1 / 44

Unit 3

Unit 3. Chapters 2 (p 64-66), 8, 9.1-9.3. Acids. H 2 SO 4. H 3 PO 4. HCl. Hydrogen-containing compounds produce hydrogen ions (H + ) when dissolved in water. They are molecular compounds that ionize in water. “aq” is short for aqueous which means water. HCl(g)  H + (aq) + Cl - (aq)

louis
Télécharger la présentation

Unit 3

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Unit 3 Chapters 2 (p 64-66), 8, 9.1-9.3

  2. Acids H2SO4 H3PO4 HCl • Hydrogen-containing compounds • produce hydrogen ions (H+) when dissolved in water. They are molecular compounds that ionize in water. “aq” is short for aqueous which means water. • HCl(g)  H+(aq) + Cl-(aq) • or, more correctly, the hydronium ion (H3O+) • HCl(g) + H2O H3O+(aq) + Cl-(aq) H2O In general, acids are hydrogen containing compounds with a formula of HX, H2X, H3X, etc. where X can be an anion from Group 6A or 7A or X can be an oxyanion.

  3. Acids hydrochloric acid HCl hydrobromic acid HBr hydroiodic acid HI hypochlorous acid HClO chlorous acid HClO2 chloric acid HClO3 perchloric acid HClO4 nitric acid HNO3 sulfuric acid H2SO4 phosphoric acid H3PO4

  4. Acid Nomenclature • If the anion in the acid ends in -ide, change the ending to -icacid and add the prefix hydro- . • HCl: hydrochloric acid • HBr: hydrobromic acid • HI: hydroiodic acid

  5. Acid Nomenclature • If the anion in the acid ends in -ate, change the ending to -icacid. • HClO3: chloric acid • HClO4: perchloric acid • H2SO4: sulfuric acid • H3PO4: phosphoric acid • HNO3: nitric acid

  6. Acid Nomenclature • If the anion in the acid ends in -ite, change the ending to -ousacid. • HClO: hypochlorous acid • HClO2: chlorous acid • H2SO3: sulfurous acid

  7. Names of Common Chemicals Containing Ions

  8. Names of Common Chemicals

  9. Chemical Bonding and Molecular Geometry Molecules (naming and writing formulas) Lewis Symbols and the Octet Rule Ionic Bonding Covalent Bonding Molecular Geometry

  10. Chemical Bonds • Three basic types of bonds • Ionic • Electrostatic attraction between ions (cations-anions), transfer of electrons • Covalent • Sharing of electrons between non-metal atoms • Metallic • Metal atoms bonded to several other atoms by free moving electrons

  11. Think about it! Salt (eg NaCl, KCl) vs. Sugar (C12O22H11) solutions conduct solutions don’t electricity conduct electricity electrolyte non-electrolyte ionic molecular • Sugar and salt differ in the type of attractive forces between the atoms/ions in the compound. • Chemical bond:strong attractive force that exists between atoms (or ions) in a compound

  12. Chemical Bonds • Covalent Bonds:the attractive force between atoms in a molecule that results from sharing of one or more pairs of electrons • non-metals • H2O : O H H • Cl2 : Cl Cl H-O and Cl-Cl bonds result from sharing of electrons

  13. Molecules • Chemistry happens among the electrons of atoms. • Bonds occur between atoms as a result of interactions among the electrons. • When the interaction results in shared electrons, the bond is said to be covalent and the entity formed is a molecule. • Formed from nonmetals.

  14. Naming and writing formulas of binary molecular compounds A molecular formula shows the number and type of elements in a molecule. The molecular formula for water is H2O. Molecule: non-metal elements only! Proper chemical name includes individual element names. Yes, there are RULES for naming!

  15. Naming Binary Molecular Compounds • Name the element farther to the left on the periodic table. (Exception: O is written last unless it’s bonded to F.) • If both elements are in the same group, name the lower one first. • Give the anion name of the second element (ie replace the end of element name with “ide” ending, eg oxygen -> oxide, chlorine->chloride). • Greek prefixes give the number of each atom in the formula. mono- is not used on the first element listed

  16. Nomenclature of Binary Compounds • If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one. N2O5: dinitrogen pentoxide pentaoxide

  17. Name them! • Name the following compounds: • (a) SO2 • (b) PCl5 • (c) N2O3

  18. Chemical bonds involves electrons, valence electrons • Electron configuration for sodium: [Ne]3s1 • The 3s electron of a sodium atom is a valence electron. • electrons residing in the incomplete outer shell of an atom • involved in chemical bonding and ion formation • For elements with atomic # <30, all outer shell electrons are valence electrons (those electrons beyond noble gas core!)

