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The Basics. Elements, Molecules, Compounds, Ions Parts of the Periodic Table How to Name. Classification of Elements. Metalloids (or semi-metals) – along the stair-step line Properties are intermediate between metals and nonmetals.

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The Basics


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    1. The Basics • Elements, Molecules, Compounds, Ions • Parts of the Periodic Table • How to Name

    2. Classification of Elements • Metalloids (or semi-metals) – along the stair-step line • Properties are intermediate between metals and nonmetals • Nonmetals – found on the right-hand side of the Periodic Table • Metals – found on the left-side of the Periodic Table

    3. Elements , Molecules, Compounds, and Ions • Element – one single type of atoms • Al Cu He • Naturally occurring elements that are Diatomic are still elements • N2 O2 F2 Cl2 Br2 I2 H2 • How many elements are in Mn(SO4)2 ? 3 • How many Atoms? 9 • Molecule - smallest electrically neutral unit of a substance that still has the properties of that substance, 2 or more different elements

    4. Elements , Molecules, Compounds, and Ions • Compounds – groups of atoms • Ionic and Molecular • Molecular Compounds – share electrons typically 2 or more non-metals (hydrocarbons) • Example H2S CO2 C5H10 • Ionic Compound (salts) – transfer electrons typically metal and non metal (watch for poly atomic ions) • FeS Mg(OH)2 (NH4)3PO4

    5. Ions • Ions have either lost or gained electrons • Typically Metals lose electrons to become positive • Example cations • Mono-valent Mg -> Mg2+ + 2e- • Group 1A = Metal1+, 2A =Metal2+ and 3A = metal3+ • Multi-valent Fe --> Fe2+ + 2e- and Fe --> Fe3+ + 3e- • Non-metals gain electrons - anion • S + 2e- -> S2-

    6. Poly Atomic Ions • A molecular compound with a charge • NH4+ • CO32- • SO42- • NO3- • OH- • H3O+ • Ammonium • Carbonate • Sulfate • Nitrate • Hydroxide • Hydronium

    7. Acids and Bases From Ions H+ or OH- • Acids look for hydrogen up front (HA) or as COOH • Example HF H3PO4 C4H6COOH • Strong Acids • HCl, H2SO4, HBr, HI, HClO3,HNO3 • Base look for hydroxide or NHgroup • Example KOH C4H6NH2

    8. Naming Compounds 1. Ionic or Covalent 2. Ionic – two ions or Poly atomic ions Covalent 2 non-metals Or a hydrocarbon Type of Metal Mono-valent metals groups 1A 2A 3A Name the Metal Name the Nonmetal + ide (if PAI use its name) CaF2 calcium fluoride RbNO3 rubidium nitrate Al(OH)3 aluminum hydroxide Multi-valent Metal Transitions metals and under the stairs Find the Charge on the Metal To make the compound neutral Write the Charge with roman numeral Name nonmetal + ide (if PAI use its name) Ni+Cl- nickel(I) chloride Pb2+SO42- lead(II) sulfate Pb4+(SO4)22- lead(IV) sulfate

    9. Ionic naming • Name to Formula– Final compound must be neutral based on subscripts and charges • Magnesium Fluoride Mg2+ + F- MgF2 • Ammonium Sulfide NH4+ + S2- (NH4)2S • Tin(II) Carbonate Sn2+ + CO32- SnCO3 • Iron(III) Oxide Fe3+ + O2- Fe2O3 • Iron(II) Oxide Fe2++ O2- FeO

    10. Ionic or Covalent Ionic – two ions or Poly atomic ions Covalent 2 non-metals Or a hydrocarbon Hydrocarbons Look for how many carbons One – methane CH4 Two – ethane C2H6 Three – propane C3H8 Four – butane C4H10 Use the prefix to tell how many of each atom there are Mono is never used with the first element Example PBr3 Phosphorous tribromide CCL4 Carbon tetrachloride P2O5 diphosphorouspentoxide CO carbon monoxide Hydrogen up front Most likely an Acid you should have memorized HCl - Hydrochloric acid HI - Hydroiodic acid HBr - Hydrobromic acid HNO3 Nitric Acid H2SO4 - Sulrufic acid HClO- Hypochlorous acid

    11. Molecular Compounds • Name to formula – charge does not matter this time, just use the prefixes or memorize • Tetraarsenichexoxide As4O6 • Sulfur hexafluroide SF6 • Butane C4H10 • Nitric acid HNO3

