3.1 Classifying Compounds

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CHAPTER 3 BONDING

mixing / reacting
element + element → mixture / element + element → compound
Each element in the mixture retains ______
properties. / New compound has ______properties unlike the properties of either element. (Not the “average” of the properties)
Properties of the compound largely depend on the ______in the compound.
A knowledge of Bonding helps us to predict and explain the properties we observe in compounds.

3.1 CLASSIFYING COMPOUNDS

IONIC COMPOUNDS (SALTS) / MOLECULAR (COVALENT) COMPOUNDS

CHEMICAL BONDS IN GENERAL:

Ø  forces that attract ______to each other in compounds

Ø  involve interaction of ______

Ø  usually create a compound that is ______than either element by itself

IONIC BONDS / COVALENT BONDS
Valence electrons are ______from one atom to another / Valence electrons are ______
Occurs between
metals (______IE, ______EA, ______electrons) and
nonmetals (______IE, ______EA, ______electrons) / Occurs between ______

Difference between ionic and covalent compounds based on properties is NOT always clear cut.

eg. HCl: gas, solution in water conducts

THOUGHTLAB: ionic or covalent?

Ethanol (C2H5OH), cyclohexane (C6H12), glucose (C6H12O6), table salt (NaCl), water, potassium permanganate (KMnO4)

sample / compound / Dissolves in H2O? / Conductivity as a liquid or solution / mp (°C) / Appearance / Ionic or Covalent?
1.  / y / high / 801 / clear, white crystalline solid
2.  / y / low / 0 / clear colourless liquid
3.  / y / high / 240 / purple crystalline solid
4.  / y / low / 146 / white powder
5.  / n / low / -23 / clear colourless liquid
6.  / y / low / -114 / clear colourless liquid

Which is the most definitive property for classifying ionic and covalent compounds?

ELECTRONEGATIVITY:

Ø  indicate an atoms ability to ______within a chemical bond

Ø  a rating scale (no units) that takes into account atomic properties such as ______, ______and ______.

Ø  The element ______has the highest electronegativity (e.n. = )

Ø  De.n. = difference in electronegativity

CLASSIFYING BONDS BY THE De.n:

BOND TYPE
De.n.
DESCRIPTION
EXAMPLE:

Ø  the scale above is like a “spectrum” – any boundaries between bond types are somewhat arbitrary and artificial

(bonds are not “black” or “white”, but many differerent shades of “gray”)

Ø  De.n is just one criteria used in assessing bond type and is not always 100% reliable. Properties of the compound should also be considered

Ø  % ionic character can be determined from ______using the chart provided on your periodic table and tells how close to the ionic end of the spectrum a bond is.

eg.

Compound / De.n. / Bond type / % ionic character
BaBr2
H2S

Bond Classification Practice:

Bond / De.n. / Bond type / % ionic character
a) O-H
b) C-H
c) Mg-Cl
d) B-F
e) Cr-O
f) C-N
g) Na-I
h) Na-Br

Ionic Bonding

Recall that metals have ______(lose electrons easily)

Nonmetals have high ______(readily accept electrons)

After electron transfer and ion formation:

(Video-Answer During the video)

1. Properties of Salts (Ionic Compounds): (that differs from metals)

2. Salts are made of small crystals that have a regular shape which suggests that they are made of a regular ______

Consider table salt: sodium chloride

3. Na  ______(Which is bigger, the sodium atom or the ion?)

Sodium atom ______to become a ______(positive ion)

4. Cl  ______(Which is bigger, the chlorine atom or the ion?)

Chlorine ______to become an ______(negative ion)

5. What are the electrostatic forces that push and pull the ions into the lattice arrangement: ______attract ______, ______repel ______repel ______

6. Electrostatic forces of attraction and repulsion push and pull the ions into a symmetrical 3-dimensional ______which is held together by the mutual attraction of ______charged ions.

7. There is no limit to the ______of the NaCl lattice.

8. The 3-d ion packing arrangement is dependent on:

a)

b)

9. In NaCl, each Na+ is surrounded by ______chlorine anions

(Likewise each Cl- is surrounded by ______sodium cations.

10. If an anion is large (like I-), ______limits the number of cations that can be packed around each anion.

(To be filled in after the video)

11. When a strong enough force is applied to an ionic crystals what happens?

/ a) strong force applied to crystal lattice
b)
c)

12. When an ionic crystal is heated strongly to a high temperature what happens?

/
b) As salt is heated: /
a) ions can only vibrate in their places in the crystal lattice / c)

13. How do molten salts conduct electricity when solid salts cannot?

Practice:

Show in three steps how ionic bonds form between Strontium and Nitrogen atoms. Show how energy is involved in each step.

