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Chemical Compounds and Bonding

Chemical Compounds and Bonding. Classifying Chemical Compounds. Bonding Basics. Review of Gr.10 Science. Section A: Complete the chart using a periodic table to help you. CATION. ANION. “Cat-Eye-On”. “An-Eye-On”. Answer these questions:

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Chemical Compounds and Bonding

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  1. Chemical Compounds and Bonding Classifying Chemical Compounds

  2. Bonding Basics Review of Gr.10 Science

  3. Section A: Complete the chart using a periodic table to help you.

  4. CATION ANION “Cat-Eye-On” “An-Eye-On” Answer these questions: An atom that gains one or more electrons will have a ____________________ charge. An atom that loses one or more electrons will have a ____________________ charge. An atom that gains or loses one or more electrons is called an ____________. A positive ion is called a ______________ and a negative ion is called an _______________. NEGATIVE POSITIVE ION

  5. What is an ionic bond? Atoms will transfer one or more ________________ to another to form the bond. Each atom is left with a ________________ outer shell. An ionic bond forms between a ___________ ion with a positive charge and a ________________ ion with a negative charge. Example B1: Sodium + Chlorine Example B2: Magnesium + Iodine ELECTRONS COMPLETE METAL NONMETAL

  6. Example B3: Potassium + Iodine Example B4: Sodium + Oxygen Example B5: Calcium + Chlorine Example B6: Aluminum + Chlorine

  7. What is a covalent bond? Atoms ___________ one or more electrons with each other to form the bond. Each atom is left with a ________________ outer shell. A covalent bond forms between two _________________. Example C1: Hydrogen + Hydrogen Example C2: 2 Hydrogen + Oxygen SHARE COMPLETE NONMETAL

  8. Example C3: Chlorine + Chlorine Example C4: Oxygen + Oxygen Example C5: Carbon + 2 Oxygen Example C6: Carbon + 4 Hydrogen

  9. Elements • There are about 90 naturally occurring elements. • Most are not found as pure elements. • The majority of elements are found combined with other elements to form compounds. • Gold, silver and platinum are rare examples of metals found in elemental form (precious metals).

  10. Classification Systems • Since the 90 elements can form thousands of different compounds, classification systems have been developed. • The classification of compounds is based on their properties to help our understanding of compounds. • Example: melting point, boiling point, hardness, etc.

  11. Properties of Ionic and Covalent Compounds

  12. What is a crystalline solid? What does it look like? • A crystal or crystalline solid is a solid material whose constituent atoms, molecules, or ions are arranged in an orderly repeating pattern extending in all three spatial dimensions.

  13. Thought Lab: Ionic or Covalent? • With the person sitting next to you, take a look at the following chart • See if you can figure out what compound belongs with what sample number. • The compounds used were: • Ethanol • Carbon tetrachloride • Glucose • Table salt (sodium chloride) • Water • Potassium permanganate

  14. Experimental Results ionic covalent ionic covalent covalent covalent

  15. What is Bonding?

  16. Chemical Bonds • forces that attract atoms to each other to form compounds • involves the interactions of valence electrons between atoms • usually the bond forms a compound that is more stable than the atoms individually.

  17. Ionic and Covalent Bonding: The Octet Rule

  18. The Octet Rule • atoms bond in order to achieve the same number of electrons as the nearest noble gas in the periodic table. • bonded atoms are said to be isoelectronic to the nearest noble gas. (All noble gases have 8 valence electrons except He.)

  19. Ionic Bonds • a chemical bond formed when one atom loses valence electron(s) an another atom gains electron(s). • occurs between metal and nonmetal atoms

  20. e– 1) 2) Na Na+ Cl Cl– Cl– Na+ Ionic Bonding • Ionic bonding involves 3 steps • 1) loss of an electron(s) by one element2) gain of electron(s) by a second element3) attraction between positive and negative 3)

  21. Ionic Bonding

  22. Ionic Bonding Examples Using Lewis Structures In some bonds, multiple electrons can be exchanged:

  23. Examples Continued Some ionic compounds involve bonds between more than two elements.

  24. Structure of Ionic Compounds • an ion of one charge is totally surrounded by ions of a opposite charge in the solid state. • the charged ions that make the ionic substance also explain why these compounds are soluble in polar solvents such as water. • also explains conductivity of ionic solutions

