UNIT 4- BONDING (Ch

S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 1 NAME______PERIOD______UNIT 6 NOTES: BONDING

STUDENT OBJECTIVES: Your fascinating teachers would like you amazing learners to be able to… 1. List the types of ions involved in ionic bonding. 2. Describe the basic structure of ionic crystals. 3. Predict the presence of ionic bonding based up the types of elements involved. 4. Calculate the likelihood of ionic bonding based on differences in electronegativity. 5. Determine the number of valence electrons in an atom. 6. Draw the Lewis dot structure of atoms, ions and ionic compounds. 7. Predict the oxidation numbers of elements in the “s” and “p” blocks of the periodic table. 8. Memorize the common oxidation numbers of common transition metals. 9. Analyze typical physical properties of ionic properties based upon the predicted structure. 10. Predict the presence of covalent bonding based upon the types of atoms involved in the molecule. 11. Calculate the bond polarity in a covalent bond using differences in electronegativity. 12. Memorize the names and formulas of the diatomic elements. 13. Illustrate the polarity of a bonding using the symbols — and + to show the electron distribution in the bond. 14. Draw Lewis dot structures for covalent molecules and polyatomic ions that have a central atom. 15. Draw Lewis dot structures for central atoms that have incomplete and expanded octets. 16. Explain the concept of resonance in covalent bonding and use Lewis dot structures to support your explanation. 17. Use Valence Shell Electron Pair Repulsion (VSEPR) theory to predict the three dimensional shape of covalent molecules with a central atom. 18. Predict the molecular polarity of a covalent molecule from its VSEPR structure. 19. Explain the common physical properties of covalent molecules based upon the predicted structure. 20. Outline the structure and function in metallic bonding. 21. Define the type of intermolecular forces involved in ionic and covalent substances. 22. Predict the type of intermolecular forces involved in ionic and covalent substances. 23. Compare and contrast the properties of covalent and ionic substances. S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 2

I. INTRODUCTION – THE THREE TYPES OF BONDS We have THREE types of bonds that we will discuss that vary, depending on those elusive little electrons!  Ionic bonding electrons are EXCHANGED: Metals are LOSERS and non-metals are WINNERS in the electron world!

 Covalent bonding electrons are SHARED: sometimes evenly and sometimes one element gets a little greedy with the electrons!

 Metallic bonding electrons live in a little “SEA” where they move from metal ion to metal ion, swimming between the ions like fish! S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 3

II. IONIC BONDING STRUCTURE a. General Structural Unit : Ionic structure has alternating positive and negative charges in a rigid ______. Electrons are ______from one species to another, not shared.

b. Prediction by Substance Type: If you see the following types of particles present in the formula of a compound, we usually predict the compound to be ionic.

Cation (+ ion, lost e-‘s ) Anion (− ion, gained e-‘s) Metal Non-metal Metal Polyatomic Ammonium Non-metal Ammonium Polyatomic

c. Prediction by Electronegativity Calculation: Electronegativity (the pull an atom has to gain an electron) can be used to verify whether a predicted bond is actually ionic or not. If the ΔEN is ______, the bond exhibits 50% or more ionic character and definitely considered to be ionic.

Example 6-1. Each of the following is predicted to have ionic bonds using their particles present as a predictor. Calculate the difference in electronegativity to determine if the bonds are actually >50% ionic.

(a) NaCl (b) BeI2 (c) AlBr3 (d) CuCl2

NOTE: YOU DO NOT NEED TO PERFORM A BOND CALCULATION UNLESS INSTRUCTED! GO BY YOUR PREDICTIONS!

LEWIS DOT STRUCTURES OF ELEMENTS AND IONS d. ______Electrons: Outer _____ and _____ electrons – the electrons most likely to be involved in ______. S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 4

Valence electrons are most helpful for predicting the ______(often referred to as ______) for elements in the s and p blocks. The oxidation number is a method used for tracking electrons, and is equal to the ionic charge for cations (+ ions) and anions (— ions).

Below are the common oxidation numbers for the representative (main group) elements: Group # Valence Typical Oxidation Number Number Electrons 1 1 +1 (when ______one e—) 2 2 +2 (when ______two e—) 13 3 +3 (when ______three e—) 14 4 Non-metals: -4 (when ______four e—) Metals: +2 when they lose their p2 electrons +4 when they lose their s2p2 electrons 15 5 Non-metals: -3 when with metals Metals: +3 when they lose their p3 electrons +5 when they lose their s2P3 electrons 16 6 -2 (when ______two e—) 17 7 -1 (when ______one e—)

ALSO: The oxidation number for pure elements (elements by themselves shown without any kind of charge) is always ______.