IONIC BONDING PRACTICE: Show the three steps for

a) Mg and Br b) Ca and N c) Li and S d) Ba and O e) Cs and P

COVALENT BONDING- Simple Lewis Theory

·  Occurs between ______atoms

·  The Valence shell consists of ______; each orbital holds ______.

Lewis dot diagram of a Cl atom:

Lewis dot diagram of a Cl2 molecule:

·  The ______pair of electrons spends time around ______Cl nuclei

·  The ______of a pair of electrons for ______is the “glue” that holds the atoms together in a covalent bond.

Drawing diagrams of Molecules:

Eg. CF4

Lewis Dot diagram Structure

Guidelines for Drawing:

1.  Determine ______of each element by drawing the Lewis dot diagram of each element (This is the number of ______, bonding electrons)

2.  Start with the atom(s) with the ______and build on them.

3.  Add atoms that can form ______last

4.  When finished check:

that you have the correct number of each atom

that each atom has fulfilled its ______.

Example: eg. C2Cl6 S

Note: These diagrams do NOT show the actual 3-d geometrical shape of the molecules. They only show how the atoms are ______to each other.

Suggestion: Try drawing the structure first. Then work backwards to the Lewis Dot diagram; substitute pairs of dots for each bond. Then add all the lone pairs.

Drawing Molecules Exercise I:

a)  H2O2 b) CBr2Cl2 c) C2H6 d) Cl2O e) NH3 f) CH5N g) P2Cl4

b)  h) CH4O i) C2H6O (draw 2 completely different molecules) j) C2H4(NH2)2

Multiple Covalent Bonds:

Consider the O2 molecule:

Lewis dot diagram of O atom: Lewis Diagram of O2 molecule: Structure of O2 molecule:

Atoms can share ______in a ______covalent bond.

Consider the N2 molecule:

Lewis dot diagram of N atom: Lewis Diagram of N2 molecule: Structure of N2 molecule:

Atoms can share ______in a ______covalent bond.

Drawing Structures and Lewis Dot diagrams:

Some molecules may require ______or ______bonds. Do not add these multiple bonds until the end of the drawing process, if necessary, to ensure that each atom fulfills its ______.

Examples:

a) CNH3 b) Si2H2

Drawing Molecules Exercise II:

a)C2H4 b) Si2Cl2 c) BCl3 d) N2H2 e) HNO f) CH2O g) HCN h) C3 H4 i) HSCN (draw 2 different molecules for this one) j) C2O4H2 k) CS2

Polar Covalent Bonding

Consider HCl:
diagram
Electronegativity
Electron attracting ability
Charge

“δ” (a lower case Greek letter delta means “partial” )

·  A bond that has a positive and negative end is described as ______(an adjective)

·  This charge separation is called a ______(a noun)

·  The greater the Δen, the greater the charge separation, the more ______the bond is and the bigger the ______.

·  Dipoles are symbolized with arrows; the arrowhead always points to the ______end of the dipole.

·  The size of the arrow represents the size of the ______as indicated by the ______.

Practice with Bond dipoles: Draw the structure of each molecule. Use arrows to indicate the direction of any bond dipoles. Use the size of the arrow to indicate the relative size of the bond dipoles.

a)  CHF3 b) C2Cl3F

Polar and Nonpolar Molecules:

·  Recall that a Δen results in a ______bond or a bond ______; a bond with a positive and negative end.

·  A ______is a molecule with a positive and negative end; where the electrons are ______more in one direction.

·  To determine whether a molecule is polar we must consider both:

The ______(magnitude) of any bond dipoles and

The ______of the bond dipoles

·  Deciding polarity is like vector addition, but we will add up the bond dipoles visually rather than mathematically.

·  Sometimes we can tell the polarity of a molecule from the ______; other times we must know the true ______to correctly determine polarity.

Predicting Molecular Polarity – examples: / BCl3 / NH3
structure
Actual 3-d geometry (shape) of molecule with bond dipoles
Predicted Polarity

Note: At the grade 11 level you are NOT expected to be able to predict the 3-d shapes of molecules. You will learn VSEPR in the grade 12 course. You will either be given the 3-d shape (if needed) by the teacher or use molecular model kits to obtain it.

