Bonding Theory, Lewis Diagrams, & Electronegativity

1. Bonding Theory, Lewis Diagrams, & Electronegativity Pages 78-84 Be able to create Lewis Diagrams Understand electronegativity and how it influences the different types of bonds

2. Bonding Theory • Bonding is one of the most theoretical concepts in chemistry • No direct evidence of bonds between atoms • Created from indirect evidence and logic • Most atoms are highly reactive and combine with other atoms to form stable molecules and compounds with uniques properties • This reactivity is due to the electron structure: atoms with incomplete outer electron shells gain, lose, or share electrons in order to fill it • This gaining, losing, or sharing results in a bond, which is an attractive force that holds the atoms together

3. Continued... • Bonding theory is an attempt to explain the physical and chemical properties of matter • Hardness • Chemical reactivity • melting/boiling points • The most complete theory was developed by Gilbert Lewis in 1916 and Linus Pauling in 1939

4. Continued. . . • Gilbert Lewis • Atoms achieve stable electron configurations by sharing or transferring electrons • This only works if the atoms sharing stay close together, resulting in an attractive force - a covalent bond • His ideas also explain ionic bonds which results from the attraction between positive and negative ions resulting from the transfer of electrons • Linus Pauling • Explained why certain electron arrangements are stable • Reasoned that electron sharing is a continuum from equal attraction to total transfer

5. Valence Electrons and Orbitals • As mentioned before, bonding involves changing the electron arrangement of atoms • Evidence shows: • Electrons are found in energy levels - the closer to the nucleus, the lower the energy level • Electrons in the highest energy level (outermost) are calledvalence electrons • According to the Bohr Model, the maximum number of electrons for the first three energy levels are 2, 8, and 8 • A filled outer energy level makes an atom stable or unreactive(inert) – • Noble Gases

6. To describe where electrons exist in the atom, chemists created the concept of an orbital • Orbital - is a specific volume of space in which an electron of a certain energy is likely to be found • Can be thought of as a 3-D space that defines where an electron may be (e.x., raindrop in a cloud or blades of a fan) • For bonding, we are only concerned with the valence orbitals - which are the volumes of space that can be occupied by electrons in the atom’s highest energy level • Why? → electrons in lower energy levels are held so strongly by the nucleus that they remain unchanged & the valence orbitals are not completely filled

7. According to Bonding Theory, valence electrons are classified in terms of orbital occupancy • No electrons = empty • One electron = half filled • Two electrons = full • An atom that has a single electron in its valence orbital can theoretically share that electron with another atom • This electron is called a bonding electron • A full valence orbital (2 electrons) has a repelling effect on nearby electrons and want to be alone • These two electrons in the same orbital are called a lone pair

8. Four Rules of Bonding Theory 2e- 2 p+ He • The first energy level only has one orbital with a maximum of two electrons • The second and higher energy levels have room for four orbitals meaning these orbitals hold a maximum of eight electrons • The Noble Gases have eight electrons in their valence orbitals → indicates that this structure is highly stable • This is known as the OCTET RULE • BUT, there are exceptions to the octet rule • B - stable with 6 electrons (3 orbitals) • P - stable with 10 electrons (5 orbitals) • S - stable with 12 electrons (6 orbitals) • And, C, N, O, F atoms are the ALWAYS obey the octet rule 8e- 8e- 2e- 18 p+ Ar

9. 3. Pauli Exclusion Rule - an orbital can be unoccupied, or it may contain one or two electrons - but never more than two can be in the same region of space at the same time 4. Electrons will “spread out” to occupy empty valence orbitals before forming electron pairs “Aluminum has three half-filled orbitals and one vacant orbital.” How would you describe Sulfur? Never more than 2e- in an orbital

10. Atomic Models: Lewis Symbols • Also known as Lewis Dot Diagrams, Electron Dot Diagrams, LDD, Lewis Models • Named after Gilbert Lewis who is responsible for the octet rule - reasoned that all atoms strive to be like the Noble Gases • The innermost electrons and nucleus are represented by the element symbols • Dots or “x’s” are used to represent the valence electrons How to Draw Lewis Symbols: • Write the element symbol • Add a dot or an x to represent each valence electron • Start by placing valence electrons singly into each of the four valence orbitals (4 sides) • If additional electrons need to be placed, start filling each of the orbitals with a second electron up to 8