Example 6-2. Notice that we have not discussed any common charges for Group 18 (Noble/Inert Gas) elements. Why do you think this is so?

e. Transition Metals The transition elements are notoriously independent! Some follow valence electron predictions and some don’t. Sometimes, you have to use the other ion’s oxidation number in a compound to figure out the transition element’s oxidation number… or be told what its oxidation number will be. There are a few to memorize… I will refer to them as the “staircase”

Al 0, +3

Zn Ga 0, +2 0, +3

Ag Cd In 0, +1 0, +2 0, +3

Example 6-3. Label the Periodic Table below with the number of valence electrons and the common oxidation numbers for the representative elements, as well as for zinc, cadmium, and silver. S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 5

f. LDS of ATOMS: To draw the LDS (Lewis Dot Structure) for atoms, simply imagine a square around the element symbol and then place the valence electrons around the symbol. Remember Hund’s Rule: put one in each spot, then double up! LDS of IONS: (1) Determine how many valence electrons the ion will have! Most metals will lose all their valence (“s” and “p”) electrons, or you will be told what the charge will be. Non-Metals will gain enough electrons to fill their “p” sublevel. (2) Place brackets around your Lewis Dot Structure (3) Add the charge as a superscript on the upper right side outside of the brackets.

ELEMENT LEWIS DOT ION LEWIS DOT Tin(II) Tin 2+ (Sn2+) [• Sn •]

Tin(IV) 4+ (Sn4+) [ Sn ]

NOTICE A PATTERN? Tin has two possible charges/oxidation numbers. How were these indicated in the name? Example 6-4. Draw the Lewis dot structures for the following atoms and ions.

ELEMENT LEWIS DOT ION LEWIS DOT Lithium Lithium ion

Boron Boron ion

Oxygen Oxide ion S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 6

Bismuth Bismuth (+3) ion

Antimony Antimony (+5) ion

Phosphorus Phosphide ion

Iodine Iodide ion

Xenon Helium (THIS ONE IS TRICKY…)

Example 6-5. NOTICE A PATTERN? How did the suffix/ending of a non-metal change when it became an ion?

III. LEWIS DOT STRUCTURES OF IONIC COMPOUNDS (1) Draw the Lewis dot structure for the positive ion with charge and brackets. (2) Draw the Lewis dot structure for the negative ion with charge and brackets. (3) Continue to draw positive and negative ions until the charges cancel themselves out to zero. It is proper notation to alternate positive and negative ions.

IONIC COMPOUND LEWS DOT A GLIMSE OF THE FUTURE: WRITE THE FORMULA UNIT! magnesium chloride

Example 6-6. Draw the Lewis Dot Structure for the following ionic compounds: IONIC LEWS DOT A GLIMSE OF THE COMPOUND FUTURE: WRITE THE S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 7 FORMULA UNIT! sodium chloride

calcium fluoride

aluminum oxide

LEGGETT PreAP Bonding 6-4 Ionic Properties (4:05) https://vimeo.com/53288361 http://youtu.be/bGjlZeem9Lk

IV. IONIC COMPOUND PROPERTIES Ionic bonds have the following characteristics: PROPERTY EXPLANATION

High melting (Tm) and boiling (Tb) points – most are ______at In order for a phase change to occur, strong room temperature ion-ion attractions that are present must be broken. This takes a significant amount energy! Brittle and Cleave when struck Caused by ion-ion repulsion When soluble, ions are attracted to the partial Many (but not all) are ______in water charges on water Movement of electrons require mobile ions - Non-conductive as ______ionic solids are rigid Ions become mobile and allow electron Conductive in ______and ______(dissolved in water) states movement

LEGGETT PreAP Bonding 6-5 BOND POLARITY (9:34) https://vimeo.com/53288362 http://youtu.be/PUNoV6zFosY

V. COVALENT BONDING STRUCTURE a. General Structural Unit: Sometimes the more electronegative element is not ______enough to actually take another atom’s electron away. The electrons of both atoms are placed between the two and are ______, which forms a covalent bond. Covalent bonds are found in molecular compounds and within (not between) polyatomic ions. The sharing that takes place can happen ______(Non-Polar Covalent) or ______(Polar Covalent).

b. Prediction by Substance Type: Covalent bonds are always between two ______elements!

c. Prediction by Electronegativity Calculation:

1. Non-polar covalent: electrons are shared approximately ______and ΔEN < 0.3 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 8

Now is a good time to memorize the ______. These are elements found in nature with 2 atoms bonded together. They are the most common examples of non-polar covalent molecules.