To Think about:

a)  Can a molecule with polar bonds be a non-polar molecule? Explain.

b)  Can a molecule with nonpolar bonds be a polar molecule? Explain.

Molecular Model Kits Lab Exercise

In this exercise, you will compare the Lewis structures of molecules with their models, using molecular model kits.

Instructions:

1.  Look at the chemical formula and draw the structural formula or Lewis structure of the molecule you are supposed to build.
2.  Construct the molecule using the model kit
a.  Use short connectors for single bonds and springs for double or triple bonds
b.  All holes must be filled with bonds
3.  Draw the 3-D shape of the molecule
4.  Add arrows (à) to the bonds to indicate the direction and
relative size of any bond dipoles. Pay attention to the
direction and the size of the arrows.
Look at the molecule you have constructed. Consider the following:
Add up the effects of all the bond dipoles.
Consider both the size and the direction of any bond dipoles to see if there is a molecular dipole.
5.  Decide if the molecule is polar or non-polar. (fill in last column of chart)
6.  If the molecule is polar, indicate the direction of the molecular dipole ( ) beside or on the 3-d structure.
7.  Show the δ+ and δ- to represent any molecular dipole beside the 3-d structure of the molecule if applicable. /

Please note the following:

C/Si = black O/S = red F, Cl, Br, I = orange, purple or green H = yellow N or P = sky blue

Instructions:

Compound / Structural formula or Lewis Diagram / 3-d Structure of Molecule / Polarity of Molecule
H2O
CCl4
CH4O
Compound / Structural formula or Lewis Diagram / 3-d Structure of Molecule / Polarity of Molecule
CH2O
CO2
HCN
CH3Cl
C2F2H2
C2F2H2
(change position of H’s and F’s)
C2F2H2
(change position of H’s and F’s)

INTERMOLECULAR BONDS

Intermolecular Forces or Bonds:

·  bonds ______molecules that hold a ______or ______together.

·  usually very ______compared to intramolecular bonds. (due to increased distance)

·  ______melting points and boiling points mean it requires a lot of ______to break the intermolecular bonds.

·  Often referred to loosely as ______forces.

Types of Intermolecular bonds:

1. London Dispersion forces

·  exist between ______molecules, both polar and nonpolar.

·  Cause: Random ______in electrons can produce temporary ______.

This temporary ______can induce dipoles in ______molecules. ______charged sides attract each other. These dipoles are ______and ______but the collective effect is significant.

·  example:

halogen / number of electrons / boiling point (K)
F2
Cl2
Br2
I2

·  Strength: ______electron clouds are more easily distorted leading to more temporary

______. \ The greater the number of ______in an atom or molecule,

the ______the London dispersion forces and the ______the m.p. and b.p.

2. Dipole-Dipole Forces

·  present only in ______substances.

·  Cause: ______charged ends of ______dipoles attract each other.

example: Account for the bp:

C2H6 bp = -87 °C

CH3Cl bp = -24 °C

Hydrogen Bonding - a special type of dipole-dipole force:

·  only occurs when H is bonded to a very ______atom.

( _____, ___, ____)

·  O-H, N-H and F-H bonds are extremely ______.

·  Cause: The hydrogen atom has an electron ______and therefore has a slight

______charge. Recall that Hydrogen is unique because it has no ______electrons. When H is deprived of its ______it becomes an

______nucleus.

\the hydrogen nucleus feels a strong attraction for the ______on nearby atoms.

example: Account for the bp below:

C3Cl8 b.p. = 543 K

C2H4(OH)2 b.p. = 471 K

C4H9OH b.p. = 391 K

INTERMOLECULAR BONDING EXERCISE

1. Rank from lowest to highest boiling point: CCl4, CI4, CBr4 , CF4. explain why.

2. Group VI elements include O, S, Se and Te. Hydrogen can form V-shaped compounds with all of these elements. The graph below shows you their boiling points.

a) Rank the four compounds according to their polarity (most polar to least polar)

b) Why does the b.p. increase from H2S, to H2Se, to H2Te?

c) Why is the b.p. of H2O such a surprisingly high one?

3. Indicate the type of intermolecular bonding present in each of the following substances:

Compound / 3-d structure / London Dispersion / Dipole-Dipole Forces / Hydrogen bonding
a) I2 / n/a
b) CH3OH /
c) CH3Cl /
d) CH2F2 /
e) HBr / n/a
f) CO2 /
g) CCl4 /
h) NH3 /

4. Consider the following boiling points: Kr , b.p. = -152°C HBr, b.p. = - 67°C