    12. Lewis Dot Structures • Rules • Fewest number of atoms goes in the middle or C if present • Connect remaining elements with single bonds • Make sure all elements have 8 electrons (H only 2) • Count the number of electrons in structure • Add up valence electrons from PT • Too many e- in structure: remove 2 adjacent pairs fill in with one bond • Too few e- in structure: add to central atom

    13. EXAMPLES: • CH4 • 1. (1) C + (4) H (1)(4e-) + (4)(1e-) = 8e- • 2. Spatial order • 3. Draw bonds • 4. Octet rule satisfied? • 5. # of e- match? H C H H H

    14. EXAMPLES: • CO2 • 1. (1) C + (2) O (1)(4e-) + (2)(6e-) = 16e- • 2. Spatial order • 3. Draw Bonds • 4. Octet rule satisfied? • 5. # of e- match? O C O

    15. EXAMPLES: • NH3 • 1. (1) N + (3)H (1)(5e-) + (3)(1e-) = 8e- • 2. Spatial order • 3. Draw bonds • 4. Octet rule satisfied? • 5. # of e- match? N H H H

    16. EXAMPLES: • CCl4 • 1. (1) C + (4) Cl (1)(4e-) + (4)(7e-) = 32e- • 2. Spatial Order • 3. Draw bonds • 4. Octet rule satisfied? • 5. # of e- match? Cl Cl Cl C Cl

    17. EXAMPLES: • NH4+ • 1. (1) N + (4) H - (1)(+) (1)(5e-)+ (4)(1e-) - (1)(1e-) = 8e- • 2. Spatial order • 3. Draw bonds • 4. Octet rule satisfied? • 5. # of e- match? H [ ]+ N H H H

    18. EXAMPLES: • SO42- • 1. (1) S + (4) O + (2)(-) (1)(6e-)+ (4)(6e-) + (2)(1e-) = 32e- • 2. Spatial Order • 3. Draw bonds • 4. Octet rule satisfied? • 5. # of e- match? O [ ] 2- O S O O

    19. EXAMPLES: • CN- • 1. (1) C + (1) N + (1)(-) (1)(4e-) + (1)(5e-)+ (1)(1e-) = 10e- • 2. Spatial order • 3. Draw Bonds • 4. Octet rule satisfied? • 5. # of e- match? [ ]- C N

    20. EXAMPLES: • CO32- • 1. (1) C + (3) O + (2)(-) (1)(4e-)+ (3)(6e-) + (2)(1e-) = 24e- • 2. Spatial Order • 3. Draw bonds • 4. Octet rule satisfied? • 5. # of e- match? [ ] 2- O C O O

    21. VSEPR: • Regions of electron density (where pairs of electrons are found) can be used to determine the shape of the molecule. • CO2 • Here there are two regions of electron density. • The regions want to be as far apart as possible, so it is linear. O C O

    22. EXAMPLES: 1 4 H • CH4 • There are four electron pairs. • You would expect that the bond angles would be 90° but… • Because the molecule is three-dimensional, the angles are 109.5°. • The molecule is of tetrahedralarrangement. C H H 2 H 3

    23. EXAMPLES: 1 4 • NH3 • Four regions of electron density • But one of the electron pairs is a lone pair • The shape is called trigonal pyramidal N H H 2 H 3

    24. EXAMPLES: 1 4 • H2O • Four regions of electron density • But two are lone pairs • This structure is referred to as bent O H H 2 3

    25. EXAMPLES: 1 3 [ ] 2- • CO32- • Three regions of electron density • This structure is referred to as trigonal planar C O O 2 O

    26. Practice determining molecular shape: • H2S • 4 regions of e- density • 2 lone pairs • bent S H S H H H

    27. Practice determining molecular shape: • SO2 • 3 areas of e- density • 1 lone pair • bent S O O S O O

    28. Practice determining molecular shape: • CCl4 • 4 areas of e- density • tetrahedral Cl 3d Cl Cl C Cl

    29. Practice determining molecular shape: • BF3 • 3 areas of e- density • trigonal planar F B F F B F F F

    30. Practice determining molecular shape: • NF3 • 4 areas of e- density • 1 lone pair • pyramidal 3d N F F F

    31. 16.3Polar Bonds and Molecules • In covalent bonds, the sharing of electrons can be equal • or it can be unequal.