Hydrogen (H2), Nitrogen (N2), Oxygen (O2), Fluorine (F2), Chlorine (Cl2), Bromine (Br2), Iodine (I2) A trick for memorizing them is the phrase “I Bring Clay For Our New House.”

2. Polar Covalent: electrons are ______shared and 0.3 ≤ ΔEN < 1.7 When you determine that a bond is “polar”, you indicate its direction of pull (also known as electron distribution) by drawing an arrow OVER the bond pointing to the atom which is the most electronegative. This shows that the electron is being shared MUCH closer to the more electronegative element, causing that side of the molecule to become “partially” negative. Because there is a shift in charge within the molecule it causes the molecule to become “partially” positive on the other side. Since these are not true charges, we cannot use a + or – sign. Instead, we use the lower-case Greek letter delta + to indicate “partially positive” or — to indicate “partially negative”.

H ─ O

Example 6-7. Show the direction of pull (electron distribution) for the bonds within F2 and HCl.

Example 6-8. How many polar bonds and how many non-polar bonds are there in C2Cl6?

LEGGETT PreAP Bonding 6-6 Covalent Lewis Dot (10:34) https://vimeo.com/53288363 http://youtu.be/kY_xQwAVwe0

VI. LEWIS DOT STRUCTURES OF COVALENT MOLECULES We aren’t very interested in which element brought which electron to the table. Think of the process like a potluck: each element brings its valence electrons. Guidelines : (1) Add up the total number of valence electrons in the particle. THIS IS THE MAXIMUM NUMBER OF DOTS YOU ARE ALLOWED TO DRAW!!! (2) Decide which atom is the central atom—this is usually the atom present in fewest number, or, if there is the same number of all atoms, it is the ______electronegative. However, hydrogen and fluorine can never be central atoms! Write down the central atom, and place all of the other elements around it. (3) Form covalent bonds between the central atom and the peripheral atom using ______electrons per bond. Sometimes a covalent bond is designated with “ ─ “ (dash) which indicates 2 electrons being shared. S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 9 (4) Start filling up atoms to satisfy the ______Rule, which states that most elements need ______electrons to be stable! There are a few exceptions to the Octet Rule, WHICH YOU MUST MEMORIZE:  H only want 2 electrons (one bond)  Be only wants 4 electrons (two bonds)  B is satisfied with 6 electrons (three bonds)  Sometimes S and P have more than 8 electrons (called an “Expanded Octet”) (5) Place any left-over electrons on the central atom in pairs. If the central atom has extra pairs of electrons beyond the typical “4”, draw a line from the extra pair to the central atom… this is called an “Expanded Octet”. While S and P are the most common two elements to have expanded octets, other elements can occasionally be forced to have them as well. This includes Group 18 elements! (6) After placing all of the electrons, if one of the atoms does not have a complete octet, atoms will share a pair of electrons to make a double bond.

RANDOM THOUGHT: We are going to have you write a general formula for each of these. Trust us…you will understand the purpose in just a little while! Use the letter “A” to represent the central atom, the letter “B” to represent peripheral atoms, and the letter “X” to represent non-bonded electron pairs on the central atom. This is the “ABX” Formula.

Example 6-9. Draw the LDS for the following compounds containing single bonds.

GENERAL MOLECULE LEWIS DOT STRUCTURE “ABX” FORMULA

NCl3 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 10

http://vimeo.com/5352896 LEGGETT PreAP Bonding 6-7 Covalent Lewis Dot (11:26) http://youtu.be/4Zqh9olD2Mw 2

Example 6-10. Draw the LDS for the following compounds containing double or triple bonds.

GENERAL MOLECULE LEWIS DOT STRUCTURE “ABX” OR ION FORMULA

2− CO3

Example 6-11. Draw the LDS for structures that violate the Octet Rule (Expanded Octets). Put any extra electrons on the central atom.

GENERAL MOLECULE LEWIS DOT STRUCTURE “ABX” OR ION FORMULA

ClF3 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 11

LEGGETT PreAP Bonding 6-8 Resonance (6:17) http://vimeo.com/53528961 http://youtu.be/XA6idB8LRB8

a. LDS for ______Structures: Having resonance structures means that there is a double bond that can be in more than one possible location… but in actuality, the extra bond is shared equally among the multiple locations!

-2 For CO3 , there are actually three equivalent ways to draw the structure. The final structure is actually in between.

When drawing, you would show this….

