Section 3.1 (pg. 78‐84) Homework: Lewis Symbols Pg. 82 #2 – 4 Pg. 84 # 2, 4, 5, 7‐10

Objectives: 1) Define valence electron, electronegativity, and ionic bond

2) Use the Periodic Table and Lewis structures to support and explain ionic bonding

3) Explain how an ionic bond results from the simultaneous attraction of oppositely charged ions. Bonding Theory: Valence Electrons & Orbitals

 To describe where electrons exist in the atom, chemists created the concept of an orbital.  Orbital – region of space around an atom’s nucleus where an electron may exist  An “orbital” is not a definite race track, it is a 3‐D space that defines where an electron may be (like a rain drop in a cloud)

 For bonding study we are only concerned with an atom’s valence orbitals (the volume of space that can be occupied by electrons in an atom’s highest energy level)

 WHY? Bonding only involves valence e‐’s because lower energy levels are held so strongly by their positively charged nucleus

FYI Read pg. 78‐79 for the history on Bonding Theory Bonding Theory: Valence Electrons & Orbitals

 According to bonding theory, valence electrons are classified in terms of orbital occupancy. (0 = empty, 1 = half filled , 2 = full)

 An atom with a valence orbital that has a single electron can theoretically share that electron with another atom  Such an electron is called a BONDING ELECTRON

 An atom with a full valence orbital (2 e‐’s), repels nearby orbitals and wants to be alone  Such a pairing is called a LONE PAIR The Four Rules of Bonding Theory

1. The first energy level has room for only one orbit 2e‐ 2e‐ + ‐ 2 p ‐ can only hold 2 e ’s max He 2. Energy levels above the first have room for four orbitals = 8 electrons max ‐ Noble gases have this structure; their lack or reactivity indicates that eight electrons filling a valence orbital is very stable (Remember the OCTET RULE)

FYI: Only C, N, O, 8e‐ 2e‐ 2e‐ 2e‐ 2e‐ and F atoms 8e‐ always obey the 2e‐ EXCEPTIONS: B = stable with 6 valence e‐ (3 orbitals) octet rule when 18 p+ bonding P = stable with 10 valence e‐ (5 orbitals) Ar S = stable with 12 valence e‐ (6 orbitals) The Four Rules of Bonding Theory

3. An orbital can be unoccupied, or it may contain one or two electrons – but never more than two (Pauli Exclusion Principle)

4. Electrons “spread out” to occupy any 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 Atomic Models: LEWIS SYMBOLS (aka Lewis Dot Diagrams, Electron Dot Diagrams, LDD, Lewis Models

• Named after Gilbert Lewis who is responsible for the Octet Rule. He reasoned that all atoms strive to be like the nearest noble gas. • Used dots or ‘x’ to represent the valence electrons • The inner electrons and the nucleus are represented by the element symbol

How to draw Lewis Symbols: 1. Write the element symbol 2. Add a dot to represent each valence e‐ 5. Start by placing valence e‐’s singly into each of the four valence orbitals (4 sides) 6. If additional e‐’s need to be placed, start filling each of the orbitals with a second e‐ up to 8

Q: Which element has 4 bonding e‐’s? Which has 3 lone pairs and 1 bonding e‐? Pracce  Draw the Lewis Symbols for the elements in Period 3  For each one indicate how many lone pairs or bonding electrons are present

 It is important to remember that the Lewis symbols do not mean that electrons are dots or that they are stationary.

 The four sides represent the four orbitals that may be occupied by electrons; it is a simplistic 2‐D diagram of a complex 3‐D structure Electronegativity ‐A measure of the force that an atom exerts on electrons of other atoms; (the “pull” on bonding electrons)

‐Each atom is assigned a value between 0.0 – 4.0; the larger the number the greater the “pulling” force

‐Example: Fluorine has an EN = 4.0 and francium has an EN = 0.7 ‐This means fluorine wants to pull on other electrons very strongly ‐This means francium doesn’t want to pull on other electrons

Q: Does lithium (EN = 1.0) want to lose or gain an electron to be stable?

Q: Does fluoride (EN = 4.0) want to lose or gain an electron to be stable?

Do you see any relation to their electronegativity numbers? So how do we assign each atom an electronegativity number?

a) The farther away from the nucleus that electrons are, the weaker their attraction to the nucleus Cesium's valence electrons are not EN = 0.8 held as tightly by its nucleus EN = 2.6 . because the atom is larger

b) Inner electrons shield valence electrons from the attraction of the positive nucleus 1 e‐ Potassium’s valence electrons are not 8e‐ 5e‐ attracted to its nucleus as much as 8e‐ EN = 0.8 2e‐ 2e‐ Nitrogen’s valence electrons because their EN = 3.0 7p+ 19p+ are more inner electrons present in K N K

c) The greater the number of protons in the nucleus, the greater the attraction for more electrons Bromine has more protons (+ charge) 14p+ EN = 3.0 35p+ Si EN =1.9 which attracts the negative charge of Br electrons more so than silicon’s 14 protons Electronegativity

In this 3‐D image, the electronegativity scale is vertical.

Q: What is the EN trend within a period and a group?

Q: Which element has the highest EN? Give three reasons why? Why do we care about electronegativity??

 Imagine that two atoms, each with an orbital containing one bonding electron, collide in such a way that these half‐filled orbitals overlap.

 As the two atoms collide, the nucleus of each atom attracts and attempts to “capture” the bonding electrons of the other atoms

 A “Tug of War” over the bonding electrons occurs

 Which atom wins?  By comparing the electronegativities of the two atoms we can predict the result of the contest = 3 different types of bonds result Covalent Bonding  Both atoms have a high EN so neither atom “wins”  The simultaneous attraction of two nuclei for a shared pair of bonding electrons = covalent bond

Cl2 = diatomic  EN difference can be zero = Cl – Cl EN = 3.2 EN = 3.2

 EN difference can be small = H ‐ Cl EN = 2.2 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 in Section 3.3 Ionic Bonding  The EN of the two atoms are quite different  The atom with the higher EN will remove the bonding e‐ from the other atom electron transfer occurs

 Positive and negative ions are formed which electrically attract each other

EN = 0.9 EN = 3.2 Metallic Bonding  Both atoms have a relatively low EN so atoms share valence electrons, but no actual chemical reaction takes place

 In metallic bonding: a) e‐’s are not held very strongly by their atoms b) the atoms have vacant valence orbitals ‐ This means the electrons are free to move around between the atoms and the (+) nuclei on either side will attract them

Analogy: The positive nuclei are held together by a glue of negative e‐’s 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.) Summary of Bonding Theory: Chemical Bond = competition for bonding electrons

1) Atoms with equal EN = electrons shared equally If both have high EN = covalent bond (equal = non‐polar) If both have a low EN = metallic bond 2) Atoms with unequal EN = covalent bond (unequal = polar) 3) Atoms with unequal EN = ionic bond

metallic The nature of chemical bonds changes in a continuous way, creating a broad range of characteristics. PRACTICE

Copy pg. 84 – Bonding Theory Summary into your Notes

Pg. 82 #2 ‐ 4

Pg. 84 # 2, 4, 5, 7‐10