms symmetrically around the central atom. Hydrogen, which contains only one bonding electron, can form only one covalent bond. Oxygen atoms usually have a maximum of two covalent bonds (two single bonds, or one double bond).

Procedures for Writing Lewis Structures Step 3. Subtract two electrons for each single bond you used in Step 2 from the total number of electrons calculated in Step 1. This gives you the net number of electrons available for completing the structure by adding lone pairs of electrons to the terminal atoms until they have an octet. Any remaining electrons become lone pairs on the central atom.

Procedures for Writing Lewis Structures Step 4. Check that each atom is satisfied. If one atom doesn’t have an octet, move lone pairs of electrons in as bond pairs to make multiple covalent bonds. Do this symmetrically. Step 5. Check the total number of electrons in the structure and make sure it matches the number of valence electrons in step 1. (Also learn the number of bonds an atom prefers to make: H and F always 1 bond and terminal atom; C mostly 4 (and usually a central atom); halogens mostly 1; O and S mostly 2; N and P mostly 3)

Write the Lewis structure for H 2 O. The total number of valence electrons is eight, two from the two hydrogen atoms and six from the oxygen atom. The two hydrogen atoms are connected to the oxygen atom which is central. Write the skeletal structure: : H: O H or H: O: H Place two dots between the hydrogen and oxygen atoms to form the covalent bonds. Subtract the four electrons used from eight valence electrons to obtain four electrons yet to be used around the oxygen. (Why not the H? )

Distribute the four remaining electrons in lone pairs around the oxygen atom. (Hydrogen atoms cannot accommodate any more electrons. NEVER have more than 1 bond to H or have lone pairs around H. ) or H: O: H : : H: O: H Thearrangements shape of the molecule not shown because by the Lewis These are Lewisisstructures each structure. atom has a noble gas electron structure.

Covalent bonding with equal sharing of electrons occurs in diatomic molecules formed from one element. hydrogen chlorine iodine A dash may replace a pair of dots that represent a bond: H-H nitrogen

Practice Lewis Structures Look in packet for practice sheet and work with one partner to draw the Lewis structures on separate paper. Bring them up to show on the document camera.

Complex Lewis Structures Do the Lewis structures for the following with a partner: HCN, CH 4, SO 3, CH 3 OH, SF 6, PCl 3, NO 2 Some will have EXCEPTIONS to the Octet Rule.

Complex Lewis Structures Exceptions to Octet Rule: Expanded valence shell: any central atom with outermost e-s in period 3 or below has d orbitals available for bonding and can hold 10 or 12 e-s Electron deficient or free radical structures: have less than 8 e-s and will be very reactive compounds

Complex Lewis Structures There are some molecules and polyatomic ions for which no single Lewis structure consistent with all characteristics and bonding information can be written. When more than one structure satisfies the rules, we call them resonance structures. Real molecule is a hybrid of all possible Lewis structures. Resonance stabilizes the molecule. Try O 3.

DRAWING LEWIS STRUCTURES Multiple Bonds: O 2 and N 2 Multiple Central Atoms: C 2 H 6, N 2 H 4, C 3 H 8, C 6 H 6, CH 3 NH 2, CH 3 COOH

Compounds Containing Polyatomic Ions A polyatomic ion is a stable group of atoms that has either a positive or negative charge and behaves as a single unit in many chemical reactions. Practice: NH 4+, SO 32 -, NO 3 -, I 3 -,

A scale of relative electronegativities wasacross developed by Electronegativity generally increases left to right a period. Linus Pauling. Electronegativity decreases down a group for representative elements. Metals are low in EN and nonmetals are high.

ELECTRONEGATIVITY Electronegativity: The relative attraction that an atom has for a pair of shared electrons in a covalent bond. If the two atoms that constitute a covalent bond are identical then there is equal sharing of electrons. This is called nonpolar covalent bonding. Ionic bonding and nonpolar covalent bonding represent two extremes.

ELECTRONEGATIVITY If the two atoms that constitute a covalent bond are not identical then there is unequal sharing of electrons. This is called polar covalent bonding. One atom assumes a partial positive charge and the other atom assumes a partial negative charge. This charge difference is a result of the unequal attractions the atoms have for their shared electron pair.

ELECTRONEGATIVITY H-H are the same atom, and have the same “greediness, ” so the two atoms are forced to share equally. F-F same - forced to share equally. If two atoms have diff EN, the one with higher EN will “take” the e-s in the pair more often than the other atom. H-F are not the same atoms, and are not equal in greediness, F is far greedier, takes the e-s more than half the time.

Partial positive Partial negative charge on hydrogen. charge on fluorine. Polar Covalent Bonding in HF + : : H F - Fluorine has a greater attraction for the The shared electron Shared electron pair is shared electron pair than hydrogen. closer to fluorine than to hydrogen.