11. Practice Time Please draw the Lewis Symbols for the following elements. Indicate the number of bonding electrons (BE) and lone pairs (LP) are present: 1. Fluorine 2. Magnesium 3. Hydrogen 4. Argon 5. Silicon 6. Sulfu **Lewis Structures do not mean that electrons are dots or stationary. Also, the four sides represent the four orbitals that may be occupied by electrons - it is a simple 2-D representation of a 3-D structure**

12. Electronegativity • Electronegativity is used to describe the relative ability of an atom to attract a pair of bonding electrons in its valence level • IN OTHER WORDS: A measure of the force that an atom exerts on electrons of other atoms; (the “pull” on bonding electrons) • It is assigned on a scale developed by Linus Pauling - atoms are assigned a number from 0.0 to 4.0; larger values means a greater “pulling” force • Example: Fluorine has an EN of 4.0 while Francium has an EN of 0.7. What does this mean about these two elements? → F wants the electrons & Fr wants to give them up Q: Does Lithium (EN = 1.0) want to gain or lose an electron to be stable? Q: Does Oxygen (EN = 3.4) want to gain or lose an electron to be stable?

13. How do we assign atoms an electronegativity number? Cesium's valence electrons are not held as tightly by its nucleus because the atom is larger • The farther away from the nucleus the valence electrons are, the weaker their attraction to the nucleus • Inner electrons shield the valence electrons from the attraction of the positive nucleus dfgdfgdfgdfgdf • The more protons are in the nucleus, the greater the attraction there is for electrons EN = 0.8 EN = 2.6 1 e- 8e- 8e- 2e- 19p Potassium’s valence electrons are not attracted to its nucleus as much as Nitrogen’s valence electrons because their are more inner electrons present in K 5e- 2e- 7p EN = 0.8 EN = 3.0 Bromine has more protons (+ charge) which attracts the negative charge of electrons more so than silicon’s 14 protons 35p+ Br 14p+ Si EN = 3.0 EN =1.9

14. Electronegativity Trends In this 3D image, the electronegativity scale is vertical (the higher the atom, the higher the electronegativity) Q: What is the EN trend within a period and a group? Q: Which element has the highest EN? Give three reasons why?

15. Why is Electronegativity important? • Electronegativity is used to help determine how atoms bond → imagine that two atoms collide in such a way that the orbitals containing one bonding electron overlap • The nucleus of each atom will attract and attempt to “capture” the bonding electron from the other atom • A “Tug-of-War” will result over the electrons and the electronegativities of the atoms will determine who “wins” • There are three types of bonds which can result from this collision and resulting “battle”

16. Covalent Bonding • Both atoms have a high EN so neither one “wins” and they share the bonding electron • EN difference can be zero • E.g. Cl-Cl • 3.2-3.2=0.0 • C • EN difference can be small • E.g. H-Cl • 3.2-2.2=1.0 EN = 3.2 EN = 3.2 This is called a polar covalent bond – because one side pulls on the electrons more but we will learn more about this later

17. Ionic Bonding • There is a large difference between the EN of the two atoms so one of them “wins” and takes the bonding electron (electron transfer) • A positive and negative ion are formed and they bond because they electrically attract one another • In ionic compounds, there are enormous numbers of both types of ions → these ions arrange themselves in positions where the maximum total attraction between positive and negative charge occurs which determines the numerical ratio of ions in a compound. • Always form a regular repeating 3-D pattern known as a crystal lattice EN = 0.9 EN = 3.2

18. Metallic Bonding • Both atoms have a relatively low EN so atoms share valence electrons, but no actual chemical reaction takes place • In metallic bonding: • e-’s are not held very strongly by their atoms • the atoms have vacant valence orbitals This means the electrons are free to move around betweenthe atoms and the (+) nuclei on either side will attract them Analogy: The positive nuclei are held together by a glue of negative e-’s

19. Metallic Bonding Visual This diagram represents Mg atoms that have released their electrons and are embedded in a sea (or glue) of electrons. Note: These metal atoms don’t have to be in a particular arrangement to attract each other therefore they are flexible, malleable and ductile = useful alloys (Brass, Stainless Steel, etc.)

20. Summary of Bonding Theory A chemical bond is a competition for bonding electrons • Atoms with an equal EN: • Both have a high EN = covalent bond (equal = nonpolar) • Both have a low EN = metallic bond • Atoms have an unequal EN • Atoms have a similar EN = covalent bond (not equal = polar) • Atoms have significantly different EN = ionic bond

21. Practice Questions: Page 83, # 2-4 Page 84, # 2, 4, 5, 7-10