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Unit 5: Bonding

Unit 5: Bonding. An example of an iconic bond would be James. Castagno Chemistry Challenge V. Rules: 1) You are working as a CLASS ! 2) Each question is multiple choice. 3) There are 7 questions so the highest score gets the points. POINTS

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Unit 5: Bonding

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  1. Unit 5: Bonding An example of an iconic bond would be James.

  2. Castagno Chemistry Challenge V • Rules: • 1) You are working as a CLASS! • 2) Each question is multiple choice. • 3) There are 7 questions so the highest score gets the points. • POINTS • Class:1st – 2pts, 2nd – 1pt, 3rd – 0pts, 4th – 0pts • Questions?

  3. Castagno Chemistry Challenge V • There are 4 types of bonds and 3 types of forces. A substance will be shown on the board and you must choose which bond or force is present. • Ready?

  4. Castagno Chemistry Challenge V Q1 • A) Ionic Bond • B) Polar Covalent Bond • C) Metallic Bond • D) Nonpolar Covalent Bond

  5. Castagno Chemistry Challenge V Q2 • A) Dispersion Force • B) Dipole – Dipole Force • C) Hydrogen Bond

  6. Castagno Chemistry Challenge V Q3 • A) Ionic Bond • B) Polar Covalent Bond • C) Metallic Bond • D) Nonpolar Covalent Bond

  7. Castagno Chemistry Challenge V Q4 • A) Dispersion Force • B) Dipole – Dipole Force • C) Hydrogen Bond

  8. Castagno Chemistry Challenge V Q5 • A) Ionic Bond • B) Polar Covalent Bond • C) Metallic Bond • D) Nonpolar Covalent Bond

  9. Castagno Chemistry Challenge V Q6 • A) Dispersion Force • B) Dipole – Dipole Force • C) Hydrogen Bond

  10. Castagno Chemistry Challenge V Q7 • A) Ionic Bond • B) Polar Covalent Bond • C) Metallic Bond • D) Nonpolar Covalent Bond

  11. Castagno Chemistry Challenge V • Answers on next slide

  12. Castagno Chemistry Challenge VAnswers • 1) Metallic Bond • 2) Dipole-Dipole Force • 3) Nonpolar Covalent Bond • 4) Hydrogen Bond • 5) Polar Covalent Bond • 6) Dispersion Force • 7) Ionic Bond

  13. Unit 5 Objectives • The following is what you can expect to learn and understand: • Define ionic bond, covalent bond, Lewis dot structure, lone pair, octet rule, electronegativity, polar covalent bond, non-polar covalent bond, VSEPR theory, intermolecular forces (dispersion, dipole, hydrogen) • How to draw a Lewis dot structure. • How to illustrate an ionic bond and covalent bond. • The difference between a lone pair of electrons and a bonding pair of electrons. • The octet rule and the elements that do not obey the octet rule. • The importance of electronegativity. • The use of VSEPR theory to determine molecular shape. • Determine the intermolecular force(s).

  14. Unit 5 Essential Questions • The questions to ask along the way: • What is the comparison between ionic and covalent bonds? • How are ionic and covalent bonds illustrated using Lewis dot structures? • How does a lone pair of electrons compare to a bonding pair? • What is the octet rule? • What is the relationship between electronegativity difference and the type of chemical bond formed? • How does a polar covalent bond compare to a pure covalent bond? • What is VSEPR theory and how is it used? • What causes London dispersion, dipole-dipole, and Hydrogen bond forces?

  15. Your Take • What would you like to learn? • Why can we see our breath in the winter? (P2) • What is the rarest compound? • “Clean” energy compounds • Are things bonded together considered “touching?” (p4) • Why is gold soft compared to other metals? • What makes silver such a good conductor? • How do bonds affect compound’s properties? • Polarity of water (P6) • Smaller compounds combining to form larger ones • Ions and their properties • The most common compound • What element can make the strongest bond? (P7) • What is my favorite chemical bond? • What is the most common chemical bond? • What is the most dangerous chemical bond and how is it made? • Why do people’s lips get chapped in the winter?