Types of Covalent Bonding The polarity of a bond is determined by the difference in electronegativity values of the atoms forming the bond. If the electronegativity difference between two bonded atoms is greater than 1. 9 to 2. 0, the bond will be more ionic than covalent. If the electronegativity difference is greater than 2, the bond is strongly ionic. If the electronegativity difference is less than 1. 9 but greater than 0. 5, the bond is polar covalent. If the electronegativity differences is 0. 5 or less, the bond in nonpolar covalent.

Types of Covalent Bonding Estimate whether a bond is polar cov, mostly pure or nonpolar cov or ionic by finding the absolute value of the difference in EN between the two atoms in the bond. _____________________ | | | | 0. 2. 4. 6. 8 1 1. 2 1. 4 1. 6 1. 8 2. 0 2. 2 2. 4 2. 6 nonpolar | polar cov | mostly ionic covalent Practice: H-Cl, C-O and C=O, N-Cl, Ca-N H-Cl C-O C=O N-Cl Ca-N 0. 9 0. 5 1. 0 0. 0 2. 0 pol cov “ “ “ nonpol ionic

Molecular Geometry (bent) = bent One more geometry is trigonal pyramidal.

Some Geometric Figures Linear 2 atoms on opposite sides of central atom 180° bond angles 180° Trigonal Planar 3 atoms form a triangle around the central atom Planar 120° bond angles Tetrahedral 4 surrounding atoms form a tetrahedron around the central atom 109. 5° bond angles 120° 109. 5°

Some Geometric Figures Trigonal Pyramidal 3 atoms form a triangular pyramid beneath the central atom Not planar ~109° bond angles Derivative of tetrahedral geometry

The Valence Shell Electron Pair Repulsion (VSEPR) Model The VSEPR model is based on the idea that electron pairs will repel each other electrically and will seek to minimize this repulsion. To accomplish this minimization, the electron pairs will be arranged as far apart as possible around a central atom. The 3 -dimensional arrangement of the atoms within a molecule determines molecular interactions (physical properties and chemical reactions).

Be. Cl 2 is a molecule with only two pairs of electrons around beryllium, its central atom. Its electrons are arranged 180 o apart for maximum separation. LINEAR EP ARRANGEMENT & MOLECULAR GEOMETRY

BF 3 is a molecule with three pairs of electrons around boron, its central atom. o apart for Its electrons are arranged 120 This arrangement of atoms is called maximum separation. trigonal planar.

CH 4 is a molecule with four pairs of electrons around carbon, its central atom. An obvious since choicethe for molecule its atomic is arrangement However, 3 o angle between its atoms with all of is a 90 dimensional the molecular structure is its atoms in a single plane. o tetrahedral with a bond angle of 109. 5.

Ball and stick models of methane, CH 4, and carbon tetrachloride, CCl 4.

Tetrahedral Shapes Tetrahedral 4 areas of electrons around the central atom 109. 5° bond angles All Bonding = tetrahedral 3 Bonding + 1 Lone Pair = trigonal pyramid 2 Bonding + 2 Lone Pair = bent

Tetrahedral Derivatives

Ammonia, NH 3, has four electron pairs around nitrogen. The arrangement of electron pairs around nitrogen is tetrahedral.

NH 3 has one lone pair of electrons. The NH 3 molecule is trigonal pyramidal.

Water has four electron pairs around oxygen. The arrangement of electron pairs around oxygen is tetrahedral.

H 2 O has two lone pairs of The H 2 O molecule electrons. is bent.

The VSEPR Model Summary: electron pair arrangement depends upon # of bonded atoms and lone pairs around the central atom. Lone pairs exert more repulsion that bond pairs. Count # of bonded atoms (B) and # of LPs (E). EP Arrangements: 2 = Linear; 3 = Trigonal Planar; 4 = Tetrahedral Within these electron pair arrangements, the molecular geometry is based only on “seeing” the atoms. I call the molecular geometry the family members and the electron pair arrangement is the electron pair family the members are in. Electron Pair Families and their Molecular Geometry members: Linear: AB 2, linear only Trigonal Planar: AB 3 trigonal planar; AB 2 E bent Tetrahedral: AB 4 tetrahedral; AB 3 E trigonal pyramidal, and AB 2 E 2 bent

VSEPR Practice Use practice sheets in packet to fill in Electron Pair Arrangement and Molecular Geometry and Bond Angles.

Dipole Moments A dipole is a molecule with positively and negatively charged ends Polar covalent bonds or molecules have one end slightly positive, +; and the other slightly negative, (not “full” charges, come from nonsymmetrical electron distribution) Dipole Moment is a measure of the size of the polarity. (We are NOT going to worry about the Debye unit or actual numbers for dipole moment, just whether a molecule has a dipole or not. )

Polarity of Molecules For a molecule to be polar it must - have polar bonds electronegativity difference - theory bond dipole moments – measured values - have an unsymmetrical shape vector addition Polarity affects the intermolecular forces of attraction

DIPOLE Polar covalent bonds and unsymmetrical shape cause molecule to be polar NO DIPOLE Polar covalent bonds, but nonpolar molecule, because vectors cancel

Adding Dipole Moments Table 10. 3 64