But in actuality, the real structure is like this…

Questions to ask yourself when determining if a structure shows resonance: (1) Do I have any double bonds? (2) Can the double bond be moved to other locations?

─ Example 6-12. Draw Lewis structures for NO3 ion. Don’t forget brackets and charges, since it is an ion! Give the general “ABX” structure. S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 12

Example 6-13. Which of these examples has resonance structures?

b. LDS for ______: You might have noticed that polyatomic ions tend to have covalent bonds within them. However, these polyatomic ions can form ionic bond between the polyatomic anion and other cations (either a metal ion or ammonium). Therefore, the resulting structure will have the LDS shown as covalent within the polyatomic ion, and the LDS shown as ionic overall.

Example 6-14. Draw the LDS for the following ionic compounds containing polyatomic ions.

Na2SO4 NH4NO

LEGGETT PreAP Bonding 6-10 VSEPR (6:13) http://vimeo.com/53528959 http://youtu.be/BXF7SqPni9A

VII. VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORY – FOR COVALENTS When we draw a Lewis Dot Structure, we are representing the molecule in a 2-D way… but in actuality, the molecule is 3-D! The VSEPR Theory refers to the actual 3-D shape that a covalent molecule has. The VSEPR Theory is based on the premise that the valence electrons of each peripheral atom will repel each other strongly, and therefore cause the peripheral atoms to move as far from each other as possible. Placing the molecules as far away from each other as possible minimizes the electrostatic repulsion between them.

The VSEPR Theory assigns a shape “name” to a covalent molecule based on how many bonded atoms and how many non-bonded electron pairs are present around the central atom of the molecule. You must memorize these shape names, and know how to assign them! We will be using the “ABX” Structure to help us do this!

Example 6-15. Fill in the following chard about VSEPR shapes. Wait on the last column… we will finish it soon! # bonded # non- General VSEPR General Exampl Lewis Dot Structure atoms on bonded pairs ABX Shape Polarity e CA of e- on CA structure Name of Shape S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 13

SeF6 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 14 LEGGETT PreAP 6-11 Molecular Polarity (9:02) http://vimeo.com/54341122 http://youtu.be/WRTxUY4L0aI

VIII. COVALENT MOLECULE PROPERTIES & POLARITY OF COVALENT MOLECULES In order to discuss covalent compound properties, we need to discuss the two types of covalent molecules – Non-Polar Covalent Molecules and Polar Covalent Molecules. MOLECULE POLARITY IS DIFFERENT THAN BOND POLARITY!!! You can have polar bonds in a molecule, but the overall molecule can still be non-polar!

a. ______Molecules: Central atom has no non-bonded pairs of electrons and is surrounded by the same atoms. KEY: there is not preference for electron density on one part of the molecule compared to any other. We say the molecule is symmetrical around the central atom. Having this structure means that there is no “pull” holding one molecule to another molecule, meaning that the molecules are held together very weakly. We will discuss more about these “intermolecular forces” later. This gives Non-Polar Covalent Molecules the following properties:

 Low melting (Tm) and boiling points (Tb)  Most are not solids, but rather liquids or gases at room temperature  Does not conduct electricity in ANY state (solid, liquid, or gas) as there is no pathway for mobile electrons  Insoluble in water and other polar solvents, but soluble in other non-polar solvents  Exist as a MOLECULE, not as ions

b. ______Molecules: Central atom has non-bonded pair(s) of electrons OR is surrounded by different atoms. We say the molecule is unsymmetrical around the central atom. Being polar means that there is some pull between the molecules, which holds the molecules together (intermolecular forces). This gives Polar Covalent Molecules the following properties:

 Higher Tm & Tb than non-polar molecules, but much lower than ionic compounds (we’ll discuss more later!)  Most are liquids at room temperature (a few with large mass are solids)  Does not conduct electricity in ANY state (solid, liquid, or gas) as there is no pathway for mobile electrons  Typically soluble in water and other polar solvents, but not soluble in non-polar solvents  Exist as a MOLECULE, not as ions The presence of polar bonds in a molecule MAY OR MAY NOT cause the entire molecule to be polar. The basic question to ask yourself when determining polarity is: DOES THE MOLECULE LOOK THE SAME ALL THE WAY AROUND THE CENTRAL ATOM? If “yes”, the molecule is ______. If “no”, the molecule is ______. (Note: We only look at the sides that have something there… either a bonded atom or a non-bonded pair of electrons.) S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 15

(1) All four sides have the same element? NON-POLAR MOLECULE. (2) Three sides have one element, and the fourth side has a different element? POLAR MOLECULE. (3) Some sides have an element, and the other side(s) has a non-bonded electron pair(s)? POLAR MOLECULE.