    32. Nonpolar Covalent Bonds • Nonpolar covalent bond - This is a covalent bond in which the electrons are shared equally. • EXAMPLES: • H2 • Br2 • O2 • N2 • Cl2 • I2 • F2 Br Cl H N F I O O Br Cl H N F I

    33. Polar Bonds and Molecules • If the sharing is unequal, the bond is referred to as a dipole. • A dipole has 2 separated, equal but opposite charges. • “∂” means partial _ +

    34. Polar Bonds and Molecules • Polar covalent bond - a covalent bond that has a dipole • It usually occurs when 2 different elements form a covalent bond. • EXAMPLE: • H + Cl  H Cl

    35. Polar Bonds and Molecules • Electronegativity - This is the measure of the attraction an atom has for a shared pair of electrons in a bond. • Electronegativity values increase across a period and up a group.

    36. Examples: • Identify the type of bond for each of the following compounds: • HBr Br = 2.8 H = 2.1 0.7 .1 < < 1.9 Polar Covalent H Br

    37. Examples: • NaF F = 4.0 Na = 0.9 3.1 • N2 N = 3.0 N = 3.0 0.0 Na F > 2 Ionic Non-Polar Covalent N N

    38. Molecular Polarity • If there is only one bond in the molecule, the bond type and polarity will be the same. • If the molecule consists of more than 2 atoms, you must consider the shape. To determine its polarity, consider the following: • Lone pairs on central atom • If so… it is polar • Spatial arrangement of atoms • Do bonds cancel each other out (symmetrical)? • If so… nonpolar • Do all bonds around the central element have the same difference of electronegativity? • If so… nonpolar

    39. Polar Molecules • If the molecule is symmetrical it will be nonpolar. • Exception hydrocarbons are nonpolar • If the molecule is not symmetrical it will be polar, with a different atom or with lone pair(s) Br Cl Cl C Cl

    40. Attractions Between Molecules • Van der Waals forces – the weakest of the intermolecular forces. These include London dispersion and dipole-dipole forces. • London dispersion forces – between nonpolar molecules and is caused by movement of electrons

    41. Attractions Between Molecules • van der Waals forces(cont.)- • Dipole interactions – between polar molecules and is caused by a difference in electronegativity.

    42. Attractions Between Molecules • Hydrogen bonds – attractive forces in which hydrogen, covalently bonded to a very electronegative atom (N, O, or F) is also weakly bonded to an unshared (lone) pair of electrons on another electronegative atom. O O H H H H O H H

    43. Attractions Between Molecules • Ionic Bonding-occurs between metals and nonmetals when electron are transferred from one atom to another. • These bonds are very strong.

    44. Summary of the Strengths of Attractive Forces Ionic bonds hydrogen bonds dipole-dipole attractions LDF

    45. Writing and Balancing Chemical Equations Example: Write the equation for the formation of sodium hydroxide and hydrogen, from the reaction of sodium with water.

    46. Write the equation for the formation of sodium hydroxide and hydrogen, from the reaction of sodium with water. • Write the formulas of all reactants to the left of the arrow and all products to the right of the arrow. Sodium + water sodium hydroxide + hydrogen Translate the equation and be sure the formulas are correct. Na + H2O NaOH + H2

    47. Write the equation for the formation of sodium hydroxide and hydrogen, from the reaction of sodium with water. • Once the formulas are correctly written, DO NOTchange them. Use coefficients (numbers in front of the formulas), to balance the equation.DO NOT CHANGE THE SUBSCRIPTS! _____Na + _____H2O ____NaOH + _____H2

    48. Begin balancing with an element that occurs only once on each side of the arrow. • Ex: Na 2 2 2 _____Na + _____H2O ____NaOH + _____H2 When you are finished, you should have equal numbers of each element on either side of the equation Na H O 2 Na H O 2 4 4 2 2

    49. To determine the number of atoms of a given element in one term of the equation, multiply the coefficient by the subscript of the element. Ex: In the previous equation (below), how many hydrogen atoms are there? 4 2 2 2 ____Na + _____H2O ____NaOH + _____H2

    50. Helpful Hints: • Balance elements one at a time. • Balance polyatomic ions that appear on both sides of the equation as single units. (Ex: Count sulfate ions, not sulfur and oxygen separately) • Balance H and O last. Save the one that is in the most places for last… • Use Pencil! (NH4)2SO4 (aq) + BaCl2 (aq)  BaSO4 (s) + 2NH4Cl (aq)