![IONic bonding
• Always formed between metals and
non-metals
[METALS ]+
[NON-METALS ]-
Lost e-
Gained e-](https://image.slidesharecdn.com/2-250503094925-01df9aed/85/Different-types-of-bonding-in-materials-for-6-320.jpg)
![NaCl
• This is the finished Lewis Dot
Structure
[Na]+
[ Cl ]
-
How did we get here?](https://image.slidesharecdn.com/2-250503094925-01df9aed/85/Different-types-of-bonding-in-materials-for-14-320.jpg)
Different types of bonding in materials for
- 1. Chemical Bond • A bond results from the attraction of nuclei for electrons – All atoms trying to achieve a stable octet • IN OTHER WORDS – the p+ in one nucleus are attracted to the e- of another atom • Electronegativity
- 2. Two Major Types of Bonding • Ionic Bonding – forms ionic compounds – transfer of e- • Covalent Bonding – forms molecules – sharing e-
- 3. One minor type of bonding • Metallic bonding – Occurs between like atoms of a metal in the free state – Valence e- are mobile (move freely among all metal atoms) – Positive ions in a sea of electrons • Metallic characteristics – High mp temps, ductile, malleable, shiny – Hard substances – Good conductors of heat and electricity as (s) and (l)
- 4. It’s the mobile electrons that enable me- tals to conduct electricity!!!!!!
- 5. IONic Bonding • electrons are transferred between valence shells of atoms • ionic compounds are made of ions • ionic compounds are called Salts or Crystals NOT MOLECULES
- 6. IONic bonding • Always formed between metals and non-metals [METALS ]+ [NON-METALS ]- Lost e- Gained e-
- 7. IONic Bonding • Electronegativity difference > 2.0 – Look up e-neg of the atoms in the bond and subtract NaCl CaCl2 • Compounds with polyatomic ions NaNO3
- 9. • hard solid @ 22o C • high mp temperatures • nonconductors of electricity in solid phase • good conductors in liquid phase or dissolved in water (aq) SALTS Crystals Properties of Ionic Compounds
- 10. Covalent Bonding • Pairs of e- are shared between non-metal atoms • electronegativity difference < 2.0 • forms polyatomic ions molecules
- 11. Properties of Molecular Substances • Low m.p. temp and b.p. temps • relatively soft solids as compared to ionic compounds • nonconductors of electricity in any phase Covalent bonding
- 12. Covalent, Ionic, metallic bonding? • NO2 • sodium hydride • Hg • H2S • sulfate • NH4 + • Aluminum phosphate • KH • KCl • HF • CO • Co Also study your characteristics!
- 13. Drawing ionic compounds using Lewis Dot Structures • Symbol represents the KERNEL of the atom (nucleus and inner e-) • dots represent valence e-
- 14. NaCl • This is the finished Lewis Dot Structure [Na]+ [ Cl ] - How did we get here?
- 15. • Step 1 after checking that it is IONIC – Determine which atom will be the + ion – Determine which atom will be the - ion • Step 2 – Write the symbol for the + ion first. • NO DOTS – Draw the e- dot diagram for the – ion • COMPLETE outer shell • Step 3 – Enclose both in brackets and show each charge
- 16. Draw the Lewis Diagrams • LiF • MgO • CaCl2 • K2S
- 17. Drawing molecules using Lewis Dot Structures • Symbol represents the KERNEL of the atom (nucleus and inner e-) • dots represent valence e-
- 18. Always remember atoms are trying to complete their outer shell! The number of electrons the atoms needs is the total number of bonds they can make. Ex. … H? O? F? N? Cl? C? one two one three one four
- 19. Methane CH4 • This is the finished Lewis dot structure How did we get here?
- 20. • Step 1 – count total valence e- involved • Step 2 – connect the central atom (usually the first in the formula) to the others with single bonds • Step 3 – complete valence shells of outer atoms • Step 4 – add any extra e- to central atom IF the central atom has 8 valence e- surrounding it . . YOU’RE DONE!