  16. Atoms in Nature • Every atom consists of protons and neutrons surrounded by a “sea” or “cloud” of electrons. • With the exception of the Noble Gases, atoms join in a (nearly) infinite number of combinations. • Combinations of atoms are formed by bonds, chemical bonds.

  17. Atoms Behaving Badly* • So Noble Gases don’t react (or form bonds) • But they actually can. • Xenon is pretty famous for forming compounds (first one created was XeF4) but Helium, Neon, and Argon are extreme loners.

  18. Atoms Behaving Badly Still* • But even a Helium ion (more later) known as HeH+ has been discovered (in a lab, not nature…yet) • This means a helium atom is bonded with a proton from hydrogen. • The electron left. Possibly because of jealousy.

  19. Atoms Behaving Badlier Still* • Argon, though, is part of a famous compound known as HArF. • If the temperature rises above -256 C, it decomposes into HF and Ar. • Neon has yet to get in on the act. • But what about NATURALLY occurring un-Noble-like behavior?

  20. 12/12/13 – Argon hydride* • ArH+ • Discovered within the cold, dusty regions of the Crab Nebula. • Studied infrared radiation from the nebula • Infrared radiation causes molecules to become excited, or spin or vibrate. • This energy is specific for every molecule. • A comparison to behavior of known molecules showed that the combination had to be Argon and Hydrogen. • Interestingly, the scientists could even determine the ISOTOPE of argon responsible: Argon – 36. • What’s weirder is argon – 40 is the most common on earth (due to the decay of potassium)

  21. Chemical Bonds • Why would atoms prefer to bond instead of remain isolated in element form? • The atoms are more stable! • The mutual electrical attraction between nuclei and valence electrons of different atoms that binds the atoms together. • Protons and neutrons are not shared nor donated in chemical bonds. • Electrons are far from the nucleus and can easily move around and be shared or transferred between nuclei.

  22. Back to the Future: Electronegativity • “a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound” • Think of it as how much an atom will share. • In the formation of chemical bonds, difference in electronegativity will tell you the type of bond.

  23. Ions • The exchange of valence electrons make atoms more stable. • From the previous unit, we discussed ionic radii – the size of ions. • Some elements preferred to lose electrons, others gain. • Cations – positive ions (atom that lost electrons) • Anions – negative ions (atom that gained electrons)

  24. Ionic Bonds • A bond that results from the electrical attraction between cations and anions. • Think of the north and south ends of a magnet attracting.

  25. Electronegativity – bonds • A measure of the ability of an atom to attract electrons. • Recall that electronegativities range from 4.0 (F) to 0.7 (Fr). • We can use these values to determine if a bond between two atoms is ionic. • Ionic Bonds • When the difference in electronegativity values is greater than 1.7*.

  26. Ionic Bonds? • Determine if the bonds formed are ionic or not. • H and S • No (0.4) • Cs and S • Yes (1.8) • Cl and S • No (0.5) • H and C • No (0.4) • K and Br • Yes (2.0) • K and O • Yes (2.7) • Cl and Cl • No (0.0) • N and O • No (0.5) • C and O • No (1.0) • Li and F • Yes (3.0) • B and F • No (2.0) ?!?!?!?!?!?!?

  27. Less than 1.7 • Not every compound is formed of cations and anions with ionic bonds. • Covalent bonds • When the difference in electronegativity is less than 1.7 • This results in a bond where electrons are shared, but not always equally.

  28. Polar Covalent Bonds • Ions do not share electrons at all. • Covalent bonds have shared electrons. • There are, however, plenty of bullies! • Bonded atoms that have an unequal attraction for the shared electrons. • If the difference in electronegativity ranges from 1.7 – 0.3.