  19. Lewis Symbols • Since valence electrons are involved in the formation of chemical bonds, it’s important to keep track of them. • Draw elements with valence electrons as dots: • electron-dot symbols or Lewis symbols • simple way to depict valence electrons and track them during the formation of chemical bonds

  20. To write Lewis Symbols • You must be able to determine the number of valence electrons for the main group elements. • For main group elements: number of valence electrons for an element = group number of the element e.g. N (group 5A) has 5 valence electrons Br (group 7A) has 7 valence electrons

  21. Lewis Symbols • Lewis symbol has two components: • chemical symbol for the element • plus a dot for each valence electron • dots are placed on all 4 sides of the chemical symbol • all four sides of the symbol are equivalent • up to 2 dots (electrons) per side (start with single dot per side of square, then double up)

  22. O Lewis Symbol for Oxygen Chemical symbol: O Group number: 6A # of valence electrons: 6

  23. Si Lewis Symbols Draw the Lewis symbol for silicon. Chemical symbol: Si Group number: 4A # of valence electrons: 4

  24. Ar Lewis Symbols Draw the Lewis symbol for argon. Chemical symbol: Ar Group number: 8A # of valence electrons: 8

  25. Octet Rule • The noble gases are particularly stable because their outer shell is full of electrons. • With the exception of He, all noble gases have 8 valence electrons. ns2np6 • Octet Rule:Atoms tend to gain, lose, or share electrons until they are surrounded by 8 valence electrons

  26. Octet Rule • The octet rule can be used to predict the charge of ions formed by main group elements as well as the structure of molecular compounds. • BUT: There are many exceptions to the octet rule: H, He….

  27. Covalent Bonding • Molecular compounds are held together by covalent bonds that result from the sharing of electrons. • Simplest example of a covalent bond is H H Indicates 2 shared electrons = 1 bond

  28. Covalent Bonding • When 2 H atoms approach each other, electrostatic interactions occur between their respective electrons and their nuclei. • The two nuclei repel each other • The two electrons repel each other • The nuclei and the electrons attract each other.

  29. Covalent Bonding • The attractions between the nuclei and the electrons cause the electron density to concentrate between the two nuclei. • The atoms in H2 are held together by the electrostatic attraction of the two nuclei for the concentration of negative charge between them. • The shared pair of electrons between the two nuclei acts as “glue”.

  30. Covalent Bonding • Lewis structures (also called electron-dot structures) can be used to represent the covalent bonds that are present in a molecule. • Symbol for each atom • Bond between atoms depicted using a solid line • Unshared electron pairs are shown around the appropriate atom

  31. Lewis Structures Lewis structures are representations of molecules showing all valence electrons, bonding and nonbonding. NOTE: Octet of electrons around each Cl atom!

  32. H + F H F = H F Covalent Bonding • The formation of H2: H + H  H H or H H • The Lewis structure for HF:

  33. Covalent Bonding • The bond between H and F in HF is called a single bond: • sharing of one pair of electrons • In some molecules, atoms attain an octet of electrons by sharing more than one pair of electrons. • Double bond • Triple bond

  34. Covalent Bonding • Double bond:two electron pairs are shared between atoms • depicted using two lines to represent the two shared electron pairs O C O or O C O Carbon dioxide

  35. Covalent Bonding • Triple bond: three electron pairs are shared between atoms • depicted using three lines to represent the 3 pairs of shared electrons N + N N N or N N Nitrogen (N2)

  36. Covalent Bonding • In some molecular compounds, the bonding electrons are shared equally between the atoms in the molecule: H2 F2 N2 • Nonpolar covalent bond:bonding electrons are shared equally

  37. Covalent Bonding • In many molecular compounds, however, one atom attracts the bonding electrons more strongly than the other. d+ d- Fluorine attracts electrons more strongly than hydrogen. H – F The fluorine end of the molecule has greater electron density than the hydrogen end. d+ d- The H – F bond is a polar covalent bond.

  38. Covalent Bonding • Polar covalent bond: • a chemical bond in which the electrons are not shared equally • one atom attracts the bonding electrons more strongly • The polarity of a covalent bond can be determined using the difference in electronegativity between the two atoms.

  39. Covalent Bonding • Electronegativity: • The ability of an atom in a molecule to attract electrons to itself • Range: 0.7 (Cs) - 4.0 (F) • As electronegativity increases, the attraction that an atom has for electrons increases.

  40. Covalent Bonding • Trends to know: • Electronegativity increases: • From left to right across a row • From bottom to top of a column • The four most electronegative elements are: • F (4.0) • O (3.5) • N (3.0) • Cl (3.0)

  41. Covalent Bonding • Chemicals bonds exist along a continuum: • The greater the difference in electronegativity between two atoms, the more polar their bond. F – F D en = 0 nonpolar covalent (en<0.5) H – F D en = 1.9 polar covalent (0.5<en<2.0) Li – F D en = 3.0 ionic (en>2.0) Polar Covalent Bonds Nonpolar Covalent Bonds Ionic Bonds

  42. Testing Which of the following bonds is a polar covalent bond? Br-Br C-H O-H Al-Cl 0 1.34 1.55 0.45 DEN= Polar covalent (0.5< en < 2.0)

  43. Polar Covalent Bonds • When two atoms share electrons unequally, a bond dipole results. • Indicated by d+ and d- or • by an arrow pointing toward the more electronegative atom. d+ d- H - F

  44. Polar Covalent Bonds The greater the difference in electronegativity, the more polar is the bond.

More Related