- 21. Sometimes . . . • You only have two atoms, so there is no central atom, but follow the same rules. • Check & Share to make sure all the atoms are “happy”. Cl2 Br2 H2 O2 N2 HCl
- 22. • DOUBLE bond – atoms that share two e- pairs (4 e-) O O • TRIPLE bond – atoms that share three e- pairs (6 e-) N N
- 23. Draw Lewis Dot Structures You may represent valence electrons from different atoms with the following symbols x, , CO2 NH3
- 24. Draw the Lewis Dot Diagram for polyatomic ions • Count all valence e- needed for covalent bonding • Add or subtract other electrons based on the charge REMEMBER! A positive charge means it LOST electrons!!!!!
- 25. Draw Polyatomics • Ammonium • Sulfate
- 26. Types of Covalent Covalent Bonds Bonds • NON-Polar bonds –Electrons shared evenly in the bond –E-neg difference is zero Between identical atoms Diatomic molecules
- 27. Types of Covalent Bonds Polar bond –Electrons unevenly shared –E-neg difference greater than zero but less than 2.0 closer to 2.0 more polar more “ionic character”
- 28. non-polar MOLECULES • Sometimes the bonds within a molecule are polar and yet the molecule is non-polar because its shape is symmetrical. H H H H C Draw Lewis dot first and see if equal on all sides
- 29. Polar molecules (a.k.a. Dipoles) • Not equal on all sides –Polar bond between 2 atoms makes a polar molecule –asymmetrical shape of molecule
- 30. H Cl - +
- 32. Water is a bent molecule O H H H H
- 33. Making sense of the polar non-polar thing BONDS Non-polar Polar Identical Different MOLECULES Non-polar Polar Symmetrical Asymmetrical
- 34. IONIC bonds …. Ionic bonds are so polar that the electrons are not shared but transferred between atoms forming ions!!!!!!
- 35. C. Johannesson VSEPR Theory • Valence Shell Electron Pair Repulsion Theory • Electron pairs orient themselves in order to minimize repulsive forces.
- 36. C. Johannesson VSEPR Theory • Types of e- Pairs – Bonding pairs - form bonds – Lone pairs - nonbonding e- Lone pairs repel more strongly than bonding pairs!!!
- 37. 4 Shapes of molecules
- 38. 1. Linear (straight line) Ball and stick model Space filling model
- 39. 2. Bent Ball and stick model Space filling model
- 40. 3.Trigonal pyramid Ball and stick model Space filling model
- 41. 4.Tetrahedral Ball and stick model Space filling model
- 42. • Attractions between molecules – van der Waals forces • Weak attractive forces between non- polar molecules – Hydrogen “bonding” • Strong attraction between special polar molecules Intermolecular attractions
- 43. van der Waals • Non-polar molecules can exist in liquid and solid phases because van der Waals forces keep the molecules attracted to each other • Exist between CO2, CH4, CCl4, CF4, diatomics and monoatomics
- 44. van der Waals periodicity • increase with molecular mass. • increase with closer distance between molecules – Decreases when particles are farther away
- 45. Hydrogen “Bonding” • Strong polar attraction – Like magnets • Occurs ONLY between H of one molecule and N, O, F of another H “bond”
- 46. H is shared between 2 atoms of OXYGEN or 2 atoms of NITROGEN or 2 atoms of FLUORINE Of 2 different molecules
- 47. Why does H “bonding” occur? • Nitrogen, Oxygen and Fluorine – small atoms with strong nuclear charges • powerful atoms – very high electronegativities
- 48. Intermolecular forces dictate chemical properties • Strong intermolecular forces cause high b.p., m.p. and slow evaporation (low vapor pressure) of a substance.
- 49. Which substance has the highest boiling point? • HF • NH3 • H2O • WHY? Fluorine has the highest e-neg, SO HF will experience the strongest H bonding and needs the most energy to weaken the i.m.f. and boil