  29. Polar Covalent Bonds II • What is the implication of this statement? • The unequal sharing means the negative electrons surround one atom more than another. • One “side” must be more negative than the other! • Which means the other side is more positive than the other • Like a battery, these compounds have a positive and negative “area.”

  30. Polar Covalent Bonds? • Determine if the bonds formed are polar covalent. • H and S • Yes (0.4) • Cs and S • No (1.8) • Cl and S • Yes (0.5) • H and C • Yes (0.4) • K and Br • No (2.0) • S and O • Yes (1.0) • Cl and Cl • No (0.0) • N and O • Yes (0.5) • C and O • Yes (1.0) • Li and F • No (3.0) • P and O • Yes (0.9)

  31. Non-Polar Covalent Bonds • Covalent bonds have shared electrons. • Some however, share evenly, or close to it. • Bonded atoms that share bonding electrons equally, resulting in a balanced distribution of charge. • The electronegativity difference is 0.3 – 0.0. • Hydrogen gas (H2) • H – H, 2.1 – 2.1 = 0.0 • Bonded atoms of the same (non-metal) element have non-polar bonds.

  32. Formation of a covalent bond

  33. Formation of H – H bond • A) Two hydrogen atoms are approaching each other. • B) The atoms are close enough that the electron in one is attracted to the nuclei of the other. At the same time, the electron clouds and nuclei repel. • C) The repulsion equals the attraction, potential energy is at a minimum, and a stable hydrogen molecule forms.

  34. Characteristics of Covalent Bonds • Covalent bonds involve a sharing of electrons. • This means that electron “clouds” overlap.

  35. Characteristics II • The average distance between nuclei is the bond length. Bond Length (75pm)

  36. Characteristics III • The separate H atoms have more energy than the molecule, exactly 436 kJ/mol more. • To break the bond of the molecule, you have to add 436 kJ/mol of energy. • The energy to break a bond to form isolated atoms is bond energy.

  37. Bond Energy • The table below represents bond energy for H bonded to a halogen. Notice anything? • The shorter the bond length, the greater bond energy!

  38. Energy of Ionic Bonds • Since there is no sharing of electrons, the bonds are stronger than covalent bonds. • This also means ionic compounds have higher melting and boiling points than molecular compounds. • Lattice energy* – energy released when bonds form

  39. Metallic Bonds • A network of metal atoms surrounded by a “sea” of electrons. • Because of this network, metals are malleable (sheets) and ductile (wire).

  40. Electron-dot notation • Indication of the valence electrons that surround an atom

  41. Lewis Structures • Similar to a chemical formula, indicating both valence electrons and bonded electrons. • The octet rule* must be obeyed where an atom must be surrounded 2 or 8 electrons • * Exceptions: B (6), P (10), S (12) • Fluorine (F2) • With bond

  42. Lewis Structures Example 1 • Tips • Carbon is always a central atom. • Hydrogen is never a central atom. • CH3I (hint: single bonds only)

  43. Lewis Structures Example 2 • Tips • Carbon is always a central atom. • Hydrogen is never a central atom. • CH2O (formaldehyde)

  44. Multiple-Covalent Bonds • In some cases (like C and O in formaldehyde), it is necessary for a pair of atoms to share more than 2 electrons in a bond to form an octet. • Double bonds – share 4 electrons • Triple bonds – share 6 electrons

  45. Lewis Structures Practice • ethene (C2H4) • Acetylene (C2H2) • Ammonia (NH3)

  46. Lewis Structures Practice II • Hydrogen sulfide (H2S) • silane(SiH4) • Phosphorus trifluoride (PF3)

  47. Lewis Structures Practice III • Nitrogen gas (N2) • Carbon dioxide (CO2) • Hydrogen cyanide (HCN)

  48. Lewis Structures Practice IV • Bromomethane (CH3Br) • Silicon tetrachloride (SiCl4) • Oxygen difluoride (OF2)

  49. Lewis Structures Practice V • CFC (CCl2F2) • Beryllium chloride* (BeCl2) • Ozone** (O3)

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