ALSO: The reason why you saw the trends of solubility (as listed above) is that LIKE DISSOLVES LIKE… polar substances dissolve in polar solvents (like water). Non-polar substances dissolve in non-polar solvents (like oil).

Example 6-16. How many polar bonds are in CF4? Is the molecule itself polar? Would it more likely be soluble in water or oil?

Example 6-17. How many polar bonds are there in CO2? Is the molecule itself polar? Would it more likely be soluble in water or oil?

Example 6-18. How many polar bonds are there in ammonia (NH3)? Is the molecule itself polar? Would it more likely be soluble in water or oil?

LEGGETT PreAP 6-12 IMF (10:12) http://vimeo.com/54341121 http://youtu.be/1bQ2LbRQ1WU

IX. METALLIC BONDS Structure of Metallic Bonds: These bonds result when metal atoms donate their valence electrons to an ______that binds the atoms together. This actually forms a very strong bond! S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 16 Having this “sea of electrons” where electrons can move around gives metallic bonds some very specific properties. Properties of Metallic Bonds

(1) Have extremely high Tm & Tb due to the strong bonds present (2) Conduct electricity in their solid and liquid (molten) forms, as electrons are free to move (3) Can conduct heat very well, again because of all of the movement from the sea of electrons (4) Have ______(shine) due to light reflecting off the sea of electrons (5) Are ______(can be hammered into a sheet) (6) Are ______(can be pulled into a wire)

While metallic bonds are present in pure-element metals, they can also be present when you have a homogeneous mixture (solution) of different metal elements. A mixture of metal elements is referred to as an ______. Interesting Tidbit: Some common examples of alloys include… (1) Bronze: Cu and Sn (2) Brass: Cu and Zn (3) Sterling Silver: Ag and Cu (4) Steel: Fe, C, and Cr X. INTERMOLECULAR FORCES (IMF) ______bonding is what occurs between two atoms. When we have called a compound ionic or covalent, we have been referring to its intramolecular bonding.

However, for compounds to remain in a solid or liquid state, there must be some sort of pull between the molecules. This pull between the molecules is what we call ______forces.

In these notes, you’ve seen a lot about the different melting and boiling points of different types of bonds. These temperatures are greatly influenced by intermolecular forces. The stronger the type of intermolecular force, the higher the melting or boiling point will be.

The following are the types of intermolecular interactions in order of decreasing strength:

A. Ion-Ion: These are the dominant forces felt within ______substances. Ion-ion interactions are very strong, which is why ionic compounds have high melting & boiling points. B. Hydrogen Bonding: These are the dominant forces felt in ______molecules in which a H-F, H-O, or H-N bond is present somewhere in the molecule. This is a special type of dipole-dipole interaction. C. Dipole-Dipole: These are the dominant forces felt in ______molecules that don’t fit into the Hydrogen Bonding category. The partial negative charge on one molecule can interact with the partial positive charge on another molecule, causing a pull between the molecules. D. London Dispersion: These are the dominant forces felt in ______molecules. However, ALL molecules are capable of London Dispersion forces, but other types of IMF’s may dominate. The electron cloud on non-polar molecules can be temporarily distorted to cause an instantaneous dipole moment. These instantaneous dipole moments can interact, but the temporary interactions are very weak. Since it is easier to shift electrons on bigger molecules, the strength of the London Dispersion interactions increases with size (molar mass). This means that for all substances – if 2 compounds have the same dominant type of IMF, the substance with the higher molar mass will have the stronger IMFs. S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 17

However, it is very important to remember that INTRAMOLECULAR bonding (ionic & covalent bonds) is WAY STRONGER than INTERMOLECULAR forces.

LEGGETT PreAP 6-13 Summary (10:02) http://vimeo.com/54341123 http://youtu.be/NiNq9eNMTx4 Example 6-19. For each of the following: determine the dominant type of IMF between the molecules or compounds in the pure state. Then, put them in order of increasing melting/boiling points. RANK of MP/BP SUBSTANCE DOMINANT IMF (5=LOWEST)

CBr4 S T U D E N T N O T E S P r e - A P C h e m i s t r y U N I T 6 | Page 18 BOND COMPARISONS: Let’s Review! Non-Polar Covalent Ionic Bond Polar Covalent Bond Bond

General Formula

& Structure

Difference in

Electronegativity

Usually called a…

How can elements reach 8 electrons to be stable? (OCTET RULE)

How do we show their Lewis Dot Structure?

Basic Properties…