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The Relative Strength of Chemical Bonds Name _______________________________________ Coursework Stamp #3
Can we use macroscopic properties like boiling points and melting points to infer the strength of chemical bonds?
Why? Chemical bonds can be categorized into four main groups: covalent, ionic, intermolecular, and metallic. Understanding these four categories of bonds allows chemist to explain and more importantly predict the properties of substances and possible chemical reactions. In addition, chemistry deals predominantly with things that we cannot see directly. We have to infer the properties from evidence that we can see and measure. Today you will be practicing that skill. Complete these summaries AFTER you have finished this POGIL Summarize what you learned from Model 1: Summarize what you learned from Models 2‐5: Summarize what you learned from Model 6: Use the scale below to rank the relative strength of the bonds.
Strongest Weakest
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Model 1: Review of Bond Types So far you have learned that differences in electronegativity are used to determine some of the types of bonds that forms between two atoms. Below are diagrams of ionic and covalent compounds. Sodium Chloride Water
Type of Bond
Difference in electronegativity is large, medium,
or small
Electrons are shared evenly,
shared unevenly, or transferred
Types of Atoms that typically
form this kind of bond
Ionic
Polar Covalent
Non‐polar Covalent
1. Using as much detail as you can, describe what holds atoms together in an
ionic bond like NaCl. Use the diagram above to help. 2. Using as much detail as you can, describe what holds atoms together in a
covalent bond like H2O.
Na+ Cl‐ Na+ Cl‐ Na+ Cl‐
Na+ Cl‐ Na+ Cl‐ Na+ Cl‐
Na+ Cl‐ Na+ Cl‐ Na+ Cl‐
Na+ Cl‐ Na+ Cl‐ Na+ Cl‐
Ionic Bonds
Covalent Bonds
Intermolecular Forces(hydrogen bond)
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23. Explain what holds the atoms together in a metallic bond.
24. How would you expect the strength of metallic bonds to compare with the strength of covalent bonds? Explain your reasoning.
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Read This! When metal atoms forms bonds they shared electrons somewhat like covalent compounds. However, unlike non‐metals which have large electronegativities, metals have a relatively weak hold on their valence electrons. The result of this weak hold on electrons is that the valence electrons are not pulled into specific bonds between atoms. Rather the valence electrons are shared between all the atoms. The valence electrons are often described as a negative “sea of electrons” moving around the atoms.
21. Describe how metallic bonds are similar and different from covalent bonds? Similarities Differences
22. Use the electronegativities to explain why non‐metals form specific bonds
between atoms and metals form a “sea of electrons.”
+ +
+ +++++++
Electrons are shared
between two atoms, forming a specific bond
Valence electrons are
shared between all metal atoms
The positive metal atoms are attracted to the negative “sea of
electrons.”
Metallic Bonds Covalent Bonds
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3. What type of bond do the dotted lines between the water molecules represent? There is a general name and a specific name.
a. Describe what causes this force. Be sure to indicate which atom is slightly negative and which is slightly positive.
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Model 2: Melting and Boiling Points Name formula
Types of Bonds
Melting Point (˚C)
Boiling Point (˚C)
Sodium chloride NaCl Ionic 801 1,413
Copper(II) sulfide CuS Ionic 550 ‐‐‐
Magnesium oxide MgO Ionic 2,852 3,600
Calcium bromide CaBr2
Ionic 750 1,935
Iron(III) sulfate Fe2(SO4)3
Ionic & Covalent 480 ‐‐‐
Sodium phosphate Na3PO4
Ionic & Covalent 622 ‐‐‐
Water H2O
Very Polar Covalent 0 100
Polyvinyl alcohol (CHCOH)n
Very polar covalent 200 228
Methanol CH3OH
Polar Covalent ‐98 64
Methane CH4
Non‐polar Covalent ‐182 ‐161
Oxygen O2
Non‐polar Covalent ‐219 ‐182
‐‐‐ Indicates the compound decomposes into a different compound prior to boiling.
4. The melting point of magnesium oxide is _____ and the boiling point is
_____. Magnesium oxide is a fine white powder. Describe what you think you would see happen if you heated a small pile of magnesium oxide to a temperature of 2,853˚C?
5. There are three bond types listed in the table. Are there any patterns
between the types of bonds and the boiling and melting points? Describe as many patterns as you can see.
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18. Does the amount of charge or the distance between charges have a greater impact on the strength of the Coulombic Force? (Remember that Coulombic force is just a fancy name for the attraction and repulsion between charged particles.)