  25. Structure of Ionic Compounds- Crystal Lattice • the ions in these compounds are not free to move around in the solid state • this explains the high melting point and non-conductivity of these compounds • the structured lattice also explains why the crystals are brittle

  26. Covalent Bonds • form when two atoms share valence electrons • occurs between two non metal atoms

  27. Covalent Bonds • consist of a shared pair of electrons between two nonmetals • most often each atom of a bond contributes one electron to the bond • atoms share electrons so that they may achieve a stable octet in their valence shell

  28. Pure Covalent Bond • two atoms of the same element share their electrons equally • these are called diatomic elements. • ex. H 2 , N 2 , O 2 , F 2 , Cl 2 • when atoms of two different elements share electrons and the ΔEN is less than 0.5.

  29. Multiple Covalent Bond • some atoms will in order to complete octets will share more than one electron and form double or triple bonds.

  30. Cl Cl C Cl Cl H H H Cl N H Covalent Bonding Examples In covalent bonding one atom does not take electrons from another atom, but rather the two atoms share electrons. HCl CCl4 NH3

  31. Properties of Covalent Compounds • wide range of properties • ex. some dissolve in water and some do not • the atoms in the bond are held together very tightly and do not break apart during change of state or during the dissolving process, and thus these compounds are called molecular compounds.

  32. Intra- and Inter- molecular Forces • Intramolecular Forces: forces that bond an atom together in a molecule ex. covalent bonds • Intermolecular Forces: forces that keep the molecules close to each other in the solid and liquid state these forces are weaker than intramolecular forces and are responsible for the low mp and bp of covalent compounds.

  33. Polar Covalent Bonds

  34. Electronegativity (EN) • a measure of an atoms ability to attract electrons in a chemical bond. • a property of an atom involved in a bond

  35. How can we use electronegativity to predict bond type? When two atoms form a bond the difference in electronegativity (ΔEN) can help to determine the bond type. | | | 0 0.5 1.7 3.3

  36. Examples • HCl 3.16- 2.2= 0.96 Polar Covalent • CrO 3.44- 1.66= 1.78 Ionic • Br2 2.96-2.96= 0 Covalent

  37. Polar Covalent Bonds and Polar Molecules • The variety of differences in electronegativities of covalent bonds accounts for the variety of properties observed in covalent compounds. • ex. -some dissolve in water, others not • some are solids while others are liquids or gases at room temperature

  38. Polar Covalent Bonds- The “In Between Bonds” • electrons shared between two atoms unequally • ΔEN is between 0.5 and 1.7 • the atom with the higher electronegativity is partially negative ( δ- ) • the atom with the lower electronegativity is partially positive ( δ+ )

  39. What does this look like?

  40. Molecular Shape • To understand why molecules have different shapes, consider how electron arrangement affects shape. • Electron pairs that are not involved in bonding are called lone pairs. • Electron pairs that are involved in bonding are called bonding pairs. • Electron pairs are arranged around molecules so that they are a maximum distance from each other. • The shape that allows four electron pairs to be a maximum distance from each other around an atom is a tetrahedron.

  41. Dipoles • The polarity of a bond in a molecule can be shown using a vector called a dipole.  • The magnitude of the dipole is determined by the ΔEN • The direction is determined by EN of each atom in the bond. • ex. Water

  42. Dipoles • A polar molecule is identified by the vector addition of all dipoles. If the resultant vector, call the dipole moment, is nonzero, then the molecule is polar. This mean one end of the molecule carries a partial positive charge and the other a partial negative charge.

  43. Example: Carbon Dioxide • Through vector addition we see that the resultant vector (dipole moment) equals zero, therefore carbon dioxide is a nonpolar molecule.

  44. Example: Hydrogen Chloride

  45. Example: Ammonia

  46. Example: Carbon Tetrafluoride

  47. In General • symmetrical molecules with no lone pairs of electrons on the central atom are nonpolar • molecules with lone pairs of electrons and polar bonds are polar.

  48. Properties

  49. + – H Cl + – + – + – + – Determining Molecular Polarity • Depends on: • dipole moments • molecular shape

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