19. Should covalent or ionic bonds be stronger? Explain your reasoning in terms of the separation (distance between) of the charges.
Model 6: Metallic Bonding ‐ A Fourth Bond Type Read This! We have seen that two non‐metals tend to share electrons in a covalent bond because their electronegativities are so similar. When a metal and non‐metal bond, the large difference in electronegativities typically results in the transfer of electrons and the formation of an ionic bond. Metals tend to have similar electronegativities as other metals.
20. Do you think metallic bonds (bonds between metal atoms) will be more like ionic or covalent bonds? Explain your reasoning.
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Model 4: Explaining the Relative Strength of Ionic Bonds and Intermolecular Forces
14. Ionic bonds form from the attraction between a ( full / partial ) positive charge on the cation and ( full / partial ) negative charge on the anion. Intermolecular forces form ( full / partial ) charges on adjacent molecules. a. Which force should be stronger, the intermolecular force (hydrogen
bond) or an ionic bond? Explain your reasoning. Model 5: Explaining the Relative Strength of Ionic and Covalent Bonds Typical Ionic Bond Typical Covalent Bond 15. How far apart are the ions in a sodium chloride crystal? 16. How far are the hydrogen nuclei from the middle of the shared electrons? 17. Ionic bonds and covalent bonds are really just attractions between positive
and negative charges. How does the distance between the charges of a typical ionic bond compare to the distances in a covalent bond?
Na+ Cl‐
Na+ Cl‐
282pm
+ +
+ +++++++
Na+ Cl‐
Na+ Cl‐
Ionic Bond
+ +
+ +++++++
Intermolecular force(hydrogen bond)
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6. Rank the three types of bonds from lowest to highest melting points and then lowest to highest boiling points. Bonds with lowest Bonds with highest melting points ________ _______ ________ melting points
Bonds with lowest Bonds with highest
boiling points ________ _______ ________ boiling points
a. Explain your rankings in the space below.
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Model 3: Understanding Melting The diagram below show what happens to particles of sodium chloride, an ionic solid as it melts. The diagram below show what happens to H2O molecules as they melt from ice to water.
7. In order for an ionic compound to melt, the ions must overcome (break) the
______________ bonds holding them in place.
8. In order for a covalent compound to melt, the ions must overcome the
___________________ forces holding them in place.
Solid – Ions are held in place by in a rigid crystal structure by ionic
bonds.
Liquid – Ions vibrate violently enough to break the ionic bonds
and start moving around as individual ions.
Melting
Cl‐
Na+
Cl‐ Na+
Cl‐ Na+
Cl‐ Na+
Cl‐ Na+
Cl‐ Na+
Na+ Cl‐ Na+
Na+ Cl‐ Cl‐
Na+ Cl‐ Na+
Na+ Cl‐ Cl‐
+ +
+ + + + + +
+ +
+ +
+ + + + + +
+ + + +
+ + + + + +
+ +
+ +
+ + + + + +
+ + + +
+ +++++++
Solid – In ice the H2O molecules are moving so slowly, the
intermolecular forces (hydrogen bonds) are strong enough to hold
the H2O molecules in place, forming a rigid crystal structure.
Liquid – In water the H2O molecules are moving so fast, the intermolecular
forces (hydrogen bonds) are NOT strong enough to hold the H2O molecules in place, which means they can move around as individual H2O molecules.
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9. Does melting or boiling break the covalent bonds between the oxygen and hydrogen in water? Explain your reasoning using information from the model and your experience with steam.
10. Look back at Model 2. Which type of compound was easier to melt, ionic or
covalent? 11. What does this tell you about the relative strength of ionic bonds and
intermolecular forces? Explain your reasoning as clearly as you can. 12. Using the melting point data from Model 1 can you tell whether covalent
bonds are stronger than ionic bonds? Explain your reasoning. Read This! A glass of water is made of trillions and trillions of individual water molecules that are held together by intermolecular forces called hydrogen bonds. Like water, a diamond is a covalent compound but instead of being made of a bunch of individual molecules, every carbon atom in the diamond is covalently bound to every other carbon atom. The entire diamond is one giant molecule of carbon atoms. This means that in order to melt a diamond you have to overcome covalent bonds rather than just intermolecular forces. The melting point of diamond is 3,550˚C.
13. What does the melting point of diamond tell you about the relative strength of covalent and ionic bonds? Use melting points from the Model 1 to help you answer this question. Explain your reasoning.
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